Name_____________________________________ Period__________

Honors Chemistry Semester 1

Ms. Regli

2

Table of Contents

Title Page Periodic Table.............................................................................................3 Equations Sheet.........................................................................................4 Ion Sheet.....................................................................................................5 Activity Series of Metals / Solubility Rules..............................................7 Unit 1: Measurements (Chapter 1)...........................................................8 Unit 2: Density/Atomic Theory/Nuclear (Ch. 1, 2, 23)............................20 Unit 3: Nomenclature (Chapter 2)............................................................44 Unit 4: The Mole (Chapter 3)....................................................................52 Unit 5: Stoichiometry (Chapter 3, 4)........................................................66 Unit 6: Atomic Theory II/Quantum (Chapter 7, 8)...................................82 Unit 7: Bonding (Chapter 9, 10).............................................................106

3

4

General Conversion Factors Length Mass Volume Area 1 inch = 2.54 cm 1 lb. = 16 oz. = 453.6 g 1 L = 1.057 qt. 1 acre = 43,560 ft2 1 meter = 39.37 in 1 kg = 2.205 lb. 1 gal. = 4 qt. = 8 pt. 1 mile = 1.609 km = 5280 ft 1 oz. = 28.35 g 1 pt. = 2 cups = 16 fl. oz. Energy 1 foot = 30.48 cm 1 mL = 1 cm3 1.60 x 10-19 J = 1 eV 1 cal = 4.184 J Pressure 1 atm = 14.7 psi = 760 mmHg = 760 torr = 101.3 kPa General Information D = m/V Density of water = 1 g/mL melectron = 9.11 x 10-31kg 1 mol = 6.022 x 1023 E = mc2 1 mol (gas) = 22.4 L @ STP

Ea =⏐O-A⏐ Er = (Ea / A) x 100 Geometry Equations Vol. of cylinder = πr2h Vol. of sphere = (4/3)πr3 Vol. of cube = s3 c = 2πr = πd Atomic Theory Equations Atomic Theory Constants E = hf = hc/λ h = 6.626 x 10-34 Js ΔE = Ef - Ei c = 3.00 x 108 m/s En = -RH/n2 RH = 2.18 x 10-18 J = 13.6 eV λ = h/mv rn = n2(5.3 x 10-11m) Gas Law Equations Gas Law Constants

Thermochemistry Equations Thermochem Constants for H2O Q = mcΔT Q = CcalΔT mHcH(TH-Tf) = mccc(Tf-Tc) cice = 2.06 J/g°C ΔHrxn = ΣΔH°f products - ΣΔH°f reactants ΔG = ΔH - TΔS cwater = 4.184 J/g°C = 1 cal/g°C ΔHrxn = Σ E reactants – Σ E products csteam = 2.02 J/g°C Q = mlf Q = mlv c = kP lf = 334 J/g lv = 2260 J/g Solutions Equations M = mol / L m = mol / kg solvent ΔTf = Kfm N = Equivalents / L % by mass = mass solute / mass solution x 100 ΔTb = Kbm M1V1 = M2V2 X = mol solute / mol solute + mol solvent π = MRT P1 = X1P°1 ΔP = X2P°1 Acids / Bases Equations Acids / Bases Constants Kw = [H+][OH-] pH = - log [H+] pOH = - log [OH-] Kw (25˚C)= 1.0 x 10-14 pH + pOH = 14 NaVa = NbVb Ka ·Kb = Kw Redox Equations Redox Constants ΔG = -nFEcell E°cell = E°red + E°ox 1 e- = 1.6 x 10-19 C 1 mol e- = 96,487 C = 1 F

5

Inorganic Nomenclature Ion Sheet Directions: Make flashcards of the following ions. The name should be written on one side, the chemical formula on the other.

Positive Ions (Cations)

Positive Ions with a fixed charge: Group IA ions: Group IIA ions: Group IIIA ions: Misc. cations: Hydrogen H+ Beryllium Be2+ Aluminum Al3+ Silver Ag+ Magnesium Mg2+ Boron B3+ Lithium Li+ Calcium Ca2+ Ammonium NH4

+ Sodium Na+ Strontium Sr2+ Potassium K+ Barium Ba2+ Zinc Zn2+ Rubidium Rb+ Cesium Cs+ Cadmium Cd2+ Positive Ions with multiple charges: (one card for each ion) IUPAC Name Common Name Ion Iron (II) Ferrous Fe2+

Iron (III) Ferric Fe3+

Copper (I) Cuprous Cu+

Copper (II) Cupric Cu2+

Cobalt (II) Cobaltous Co2+

Cobalt (III) Cobaltic Co3+

Mercury (I)** Mercurous Hg22+

Mercury (II) Mercuric Hg2+

Manganese (II) Manganous Mn2+

Manganese (III) Manganic Mn3+

Tin (II) Stannous Sn2+

Tin (IV) Stannic Sn4+

Lead (II) Plumbous Pb2+

Lead (IV) Plumbic Pb4+

**Mercury (I) exists as a pair or dimmer as shown. No single Hg+ ions are known to exist.

6

Negative Ions (Anions) Monatomic anions: Group IVA ions: Group VA ions: Group VIA ions: Group VIIA ions: Carbide C4- Nitride N3- Oxide O2- Fluoride F- Silicide Si4- Phosphide P3- Sulfide S2- Chloride Cl- Bromide Br- Group IA ions: Iodide I- Hydride H- Polyatomic anions: Oxyhalogens: Hypochlorite ClO- Hypobromite BrO- Hypoiodite IO- Chlorite ClO2

- Bromite BrO2- Iodite IO2

- Chlorate ClO3

- Bromate BrO3- Iodate IO3

- Perchlorate ClO4

- Perbromate BrO4- Periodate IO4

- Miscellaneous polyatomic ions: (one card for each ion) Peroxide O2

2- Cyanide CN-

Hydroxide OH- Thiocyanate SCN-

Carbonate CO32- Cyanate OCN-

Hydrogen carbonate / bicarbonate HCO3- Nitrite NO2

-

Monohydrogen phosphate HPO42- Nitrate NO3

-

Dihydrogen phosphate H2PO4- Sulfite SO3

2-

Arsenite AsO33- Sulfate SO4

2-

Arsenate AsO43- Thiosulfate S2O3

2-

Chromate CrO42- Permanganate MnO4

-

Dichromate Cr2O72- Oxalate C2O4

2-

Phosphite PO33- Phthalate C8H4O4

2-

Phosphate PO43- Silicate SiO4

4-

Acetate C2H3O2- Borate BO3

3-

Binary Covalent Greek Prefixes Prefix Number Prefix Number Mono- 1 Hexa- 6 Di- 2 Hepta- 7 Tri- 3 Octa- 8 Tetra- 4 Nona- 9 Penta- 5 Deca- 10

7

(Fig. 4.15) Activity Series of Metals Li K Ba Ca Na React with cold water to produce H2 Mg Al Zn Cr Fe Cd React with steam to produce H2 Co Ni Sn Pb React with acids to produce H2 H Cu Ag Hg Pt Au Do NOT react with water or acids

Solubility Rules: 1. ALL alkali metal compounds are soluble. 2. ALL NH4

+ compounds are soluble. 3. ALL compounds containing NO3

-, ClO3-, and ClO4

- are souble. 4. Most OH- compounds are INSOLUBLE. Alkali metal hydroxides and Ba(OH)2 are exceptions. 5. Most Cl-, Br-, and I- compounds are soluble. The exceptions are: Ag+, Hg2

2+, Pb2+. 6. All CO3

2-, PO43-, and S2- compounds are INSOLUBLE. Exceptions are alkali metal

and ammonium compounds. 7. Most SO4

2- compounds are soluble. Exceptions are Ba2+, Hg2+, and Pb2+.

8

Unit 1: Measurements I. Algebra In Chemistry A. Many relationships used in chemistry require working with algebraic equations. The equation or equations you start with are the _______________ equations while the equation that you solve for is called the ________________ equation. The working equation is formed by ____________ the desired variable on one side of the equation. The basic rule for doing this is …. “What you do to one side of the equation you must do to the _____________” Examples: 1. Find x for the following: (a) x + y = z (b) x – y = z (c) 2x + y = z (d) 3x + 2y = z (e) x2 + y = z (f) x3 + 3y = z (g) 2x – 4 = 14 2. D = m/V is the given equation for the density of an object.

(a) If the mass and density of an object were given, what would be the equation used to find the volume?

(b) If the density and volume of an object were known, what equation could be used to find its mass?

3. E = mc2; Solve for (a) m (b) c 4. PV = nRT; Solve for (a) V (b) T 5. The formula for the volume of a cylinder is V = πr2h; Solve for (a) h & (b) r 6. The formula for the volume of a sphere is V = 4/3 πr3. What equation could be used to solve for the radius of a spherical container of known volume?

7. Solve for V2:

9

B. It is often required to substitute one or more equations into another equation to form the working equation. 1. For example: If the radius, mass, and density of a cylinder are known, the equation for the height of the cylinder can be formed by combining and rearranging the density and volume formulae. 2. If the length, width, height, and density of a block of metal are known, show the equation that can be used to calculate its mass. 3. If the radius and mass of a metal ball are known, show the equation that can be used to determine its density. 4. Derive an equation that can be used to determine the diameter of a crystal ball with a known mass and density. 5. If the radius and height of a cylinder are known, determine the formula that can be used to find the side of a cube that has twice the volume as the cylinder. C. Algebra can also be used to solve problems …… 1. A chemist has 300 g of 60.0 % acid. How much water should be added so that it is only 45.0 % acid?

10

II. Useful Geometry Formulas Area of a square Area of a rectangle Area of a circle A = A = A = Circumference of a circle c = c = Volume of a square Volume of a rectangular solid

V = V = Volume of a sphere V = Volume of a cylinder

V =

Density of an Object: D =

11

III. Significant Figures:

A. Example: 1. Calculate the area of the dark rectangle. 2. Calculate the volume of the object. 3. Calculate the sum of the length, width, and height. 4. What is the length of each segment?

h = 0.05 cm

l = 17.9 cm

w = 6.87 cm

11 cm 10 cm

A B C D

A = ____________________

B = ____________________

C = ____________________

D = ____________________

12

B. Introduction: When making measurements or doing calculation, you should not keep more digits in a number than is _________________. These rules of significant figures will show you how to determine the correct number of digits. C. What is a significant figure? Significant figures in a measurement are all values (digits) known precisely, plus ______ digit that is estimated. Example: Make the measurement with the correct significant figures. D. How do you determine sig figs in a measurement that has already been recorded?

Sig Figs: The Rules 1. Every nonzero digit in a recorded measurement is significant. Examples: 47,357 5 sig figs 25 _______ 2. Zeroes appearing between nonzero digits are significant. (Sandwich rule) Examples: 1.007 4 sig figs 305 _______ 3. Zeroes appearing in front of all nonzero digits are NOT significant. They are acting as place holders (leading zeroes). Examples: 0.00238 3 sig figs 0.98 _______ 0.000006 _______ 4. Zeroes at the end of a number AND to the right of a decimal point are significant. Examples: 426.00 5 sig figs 2.060 _______ 0.8080 _______ 5. Zeroes at the end of a measurement where there is no decimal point are ambiguous. To clearly show the correct number of sig figs, these measurements should be written in scientific notation. Examples: 120 2-3 sig figs 3000 1-4 sig figs 1,000,000 _______ Examples: Write the number 100,000 with (a) 1 sig fig, (b) 3 sig figs, and (c) 5 sig figs.

a. _____________

b. _____________

c. _____________

d. _____________

e. _____________

f. _____________

g. _____________

h. _____________

a. b. c. d. e. f. g. h.

9 cm 10 cm 9 cm 10 cm 9 cm 10 cm 0 cm 1 cm

13

E. Practice: 1. Determine the number of significant figures in each of the following measurements. (a) 54320.0 (b) 0.004550 (c) 151309 (d) 10.54 (e) 5.20 x 105 (f) 15,000 (g) 10.04 (h) 0.0750 2. When completing calculations, it is often necessary to round the final answer to a particular number of significant figures (round up for 5 and above; keep digits the same for 4 and below). Round the above measurements to 2 significant figures.

Example: 0.0753 = 0.075 107.0 = ________________

3. Determine the number of sig figs for each measurement. Round the measurements to 2 sig figs. If original measurement only contains 1 or 2 sig figs, leave the second line blank.

# sig figs Rounded Answer # sig figs Rounded Answer

1. 0.0037 _______ ______________ 11. 14.04 _______ ______________ 2. 134.1 _______ ______________ 12. 5.401 _______ ______________ 3. 1,000,000 _______ ______________ 13. 1340 _______ ______________ 4. 5.730 x 102______ ______________ 14. 0.00566 _______ ______________ 5. 410.50 _______ ______________ 15. 0.8120 _______ ______________ 6. 79500 _______ ______________ 16. 18.009 _______ ______________ 7. 3071.04 _______ ______________ 17. 100.5 _______ ______________ 8. 4.08 x 10-6_____ __ ______________ 18. 3008 _______ ______________ 9. 0.998 _______ ______________ 19. 112040.0_______ ______________ 10. 1.570 _______ ______________ 20. 43.05 _______ ______________

=========================================================================== 4. Rules for Significant Figures in Calculations a. Multiplication or Division: The number of sig figs in the result is the same as the number in the least precise (least sig figs) measurement.

Example: 4.56 m x 1.4 m = 6.38 m2 (Rounds to TWO sig figs) = 6.4 m2

(a) 17.24 x 0.52 (b) 118.24 x 3.5 (c) 1.007 x 14.40 7.58

b. Addition and Subtraction: The answer should be rounded off to retain digits only as far as the first column containing estimated digits (remember that the last digit is estimated).

Example: 12.11 m + 8.0 m + 1.013 m = 21.123 m (Rounds to ONE place after the decimal) = 21.1 m

(a) 21 cm – 18.3 cm = (b) 10000.00 mm + 25.116 mm =

14

IV. Scientific Notation (Exponential Notation)

A. Chemistry examples: 1. Avogadro’s Number 2. Mass of an electron

B. Technique to change from positional notation to scientific notation: 1. Leave ________ number to the ____________ of the decimal. 2. When the decimal is moved to the _________, the exponent is_______________. 3. When the decimal is moved to the _________, the exponent is ______________. 4. Number must contain the same number of __________________ as the original

value.

C. Convert the following to scientific notation: 1. 135000 (3 sig figs) ____________________

2. 0.005500 ____________________

3. 120,000,000,000 (2 sig figs) ____________________

4. 0.00000004441 ____________________

D. Use of calculator with scientific notation:

Example: 1.61 x 10-19 What you see on your calculator

Step 1: Enter the number

Step 2: Press the Exponent button or

Step 3: Enter the exponent

Step 4: If negative exponent, use key

E. Exponent problems (Use correct sig figs!) 1. (3.5 x 10-2)(4.44 x 1012) (1.280 x 10-22)

2. (1.76 x 10-2)(4.2 x 10-4) (6.99 x 106)(8 x 1014) Raising to a power Taking a root Step 1: Enter number Step 1: Enter number Step 2: Press Step 2: Press

Step 3: Enter power Step 3: Enter root

Step 4: Press Step 4: Press

Example: (a) (14.5)6 Example: (a)

(b) (1.72 x 105)4 (b)

15

V. Metric System

A. Based on powers of 10 Ex. 1 m = _______ dm = _______ cm = _______ mm

B. Uses “________________” and “____________________.” Measurement Metric Base Unit

1.

2.

3.

4.

5.

C. Metric Prefixes: MEMORIZE this table!

Prefix Symbol Multiplier/Factor

1.

2.

3.

4.

5.

6.

7.

Base Unit

8.

9.

10.

11.

12.

13.

D. Examples: Multiplier ALWAYS goes with the _______________________.

1 Mm = _______ m 1 µg = _______ g 1 Ts = _______ s 1 pm = _______m

16

E. Converting within metric system using dimensional analysis:

1. Convert to base unit by canceling units (Top unit cancels with __________ unit).

2. Place the multiplier with the __________________.

3. Place a _______ in front of the unit with ________________.

*** 4. To enter multiplier into the calculator, use a _____ before the exponent key.

Example: 10-6

F. Metric dimensional analysis examples:

1. Convert 3.6 nm to m.

2. Convert 55.6 g to Tg

3. Convert 575 cm to Mm.

4. Convert 0.456 dag to pg.

5. Convert 78.5 km to µm.

6. Convert 0.000590 mL to GL.

Metric / English Conversion Factors (given on test):

Length Mass Volume Time

1 inch = 2.54 cm 1 lb. = 16 oz. = 256 drams 1 L = 1.057 qt. 1 fortnight = 2 weeks

1 meter = 39.37 in 1 kg = 2.205 lb. 1 gal. = 4 qt. = 8 pt.

1 mile = 1.609 km 1 lb = 453.6 g 1 pt. = 2 cups Density of water

1 furlong = 220 yd. 1 oz. = 28.35 g 1 mL = 1 cm3 1 g = 1 mL

1 pt. = 16 fl. oz. 1 g = 1 cm3 VI. Conversion Factors:

A. Whenever two measurements are equal, or __________________, a ratio of these two measurements will equal ________.

Example: ______ ft. = ______ in. Can be written as the following ratios:

17

B. Conversion factor: ratio of ____________________ measurements.

C. Write conversion factors for the following pairs of units:

1. miles and feet

2. days and year 3. yard and feet

D. Assume all conversion factors are _________________ significant. (Use initial number to determine sig figs).

VII. Dimensional Analysis I

Units (______________________) are used to solve a problem.

Examples: A. The average human brain weighs 8.13 lb. What is the mass in ng?

B. How many microseconds in 8.37 years? Write answer in scientific notation.

C. A container contains 15 kL. Convert this to cm3.

D. Apollo 13 re-entered the Earth’s atmosphere at a speed of 32,805 ft/s. What was the speed in miles per hour (mph)?

E. An arrow moves towards you at 235 m/s. How many miles could the arrow move in one day? (Assume the arrow never falls to the Earth).

18

F. (a) Determine the number of cm3 in a 20.0 fl. oz. bottle of Coke. (b) What is the mass of the Coke in pounds, assuming that it is the density of water (1 g / mL)?

G. The speed of light is 3.00 x 108 m/s. How many miles does light travel per year?

H. Carl Lewis set the world record for the 100.0 m dash on August 25, 1991 in the finals of the World Track Championships with a time of 9.86 seconds. What was his average speed in miles per hour?

VIII. Dimensional Analysis II: Square and cubic units

A. Convert 3.7 ft2 to in2.

B. The engine in a Jeep Cherokee is 4.0 L. Calculate the engine volume in (a) in3 and (b) ft3.

C. The density of gold is 19.3 g/mL. Calculate the density of gold in (a) lb/ft3, (b) kg/m3.

19

D. A spherical container with a diameter of 2.85 dam is filled with water. (a) Determine the volume of the sphere in cm3. (b) Determine the mass of the water in kilograms.

E. The dimensions of a swimming pool are 13.5 ft. x 22 m x 225 cm. (a) Determine the volume of the pool in m3. (b) Determine the mass of the water in pounds.

F. A 12.0 fl. oz. soda spilled onto the floor into a cylindrical puddle with a 15.4 inch diameter. Calculate the depth (height) of the puddle in µm.

G. The volume of a red blood cell is 90.0 µm3. What is its diameter in mm? Assume it is spherical.

H. The lid of a soup can is 5.40 cm across and the can is 12.2 cm high. What is the volume of the can in fluid ounces?

20

Unit 2: Density, Atomic Theory & Nuclear I. Density

A. Density is defined as _______________ per unit _______________ and has the symbol ______ (or the Greek letter “rho” = _______).

B. The formula for density is D = _______ Density = _________________ C. Density can have many different units but the most commonly used units are as follows: English system Metric system

for solids & liquids

for gases

D. Converting from between units is easy if you use ____________________ analysis. Example1: Gold has a density of 19.3 g/cm3. What is its density in…. (a) kg/m3 (b) g/mL

(c) lb/ft3 II. Specific Gravity

A. Specific gravity is a measure of the mass of an object compared to the mass of an equal volume of ___________________.

B. It has the formula: sp. gr. =_____________ = mass of object mass of equal volume of water

C. Since the density of water is 1.00 g/mL, we see that specific gravity is equal numerically to the ______________ but has ______ __________!!!

Example 2: Gold has a density of 19.3 g/mL. Calculate the specific gravity of gold.

D. The specific gravity has the same value in any _____________ of units, since it expresses the quotient of the mass of the substance divided by the mass of an equal volume of _____________. It is expressed by a pure number ___________________________. Specific gravity is usually used with _________________. E. The water standard is sometimes taken at the temperature where it has maximum density. At typical laboratory values (0 oC to 30 oC) the density of water does not vary much and the rounded value ________________________ can be used.

21

F. Steps to calculate the density of objects:

1. Find the mass (balance).

2. Find the volume (dimensions of displacement).

3. Use the equation D = m / V

Example 3: Geometric shapes Step 1: Measure the mass of the object.

Step 2: Measure the dimensions of the object.

Step 3: Calculate density (D = m / V).

Other formulae: Vcyl = πr2h Vsphere = 4/3 πr3 Vcube = s3

Example 4: Irregular shapes “Water displacement”

Step 1: Measure mass of object

Step 2: Put water in a graduated cylinder and record initial volume. Step 3: Carefully put object in and record the final volume. Step 4: Calculate density.

l =______cm

h =______cm

w =______cm

mass = _____________ g

Object

Initial measurement

Final measurement

mass = ________ g

Vi = ________ mL

Vf = ________ mL

22

III. Density Calculations: 1. An object has a mass of 1.00 kg and a density of 4.00 g/mL. What is its volume in liters? 2. What is the mass of a 500 cm3 object that has a specific gravity of 0.800? 3. If an object has a density of 7.2 g/mL, what is its sp. gr.? 4. What is the density (g/cm3) of solid 75.84 gram cylinder that is 4.00 cm tall and 2.00 cm

wide?

5. Calculate the diameter of a 16.0 pound shot put. A shot put is a solid metal ball thrown by track athletes and it has a sp. gr. of 7.78. Calculate the diameter in inches.

23

IV. Accuracy and Precision: Methods of expressing laboratory error: In many of the laboratories done in a science class, you ultimately end up measuring or determining something that is already ______________. For example, a student may be given an aluminum cylinder and be asked to measure it so that they can determine its density. Why don’t they just look up the “accepted” value of the density of aluminum in the book? Well, the purpose of most introductory laboratories is to teach you the proper experimental techniques, not to discover new information. Once a student has obtained specific values, how do they know how close their measurement is to the true (accepted) value that they are trying to obtain? This is done most simply by determining ________________ and ________________ (percent) error.

A. Introduction: Chemistry is a quantitative science…it uses lots of numbers!! These numbers come from experimental measurements and each measurement has some degree of uncertainty in it.

1. Reasons for UNCERTAINTY in measurements:

(a) ___________________________ = Construction of the device.

(b) ___________________________ = Incorrect usage of the device.

(c) ___________________________ = Temperature, pressure, etc.

2. Methods of expressing uncertainty:

(a) +/- notation

(b) Accuracy and Precision: “The Dartboard Analogy”

accurate but precise but accurate not accurate not precise not accurate and precise not precise

3. Explain how a student can get precise inaccurate measurements.

B. Accuracy: The closeness of a measurement (calculation) to the ________ (True or Accepted) value. It is expressed in terms of _____________.

O = _________________, experimental value A = _________________ (True) value

1. Absolute error (Ea)

a. Ea =

2. Relative error (Er) (Percent % error)

a. Er =

24

Example # 1: A student measures the length of an object to be 7.45 cm. If the actual length is 8.000 cm, calculate the absolute and relative error of his measurement.

Question: Explain why relative error is more useful than absolute error? Give an example.

Example # 2: Calculate the density of the aluminum block.

If the accepted value for the density of aluminum is 2.70 g/cm3, calculate the absolute and relative error of this value.

Example # 3: Calculate the density of this cylinder: m = _______________ g h = _______________ cm d = _______________ cm

If the accepted value of this cylinder is __________ g/cm3, calculate the absolute and relative error of this value.

Example # 4: Calculate the density of this sphere:

m = _______________ g d = _______________ cm

If the accepted value of this cylinder is __________ g/cm3, calculate the absolute and relative error of this value.

l =______cm

h =______cm

w =______cm

mass = _____________

25

V. Elements: A. A substance in which all of the ____________ have the same number of protons in the

nucleus. B. ______________________ (a Swedish chemist): generally given credit for creating the

modern symbols of elements. C. Atomic number: Number of _________________. D. Approximately __________. E. 90 elements occur _________________ (elements 43 & 61 are man-made) F. 93 and beyond are ___________________ transuranium elements.

G. Names & symbols: 1. Carbon – _____; Calcium – _____; Chlorine – ______ 2. 104 and beyond Unnilquadium (Unq) Un = _____ nil = _____ quad = _____ 3. Elements named after _____________ (old). Examples: a. Sodium __________ _______________________

b. Gold __________ _______________________

c. Silver __________ _______________________

d. Potassium __________ _______________________

e. Lead __________ _______________________

f. Antimony __________ _______________________

g. Iron __________ _______________________

h. Tungsten __________ _______________________

i. Tin __________ _______________________

j. Copper __________ _______________________

k. Mercury __________ _______________________

H. States of matter 1. solid (s) 2. ________________ 3. gas (g) I. ___________________: 2 or more elements or compounds physically joined.

1. Alloy: Brass: _______________

Bronze: _______________ J. ___________________: 2 or more elements chemically bonded to one another. Ex. : O : H H

26

K. Elements: ______________________________________ into simpler substances by chemical or physical means (Excluding nuclear processes).

L. Periodic Table:

1. Vertical columns – ____________________________________ (Chemically similar)

2. Rows - ______________________

3. Important Group Names: a. Group IA(1): ___________________________ (except Hydrogen) b. Group IIA(2): alkaline earth metals c. Group VIIA(17): ____________________________ d. Group VIIIA(18): noble gases e. _______________ ; Metalloids; ________________ f. Most elements are _________________ ( sulfur, carbon, and sodium) g. Some are gases, & a few are liquids at room temperature (i.e. Br & Hg) VI. Properties of Matter A. Substances: 1. Substances can be either pure ______________ or _______________. 2. Substances are ________________ by enumerating their physical and chemical

properties. 3. All specimens of a given substance will have the ____________ chemical and

physical properties.

4. Examples of substances: elements compounds B. Mixtures: 1. Mixtures: Two or more pure _________________. a. Homogeneous mixtures __________________________________________ b. Heterogeneous mixtures__________________________________________

2. Examples of mixtures: Homogeneous Heterogeneous

27

C. Physical Properties of substances: Include those features, which definitely distinguish one substance from another.

1. Density- ____________________________ 2. Specific gravity- ______________________________

3. Hardness- Ability to resist scratching. MOH hardness scale is used as a basis of comparing the hardness of substances.

a. Low numbers = relatively ________ substances. b. High numbers = relatively ________ substances.

MOH scale: Talc 1 Gypsum 2 Calcite 3 Fluorite 4 Apatite 5 Feldspar 6 Quartz 7 Topaz 8 Corundum 9 ___________ 10

4. Odor- a. Good smells: ___________________________________________________ b. Bad smells: ____________________________________________________

5. Color- ______________________________________________________________

6. Taste- __________________________________-___________________________

7. Solubility in solvents Example: a. ____________ soluble vitamins b. ____________ soluble vitamins

8. Physical state: _______________________________________________________

M.P. _________________ F.P. ___________________B.P. ___________________ 9. Properties of metals- a. Malleability ____________________________________________________ b. Ductility _______________________________________________________ c. Conduction of heat ______________________________________________

10. Accidental physical properties: Not used to _________________ a substance. Examples: _________________________________

28

D. Chemical Properties: Describe the ability of a substance to form ___________________ under given conditions.

1. Chemical change: A change from one substance into another.

2. Chemical properties may be considered to be a listing of all the chemical ____________ of a substance.

E. Changes in Matter: Matter undergoes ______________ and ______________ changes.

1. Energy may be defined as the ability to ___________________.

2. Forms of energy: _____________________________________________________

3. Matter always possesses ______________ in one form or another.

F. Physical Change: The composition of a __________________ is not ________________ and the substance _________________its own identity.

Examples: (1) _______________________________

(2) _______________________________

(3) _______________________________

(4) _______________________________

(5) _______________________________

(6) _______________________________

G. Chemical Change: The substance loses its _______________, and the new substance formed has new _______________ and _______________ properties.

Example reaction: Evidence of chemical changes: (1) ___________________ (2) ___________________

(3) ___________________ (4) ___________________

Examples: (1) ______________________________________

(2) ______________________________________

(3) ______________________________________

(4) ______________________________________

(5) ______________________________________

(6) ______________________________________

29

ATOMIC THEORY ________________ is composed of tiny _____________________ called ___________. Chart: I. Historical Perspective: A. Early Ideas

1. Democritus (400 B.C.) a. ____________ philosopher … b. Matter is composed of _______________ _________________ called atoms. c. “atomos” means ____________________. d. His theory was forgotten because … 1. _________________ and __________________ disagreed. 2. He had no _____________________ _____________________ a. He used _____________ not _______________________.

2. Alchemy a. 2 goals: 1. Elixir of ____________. 2. Transmutation: Lead into ___________.

b. Importance of alchemists is that they shifted from _____________ to observation and _________________!

3. 1500’s – 1700’s: With the development of true _______________ methods, scientists discovered many important things about matter.

a. Such as: electricity, magnetism, chemical reactions; this information would help establish important scientific principles that would be used to develop the ________________ Theory.

4. Dalton’s Atomic Theory (1808) a. ____________ Dalton reproposes the atomic _________________ and supports his ideas with chemical ________________. (Look on study guide for postulates).

b. Three Important Laws Dalton based his Atomic Theory:

i. Law of ____________________ of Mass: Mass is neither ________________nor ________________.

ii. Law of Constant Proportions (__________________): Compound always contains the same elements in the same __________________ by mass. Example: H2O

iii. Law of Multiple Proportions (______________________): When two elements can form multiple compounds, the ratio of masses will remain constant for each compound. Example: CO : CO2

30

5. Thomson Model of the Atom Theory

a. _______ Thomson discovered that atoms are made of _______________, in other words they are made of ___________ things.

b. J.J. did experiments with __________________________ (CRT) and he found that:

i. Cathode rays are _____ particles that he called _______________. ii. This showed that atoms are not _____________________! iii. He determined the _______________________ ratio of the electron. (_____________________)

iv. He knew atoms were _______________ so he proposed a model of the atom called ____________________. The plums were _________. v. Robert A. _______________ determined the charge of the _____ with his oil _____________ experiment.

6. Rutherford Model of the Atom Theory a. The discovery of _____________ led to further advances in the atomic theory. b. New Zealand physicist Ernest B. __________________ and his associates

(Geiger & Marsden) used radioactive ______ particles to probe the _________. c. He discovered the ________________________ with the __________ scattering experiment.

d. Diagram of the experiment:

(a) Experimental setup for Rutherford’s experiment: α particles emitted by radium strike metallic foil and some rebound backward; (b) backward rebound of α particles explained as repulsion from heavy positively charged nucleus.

31

e. His calculations regarding the deflected particles indicated that atoms have a _________, ______________, _______________________.

f. If the atom were a ________ in diameter, the nucleus would be the size of a

___________…yet the nucleus contains virtually all of the atom’s __________. In other words most of the atom is made up of _________ _________.

g. He proposed the still popular (yet wrong) planetary model: i. electrons _________ the positive _____________.

ii. like the ______________ around the _________.

h. Problems with a planetary atomic model; It could not explain: i. Electron ___________: classical physics theory says that a charged

particle (like an electron) moving in a circular orbit would _________ energy and slow down, eventually collapsing out of its _________ and ___________ into the_________.

ii. Periodic _____________________ behavior.

iii. Atomic ___________ spectra.

7. The ___________ model of the atom: proposed by Danish physicist, _________________, would first explain some of these problems.

II. Atomic Structure:

A. 3 particles

1. Proton (p) a. located in the _____________ b. unit charge = _______ c. mass = 1.67265 x 10-24 g d. relative mass = _____ (relative to the other particles) e. charge = + 1.6022 x 10-19 C (Coulomb) f. the mass is 1,836 x the mass of a _____________. g. discovered by ___________________ h. made of _____ quarks:

2. Electron (e)

a. located outside the nucleus in ____________ levels (shells). b. unit charge = _____ c. mass = 9.11 x 10-28 g (9.10953 x 10-28 g) d. relative mass = ______ (tiny compared to n & p) e. charge = −1.6002 x 10-19 C (Coulomb) f. the mass is 1/1,836 the mass of a ___________. g. discovered by ______________.

u

u

d

(Proton)

32

3. Neutron (n) a. located in the _____________. b. unit charge = _______ “neutral” c. mass = 1.67495 x 10-24 g slightly more massive than a proton. d. relative mass = ____ e. discovered by ________________ in 1932 (Why so late? ________________) f. made of ____ quarks:

B. Nucleus 1. The central ________ of the atom. 2. Contains the neutrons and protons:____________ = a particle in the nucleus (n or p) 3. _________ = the nucleus of an atom. 4. Contains almost all the _________ of the atom. 5. Has a __________ volume compared to the _________ atomic volume.

Ping pong ball in the _______________.

6. Density of nucleus = _________________ g/cm3. 7. Radionuclide: an unstable __________. Why are some nuclei stable and others

are unstable? ________

8. Nuclear stability is due to: a. Nuclear binding _____________________ ∗Holds protons and neutrons together. b. n/p ratio = stable atoms have a favorable n/p ratio.

Examples: _________________________________________

_________________________________________

C. Atomic Size 1. Atomic diameters = ____________________________________

1 Angstrom = ___________; 1 picometer = __________

2. Diameter of atomic nuclei = _________________________ 3. How many carbon atoms are there in a pencil line 1.00 inch long?

(radius = 0.77 Angstroms; assume 1 atom wide)

d

u

d

(Neutron)

33

III. Mass Relationships in atoms A. Atomic Number (Z) 1. Equal to the number of ___________. a. Equals the number of ________________ in a ________________________. i. Example: (a) Oxygen (b) Oxide ion (ions: charged particles) b. Each type of ______________ has a specific number of protons…this

determines the element’s _________.

B. Mass Number (A) 1. Equals the number of _____________ + the number of ______________. a. It is the number of ____________. 2. # of neutrons = ________________________________________________ 3. Correct notation: C. Isotopes: Atoms with the same _________ number but different ___________numbers. 1. Same number of ___________ 2. Different number of _____________. 3. Isotopes are named by their ________________. a. Carbon-14 b. Uranium-238 4. Example: The three isotopes of Hydrogen 5. Example: The three oxygen isotopes 6. Ions: _____________ atoms (or groups of atoms) that have lost or gained electrons. Examples: IV. Nuclear Chemistry: Chemistry of the nucleus (______ + ______). A. Radioactivity: The ___________________ emission of _______________ or EMR (_____________________) from the nucleus. 1. Henri Becquerel: Discovered radioactivity (1896) using _____________________

and uranium ore. 2. Types of Radioactivity: a. Alpha (_____): Nucleus of a helium atom. (_________) b. Beta (_____): High speed electron emission from the __________. (______) c. Gamma ray (______): Photon of high energy light. (_________)

34

3. Penetrating power of radiation: a. Alpha can go through _____________. b. Beta particles can go through _______________________. c. Gamma rays can go through ________________________. ________>________>________ B. Nuclear equations: Must obey Law of Conservation of __________ (top line) and Law of Conservation of _________ (bottom line). 1. Alpha emission (__________)

2. Beta emission (________or________)

a. Electron is formed in the _______________. b. Nucleus has too many __________________.

c. Neutron spontaneously becomes a ______________, which causes a high energy electron to be ejected from the nucleus.

3. Positron emission: (________or________) a. Positron is formed in the _______________. b. Nucleus has too many ___________________.

c. Proton spontaneously becomes a ________________, which causes a high energy positron to be ejected from the nucleus.

35

4. Gamma emission (_________) a. Nucleus is in an _______________ energy state (excess energy from another decay).

b. As nucleus loses energy, a _________________ is emitted. c. Asterisk (*) is used to symbolize __________________________. 5. Electron Capture a. An electron from the innermost energy level “falls” into the nucleus. b. Electron capture is more common with __________ nuclei. c. The product is the same as that of ______________ emission. 6. General examples: a. β, β decay

b. Uranium decay series: Radon-222 - α,α,β,β,α decay C. Decay Series: A radioactive decay often results in a ________________ nucleus that is

also radioactive. A radioactive _____________ _____________ refers to successive decays, which starts with one parent isotope and proceeds through a number of daughter isotopes. The series ends when a stable, _____________________ isotope is produced.

36

D. Nuclear Stability 1. Belt of Stability (FIGURE 1)

2. Binding Energy per Nucleon versus Atomic Mass (FIGURE 2)

3. Why are some nuclei stable and others are not?

Rule 1: The greater the binding energy per nucleon (energy holding the nucleus together), the greater the stability. (See Figure 2).

Which isotope is the most stable, according to Fig. 2? ___________

Figure 1: Stable and unstable nuclides

Figure 2: A plot of nuclear binding energy per nucleon versus mass number. Nuclei with large binding energies per nuclei are the most stable.

37

Rule 2: Nuclei of low atomic numbers with a 1:1 ratio of neutrons to protons are very stable. (See Fig. 1) Example: Carbon-12 Helium-4

a. Isotopes decay with ________. b. Isotopes decay with ________. c. Isotopes decay with ________. ** Radioactive isotopes decay until they reach the “Belt of Stability.”

Rule 3: The most stable nuclei tend to contain an ___________ number of both protons and neutrons.

Example: Iron-56 Oxygen-16

E. How is Binding Energy determined? 1. Nuclear binding energy: The energy required to break up a nucleus into its component ____________ and ___________ (nucleons).

2. Binding energy comes from the mass defect of the nucleus.

3. Mass Defect: The total mass of the stable nucleus _________ the sum of the masses of the nucleons. The “missing” mass has converted into energy! (________________)

Masses of subatomic particles e- 0.00054858 amu p+ 1.007276 amu n 1.008665 amu (p+ + e-) 1.007825 amu

4. Information for conversions between mass and energy:

1 g = 6.02 x 1023 amu 1 Joule (J) = c = 3.00 x 108 m/s

Ex. 1: Determine the energy in Joules of a mass defect of 1.00 amu.

Refer to Figure 1

Note: The whole atom is NOT the sum of its parts. MISSING MASS

Note: When calculating masses, include the mass of the electrons because the mass of the whole atom includes electrons.

Mass defect = Atomic mass of the isotope - Σ mass of subatomic particles.

Relationship between mass and energy: E = mc2

38

Ex. 2: a. Determine the mass defect of Fluorine-19.(The atomic mass is 18.9984 amu).

b. Calculate the binding energy in J (use conversion factor between mass and

energy). c. Calculate the binding energy per nucleon. 5. As nuclear binding energy per nucleon increases, the stability of the nucleus

_____________.

6. The element with the greatest binding energy per nucleon is ______________. F. Half Life: The average time it takes for ____________ of the unstable atoms in a sample

to decay. 1. Exponential decrease of atoms _________________________________________. 2. How much of a 500 g sample of Uranium-235 would be left after five half-lives?

39

3. A 16.00 mg sample of Radon-222 decays to 0.250 mg after 24 hours. Determine the half-life.

4. The half-life of molybdenum-99 is 67 hours. How much of a 1.000 mg sample is left

after 335 hours? G. Transmutation Reactions 1. Bombardment: Bombard a target nucleus with other _________________ or nuclei. Example particles: _____,_____,_____,_____,_____ 2. Particle + nucleus → usually more than one _______________ formed. 3. Examples: (a) (b) *Chadwick – Discovery of the neutron (c) *Rutherford (1919)-First artificial transmutation (d) 4. Shorthand notation: (d = deuterium = ; T = tritium = ) (a) (b) (c) (d) H. Fusion: Combining smaller nuclei to form _______________ nuclei. 1. The combining of small nuclei to form larger nuclei increases the ____________ per

nucleon. Therefore, stability _______________.(See Figure 2: binding energy)

2. Nuclei smaller than _____________ will give off energy when they fuse together. 3. Example: Sun reactions (made up mostly of H and He; T in interior ≈15 million °C) ______ ______ ______

4. When products are formed in a fusion reaction, mass is __________. This mass __________ turns into _________________.

5. Fusion reactions take place at very high _________________.

40

6. They are often called _____________________. 7. Reasons for high temperature: a. Positive nuclei ___________ one another. b. High _____ or speed needed to fuse the elements to overcome the repulsion. c. High _______ = High temperature (Millions of degrees Celsius) I. Fission: Large nucleus is ________________ into smaller nuclei + one or

more_______________.

1. Example: 2. Many fission products are formed. More than 30 different elements have been found

among the products.

3. Diagrams: 4. The ______________ formed in the initial stages can induce ______________ in

other .

5. Nuclear fission takes place for elements with very large nuclei. Elements larger than _______can undergo fission. (See Figure 2).

6. Practical applications of fission: naturally occurring __________________ and artificial _______________________.

7. The mass of the starting material is __________ the sum of the masses of the products. The mass defect (missing mass) turns into _____________. 8. Nuclear chain reaction: a _____________________ sequence of nuclear fission reactions.

Critical mass: the minimum mass of fissionable material needed for a ________________ chain reaction.

Subcritical mass: Less than the minimum mass. Too many neutrons will escape (will not be absorbed by the fissionable material), so _______ chain rxn will occur.

41

9. Applications: a. Atomic Bomb: Self-sustaining nuclear chain reaction. i. Must contain __________ critical mass of U-235 (or Pu-239). ii. Critical mass must be kept in ____________ places before detonation! iii. Neutrons originate from a source at the center of the bomb_________

the rxn.

b. Nuclear Reactors: Controlled fission reaction. (8% of electricity in the U.S.). i. Advantage: Source of energy other than fossil fuels. No global

warming gases. ii. Disadvantages: Produces highly radioactive products. Example: Strontium-90 (radioactive) = chemically like

___________. (concentrates in the _____________).

Iodine-131 = ______________ cancer. (within 10 mi. radius of nuclear power plants → I pills) J. Nuclear reactors: A device in which the controlled ____________ of a certain substance

is used to produce new substances and ___________________!!! 1. _________________: Using the neutrons released during fission to cause other

nuclei to undergo fission.

2. Self-sustaining ___________ reaction keeps on going once it has begun. 3. The neutrons released in a typical fission can be used to cause other nuclei to

fission. If enough _________ or _________ is available, it is possible to create a __________________________. The minimum amount of fissionable material required for a self-sustaining chain reaction is known as the _______________ ___________. The critical mass depends on a number of factors, among the factors:

a. ________________: slows neutrons down which increases the # of favorable collisions. The most common are _________________ (Chernobyl) and __________________ (D2O)(Used in Canada)(No need for enriched uranium). b. _________________ ie. 235U & 239Pu

c. _______________________: The % of fission material is increased (by diffusion or centrifugation). d. In _____________ _____________ the chain reaction is controlled and the energy is released gradually. In an “atomic bomb”, the chain reaction is uncontrolled and the energy release occurs in a few moments of time.

4. __________ Power Plant

a. _______________(Cadmium/Boron) = Absorb neutrons (Stops rxn—Applying the brakes).

b. Many use Uranium – 235 for fuel; fuel rods contain fissionable material.

42

c. Where does the electricity come from? __________occurs in the nuclear reactor core → __________ is produced → heat causes liquid water to turn to ____________ → steam causes the _____________to turn → ________________ is generated → steam is turned back into liquid water in the _______________ (which uses cold water from a river, lake, or ocean)→ warm water (non-radioactive) is sent back into the river, lake or ocean.

Demo: Turn turbine—electricity is generated. 5. ______________ Reactor: One in which fissionable material is produced at a

greater rate than it is consumed. a. Converts ______________________ into ___________________ b. Where does the ___________ come from? 6. Problems of nuclear power. a. Thermal pollutions: Hot water exits into the ocean. b. Leakage of fission fragments: Highly radioactive. Examples: Chernobyl (Ukraine), Three Mile Island (USA) c. Spent fuel disposal: Yucca Mtn, NV. d. Only lasts for 30 years. Why?

i. The build up of __________________. ii. The structural materials ________________.

43

K. Radioactive Dating: 1. If the half-life of a radioactive isotope is known, an estimate of the _________ of an

object often times can be made. For example, the ratio of _____________ to carbon- 12 in a living object is relatively constant. However, a living object _________ ____________ carbon-14 when it dies. Thus, knowing the half-life of carbon-14 (_____________) and the object’s 14C/12C ratio, an estimate of the object’s age can be made. Willard __________ won the 1960 Nobel Prize in chemistry for development in 1947 of radiocarbon dating.

2. Formation of Carbon-14: _______________________________________________ 3. Beta minus decay of Carbon-14: ________________________________________ 4. Useful for dating objects about ___________________ years old. 5. How about older things? ___________________

6. Geologists determine the amount of __________ remaining in a rock relative to the amount of daughter nuclei present to estimate the passage of time since the rock solidified from molten material.

44

Unit 3: Inorganic Nomenclature I. Background: A. Periodic Table

1. Column: __________________ or __________________. (Similar properties)

2. Row: __________________.

3. _____________: Left of staircase (Majority of the elements).

4. ___________________: Right of staircase.

• Exception: ______

5. ___________________: Touching the staircase.

• Exception: ______ (metal).

B. Ions (Charged atoms)

1. _____________: positively charged (lost e-).

2. _____________: negatively charged (gained e-).

C. Trends in the periodic table

1. Using the planetary model – (simplified model of atom)

2. Energy levels can contain a maximum of:

1st energy level: _____ 2nd energy level: _____ 3rd energy level: _____ (_____)

3. _____________ are the keys to chemical bonds

Ex.

Column IA (1) (_______________) Column VIIIA (18) (__________) H (____ e-) He (____ e-) Li (____ e-) Ne (____ e-) Na (____ e-) Ar (____ e-) Similarities:_______________ _____________________

45

4. Atoms can gain or lose _____ to achieve a full outer shell (more stable).

5. Atoms will do what is ______________ (least energy) i.e. Oxygen has 6 valence e-:

easier to ___________ than to ____________.

Group # of valence e- Gain or lose e- Charge IA

IIA

IIIA

IVA

VA

VIA

VIIA

VIIIA

6. Label your periodic table!

II. Binary Ionic Compounds A. Background info

1. Metal / ________________ (_______________ is always written first!)

2. One element ______________ and the other _________________.

3. _________________ of e-

4. Charged ions ______________ one another (opposites attract).

5. Compound is ______________.

B. Examples of compound formation of binary ionic compounds:

Ex. Sodium & chlorine

Ex. Calcium & bromine

Ex. Lithium & oxygen Ex. Aluminum & sulfur

46

C. Shortcut to determining binary ionic compounds (Criss-Cross method):

1. ________________ from charge becomes the subscript.

2. All ionic compounds are ______________ (no + or -).

3. Subscripts are written in the _______________ possible ratio.

4. The number “1” is never written (It is implied).

5. Examples:

Lithium & oxygen Aluminum & oxygen

Calcium & oxygen Magnesium & nitrogen

D. Nomenclature of binary ionic compounds (bi = 2).

1. ____________ is named first (name of atom).

2. ____________ is named second, ending changed to ___________.

3. If the metal (cation) can have multiple charges, the charge is written as a roman

numeral (IUPAC) or as the common name. (Fe, Cu, Co, Hg, Mn, Sn, Pb)

4. Formula to name:

a. Li2O _______________________________________

b. Al2O3 _______________________________________

c. CaO _______________________________________

d. Mg3N2 _______________________________________ e. Fe2O3 _______________________________________

f. SnO2 _______________________________________

g. CuCl _______________________________________

h. MnN _______________________________________

47

5. Name to formula:

a. Beryllimum fluoride _______________ _______________

b. Potassium bromide _______________ _______________

c. Tin (II) oxide _______________ _______________

d. Cobaltic sulfide _______________ _______________

e. Strontium iodide _______________ _______________

6. Polyatomic Ion: A group of atoms with a _________ charge.

Ex. (1) CN- =

(2) NH4+ =

(3) OH- =

a. Polyatomic ions will _______________ stay together as a group.

b. If there is more than one polyatomic ion, it must be placed in_____________.

Examples:

Ions Formula Name

48

III. Helpful Hints to Memorize Oxyanions: A. In learning the formulas and charges of common oxyanions, start with the –ate form. From

there it follows that: hypo______ite = 2 less oxygens

______ite = 1 less oxygen

______ate

per_____ate = 1 more oxygen

**ALL forms have the SAME charge!**

A Guide to Determine Whether the –ate Formula is –XO3 or –XO4

A Guide to Determine What the Charge of the Oxy-Anion is:

49

B. Examples:

Borate = _____________ Carbonate = _____________

Nitrate = _____________ Chlorate = _____________

Nitrite = _____________ Perhlorate = _____________

C. “Thio-“ = Sulfur replaces an oxygen.

Ex. Sulfate = _____________ Thiosulfate = _____________

Ex. Cyanate = _____________ Thiocyanate = _____________

IV. Ternary Ionic Compounds: (compounds containing _____ or more elements).

1. Name the _______________

2. Find the appropriate name of the ____________________________________.

3. Formula to name:

a. Li2SO4 ____________________________

b. Fe(NO3)3 ____________________________

c. CdC2O4 ____________________________

d. Cu3AsO3 ____________________________

e. Mn2SiO4 ____________________________

f. (NH4)2SO4 ____________________________

4. Name to formula:

a. Potassium thiocyanate _______________ _________________________

b. Aluminum permanganate _______________ _________________________

c. Plumbic acetate _______________ _________________________

50

d. Cobalt (III) oxalate _______________ _________________________

e. Sodium hypochlorite _______________ _________________________

V. Nomenclature of Hydrates A. Hydrate: Ionic compound with ______________ molecules stuck in the ___________

lattice. The water is included in the ____________ and formula.

1. ZnSO4 · 7 H2O _______________________________________

2. CaCO3 ·3 H2O _______________________________________

3. Cu2C2O4 · 2H2O _______________________________________

4. Calcium chloride pentahydrate _______________ _________________________

5. Cupric acetate monohydrate _______________ _________________________

VI. Binary Molecular Compounds A. Molecular (___________________) compounds

1. Non-metal to _______________________.

• ____________________ of staircase including hydrogen

2. ___________________ of electrons.

Ex

3. Non-metals can often combine in several different ways.

Ex.

B. Nomenclature of binary molecular compounds:

1. Greek prefixes are used:

mono = hexa = di = hepta = tri = octa = tetra = nona = penta = deca =

2. The prefix “___________” is omitted for the 1st element.

Ex. CO = __________________________

3. For oxides the ending “_____ or _____” is omitted.

a. N2O = ____________________________

b. N2O3 = ____________________________

c. N2O4 = ____________________________

d. NO = ____________________________

e. NO2 = ____________________________

f. NO5 = ____________________________

51

Compound

Ionic Covalent

1. 1. 2. 2. 3. 3.

Ex. 1. P2O5 ______________________________

2. NCl3 ______________________________

VII. Nomenclature of Acids A. Acids: Compounds that contain _________________ as the positive ion (H+).

B. Exceptions: ___________ (water) & ___________ (hydrogen peroxide).

C. Binary Acids: Acids that ______ ________ contain oxygen.

1. Use prefix “___________”

2. Add stem or full name of _______________.

3. Add suffix “_______”.

4. Add the word ________________.

Ex. HBr = __________________________________

HCl = __________________________________

HCN = _________________________________

D. Ternary Acids: Contain ____ or more elements __________________ oxygen.

1. Acids formed with anions that contain ____________ become __________ acids.

HNO3 (NO3- = ______________) __________________________

HClO4 (ClO4- = ______________) __________________________

H2SO4 (SO42- = ______________) __________________________

H3PO4 (PO43- = ______________) __________________________

2. Acids formed with anions that contain _________ become ___________ acids.

HNO2 (NO2- =______________) __________________________

HClO2 (ClO2- =______________) __________________________

H2SO3 (SO32- =______________) __________________________

3. Name to formula:

a. cyanic acid _______________ _________________________

b. dichromic acid _______________ _________________________

c. hypochlorous acid _______________ _________________________

d. hydrosulfuric acid _______________ _________________________

52

Unit 4: The Mole I. Atomic Weight (Mass)

A. Atomic Mass Scale: 1. Reference Unit: The _______________________ atom. a. Symbol b. 1 ________ atom = ______________ (exactly) c. amu = ______________________________ u (_______________________) d. 1 amu = __________ the mass of a _______________ atom.

B. Mass of atomic particles in amu 1. electron 9.11 x 10-31 kg or ____________________ amu. 2. proton = __________________ amu. 3. neutron = _________________ amu.

C. Atomic mass (Weight): An atom’s mass (_____________) in atomic mass units (amu). i.e. Carbon-12: 12.00000 amu 12.011 amu (Chart value)

1. The atomic weight found on the periodic table is the ______________ of the naturally occurring ___________ of an element. (Look on chart)

2. Calculation of Atomic Mass (Weight) you must be given …

a. The _________ of each ____________. b. Abundance (%) in _________ (varies slightly by location) 3. Calculate the atomic mass of carbon. Isotopes % abundance Atomic mass Carbon-12 98.890 12.00000 amu Carbon-13 1.1100 13.00335 amu 4. The mass and % abundance are determined with an instrument called a

__________ _____________. The abundance can _________ depending on where the sample is from. This is why the sig. figs. vary on the chart.

D. Other calculations:

1. Lithium has two isotopes. If lithium-6 has a mass of 6.015123 and 7.42 % occurrence, what is the % abundance and mass of lithium -7?

2. What is the % abundance of the two isotopes of chlorine if the atomic weight of chlorine on the periodic table is 35.453? Chlorine-35 has a mass of 34.968853 amu and Chlorine-37 has a mass of 36.965903 amu.

53

E. Mass Spectrometry: Is the process involving the production of _________________________ from a sample, and their resulting ________________ (and __________________) according to their __________ to __________ ratio. (m/q)

1. 3 distinct operations involved in mass spectrometry.

a. __________________ of _______________ why? (________________) b. _________________ (by ____________ or _____________ fields) c. __________________ of + ions (photographically or electronically)

2. Key Ideas: __________ + ions are deflected ___________. Results give the _________ and % ________________ of each isotope. (See diagrams & book for more info).

II. Molar Mass & Avogadro’s Number

A. Definitions: 1. Molar Mass (Gram atomic weight): the mass in _________ that is numerically equal to the mass in _____.

a. mass of 1 carbon atom _______________ b. mass of 1 mole of C _______________ 2. Mole = ___________________________ of anything. 3. Avogadro’s number = _________________ 4. Molar Mass = The mass in grams of ____________ (6.02 x 1023 atoms or

molecules).

5. What is the relationship between atomic mass units and grams?

B. Calculations based on the mole. 1. What is the mass of a Titanium atom in amu? Grams? 2. What is the mass of one mole of Titanium?

54

3. How many (a) moles and (b) atoms of sulfur are there in 255 g of Sulfur? 4. How many (a) moles and (b) grams of Uranium are there in a billion U atoms? 5. How many silver atoms are there in a solid silver ball that is 2.00 cm in diameter?

(DAg = 10.5 g/cm3) 6. How many Cu atoms are there in a solid copper penny that has a mass of 3.06 g? 7. What is the mass of one atom of silicon in (a) amu (b) grams (c) pounds? 8. What is the mass of 1500 sulfur atoms in grams? 9. A student measures a solid aluminum block: l = 4.84 cm, w = 3.68 cm, h = 2.14 cm.

If the specific gravity of aluminum is 2.70, how many aluminum atoms does it contain?

55

10. Mercury has a specific gravity of 14.7. How many moles of mercury are in a bottle that contains 3.55 pints of mercury?

III. Molecular Mass & Percent Composition

A. Chemical Formulas: Express the composition of molecules (and ionic compounds) in terms of the ___________ for the elements they contain. B. Molecular Formula: Show the ________ and _________ numbers of ________. They are the _________ formulas of molecules! C. Diatomic Molecules: Memorize the 7 elements that occur as diatomic molecule…

D. Molecules containing more than two atoms are called _____________________________.

Examples:

E. Different forms of the same element are called _________________ Examples:

F. Empirical (simplest) Formula: Show the ________ and ________ of atoms.

Water Glucose Hydrogen peroxide Benzene MF

EF

G. Ionic compounds only have ______________ formulas, they exist as 3D lattice ______________ but not as individual ______________. When _______________ are removed or added to a neutral atom (or group of atoms) a charged particle is formed called an _______. A positive ion is called a __________ while a negative ion is called a ________. Substances made of cations and anions are called ___________ _______________. Ions that contain only one atom are ______________ while those containing two or more are called __________________. H. Structural Formula: Show the kind, number, and arrangement of atoms.

water ethane

56

IV. Molecular Mass (weight) A. The molecular mass is the ________ of the atomic masses in (amu) of all the _________ in a molecule. The molar mass is the same mass but in _________.

1. What is the molecular mass of benzene, C6H6? What is the molar mass?

B. Substances that do not form molecules such as ionic solids do not have a molecular weight; they have formula weights.

1. What is the formula mass of aluminum sulfite? The molar mass? 2. Calculate the molecular mass of TNT, C7H5N3O6. 3. What is the formula mass (amu) of Cr(OH)2 • 18 H2O?

C. Percent Composition: The percent by ________ of each ___________ in a compound.

1. What is the % composition of Ba(CN)2?

2. What is the % composition of ___________________, Fe2(SO4)3?

3. What % hydrogen is water?

4. % H2O in a Hydrate: Ionic substances are usually formed in water solutions. The water molecules get trapped in the crystal structure and form ____________.

Examples: Calcium chloride dihydrate Plumbic nitrate heptahydrate

57

5. Ba(CN)2 • 2H2O is what % water by mass? V. Formula Determination:

A. Review Problems: 1. What is the % composition of Ca(OH)2? FW = 74.10

2. Ca(OH)2 • 2 H2O is what % water by mass?

B. Determination of Empirical Formula: 1. Laboratory analysis shows that a compound is 5.80 % H and 94.2 % S by weight (mass). What is its E.F.?

Step 1: Calculate the moles of each element present. (Assume 100 % = 100 g) Step 2: Determine the ratio between moles (smallest whole numbers) by dividing

by the smallest value in step 1. 2. Remember that these values are measurements and need not be exact. They should be close. What would be the formulas for elements “X” and “Y”. X1.5 Y1.0 X1.00 Y1.03 X1.00 Y1.33 X1.00 Y1.25

3. (a) Laboratory analysis shows that a compound has the following composition: Zinc 52.0 %, Carbon 9.60 %, & Oxygen 38.4 %. Determine its empirical formula.

(b) What is the empirical formula of a compound that is 52.9 % Al and 47.1 % oxygen?

58

C. Determination of the molecular formula: 1. You must be given the ____________________ weight. 2. What is the relationship between the empirical formula weight and the molecular

formula weight? 3. Molecular formula = Empirical formula x __________________

4. Analysis shows that a compound is 30.4 % Nitrogen and 69.6 % Oxygen. If it has a molecular weight of 92 g, what is its Molecular formula?

Step 1: Calculate the Empirical formula. Step 2: Calculate the Empirical formula weight. Step 3: Divide the Empirical formula weight into molecular weight (must be

given). Step 4: Determine the Molecular formula.

5. Analysis of a compound shows it to be 38.71 % Carbon, 9.68 % Hydrogen, and 51.61 % Oxygen. If further data reveals that it has a molecular weight of about 93 g, determine the Empirical and Molecular formulas of the compound?

D. Determining the Formula of a Hydrate: 1. Hydrate: Is an __________ compound that has _______ molecular trapped in its

crystal _________ structure. The formula of the hydrate can be determined by heating up a known mass of the hydrated salt in a crucible and then measuring the mass of _____________ salt formed.

59

2. A barium iodide hydrate is heated in a crucible. What is the formula of the hydrate? (BaI2 • x H2O)

10.407 g BaI2 • x H2O (Hydrated) + Flame → “Anhydrous Salt” 9.520 BaI2 + H2O

Step 1: Find the grams of water lost. Step 2: Find the moles of water and anhydrous salt. Step 3: Find the mole ratio.

3. In lab a student heats 5.20 grams of hydrated zinc sulfite in a 20.40 g empty crucible. After heating, the anhydrous ZnSO3 compound and crucible have a combined mass of 23.00 g. What is the formula of the hydrate?

VI. MOLES II: Moles & Formula

A. General Information: 1. 1 mole = _________________________ number of anything. Avogadro’s number = ________________________________. 2. 1 mole of atoms = ____________________________ atoms. 1 mole of molecules = ____________________________ molecules.

° ° °

°

°

H2O

°

60

3. The atomic weight of oxygen is _________________ amu. a. One oxygen atom has a mass of ________________ amu. b. One mole of oxygen atoms (6.02 x 1023 atoms) has a mass of

_____________. This is called the __________ atomic weight.

4. The molecular weight (mass) expressed in ________ represents the mass of a ___________ molecule, while the molecular weight expressed in ____________ represents the mass of __________ of molecules; this is called the gram molecular weight.

5. Substance the do not form molecules, such as ionic solids, do not have a molecular weight. They have a formula weight that when given in grams is called the gram formula weight and represents the mass of ___________ or the compound.

6. Road Map:

B. Sample Problems 1. 0.500 moles of Fe2(CO3)3 (________________________) contains how many moles of a. Fe b. C c. O

2. Calculate the mass of one water molecule in (a) amu (b) grams 3. (a) How much does 4.50 moles of water weigh in grams? (b) How many molecules is

this?

4. 50.0 g of methanol, CH3OH, are (a) how many moles? (b) contain how many molecules?

5. How much does a mole of _____________ (C6H12O6) weigh in kilograms?

61

6. How many water molecules are there in a pint of water?

7. 50.0 g of ___________________ (Ca(OH)2) contain how many (a) moles of oxygen? (b) oxygen atoms?

8. What is the volume in mL of 5.55 x 1024 water molecules? 9. A spherical metal tank contains carbon monoxide gas under pressure. If the tank

has a diameter of 4.00 ft. and the density of the gas under pressure is 7.50 g/L, how many molecules are there in the tank?

VII. Balancing Chemical Equations

A. Chemical equations must be balanced due to the Law of Conservation of ________. (_______________). Reactants are written on the left of the arrow, ________________ are written to the right of the arrow.

______________________→______________________

B. Equations are balanced by changing ____________________, not by changing ________________. For now all equations will be balanced using ____________ numbers.

C. Certain elements will are found as __________________ molecules in nature. In a chemical equation, they will be written in this form. The elements are easily remembered by learning the name of the German guy:

Mr. __________________________. The seven diatomic molecules are:_____, _____, _____, _____, _____, _____, and _____.

62

Example: Lighting a match: 1. Observations: 2. Word Equation: phosphorus + potassium chlorate → potassium chloride + diphosphorus pentoxide 3. Skeleton Equation: P4 + _________________→__________________+__________________ 4. Balanced Equation:

D. Coefficients: Small, ________________ numbers placed before the chemical formula. They are multiplied by each atom in the compound.

E. Law of Conservation of Mass: Matter cannot be _______________ nor ______________, so reactions must be ________________.

F. Tips for balancing equations: 1. Number of atoms of _______________ must equal number of atoms of __________.

2. Coefficients are whole numbers written at the __________ of the substances. 3. All atoms are ___________________ by the coefficients. 4. Subscripts are ________________ changed. 5. Keep polyatomic ions together as a ___________________ if unchanged from

reactants to products. 6. Balance single elements __________. 7. Use the even/odd rule.

8. If an element is in ________________ compounds, balance that element last. G. Example:

1. _____NaClO3 → _____NaCl + _____O2

2. _____Fe3O4 + _____H2 → _____Fe + _____H2O

H. Sample Problems:

1. Hydrogen + oxygen → water

2. Zinc + hydrochloric acid → zinc chloride + hydrogen

63

3. Copper + silver nitrate → cupric nitrate + silver

4. Ferric hydroxide → iron (III) oxide + water

5. Ethane (C2H6) + oxygen → carbon dioxide + water

6. Calcium + water → calcium hydroxide + hydrogen

7. Potassium + sulfuric acid → potassium sulfate + hydrogen

8. Calcium nitrate + aluminum sulfite → calcium sulfite + aluminum nitrate

9. Phosphoric acid is formed when crystalline diphosphorus pentoxide is dissolved in water.

10. When rust (ferric oxide) is dissolved in hydrochloric acid, it dissolves forming aqueous ferric chloride and water.

11. Benzene, C6H6, is an organic solvent used to dissolve many organic compounds. Write an equation for the combustion of liquid benzene (carbon dioxide and water are formed).

12. When a solution of aqueous plumbic nitrate and aqueous barium hydroxide are mixed, solid plumbic hydroxide and aqueous barium nitrate are formed.

64

VIII. Balancing Chemical Equations: Algebraic Technique

A. Balance the following equation: Ca3(PO4)2 + H2SO4 → CaSO4 + H3PO4

STEP 1: Assign letter to unknown coefficients: a Ca3(PO4)2 + b H2SO4 → c CaSO4 + d H3PO4

STEP 2: Make a grid indicating the appearance of element or ion in each species of the equation. Use whole number and the coefficient to indicate the appearance.

Ca 3a + 0 = c + 0 PO4 2a + 0 = 0 + d H 0 + 2b = 0 + 3d SO4 0 + b = c + 0

STEP 3: Reduce the equations: 3a = c 2a = d 2b = 3d b = c

STEP 4: Assume a = 1; solve for coefficients (**any variable can assumed to be 1) a = 1 b = 3 c = 3 d = 2

STEP 5: Write equation with coefficients:

Ca3(PO4)2 + 3 H2SO4 → 3 CaSO4 + 2 H3PO4 B. Examples: 1. ______Na2CO3 + ______C + ______Sb2S3 → ______Sb + ______Na2S + ______CO2 2. ______C6H6 + ______O2 → ______CO2 + ______H2O

65

3. ______HClO4 + ______P4O10 → ______H3PO4 + ______Cl2O7 4. ____MnO4 + ____CaC2O4 + ____H2SO4 → ____CaSO4 + ____Mn + ____CO2 + ____H2O 5. _____FeS2 + ______O2 → ______Fe2O3 + ______SO2 6. ______C7H6O2 + ______O2 → ______CO2 + ______H2O

66

Unit 5: Stoichiometry I. Stoichiometry = The ____________ relationships among reactants and products in

chemical reactions.

A. A chemical equation is a statement of __________________ fact. On the left

side of the reaction are the __________________, and on the right side the

________________ of the reaction. Since no atoms are created nor destroyed in a

__________________chemical reaction, the equation must be ________________.

This means that the combined weight of the reactants is exactly ___________ to

the combined weight of the products. In terms of chemical laws, this is the

Law of _______________________ of matter (_____________________).

B. Stoichiometric “road map” (Use the balanced chemical equation)

Examples: (1) Calcium reacts with oxygen to form calcium oxide. If you are given 80.0 g of calcium, (a) how many grams of calcium oxide can be made? (b) How many moles of oxygen gas are required?

67

(2) Nitrogen and oxygen react to form dinitrogen pentoxide. If 112 g of nitrogen react with unlimited oxygen, (a) how many moles of oxygen react and (b) how many grams of dinitrogen pentoxide are formed? (3) Aqueous sodium chloride and aqueous silver nitrate react forming solid silver chloride (a precipitate) and aqueous sodium nitrate. If one solution contains 100.0 g of silver nitrate and it reacts with unlimited sodium chloride, how many grams of silver chloride can be made? (4) How many (a) grams and (b) molecules of carbon dioxide gas are produced when 100.0 g of methane gas (CH4) are burned. Methane reacts with oxygen to form carbon dioxide and water. (c) If air is 22.0% oxygen by mass, what mass of air is needed to supply enough oxygen? (5) Nitrogen reacts with hydrogen gas forming ammonia (nitrogen trihydride). How many (a) grams of hydrogen and (b) moles of nitrogen are required to form 1.00 pound of ammonia? (6) How many oxygen molecules are produced if 1.00 gallon of water is electrolyzed into hydrogen and oxygen gas?

68

Review Question: Nitrogen and hydrogen gas combine to form hydrazine (dinitrogen tetrahydride) which is used for rocket fuel. How many grams of nitrogen gas are needed to form 155 g of hydrazine? II. Limiting Reagent A. Stoichiometric amounts: The proportions indicated in the ________________ rxn.

B. Most reactions do not have stoichiometric amounts. Generally, one reactant will be

__________________ before the other. The reactant that is depleted first is known as

the _________________________. The reactant that is left at the end of the reaction is

called the __________________________.

C. Analogy: How to make a cheese sandwich. 2 slices of bread + 1 slice of cheese → 1 cheese sandwich If you have 8 slices of bread and 6 slices of cheese, how many sandwiches can you make? _________ (theoretical yield) What is the limiting reagent? _____________ What is the excess reagent? _____________ How much of the excess reagent is left at the end of the rxn? ___________ D. Theoretical yield: The amount of product that forms if all of the __________________ reagent has reacted. (This number is CALCULATED!) E. Actual yield: The amount of product that is actually made (Experimental). F. Percent yield: The comparison of the actual yield to the theoretical yield. G. Limiting reagent problems (1) How much hydrogen peroxide can be formed from 10.0 g of hydrogen and 125 g of oxygen? What is the limiting reagent?

69

(2) Determine the theoretical yield of water that can be formed from 10.0 g of hydrogen and 100.0 g of oxygen. What is the limiting reagent? (3) How many grams of aluminum oxide can be formed from 100.0 g of aluminum and 100.0 g of oxygen gas? (4) A student does a lab in which she makes 2.17 g of plumbous nitrate. From her limiting reagent she calculated that she should have made 2.55 g. Determine: (a) theoretical yield, (b) actual yield, (c) % yield (d) absolute and relative error. (5) A sample of 100.0 g of hydrazine (N2H4) burns in 280.0 g of oxygen gas and 97.2 g of water are formed. Calculate the % yield of water. (Nitrogen dioxide is also produced.)

70

(6) In lab a student mixes a solution that contains 1.70 g silver nitrate and a solution that contains 0.832 g calcium chloride. Aqueous calcium nitrate and solid silver chloride are produced. After washing, filtering, and drying the precipitate, he finds that he has obtained 1.25 g of silver chloride. What is the % yield? III. Reaction Types Symbols: Solid (___) or (___); Liquid (____); Gas (____); Aqueous (dissolved in water) (____) A. Synthesis (__________________________) General form: 1. element + element → compound (Use ions to form compound) a. sodium + chlorine → b. calcium + oxygen → c. lithium + sulfur → d. aluminum + oxygen → 2. metal oxide + water → metal hydroxide a. Na2O(s) + H2O(l) → b. CaO(s) + H2O(l) → c. Al2O3(s) + H2O(l) →

71

B. Decomposition General form: 1. Compound → ________________________+_________________________ a. FeCl3(s) → b. CuBr(s) → c. MgO(s) → 2. Metal hydroxide → ___________________ + ___________________ a. Be(OH)2(s) → b. Mn(OH)2(s) →

c. CuOH(s) → 3. Metal chlorate → ___________________ + ___________________ a. KClO3(s) → b. Zn(ClO3)2(s) → C. Double Displacement (___________________) General form: **Use solubility table for determining Ionic and Net Ionic equations.** Solubility Rules: (MEMORIZE!!) 1. ALL alkali metal compounds are soluble. 2. ALL ammonium compounds are soluble. 3. ALL compounds containing nitrate are soluble.

4. Most hydroxide compounds are INSOLUBLE. Alkali metal hydroxide are exceptions. 5. Most chlorides, bromides, & iodides are soluble. The exceptions are:Ag+, Hg2

2+, Pb2+

72

1. HCl(aq) + NaOH(aq) → Ionic: Net Ionic: ________________________________________________________________________________ 2. AgNO3(aq) + BaCl2(aq) → Ionic: Net Ionic: ________________________________________________________________________________ 3. FeCl3(aq) + KOH(aq) → Ionic: Net Ionic: ________________________________________________________________________________ 4. Pb(ClO3)2(aq) + NH4Cl(aq) → Ionic: Net Ionic: ________________________________________________________________________________ 5. Ca(C2H3O2)2(aq) + (NH4)2CO3(aq) → Ionic: Net Ionic: ________________________________________________________________________________

73

D. Single Displacement (____________________)

(Fig. 4.15) Activity Series of Metals Li K Ba Ca Na React with cold water to produce H2 Mg Al Zn Cr Fe Cd React with steam to produce H2 Co Ni Sn Pb React with acids to produce H2 H Cu Ag Hg Pt Au Do NOT react with water or acids 1. Cation displacement - ****MUST use the Activity Series of Metals**** General form: (Single element must be ____________ reactive than the element it replaces). a. Sn(s) + NaNO3(aq) → b. Zn(s) + H2SO4(aq) → c. K(s) + CaCl2(aq) → d. Na(s) + H2O(l) → e. Cu(s) + HCl(aq) →

74

2. Anion displacement - ****MUST use the periodic table**** (Single element must be __________ reactive than the element it replaces. General form: Halogens (Column _______) Reactivity: _______>_______>_______>_______ a. Cl2(g) + NaBr(aq) → b. I2(s) + KBr(aq) → c. F2(g) + CaCl2(aq) → E. Combustion (_________________) General form: 1. C2H6(g) + O2(g) → 2. C3H8O(l) + O2(g) → IV. Concentration of a solution A. Concentration units 1. Percent by mass = _____________________ x 100 2. Percent by volume = _____________________ x 100 3. Molarity = ____________________ B. Example calculations (1) If 3.50 g of sugar is dissolved per 85.5 g of solution, what is the % sugar by mass?

75

(2) If there is 12.5 mL of alcohol dissolved in 128.4 mL of solution, what is the % alcohol by volume? (3) If 4.44 g of sodium chloride is dissolved in 122 mL of solution, what is the molarity? (4) Compare the ion concentrations of 1.20 M glucose, 1.20 M KCl, and 1.20 M Ba(NO3)2. V. Preparation of a solution of known molarity A. Steps: 1. Calculate the mass required. 2. Weight this amount and put it into a __________________ flask. 3. Add _______________ water and swirl to dissolve. 4. After all the ________________ has dissolved, add water to the line. B. Example: Prepare 500 mL of 0.200 M potassium dichromate VII. Molarity calculations Use dimensional analysis; start with the number other than the molarity (mass or vol.) (1) What is the molarity of a solution that contains 3.45 g of sodium hydroxide dissolved in 455 mL of solution?

76

(2) How many grams of ferric nitrate are there in 125 mL of 0.150 M solution? (3) What volume of 0.200 M silver nitrate solution is required to obtain 1.45 g of silver nitrate? (4) Explain how you would make 750 mL of 0.125 M barium chloride solution. (5) What is the concentration of each ION in 0.550 M solutions of: (a) potassium dichromate (b) ammonium phosphate (c) hydrochloric acid

77

VIII. Dilution of solutions: Concentrated solutions are often stored in chemical stockrooms and are called ________________ solutions. Preparation of less concentrated solutions from a concentrated stock solution is done by the process of __________________. This is accomplished by increasing the amount of ______________(________________). Equation: Examples: (1) Explain how you would prepare 1.50 L of 0.0500 M potassium permanganate from a stock 1.50 M potassium permanganate solution. (2) Explain how you would prepare 250 mL of 0.250 M sulfuric acid solution from a concentrated 10.5 M solution. (3) If 125 mL of distilled water is added to 115 mL of 0.100 M hydrochloric acid, what will be the final concentration of the solution? (4) What is the [Cl-] if 125 mL of 0.100 M sodium chloride is mixed with 225 mL of calcium chloride.

78

IX. Quantitative Aspects of Reactions in Aqueous Solutions A. Quantitative Analysis: the determination of the _________ or ___________________ of a substance in a sample. B. Qualitative Analysis: The determination of the types of ___________ present in a solution. C. Gravimetric Analysis: An experimental procedure based on the measurement of ___________. D. Precipitation Analysis: Finding the composition of a component of a compound by depositing it from a solution as an _________________ compound called a __________. E. Quantitative Analysis Example: Determination of % by mass of Cl- in a sample.

1. Weigh the sample; dissolve it in _____________ water. (Why distilled water?)

2. Add enough silver nitrate of known concentration to ___________________ all

of the Cl- as _____________.

3. The silver chloride precipitate is separated from the mixture by _____________

and then it is ______________ and _______________.

4. From the amount of AgCl, determine the amount of _______ present.

5. The % by mass of Cl- in the original sample is calculated by:

F. How do you know which solution to use? Solubility rules!!! Solubility Rules: 1. ALL alkali metal compounds are soluble. 2. ALL NH4

+ compounds are soluble. 3. ALL compounds containing NO3

-, ClO3-, and ClO4

- are souble. 4. Most OH- compounds are INSOLUBLE. Alkali metal hydroxides and Ba(OH)2 are exceptions. 5. Most Cl-, Br-, and I- compounds are soluble. The exceptions are: Ag+, Hg2

2+, Pb2+. 6. All CO3

2-, PO43-, and S2- compounds are INSOLUBLE. Exceptions are alkali metal

and ammonium compounds. 7. Most SO4

2- compounds are soluble. Exceptions are Ba2+, Hg2+, and Pb2+. 1. Example: How could you precipitate PbI2?

79

2. Example: How could you precipitate Mn(OH)2? 3. Example: How could you separate a mixture containing Cu2+, Ba2+, and Ag+? Quant example problems: (1) In lab a student forms 5.34 g of silver chloride precipitate by reacting 1.775 g of an unknown chloride sample with an excess of silver nitrate solution. (a) Write the net ionic equation. (b) Calculate the % by mass of Cl-. (c) If the formula is XCl2, what is element X’s molar mass and identity? (2) A student mixes 125 mL of 0.124 M aluminum nitrate and 175 mL of 0.104 M barium hydroxide. (a) Write the molecular, ionic, and net ionic equations. Identify the precipitate. (b) Determine the mass of precipitate formed.

80

(c) Calculate the concentration of excess cation (M). (3) A student wants to determine the molarity and % by mass of an NaCl solution. He slowly adds 0.100 M AgNO3 until he observes that no more precipitate forms. He has added 12.45 mL of the silver nitrate to 100.0 mL of the NaCl solution. (sp. gr. = 1.00) (a) What is the molarity of the NaCl solution? (b) What is the % by mass of the NaCl solution (4) A 0.08750 g sample of an unknown compound containing phosphate ions is dissolved in water and treated with excess calcium nitrate and 0.6764 g of precipitate forms. (a) What is the precipitate? Write the net ionic equation. (b) Calculate the % by mass of phosphate in the original sample.

81

(5) A student wishes to determine the amount of NaCl and KCl in a 1.557 g sample that is a mixture of both. She dissolves the sample in distilled water and then reacts it with excess silver nitrate. He obtains 3.408 g of dry precipitate. Picture: (a) What is the precipitate? Write the net ionic equation. (b) What is the mass of chloride in the mixture? (c) What is the % of KCl and NaCl originally present in the sample?

82

Unit 6: Atomic Theory II /Periodic Trends

Chapter 7: Atomic Theory II I. History of the Atom A. ________________: Greek philosopher from 400 BC, “atomos” B. ________________: Atomic Theory, based on chemical behavior, around 1800 C. ________________: Plum pudding model of the atom (First atomic model), after he discovered the _______ with the _______________________. D. ________________: Planetary model of the atom, after he discovered the ______ with the _________________________ and ______ particles. Problems: The planetary model does not explain… 1. why electrons do not ________ energy and _______ into the nucleus. 2. the atomic line __________. 3. ____________ trends of the elements. E. ________________: The Bohr Model of the atom F. ______________________________: Our current model of the atom. II. Radiant Energy A. What is light? (1600s) 1. Particle Theory: ___________________ 2. Wave Theory: _____________________ B. What is light? (today) 1. ________________________ duality III. Electromagnetic Radiation (Light behaves as a wave) A. Electromagnetic Theory 1. James Clerk ________________ (~1870) proposed that accelerating _____________ produce a changing ____________ field, which produces a changing ________________ field, creating electromagnetic radiation (____________). 2. Maxwell’s proposal was experimentally verified by _____________________ using a ______________________. Diagram: 3. Hertz created the first ________________ transmitter and receiver.

83

B. Electromagnetic Spectrum: The range of ________________ or ________________ over which electromagnetic waves occur.

Blue Red C. Characteristics of Electromagnetic Waves 1. Speed: All EM waves travel at the speed of ___________ (in a vacuum) c = ____________________ m/s or ___________________ mi/s Example: How many kilometers is the sun from earth if it takes 8.1 min for light to reach the earth from the sun? 2. Frequency: The number of wave _____________ per second. The symbols commonly used are ______ or ______. The units are _______ = _______ = ______. 3. Wavelength: The length of one complete wave cycle, often measured from crest to crest, or trough to trough. The symbol used is ______. It is generally measured in __________. Diagram:

4. The Wave Equation: Velocity = frequency x wavelength

84

D. Sample Calculations: 1. If a red light has a wavelength of 725 nm, what is its frequency? 2. What is the wavelength (in meters and Angstroms) of an FM radio station that has a frequency of 101.5 MHz? IV. Quantum Theory (Light behaves as a particle) A. In 1900, ___________________ proposed that energy comes in ___________ amounts called _______________. 1. A quantum is a bundle, packet, or _________________ of energy. 2. A quantum of light energy is called a ______________. B. The energy of a photon (“particle” of light): E = h = f = C. Graph of energy versus frequency: E f D. Another unit of energy is the __________________(_____). The conversion factor is: E. Using the wave equation (____________), the energy of a photon can be calculated in terms of wavelength:

85

F. Electromagnetic Spectrum _________λ _________λ _________f _________f _________E _________E ________________________________________________________________________________ G. Sample Problems: 1. Calculate the energy, speed, and frequency of monochromatic blue light that has a wavelength of 490 nm. 2. If a photon has an energy of 7.55 x 10-19 J, what is its wavelength in nm? H. ____________________ further showed that light behaved as a ______________ with his theory of the _______________________. He suggested that a beam of light is really a stream of ___________________. He concluded that a photon of light contained a quantum of energy, E = hf. This work won him the Nobel Prize in Physics in 1905. V. Bohr Model of the Hydrogen Atom A. The ____________________ model (by __________________) did not explain: (1) the ________________ collapse problem, (2) periodic __________________, and (3) atomic _______ spectra. B. Physicist ______________________ applied the newly developed quantum idea (light travels as a packet of energy that we call a ______________) to the hydrogen atom. 1. The electrons are found in _______________ levels outside the nucleus called _______________ or orbits. 2. Electrons could only be found in these energy levels, and they could not lose energy and collapse into the nucleus. Why? _______________________________. 3. Energy levels were designated by the n, which referred to the ______________quantum number.

86

C. Diagram:

D. Atomic Line Spectrum: The Bohr Model was able to explain the _________________ line spectrum. 1. When ______________ is added to an electron, it will jump to _______________ energy levels, known as the ____________________ energy state. 2. When the excited electron falls down to a ____________ energy level, it releases ______________ in the form of a ______________ of light. The frequency or wavelength of that photon could be determined by the equations: 3. The larger the gap in energy levels, the larger the _____________ of the photon, therefore the _______________ the wavelength, and the _______________ the frequency. 4. The spectral lines are due to light emitted by the falling electron. Since the spacing between energy levels is constant, the light emitted for a particular electron jump will always be the same _________________ and _______________.

87

Spectral lines: Lyman Balmer Paschen Brackett Series Series Series Series Light: ______ ______ ______ ______ Jumps from excited states to: ______ ______ ______ ______ **Note: n = 1 is known as the ___________________ state. All others are __________________ states.

Line spectrum (_______________________)

E. Bohr Calculations 1. Each spectral line corresponds to difference in ________________(_____) between two permissible ___________________. a. Calculation of energy, f, wavelength of photon emitted (or absorbed):

88

b. Calculation of the radius of orbits (Rn): c. Calculation of the energy of each energy level (shell), En: F. Sample Problems: (1) Calculate the radius and energy of the 3rd shell of the hydrogen atom. (2) What are the frequency and wavelength of the photon emitted when an excited electron jumps from the 3rd shell to the ground state in a hydrogen atom? (3) (a) What is the frequency and wavelength of the photon absorbed when an electron jumps from the 2nd to the 6th shell? Given: E6 = -0.378 eV, E2 = -3.40 eV.

89

(b) How would the frequency and wavelength of the emitted photon compare to the absorbed photon if the electron jumped back from the 6th to the 2nd? (4) What frequency photon is needed to cause an electron in the hydrogen atom to jump from the ground state to the 6th shell? (5) An excited hydrogen atom emits a photon of 95.0 nm when an electron jumps from a higher energy level to the ground state. From what shell did the excited electron jump? VI. Quantum Mechanical Model of the Atom A. Problems with the Bohr Model 1. It could only explain the atomic line spectrum for _______________________ elements: 2. There was no proof of _________________ orbits. 3. It did not explain the electron _______________ problem of circular orbits. 4. Could not explain chemical __________________. B. Changes from the Bohr Model 1. Treats the electron as a ______________ and as a ________________. 2. Develops a model for the _____________________ of finding an electron in a space outside the nucleus, called an ___________________ or _________________________. 3. We now think of an electron as: ____________, _____________, ____________

90

C. Scientists who contributed to the Quantum Mechanical Model 1. Louis __________________(1929) a. Wave-particle duality of _________________. b. If a wave can behave like a particle, then a ________________ can behave like a wave. c. Electrons show diffraction patterns very similar to _____________, which verifies the wave properties of matter. See Figure 7.14 (p. 277). d. Derivation of de Broglie’s wave equation: e. The wave equation could be applied to all systems, but is only _____________________ for very small objects. For large objects, like a baseball, the wavelength would be so small, that it would be undetectable. f. Example: Calculate the wavelength of a 2.00 kg ball moving 5.00 m/s. g. Example: Calculate the de Broglie wavelength of an electron moving 10.0% the speed of light. (Write the answer in meters and Angstroms). h. de Broglie’s idea explains why electrons can only be present in specific ____________________ (e.g. electrons can only be on the rungs of the ladder, but never between), and why electrons in an energy level do not lose ________________.

91

i. de Broglie thought that electrons behaved like _________________ waves, and that the wave must fit the _____________________ of the orbit exactly. Otherwise, the wave would _______________ itself out.

Standing waves (a) circumference = integral # of λ (b) circumference ≠ integral # of λ, so the wave would cancel itself out. j. Each orbit corresponds to a _______________ wavelength, therefore electrons can only have particular amounts of energy. Since a wave is continuous, electrons do not ____________ energy. 2. Werner Heisenberg: The Heisenberg __________________ _________________ a. It is not possible to know the exact ___________________ and ___________________ of an electron at the same time. b. As the uncertainty of the position decreases, the uncertainty of the momentum (mass x velocity) __________________. c. This mathematical equation shows that an electron cannot move in a well-defined ___________ as Bohr thought. If an electron moved in a circular orbit, we could know the position and momentum at the same time. d. We do not know the exact path of an electron, but we can call the region of space where the electron has a high _________________ of being found the ____________________ or electron cloud. Diagram:

92

3. Erwin Schrödinger: The Schrödinger ______________ Equation a. The equation (which involves advanced calculus) describes the behavior and _______________ of electrons. b. It incorporates both the __________________ and __________________ nature of matter. c. The wave equation can only be solved for _______ electron systems. All others are only approximated. d. The wave function = _______ (psi). e. ψ 2 = _________________ of finding an electron in a particle region of space. f. The solutions to the wave equation yields the ______________numbers. D. Quantum Numbers 1. The ________ quantum numbers are used to describe the most _____________ location of the electron in the atom. 2. Analogy: Where is the most probable place to find you at 3 a.m.? ______________ _______ ___________ ________ Principal Energy Level Sublevel Orbital Spin (n) → (l) → (m) → (s) 3. The four quantum numbers: a. Principal Quantum Number (n) i. Gives the principal ________________ of the electron ii. Possible values: n = 1, 2, 3, …_____ iii. Represents the average ______________ from the nucleus. The larger the number the _______________ the electron is from the nucleus. b. Angular Momentum Quantum Number (l) i. Determines the ______________ of the electron. ii. Gives the “shape” of the ___________ iii. The value of l depends on the value of n: l = 0, 1, 2, …______

93

iv. Examples of shapes:

l = 0 (s orbital)

l = 1 (p orbital)

l = 2 (d orbital)

l = 3 (f orbital)

8 – lobed orbital of complex shapes (not shown in our textbook) l = 4 (g orbital), l = 5 (h orbital), l = 6 (i orbital) **** There are currently no atoms large enough for electrons to be present in any of these sublevels. v. Chart of sublevel types for each energy level

94

Value of n Possible Values of l Types of Orbitals

1

2

3

4

5

c. Magnetic Quantum number (ml) i. Gives the _____________________ of the angular momentum. ii. It describes the magnetic ____________ generated by an electron as it moves around the nucleus. iii. The values of m depend on l: __________ to __________. iv. It gives the ___________________ of the orbitals in space. v. Chart of Orbital Orientations for Each Sublevel

Value of l Type of Sublevel Possible Values of ml # Orbitals

0

1

2

3

d. Spin Quantum Number (s) i. Every electron behaves in certain respects as though it is a spinning charged sphere. There are _______ possible directions of spin. ii. Possible values of s: ________ and ________. iii. Electrons within an orbital must have _________________ spins. iv. Diagram showing electron spin

95

****Electrons are represented with arrows to indicate spin direction.

v. Stern and Gerlach experimentally verified spin with the following experiment:

vi. A beam of hydrogen atoms is passed through a ______________ field. Half of the atoms are deflected ____________, the other half are deflected downward. This means that on average, half of the atoms have electrons spinning in one direction, the other half are spinning in the opposite direction. 4. Pauli Exclusion Principle: No two electrons can exist in exactly the ___________ quantum state. Every electron in atom will have a _______________ set of quantum numbers. E. Sample Questions: 1. What are the four subshells? Which quantum number represents subshell? 2. How many (a) orbitals, and (b) electrons maximum are there in each subshell? (i) s (ii) p (iii) d (iv) f 3. Complete the table:

96

Value of n Values of l Subshells # Orbitals Max. # electrons

1 0 1s 1 2

2

3

4

4. What is the maximum number of electrons that could occupy: (a) a 1s orbital? (b) a 2p subshell? (c) the 3rd energy level? (d) a 2p orbital? 5. How many orbitals are there in: (a) a 2p subshell (b) a 3f subshell? (c) the 4th energy level? (d) the 1st energy level? 6. Give the four quantum numbers for all of the electrons in a full: (a) 6s orbital (b) 3p subshell (c) 4f orbital 7. In which subshell is each of the following electrons found? (a) n=5, l=1, m=0, s=+ ½ (b) n=4, l=3, m=2, s=- ½ (c) n=3, l=4, m=5 VII. Electron Configuration: Shows how the electrons are arranged in an atom.

97

A. The electron configuration is written in the following format: C (______electrons): 1s22s22p2 number of electrons in that sublevel sublevel (l) Energy level (n) (The sum of all of the superscripts should equal the __________ # of electrons.) B. The orbital diagram uses arrows to designate the electrons: C C. How to write an electron configuration: 1. Start filling electrons in the subshells of _______________ energy first, then build up (aufbau principle). 2. Only ________ electrons per orbital. 3. Subshell # orbitals s p d f 4. Electrons within an orbital much have ______________ spins to follow the ___________ Exclusion Principle. 5. Hund’s Rule: If there are multiple orbitals of the same energy, each orbital is first filled with only ______ electron, before they begin to _________ up. example: (4 electrons) ___ ___ ___ ___ ___ ___ D. Auf Bau Diagram

E. Examples:

98

1. Boron 2. Phosphorus 3. Potassium 4. Iron F. Using the Periodic Table to determine electron configurations:

***Label your periodic table with the energy subshells*** G. Kernal Configurations

99

1. Use the preceding _______________ as the kernel and place it in brackets. 2. Examples: Nickel Silicon Tungsten H. Irregularities with the Transition Metals: Half-filled and filled ______ subshells are more stable than other configurations, so some atoms have irregular electron structures. MEMORIZE THESE EXCEPTIONS!! 1. Examples: Chromium Copper Silver Gold I. Valence Electrons: The electrons in the outermost _____ and _____ subshells. 1. The value of _____ must be the same for the s and p. 2. The number is never greater than _______ (___________) 3. Elements can be represented with the dot formula:

Li Be B C N O F Ne

100

VIII. Review of Quantum Mechanical Model A. Example: Ne B. n2 = # __________ in an energy level. 2n2 = # __________ in an energy level. C. Electron configurations of transition metal cations (lose e- from VALENCE first!) Fe Fe2+ Fe3+ Sn Sn2+ Sn4+

Chapter 8: Periodic Trends I. History of the Periodic Table A. In 1869, ___________________________ organized the elements by increasing ______________________. The organization showed similar ________________ for elements in the columns. He left gaps where some elements had not yet been discovered and predicted their __________________ based on the properties of other elements in the column. For example, there was a gap below Aluminum. He called the undiscovered element “___________________” (first element below aluminum). He predicted its properties. When the element was discovered (and called __________________), its true properties were compared to the predicted properties of eka-aluminum: Eka-Aluminum Gallium Atomic mass 68 amu _______

Melting Point Low _______

Density 5.9 g/cm3 _______

Formula of oxide Ea2O3 _______

101

II. Periodic Classification of the Elements

III. Electron Configuration of Cations and Anions A. Examples: Na Na+ Al Al3+ F F- O O2- B. Isoelectronic: Atoms and ions that have the _______ number of electrons and the ________ ground-state electron configuration. IV. Atomic Radius A. Atoms do not have __________________ boundaries, so it is difficult to determine atomic size. B. Atomic radius is defined as _______________ the distance between two nuclei in adjacent atoms.

102

C. Periodic trend in atomic radius:

1. Within a group: Atoms increase in size going ___________ the column because elements in each successive period have an extra energy level. Examples: Na K 2. Within a period: Atoms increase in size going from _________ to _________ in a row!! a. WHY? As the number of ____________ in the nucleus increases within a period, the electrons are pulled inward with a greater __________. b. As Zeff _______, atomic radius _______. c. Zeff = Effective nuclear charge = Z – inner electrons d. For neutral atoms: Zeff = # of ____________________. e. Example: Calculate the Zeff of Li and O. Which atom would be larger? 3. General trend of atomic radius:_______________

V. Ionic Radius A. Cations (___________): Are _______________ than the corresponding ____________ atom. 1. Example: Calculate Zeff for Na and Na+ 2. As Zeff ______, ionic radius _______.

B. Anions (___________): Are _______________ than the corresponding neutral atom.

103

1. Example: Calculate Zeff for F and F-

2. Zeff remains _____________________ as electrons are added to an atom.

3. Why do atoms get larger when they gain electrons? More e- ______________.

C. Example: Circle the larger atom.

(1) Br, I (2) O, O2- (3) Ca2+, Ca (4) V, Fe (5) C, Si

VI. Chemical Reactivity

A. Nonmetal reactivity: _______________ Most reactive nonmetal:_______________ Same trend for Nonmetallic character. B. Metal reactivity: _______________ Most reactive metal:________________ Same trend for Metallic character. WHY is cesium more reactive than lithium?

104

VII. Ionization Energy (IE) A. The minimum energy needed to remove an _____________ from a gaseous state atom, generally measured in kJ/mol. X(g) + ionization energy _______(g) + _______(g) B. Gaseous atoms are used because they are uninfluenced by their neighbors. C. In a multi-electron atom, more than one electron can be removed from the atom. When the second electron is removed, we call it the ________________ IE, etc. X+

(g) + IE2 _______(g) + _______(g) Second IE X2+

(g) + IE3 _______(g) + _______(g) Third IE D. It takes ___________ energy to remove the second electron than the first, because the positively-charged ion has a larger attractive force towards the remaining electrons. E. Ionization Energy Data

F. Ionization Energy Periodic Trend 1. Within a group: IE _____________ going down a column. Why? As an atom gets larger, the electron is ______________ to remove because there is less pull to the nucleus (more e- shielding and larger distance from nucleus). 2. Within a period: IE generally increases going from __________ to ________ on the periodic table. Which group has the lowest IE?_____________Why? Which group has the highest IE?_____________Why?

105

3. Within a period there are _____ exceptions to the general trend. a. Group IIIA has a LOWER ionization energy than Group ______.

b. Group VIA has a LOWER ionization energy than Group ______. Examples: Be vs. B N vs. O *** Half-filled and Full subshells are relatively stable *** G. By studying the IE of an element, you can determine the number of ___________ e- of an atom. Outer (valence) electrons take significantly ________ energy to remove than an inner electron. Example: Be VIII. Electron Affinity (EA) A. The negative of the energy change that occurs when an electron is ___________ by an atom in the gaseous state to form an ____________.

X(g) + e- X-(g)

F(g) + e- F-

(g) Energy = -328 kJ/mol; EA = +328 kJ/mol

B. The more POSITIVE the EA, the more easily an atom accepts an electron. Negative electron affinities are for atoms that do not want to accept an electron.

C. Which group has the highest EA? ______________ Why? D. Why do the Alkali Metals have a relatively high EA?

106

Unit 7: Chemical Bonding

Chapter 9 I. Introduction A. What is a chemical bond? A _________ that holds atoms together. B. Why do atoms bond? To become more ___________. Maximum ____________ occurs when electrons are the same as a _____________. C. Forces of _________________> Forces of __________________ D. Types of bonds: 1. Ionic: _____________ of electrons. 2. Covalent: ________________ of electrons. 3. Metallic: _______________ electrons. II. Valence electrons A. The ________________ electrons (usually the ____ and ____ of the highest shell). B. These are the electrons involved in _____________. C. Transition elements sometimes use ____ subshell electrons for bonding. D. Octet rule: Atoms tend to form bonds until the yare surrounded by _____ valence e-. Exceptions: He Li+ Be2+ III. Ionic Bonding A. The _____________________ of electrons from a ____________ to a ____________. Cations and __________ are formed which _______________ one another with an ________________ force.

B. Lewis dot formula for ions: Sodium Sodium ion Fluorine Fluoride Magnesium Magnesium ion Oxygen Oxide B. Examples: Sodium chloride Magnesium chloride

107

C. Ionic Structure 1. Crystal ___________: 3-D arrangement of ____ and ____ ions in an ionic solid.

2. Lattice energy: The energy _____________ when 1 mol of an ionic solid is ________________ into its ions. a. Always ____________________ (energy absorbed). b. Lattice energies are generally large values, which shows that: i. ionic bonds are _____________. ii. ionic solids have relatively _______ melting points. iii. ionic compounds are generally ____________ at room temperature. c. Amount of energy depends on: i. the ___________ between ions. ii. the ___________ of the ions. **(This is has a larger effect)

d. Coulomb’s Law: + r - e. Which compound has the highest lattice energy? (i) CaO or CaS? (ii) LiH or CaH2?

f. Why is the order of lattice energies: LiF > LiCl > LiBr > LiI? g. Which has a higher lattice energy: MgCl2 or NaCl? Why?

108

D. Ionic compounds are generally __________________ in water, and the solution ___________ electricity because the compounds are __________________. E. Molten (liquid) ionic compounds are also ___________ conductors of electricity, because in the liquid state, charged particles (________) can flow. IV. Covalent Bonding A. The ____________ of electrons between __________________.

B. Structure:

1. individual _______________ (e.g. H2O) 2. 3-D lattices in network covalent bonds (e.g. _________________)

C. Covalent compounds have attractive forces between atoms (______________) and attractive forces between molecules (________________________). D. Intermolecular forces between molecules are generally ___________, so covalent compounds are generally in the _____________ or __________ phases, or possibly a low- melting __________. E. Covalent compounds are generally ___________ conductors of electricity because no ions are present. F. Types of covalent bonds:

(shared pairs of electrons are shown as _____ or _____.) 1. Single bond: _____ pair of shared electrons. Example: H2 2. Double bond: _____ pairs of shared electrons. Example: O2

3. Triple bond: ______ pairs of shared electrons. Example: N2 G. Bond length and energy:

Bond Type N-N N=N N≡N

Bond Length (Å) 1.47 1.24 1.10

Bond Energy (kJ/mol) 163 418 941 1. As the number of bonds _____, the bond length ______ (______________ relationship) 2. As the number of bonds _____, the bond energy ______ (_____________ relationship) 3. Why can’t N2 be directly used by plants? What forms of nitrogen can be used as a plant nutrient?

109

V. Electronegativity A. The ability of an atom to ______________ a shared pair of electrons toward itself. B. Elements with high _______ (accept electrons easily) and high _______ (do not lose electrons easily) also have _________ electronegativity. C. The general trend of electronegativity is the same as IE: _______________ D. Electronegativity is a relative concept and can only be measured relative to other elements. Therefore it has no ___________, and the highest value was set at _______. E. Chart of electronegativities:

F. Element with highest electronegativity: _______ Lowest electronegativity: _______ VI. Bond Polarity A. Non-Polar Covalent (_____________ Covalent) Bond 1. Electrons are shared ________________. 2. The bond has 100% covalent character and 0% ___________ character. 3. The electronegativity difference between atoms in the bond is __________. 4. Example: H2 B. Polar Covalent Bond 1. Electrons are shared _________________. 2. The electronegativity difference between atoms in the bond is < _______. 3. Example: HCl

110

C. Ionic Bond 1. Electrons are transferred from one atom to the other. 2. The electronegativity difference between atoms in the bond is ≥ _______. 3. Example: LiF

D. Examples: What type of bonds are each of the following: (a) KF (b) HBr (c) Br2 (d) H2S (e) CsCl VII. Ionic or Covalent Character A. Only bonds between identical elements are purely covalent. All others have some % ionic character and some % _______________ character. B. Example from period 2 of the types of bonding between atoms:

VIII. Oxidation Numbers (See page 129 in text): The charge an atom would have if electrons were completely transferred to the __________ electronegative atom in the bond. The “relative” charge of an atom. A. Rules to assign oxidation numbers: 1. In free elements, each atom has an oxidation state of _______. (i.e. H2, Br2, Li, etc.) 2. For monatomic ions, the oxidation state is equal to the __________ on the ion. (i.e. Li+ has an oxidation state of +1) 3. The oxidation state of oxygen in most compounds is ________, but in peroxides its oxidation state is _________. 4. The oxidation state of hydrogen when it is bonded to a non-metal is ______. When it is bonded to a metal its oxidation state is ______. (i.e. HCl = +1; LiH = -1)

111

5. Fluorine has an oxidation state of ______ in ALL its compounds. Other halogens can have ____________ oxidation states when bonded to oxygen. 6. In a neutral compound, the sum of the oxidation states is ______. In a polyatomic ion, the sum of the oxidation states is the ___________ of the ion. B. Examples: Determine the oxidation states of each element in the following: LiH BeH2 H2O HF OF2 H2O2 CaCl2 Cl2O5 Cl2O7 ClO2 PbO2 SiO2 NO2 N2O NO3

- NO2- MnO4

- C2O42-

IX. Periodic Table Relationships:

Column # # Valence Lewis

dot formula Most likely #

of bonds # lone pairs

IA 1

IIA 2

IIIA 3

IVA 4

VA 5

VIA 6

VIIA 7

VIIIA 8

X. Lewis Structures of Covalent Compounds: A representation of the covalent bonding in a molecule. A. Covalent bonds are shown as __________. Example: H2 B. Lone pairs of electrons are shown as _________. Example: O2 C. ONLY ____________ electrons are shown.

112

D. General steps for drawing Lewis structures: 1. Sum the valence electrons in the compound. 2. Add _____ for each negative charge. Subtract 1 for each _____ charge. 3. Generally place the element that makes the _________________ number of bonds in the center. 4. Draw ____________ bonds to the other atoms off of the central atom. 5. Place electrons around the ___________ atoms until an __________ is reached. 6. If you run out of electrons, start forming _____________ or ___________ bonds. 7. If you have EXTRA electrons after all have octets, place them on the __________ atom. 8. In the end, all atoms should have an octet that need an octet (____ is an exception), and the total number of electrons should be placed on the molecule. EXAMPLES: 1. CH4 2. O3 3. CO2 4. CO3

2- PRACTICE!!! 1. HOCl 2. NF3 3. TeCl2 4. ClO3

- 5. C2H2 6. XeF2 7. N2O4 8. HOCN 9. PO3

3- 10. HNO2 11. CH2O 12. CH3COOH

113

XI. Formal Charge A. (Determined for each atom) 1. Formal charge = ___________________________________. 2. Use: To decide between _____________ Lewis structures. (Smaller formal charge =____________) 3. For neutral compounds,____________________________________________________. 4. For cations,_________________________________. 5. For anions, _________________________________.

6. Examples: CO2 Formal Charge

Structure 1:

O C

O Net Charge = Structure 2:

O C

O Net Charge =

Ammonium ion N all H Net Charge = Cyanate ion Which structure is more plausible? Why? XII. Resonance Forms A. Two or more equally ____________ structures of the same compound. B. Example: Ozone O3 C. Neither form alone accurately describes the _______________ of ozone. D. The bond lengths between each oxygen atom are the _________, which means that neither bond is a ___________ bond or a ____________ bond, but a hybrid of the two.

114

D. Example: benzene (C6H6) XIII. Exceptions to the Octet Rule A. Odd number of ____________________ 1. Example: Nitrogen monoxide _________ Valence electrons:____________ 2. Free radical = ________________________! (SMOG) B. Incomplete octet (_________on central atom) : Elements with 3 or less valence electrons. 1. BF3 2. BeCl2 3. HgCl2

C. Expanded octet (_________ on central atom): Occurs when (1)_________central atom and (2) surrounding atoms are very ___________________ (ex. ____, ____, ____)

1. P

2. S XIV. Coordinate Covalent Bonds: A covalent bond in which ____ atom provides both of the _________ electrons. A. Ammonia (________) Ammonium ion (________) Formal charge of N? B. Water (_________) Hydronium ion (_________) Formal charge of O?

115

Chapter 10

I. Molecular Geometry A. Ideal Geometries: The ____________ atom has NO lone pairs of electrons.

Type Geometry Example Compound

Example Structure

Bond Angle

AX2

AX3

AX4

AX5

AX6

II. VSEPR Theory: Valence Shell Electron ___________ _________________ Theory A. Double or ___________ bonds can be treated like single bonds. B. If a molecule has _________________ structures, VSEPR can be applied to any of them. C. Lone pairs of electrons repel atoms ____________ than bonding pairs. Therefore, lone pairs take up more __________ than atoms.

Type Geometry Example Compound

Example Structure

Bond Angle

AX3E

AX2E2

AXE3

116

MEMORIZE THE SHAPES FOR THE TEST!

III. Polar Molecules: Molecules with a ________________ moment (dipole arrows don’t cancel out!) A. Determination of molecular polarity 1. Draw Lewis structure. 2. Determine ______________ of molecule. 3. Determine polarity of ____________. 4. Draw dipole arrows. 5. Determine net dipole (_________________________________________________).

117

B. Examples: 1. CO2 Type = _____________ Shape = _____________________ 2. H2O Type = _____________ Shape = _____________________ 3. BF3 Type = _____________ Shape = _____________________ 4. CH2F2 Type = _____________ Shape = _____________________ 5. XeF2 Type=______________ Shape = ____________________ 6. NH3 Type = _____________ Shape = _____________________

118

IV. Hybridization of Atomic Orbitals: The _________________ of atomic orbitals in an atom (generally the_______________ atom) to generate a set of _______________ orbitals. A. Hybrid orbitals: a combination of _______ or more atomic orbitals. Example: Carbon generally forms ____ bonds, but only has ____ lone electrons. Why?? Electron configuration of carbon: Atomic orbitals of valence electrons: Hybridization of carbon: When carbon has four single bonds, it is _________ hybridized. B. Procedure for Hybridizing Atomic Orbitals: 1. Hybridization of orbitals only applies to _______________. 2. Hybridization is the mixing of at least _______ or more nonequivalent atomic orbitals. Hybrid orbitals have different _________ than the atomic orbitals: Atomic Orbitals Hybrid Orbitals: s p sp sp 3. The number of hybrid orbitals is equal to the number of pure _________ orbitals. 4. Hybridization requires an input of _____________. However, this energy is more than recovered during __________ formation (an ________________ process). 5. Covalent bonds are formed by an __________________ of _____________ orbitals with atomic orbitals: C. Example hybridizations:

Hybrid # Single Bonds Example Atomic Orbitals Hybrid Orbitals Shape

sp

sp2

sp3

sp3d

sp3d2

119

D. Example: Predict the hybridization of the central atom of each of the following: 1. HgCl2 2. AlI3 3. PF3 E. Hybridiztion of s, p, and d orbitals: For elements in the _____________ period and beyond, the s, p, and d orbitals all contribute to hybridization. Expanded ____________ can be explained by the hybridization of the d orbitals. Because there is no ________ level, elements of the ____________ period cannot form expanded octets. 1. Example: Phosphorus in PBr5 2. Example: Sulfur in SF6 V. Diagram of Hybrid Orbitals

120

VI. Hybridization in Molecules Containing Double and Triple Bonds: A. Types of covalent bonds: 1. Sigma bond (_______): The bonds formed when the ______________ orbitals overlap with the atomic orbitals. 2. Pi bond (_______): The bonds formed when ______ orbitals overlap above and below the plane. Example: C2H4 B. Types of Hybridization for Carbon:

Bonding Hybrid Atomic Orbitals Hybrid Orbitals Sigma/Pi bonds

C. Example: Determine the hybridization of each carbon and the shape of each molecule: (1) CH2O, (2) C2H2.