Final Exam Review Day 1 (revised 2017) Name __________________Period__________________

Classify the following matter as an element, a compound or a mixture.

1. ____HCl 3. ____ Cl2 (g) 5. ____ barium

2 ____ HCl (aq) 4. ____ diet coke 6. ____ calcium nitrate

Classify the following matter as a pure substance or a mixture

1. ____ Na2CO3 (s) 3. ____ air 5. ____ dirt

2. ____ Na2CO3 (aq) 4. ____ helium 6. ____ H2O (l)

Classify the following matter as homogeneous or heterogeneous.

1. ____ tossed salad 3. ____ stainless steel 5. ____ KOH (aq)

2. ____ Cl2 4. ____ saline solution 6. ____ MgSO4 (s)

Classify the following as a physical change or a chemical change.

1. __________ dissolving sugar into water

2. __________ burning methane gas

3. __________ decomposing water into hydrogen and oxygen

4. __________ evaporating water

Units of measurement (list what these quantities measure)

cm3, L or mL are for measuring __________________

g or kg are for measuring_______________________

g/cm3, g/l or g/mL are for measuring ________________

joules , calories, kcal are for measuring ______________________

K or ºC are for measuring _____________________________

Temperature conversions

ºC + 273 = K 20º C = __________K

Metric Conversions

________ mm = 1 m ___________ cm = 1m __________ m = 1 km

Significant Figures

1. Reading Instruments Properly – give all figures known and estimate the last one

2. Determine the number of significant figures in the following measurements.

a) 87 000 000 000 __________

b) 0.000 607 0 __________

c) 320.00 __________

3. Calculate each of the following. Report answers using significant figures.

a) (4.15 × 102) m × (6.024 × 103) m

b) __________3.56_g__________ 3.6 cm × 2.5 cm × 5.2215 cm

c) 18.63 g + 5.2 g

Density D = M /V

Calculate the density of an object measuring 21 cm by 6.0 cm by 2.12 cm and having a mass of 522.2 grams

A stone has a mass of 55 g. When I place it into a graduated cylinder that has 50.0 mL of water in it the level changes to 70.0 mL. What is the density of this piece of granite?

Conversions

Given: 1 joule = 0.239 calorie. Convert 630 joules to calories.

Subatomic Particles and their Jobs# protons = atomic number (identifies the element)neutrons define the mass of the atom # protons + neutrons = mass number# electrons determines the chargeComplete following table. List as much information with the symbol column as possible.

Symbol Atomic Number

Protons Neutrons Electrons Mass Number

Xe 77

Ba-137

27 33 27

Na+ 22

55Mn 25

17 18 36

15O-2

Write the complete chemical symbol for the ion with 31 protons, 39 neutrons, and 28 electrons.

Iron consists of four natural isotopes:Isotope

Mass (amu) Percent Abundance

54Fe 53.9696 5.8256Fe 55.9349 91.6657Fe 56.9354 2.1958Fe 57.9333 0.33

Calculate the atomic mass of iron to five significant digits.

Convert the following

1. 48.6 g Mg to atoms

2. 4.59 moles gold to g

3. 414.4 g Pb to atomsFinal Exam Review Day 2 Name_________________

Electron Configuration – draw diagonal lines in the chart below, then draw a small periodic table in the right box below labeling the s, p, d, and f blocks numbering the periods (1-7 down the left side)

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6s2 6p6 6d10 7s2 7p6

Draw the electron configuration for the following:Potassium

Sulfur

Cobalt

P−3

Mg+2

_________________ electrons are electrons in the outer shell.

Draw the dot diagrams for the following:

Potassium Sulfur P−3 Mg+2

Draw the shortcut electron configuration for each of the following:

Potassium

Sulfur

Cobalt

Identify the following with the correct element symbol:

[Xe]4f 46s2 ( or [Xe] 6s24f 4 ) __________This atom’s last electron fills the 3d sublevel __________ This atom’s last electron fills the p sublevel in period 2 __________This atom in period 2 and has 3 paired valence electrons in its dot picture____________This atom has two dots in its dot picture, but isn’t a metal ___________

Fill in the blank: each word/word group in the word bank may only be used 1x.

seven Alkali Metal atomic number bright line spectrumfamilies groups noble gases prinicipal energy levellose gain Alkaline Earth valence electronsright left staircase transition metalsHalogen anions cations stable octetPeriods s, p, d, f Mendeleev continuous spectrum

The ____________ _______________ ___________ is the same as the period number. There are ____________ main energy levels. The _______________________________is caused when electrons emit energy as they fall back to a lower energy level. There are four “blocks” on the periodic table represented by the letters _____________.

Periodic TableKnow the name and location of the following groups

Group 1 ___________________ Group 2 _____________________Group 3-13 ___________________ Group 17 _____________________Group 18 ___________________

The Noble gases are stable because they have a _______________________________ Elements are placed in order of ______________ _______________ and placed in groups according to their number of __________________ __________________.

Horizontal rows are called __________________________Vertical columns are called _________________ or _________________Metals are located on the ___________ side of the periodic tableNonmetals are located on the ____________ side of the periodic tableMetalloids are located __________________________________

Metals _____________ electrons. They form __________________.Nonmetals ______________________ electrons. They form __________.

Trends – on small, hand drawn periodic tables, illustrate the trends using arrows

Atomic radius (size) ionization energy (energy to remove an electron)

Electronegativity (ability to gain an electron) Reactivity for metals (metallic character)

Reactivity for nonmetals Ionic radius

Choose the correct answer by circling the chemical symbol on both sides – don’t circle either answer if there isn’t a correct answer – circle both if both answers are correct

F Se is more reactive. Li KK Na is the smaller atom. Na MgC O has the higher ionization energy. Si CBa B has only s electrons in its outer shell P SSi I has 4 valence electrons Pb SnN Cs is a metal Si GeMn Cl last electron is d5 Tc ReMg S goes to a charge of −2 Fe CuGe Br has higher electronegativity Br ClFe Sb is a metalloid As BBe F is a halogen Cl BrCa Mg is more reactive K CaNa I is more likely to combine with oxygen Mg ClRb Sr loses electrons more readily I FAu Pb is a transition metal K Ca

Chemical BondingFill in the blank – words/word phrases can be used multiple times

Metallic bonding metal nonmetal sea of electronsIonic bonding shared transfer polarCovalent bonding low high same nonmetalpolyatomic ion evenly unevenly nonpolardifference in electronegativity

1. _________________ bonding usually occurs between a ___________ and a ___________ or a __________________ and a ____________________________. There is a _______________ of electrons. There is a high _________________________________________ values.

2. ___________________ bonding usually occurs between a ___________ and a _______________. The electrons are __________________ between the atoms.a) _____________ __________________________ occurs between 2 of the same nonmetals. Electrons are _________________ shared.b) _____________ ___________________________ occurs between 2 different nonmetals. Electrons are _________________ shared.

3. _____________________ occurs between atoms in a single element which is made up of only type of ________________________. Many electrons are shared. This type of bonding is described by the ________________________________ theory.

Characteristics of ionic vs covalent. Fill in the chart below comparing properties of ionic/covalent chemicals – use the words yes, no, higher, lower

Ionic covalentconduct electricity as solidsconduct electricity in watervapor pressuresmellmelting and boiling points

What is an intermolecular force?

What are the three types of intermolecular forces? Order them from left to right (strongest to weakest)

Lewis Structures for covalent compounds/polyatomic ionsDraw Lewis structures for molecules or polyatomic ion. Predict the shape names. Predict the intermolecular force for the molecules.

H2 Molecular geometry:

IMF:

HBr Molecular geometry:

IMF:

HF Molecular geometry:

IMF:

H2O Molecular geometry:

IMF:

H2S Molecular geometry:

IMF:

NCl3 Molecular geometry:

IMF:

NH3 Molecular geometry:

IMF:

CH4 Molecular geometry:

IMF:

CH3Cl Molecular geometry:

IMF:

Ionic Bonding Pictures and polyatomic ions: Draw in the boxes belowJust polyatomic ionsOH-1 NH4

+1

PO4-3 CO3

-2

Ionic Bonding PicturesNaF CaF2

AlF3 Na2O

Ionic bonding pictures with polyatomic ionsNaOH CaCO3

AlPO4 (NH4)2O

Final Exam Review Day 3 Name_________________

Calculate the oxidation numbers (charge) of the atoms in each of the following compounds or polyatomic ions:

FeCl2 FeCl3

TiO2 H3AsO4

H2Cr2O7 SO3−2

Calculate the molar mass of the following:

CaCl2 (NH4)2S Li3N

Convert each of the following.

a. 67.2 L of CH4 to grams CH4

b. 69.4g of lithium nitride to moles

c. 1.204 × 1024 formula units of ammonium sulfide to moles

d. 0.500 moles of calcium chloride to grams

Find the % composition by mass of each element in Be(OH)2

A compound contains 68.04 g N and 155.52 g O. Find its empirical formula.

A compound contains 68.04 g N and 155.52 g O. Its molar mass is 92 g. Find its molecular formula.

Naming Compounds and writing formulas

Type of Compound Compound NameBinary Ionic MgCl2

CuCl2

Tertiary Ionic MgSO4

Covalent CO2

Binary acids HBr (aq)

Oxyacids (ternary acids) HNO3 (aq)

HNO2 (aq)

Write the formula for each of the following substances:

strontium chloride phosphorus trichloride

sodium sulfide dinitrogen pentasulfide

magnesium nitrate aluminum nitrite

iron (II) hydroxide iron (III) hydroxide

hydrosulfuric acid sulfuric acid

sulfurous acid ammonium phosphate

write the names and formulas of the seven diatomic elements below:

Name the following substances:

HCl (aq) CaO

F2 N2O4

H3PO4 (aq) NaClO3

CO CO3−2

Label as an acid, a base, a salt, or other:

MgSO4 __________ C6H12O6 _________ HCl ____________ NaOH ____________

Chemical Reactions: Directly below are listed the types of reactions, how to identify the types, and examples of reactions showing reactants:

Synthesis (element + element → compound) Mg (s) + O2 (g) →

Decomposition (compound → element + element) HgO (s) →

Single replacement (element + cpd→element + cpd) K + Ni(NO3)2 →

Double replacement (cpd + cpd → cpd + cpd) Pb(NO3)2 + NaBr →

Combustion (CxHy + O2 → CO2 + H2O) C3H8 + O2 →

Neutralization (acid + base → salt + water) LiOH + H2SO4 →

Classify and balance the following.

________ 1. BaCl2 + (NH4)2CO3 → BaCO3 + NH4Cl

________ 2. KClO3 → KCl + O2

________ 3. Na2O + P4O10 → Na3PO4

________ 4. C6H6 + O2 → CO2 + H2O

________ 5. FeCl2 (aq) + Na3PO4 (aq) → Fe3(PO4)2 (s) + NaCl (aq)

What do the symbols (s), (l), (g) and (aq) mean?

Which substances in question 5 are solutions?

What are the reactants in question 5?

What are the products in question 5?

What is the subscript of oxygen in question #4?

What is the coefficient of oxygen in question #4?

Stoichiometry:

1. Given the following: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)

a) How many moles of C3H8 are needed to produce 9 moles of CO2?

b) How many grams of CO2 are produced when 2.0 moles of C3H8 react with excess oxygen?

c) How many liters of oxygen are needed to react with 22.0 g C3H8?

d) How many grams of oxygen are needed to react with 44.8 liters of C3H8 at STP?

2. Given the following: 2 Na + Cl2 → 2 NaCl

(a) If 23.0 g Na and 71.0 g Cl2 react, what is the theoretical mass of NaCl produced? What is the limiting reactant? What is the excess

reactant?

b) An excess of sodium is reacted with 142 g Cl2. If 199 g of NaCl are actually produced, what is the percent yield?

3. Aluminum reacts with hydrochloric acid according to the following unbalanced equation: 2 Al + 6 HCl → 2 AlCl3 + 3 H2

(a) If 27.0 g Al are combined with 73.0 g HCl, which reactant is the limiting reactant?

(b) How many moles of AlCl3 will be theoretically formed?

(c) An excess of HCl is reacted with 54.0 g of Al. If 4.0 g H2 are actually produced, what is the percent yield?

Final Exam Review Day 4 Name_________________

Gas Laws

Fill in the blank using words/word phrases only 1x each. Some words are not used.ideal gases low real gases Small particlesstraight lines elastic light Highnonpolar heavy polar water vapor helium neon ammonia hydrogencarbon dioxide powerful absolute zero fluid

Kinetic Molecular Theory (or The Kinetic Theory of gases) assumes gases are made up of _________ ___________ moving in _____________ ___________, colliding into each other with ______________ collisions.

There is no such thing as an ideal gas, but some gases are more ideal than others. Gases that are more ideal are ______________ and _________________. Some examples of gases that would be more ideal (write their formulas) would be _________ , __________ , and ___________.Some examples of gases that would be less ideal (write their formulas) would be _______________, ___________, and _______________.

Gases obey most ideally when there are certain conditions in the experiment regarding temperature and pressure. Ideal conditions are when the temperature is ___________ and the pressure is ____________.

Know these gas laws: combined gas law, ideal gas law, Dalton’s law of partial pressures

P1V1 = P2V2 PV = nRT Ptot = P1 + P2 + P3 … T1 T2

STP means:standard pressure1 atm, 760 mm Hg or 760 torr, 101.3 kPa and…standard temperature: 0 oC = 273 K (you must use Kelvin in all gas law problems)

1. A sample of gas occupies 400.0 mL at a pressure of 1.0 atm. What will the volume be if the pressure is changed to 2.0 atm while the temperature remains constant?

2. Calculate the volume occupied by 64.0 grams of O2 at a pressure of 1.0 atm torr and a temperature of 0C.

3. A gas occupies 250 mL at STP. It expands to 750 mL as it is heated to 273ºC. What is the new pressure in mm Hg?

4. A gas sample is collected over water when at a pressure of 760.0 mm Hg. What is the pressure of the dry gas if the partial pressure of the water vapor is 11.5 mm Hg?

5. Calculate the number of grams of N2 in a 5.6 liter cylinder at 273 K and 1520 torr.

Gas Stoichiometry

1. Given: 2 H2(g) + O2(g) → 2 H2O(g)If this reaction occurs at 200C and 1500 torr, how many liters of O2 will react with 20.0 L of H2?

2. Given: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)How many liters of CO2 at 273C and 380. torr are produced from 88.0g C3H8?

Phases and Phase Diagrams: Use the word bank descriptors or combinations of those descriptors to fill in the boxes below: high medium low fast slowclose far yes novery not very strong weak

Solids Liquids Gases

Definite shape

Definite Volume

Fluidity

Particle movement (KE)

Particle arrangement

Attractive forces

Density

Diffusion rate

Compressibility

Draw a phase diagram for most substances in the left box. (label each axis -list units of your choice, labeling the solid/liquid/gas sections, and the triple point and critical point. Draw a phase diagram for water in the right box.Below the boxes, write everything you know about phase diagrams

________________________________________________________________________

________________________________________________________________________

________________________________________________________________________

_______________________________________________________________________

Heat Calculations:

q = m∙Cp∙∆T q = m∙∆Hfus q = m∙∆Hvap 1.00 calorie = 4.18 J

Values for water Cp = 1.00 cal/g·°C ∆Hvap = 540. cal/g; ∆Hfus = 80.0 cal/g

1. Suppose you have 100.0 grams of ice. Calculate:a) The quantity of heat energy (to the nearest calorie) required to convert the ice at 0°C to liquid water.

b) The quantity of heat energy (to the nearest calorie) required to convert the melted ice (liquid water) at 0°C to liquid water at 100°C.

c) The quantity of heat energy (to the nearest calorie) required to convert the (liquid water) at 100°C to steam at100°C.

2. In an experiment, 35.1 joules, is required to raise the temperature of 22.5 g of lead from 17°C to 29°C. Calculate the specific heat of lead in J/g·°C.

3. Convert

a) 41.8 J to calories and kilocalories

b) 1000.0 calories to Joules and kiloJoules

Final Exam Review Day 5 Name_________________

Concentration% by mass = g solute × 100 Molarity (M) = moles solutes g solution Liter solution

Dilution formulas: C1V1 = C2V2 (MoVo=MnVn)

1. Calculate the molarity of 117.0 grams of NaCl dissolved in 500. mL of solution.

2. If you have a beaker with 250 mL of a 4.0 M NaCl solution, how many grams of salt are in the beaker?

3. Calculate the % of a solution containing 35 g NaCl in 100 g of solution

4. Calculate the % concentration by mass of a solution in which 20 g Mg(OH)2 is dissolved in 80 g H2O.

5. How many mL of a 12.0 M HCl solution are needed to make 100 mL of a 1.0 M HCl solution?

Solubility

1. What is the solubility of KClO3 at 30 ºC?

2. How many grams of the salt could dissolve at that temperature in 40 g of water?

3. 50 g of KCl are added to 100 g of water at 40°C. Is the solution saturated or unsaturated? If saturated, how many more grams of KCl remain undissolved?

If unsaturated, how many more grams of KCl could be dissolved?

4. 15 g of NaCl are added to 100 g of water at 100°C. Is the solution saturated or unsaturated? If saturated, how many more grams of NaCl remain undissolved?If unsaturated, how many more grams of NaCl could be dissolved?

5. How does temperature affect the solubility of a solid in a liquid?

6. How does temperature affect the solubility of a gas in a liquid?Ions in solution

1. Write an equation with phases showing the dissociation of AlCl3 in water.

2. For the following write (a) the ionic equation; and (b) the net ionic equation.

formula equation: 2 AlCl3(aq) + 3 Na2CO3(aq) → Al2(CO3)3(s) + 6 NaCl (aq)

ionic equation: →

net ionic equation: →

3. Write the equations showing the two-step ionization of H2SO4 in water.

4. Label the chemicals below as acids, bases, salts, or other. Circle the chemicals below that are electrolytes.

CaBr2 CH4 Ca(OH)2 Na3PO4 HF

__________ _________ _________ _________ _________

5. Write a few sentences explaining how you arrived at the answers in question #4.

Acids and Bases

1. Calculate the Normality of: Normality (N) = M × n n = # of H or OH

a. 0.45 M H2SO4 b. 0.32 M Al(OH)3

2. Calculate the pH of a 0.0052 M solution of HCl. pH = - log [H3O+1]

3. In an aqueous solution if H3O+ = 1 10 −5, then OH− = ________,

pH= ______, and pOH =______.

pH + pOH = 14 and kw = [H3O+1] [OH-1] = 1 x 10-14

Know pH scale |-----------------|-----------------|

acid ← 7 → base

Properties of acids Properties of bases1. 1.

2. 2.

3. 3.

4. 4.

5. 5.

Titration: N1V1 = N2V2 or NaVa = NbVb

How many mL of 0.15 N KOH are needed to neutralize 50.0 mL 0.25 N H2S?

Final Exam Review Day 6 Name____________________

Equilibrium Keq= Products only include (g) and (aq) reactants

1. Given: N2(g) + 3 H2(g) ↔ NH3(g) + heat

Write the equation for the equilibrium constant, Keq.

If at equilibrium the concentrations are: [N2] = 1.2 × 10−2 M; [H2] = 2.0 × 10−2 M;[NH3] = 1.5 × 10−5 M ; calculate the value of the equilibrium constant.

Is the reaction endothermic or exothermic?

2. For the same equation given above, predict the effect of the following (get more products, get more reactants, or no effect on the amount of products or reactants)

Adding more B Removing some A

Heating decreasing the pressure

Adding a catalyst removing AB2

3. For the equation: Al(OH)3 (s) + H2O(l) ↔ Al+3 (aq) + 3 OH−1(aq)

Write the expression for the solubility product, Ksp.

4. Name 4 ways to increase reaction rate.

1) 2)

3) 4)

5. According to collision theory, what two requirements need to be met in order to have an effective collision between molecules?

Time

Energy

Time

Energy

Heat in reactions: Draw graphs and label energy axis with numbers that are multiples of 10 Joules. Make the left graph exothermic. Label PE of reactants and products, ΔH, activation energy, and activated complex.

1. What is the value for the activation energy for the forward reaction?

(a) (b)

2. What is the value of ∆H for the reaction?

(a) (b)

Oxidation and Reduction

Oxidation = _______ of electrons

Reduction = _______ of electrons

The oxidizing agent caused the oxidation. (It was reduced)The reducing agent caused the reduction. (It was oxidized)Assign oxidation numbers to the elements in each substance.

KNO3 + CO → CO2 + NO2 + K2O

___________ was oxidized.(which element)

___________ was reduced. (which element)

___________ was the reducing agent. (which chemical)

___________ was the oxidizing agent. (which chemical)