Exam Review

Answer all questions completely. Show work where necessary. Work neatly.

Labs

1. Which of the following represents an Erlenmeyer flask? C

2. The reaction times for three trials of an experiment are 90.3, 90.2, and 90.5 seconds.

Which average time is expressed using the correct number of significant figures in

seconds? A

a. 90.3 b. 90.33

c. 90

d. 90.333

3. A student performed an analysis of a sample for the calcium content and got the

following results: 14.8% 14.9% 15.0% 14.9% The actual percentage of calcium in the

sample is 15.7%. Characterize these data in terms of accuracy and precision. C a. These data are accurate and precise

b. These data are accurate but not precise.

c. These data are precise but not accurate. d. These data are neither precise nor accurate.

4. Which piece of lab equipment would most precisely measure 13 mL of water? B a. 25 mL test tube

b. 25 mL graduated cylinder c. 25 mL beaker

d. 25 mL Erlenmeyer flask

5. Which is the safest practice when heating the contents of a test tube over a flame? C a. Wearing long hair down

b. Having safety goggles within reach

c. Pointing the test tube away from people

d. Keeping the test tube securely stoppered

6. A student spills a diluted acid solution on his hand. He should — D a. wipe it off with a paper towel

b. let it air dry

c. apply a base solution to neutralize it

d. rinse it off with running water

Mr. Buchanan

Chemistry

Name ___________________________________

Block ___________ Date____________________

7. What is the name of the lab equipment shown below? B

a. Watch glass

b. Crucible _

c. Beaker

d. Evaporating dish

8. A student used the below ruler to measure the length of a nail. The length of this nail,

according to the precision of the ruler, is — D

a. 3.5 cm

b. 3.55 cm

c. 3.7 cm

d. 3.75 cm

Scientific notation and significant digits:

1. Express the following numbers in standard exponential form:

a) 89000 8.9 x 104

b) 0.000 000 643 6.43 x 10-7

c) 7800 x 10-7

7.800 x 10-4

d) 0.00043 x 10-3

4.3 x 10-7

e) 0.005 x 1015

5 x 1012

2. Perform the following mathematical operations, write all answers in standard exponential

form and be sure to round correctly.

a) (9 000 000) (0.00853) (542 800) = 4.1670756 x 1010

= 4 x 10

10

b) 56.8900 — 0.039 = 56.851 = 57 = 5.7 x 101

c) ( 4.861 x 105 ) ( 8.20 x 10

-6 ) ( 0.005712 ) = 1.18657442 x 10

14 = 1.2 x 10

14

( 8.03 x 10-17

) ( 730 ) ( 0.00329 )

d) 3.89 x 10823

+ 9.7623 x 10825

= 9.8012 x 10825

= 9.80 x 10825

Metric conversions:

1. Perform the following conversions:

a) 3.5 L = __3500_ cm3 d) 436 ng = __4.36 x 10

-7__ g

b) 0.75 kg = __7.5 x 10 5__ mg e) 7800 µm = __7.8 x 10

-2__ dm

c) 5000 cm3 = _5 x 10

6__ mm

3 f) 4.89 x10

-7 cm = _ 4.89 _nm

2. The angstrom ( A ) is a unit of length used for measuring small dimensions, such as the

diameters of atoms. 1 A = 1 x 10-10

m. If the diameter of an atom is 3 A , what is its

diameter in millimeters? in kilometers?

1 A = 1 x 10-10

m so 3 A = 3 x 10-10

m

3 x 10-10

m • 1000 mm = 3 x 10-7

mm 3 x 10-10

m • 1 km = 3 x 10-13

km

1 m 1000 m

3. Distinguish between accuracy and precision.

Accuracy = how close is the data to the true or real value

Precision = how close is the measurable data to itself.

4. First measurement: 6.293 g; Second measurement: 6.294 g ; Third measurement: 6.295 g A

student obtained these data after measuring the mass of an object three different times. If the

true value of the object’s mass is 5.550 g, these data are best described as — A

a) precise but not accurate _

b) accurate but not precise

c) accurate and precise

d) neither accurate nor precise

5. Four students each took three temperature readings of a sample of water. The actual

temperature of the sample was 80.0ºC. Which student’s measurements were both accurate

and precise? D

a) Student 1 c) Student 3

b) Student 2 d) Student 4

Atomic Structure:

1. Using the Periodic Table, determine the following information concerning the element

gold.

a) symbol ___Au__

b) atomic mass ___197___

c) atomic number ___79___

d) number of protons ___79____

e) number of neutrons ___118_____

f) number of electrons ___79____

2. Calculate the percent abundance of copper if the average atomic mass of copper is 64.5

amu. The two isotopes of copper are found in nature to have masses of 63.0 amu and

65.0 amu.

65 X + 63 (1-X) = 64.5 65 X + 63 - 63X = 64.5

2 X + 63 = 64.5 2 X = 1.5

X = .75

65 • 75% + 63 • 25 % = 64.5 amu

3. Determine the average atomic mass for lead given the following information:

1.37 % Pb-205; 26.26 % Pb-206; 20.82 % Pb-207; and 51.55 % Pb-208.

1.37 % 205 + 26.26 % 206 + 20.82 % 207 + 51.55 % 208 = 207.2255 amu

4. A particle of matter contains 6 protons, 7 neutrons and 6 electrons. This particle must be a A

a. neutral carbon atom b. neutral nitrogen atom

c. positively charged carbon ion

d. positively charged nitrogen ion

5. Complete the following table:

IsIsotope Symbol Atomic No. Atomic Mass # Protons # Neutrons # Electrons Hydrogen-2

H 1 2 1 1 1

Strontium 83

Sr 38 83 38 45 38

Uranium U 92 238 92 146 92

Mercury Hg 80 201 80

121 80

6. Which of the following particles has the least mass? A

a. an electron b. a proton

c. an alpha particle

d. a neutron

7. The atoms in a sample of an element must contain nuclei with the same number of B a. electrons

b. protons c. neutrons

d. nucleons

8. Which of these conclusions can be drawn from Rutherford’s experiment? D a. Each atom contains electrons.

b. The nucleus of an atom can be split.

c. Each atom contains protons.

d. Atoms are mostly empty space.

9. A strontium atom differs from a strontium ion in that the atom has a different: A

a. number of electrons b. number of protons

c. atomic number

d. mass number

10. A chloride ion (Cl-) has the same number of electrons as a neutral atom of — C

a. fluorine

b. sulfur

c. argon

d. bromine

11. After 1911, most scientists accepted the theory that the nucleus of an atom was very

dense and very small and had a positive charge. What led scientists to accept this

theory? D a. Dalton’s theory of the atom was over 100 years old.

b. Scientists before 1911 used the scientific method of inquiry improperly.

c. A new model proved that the quantum theory of the atom was inaccurate.

d. Rutherford did an experiment firing alpha particles at a thin piece of gold foil.

12. How does the radioactive isotope C-14 differ from its stable counterpart C-12? C a. It has a different number of protons and two less neutrons than C-12.

b. It has the same number of protons and two more electrons than C-12.

c. It has the same number of protons but two more neutrons than C-12. _

d. It has a different number of protons and two more neutrons than C-12.

13. Chlorine forms a 1– ion. How many electrons does a chloride ion have? D a. 1 c. 17

b. 16 d. 18

14. Which of these elements is the most chemically active? A a. F

b. Cl

c. Br

d. I

15. Which of the following orbital diagrams is incorrect because it violates Hund’s rule? H

Dimensional analysis

1. The first measurement of sea depth was made in 1840 in the central South Atlantic, where

a plummet was lowered 2425 fathoms. What is the depth in meters? Note that

1 fathom = 6 ft and 1 inch = 2.54 cm.

2425 fathoms • 6 feet • 12 inches • 2.54 cm • 1 meter = 4434.84 m or 4435 m

1 fathom 1 feet 1 inches 100 cm

2. In a laboratory exercise to determine the density of a substance, a student found the mass of the

substance to be 6.00 grams and the volume to be 0.50 mL. Expressed to the correct number of

significant figures, the density of the substance is F

a. 0.08 g/mL

b. 0.083 g/mL

c. 0.0833 g/mL

d. 12.00 g/mL

e. 12.0 g/mL

f. 12 g/mL

3. Which Kelvin temperature is equal to –73°C? (C = 273 – K or K = C + 273) C

a. 100 K

b. 173 K

c. 200 K d. 346 K

4. An over the counter medicine has 325 mg of its active ingredient per tablet. How

many grams is this? D a. 325,000 g

b. 3.25 g

c. 32.5 g

d. 0.325 g

Physical and chemical properties/changes and mixtures:

1. List five characteristics of a chemical change.

A change in color, a substance emits light, a change in the smell, emanate heat

or cool and/or decomposes

2. Distinguish between a heterogeneous mixture and a homogeneous mixture. Give at least

three examples of each type of mixture.

Homogeneous mixture is considered homogeneous if it appears uniform to the

eye. That is, they are made up of atomic-sized particles, too small to see.

Homogeneous mixtures are also called Solutions. Heterogeneous mixtures are

also known as suspensions and consist of particles larger than those in a

homogeneous mixture. These particles are visible to the naked eye and can be

easily distinguished from one another

3. When sugar is dissolved in water, with no visible sugar remaining, the resulting solution is

classified as a A

a. homogeneous mixture b. heterogeneous mixture

c. homogeneous compound

d. heterogeneous compound

4. An example of a heterogeneous mixture is A

a. soil b. sugar

c. carbon monoxide

d. hydrogen

5. Which statement describes a characteristic of all compounds? (omit question) C

a. compounds contain one element

b. compounds contain two elements

c. compounds can be decomposed by physical means

d. compounds can be decomposed by chemical means

Bonding and Periodic Trends

1. On a separate sheet of paper, set up a table with the following headings: Formula;

Bond Type (electronegativity difference); Lewis Structure (dot diagram); Shared vs.

Unshared Pairs; Total Pairs; Bonding Orbitals and

the following compounds: CO

CaCl2.

Formula Bond

Type

electronegativity

difference

CO2

PC

1

CCl4

C

0.5

H2O

PC

1.4

PCl3

C

0.9

CH2F2

C

0.4

1.5

AlCl3

PC

1.5

H2CO

PC

0.4

1.4

CaCl2

I

2.0

nd Periodic Trends

On a separate sheet of paper, set up a table with the following headings: Formula;

Bond Type (electronegativity difference); Lewis Structure (dot diagram); Shared vs.

Unshared Pairs; Total Pairs; Bonding Orbitals and Shape. Complete the table using

the following compounds: CO2 ; CCl4; H2O; PCl3; CH2F2; AlCl3; H

Lewis Structure Shared Unshared

Pairs

Total

Pairs

2

0

2

4

0

4

2

2

4

3

1

4

4

0

4

3

0

3

3

0

3

2

0

2

On a separate sheet of paper, set up a table with the following headings: Formula;

Bond Type (electronegativity difference); Lewis Structure (dot diagram); Shared vs.

te the table using

; H2CO; and

Bonding

Orbitals

Shape

p Linear

s p

Tetrahedral

p Bent

p

Trigonal

Pyramidal

s p

Tetrahedral

p

Trigonal

Planar

s p

Trigonal

Planar

s

Linear

2. Given a blank periodic table, show the periodic trends ( atomic size, ionization energy,

electron affinity, electronegativity, and shielding effect ) and how these trends change

from left to right and top to bottom.

3. What is the difference in size of a metallic ion compared to its original atom? What is

the difference in size of a nonmetallic ion compared to its original atom?

Metallic Ions are smaller than their non ions atoms.

Nonmetallic Ions are larger than their non ions atoms

4. Set up a table to show the characteristics of metallic, ionic, and covalent bonds.

Characteristics of …

Ionic Bonds Covalent Bonds Metallic Bonds • Ionic bond, also known as

electrovalent bond

• This bond is formed when one

element pulls the electron from

another element.

• A electrostatic attraction forms

between these oppositely charged

ions in a chemical compound.

• These kinds of bonds occur

mainly between a metallic and a

non metallic atom.

• Boiling Point- High

• Melting point –High

• Polarity - High

• Phase - Solids at room

temperature

• When in solid form, all ionic

compounds are crystalline solids

(the alternating positive and

negative ions are arranged in a

crystal lattice structure)

• Most ionic compounds are

soluble in water

• Covalent bonding is formed

between two non metallic atoms.

• It is characterized by the sharing

of pairs of electrons between

atoms and other covalent bonds.

• Boiling Point - Low

• Melting point – Low

• Phase – Liquids or gases at room

temperature

• Polarity – Low

Generally:

• High stability (it’s not easy to

break a bond of shared

electrons).

• High energy (when a covalent

bond is broken, it releases a

large quantity of energy), and

No net charge.

• Metallic bonding is the bonding

within metals.

• It involves the delocalized

sharing of free electrons (sea of

electrons) between a lattice of

metal atoms.

• The metallic bond is actually a

special case of an ionic bond.

• Boiling Point – High

• Melting point – High

• Phase – Solids at room

temperature

• Polarity – High

• Think of metallic bonding as a

more or less static arrangement

of positive ions within a moving

array of mobile electrons.

Metallic properties include:

• Heat conductivity (mobile

electrons can carry the kinetic

energy of heat),

• Shiny appearance (the rapidly

moving electrons emit energy in

the form of light),

• Electricity conductors

(electricity is the flow of

electrons), and

• Malleability (ability to be easily

shaped into flat sheets or drawn

into wires). .

5. Which of the following groups on the periodic table contains a metalloid? B a. 1 c. 7

b. 13 d. 18

6. Sodium chloride conducts electricity when dissolved in water. What type of bond is

present in NaCl? D a. Nonpolar covalent c. Hydrogen

b. Polar covalent d. Ionic

7. Which of the following shows the correct number of atoms of each element in the

formula Mg(NO3)2 A

a. 1 magnesium atom, 2 nitrogen atoms, and 6 oxygen atoms b. 1 magnesium atom, 2 nitrogen atoms, and 5 oxygen atoms

c. 1 magnesium atom, 1 nitrogen atom, and 6 oxygen atoms

d. 1 magnesium atom, 1 nitrogen atom, and 5 oxygen atoms

8. An alien astronaut landed on Earth and created the periodic table shown. The

astronaut was trying to determine what type of bond would be present in several

compounds. The type of bond in a compound containing G and E would be — D

a. a metallic bond

b. a nonmetallic bond

c. a covalent bond

d. an ionic bond

9. Which of the following is the correct Lewis electron-dot diagram for the sodium atom? A

10. What shape does the molecule BF3 have? D a. Bent c. Tetrahedral

b. Linear d. Trigonal planar

11. - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - C

12. The shape of PCl3 molecule is _____. E a. bent

b. trigonal planar

c. linear

d. tetrahedral

e. trigonal pyramid

13. Which of the following would not be considered a bond? B a. Covalent bond

b. Hydrogen bond c. Ionic bond

d. Metallic bond

Nuclear Chemistry

1. Iodine-131 is a radioactive isotope with a half-life of 8 days. How many grams of a 64 g

sample of iodine-131 will remain at the end of 24 days? D a. 56 g

b. 48 g

c. 32 g

d. 8 g

64g (½)3 = Mf

64g (⅛) = Mf

8g = Mf

2. A gamma ray is best described as having A

a. no charge and no mass b. a negative charge and no mass

c. a positive charge and a mass number of 2

d. a positive charge and a mass number of 4

3. In the following reaction the isotope represented by X is: ����

��� � � �

� A

a. ���

���

b. Th90

230

c. U92

236

d. Th90

233

4. Write the balanced equation for the following nuclear decays. (4 pts total)

a. Alpha decay of Astatine – 196.

b. Electron capture of Aluminum- 26

Light and Energy 1. A hydrogen lamp emits several lines in the visible region of the spectrum. One of these

lines has a wavelength of 656 nm. What are the color, frequency, and energy of this

radiation?

A Spectrum Chart will be provided for the Exam. The wavelength corresponds to Red.

The first formula is c= λν (cc == vveelloocciittyy ooff lliigghhtt == 33..0000 xx 110088 mm//sseecc,, GGrreeeekk lleetttteerr ““nnuu””,, ν,, ffoorr

ffrreeqquueennccyy aanndd == λ wwaavvee lleennggtthh))..

33..0000 xx 110088 mm//sseecc == 656 nm • ν (must covert nm to meters – 656 nm = 6.56 x 10

-7m)

33..0000 xx 110088 mm//sseecc == 6.56 x 10

-7m • ν

6.56 x 10-7

m 6.56 x 10-7

m

0.4576170732 x 1015

1/sec or sec-1

or hertz = 4.58 x 1014

1/sec or hertz

The second formula for energy is E = h ν (E = energy, h = Planck's Constant 6.626 x 10-34

Js

and ν,, ffoorr ffrreeqquueennccyy))

E = h ν E =6.626 x 10-34

Js • 4.58 x 1014

1/sec

E = 30.34708 x 10-20

J = 3.03 x 10-19

J

2. What is the frequency of light with a wavelength of 8.87 x 10-5

cm.?

The formula is c= λν 33..0000 xx 110088 mm//sseecc == 8.87 x 10

-5 cm • ν (convert units)

33..0000 xx 110088 mm//sseecc == 8.87 x 10

-7 m • ν

8.87 x 10-7

m 8.87 x 10-7

m

0.3382187148 x 1015

1/sec or hertz 3.34 x 1014

1/sec or hertz

3. Determine the wavelength and identify the color of light associated with a frequency of

5.49 x 1014

s-1

.

The formula is c= λν 33..0000 xx 110088 mm//sseecc == λ • 5.49 x 10

14 s

-1

5.49 x 1014

s-1

5.49 x 1014

s-1

0.546448087 x 10-6

m = 5.46 x 10-7

m

Quantum Numbers

1. Write electron configurations, orbital filling diagrams, and electron dot diagrams for the

following elements:

a) Al

1s

22s

22p

63s

23p

1

b) Mn

1s

22s

22p

63s

23p

64s

23d

5

c) Ag

1s

22s

22p

63s

23p

64s

23d

104p

65s

24d

9

d) Rb

1s

22s

22p

63s

23p

64s

23d

104p

65s

1

2. For the element Ag assign quantum numbers for the electrons found in the outermost

shell in the electron configuration to the end of the configuration.

Silver is a transition metal and has electrons in the 4d orbital, but its outer energy

level is 5s.

5s 4d n=5, l=0, ml=0, ms=+½ n=4, l=2, ml=-2, ms=+½

n=5, l=0, ml=0, ms=-½ n=4, l=2, ml=-1, ms=+½

n=4, l=2, ml=0, ms=+½

n=4, l=2, ml=1, ms=+½

n=4, l=2, ml=2, ms=+½

n=4, l=2, ml=-2, ms=-½

n=4, l=2, ml=-1, ms=-½

n=4, l=2, ml=0, ms=-½

n=4, l=2, ml=1, ms=-½

3. A particle of matter contains 6 protons, 7 neutrons and 6 electrons. This particle must

be a A

a. neutral carbon atom

b. neutral nitrogen atom

c. positively charged carbon ion

d. positively charged nitrogen ion

4. What is the total number of electrons in an Mg2+

ion? A

a. 10 b. 2

c. 12

d. 14

5. The mass number of an atom is always equal to the total number of its D a. electrons

b. protons

c. electrons and protons

d. neutrons and protons

6. What is the mass number of an ion that has 83 protons, 80 electrons and 126 neutrons? C a. 83

b. 206

c. 209 d. 289

7. An atom of an element has an electron configuration of 1s22s

22p

2. What is the total

number of valence electrons in this atom? D a. 6

b. 2

c. 5

d. 4

8. What are the numbers of protons, neutrons, and electrons in an isotope of titanium with a

mass number of 50? D

a. 22 p, 22 n, 28 e

b. 28 p, 22 n, 22 e

c. 50 p, 22 n, 50 e

d. 22 p, 28 n, 22 e

Nuclear Chemistry

1. Write a balanced nuclear equation for each of the following reactions.

a) Polonium (Po) is a naturally radioactive element decaying with the loss of an alpha particle.

b) Thorium-234 (Th) undergoes beta particle production.

c) Lead-210 undergoes beta particle production.

d) Potassium-38 decays by positron production.

e) Lead-212 decays by beta particle and gamma ray emission.

f) Curium-240 (Cm) decays by alpha particle production.

2. Phosphorus-32 is a commonly used radioactive nuclide in biochemical research, particularly in

studies of nucleic acids. The half-life of phosphorus-32 is 14.3 days. What mass of

phosphorus-32 is left of an original sample of 175 mg after 35.0 days?

MO (½)n = Mf

n = 35/14.3

MO = 175 g

175g (½) 35/14.3

= Mf

175 mg 14.3

½ 35

= Mf

175 mg (7.71 x 10-5

) = Mf

32.1 mg = Mf

Heat and Energy

1. If 82.4 kJ of heat is applied to a 121.2 gram block of metal, the temperature increases by 20.5 oC. Calculate the specific heat capacity of the metal in J/g

oC.

SS..HH.. == hheeaatt ((ccaall))

mmaassss((gg)) xx cchhaannggee iinn tteemmpp((ººCC))

8822..44 kkJJ •• 11000000 JJ // 11kkJJ == 8822440000 JJ

SS..HH.. == 8822440000 JJ == 3333..1166442299220044 JJ//ggooCC == 3333..22 //gg

ooCC

112211..22 gg •• 2200..55ººCC

Naming compounds and writing formulas:

1. Name the following:

a) Na2S ___Sodium Sulfide____________

b) H2O2 ____ Hydrogen peroxide______

c) HBr _____Hydrobromic Acid_______

d) H ( C2H3O2 ) ____Acetic Acid________

e) Sn ( SO4 )2 ______Tin (IV) Sulfate____

f) N2O4 ____ Dinitrogen tetroxide ________

2. Write formulas for the following:

a) perchloric acid _______ HClO4 ______

b) aluminum hyposulfite ____Al2(SO2)3___

c) carbon tetrachloride ______CCl4______

d) hydroiodic acid ________HI__________

e) ammonia _______NH3_______________

f) ferrous permanganate Fe(MnO4) ___