Final Exam Review for Chemistry Name____________________

Name the following lab equipment.

1. 2. 3.

__________________ ____________________ __________________

4. 5. 6.

__________________ ____________________ __________________

7. 8. 9.

__________________ ____________________ __________________

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Scientific Notation

Compute the following. Write answers in scientific notation.

a. (6.02 × 1023)(8.65 × 104) = ____________________

b. 5.6 ×10 − 18

8.9 × 108 = ____________________

c. (5.4 × 10 4 )(2.2 × 10 7 ) 5.4 ×105 = ____________________

d. 6a × 7b 2 × 8c 3 14 4c2 3 = ____________________

e. (1.2 × 105) + (5.35 × 106) = ____________________

f. (3.67 × 102) ̶ (1.6 × 101) = ____________________

Measurements

What is being measured?

International System Unit

Measurement Device

Mass

Length

Time

Liquid Volume

Temperature

EXAMPLE: What is the Kelvin temperature equivalent to 30⁰C?

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Remember, when using a measurement device, estimate one place further than the device measures.

________ ________ ________

Prefix Abbreviation Relationship to Base Unit

Kilo

Deci

Centi

Milli

EXAMPLE: Rank the following from largest to smallest.

10.0 m 216.50 cm 0.0445 km 0.14 dm

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

1. __________ HCl

2. __________ Cl2

3. __________ barium

4. __________ NaCl(aq)

Classify the following matter as a pure substance or a mixture.

5. __________ titanium

6. __________ air

7. __________ leaf

8. __________ helium

Classify the following matter as homogeneous or heterogeneous.

9. __________ dirt

10. __________ stainless steel

11. __________ Cl2

12. __________ acidic solution

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

13. __________ dissolving sugar into water

14. __________ burning gas

15. __________ decompose water into hydrogen and oxygen

16. __________ frost forms on a car windshield in the morning

17. __________ squeezing orange juice

18. __________ bleaching hair

Classify each of the following as a physical property or a chemical property.

19. __________ height

20. __________ texture

21. __________ is flammable

22. __________ is nonreactive with hydrogen

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1. List 4 observable indications that a chemical reaction may have occurred.

2. What is a precipitate?

Density

Density =

3. Which liquid is the most dense: A, B, C, or D?

4. What is the volume of 60.0 g of ether if the density of ether is 0.70g/mL?

5. A piece of granite 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?

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Use the factor label method to solve the following problems

6. Given: 1 cup = 236 mL and 1 cup = 8 ouncesConvert: 0.5 ounce to mL

7. Given: 946 mL = 1 quart and 4 quart = 1 gallon and 1000 mL = 1 LConvert: 55 gal to L

8. Given: 1 calorie = 4.184 joulesConvert: 785 joules to calories

Atomic Theory

1. Definitions Word Bank: Anion Cation Ion

Atomic Mass Electrons Isotope

a. Any charged particle is called a(n) _______________. A positively charged particle is a(n)

_______________. A negatively charged particle is a(n) ______________ . These are all

formed by losing or gaining _________________.

b. An atom that has a set number of protons, but varying numbers of neutrons is a(n)

_____________.

c. The average mass of all of the isotopes of an element is called _____________________.

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2. Complete the following table.

Symbol Protons Electrons Neutrons Mass Number

Atomic Number

39K

25 30

56 54 82

S-2 18

Al+3 27

How many protons are in Li+1?

How many electrons are in Li+1?

3. Given the following information, what is the atomic mass of iron?

Percent Abundance Isotope Exact Mass

5.82 Iron -54 53.9696 amu

91.66 Iron -55 55.9349 amu

2.19 Iron -57 56.9354 amu

0.33 Iron -58 57.9333 amu

4. Explain two ways that the isotopes 14C and 13C are alike. Explain two ways they are different.

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5. Draw Bohr diagrams for the following. Be sure to show the protons and electrons.

a. F b. C

c. S d. Be

6. Give 3 points of Dalton’s atomic Theory.

7. What was the name of the experiment Rutherford used to discover the nucleus? What did he discover?

Periodic Table

Know the name and location of the following groups

Group 1 ___________________ Group 2 _____________________

Group 3-12 ___________________ Group 17 _____________________

Group 18 ___________________

8. Name the family to which each element belongs:

a. K _______________

b. Kr_______________c. I________________

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The Noble gases are stable because they have ___ ______________ ____________.

Elements are placed in order of increasing ______________ _______________ and placed in groups according to their ___________________ __________________.

Rows = ______________ =___________ __________ _______

Columns =______________ = _______________

Metals are located _________________________________

Nonmetals are located______________________________

Metalloids are located __________________________________

Metals _____________ electrons. They form __________________.

Nonmetals ______________________ electrons. They form __________.

Nonmetals can also _______________ electrons.

9. What results for each of the following tests would you expect to see if we tested chromium (Cr)?

a. __________ Appearance: dull or shiny

b. __________ Flexibility: malleable or brittle

c. __________ Conductivity: Conducts or insulates

10. Identify the following as a metal (M), nonmetal (NM), or metalloid (ML)

a. ________ Has luster and conducts electricity

b. ________ B and As

c. ________ Is not ductile and is shiny

d. ________ Zr, Cu, and Ca

e. ________ An insulator

f. ________ Fe, Mg, Al and Zn

g. ________ C, H, and O

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Trends

Atomic radius (size)

Ionization energy (energy to remove an electron)

Electronegativity (ability to gain an electron)

Reactivity for metal (metallic character)

Reactivity for nonmetal

Multiple Choice

_____ 11. From left to right across the periodic table

a. Ionization energy increases

b. Atomic radius increases

c. Electronegativity decreases

d. Atomic mass decreases

_____ 12. Electronegativity

e. Generally decreases from left to right across a period

f. Is the energy change that accompanies the loss of an electron from a gaseous atom

g. Generally decreases from top to bottom within a group

h. Is generally higher for metals than for nonmetals

13. In each of the following pairs of elements circle the one which best completes the statement.

a) F Se is more active.

b) K Na is the smaller atom.

c) C O has the higher ionization energy.

d) Ba B is an alkaline earth metal

e) Si I has a higher electronegativity

f) Zn Al behaves more like Ga

g) N Cs is a metal

h) Li N is the larger atom

i) S O is in period 2 group 16

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j) Ge Br has higher electronegativity

k) Fe Sb is a metalloid

l) Be F is a halogen

m) Ca Mg is malleable

n) Ag Xe is a gas at room temperature

o) Na P is not ductile

p) Rb Cd is more reactive

q) Au Pb is a transition metal

Chemical Bonding

The three types of chemical bonds are

1. _________________ usually between a ___________ and a ___________. There is a __________ of electrons. The bond is due to attraction between _______________ and _______________.

2. _________________ usually between a ___________ and a ___________. The electrons are _____________.

a) _____________ _______________ between 2 same nonmetals. Electrons are ____________ shared.

b) _____________ _______________ between 2 different nonmetals. Electrons are ____________ shared.

3. _________________ usually between a ___________ and a ___________.

Many electrons are shared. They are ______________.

Classify the following chemical bonds as ionic, covalent or metallic:

_____ CuO _____AlBr3 _____ H2 _____ PCl5 _____ MgSO4

Classify the following chemical bonds as polar covalent or nonpolar covalent:

_____O2 _____CO _____N2 _____PCl5 _____H2O

Characteristics of ionic bonds vs covalent bonds.

Ionic bonds are strong. Ionic substances have orderly pattern and high ____________ ____________. They form ______________. Ionic substances are electrolytes. This means they ___________________ ___________________________ when dissolved in water.

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Covalent substances have __________ melting points and low _____________ _____________. They are _______________________.

1. Is the following most likely an ionic compound or a covalently bonded compound?

a. _____ Made up of P and I

b. _____ Has a boiling point of 100˚C

c. _____ Conducts electricity when dissolved in water

d. _____ Transfers electrons when bonds are formed

e. _____ Is a liquid at room temperature

f. _____ Is made from a metal and a nonmetal

g. _____ exists as a diatomic molecule

2. How do cations form?

3. How do anions form?

4. Fill in the following chart

ElementName

Number ofvalence electrons

Will it gain or lose electrons?

Number of electrons lost or gained

Symbol for ion formed

Is it a cation or anion?

magnesium

sulfur

chlorine

5. Circle the letter of the following pairs of elements that will most likely form an ionic compound. (There may be more than one correct answer.)

a. Calcium and Nitrogen

b. Potassium and Iodine

c. Hydrogen and Sulfur

d. Nitrogen and Phosphorus

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6. Which of the following electron dot structures might represent a halogen?

a.

b.

c.

d.

7. What is the octet rule?

8. Draw the Lewis structure for the following molecules. Then predict the shape of each molecule. See question 14 for the names of the possible shapes.

a. CF4 b. F2O

c. COCl2 d. I2

e. NCl3 f. HCN

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Nomenclature

1. Complete the data table identifying the composition of each ionic compound.

Cation Anion Formula Name

Pb2+ SO4–2

Cu2+ PO43–

Cadmium hydroxide

Ba(N02)2

Al2(CO3)3

Potassium sulfide

2. Complete the table for each covalent compound.

Formula Name

N2O5

Phosphorus pentachloride

Dinitrogen difluoride

SO2

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Tetraphosphorus hexoxide

1. Write the formula for each of the following.

a. Calcium carbonate

b. Tetraphosphorus decasulfide

c. Iron (III) bromide

d. Copper (I) nitrite

e. Ammonium sulfate

f. Sulfur dioxide

g. Sodium bicarbonate

h. Zinc phosphate

2. Write the name for each of the following.

a. FeCl2 ________________________________________________

b. N2O ________________________________________________

c. Cu(OH)2________________________________________________

d. KNO3 ________________________________________________

e. MgF2 ________________________________________________

f. N2F4 ________________________________________________

g. CaO ________________________________________________

h. NO _________________________________

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Moles

1. How many oxygen atoms are in Cu(ClO3)2?

2. Calculate the molar mass of the following:

a. BaCl2 b. (NH4)2SO4

Convert the following:

3. 5.6 × 1026 molecules NaOH to grams NaOH

4. 45g KClO3 to moles KClO3

5. 7.4 moles Sn(SO4)2 to molecules

6. 3.25 moles (NH4)2CO3 to grams

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7. Find the % composition by mass of each element in Fe3(PO4)2.

Chemical ReactionsSynthesis (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)

Identify the precipitate in question #5.

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?

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Predict whether the following single replacement reactions will occur.

6. ________ Ca + AgNO3 →

7. ________ Cu + Li2O →

Heat

Classify the following processes as endothermic or exothermic.

8. __________ clothes dry

9. __________ H2O(g) → H2O(l) + 44.02 kJ

10. __________ a bomb explodes

11. __________ H2O(g) + C(s) + 131.3 kJ → CO(g) + H2(g)

Stoichiometry

12. Given the equation: 2 C2H6 (g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g)

If 4.5 mole ethane (C2H6) undergo combustion, how many moles of oxygen are required?

13. Iron is generally produced from iron ore through the following reaction in a blast furnace:

Fe2CO3 + CO → 2 Fe + 2 CO2

If 4000 g of Fe2CO3 are available to react, how many moles of CO are needed?

14. Methanol is an important industrial compound that is produced from the following reaction:

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CO + H2 → CH3OH (unbalanced)

How many grams of H2 would be needed to produce 100 moles methanol?

15. Given the equation: 2 C2H6 (g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g)

If 4.5 liters ethane (C2H6) at STP undergo combustion, how many liters of oxygen are required?

States of Matter

Solids Liquids Gases

Definite shape

Definite Volume

Fluidity

Particle movement (KE)

Particle arrangement

Attractive forces

Density

Diffusion rate

Compressibility

Heat Calculations

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Q = m Cp ∆T∙ ∙1. What quantity of heat, in joules, is required to raise the temperature of 22.5 g of

lead from 18.0°C to 30.0°C? The Cp of lead is 0.129 J/g·°C.

2. A 0.3 g piece of copper is heated and fashioned into a bracelet. How much energy is required to change the temperature of the copper from 25˚C to 592˚C during the process? The Cp of copper is 390 J/g 0C.

3. What word describes each of the following phase changes?

a. Solid to liquid ____________b. Liquid to gas ____________c. Solid to gas directly ____________

d. Liquid to solid ____________e. Gas to liquid ____________f. Gas to solid directly ____________

4. What section represents the solid phase? __________

5. What section represents the liquid phase? __________

6. What section represents the gas phase? __________

7. What letter represents the triple point? __________

8. What is the substance’s melting point at 1 atm? __________

9. What is the substance’s boiling point at 1 atm? __________

10. At what temperature and pressure do all three phases coexist? __________

11. Would an increase in pressure at 60˚C cause a liquid to freeze or vaporize?

Gases

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Kinetic Molecular Theory assumes gases are made up of _________ ___________moving in _____________ ___________, colliding into each other with ______________ collisions. As temperature increases, the particle movement also _____________________.

Gases do not behave ideally when gases stop moving (or move slowly)That is at _________ temperatures and __________ pressures.

STP means ______________ (760 mm Hg) and _______________

Gas Laws P1V1T2 = P2V2T1 PV = nRT R = 0.0821 L·atm/mol·K

1. Determine the pressure exerted by 0.122 mole oxygen gas in a 1.50 L container at 298 K.

2. Determine the temperature of 4.2 moles SO2 in a 94 L container at 1.2 atm.

3. A sample of hydrogen occupies a volume of 1.2 L at a pressure of 4.2 atm. I the temperature of the gas is kept constant, what would be the new volume of the gas at 6.1 atm?

4. A rigid container of Nitrogen has a pressure of 2910 mm Hg at a temperature of 713 K. What is the pressure at 273 K?

5. What temperature is required to increase the pressure of a 450 mL sample of gas at 60˚C and 1 atm to 2.5 atm if the volume remains constant?

Multiple Choice

1. _____As the temperature of a fixed volume of gas increases, the pressure will

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a. vary inverselyb. be unchanged

c. decreased. increase

2. _____Increasing the volume of a given amount of gas at constant temperature causes the pressure to decrease because

a. the molecules are striking a larger area with the same force.b. there are fewer moleculesc. the molecules are moving more slowly.d. there are more molecules.

3. _____ The volume of a gas is doubled while the temperature is held constant. The pressure of the gas

a. remains unchanged.b. is reduced by one half

c. is doubled.d. depends on the kind of gas.

4. _____ As the temperature of the gas in a balloon decreases a. the volume increases.b. the pressure increases.c. the average kinetic energy of the gas particles decreases.d. All of the above are true.

5. _____ Which of the following would double the pressure on a contained gas at a constant temperature?

a. doubling the volume of the container.b. halving the number of particles in the container.c. doubling the number of particles in the container.d. none of the above.

6. _____ If the temperature of a gas on a close container increasesa. the pressure of the gas decreases.b. the average kinetic energy of the molecules decreases.c. the molecules collide with the walls of the container less frequently.d. the pressure of the gas increases.

7. _____ A temperature of −25⁰C is equivalent to

a. 248 K.b. 25 K.

c. −25 K.d. 298 K.

Solutions

1. Fill in whether the substances in the lab are a solution, colloid, or suspension.

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Station

Do the contents settle out?

Is there a Tyndall Effect?

Is this a solution, colloid, or

suspension?

1 Yes No

2 Yes Yes

3 No Yes

4 No No

2. Classify each of the following mixtures as a colloid, suspension, or solution.

a. ____________________ Fog

b. ____________________ Milk

c. ____________________ Sodium chloride dissolved in water

d. ____________________ Potting soil shaken with water

e. ____________________ Foam

f. ____________________ Pepsi Cola

3. You test the solubility of two solutes -A and B-in three different solvents: water, vegetable oil and hexane. Your results were:

Water Vegetable Oil Hexane

Solute A

dissolved Not dissolved Not dissolved

Solute B Not dissolved dissolved dissolved

a. Classify the solvents as polar or nonpolar.

Water __________ Vegetable Oil __________

Hexane__________

b. Which solute can you classify as nonpolar? Explain.

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Solubility Curves

1. What is the solubility of K2Cr2O7 at 20°C? __________

2. How many grams of CaCl2 will dissolve in 100 g water at 20°C? _________

3. What is the minimum temperature required to completely dissolve 50 g KNO3 in 100 g H2O? __________

5. Classify as saturated or unsaturated a solution that contains 20 g KClO3 in 100 g water at 50°C. Will undissolved salt be present?_________________

6. How does temperature affect the solubility of most solid substances?

7. What mass of KCl will dissolve in 35 g H2O at 80°C?

8. How much water is needed to completely dissolve 60 g NaCl at 30°C?

Concentration

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

9. Calculate the molarity of a solution by dissolving 95.5 g KNO3 in enough water to make 750 mL of solution.

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10. How many moles of H2SO4 would be required to prepare 750 mL of a 0.15 M H2SO4

solution?

11. Calculate the % concentration by mass of a solution in which 20.0 g Mg(OH)2 is dissolved in 80.0 g H2O.

Acids, Bases and Salts

1. List five properties of acids.

2. List 5 properties of bases.

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

acid ← 7 → base

3. Determine if the following is an acid, a base or a salt.

a. ________ Mg(OH)2

b. ________ turns red litmus paper red

c. ________ HNO3

d. ________ barium nitrate

e. ________ produces hydronium ion when dissolved in water

f. ________ pH is 9

4. Determine the pH of the following:

a) 7.3 × 10−9 M HCl ____________________

b) 0.00078 M NaOH ____________________

c) 6.2× 10−5 M LiOH ____________________

d) 0.0075 M HBr ____________________

5. If the pH is ___________ the H+ concentration is __________

a. 4.5 _____________

b. 9.0 _____________

c. 6.7 _____________

6. Name the following acids, bases and salts.

a. Al(OH)3 b. H2SO3

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c. HBr d. K3PO4

7. Write the formula for the following acids, bases and salts.

a. carbonic acid

b. Iron (III) oxide

c. Hydroiodic acid

d. Iron (III) hydroxide

8. Identify the following equations as Ionization, Dissociation, or Neutralization.

a. H2SO3 + H2O → H3O+1 + HSO3-1

H2Ob. Cu(OH)2 → Cu+2 + 2 OH-1

c. Ca(OH)2 + H2SO4 → 2 HOH + CaSO4

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