Name_____________________________
Honors Chemistry
FINAL EXAM REVIEW
MAIN TOPICS TO REVIEW
 INTRODUCTION TO CHEMISTRY AND MEASUREMENT
o Definition of chemistry
o The scientific method

Qualitative and quantitative data

Independent and dependent variables

Data analysis (tables and graphs, accepting/rejecting a hypothesis, etc.)
o Lab safety and equipment
o Measurement using various equipment (rulers, graduated cylinders, balances, etc.)
o Precision and accuracy
o Percent error
o Scientific notation, Significant figures rules
o Calculations using significant figures
o Metric prefixes
o Factor label method/Dimensional analysis, Metric to Metric; English to Metric Conversions
o Temperature conversions (Fahrenheit, Celsius, and Kelvin)
 MATTER AND ENERGY
o Definition of matter, Three main states of matter
o Physical and chemical properties of matter
o Physical and chemical changes
o Heating and cooling curves
o Phase diagrams
o Evidence of a chemical reaction
o Density calculations
o Pure substances and mixtures
o Separation of heterogeneous/homogeneous mixtures (chromatography, filtration, etc.)
o Potential and kinetic energy
o Specific heat (specific heat of water), Calorimetry (q=mcΔT)
o Conservation of energy
o Enthalpy, endothermic and exothermic reactions
 CHEMICAL FOUNDATIONS
o Atomic history
o Law of constant composition, multiple proportions, and conservation of mass
o Dalton’s Atomic Theory
o Thomson’s and Rutherford’s contributions to the Modern Atomic Theory
o The charge, location, and mass of the 3 subatomic particles
o Ions and isotopes

Cation, Anion

Relative abundance

Nuclide symbols
o Average atomic mass and mass number
 MODERN ATOMIC THEORY/QUANTUM MECHANICS
o The relationship of energy, wavelength, and frequency of light

Speed of light constant

Planck’s constant
o Electromagnetic spectrum
o Visible light spectrum
o The Bohr model of the atom
o Photon, quanta, ground state and excited state
o Atomic emission and absorption spectra
o Einstein, Heisenberg, deBroglie, and Schrodinger
o Shapes and energy of atomic orbitals
o Orbital diagrams
o Electron configurations including exceptions and ions

Handy Order Filling Chart – (AKA Aufbau Principle)

Pauli Exclusion Principle

Hund’s Rule
o Quantum numbers
o Relate an atom’s electron configuration to its position on the periodic table.
 THE PERIODIC TABLE
o Define the terms group, family, and period in terms of the periodic table.
o Families on the periodic table and their properties
o Trends of the periodic table

Atomic Radius, Ionic Radius

Valence electrons

Reactivity

Ionization Energy, Electron Affinity & Electronegativity

Electron configurations
 CHEMICAL BONDING
o The formation of chemical bonds
o Octet rule
o Lewis dot diagrams
o Ionic and molecular compounds and their properties
o Types of Chemical Bonds (Ionic Bonds & Covalent Bonds)
o Properties of bonds (polarity, length, energy, etc.)
o Attraction and repulsion forces
o Dipole moment
o Intermolecular forces (dipole-dipole, hydrogen bond, London dispersion)
o Crystal lattice
o Metallic compounds
o VSEPR, molecular shapes and bond angles
o Resonance structures
 NOMENCLATURE
o Chemical names and symbols
o Polyatomic ions
o Oxidation numbers
o Name and write the formula for each of the following:

Ionic compounds (classic and stock)

Covalent compounds (classic and stock)

Acids

Hydrates
 THE MOLE
o Define the mole
o Molar mass
o Avogadro’s number
o Conversions from grams to moles to number of atoms, formula units, molecules, etc.
o Percent Composition, Empirical and molecular formulas
 CHEMICAL REACTIONS
o Write a balanced chemical equation from a description including states
o Identify the 5 types of reactions and any specific classification for the main type
o Predict the products of a chemical reaction by name and by formula
o Use the activity series and solubility charts
o Total & Net ionic equations
 STOICHIOMETRY
o Mole ratios from balanced equation
o Reaction stoichiometry
o Limiting and excess reactants
o Percent yield
o Heat (enthalpy) of reaction’s relationship to stoichiometry
o Gas stoichiometry
 SOLUTIONS
o Aqueous solutions, electrolytes/nonelectrolytes
o Molarity
o Dilutions and preparing solutions
o Titrations
o Acids and Bases
o Chemical indicators
o pH and pOH calculations
 SOLIDS, LIQUIDS AND GASES
o Pressure: Atmospheric and Vapor
o Intermolecular forces in solids, liquids, and gases
o Factors that affect solubility, boiling point, etc.
o Characteristics of a solid, liquid, and gas
o Pressure conversions
o Dalton’s Law of partial pressure
o Water displacement by a gas
o Boyle’s Law, Charles’Law, Gay-Lussac’s Law
o Combined Gas Law, Ideal Gas Law
FINAL REVIEW PRACTICE
DIRECTIONS: Complete on a separate sheet of paper. Show all work.
INTRODUCTION TO CHEMISTRY AND MEASUREMENT
1. Is the following set of data accurate, precise, both, or neither if the actual value is 23.4 mL?
20.4 mL
20.3 mL
20.1 mL
20.5 mL
2. Indicate the number of significant figures in the following:
a. 1040
b. 600
c. 12.40
d. 0.00230
e. 0.82
3. Solve each expression using correct number of significant figures and units.
a. 24.4g + 110g
b. 145.30 mL – 12.8 mL
c. 9.8 g / .123 L
d. 1.4500 cm x 2.55 cm x 1.5 cm
4. Solve the following using dimensional analysis (factor label method).
a. Convert 23.5 yd3 to ft3.
b. A pound of coffee beans yields 50 cups (this number is obtained by counting) of coffee.
How many milliliters of coffee can be obtained from 1.00 g of coffee beans? (1 cup = 240mL)
c. Convert 23.50 m/s2 to km/min2.
5. Convert the following to the indicated temperatures
a. -34 ̊C to ̊F
b. 213 K to ̊C
6. Explain how a Bunsen burner works.
7. What is the most practical and accurate tool used to measure volume in a laboratory?
8. Explain how you would set up a laboratory apparatus for
a. heating a crucible of magnesium
b. heating a beaker of copper II oxide and water
9. A student takes an object with an accepted mass of 200.00 grams and masses it on his own
balance. He records the mass of the object as 196.5 g. What is his percent error?
10. What value does each of the following metric unit represent?
a. centi
b. nano
c. kilo
d. milli
e. micro
MATTER AND ENERGY
1. Solve the following density problems using significant figures.
a. What is the density (g/mL) of a block of metal 6cm x 10cm x 100cm that weighs 6820 dg?
What would be the volume (in L) of 2.8 kg of the metal?
b. A block of metal has a mass of 42.3 g. It is placed in a graduated cylinder containing 56.0
cm3 of water. If the final volume is 75.1 mL, what is the density of the metal block?
2. How would you separate…..
a. iron filings and pepper
3. What is the specific heat of water?
b. salt and sand
c. methanol and water
4.
Which requires more heat to warm from 22.0 ̊C to 85.0̊C if they have exactly the same mass, a
sample of aluminum (c = .902 J/ g ̊C) or a sample of water. Explain.
5. A .0234 kg sample of copper cools from 53 ̊C to - 9.4 ̊C. What is the change in its energy if the
specific heat of copper is .387 J/g ̊C?
6. Distinguish between a physical and chemical change.
7. Classify the following as either a physical or chemical change:
a. spoiling of milk
b. softening glass to bend it into a new shape
c. burning a piece of paper
d. rusting of a nail
8. Identify the following as a pure substance, heterogeneous mixture, or homogeneous mixture.
a. copper
b. sweetened tea
c. calcium carbonate
d. sandy water
9. What is the difference between an exothermic reaction and an endothermic reaction? Relate each
reaction type to enthalpy.
10. Using the potential energy diagram above, answer the following questions.
a. How much energy do the reactants have?
b. How much energy is needed to get the reaction started?
c. What type of reaction is this, endothermic or exothermic?
CHEMICAL FOUNDATIONS
1. What is the atomic mass of silver? What is silver’s mass number?
2. Determine the symbol for the atom, ion, and or isotope that fits the description below.
a. 16 p, 18 e, 16 n
b. 26 p, 26 e, 30 n
c. 13 p, 10 e, 11 n
d. 35 p, 35 e, 45 n
3. Determine the identity of an element with a mass number of 75, that contains 42 neutrons.
4. Determine the number of protons, neutrons, and electrons in the following:
a. Ar – 40
b. F-1
c.
29
Al
d. Co
e.
56 Fe
26
5. Distinguish between an alpha and beta particles.
6. If 36 grams of water are produced and 4 grams of hydrogen are used in its production, how many
grams of oxygen is needed? What law is represented here?
7. What is the difference between the mass number for Carbon 14 and carbon’s average atomic
mass of 12.011 amu?
8. Calculate the atomic mass of lithium if one isotope has a mass of 6.0151 amu and a percent
abundance of 7.59% and a second isotope has a mass of 7.0600 amu and a percent abundance of
92.41%.
9. Describe the debate between Democritus and Aristotle in terms of the atom.
10. List the five postulates to Dalton’s theory. Explain two reasons why it is no longer accepted.
11. Compare and contrast Rutherford and Thomson’s experiments and atomic models.
12. What significance does each subatomic particle have in regards to an element’s properties?
MODERN ATOMIC THEORY/QUANTUM MECHANICS
1. What is the wavelength (in m) of a photon of light containing 3.19 x 10 -13 J of energy?
2. How much energy in Joules will a photon of light with a wavelength of 345.0 nm possess?
3. What is the maximum number of electrons that can occupy the following?
a. 3d
b. 3rd energy level
c. 4s
d. one orbital
e. 4th energy level
4. Draw the orbital diagrams for the following elements:
a. argon
b. copper
c. nitrogen
5. Write the electron configuration for the following elements/ions:
a. calcium
b. arsenic
c. molybdenum
d. magnesium ion
e. sulfur ion
6. Write the noble gas configuration for the following elements.
a. Tin
b. Gold
c. Chromium
7. Answer the following questions about the electron configuration of aluminum:
a. What does the 3 in 3s2 mean?
b. What does the 2 in 3s2 mean?
c. How many valence electrons does this atom have?
d. How many electrons will an atom of aluminum (lose or gain) to form an ion?
e. What orbitals will aluminum lose its electrons from?
8. What is needed for an electron to move from one energy level to a higher energy level?
9. What can be determined from an atom’s atomic emission spectra?
10. What happens when an electron falls back to its ground state?
THE PERIODIC TABLE
1. Write the symbol of the element that fits each description:
a. alkali metal in period 3
b. alkaline earth metal with 5 energy levels
c. the only metalloid chalcogen (oxygen group)
d. halogen with four energy levels
2. How many valence electrons does each of the following possess? What ion does each form?
a. calcium
b. silver
c. phosphorus
d. lead
e. carbon
3. Which is smaller?
a. Ca or Ca+2
b. P or O
c. N or F
d. Sulfur ion or Aluminum ion
4. Does each of the following trends increase or decrease down a family or across (left to right) a
period?
Down a Family
Across a Period
Electronegativity
Ionization Energy
Atomic Radius
5. Define electron affinity. Which elements have a negative value for electron affinity? Positive
value? Why?
6. Draw the Lewis structures for the following elements.
a. P
b. Ni
c. Al
7. Why are the alkali metals one of the most reactive families of elements?
8. Which groups make up the representative or main group elements?
9. Give three physical properties that differ between metals and nonmetals.
10. What is a metalloid? List each of them.
CHEMICAL BONDING
1. What is the difference between an ionic and a covalent bond?
2. Compare the properties of an ionic compound and a molecular compound.
3. Would the following pairs most likely form an ionic or molecular compound?
a. Mg and Cl
b. I and F
c. P and Cl
d. Sn and O
e. Ag and S
d. O
e. Li
4. Draw a Lewis dot diagram for each of the following:
a. Be
b. N
c. Al
5. Using electronegativity values, write the ranges for a polar covalent, non polar covalent, and ionic
bonds. What is the significance of these electronegativity differences?
6. Describe the three types of intermolecular forces. Give one example of each.
7. Draw the Lewis structures for each of the following compounds:
a. N2
b. O3
c. H2S
d. NCl3
i. Identify the total number of valence electrons for each.
e. SiF4
ii. What is the molecular shape (linear, bent, trigonal pyramidal, trigonal planar,
tetrahedral)?
iii. What is the bond angle for each shape?
iv. Polar or nonpolar covalent bonds?
v. Polar or nonpolar molecule?
NOMENCLATURE
1. Name each of the following:
a. Ba(CN)2
m. HNO3 (aq)
b.
CoBr2
n. Al(OH)3
c.
P2O5 (STOCK)
o. HBrO4 (aq)
d.
NaH
p. KMnO4
e.
MgI2
q. Hg3(PO3)2 (CLASSIC)
f.
Ca3(AsO4)
r. Na2O2ME
g.
HCl
s. LiC2H3O2
h.
Sn(SO4)2
t. FeCl3
i.
SO2
u. CCl4
j.
BaSO3
v. AgNO3
k.
Zn(ClO4)2
w. N2O3 (STOCK)
l.
Be(BrO)2
x. Pb3N2 (STOCK)
(CLASSIC)
∙ 6 H2O (CLASSIC)
2. Write the formula for each of the following:
a.
Diboron Hexahydride
m. Mercuric Sulfate
b.
Ammonium Chloride
n.
Auric Bromide
c.
Iron (III) Nitrate
o.
Cadmium Oxide
d.
Phosphoric Acid
p.
Calcium Hypochlorite
e.
Titanium (IV) Bromide
q.
Lead (II) Phosphate
f.
Cuprous Phosphide
r.
Magnesium Sulfate Heptahydrate
g.
Selenium (VI) Fluoride
s.
Silicon Dioxide
h.
Sodium Cyanide
t.
Sulfuric Acid
i.
Hydroselenic Acid
u.
Aluminum Chlorate
j.
Phosphorus Triiodide
v.
Magnesium Sulfide
k.
Barium Chloride Dihydrate
w.
Plumbous Acetate
l.
Potassium Hydroxide
x.
Manganese (III) Sulfite
THE MOLE
1. What is the molar mass of each of the following?
a. Al(C2H3O2)3?
b. Ammonium chloride
c. Potassium bromate
2. I know that 8.401 x 1024 atoms of a certain element has a mass of 391.8 g. What element is this?
3. How many Sulfide ions are present in 12.24 mg of potassium sulfide?
4. How many carbon atoms are in 1 mole of dicarbon hexahydride?
5. How many grams are in 2.34 x 1023 molecules of C6H10OS2?
6. How many molecules can be found in a sample of .847 grams of carbon dioxide?
7. Determine the percent composition of all of the elements found in the following compounds:
a. Sodium chlorate
b. Cuprous sulfide
8. What is the difference between an empirical and molecular formula?
9. Determine the empirical formula for a compound using the following information:
a. 58.00% Rb; 9.50% N; 32.50% O
b. 49.62% C; 10.82% H; 39.56% O
10. What is the molecular formula of a compound with an empirical formula of CHOCl (chlorine) and
a molecular weight of 129 grams?
11. What are the 7 diatomic elements?
12. Determine the molecular formula for a compound containing 7.70% carbon, 1.30% hydrogen,
and 81.90% iodine. Its molar mass is 281 g/mol.
13. If you take 20.8 g of aluminum hydrate and heat it to remove all of the water, what is the formula
of the original hydrate if the mass of the remaining solid after heating is 15.4 g?
CHEMICAL REACTIONS
1. For each of the following: (NOTE: some of these are pretty advanced… use the link to help)
a. Indicate the type of the reaction
b. Predict the products by name
c. Write the equation for the complete reaction. If the reaction does not occur, write NR.
d. For DD reactions include the states
e. Balance the reaction only if it occurs
(http://schs.stvrain.k12.co.us/registration/summerassign.schedule%20changes/AP%20Chemistry/Writ
ing,%20Balancing%20and%20Predicting%20Products%20of%20Chemical.pdf)
a. silver nitrate + cupric chloride 
b. copper (II) + aluminum sulfate 
c. sodium nitrite + ammonium sulfate 
d. iodic acid 
e. hydrochloric acid + barium hydroxide 
f. sodium + chromium (II) chlorate 
g. cupric oxide 
h. C4H8 + oxygen (complete) 
i. dibromine pentoxide + water 
j. iron (II) oxide + carbon dioxide 
k. chlorous acid + magnesium hydroxide 
l. aluminum carbonate 
m. plumbic sulfide + hydrochloric acid 
n. hydrobromic acid + calcium carbonate 
2. What is the difference between oxidation and reduction?
3. Write the net ionic equations for the reaction of
a. hydrochloric acid + barium hydroxide 
b. chlorous acid + calcium hydroxide 
c. plumbic sulfide + hydrochloric acid 
d. silver nitrate + cupric chloride 
4. What are the four main signs that a chemical reaction has occurred?
5. Balance the following redox reaction in a basic environment:
MnO4- + SO2  SO4 -2 + Mn+2
STOICHIOMETRY
1. If you react 25.4 g of bromine with an excess of sodium iodide and you produce 2320 mg of
iodine, what is the percent yield of iodine?
2. How many grams of potassium sulfate are produced if 10.0 g of sulfuric acid and 7000.0 mg of
potassium hydroxide are mixed together? What is the limiting reactant?
3. How many moles of chloric acid are required to react completely with 2.05 g of sodium carbonate?
4. The reusable booster rockets of the U.S. space shuttle employ a mixture of aluminum and
ammonium perchlorate for fuel. The balanced equation is
3Al(s) + 3NH4ClO4  Al2O3(s) + 3NO(g) + 6H2O(l) NH3(g)
What mass of ammonium perchlorate should be used in a fuel mixture for every 1.00 kg of
aluminum used?
5. How many grams of magnesium chloride are produced if 420 cm3 of chlorine gas reacts with 14.8
g of Magnesium at STP?
6. If the following reaction produces 6840 J of energy, how many grams of CrO 3 is produced?
CrO3 + H2O  H2CrO4 + 5.4 kJ
7. How much heat is absorbed/released when 193 g of ammonium bromide react according to the
following equation?
NH3 + HBr  NH4Br
ΔH = 188.32 kJ
SOLUTIONS
1. If 960 g of sodium hydroxide are used to prepare 16000 mL solution. What is the molarity of the
solution?
2. If we add .15 L of a 6.02M HCl solution and some water to a beaker, creating a 2.42 M HCl
solution, what is the volume of the new solution?
3. A 2.34 M solution contains 10. g of acetic acid. What is the volume of the solution in mL?
4. What mass of sodium oxide is consumed if it reacts with 12.3 mL of a 3.45 M HCl solution?
5. How many grams of .67 M KClO3 are needed to prepare 1.00 Liters of solution?
6. Describe how 100 ml of a .10 M solution of sodium hydroxide would be made.
7. Using the pH scale, define an acid and a base.
8. Calculate the pH when [H+] = 3.9 x 10 -5. Calculate the [H+] when the pH value is 9.57.
9. Which solution is a poor conductor of electricity? 1M NaCl; 1MHC2H3O2; 1M NaOH; 1M HCl
10. What volume of 3.1M H2SO4 will neutralize 250 mL of .30 M Al(OH)3?
CHAPTER 13 – SOLIDS, LIQUIDS,GASES
1. Describe the conditions that allow a liquid to boil.
2. Compare the particles within a solid, liquid, and a gas.
3. Relate the intermolecular forces within a solid, liquid, and gas.
4. Identify and define the different phase changes of matter.
5. What is vapor pressure?
6. A balloon at 22 °C holds 2.00L of oxygen. If the gas in the balloon is heated to 90° C, what
volume will the balloon be?
7. What is the temperature of .70 moles of a gas that occupies 470 mL at a pressure of 150 kPa?
8. 500.0 mL of O2 gas is collected over water. Atmospheric pressure is 101.0 kPa. The temperature
of the water is 22 °C. What is the pressure of the gas? You will need a chart from your notes to
complete this problem!
9. A gas is at 1.33 atm of pressure and a volume of 682 mL. What will the pressure be if the volume
is reduced to 0.419 L?
10. Nitrogen gas is being held in a 14.3 m3 tank at a temperature of 62oC. What will the volume be
when the temperature drops to 24oC?
11. A gas storage tank is a 1.72 atm and 35oC. What temperature is the gas at if the pressure
increases to 2.00 atm?
12. A gas with a volume of 1.00 L is at 135oC and 844 mm Hg. What is the volume if the conditions
change to 14o C and 748 mm Hg?
13. Calculate the mass of 162 L of chlorine gas, measured at STP.
14. Find the pressure in mm Hg produced by 2.35 g of carbon dioxide in a 5.00 L flask at 18 oC.
15. How many grams of carbon monoxide must be placed into a 40.0 L tank to develop a pressure of
965 mm Hg at 23oC?
16. At what Celsius temperature will argon have a density of 10.3 g/L and a pressure of 6.43 atm?