Chapter 1: Matter and Change
1. In which form are the atoms the most closely packed together: Solid, liquid, or gas?
2. Which of the following is not a chemical property: decomposing, burning, rusting, or breaking?
3. What physical property changes as gravity changes?
4. What are three ways to know that a chemical reaction occurred?
5. What is the first step in the scientific method?
Chapter 2 Measurements and Calculations
1. Find the mass in grams of a material that occupies 2.43 mL and has a density of 1.46 g/mL.
2. Round each of the following measurements to the number of significant figures indicated.
a. 43.56789 grams to five significant figures
b. .0001 liters to two significant figures
c. 12.301 grams to three significant figures
d. .0156 moles to one significant figure
3. Identify the SI prefix equivalent of each of the numerical values
a. 100
b. 0.01
c. 0.000000001
d. 0.000000000001
e. 0.1
4. A student measures the volume and mass of a substance and calculates it density as 2.53 g/mL. The accepted
value of the density is 2.34 g/mL. What is the percentage error of the student’s measurement?
5. Complete the following conversions
a. 2.35 km =
m
b. 3.24 mL =
L
c. 1.7 g =
kg
d. 456.7 mL =
cm3
Chapter 3 Atoms
1. Which number of the following symbol is the atomic mass?
He24
2. Which number of the following symbol is the atomic number? Al1326
3. Write the following with the correct Atomic symbol? Atomic mass: 35.45, Atomic number: 17, Symbol: Cl
4. Calculate the molar mass of the following element. 152.09g(92%abundance), 150.23g(4.3%abundance),
147.73g(3.7% abundance)
5. What letter represents the element Tungsten in its atomic symbol?
Chapter 4: Arrangement of Electrons in Atoms
1. What is the difference between the ground state and the excited state of an atom?
2. Fill out the table below:
Location
Symbol
List Two Facts
Charge
otons
eutrons
ectrons
3. How did the Heisenberg uncertainty principle contribute to the idea that electrons occupy “clouds,”or
“orbitals”?
4. Write the ground-state electron configuration and orbital notation for each of the
following atoms: Oxygen, Boron, and Carbon
5. List the order in which orbitals generally fill, from the 1s to the 4f orbita
Type your answer using the following format: 1s, 2s, 3s
Use a comma between each orbital in the list.
Chapter 5 - The Periodic Law – Questions
1. Which scientist discovered a trend in elements in the order of increasing atomic weight, and later the
scientist, Henry Moseley, discovered a better way to for elements to fit their pattern better, which was
according to what characteristic?
2. Name the electron configuration for each of the three elements, with one of them using noble gas
configuration:



Magnesium
Selenium
Actinium
3. Name 2 terms of periodic trends, define them, and explain their relationship with each other (i.e how
they increase or decrease).
4. In a list, define these elements according to their atomic number and mass, and from least to greatest,
arrange them according to their ionization energy and another list for atomic radii.



Fluorine
Hydrogen
Cesium
5. Make your own definition of the periodic law and explain one reason on how useful it is to the modern
world.
Chapter 6: Chemical Bonding, Covalent
1. How do the electrons behave in a covalent bond?
2. Draw a Lewis Dot Structure between the single covalent bonded compound CF4.
3. What is the difference between a polar covalent bond and a nonpolar covalent bond?
4. How do you determine if a molecule is polar covalent or nonpolar covalent?
5. Draw a Lewis Dot Structure of N2 and name the type of bond.
Chapter 6 Chemical Bonding Ionic Problems/ Questions
1. Name each of the following compounds using the Stock System.
A. ZnI2
B. PbO2
C. AgF
D. CsBr
E. NaCl
2. Give the name of the following binary compounds.
A. BaCl2
B. Sr3P2
C. KCl
D. MgCl2
3. How many Fe3+ and S2- ions would be in one formula unit of the ionic compound formed by these specific ions
present?
4. Write formulas for the binary ionic compounds formed between the following elements.
A. Calcium and Sulfur
B. Calcium and Chlorine
C. Potassium and Bromine
5. Write the formulas for and give the names of the compounds formed by the following ions.
A. Cl- and Pb2+
B. Na+ ClC. S2- Sn4+
Chapter 6 Chemical Bonding Metallic – There are currently no questions for this section
Chapter 7: Covalent Names and Formulas
1. Fill in the prefix for the following numbers.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
2. Fill in the chart below.
Formula
1. NO
2. CF6
3. NF3
Compound
4. SO2
5. NO3
3.Fill in the chart below
Compound Name
Formula
1. phosphorous trichloride
2. sulfur hexafluoride
3. dinitrogen pentoxide
4. boron trichloride
5. carbon dioxide
4. State the octet rule and describe its exceptions. (There are two exceptions – hydrogen and boron)
5.How are Lewis dot structures useful? Provide an example in your explanation.
Chapter 7:
1. What is
Chemical Names and Formulas; Ionic
the chemical formula for the compound formed by the aluminum ion and oxide ion?
2. What is the name of the above compound?
3. Name As2S3.
4. What is the formula mass of H2O?
5. What is the probable bond that will form with K and Cl?
Chapter 8: Chemical Equations and Reactions
1.) Write the balanced chemical equation:
Potassium Nitrate + Magnesium Iodide  Potassium Iodide + Magnesium Nitrate
________________________________________________________________________________
2.) Complete the word equation:
Potassium + Sulfide  ___________________
Write the balanced chemical equation: ___________________________________________________
3.) Determine which type of reaction:
 2KBr(s) + Cl2(s)  2KCl(s) + Br2 (S)
___________________________________
 2LiF(s) + CaCl2(s)  2LiCl(s) + CaF2(s) ___________________________________
4.) Write the chemical formula for the compounds:
 Calcium Chloride _________________________
 Sodium Nitrate _________________________
 Iron(III) Chloride _________________________
5.) Balance the following equations:
 __Co(s) +__ F2(s)  __CoF3(s)
 __Ca(s) + __H2O(l)  __Ca(OH)2(aq) + __ H2(g)
 __KI(s) + __Pb(NO3)2(aq)  PbI2(s) + KNO3(aq)
Chemistry Chapter 9 Stoichiometry Review
1.) If 5.10 mol of ethane, C2H6, undergo combustion according to the unbalanced equation given below, how many
moles of oxygen are required?
C2H6 + O2
CO2 + H2O
2.) Sulfuric acid reacts with sodium hydroxide according to the following equation.
H2SO4 + NaOH
Na2SO4 + H2O
What mass of H2SO4 would be required to react with 0.89 mol NaOH?
3.) 2 H2(g) + O2(g)
2 H2O(g)
(a) According to the equation, how many moles of hydrogen would be required to produce 8.0 mol of water?
(b) How many moles of oxygen would be required?
4.) Coal gasification is a process that converts coal into methane gas. If this reaction has a percentage yield of 87.0%,
what mass of methane can be obtained from 960 g of carbon?
2 C(s) + 2 H2O(l)
CH4(g) + CO2(g)
5.) Sulfuric acid reacts with aluminum hydroxide by double replacement.
(a) If 31.0 g of sulfuric acid react with 27.0 g of aluminum hydroxide, identify the limiting reactant.
(b) Determine the mass of excess reactant remaining.
(c) Determine the mass of each product formed. Assume 100% yield.
aluminum sulfate
Chapter 10 States of Matter Questions
1. What are the four variables that affect gas properties?
2. What is happening during deposition? Is it endothermic or exothermic? Provide an example.
3.
a. Label lines, boxes, and circles.
b. Label supernatural fluids.
c. Which state is more dense? Solid or liquid?
4. Calculate the molar enthalpy of vaporization of a substance given that .225 mol of the substance absorbs 44.8 kJ
of energy when it is vaporized.
5. What are the five equations used to solve this heating curve? What do numbers 1-5 represent?
1.
2.
3.
4.
5.
Chapter 11: Gas Laws
1. How many grams of oxygen will occupy a volume of 1750 mL at 2.30 atm and 25C?
2. A gas measures 55.0 mL at 1.30 atm and 273 K. What will the volume be at 20C and 2.70 atm?
3.The occupied volume of a gas is 145 L and has a pressure of 23.0 torr. If the pressure is increased to 52.0 mmHg, what
is the volume?
4. A 250 mL sample of nitrogen is heated from 32C to 66C, with a pressure remaining constant. What is the volume
after the heating process?
5. A container with 25 mL volume has 4.00 atm of gas at 30C that is heated to become 55C. What is the pressure of the
gas after the container has been heated?
Chapter 12: Solutions questions
1. How many moles of NaCl are in a 78.0 ml of a 3.4 M solution of NaCl in water?
2. Determine the number of kilograms of H2O must be added to 65.4g of KNO3 to form a .544m solution
3. A solution is made from C2H5OH, and water is 2.43 m in C2H5OH. How many grams of C2H5OH are contained
per 160 g of water?
4. Given the balance equation: 2H3PO4 + 3Ca(OH)2 = Ca3(PO4)2 + 6H2O, determine the mass of Ca3(PO4)2
when 650ml of 5.0M H3PO4reacts to the equation.
5. If a saturated solution of AgNO3 in 100 g of H2O at 100°C is cooled to 40°C, how many grams of
the solute will precipitate out of the solution? Values of AgNO3 at 20°C,216g 40°C,311g 60°C,440g
80°C,585g 100°C,733g.
Chapter 13: Colligative Properties – there are currently no questions for this section
Chapter 14: Acids and Bases
1. How can you tell the difference between an acid and a base?
2. What happens if you react an acid and a base?
3. How can you tell the difference between monoprotic acids and polyprotic acids?
4. What are some examples of Arrhenius, Bronsted Lowrey, and Lewis acids and bases?
5. Where does acid rain come from?
Chapter 15
1. What is the pH of a 0.038 M HF solution?
2. What is the pH of a 0.0023 M LiOH solution?
3. What is the volume of titration standard at the equivalency point of this titration curve below? What is the pH of
the solution at this point? Is the acid strong or weak? Is the base strong or weak?
4. Determine whether the following acids are strong or weak: HCl, HF, C2H4O2, HBr
5. Determine the following names or chemical formulas of the following acids
HClO4:
Hydrobromic acid:
HCl:
Nitric acid:
HF:
Acetate:
HNO2:
Sulfuric acid:
Chapter 16: Reaction Energy
1. Assume there are 35.21 grams of steam at 112.4*C. Determine how much energy is needed to convert ice at 22*C.
Steps
Step 1
Step2
Step3
Equation
Calculation
Energy
Step4
Step 5
Total Energy
2. Draw the heating curve for water. Make sure to label what each step represents on a typical heating curve.
Is it exothermic or endothermic?
3. Determine the percent composition of HCL and NaCl.
H=__________________
Cl=__________________
Na=_________________
CL=__________________
4. If 3.53 g of C3H8O(l) are reacted with 6.54 g H2Cr2O7 and with more than enough HCl, how many grams of CrCl3
can be produced?
A: Balanced equation _________________________________________________________________
B: What are the limiting and excess reactants?
_______________________
________________________
C: How much excess reactant will remain after the reaction has gone to completion?
D: What is the percent yield of the reaction if 22.3 g of CrCl3 are produced in the laboratory?
5: Calculate the enthalpy change for the overall reaction based on equations 1 and 2.
Ca(s) + ½ O2 (g) + CO2 (g)  CaCO3 (s)
Rxn 1: Ca(s) + ½ O2 (g)  CaO (s)
Δ = -631.5 kJ
Rxn 2: CaCO3 (s)  CaO (s) + CO2 (g)
Δ = 178.3 kJ
Chapter 17: Questions/Problems for Reaction Kinetics
1.
2.
3.
4.
5.
Define Chemical kinetics and reaction rate
Explain the collision theory.
List 4 major factors that affect the rate of reaction and describe how each factor affects the rate.
What is the species that changes the rate of a reaction, but is neither consumed nor changed?
What are the values for the heat of reaction and the activation energy for reaction below?
Chapter 18- Chemical Equilibrium
1) How is a reversible reaction indicated?
2) Within a reversible reaction, how does one know when chemical equilibrium has been established?
3) What does it mean if:
a) K = 1 ?
b) K > 1 ? (Which reaction is favored, forward or reverse?)
c) K < 1 ? (Which reaction is favored, forward or reverse?)
eq
eq
eq
4) Write out the equilibrium expression for the equation:
3A(aq) B(aq) → ← 2C(aq) 3D(aq)
5) Consider the following equation:
2NO(g) O (g) → ← 2NO (g)
At equilibrium, [NO] = 0.80 M, [O ] = 0.50 M, and [NO ] = 0.60 M. Calculate the value of K for this
reaction.
2
2
2
2
Chapter 19 + 20: Electrochemisty (oxidation and redox) (Problem Sheet)
Directions: Determine the Oxidation # of the elements in each of the following compounds: (Q 1-3)
1. CuSO4
2. Fe2O3
3. PO4 (-3 charge)
Directions: Write the oxidation and reduction half reactions: (Q 4)
4. Zn + 2HCl -> H2 + ZnCl2
Directions: Identify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent: (Q 5)
5. Zn + 2HCl -> H2 + ZnCl2
Chapter 22 Organic Chemistry
1. Carbon atoms have a unique ability to form a diverse group of compounds and structures. Why is this?
2. Which of the following is a saturated hydrocarbon and why?
H H H H
H H H H
I I
I
I
I
H – C– C – C = C
3.
I
I
I
H–C–C–C–C–H
I I
I
I
I
I
I
H H
H
H H H H
Write the chemical formula for each hydrocarbon below
Methane = ______
Nonyne = _______
Heptene = ______
4. Name the following hydrocarbons
C3H8 = _______
C10H20 = ______
C5H8 = _______
5. Name and write out the formulas of the four functional groups that contain oxygen
Chapter 22 Organic Chemistry
1. Carbon atoms have a unique ability to form a diverse group of compounds and structures. Why is this?
2. Which of the following is a saturated hydrocarbon and why?
H H H H
H H H H
I I
I
I
I
H – C– C – C = C
3.
I
I
I
H–C–C–C–C–H
I I
I
I
I
I
I
H H
H
H H H H
Write the chemical formula for each hydrocarbon below
Methane = ______
Nonyne = _______
Heptene = ______
4. Name the following hydrocarbons
C3H8 = _______
C10H20 = ______
C5H8 = _______
5. Name and write out the formulas of the four functional groups that contain oxygen