Name
Date
Class
INTRODUCTION TO CHEMISTRY
Practice Problems
In your notebook, solve the following problems.
CHEMISTRY
1. Match the project to the appropriate field of chemistry (inorganic chemistry, organic chemistry, biochemistry,
analytical chemistry, or physical chemistry).
a. Determine the composition of a moon rock sample.
b. Do research on making a new medicine to treat high blood pressure.
c. Investigate ways to regulate the rate of gasoline burning in an automobile engine.
d. Develop a plastic that can be decomposed by bacteria.
e. Improve the method for extracting iron from iron ore.
2. Classify the following examples as examples of pure chemistry or applied chemistry.
a. developing a shampoo to be used with dry or damaged hair
b. determining the conditions required for materials to burn
c. figuring out the general structure of materials such as cotton and silk
d. designing a large-scale method for producing nylon
e. explaining why water expands when it freezes
CHEMISTRY FAR AND WIDE
1. Identify three areas of energy research that scientists are working on today.
2. The following statements are all concerned with the work chemists do. Write T for each true statement and F for
each false statement.
a. Chemists design materials to meet specific needs.
b. Oil from the soybean plant is used to make biodiesel.
c. As the world’s population increases, the amount of land available to grow food increases.
d. Many drugs are effective because they interact in a specific way with chemicals in cells.
e. The trend in crop protection is toward chemicals that are less specific.
f. The use of lead paint in houses was banned in 1978.
g. Chemists are doing research to improve batteries.
h. To study the universe, chemists gather data from afar and analyze matter that is brought back to Earth.
i. Chemists have developed a plastic “skin” that can heal itself when it cracks to help patients with burns.
THINKING LIKE A SCIENTIST
1. One cold morning your car does not start. Make two hypotheses about why the car will not start.
2. Suppose you try several experiments with your car. You try a battery jump, which does not work. There seems to be
enough gas in the car. You wiggle a wire in the engine, and the car starts on the next try. Explain how these tests
help you decide what was wrong with the car.
3. The following is a list of observations from everyday experiences:
Hummingbirds have long beaks.
Moisture forms on the outside of a cold glass.
Ice cubes float.
Oil and water don't mix.
There are fewer fish in a particular creek this year.
a. Propose one hypothesis for each observation.
b. Select one of the hypotheses and describe an experiment that you could do to test it.
4. Discuss the statement “No theory is written in stone.”
PROBLEM SOLVING IN CHEMISTRY
1. Apples are selling for $1.50 a pound. Each apple weighs, on average, 0.50 pounds. You have $6.00. How many
apples can you purchase?
a. ANALYZE (List the knowns and unknown.)
Knowns:
Unknown:
cost of apples =
number of apples purchased = ?
weight of an apple =
dollars available =
b. CALCULATE (Solve for the unknown.)
Use an expression that converts cost per pound to cost per apple.
cost per apple =
Use an expression that relates cost per apple to dollars available.
$6.00
number of apples purchased =
$0.75
number of apples purchased =
2. Describe an alternate way to solve Problem 1.
MATTER AND CHANGE
Practice Problems
In your notebook, solve the following problems.
PROPERTIES OF MATTER
1. Which of the following is not a physical change?
a. dissolving sugar in water
c. evaporating sea water to obtain salt
b. burning gasoline in an engine
d. slicing a piece of bread
2. Which of the following is not a property of a gas?
a. has a definite shape
c. assumes the shape of its container
b. has an indefinite volume
d. is easily compressed
3. Which of the following is not a physical property of sucrose?
a. solid at room temperature
c. dissolves in water
b. decomposes when heated
d. tastes sweet
4. Which of the following is in a different physical state at room temperature than the other three?
a. salt
b. sugar
c. flour
d. water
5. Complete the following table.
Physical state
Definite Shape?
Definite Volume?
Easily Compressed?
gas
no
no
yes
6. Classify the following properties as extensive or intensive.
a. color
b. volume
c. mass
d. boiling point
MIXTURES
1. How might you separate a mixture of water and salt?
2. What is a homogeneous mixture?
3. Which of the following mixtures are homogeneous? Which are heterogeneous?
a. gasoline
b. chunky peanut butter
c. oil and vinegar salad dressing
4. Which of the following are substances? Which are mixtures?
a. ethanol
b. motor oil
c. vinegar
d. neon
ELEMENTS AND COMPOUNDS
1. What elements make up ammonia, chemical formula NH3?
2. Name the elements represented by the following chemical symbols.
a. Pb
b. K
c. Au
d. Fe
3. Classify the following as elements, compounds, or mixtures.
a. table salt
b. water
c. iron
d. stainless steel
4. Write the chemical symbol for each of the following elements.
a. tin
b. sodium
c. silver
d. carbon
5. A liquid is allowed to evaporate and leaves no residue. Can you determine whether it was an element, a compound,
or a mixture?
6. Which of the following is not an element?
a. copper
b. sulfur
c. sucrose
d. helium
CHEMICAL REACTIONS
1. Which one of the following is a chemical change?
a. Gasoline boils.
c. Gasoline burns.
b. Oxygen is added to gasoline.
d. Gasoline is poured into a tank.
2. Classify each of the following changes as physical or chemical.
a. A puddle is dried by the sun.
c. Bread is toasted.
b. A dark cloth is faded by sunlight.
d. Soap is mixed with water.
3. Carbon dioxide plus water yields carbonic acid.
a. Name the product(s) of this reaction.
b. Name the reactant(s) of this reaction.
4. If 44 grams of carbon dioxide react completely with 18 grams of water, what is the mass of carbonic acid formed?
5. In an engine, octane combines with oxygen to form carbon dioxide and water. If 22.8 grams of octane combine
completely with 80 grams of oxygen to form 70.4 grams of carbon dioxide, what mass of water is formed?
6. What is the name of the chemical law on which problems 4 and 5 are based?
ATOMIC STRUCTURE
Practice Problems
In your notebook, solve the following problems.
DEFINING THE ATOM
1. According to Figure 5.2, 100,000,000 copper atoms would form a line 1 cm long. How long would a line formed by
1 x 107 copper atoms be? Express your answer in millimeters.
STRUCTURE OF THE NUCLEAR ATOM
1. A sulfur-32 atom contains 16 protons, 16 neutrons, and 16 electrons. What is the mass (in grams) of a sulfur-32
atom?
2. The mass of a neutron is 1.67 x 10-24 g. Approximately what number of neutrons would equal a mass of one gram?
3. Which statement is consistent with the results of Rutherford’s gold foil experiment?
a. All atoms have a positive charge.
b. Atoms are mostly empty space.
c. The nucleus of an atom contains protons and electrons.
d. Mass is spread uniformly throughout an atom.
DISTINGUISHING BETWEEN ATOMS
1. How many protons are found in an atom of each of the following?
a. boron
c. neon
b. sulfur
d. lithium
2. Complete the table for the following elements.
Element
Number of
Protons
Manganese
25
Number of
Electrons
11
35
Mass
Number
39
89
12
45
Yttrium
Arsenic
33
Actinium
75
227
3. How many neutrons are in each atom?
a. 23 Na
c. 81 Br
11
35
b. 238U
92
Atomic
Number
30
Sodium
Bromine
Number of
Neutrons
d. 19
9F
4. The two most abundant isotopes of carbon are carbon-12 (mass = 12.00 amu) and carbon-13 (mass = 13.00 amu).
Their relative abundances are 98.9% and 1.10%, respectively. Calculate the atomic mass of carbon.
5. Element X has two isotopes: X-100 and X-104. If the atomic mass of X is 101 amu, what is the relative abundance
of each isotope in nature?
ELECTRONS IN ATOMS
Practice Problems
In your notebook, solve the following problems.
MODELS OF THE ATOM
1. How many sublevels are in the following principal energy levels?
a. n = 1
c. n = 3
e. n = 5
b. n = 2
d. n = 4
f. n = 6
2. How many orbitals are in the following sublevels?
a. 1s sublevel
d. 4f sublevel
g. fifth principal energy level
b. 5s sublevel
e. 7s sublevel
h. 6d sublevel
c. 4d sublevel
f. 3p sublevel
3. What are the types of sublevels and number of orbitals in the following energy levels?
a. n = 1
c. n = 3
b. n = 2
d. n = 4
e. n = 5
ELECTRON ARRANGEMENT IN ATOMS
1. Write a complete electron configuration of each atom.
a. hydrogen
d. barium
g. krypton
b. vanadium
e. bromine
h. arsenic
c. magnesium
f. sulfur
i. radon
THE PERIODIC TABLE
Practice Problems
In your notebook, solve the following problems.
ORGANIZING THE ELEMENTS
1. Which element listed below should have chemical properties similar to fluorine (F)?
a. Li
b. Si
c. Br
d. Ne
2. Identify each element as a metal, metalloid, or nonmetal.
a. fluorine
b. germanium
c. zinc
d. phosphorus
e. lithium
3. Which of the following is not a transition metal?
a. magnesium
b. titanium
c. chromium
d. mercury
4. Name two elements that have properties similar to those of the element potassium.
5. Elements in the periodic table can be divided into three broad classes based on their general characteristics.
What are these classes and how do they differ?
CLASSIFYING THE ELEMENTS
1. Use the periodic table to write the electron configuration for silicon. Explain your thinking.
2. Use the periodic table to write the electron configuration for iodine. Explain your thinking.
3. Which group of elements is characterized by an s2p3 configuration?
4. Name the element that matches the following description.
a. one that has 5 electrons in the third energy level
b. one with an electron configuration that ends in 4s24p5
c. the Group 6A element in period 4
5. Identify the elements that have electron configurations that end as follows.
a. 2s22p4
b. 4s2
c. 3d104s2
6. What is the common characteristic of the electron configurations of the elements Ne and Ar? In which group would
you find them?
7. Why would you expect lithium (Li) and sulfur (S) to have different chemical and physical properties?
8. What characterizes the electron configurations of transition metals such as silver (Ag) and iron (Fe)?
PERIODIC TRENDS
1. Explain why a magnesium atom is smaller than atoms of both sodium and calcium.
2. Predict the size of the astatine (At) atom compared to that of tellurium (Te). Explain your prediction.
3. Would you expect a Cl– ion to be larger or smaller than an Mg2+ ion? Explain.
4. Which effect on atomic size is more significant, an increase in nuclear charge across a period or an increase in
occupied energy levels within a group? Explain.
5. Explain why the sulfide ion (S2–) is larger than the chloride ion (Cl–).
6. Compare the first ionization energy of sodium to that of potassium.
7. Compare the first ionization energy lithium to that of beryllium.
8. Is the electro negativity of barium larger or smaller than that of strontium? Explain.
9. What is the most likely ion for magnesium to form? Explain.
10. Arrange oxygen, fluorine, and sulfur in order of increasing electro negativity.
IONIC AND METALLIC BONDING
Practice Problems
In your notebook, answer the following.
IONS
1. For each element below, state (i) the number of valence electrons in the atom, (ii) the electron dot structure, and
(iii) the chemical symbol(s) for the most stable ion.
a. Ba
b. I
c. K
2. How many valence electrons does each of the following atoms have?
a. gallium
b. fluorine
c. selenium
3. Write the electron configuration for each of the following atoms and ions.
a. Ca
c. Na+
e. O2–
b. chlorine atom
d. phosphide ion
4. What is the relationship between the group number of the representative elements and the number of valence
electrons?
5. How many electrons will each element gain or lose in forming an ion? State whether the resulting ion is a cation
or an anion.
a. strontium
c. tellurium
e. bromine
b. aluminum
d. rubidium
f. phosphorus
6. Give the name and symbol of the ion formed when
a. a chlorine atom gains one electron.
b. a potassium atom loses one electron.
c. an oxygen atom gains two electrons.
d. a barium atom loses two electrons.
7. How many electrons are lost or gained in forming each of the following ions?
a. Mg2+
b. Br–
c. Ag+
d. Fe3+
8. Classify each of the following as a cation or an anion.
a. Na+
c. I–
e. Ca2+
b. Cu2+
d. O2–
f. Cs+
IONIC BONDS AND IONIC COMPOUNDS
1. Use electron dot structures to predict the formula of the ionic compounds formed when the following elements
combine.
a. sodium and bromine
d. aluminum and oxygen
b. sodium and sulfur
e. barium and chlorine
c. calcium and iodine
2. Which of these combinations of elements are most likely to react to form ionic compounds?
a. sodium and magnesium
c. potassium and iodine
b. barium and sulfur
d. oxygen and argon
3. What is the meaning of coordination number?
4. How is the coordination number determined?
BONDING IN METALS
1. What is a metallic bond?
2. How is the electrical conductivity of a metal explained by metallic bonds?
3. Are metals crystalline? Explain.
4. Give three possible crystalline arrangements of metals. Describe each.
5. What is an alloy?
6. Name the principal elements present in each of the following alloys.
a. brass
d. sterling silver
b. bronze
e. cast iron
c. stainless steel
f. spring steel
COVALENT BONDING
Practice Problems
In your notebook, solve the following problems.
MOLECULAR COMPOUNDS
1. Classify each of the following as an atom or a molecule.
a. Be
c. N2
b. CO2
d. H2O
e. Ne
2. Which of the following are diatomic molecules?
a. CO2
c. O2
b. N2
d. H2O
e. CO
3. What types of elements tend to combine to form molecular compounds?
4. What information does a molecule’s molecular structure give?
5. How do ionic compounds and molecular compounds differ in their relative melting and boiling points?
THE NATURE OF COVALENT BONDING
1. Draw the electron dot structure for hydrogen fluoride, HF.
2. Draw the electron dot structure for phosphorus trifluoride, PF3.
3. Draw the electron dot structure for nitrogen trichloride, NCl3.
4. Draw the electron dot configuration for acetylene, C2H2.
5. How many resonance structures can be drawn for CO32–? Show the electron dot structures for each.
BONDING THEORIES
1. Predict the shape and bond angle for the compound carbon tetrafluoride, CF4.
2. Predict the shape and bond angle for phosphorus trifluoride, PF3.
3. Predict the type of hybridized orbitals involved in the compound boron trichloride, BCl3.
4. What types of hybrid orbitals are involved in the bonding of the silicon atoms in silicon tetrafluoride, SiF4?
5. Predict the shape and bond angle of fluorine monoxide, F2O.
6. Predict the shape of the CH2CF2 molecule. What hybridization is involved in the carbon-carbon bonds?
7. How many sigma and pi bonds are used by each of the carbon atoms in the following compound?
POLAR BONDS AND MOLECULES
1. What type of bond—nonpolar covalent, polar covalent, or ionic—will form between each pair of atoms?
a. Na and O
b. O and O
c. P and O
2. Explain why most chemical bonds would be classified as either polar covalent or ionic.
3. Would you expect carbon monoxide and carbon dioxide to be polar
or nonpolar molecules?
4. Draw the structural formulas for each molecule and identify polar covalent bonds by assigning the slightly positive
(δ+) and slightly negative (δ–) symbols to the appropriate atoms.
a. NH3
b. CF3
5. Which would you expect to have the higher melting point, CaO or CS2?
CHEMICAL NAMES AND FORMULAS
Practice Problems
In your notebook, solve the following problems.
NAMING IONS
1. What is the charge on the ion typically formed by each element?
a. oxygen
c. sodium
e. nickel, 2 electrons lost
b. iodine
d. aluminum
f. magnesium
2. How many electrons does the neutral atom gain or lose when each ion forms?
a. Cr3+
c. Li+
e. Cl–
b. P3–
d. Ca2+
f. O2–
3. Name each ion. Identify each as a cation or an anion.
a. Sn2+
c. Br–
e. H–
b. Co3+
d. K+
f. Mn2+
4. Write the formula (including charge) for each ion. Use Table 9.3 if necessary.
a. carbonate ion
c. sulfate ion
e. chromate ion
b. nitrite ion
d. hydroxide ion
f. ammonium ion
5. Name the following ions. Identify each as a cation or an anion.
a. CN–
c. PO43–
e. Ca2+
b. HCO3–
d. Cl–
f. SO32–
NAMING AND WRITING FORMULAS FOR IONIC COMPOUNDS
1. Write the formulas for these binary ionic compounds.
a. magnesium oxide
c. potassium iodide
e. sodium sulfide
b. tin(II) fluoride
d. aluminum chloride
f. ferric bromide
2. Write the formulas for the compounds formed from these pairs of ions.
a. Ba2+, Cl–
c. Ca2+, S2–
e. Al3+, O2–
b. Ag+, I–
d. K+, Br–
f. Fe2+, O2–
3. Name the following binary ionic compounds.
a. MnO2
c. CaCl2
e. NiCl2
g. CuCl2
b. Li3N
d. SrBr2
f. K2S
h. SnCl4
4. Write formulas for the following ionic compounds.
a. sodium phosphate
c. sodium hydroxide
e. ammonium chloride
b. magnesium sulfate
d. potassium cyanide
f. potassium dichromate
5. Write formulas for compounds formed from these pairs of ions.
a. NH4+, SO42–
+
b. K , NO3
c. barium ion and hydroxide ion
–
d. lithium ion and carbonate ion
6. Name the following compounds.
a. NaCN
c. Na2SO4
e. Cu(OH)2
b. FeCl3
d. K2CO3
f. LiNO3
7. Name and give the charge of the metal cation in each of the following ionic compounds.
a. Na3PO4
c. CaS
e. FeCl3
b. NiCl2
d. K2S
f. CuI
NAMING AND WRITING FORMULAS FOR MOLECULAR
COMPOUNDS
1. Name the following molecular compounds.
a. PCl5
c. NO2
e. P4O6
g. SiO2
b. CCl4
d. N2F2
f. XeF2
h. Cl2O7
2. Write the formulas for the following binary molecular compounds.
a. nitrogen tribromide
c. sulfur dioxide
b. dichlorine monoxide
d. dinitrogen tetrafluoride
NAMING AND WRITING FORMULAS FOR ACIDS AND BASES
1. Name the following compounds as acids.
a. HNO2
b. H2SO4
c. HF
2. Write the formulas for the following bases.
a. calcium hydroxide
c. aluminum hydroxide
b. ammonium hydroxide
d. lithium hydroxide
d. H2CO3
THE LAWS GOVERNING FORMULAS AND NAMES
1. Write the formulas for these compounds.
a. potassium sulfide
e. hydrobromic acid
i. sulfur hexafluoride
b. tin(IV) chloride
f. aluminum fluoride
j. magnesium chloride
c. hydrosulfuric acid
g. dinitrogen pentoxide
k. phosphoric acid
d. calcium oxide
h. iron(III) carbonate
l. nitric acid
2. Complete this table by writing correct formulas for the compounds formed by combining positive and negative ions.
SO42–
NO3–
OH–
PO43–
Ca2+
Al3+
Na+
Pb4+
3. Name the following compounds.
a. K3PO4
c. NaHSO4
e. N2O5
g. PI3
b. Al(OH)3
d. HgO
f. NBr3
h. (NH4)2SO4
4. Explain the difference between the law of definite proportions and the law of multiple proportions.
CHEMICAL REACTIONS
Practice Problems
In your notebook, solve the following problems. Use the 3-step problem-solving approach you learned in Chapter 1.
DESCRIBING CHEMICAL REACTIONS
1. Write the skeleton equation for the reaction between hydrogen and oxygen that produces water.
2. Write the skeleton equation for the reaction that produces iron(II) sulfide from iron and sulfur.
3. Write the skeleton equation representing the heating of magnesium carbonate to produce solid magnesium oxide
and carbon dioxide gas.
4. Write a balanced equation for the production of HCl gas from its elements.
5. Write a sentence that completely describes the chemical reaction represented by this balanced equation.
2HCl(aq) + CaCO3(s) → CO2(g) + CaCl2(aq) + H2O(l)
6. Write the word equation for the following equation. Write a sentence fully describing the reaction. Is the
equation correctly balanced? Explain.
2Ag(s) + S(s) → Ag2S(s)
7. Write a balanced equation representing the formation of aqueous sulfuric acid from water and sulfur trioxide gas.
8. Write a balanced equation from this word equation.
aqueous silver nitrate + copper metal → silver metal + aqueous copper nitrate
9. Write a balanced equation for the following word equation.
phosphorus + oxygen → tetraphosphorous decoxide
TYPES OF CHEMICAL REACTIONS
1. Write a balanced equation representing the reaction of magnesium with oxygen gas to produce magnesium
oxide.
2. Write the balanced equation for the reaction that occurs between aluminum and fluorine.
3. Write the balanced equation for the production of oxygen gas and potassium chloride from the decomposition of
potassium chlorate.
4. Write the balanced equation for the reaction between hydrochloric acid and calcium metal. The products are
hydrogen gas and calcium chloride.
5. Write the balanced equation for the combustion of propane (C3H8) to produce carbon dioxide and water vapor.
6. Write the balanced equation for the reaction between iron(III) chloride and sodium hydroxide. The products are
iron(III) hydroxide and sodium chloride.
7. Classify each of the reactions in problems 1–6 as to type.
8. Use the activity series of metals (Table 11.2) and your knowledge of the relative reactivity of the halogens to predict
whether the following reactions will occur. Write balanced equations for those reactions that do occur.
a. Br2(l) + NaCl(aq) →
b. Ca(s) + Mg(NO3)2(aq) →
c. K(s) + H2SO4(aq) →
d. Zn(s) + NaOH(aq) →
REACTIONS IN AQUEOUS SOLUTION
1. Write the net ionic equation for the reaction between aqueous barium nitrate, Ba(NO3)2, and sodium sulfate, Na2SO4.
2. Magnesium reacts with HCl to form hydrogen and magnesium chloride. Write the balanced net ionic equation for
this reaction.
3. The double-replacement reaction below results in the formation of the precipitate lead chloride. Balance the equation
and write the net ionic equation.
Pb(NO3)2(aq) + NH4Cl(aq) → PbCl2(s) + NH4NO3(aq)
4. Identify the precipitate formed when solutions of the following ionic compounds are mixed. If no precipitate is
formed, write no precipitate.
a. Zn(NO3)2 + SnCl2 →
b. KCl + AgNO3 →
c. Cu(NO3)2 + Na2S →
d. Al2(SO4)3 + 3Mg(OH)2 →