5 Steps to a 5: AP Chemistry 2017 (2016)


Review the Knowledge You Need to Score High

CHAPTER 5   Basics

CHAPTER 6   Reactions and Periodicity

CHAPTER 7   Stoichiometry

CHAPTER 8   Gases

CHAPTER 9   Thermodynamics

CHAPTER 10   Spectroscopy, Light, and Electrons

CHAPTER 11   Bonding

CHAPTER 12   Solids, Liquids, and Intermolecular Forces

CHAPTER 13   Solutions and Colligative Properties

CHAPTER 14   Kinetics

CHAPTER 15   Equilibrium

CHAPTER 16   Electrochemistry

CHAPTER 17   Nuclear Chemistry

CHAPTER 18   Organic Chemistry

CHAPTER 19   Experimental Investigations




Summary: This chapter on basic chemical principles should serve as a review if you have had a pre-AP chemistry course in school. We assume (and we all know about assumptions) that you know about such things as the scientific method, elements, compounds, and mixtures. We may mention elementary chemistry topics like this, but we will not spend a lot of time discussing them. When you are using this book, have your textbook handy. If we mention a topic and it doesn’t sound familiar, go to your textbook and review it in depth. We will be covering topics that are on the AP exam. There is a lot of good information in your text that is not covered on the AP exam, so if you want more, read your text.

Keywords and Equations

This section of each chapter will contain the mathematical equations and constants that are supplied to you on the AP exam. We have tried to use, as much as possible, the exact format that is used on the test.

Boltzmann’s constant, k = 1.38 × 10−23 J K−1 
electron charge, e = −1.602 × 10−19 coulomb
1 electron volt per atom = 96.5 kJ mol−1

Avogadro’s number = 6.022 × 1023 mol−1

Units and Measurements

Almost all calculations in chemistry involve both a number and a unit. One without the other is useless. Every time you complete a calculation, be sure that your units have canceled and that the desired unit is written with the number.

Always show your units!


The system of units used in chemistry is the SI system (Système International), which is related to the metric system. There are base units for length, mass, etc. and decimal prefixes that modify the base unit. Since most of us do not tend to think in these units, it is important to be able to convert back and forth from the English system to the SI system. These three conversions are useful ones, although knowing the others might allow you to simplify your calculations:

As shown above, the SI unit for volume is the cubic meter (m3 ), but most chemists use the liter (L, which is equal to 1 cubic decimeter (dm3 )) or milliliter (mL). The appendixes list the SI base units and prefixes, as well as some English–SI equivalents.

We in the United States are used to thinking of temperature in Fahrenheit, but most of the rest of the world measures temperature in Celsius. On the Celsius scale, water freezes at 0°C and boils at 100°C. Here are the equations needed to convert from Fahrenheit to Celsius and vice versa:

Many times, especially in working with gases, chemists use the Kelvin scale. Water freezes at 273.15 K and boils at 373.15 K. To convert from Celsius to Kelvin:

K = °C + 273.15

Absolute zero is 0 K and is the point at which all molecular motion ceases.

The density of a substance is commonly calculated in chemistry. The density (D ) of an object is calculated by dividing the mass of the object by its volume. (Some authors will use a lowercase d to represent the density term; be prepared for either.) Since density is independent of the quantity of matter (a big piece of gold and a little piece have the same density), it can be used for identification purposes. The most common units for density in chemistry are g/cm3 or g/mL.


We deal with two types of numbers in chemistry—exact and measured. Exact values are just that—exact, by definition. There is no uncertainty associated with them. There are exactly 12 items in a dozen and 144 in a gross. Measured values, like the ones you deal with in the lab, have uncertainty associated with them because of the limitations of our measuring instruments. When those measured values are used in calculations, the answer must reflect that combined uncertainty by the number of significant figures that are reported in the final answer. The more significant figures reported, the greater the certainty in the answer.

The measurements used in calculations may contain varying numbers of significant figures, so carry as many as possible until the end and then round off the final answer. The least precise measurement will determine the significant figures reported in the final answer. Determine the number of significant figures in each measured value (not the exact ones) and then, depending on the mathematical operations involved, round off the final answer to the correct number of significant figures. Here are the rules for determining the number of significant figures in a measured value:

  1. All non-zero digits (1, 2, 3, 4, etc.) are significant.
  2. Zeroes between non-zero digits are significant.
  3. Zeroes to the left of the first non-zero digit are not significant.
  4. Zeroes to the right of the last non-zero digit are significant if there is a decimal point present, but not significant if there is no decimal point.

Rule 4 is a convention that many of us use, but some teachers or books may use alternative methods.

By these rules, 230,500. would contain 6 significant figures, but 230,500 would contain only 4.

Another way to determine the number of significant figures in a number is to express it in scientific (exponential) notation. The number of digits shown is the number of significant figures. For example, 2.305 × 10−5 would contain 4 significant figures. You may need to review exponential notation.

In determining the number of significant figures to be expressed in the final answer, the following rules apply:

  1. For addition and subtraction problems, the answer should be rounded off to the same number of decimal places as the measurement with the fewest decimal places.
  2. For multiplication and division problems, round off the answer to the same number of significant figures in the measurement with the fewest significant figures.

Remember : Carry as many numbers as possible throughout the calculation and only round off the final answer.

The use of an improper number of significant figures may lower your score on the AP exam.

Dimensional Analysis—the Factor Label Method

Dimensional analysis, sometimes called the factor label (unit conversion) method, is a method for setting up mathematical problems. Mathematical operations are conducted with the units associated with the numbers, and these units are canceled until only the unit of the desired answer is left. This results in a setup for the problem. Then the mathematical operations can efficiently be conducted and the final answer calculated and rounded off to the correct number of significant figures. For example, to determine the number of centimeters in 2.3 miles:

First, write down the initial data as a fraction:

Convert from miles to feet:

Convert from feet to inches:

Finally, convert from inches to centimeters:

The answer will be rounded off to 2 significant figures based upon the 2.3 miles, since all the other numbers are exact:

Sometimes on the AP exam, only setups will be given as possible answers. Write the correct setup to the problem and then see which one of the answers represents your answer.

Remember: The units must cancel!

Also: Make sure that the answer is legible and reasonable!

The States of Matter

Matter can exist in one of three states: solid, liquid, or gas. A solid has both a definite shape and a definite volume. At the molecular level, the particles that make up a solid are close together and many times are locked into a very regular framework called a crystal lattice. Molecular motion exists, but it is slight.

liquid has a definite volume but no definite shape. It conforms to the container in which it is placed. The particles are moving much more than in the solid. There are usually clumps of particles moving relatively freely among other clumps.

gas has neither definite shape nor volume. It expands to fill the container in which it is placed. The particles move rapidly with respect to each other and act basically independently of each other.

We will indicate the state of matter that a particular substance is in by a parenthetical s, l, or g. Thus, H2 O(s) would represent solid water (ice), while H2 O(g) would represent gaseous water (steam). For a more detailed discussion of solids, liquids, and gases see Chapters 8 and 12 .

The Structure of the Atom

Historical Development

The first modern atomic theory was developed by John Dalton and first presented in 1808. Dalton used the term atom (first used by Democritus) to describe the tiny, indivisible particles of an element. Dalton also thought that atoms of an element are the same and atoms of different elements are different. In 1897, J. J. Thompson discovered the existence of the first subatomic particle, the electron , by using magnetic and electric fields. In 1909, Robert Millikan measured the charge on the electron in his oil drop experiment (electron charge = −1.6022 × 10−19 coulombs), and from that he calculated the mass of the electron. Thompson developed an atomic model, the raisin pudding model, which described the atom as being a diffuse positively charged sphere with electrons scattered throughout.

Ernest Rutherford, in 1910, was investigating atomic structure by shooting positively charged alpha particles at a thin gold foil. Most of the particles passed through with no deflection, a few were slightly deflected, and every once in a while an alpha particle was deflected back towards the alpha source. Rutherford concluded from this scattering experiment that the atom was mostly empty space where the electrons were, and that there was a dense core of positive charge at the center of the atom that contained most of the atom’s mass. He called that dense core the nucleus .

Subatomic Particles

Our modern theory of the atom describes it as an electrically neutral sphere with a tiny nucleus at the center, which holds the positively charged protons and the neutral neutrons. The negatively charged electrons move around the nucleus in complex paths, all of which compose the electron cloud . Table 5.1 summarizes the properties of the three fundamental subatomic particles:

Table 5.1 The Three Fundamental Subatomic Particles

Many teachers and books omit the charges on the symbols for the proton and neutron.

The amu (atomic mass unit) is commonly used for the mass of subatomic particles and atoms. An amu is  the mass of a carbon-12 atom, which contains 6 protons and 6 neutrons (C-12).

Since the atom itself is neutral, the number of electrons must equal the number of protons. However, the number of neutrons in an atom may vary. Atoms of the same element (same number of protons) that have differing numbers of neutrons are called isotopes . A specific isotope of an element can be represented by the following symbolization:

X represents the element symbol taken from the periodic table. Z is the atomic number of the element, the number of protons in the nucleus. A is the mass number , the sum of the protons and neutrons. By subtracting the atomic number (p) from the mass number (p + n), the number of neutrons may be determined. For example,  (U-238) contains 92 protons, 92 electrons, and (238 − 92) 146 neutrons.

Electron Shells, Subshells, and Orbitals

According to the latest atomic model, the electrons in an atom are located in various energy levels or shells that are located at different distances from the nucleus. The lower the number of the shell, the closer to the nucleus the electrons are found. Within the shells, the electrons are grouped in subshells of slightly different energies. The number associated with the shell is equal to the number of subshells found at that energy level. For example, energy level 2 (shell 2) has two subshells. The subshells are denoted by the symbols s, p, d, f, etc. and correspond to differently shaped volumes of space in which the probability of finding the electrons is high. The electrons in a particular subshell may be distributed among volumes of space of equal energies called orbitals. There is one orbital for an s subshell, three for a p, five for a d, seven for an f, etc. Only two electrons may occupy an orbital. Table 5.2summarizes the shells, subshells, and orbitals in an atom. Chapter 10 on Spectroscopy, Light, and Electrons has a discussion of the origin of this system.

Table 5.2 Summary of Atomic Shell, Subshells, and Orbitals for Shells 1–4

Energy-Level Diagrams

The information above can be shown in graph form as an energy-level diagram, as shown in Figure 5.1 :

Figure 5.1   Energy-level diagram of an atom.

Be sure to fill the lowest energy levels first (Aufbau principle ) when using the diagram in Figure 5.1 . In filling orbitals having equal energy, electrons are added to the orbitals to half fill them all before any pairing occurs (Hund’s rule ). Sometimes it is difficult to remember the relative energy position of the orbitals. Notice that the 4s fills before the 3d. Figure 5.2 may help you remember the pattern in filling. Study the pattern and be able to reproduce it during the exam.

Figure 5.2   Orbital filling pattern.

Following these rules, the energy-level diagram for silicon (Z = 14) can be written as shown in Figure 5.3 .

Figure 5.3   Energy-level diagram for silicon.

Although this filling pattern conveys a lot of information, it is bulky. A shorthand method for giving the same information has been developed—the electronic configuration.

Electronic Configurations

The electronic configuration is a condensed way of representing the pattern of electrons in an atom. Using the Aufbau build-up pattern that was used in writing the energy-level diagram, consecutively write the number of the shell (energy level), the type of orbital (s, p, d, etc.), and then the number of electrons in that orbital shown as a superscript. For example, 1s2 2s1 would indicate that there are two electrons in the s-orbital in energy level (shell) l, and one electron in the s-orbital in energy level 2. Looking at the energy-level diagram for silicon in Figure 5.3 , the electronic configuration would be written as:

silicon: 1s2 2s2 2p2 3s2 3p2

The sum of all the superscripts should be equal to the number of electrons in the atom (the atomic number, Z). Electronic configurations can also be written for cations and anions.

Periodic Table

If chemistry students had to learn the individual properties of the 100+ elements that are now known, it would be a monumental and frustrating task. Early scientists had to do just that. Then several scientists began to notice trends in the properties of the elements and began grouping them in various ways. In 1871, a Russian chemist, Dmitri Mendeleev, introduced the first modern periodic table. He arranged the elements in terms of increasing atomic mass. He then arranged columns so that elements that had similar properties were in the same column. Mendeleev was able to predict the existence and properties of elements that were then unknown. Later, when they were discovered, Mendeleev’s predictions were remarkably accurate. Later the periodic table was rearranged to sequence the elements by increasing atomic number, not mass. The result is the modern periodic table shown in Figure 5.4 .

This is not the periodic table supplied on the AP exam. The one in this book has family and period labels. Become familiar with these labels so that you can effectively use the unlabeled one. You may wish to add labels to the one supplied with the AP exam.

Each square on this table represents a different element and contains three bits of information. The first is the element symbol. You should become familiar with the symbols of the commonly used elements. Second, the square lists the atomic number of the element, usually centered above the element. This integer represents the number of protons in the element’s nucleus. The atomic number will always be a whole number. Third, the square lists the element’s mass, normally centered underneath the element symbol. This number is not a whole number, because it is the weighted average (taking into consideration abundance) of all the masses of the naturally occurring isotopes of that element. The mass number can never be less than the atomic number.

Arrangement of Elements

There are a number of different groupings of elements on the periodic table that may be utilized. One system involves putting the elements into three main groups—metals, nonmetals, and metalloids (semimetals). Look at Figure 5.4. Notice the heavy, stair-stepped line starting at boron (B) and going downward and to the right. The elements to the left of that line (except for H, Ge, and Sb) are classified as metals. Metals are normally solids (mercury being an exception), shiny, and good conductors of heat and electricity. They can be hammered into thin sheets (malleable) and extruded into wires (ductile). Chemically, metals tend to lose electrons in reactions, to form cations.

Figure 5.4   The periodic table.

Elements bordering the stair-stepped line (B, Si, Ge, As, Sb, Te) are classified as metalloids. Metalloids have properties of both metals and nonmetals. Their unusual electrical properties make them valuable in the semiconductor and computer industry.

The rest of the elements, to the right of the metalloids, are called nonmetals. Nonmetals have properties that are often the opposite of metals. Some are gases, are poor conductors of heat and electricity, are neither malleable nor ductile, and tend to gain electrons in their chemical reactions to form anions.

Another way to group the elements on the periodic table is in terms of periods and groups (families). Periods are the horizontal rows, which have consecutive atomic numbers. The periods are numbered from 1 to 7. Elements in the same period do not have similar properties in terms of reactions.

The vertical rows on the periodic table are called groups or families . They may be labeled in one of two ways. An older and still widely used system is to label each group with a Roman numeral and a letter, A or B. The groups that are labeled with an A are called the main-group elements, while the B groups are called the transition elements . Two other horizontal groups, the inner transition elements , have been pulled out of the main body of the periodic table. The Roman numeral at the top of the main-group families indicates the number of valence (outermost shell) electrons in that element. Valence electrons are normally considered to be only the s and p electrons in the outermost energy level. The transition elements (B groups) are filling d-orbitals, while the inner transition elements are filling f-orbitals.

Four main-group families are given special names, which you should remember:

Another way to label the groups is to consecutively number the groups from left to right, 1–18. This method is newer than the other labeling method, and it has not gained wide use. Most teachers and chemists still prefer and use the older method.

Trends in Periodic Properties

Trends are useful on the multiple-choice portion of the AP exam, but simply stating a trend will not be sufficient on the free-response portion of the exam. You must give the reason behind the trend. For example, “higher on the periodic table” is a trend, but not a reason.

The overall attraction an electron experiences is due to the effective nuclear charge . This attraction is related to the positive nuclear charge interacting with the negative electrons. Electrons between the nucleus and the electron under consideration interfere with, or shield, that electron from the full nuclear charge. This shielding lessens the nuclear charge. Within a period, the shielding is nearly constant; however, the effective nuclear charge will increase with an increasing number of protons (atomic number). Within the same family or group, as the atomic number increases so does the shielding, resulting in a relatively constant effective nuclear charge.

The size of an atom is generally determined by the number of energy levels occupied by electrons. This means that as we move from top to bottom within a group, the size of the atom increases due to the increased number of shells containing electrons. As we move from left to right within a period (within the same valence shell), the atomic size decreases somewhat owing to the increased effective nuclear charge for the electrons. This increased attraction is related to the increasing number of protons within the nucleus. The size of a cation is smaller than the neutral atom, because in many cases an entire energy shell has been removed, while an anion is larger than the corresponding neutral atom since the nuclear attraction is being distributed over additional electrons. As the number of electrons changes so will the electron–electron repulsion. The greater the electron–electron repulsion, the larger the species becomes, and vice versa.

The ionization energy (IE) is the energy needed to completely remove an electron from an atom. It may be expressed in terms of 1 atom or a mole of atoms. Energy is required in this process in order to overcome the attraction of the nucleus for the electrons. There are two factors affecting the magnitude of the ionization energy. One is the size of the atom. The closer the electrons are to the nucleus, the more energy is needed to overcome the effective nuclear charge.

Therefore, ionization energy tends to decrease from top to bottom within a group, since the valence electrons (the first ones to be lost) are farther away from the nucleus.

The other factor is the magnitude of the effective nuclear charge. The greater the effective nuclear charge, the more energy is required to remove the electron. Since the effective nuclear charge increases from left to right within a period, the ionization energies will also increase from left to right. The increased effective nuclear charge results in the atom becoming slightly smaller, which also leads to a greater nuclear attraction for the electrons.

The ionization energy for the removal of a second electron is greater in all cases than the first, because the electron is being pulled away from a positively charged ion and the attraction is greater than from a neutral atom.

The electron affinity (EA) is the energy change that results from adding an electron to an atom or ion. The trends in electron affinity are not quite as regular as size or ionization energy. In general, electron affinity increases from left to right within a period (owing to the increased effective nuclear charge), and decreases from top to bottom within a group owing to increased atomic or ionic size. Noble gases are an exception—they have no EA.

Do not forget that the trends mentioned in this section may help you on the multiple-choice portion of the AP exam. However, it is the underlying reasons that you need for the free-response portion.

Oxidation Numbers

Oxidation numbers are bookkeeping numbers that allow chemists to do things like balance redox equations. Don’t confuse oxidation numbers with the charge on an ion. Oxidation numbers are assigned to elements in their natural state or in compounds using the following rules:

⊠   The oxidation number of an element in its elemental form (i.e., H2 , Au, Ag, N2 ) is zero.

⊠   The oxidation number of a monoatomic ion is equal to the charge on the ion. The oxidation number of Mg2+ is +2. Note that the charge is written with number first, then sign; for oxidation numbers it is sign, then number.

⊠   The sum of all the oxidation numbers of all the elements in a neutral molecule is zero. The sum of all the oxidation numbers in a polyatomic ion is equal to the charge on the ion.

⊠   The alkali metal ions have an oxidation number of +1 in all their compounds.

⊠   The alkaline earth metals have an oxidation number of +2 in all their compounds.

⊠   The oxidation number of hydrogen in compounds is +1, except it is −1 when combined with metals or boron in binary compounds.

⊠   The oxidation number of halogens in their compounds is −1 except when combined with another halogen above them on the periodic table, or with oxygen.

⊠   The oxidation number of oxygen is −2 in compounds, except for peroxides, in which it is −1.

Determine the oxidation number of sulfur in sulfuric acid, H2 SO4 . The sum of all the oxidation numbers must equal zero, since this is a neutral compound. The oxidation numbers of hydrogen (+1) and oxygen (−2) are known, so the oxidation number of sulfur can be determined:

2(+1) + ? + 4(–2) = 0
H2 SO4

The oxidation number of sulfur in this compound must be +6.

Nomenclature Overview

This overview covers some of the rules for naming simple inorganic compounds. There are additional rules, and some exceptions to these rules. The first part of this overview discusses the rules for deriving a name from a chemical formula. In many cases, the formula may be determined from the name by reversing this process. The second part examines situations in which additional information is needed to generate a formula from the name of a compound. The transition metals present some additional problems; therefore, there is a section covering transition metal nomenclature and coordination compounds.

Binary Compounds

Binary compounds are compounds that consist of only two elements. Some binary compounds have special names, and these special names supersede any of the rules given below. H2 O is water, NH3 is ammonia, and CH4 is methane. All other binary compounds have a name with a suffix ide . Binary compounds may be subdivided into metal type, nonmetal type, and acid type.

(a) Metal type These binary compounds begin with metals. The metal is given first in the formula. In general, metals are the elements on the left-hand side of the periodic table, and the nonmetals are on the right-hand side. Hydrogen, a nonmetal, is an exception to this generalization.

First name the metal, then name the nonmetal with the suffix ide. Examples:

The ammonium ion  is often treated as a metal, and its compounds are named under this rule. Thus, NH4 Cl is named ammonium chloride.

(b) Nonmetal type These binary compounds have formulas that begin with a nonmetal. Prefixes are used to indicate the number of each atom present. No prefixes are used for hydrogen. Naming the compounds can best be explained using the following examples:

Carbon monoxide is one of the very few cases where the prefix mono is used. In general, you should not use mono in any other compound.

Some of the prefixes used to denote the numbers of atoms in a compound are listed below:

On many occasions the terminal a or o is dropped for oxides, so they read as pentoxide, heptoxide, or monoxide.

In normal nomenclature, the nonmetal prefixes are not used if a metal is present. One of the few exceptions to this is MnO2 , sometimes called manganese dioxide.

(c) Acid type These binary compounds have formulas that begin with hydrogen. If the compound is not in solution, the naming is similar to that of the metal type. If the compound is dissolved in H2 O, indicated by (aq ), the compound takes on the prefix hydro and the suffix ic . If the compound is not in solution, the state of matter should be shown as follows:

HCl(g), HF(l)

If the formula has no designation of phase or water, either name may be used. Examples for naming these compounds are:

HCN (hydrocyanic acid) is named using these rules. However, in this case, it does not matter if the phase or water is indicated.

Ternary Compounds

Ternary compounds are those containing three or more elements. If the first element in the formula is hydrogen, it is usually classified as an acid. If the formula contains oxygen in addition to the hydrogen, the compound is usually classified as an oxyacid. In general, if the first element in the formula is not hydrogen, the compound is classified as a salt.

Ternary acids are usually named with the suffixes ic or ous . The exceptions are the acids derived from ions with an ide suffix (see HCN in the preceding section). These acids undergo many reactions to form salts, compounds of a metal, and the ion of an acid. The ions from the acids H2 SO4 and HNO3 are SO4 2− , NO3  . If an acid name has the suffix ic , the ion of this acid has a name with the suffix ate . If an acid name has the suffix ous , the ion has a name with the suffix ite . Salts have the same suffixes as the suffixes of the ions. The difference between the acid with a suffix ic and the acid with the suffix ous can many times be determined by visual inspection of the formula. The acid with the suffix ous usually has one fewer oxygen atom than the acid with the suffix ic . Examples:

When the ternary compound is not an acid, the first element is usually a metal. In these cases, the name of the compound is simply the name of the metal followed by the name of the ion. The ammonium ion is treated as a metal in these cases.

The following are examples:

Writing Formulas

To write the formula from the name of a binary compound containing only nonmetals, simply write the symbols for the separate atoms with the prefixes converted to subscripts.

In all compounds, the total charge must be zero. There are NO exceptions. Thus, to determine the formula in those cases where no prefixes are given, it is necessary to have some idea what the individual charges are. The species with the positive charge is listed and named first; this is followed by the species with the negative charge. Subscripts may be needed to make sure the sum of the charges (valances) will equal zero. Examples:

  1. Magnesium oxide

Mg2+ O2– = +2 – 2 = 0

This gives MgO.

  1. Sodium oxide

This gives Na2 O.

  1. Aluminum oxide

This gives Al2 O3 .

If a polyatomic ion must be increased to achieve zero charge, parentheses should be used. An example of this is shown as:

This gives (NH4 )2 SO4 .

One way of predicting the values of the subscripts is to crisscross the valences. This is not a rule of nomenclature, but for practice purposes in this exercise it will be referred to as the crisscross rule. It works most of the time and therefore is worth considering. Example:

If the crisscross rule is applied, you should reduce the formula if possible. For example:

Mn4+ O2− crisscrosses to Mn2 O4 , which reduces to MnO2

If a formula is given, the crisscross rule can be reversed to give the valences:

As a first approximation, the valences of the representative elements can be predicted from their position on the periodic table. Hydrogen and the metals have positive charges beginning with +1 on the left and increasing by one as you proceed to the right on the periodic table (skipping the transition metals). Nonmetals begin with 0 in the rightmost column of the periodic table and decrease by 1 as you move to the left on the periodic table. Metalloids may be treated as metals or nonmetals. Examples are:

Na+ Al3+ Pb4+ N3− Se2− I 
Na+ Mg2+ Al3+ Si4+ P3– S2– Cl Ar0

Transition Metals

Many transition metals and the group of six elements centered around lead on the periodic table commonly have more than one valence. The valence of these metals in a compound must be known before the compound can be named. Modern nomenclature rules indicate the valence of one of these metals with a Roman numeral suffix (Stock notation). Older nomenclature rules used different suffixes to indicate the charge. Examples:

  1. FeCl3
    Fe3+Cl3 1− (crisscross rule)

The compound is named iron(III) chloride or ferric chloride.

  1. FeCl2

If chloride is −1, two chloride ions are −2. Fe has a valence of +2, to give a total charge of zero. The name is iron(II) chloride or ferrous chloride.

  1. MnO2

Mn4+ (found previously)
The name would be manganese(IV) oxide, although it is often named manganese dioxide.

The Roman numeral suffix is part of the name of the metal. Thus iron(III) is one word.

Stock notation should be used for all metals that have a variable valence. This includes almost all the transition elements and the elements immediately around lead on the periodic table. Stock notation is often omitted for Zn, Cd, and Ag, as they do not have variable valences.

The valences of some common metals and acids are listed in the appendixes.

Coordination Compounds

Coordination compounds contain a complex. In general, a complex may be recognized because it is enclosed in square brackets [ ]. The square brackets are omitted when the actual structure of the complex is uncertain.

A complex is composed of a central atom, normally a metal, surrounded by atoms or groups of atoms called ligands. One way of forming a complex is illustrated below:

Ni2+ + 6 H2 O → [Ni(H2 O)6 ]2+

In this reaction, the metal behaves as a Lewis acid and accepts a pair of electrons from the Lewis base (ligand). In this case, the ligand is water, with the oxygen atom donating one of its lone pairs to the nickel. The oxygen atom is called the donor atom. In this complex, there are six donor atoms.

A complex may be ionic or neutral. An ionic complex is called a complex ion. A neutral complex is a type of coordination compound. The only difference in naming coordination compounds or complex ions is that anionic complex ions have an ate suffix.

A coordination compound may contain more than one complex ion or material that is not part of the complex, but it must have an overall neutral charge. Examples of coordination compounds are: [Pt(NH3 )2 Cl2 ], K2 [Mn(C2 O4)3 ], and [Ni(H2 O)6 ]SO4 .

When writing formulas, the metal (central atom) is always listed first within the brackets. However, when writing names, the metal name is always given last. Any material not listed within the brackets is named separately.


If everything in the formula is enclosed within one set of brackets, the entire name will be one word. If there is material outside the brackets, this outside material is named separately.

Just as with simpler compounds, cations are always named before anions. Thus, a cationic complex would be the first word in the name, and an anionic complex would be the last word in a name (with an ate ending).


When naming a complex, or when writing the formula for a complex, the ligands are listed alphabetically. Again, do not forget that metals are first in the formula and last in the name.

The names of anionic ligands always end in an o . Neutral ligands are basically unchanged. Two common exceptions in the case of neutral ligands are NH3 = ammine (note the double m ), and H2 O = aqua. Other common ligands and their names are listed in the appendixes.

Multiple identical ligands have prefixes added to designate the number of such ligands:


If the ligand name contains a prefix or begins with a vowel (except ammine and aqua), alternate prefixes should be used:

When using the alternate prefixes, it is common practice to enclose the name of the ligand within parentheses. Either type of prefix is added after the ligands have been alphabetized.


Anionic complexes always have names ending in ate . This will require a change in the name of the metal. Thus, aluminum would become aluminate, and zinc would become zincate. The only exceptions to this are some of the metals whose symbols are based on Latin or Greek names. These exceptions are:


If the metal ion may exist in more than one oxidation state, this oxidation state should be listed, in Roman numerals, immediately after the name of the metal ion. The Roman numeral is enclosed in parentheses and is considered part of the same word, and not a separate grouping. If the metal occurs in only one oxidation state, no such indicator is used. This notation is the Stock system discussed earlier.


Experiments involving the basic material covered in this chapter have been placed in the in-depth chapters throughout the remainder of this book.

Common Mistakes to Avoid

Between the two of us, we have almost 60 years of teaching experience. We’ve seen a lot of student mistakes. We will try to steer you clear of the most common ones.

  1. Always show your units in mathematical problems.
  2. In the conversion from °F to °C, be sure to subtract 32 from the Fahrenheit temperature first, then multiply by 5/9.
  3. In the conversion from °C to °F, be sure to multiply the Celsius temperature by 9/5, then add 32.
  4. No degree sign is used for Kelvin.
  5. Only consider measured values for significant figures.
  6. When considering whether or not zeroes to the right of the last non-zero digit are significant, pay attention to whether or not there is a decimal point.
  7. Round off only your final answer, not intermediate calculations.
  8. In working problems, be sure that your units cancel.
  9. If you are solving for cm, for example, be sure you end up with cm and not 1/cm.
  10. Make sure your answer is a reasonable one.
  11. Don’t confuse the mass number (A) with the atomic number (Z ).
  12. When determining valence electrons, only the s and p electrons are considered.
  13. Don’t put more than 2 electrons in any individual orbital.
  14. Always fill lowest energy levels first.
  15. Half fill orbitals of equal energy before pairing up the electrons.
  16. In writing the electronic configuration of an atom, make sure you use the correct filling order.
  17. Don’t confuse the periods with the groups on the periodic table.
  18. Don’t confuse ionization energy with electron affinity.
  19. Don’t confuse oxidation numbers with ionic charge.
  20. In naming compounds, don’t confuse metal and nonmetal type binary compounds. Prefixes are used only with nonmetal types.
  21. Be careful when using the crisscross rule to reduce the subscripts to their lowest whole-number ratio.
  22. Be sure to report the proper number of significant figures.
  23. Simply knowing a periodic trend will allow you to pick the correct multiple-choice answer, but be prepared to explain the trend in free-response questions.

 Review Questions

Here are questions you can use to review the content of this chapter and practice for the AP Chemistry exam. First are 25 multiple-choice questions similar to what you will encounter in Section I of the AP Chemistry exam. Following those is a four-part free-response question like the ones in Section II of the exam. To make these questions an even more authentic practice for the actual exam, time yourself following the instructions provided.

Multiple-Choice Questions

Answer the following questions in 30 minutes. You may not use a calculator. You may use the periodic table and the equation sheet at the back of this book.

1 . In most of its compounds, this element exists as a monatomic cation.

(A) O

(B) Cl

(C) Na

(D) N

2 . This element may form a compound with the formula CaXO4 .

(A) Se

(B) Cl

(C) P

(D) Na

3 . Which of the following elements may occur in the greatest number of different oxidation states?

(A) C

(B) F

(C) O

(D) Ca

4 . Choose the group that does NOT contain isotopes of the same element.

5 . Which of the following groups has the species correctly listed in order of increasing radius?

(A) Mg2+ , Ca2+ , Ba2+

(B) K+ , Na+ , Li+

(C) Br , Cl , F

(D) Na, Mg, Al

6 . Which of the following elements has the lowest electronegativity?

(A) C

(B) K

(C) Al

(D) I

7 . Choose the ion with the largest ionic radius.

(A) F

(B) Al3+

(C) K+

(D) I

8 . What is the name of the energy change when a gaseous atom, in the ground state, adds an electron?

(A) ionization energy

(B) sublimation energy

(C) atomization energy

(D) electron affinity

9 . The following ionization energies are reported for element X . (All the values are in kJ/mol.)

Based on the above information, the most likely identity of X is:

(A) Mg

(B) Cl

(C) Al

(D) Na

10 . In general, as the atomic numbers increase within a period, the atomic radius:

(A) decreases

(B) increases

(C) first decreases and then increases

(D) does not change

11 . Which of the following elements is a reactive gas?

(A) chlorine

(B) gold

(C) sodium

(D) radon

12 . Which of the following elements is an unreactive metal?

(A) chlorine

(B) gold

(C) sodium

(D) radon

13 . Which of the following represents the correct formula for potassium trisoxalatoferrate(III)?

(A) P3 [Fe(C2 O4 )3 ]

(B) K3 [Fe(C2 O4 )3 ]

(C) KFe3 (C2 O4 )3

(D) K3 [Fe3 (C2 O4 )3 ]

14 . Which of the following substances will produce a colorless aqueous solution?

(A) Zn(NO3 )2

(B) CuSO4

(C) K2 Cr2 O7

(D) Co(NO3 )2

15 . This element is a liquid at room temperature.

(A) Hg

(B) Th

(C) Na

(D) Cl

16 . Which of the following elements is present in chlorophyll?

(A) K

(B) Ga

(C) Al

(D) Mg

17 . What is the symbol for the element that forms a protective oxide coating?

(A) K

(B) Ga

(C) Al

(D) Mg

18 . Which of the following elements is important in the semiconductor industry to improve the conductivity of germanium, Ge?

(A) K

(B) Ga

(C) Al

(D) Mg

19 . Which of the following aqueous solutions is blue?

(A) CuSO4

(B) Cr2 (SO4 )3

(C) NiSO4

(D) ZnSO4

20 . In order to separate two substances by fractional crystallization, the two substances must differ in which of the following?

(A) solubility

(B) specific gravity

(C) vapor pressure

(D) viscosity

21 . In a flame test, copper compounds impart which of the following colors to a flame?

(A) red

(B) orange

(C) blue to green

(D) violet

22 . What should you do if you spill sulfuric acid on the countertop?

(A) Neutralize the acid with vinegar.

(B) Sprinkle solid NaOH on the spill.

(C) Neutralize the acid with NaHCO3 solution.

(D) Neutralize the acid with an Epsom salt (MgSO4 ) solution.

23 . Which of the following can be achieved by using a visible-light spectrophotometer?

(A) Run a flame test to determine if Na+ or K+ is in a solution.

(B) Find the concentration of a KMnO4 solution.

(C) Detect the presence of isolated double bonds.

(D) Measure the strength of a covalent bond.

24 . You have an aqueous solution of NaCl. The simplest method for the separation of NaCl from the solution is:

(A) evaporation of the solution to dryness

(B) centrifuging the solution

(C) filtration of the solution

(D) electrolysis of the solution

25 . The determination that atoms have small, dense nuclei is attributed to:

(A) Rutherford

(B) Becquerel

(C) Einstein

(D) Dalton

Answers and Explanations for the Multiple-Choice Questions

1 . C —All the other elements are nonmetals. Nonmetals usually form monatomic anions.

2 . A —The element cannot be a metal (Na). A nonmetal that can have a +6 oxidation state is necessary. P has a maximum of +5. Cl may be +5 or +7. Se, in column 16, can easily be +6.

3 . A —Based on their positions on the periodic table:

4 . D —Isotopes MUST have the same number of protons. Different isotopes of an element have different numbers of neutrons.

5 . A —All the others are in decreasing order. Ions in the same column and with the same charge increase in size when going down a column the same as atoms. Atoms in the same row increase in size toward the left side. This argument is not sufficient on the free-response portion of the exam.

6 . B —In general, the element farthest from F on the periodic table will have the lowest electronegativity. There are exceptions, but you normally do not need to concern yourself with exceptions.

7 . D —The very large iodine atom (near the bottom of the periodic table) gains an electron to make it even larger. This reasoning is not sufficient on the free-response portion of the exam.

8 . D —The definition of electron affinity is the energy change when a ground-state gaseous atom adds an electron.

9 . D —The more electrons removed, the higher the values should be. The large increase between the first and second ionization energies indicates a change in electron shell. The element, X , has only 1 valence electron. This is true for Na. For the other elements the numbers of valence electrons are: Mg – 2; Cl – 7; and Al – 3.

10 . A —The increase in the number of protons in the nucleus has a greater attraction (greater effective nuclear charge) for the electrons being added in the same energy level. Thus, the electrons are pulled closer to the nucleus and the size slightly decreases. This thought process should be used on the free-response portion of the AP exam; however, simply remembering that radii decrease across a period is sufficient for most multiple-choice questions.

11 . A —The only other gas is radon, and it is inert.

12 . B —Sodium is a metal on the left side of the periodic table. Metals on the left side of the periodic table are very reactive. Radon is not a metal.

13 . B —Ferrate(III) means Fe3+ , while trisoxalato means (C2 O4 )3 6– ; three potassium atoms are needed to balance the charge.

14 . A —B is blue; C is orange; and D is pink to red.

15 . A —Chlorine is a gas; all the others are solid metals.

16 . D —Magnesium is present in chlorophyll.

17 . C —Aluminum forms a protective oxide coating.

18 . B —Gallium, adjacent to Ge on the periodic table, is one of the elements that will improve the conductivity of germanium.

19 . A —B is purple; C is green; and D is colorless.

20 . A —Fractional crystallization works because the less soluble material separates first.

21 . C —A could be Li or Sr; B is Ca; and D is K.

22 . C —Adding a weak base solution, such as NaHCO3 , which will not only neutralize the acid but will help to disperse the heat, is the best choice.

23 . B —A solution containing a colored substance is necessary.

24 . A —Separation of materials in solution is normally not simple; therefore, removal of the solvent through evaporation is the best choice.

25 . A —Rutherford, and his students, demonstrated the existence of the nucleus.

Free-Response Question

Both authors have been AP free-response graders for years. Here is a free-response question for practice.

You have 10 minutes to do the following question. You may use a calculator and the tables in the back of the book.

Question 1

Use the periodic table and other information concerning bonding and electronic structure to explain the following observations.

(a) The radii of the iron cations are less than that of an iron atom, and Fe3+ is smaller than Fe2+ .

(b) When moving across the periodic table from Li to Be to B, the first ionization energy increases from Li to Be, then drops for B. The first ionization energy of B is greater than that of Li.

(c) The electron affinity of F is higher than the electron affinity of O.

(d) The following observations have been made about the lattice energy and ionic radii of the compounds listed below. Compare NaF to CaO, and then compare CaO to BaO. All of the solids adopt the same crystal structure.

Answer and Explanation for the Free-Response Question

Notice that all the answers are very short. Do not try to fill all the space provided on the exam. You score points by saying specific things, not by the bulk of material. The graders look for certain keywords or phrases. The answers should not contain statements that contradict each other; otherwise, there may be a penalty. Contradictions most commonly occur when the student tries to say too much. On the AP exam, the different parts of the free-response questions tend to be more diverse than this one, as this question focuses on this chapter, whereas the AP free-response questions focus on the entire course.

(a) The observed trend of radii is Fe > Fe2+ > Fe3+ . There is an increase in the effective nuclear charge in this series. As electrons are removed, the repulsion between the remaining electrons decreases. The larger the effective nuclear charge, the greater the attraction of the electrons toward the nucleus and the smaller the atom or ion becomes.

Give yourself 1 point for “effective nuclear charge” and 1 point for the effective nuclear charge, and give yourself 1 point for the remainder of the discussion.

(b) When moving across a period on the periodic table, the value of the effective nuclear charge increases with atomic number. This causes a general increase from Li to Be to B. DO NOT use the argument that ionization energies increase to the right on the periodic table, unless you also discuss effective nuclear charge.

The even higher value of Be (greater than B) is due to the increased stability of the electron configuration of Be. Beryllium has a filled s-subshell. Filled subshells have an increased stability, and additional energy is required to pull an electron away.

This effective nuclear charge argument is worth 1 point. Give yourself 0 points if you say that the ionization energy increases to the right on the periodic table. This is an observation; it is not an explanation. Give yourself 1 more point for the filled subshell discussion.

(c) The effective nuclear charge in F is greater than the effective nuclear charge in O. This causes a greater attraction of the electrons. DO NOT use the argument that electron affinity increases to the right on the periodic table, unless you also discuss effective nuclear charge.

You get 1 point for this answer.

(d) Because all these solids adopt the same structure, the structure is irrelevant. The sizes of the anions are similar; thus, anion size arguments are not important. Two factors, other than structure and anion size, are important here. The two compounds with the highest lattice energies contain divalent ions (+2 or –2), while NaF contains univalent ions (+1 or –1). The higher the charge is, the greater the attraction between the ions is. The lattice energy increases as the attraction increases.

You get 1 point for correctly discussing the charges. The difference between the CaO and BaO values is because the larger the ion is, the lower the attraction is (greater separation). The lower attraction leads to a lower lattice energy. This size argument will get you 1 point.

Total your points. The maximum is 7.

 Rapid Review

Here is a brief review of the most important points in the chapter. If something sounds unfamiliar, study it in the chapter and your textbook.

  • Know the metric measurement system and some metric/English conversions.
  • Know how to convert from any one of the Fahrenheit/Celsius/Kelvin temperature scales to the other two.
  • The density of a substance is mass per unit volume.
  • Know how to determine the number of significant figures in a number, the rules for how many significant figures are to be shown in the final answer, and the round-off rules.
  • Know how to set up problems using the factor label method.
  • Know the differences between a solid, a liquid, and a gas at both the macroscopic and microscopic levels.
  • Know what part Dalton, Thompson, Millikan, and Rutherford had in the development of the atomic model.
  • Know the three basic subatomic particles—proton, neutron, and electron—their symbols, mass in amu, and their location.
  • Isotopes are atoms of the same element that have differing numbers of neutrons.
  • Electrons are located in major energy levels called shells. Shells are divided into subshells, and there are orbitals for each subshell.
  • Know the electron capacity of each orbital (always 2).
  • Be able to write both the energy-level diagram and the electronic configuration of an atom or ion by applying both the Aufbau build-up principle and Hund’s rule.
  • Know how the modern periodic table was developed, including the differences between Mendeleev’s table and the current table.
  • Periods are the horizontal rows on the periodic table; the elements have properties unlike the other members of the period.
  • Groups or families are the vertical rows on the periodic table; the elements have similar properties.
  • Know the properties of metals, nonmetals, and metalloids and which elements on the periodic table belong to each group.
  • Valence electrons are outer-shell electrons.
  • The IA family is known as the alkali metals; the IIAs are the alkaline earth metals; the VIIAs are the halogens; and the VIIIAs are the noble gases.
  • Know why atoms get larger as we go from top to bottom in a group and slightly smaller as we move from left to right on the periodic table. Remember that on the free-response section, simply quoting a trend is not sufficient in answering the question. This is true for all trends.
  • Ionization energy is the energy it takes to remove an electron from a gaseous atom or ion. It decreases from top to bottom and increases from left to right on the periodic table. Much the same trend is noted for electron affinity, the energy change that takes place when an electron is added to a gaseous atom or ion. The trends depend on the size of the atom or ion and its effective nuclear charge.
  • Oxidation numbers are bookkeeping numbers. Know the rules for assigning oxidation numbers.
  • Be able to name binary metal type and nonmetal type compounds, as well as ternary compounds, oxyacids, simple coordination compounds, etc.