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CHAPTER 2

Atoms, Molecules, and Ions

CHAPTER TERMS AND DEFINITIONSNumbers in parentheses after definitions give the text sections in which the terms are explained. Starred terms are italicized in the text. Where a term does not fall directly under a text section heading, additional information is given for you to locate it.

atomic theory explanation of the structure of matter in terms of different combinations of very small particles (2.1)

atom minute particle of which matter is composed; the smallest part of an element that can enter into chemical reaction (2.1)

element substance whose atoms all have the same atomic number (2.1, 2.3)

compound type of matter composed of atoms of two or more elements chemically combined in fixed proportions (2.1)

chemical reaction rearrangement of atoms present in reacting substances to give new chemical combinations present in the substances formed (2.1)

atomic symbol one- or two-letter notation used to represent an atom corresponding to a particular element (2.1)

law of multiple proportions when two elements form more than one compound, the masses of one element in these compounds for a fixed mass of the other element are in ratios of small whole numbers (2.1)

nucleus positively charged central core of an atom; contains most of the atom’s mass (2.2)

electron very light, negatively charged particle that exists in the region around the positively charged nucleus (2.2)

cathode* negative electrode (2.2)

anode* positive electrode (2.2)

cathode rays* rays that originate from the cathode, or negative electrode, in a gas-discharge tube (2.2)

coulomb (C)* unit of electric charge (2.2)

nuclear model* most of the mass of an atom is concentrated in a positively charged center, called the nucleus, around which negatively charged electrons move (2.2)

proton nuclear particle having a positive charge equal to +e (e being the charge on an electron) and a mass more than 1800 times that of an electron (2.3)

atomic number (Z) number of protons in an atomic nucleus; identifies the element (2.3)

neutron neutral particle of mass almost identical to that of a proton but without electric charge (2.3)

mass number (A) total number of protons and neutrons in a nucleus (2.3)

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22 Chapter 2: Atoms, Molecules, and Ions

nuclide an atom characterized by a definite atomic number and mass number (2.3)

nuclide symbol* symbol for a nuclide in which the mass number is written as a superscript and the atomic number as a subscript on the left of the symbol for the element (2.3)

isotopes atoms whose nuclei have the same atomic number but different mass numbers (2.3)

mass spectrometer* instrument used to determine atomic mass (2.4)

atomic mass average atomic mass for the naturally occurring element, expressed in atomic mass units (2.4)

atomic mass unit (amu) mass unit equal to exactly one-twelfth the mass of a carbon-12 atom (2.4)

mass spectrum* chart recording from the mass spectrometer that shows the relative numbers of atoms for various masses (2.4)

fractional abundance fraction of the total number of atoms that is composed of a particular isotope (2.4)

periodic table tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements (2.5)

period (of periodic table) elements in any one horizontal row of the periodic table (2.5)

group (of periodic table) elements in any one column of the periodic table (2.5)

main-group (representative) elements* elements in the A groups of the periodic table (2.5)

transition elements* elements in the B groups of the periodic table (2.5)

inner-transition elements* two rows of elements at the bottom of the periodic table (2.5)

lanthanides* first of the two rows of inner-transition elements (2.5)

actinides* second of the two rows of inner-transition elements (2.5)

alkali metals* elements in Group IA of the periodic table (2.5)

halogens* elements in Group VIIA of the periodic table (2.5)

metal substance or mixture that has a characteristic luster or shine and is generally a good conductor of heat and electricity; elemental metals are to the left of the staircase line on the periodic table (2.5)

malleable* able to be hammered into sheets (2.5)

ductile* able to be drawn into wire (2.5)

nonmetal element to the right of the staircase line on the periodic table; exhibits characteristics different from those of metals (2.5)

metalloid (semimetal) element bordering the staircase line on the periodic table; exhibits both metallic and nonmetallic properties (2.5)

semiconductors* elements that, when pure, are poor conductors of electricity at room temperature but become good conductors at higher temperatures (2.5)

doping* adding small amounts of other elements to pure semiconductor elements to make them very good electrical conductors (2.5, margin note)

chemical formula notation that uses atomic symbols with numerical subscripts to convey the relative proportions of atoms of the different elements in the substance (2.6)


Chapter 2: Atoms, Molecules, and Ions 23

molecule definite group of atoms that are chemically bonded together and, as a group, electrically neutral (2.6)

molecular substance* substance composed of molecules all of which are alike (2.6)

molecular formula gives the exact number of different atoms of an element in a molecule (2.6)

structural formula* chemical formula that shows which atoms are bonded to one another in a molecule (2.6)

polymers very large molecules that are made up of a number of smaller molecules repeatedly linked together (2.6)

monomers the small molecules that are linked together to form a polymer (2.6)

ion electrically charged particle obtained from an atom or chemically bonded group of atoms by addition or removal of one or more electrons (2.6)

anion negatively charged ion (2.6)

cation positively charged ion (2.6)

ionic compound compound composed of cations and anions (2.6)

crystal* solid having a regular three-dimensional arrangement of either ions, atoms, or molecules (2.6)

formula unit group of atoms or ions explicitly symbolized in the chemical formula (2.6)

organic compounds molecular substances that contain carbon combined with other elements, such as hydrogen, oxygen, and nitrogen (2.7)

hydrocarbons compounds containing only hydrogen and carbon (2.7)

functional groups reactive portion of a molecule that undergoes predictable reactions (2.7)

alcohol* molecule that contains an –OH functional group (2.7)

ether* organic molecule that contains an oxygen atom between two carbon atoms (2.7)

chemical nomenclature systematic naming of chemical compounds (2.8)

inorganic compounds compounds composed of elements other than carbon (2.8)

monatomic ion ion formed from a single atom (2.8)

Stock system* system for naming compounds in which a Roman numeral within parentheses follows the first-named element to indicate its charge or oxidation number (2.8)

oxidation state* or oxidation number* hypothetical charge assigned in accordance with certain rules; denoted with a Roman numeral following the name of the metal atom (2.8, margin note)

-ous* in an older system of nomenclature, a suffix added to the stem name of an element to indicate the cation of lower charge; also indicates the oxoacid with fewer oxygen atoms (2.8)

-ic* in an older system of nomenclature, a suffix added to the stem name of an element to indicate the cation of higher charge; indicates the oxoacid with more oxygen atoms; also indicates an acid solution obtained from binary compounds of hydrogen and nonmetals (2.8)

-ide* suffix added to the stem name of the element to name monatomic anions or the more electronegative element in binary compounds (2.8)

polyatomic ion ion consisting of two or more atoms chemically bonded together and carrying a net electric charge (2.8)



oxoanion (oxyanion)* anion composed of oxygen with another element, which is the central element (2.8)

acid* molecular compound that can yield one or more hydronium ions, H3O+, and an anion for each acid molecule when the acid dissolves in water (2.8)

oxoacid acid containing hydrogen, oxygen, and another element (2.8)

-ate* suffix denoting the oxoanion with the greater number of oxygen atoms (2.8)

-ite* suffix denoting the oxoanion with the lesser number of oxygen atoms (2.8)

hypo-* prefix denoting the oxoacid or oxoanion with the least number of oxygen atoms in the series (2.8)

per-* prefix denoting the oxoacid or oxoanion with the greatest number of oxygen atoms in the series (2.8)

acid anions* anions that have hydrogen atoms they can lose as hydronium ions, H3O+ (2.8)

di-* Greek prefix meaning two (2.8)

thio-* prefix meaning an oxygen in the root ion name has been replaced by a sulfur atom (2.8)

binary compound compound composed of only two elements (2.8)

hydro-* prefix added to the stem name of a nonmetal to name the acid solution obtained from binary compounds of hydrogen and nonmetals (2.8)

hydrate compound that contains water molecules weakly bound in its crystals (2.8)

chemical equation symbolic representation of a chemical reaction in terms of chemical formulas (2.9)

reactant starting substance in a chemical reaction; appears to the left of the arrow in a chemical equation (2.9)

product substance that results from a chemical reaction; appears to the right of the arrow in a chemical equation (2.9)

coefficient* number that appears in front of a formula in a chemical equation and gives the relative number of molecules or formula units of a substance involved in the reaction (2.9)

(g)* phase label placed after a formula in a chemical equation to indicate that the substance is a gas (2.9)

(l)* phase label placed after a formula in a chemical equation to indicate that the substance is a liquid (2.9)

(s)* phase label placed after a formula in a chemical equation to indicate that the substance is a solid (2.9)

(aq)* phase label placed after a formula in a chemical equation to indicate that the substance is in aqueous (water) solution (2.9)

catalyst* substance that speeds up a reaction without undergoing any net change itself (2.9)

balanced* describes a chemical equation having correct coefficients (2.10)

balancing by inspection* trial-and-error method of balancing a chemical equation by writing appropriate coefficients until there is the same number of any one elemental atom on each side of the arrow (2.10)



CHAPTER DIAGNOSTIC TEST1. Dalton’s atomic theory postulated that matter

a. is in continuous motion.

b. is continuous (infinitely divisible) in nature.

c. changes in mass when heated to combustion.

d. can exist in three states—gas, liquid, and solid.

e. is composed of small particles called atoms.

2. Robert Millikan’s _______________ experiment, in conjunction with Thomson’s value for m/e of the electron, allowed an accurate calculation of the mass of the electron.

3. In a mass spectrograph, the natural isotopes of iron were observed to have the following atomic masses (and percentage abundances): 53.94 (5.84%), 55.94 (91.68%), 57.94 (2.17%), and 57.93 (0.310%). From these data, give the average atomic mass of iron.

4. Explain the difference in meaning between the symbols H, H2, and H2O.

5. Complete the following statements about aluminum nitrate, Al(NO3)3.

a. One formula unit of aluminum nitrate contains _____ Al atom(s).

b. One formula unit of aluminum nitrate contains _____ N atom(s).

c. One formula unit of aluminum nitrate contains _____ O atom(s).

6. A nucleus consists of 32 protons and 41 neutrons. What is the nuclide symbol for this nucleus?

7. In a sample of hydrogen gas there are 4.92 1018 hydrogen atoms. How many methane molecules (formula CH4) could be formed?

8. When the ions Na+ and CO32 combine chemically, the compound sodium carbonate (called soda

ash) is formed. Keeping in mind that chemical formulas are written as electrically neutral species, write the correct formula for sodium carbonate.

9. Write the molecular formulas for the molecules having the following structural formulas.

a.

b.



c.

d. F—O—O—F

10. Match each term in the left-hand column with a descriptive example in the right-hand column.

___ 1. compound a. Sn

___ 2. Rutherford b. characterized by Z and A

___ 3. atomic symbol c. 1/12 the mass of a C-12 atom

___ 4. J. J. Thomson d. gold-foil experiment

___ 5. nuclide e. cathode-ray-tube experiments

___ 6. amu f. methane (CH4)

11. On the following diagram of the periodic table, indicate the regions where nonmetals, metals, and metalloids are to be found.

12. Fluorine, chlorine, and bromine are _______________ (metals/metalloids/nonmetals) that belong to the Group VIIA family, commonly referred to as the ____________________.

13. Carbonic acid, H2CO3, exists only in aqueous solution and is formed when CO2 dissolves in water. This is what gives carbonated drinks their “sparkling” taste. Give the formula and name of the oxoanion of carbonic acid.

14. Tell whether you would expect the following compounds to be ionic or molecular in nature.

a. XeF4

b. CS2

c. NaI



15. Complete the following chart with the appropriate numbers or symbols.

Symbol C Al3+ ? ? ?

Atomic number ? ? 1 ? 47

Protons ? ? ? 16 ?

Neutrons ? 14 ? ? 61

Mass number 14 ? 1 32 ?

Electrons ? ? 1 18 46

16. Write the formula for each of the following substances.

a. zinc hydrogen carbonate

b. copper(II) dichromate dihydrate

c. manganous hydroxide

d. bismuth nitride

e. plumbous iodide

f. periodic acid

17. Write the name of each of the following substances. Where appropriate, give the Stock system name and the common name.

a. Cr(NO3)3 ∙ 9H2O

b. AlPO4

c. Fe2S3

d. PbSO3

e. Co(CN)2

f. P2O5

g. CO

18. Fill in the blanks, referring to the following:

Fe3O4 + 4H2 3Fe + 4H2O

The preceding expression is called a chemical (a) ______. It describes a chemical (b) ______ in which a chemical (c) ______ occurs in the identity of the reacting molecules. Write a description of the information it gives you: (d) ______.

ANSWERS TO CHAPTER DIAGNOSTIC TESTIf you missed an answer, study the text section and problem-solving skill given (PS Sk.) in parentheses after the answer.

1. e (2.1)

2. oil-drop (2.2)

3. 55.9 amu (2.4, PS Sk. 2)



4. H is the symbol for the element hydrogen. It represents one atom of hydrogen. H2 represents a molecule of the element hydrogen, which is made up of two atoms of hydrogen. H2O represents a molecule of a compound (two or more different elements combined together) that contains two atoms of hydrogen and one atom of oxygen. (2.6)

5.

a. 1

b. 3

c. 9 (2.6)

6. (2.3, PS Sk. 1)

7. 1.23 1018 molecules (2.6)

8. Na2CO3 (2.6, PS Sk. 3)

9.

a. C3H6

b. CH4NCl

c. N2F4

d. O2F2 (2.6)

10. (1) f (2.1), (2) d (2.2), (3) a (2.1), (4) e (2.2), (5) b (2.3), (6) c (2.4)

11.

N o n m eta ls

M eta ls

(2.5)

12. nonmetals, halogens (2.5)

13. CO32– is the carbonate ion. (2.8, PS Sk. 6)

14.

a. molecular

b. molecular

c. ionic (2.8, PS Sk. 5)



15.

Symbol C Al3+ H S2 Ag+

Atomic number

6 13 1 16 47

Protons 6 13 1 16 47Neutrons 8 14 0 16 61Mass number 14 27 1 32 108Electrons 6 10 1 18 46

(2.3, PS Sk. 1)

16.

a. Zn(HCO3)2

b. CuCr2O7 ∙ 2H2O

c. Mn(OH)2

d. BiN

e. PbI2

f. HIO4 (2.8, PS Sk. 4)

17.

a. chromium(III) nitrate nonahydrate

b. aluminum phosphate

c. iron(III) sulfide or ferric sulfide

d. lead(II) sulfite or plumbous sulfite

e. cobalt(II) cyanide or cobaltous cyanide

f. diphosphorus pentoxide

g. carbon monoxide (2.8, PS Sk. 4)

18.

a. equation

b. reaction

c. change

d. 1 formula unit of Fe3O4 (iron oxide) reacts with 4 molecules of H2 (hydrogen) to form 3 atoms of Fe (iron) and 4 molecules of H2O (water). (2.9, 2.10)

SUMMARY OF CHAPTER TOPICSThe definitions presented in this chapter are central to the language of chemistry. Work at mastering them as soon as possible. The descriptions of the terms atom, element, and compound are theoretical explanations that form the basis of our understanding of chemistry. Besides knowing their definitions, it is equally important that you know how atoms, elements, and compounds behave in laboratory work. Both theoretical and practical descriptions are given in the list of chapter terms and definitions.



To keep ion names and charges on the tip of your tongue, use text Tables 2.4 and 2.5 and study guide Table 2.1 to make flip cards on 3 × 5 in index cards. Write the name of the ion or element on one side and the symbol (with charge, if an ion) on the other side. Flip through the cards, putting the ones you don’t know in a separate pile. Work on the ones you don’t know in your spare time.

2.1 Atomic Theory of Matter

Learning Objectives List the postulates of atomic theory.

Define element, compound, and chemical reaction in the context of these postulates.

Recognize the atomic symbols of the elements.

Explain the significance of the law of multiple proportions.

2.2 The Structure of the Atom

Learning Objectives Describe Thomson’s experiment in which he discovered the electron.

Describe Rutherford’s experiment that led to the nuclear model of the atom.

In the cathode-ray tube, the electrode through which the electrons enter the tube, the cathode, is designated the negative electrode. The electrode at which the electrons leave the tube, the anode, is electrically positive with respect to the cathode. In Chapter 18 you will learn more about electricity as it relates to chemistry.

J. J. Thomson established the mass-to-charge ratio of the electron, and Robert Millikan determined the electric charge on an electron. Shortly thereafter, Ernest Rutherford, on the basis of his gold-foil experiment, postulated a nuclear model of the atom.

2.3 Nuclear Structure; Isotopes

Learning Objectives Name and describe the nuclear particles making up the nucleus of the atom.

Define atomic number, mass number, and nuclide.

Write the nuclide symbol for a given nuclide.

Define and provide examples of isotopes of an element.

Write the nuclide symbol of an element. (Example 2.1)

Problem-Solving Skill1. Writing nuclide symbols. Given the number of protons and neutrons in a nucleus, write its

nuclide symbol (Example 2.1).

The atomic number (the number of protons in the nucleus) tells what element we are dealing with. Every calcium atom has 20 protons in its nucleus. The calcium atom loses two electrons to become the Ca2+ ion. This charged atom still has 20 protons in the nucleus. The bromine atom gains one electron to become the bromide ion, Br, with 36 electrons. This charged atom still has 35 protons in its nucleus.



Exercise 2.1A nucleus consists of 17 protons and 18 neutrons. What is its nuclide symbol?

Known : The complete symbol includes the symbol for the element, the mass number, and the atomic number. Atomic number = number of protons; mass number = number of protons + number of neutrons. (See table on inside back cover of the text.)

Solution: Atomic number = 17; mass number = 17 + 18 = 35. The symbol is

The atomic numbers and average atomic masses of the elements are given in the periodic table on the inside front cover of the text. Each element is represented by a square with the atomic symbol in it. The number above the atomic symbol is the atomic number.

2.4 Atomic Masses

Learning Objectives Define atomic mass unit and atomic mass.

Describe how a mass spectrometer can be used to determine the fractional abundance of the isotopes of an element.

Determine the atomic mass of an element from the isotopic masses and fractional abundances. (Example 2.2)

Problem-Solving Skill2. Determining atomic mass from isotopic masses and fractional abundances.

Given the isotopic masses (in atomic mass units) and fractional isotopic abundances for a naturally occurring element, calculate its atomic mass (Example 2.2).

If you look at text Table 2.1, you will see that the relative masses of protons and neutrons are not exactly 1 amu. Thus the mass number of an isotope (the sum of the number of protons and neutrons) always will be a whole number, but the mass of the isotope (in amu) will not be. You will notice from Example 2.2 that we can get the mass number by rounding off the isotopic mass to a whole number.

To take an average, we usually add the given values and divide by the number of them. For instance, the average of 5, 7, and 9 is 5 + 7 + 9 = 2 1/3 = 7. This method gives us the correct answer because each value has equal weight or representation. We could have gotten the same average, 7, by taking ⅓ of each value and then adding them:

+ + = + + = = 7

This second method exemplifies a weighted average. We must use this method for determining average relative atomic masses because there is never an equal number of atoms of each isotope in any naturally occurring sample of an element. We could give the amount of each isotope as a percentage of the total, but it is more useful to give the value in decimal form, which we call the fractional abundance. For example, in Example 2.2 the fractional abundance of chromium-50 is 0.0435. This means that 4.35% of the atoms in any naturally occurring sample of chromium are the chromium-50 isotope. You should memorize this method of finding atomic masses.



Exercise 2.2Chlorine consists of the following isotopes:

Isotope Mass (amu) Fractional AbundanceChlorine-35 34.96885 0.75771Chlorine-37 36.96590 0.24229

What is the atomic mass of chlorine?

Known: The atomic mass is the weighted average of isotopic masses.

Solution: 34.96885 amu 0.75771 = 26.4962 amu

36.96590 amu 0.24229 = 8.95647 amu

Atomic mass of chlorine = 35.453 amu

Although the actual masses of atoms are known, the relative atomic masses (called atomic masses) are much easier to use in calculations. For example, the actual average atomic mass of a calcium atom is 6.656 1023 g, whereas its relative average atomic mass (atomic mass) is 40.08 amu. Be sure you understand the difference between the two and the relationship between them.

2.5 Periodic Table of the Elements

Learning Objectives Identify periods and groups on the periodic table.

Find the main-group and transition elements on the periodic table.

Locate the alkali-metal and halogen groups on the periodic table.

Recognize the portions of the periodic table that contain the metals, nonmetals, and metalloids (semimetals).

There are many periodic phenomena in our experience. The phases of the moon are periodic; every 28 days we see a full moon. The seasons of the year are periodic; there is a regular repetition of winter, spring, summer, and fall.

Exercise 2.3By referring to the periodic table (Figure 2.15 or inside front cover of the text), identify the group and period to which each of the following elements belongs. Then decide whether the element is a metal, nonmetal, or metalloid.

a. Se

b. Cs

c. Fe

d. Cu

e. Br

Known: Elements to the left of the periodic table staircase line with characteristic properties are metals; elements to the right of the line with characteristic properties are nonmetals; elements bordering the line with properties of metals and nonmetals are metalloids.



Solution:

a. Selenium is a nonmetal. It is to the right of the staircase line, in Period 4 and Group VIA.

b. Cesium is a metal. It is to the left of the line, in Period 6 and Group IA, the alkali metals.

c. Iron is a metal. It is in Period 4 and Group VIIIB.

d. Copper is a metal. It is in Period 4 and Group IB.

e. Bromine is a nonmetal. It is in Period 4 and Group VIIA, the halogens.

2.6 Chemical Formulas; Molecular and Ionic Substances

Learning Objectives Determine when the chemical formula of a compound represents a molecule.

Determine whether a chemical formula is also a molecular formula.

Define ion, cation, and anion.

Classify compounds as ionic or molecular.

Define and provide examples for the term formula unit.

Specify the charge on all substances, ionic and molecular.

Write an ionic formula, given the ions. (Example 2.3)

Problem-Solving Skill3. Writing an ionic formula, given the ions. Given the formulas of a cation and an anion, write the

formula of the ionic compound of these ions (Example 2.3).

Exercise 2.4Potassium chromate is an important compound of chromium (see Figure 2.23). It is composed of K+ and CrO4

2 ions. Write the formula of the compound.

Known: Compounds must be electrically neutral.

Solution: There must be two K+ ions to provide two positive charges to neutralize the 2 charge on the chromate ion. The formula is K2CrO4.

2.7 Organic Compounds

Learning Objectives List the attributes of molecular substances that make them organic compounds.

Explain what makes a molecule a hydrocarbon.

Recognize some functional groups of organic molecules.



2.8 Naming Simple Compounds

Learning Objectives Recognize inorganic compounds.

Learn the rules for predicting the charges of monatomic ions in ionic compounds.

Apply the rules for naming monatomic ions.

Learn the names and charges of common polyatomic ions.

Name an ionic compound from its formula. (Example 2.4)

Write the formula of an ionic compound from its name. (Example 2.5)

Determine the order of elements in a binary (molecular) compound.

Learn the rules for naming binary molecular compounds, including the Greek prefixes.

Name a binary compound from its formula. (Example 2.6)

Write the formula of a binary compound from its name. (Example 2.7)

Name a binary molecular compound from its molecular model. (Example 2.8)

Recognize molecular compounds that are acids.

Determine whether an acid is an oxoacid.

Learn the approach for naming binary acids and oxoacids.

Write the name and formula of an anion from the acid. (Example 2.9)

Recognize compounds that are hydrates.

Learn the rules for naming hydrates.

Name a hydrate from its formula. (Example 2.10)

Write a formula of a hydrate from its name. (Example 2.11)

Problem-Solving Skills4. Writing the name of a compound from its formula, or vice versa. Given the formula of a

simple compound (ionic, binary molecular, acid, or hydrate), write the name (Examples 2.4, 2.6, and 2.10), or vice versa (Examples 2.5, 2.7, and 2.11).

5. Writing the name and formula of an anion from the acid. Given the name and formula of an oxoacid, write the name and formula of the oxoanion, or given the name and formula of an oxoanion, write the formula and name of the oxoacid (Example 2.9).

A Chemist Looks at: Thirty Seconds on the Island of Stability

Questions for Study1. Which elements are unstable and fall apart by radioactive decay?

2. What is the heaviest naturally occurring element?

3. How are the transuranium elements generally made in the laboratory?

4. How are elements represented in nuclear equations?



5. What would you predict to be the nature of the yet-to-be-discovered elements 113, 115, and 117?

Answers to Questions for Study1. Those elements with nuclides having atomic numbers greater than 83 are unstable and undergo

radioactive decay.

2. The heaviest naturally occurring element is uranium (element 92), which is unstable and undergoes radioactive decay.

3. The transuranium elements are made by bombarding elements of fairly high atom mass with particles of lower mass.

4. In nuclear equations, the elements are represented using their nuclide symbols, which include information on the atomic mass number (number of protons and neutrons) and atomic number (number of protons).

5. Elements 113, 115, and 117 would be seventh-row elements in the periodic table and should exhibit properties consistent with Groups 3A, 5A, and 7A, respectively.

You must learn the element names and symbols and the ion names, formulas, and charges to master naming and writing the formulas of compounds. Refer to text Tables 2.4 and 2.5 or to study guide Table 2.1. Make flip cards to help you.

It also will be useful for you to know the formulas and names of common acids. They are given in text Table 2.7 and study guide Table 2.2.

Exercise 2.5Write the names of the following compounds.

a. CaO

b. PbCrO4

a. Known: CaO is a binary compound and will end in -ide. Ca (Group IIA) forms the 2+ ion; thus the oxygen ion is O2–.

Solution: Calcium oxide

b. Known: The chromate ion has a charge of 2–. The lead ion is thus Pb2+.

Solution: Lead(II) chromate, or plumbous chromate

Exercise 2.6A compound has the name thallium(III) nitrate. What is the formula? (The symbol of thallium is Tl.)

Known: Tl (Group IIIA) has a charge of 3+; nitrate is NO3; the compound must be neutral.

Solution: Tl(NO3)3



Table 2.1 Charges, Formulas, and Names of Some Common Ions

Cations Anions

Charge Formula Name Charge Formula Name1+ NH4

+ ammonium 1 C2H3O2 acetate

Cs+ cesium Br bromideCu+ copper(I); cuprous Cl chlorideH3O+ hydronium ClO hypochloriteLi+ lithium ClO2

chloriteK+ potassium ClO3

chlorateRb+ rubidium ClO4

perchlorateAg+ silver CN cyanideNa+ sodium H2PO4

dihydrogen phosphateF fluoride

2+ Ba2+ barium H hydrideBe2+ beryllium HCO3

hydrogen carbonate;Cd2+ cadmium bicarbonateCa2+ calcium HSO4

hydrogen sulfate;Co2+ cobalt(II); cobaltous bisulfateCu2+ copper(II); cupric HSO3

hydrogen sulfite;Fe2+ iron(II); ferrous bisulfitePb2+ lead(II); plumbous OH hydroxideMg2+ magnesium I IodideMn2+ manganese(II); NO3

nitratemanganous NO2

nitriteHg2

2+ mercury(I); mercurous

MnO4 permanganate

2+ Hg2+ mercury(II); mercuric

2 CO32 carbonate

Ni2+ nickel(II); nickelous C2O42 oxalate

Sr2+ strontium CrO42 chromate

Sn2+ tin(II);stannous Cr2O72 dichromate

Zn2+ zinc HPO42 monohydrogen phosphate

O2 oxide3+ Al3+ aluminum O2

2 peroxideBi3+ bismuth SO4

2 sulfateCr3+ chromium(III); S2 sulfide

chromic SO32 sulfite

Fe3+ iron(III); ferric S2O32 thiosulfate

4+ Sn4+ tin(IV); stannic 3 AsO43 arsenate

N3 nitridePO4

3 phosphate



Table 2.2 Formulas and Names of Some Common Acids

Formula Name Formula NameHF(aq) hydrofluoric acid HClO2 chlorous acidHCl(aq) hydrochloric acid HClO3 chloric acidHBr(aq) hydrobromic acid HClO4 perchloric acidHI(aq) hydroiodic acid HNO3 nitric acidH2S(aq) hydrosulfuric acid HNO2 nitrous acidHCN(aq) hydrocyanic acid H2SO4 sulfuric acidHC2H3O2 acetic acid H2SO3 sulfurous acidH2CO3 carbonic acid H3AsO4 arsenic acidHClO hypochlorous acid H3PO4 phosphoric acid

Following are practice grids and answers for writing chemical formulas and names for ionic and covalent compounds.

Formula/Nomenclature Practice Grid for Ionic Compounds

AnionCation Fluoride Chloride Oxide Sulfide Sulfate Sulfite Carbonate Chlorate

SodiumPotassiumMagnesiumCalciumAluminumChromium(II)Iron(II)Iron(III)Copper(I)Copper(II)SilverAmmonium

This grid is designed to help you learn to name a compound when the formula is given and to write the formula when the name of the compound is given. To use this grid effectively:

1. read aloud the name of the compound as you write the formula, and

2. once all formulas are written, cover up the names of the cations and anions and recite the names of the compounds from the formulas.

Note: This is a naming exercise; some of the named compounds do not exist.



Formula/Nomenclature Practice Grid for Molecular Compounds

Fluorine Chlorine Oxygen SulfurFormula Name Formula Name Formula Name Formula Name

Boron(III)

Silicon(IV)

Phosphorus(V)

Tungsten(VI)

This grid is designed to help you learn to name molecular compounds. To use this grid effectively:

1. write the formula for the compound formed between the element in the far left column (which is written as if it were an ion) and the nonmetal element listed across the top of the grid, and

2. write the name using Greek number prefixes.



Answers

Formula/Nomenclature Practice Grid for Ionic Compounds

AnionCation Fluoride Chloride Oxide Sulfide Sulfate Sulfite Carbonate Chlorate

Sodium NaF NaCl Na2O Na2S Na2SO4 Na2SO3 Na2CO3 NaClO3

Potassium KF KCl K2O K2S K2SO4 K2SO3 K2CO3 KClO3

Magnesium MgF2 MgCl2 MgO MgS MgSO4 MgSO3 MgCO3 Mg(ClO3)2

Calcium CaF2 CaCl2 CaO CaS CaSO4 CaSO3 CaCO3 Ca(ClO3)2

Aluminum AlF3 AlCl3 Al2O3 Al2S3 Al2(SO4)3 Al2(SO3)3 Al2(CO3)3 Al(ClO3)3

Chromium(II) CrF2 CrCl2 CrO CrS CrSO4 CrSO3 CrCO3 Cr(ClO3)2

Iron(II) FeF2 FeCl2 FeO FeS FeSO4 FeSO3 FeCO3 Fe(ClO3)2

Iron(III) FeF3 FeCl3 Fe2O3 Fe2S3 Fe2(SO4)3 Fe2(SO3)3 Fe2(CO3)3 Fe(ClO3)3

Copper(I) CuF CuCl Cu2O Cu2S Cu2SO4 Cu2SO3 Cu2CO3 CuClO3

Copper(II) CuF2 CuCl2 CuO CuS CuSO4 CuSO3 CuCO3 Cu(ClO3)2

Silver AgF AgCl Ag2O Ag2S Ag2SO4 Ag2SO3 Ag2CO3 AgClO3

Ammonium NH4F NH4Cl (NH4)2O (NH4)2S (NH4)2SO4 (NH4)2SO3 (NH4)2CO3 NH4ClO3



Answers

Formula/Nomenclature Practice Grid for Molecular Compounds

Fluorine Chlorine Oxygen SulfurFormula Name Formula Name Formula Name Formula Name

Boron(III) BF3

borontrifluoride BCl3

borontrichloride B2O3

diborontrioxide B2S3

diborontrisulfide

Silicon(IV) SiF4

silicontetrafluoride SiCl4

silicontetrachloride SiO2

silicondioxide SiS2

silicondisulfide

Phosphorus(V) PF5

phosphoruspentafluoride PCl5

phosphoruspentachloride P2O5

diphosphoruspentoxide P2S5

diphosphoruspentasulfide

Tungsten(VI) WF6

tungstenhexafluoride WCl6

tungstenhexachloride WO3

tungstentrioxide WS3

tungstentrisulfide

Exercise 2.7Name the following compounds: (a) Cl2O6, (b) PCl3, (c) PCl5.

Known: For binary compounds, we name the first element and add a suffix to the stem name of the second element. For elements forming more than one compound, we use the Greek prefixes listed in text Table 2.6, omitting mono- for a first-named element.

Solution: (a) The text list of compounds in which the Greek prefixes are used shows that chlorine and oxygen form more than one compound. The name would be dichlorine hexoxide. (b) phosphorus trichloride (c)phosphorus pentachloride

Exercise 2.8Give formulas for the following compounds.

a. carbon disulfide

b. sulfur trioxide

Known: We change the names of the elements to symbols and translate the prefixes to subscripts.

Solution:

a. CS2

b. SO3

Exercise 2.9Using the molecular models, name the following chemical compounds: (a) BF3 (b) H2Se.

Known: Because these are binary molecular compounds, we use Greek prefixes listed in text Table 2.6 and name the elements in the order in which they appear in the chemical formula.

Solution: (a) boron trifluoride (b) dihydrogen selenide



Exercise 2.10What is the name and formula of the anion corresponding to perbromic acid, HBrO4?

Known: We remove H+ from the acid to get the ion. We name the ion by replacing -ic in the acid name with -ate.

Solution: The ion is BrO4–, and it is the perbromate ion.

Exercise 2.11Washing soda has the formula Na2CO3 ∙ 10H2O. What is the chemical name of this substance?

Known: The formula indicates that the compound is a hydrate. We name hydrates from the anhydrous compound, followed by the word hydrate with a Greek prefix to indicate the associated number of water molecules.

Solution: The name is sodium carbonate decahydrate.

Exercise 2.12Photographers’ hypo, used to fix negatives during the development process, is sodium thiosulfate pentahydrate. What is the chemical formula of this compound?

Known: The anhydrous compound is composed of sodium ions (Na+) and thiosulfate ions (S2O3

2); pentahydrate means there are five associated molecules of water.

Solution: The formula is Na2S2O3 ∙ 5H2O.

2.9 Writing Chemical Equations

Learning Objectives Identify the reactants and products in a chemical equation.

Write chemical equations using appropriate phase labels, symbols of reaction conditions, and the presence of a catalyst.

2.10 Balancing Chemical Equations

Learning Objectives Determine if a chemical reaction is balanced.

Master techniques for balancing chemical equations. (Example 2.12)

Problem-Solving Skill6. Balancing simple equations. Given the formulas of the reactants and products in a chemical

reaction, obtain the coefficients of the balanced equation (Example 2.12).

Balancing equations does not have to be a difficult process. The principle behind it is the law of conservation of matter. Usually, balancing equations is a trial-and-error process. Always check your final answer to be sure that you have the same number of atoms of each element on each side of the arrow.



Exercise 2.13Find the coefficients that balance the following equations.

a. O2 + PCl3 POCl3

b. P4 + N2O P4O6 + N2

c. As2S3 + O2 As2O3 + SO2

d. Ca3(PO4)2 + H3PO4 Ca(H2PO4)2

Solution:

a. O2 + 2PCl3 2POCl3

b. P4 + 6N2O P4O6 + 6N2

c. 2As2S3 + 9O2 2As2O3 + 6SO2

d. Ca3(PO4)2 + 4H3PO4 3Ca(H2PO4)2

ADDITIONAL PROBLEMS1. Supply the missing information in the following table.

Symbol C ? Al3+

Atomic number ? ? ?Number of protons ? 16 ?Number of neutrons 8 16 ?Number of electrons ? 18 ?Mass number ? ? 27

2. Explain the nature of isotopes of elements in terms of the subatomic particles of matter.

3. Write the symbol for

a. a nucleus containing 8 protons and 8 neutrons.

b. the carbon-14 nucleus.

4. A sample of neon always contains the three isotopes of neon: neon-20, neon-21, and neon-22. The natural abundances of these isotopes are 90.92%, 0.257%, and 8.82%, respectively. Their isotopic masses are 19.99244, 20.99395, and 21.99138 amu, respectively. Calculate the atomic mass of neon.

5. Write the formula for the compound of each of the following ion pairs.

a. Ca2+ and H–

b. Al3+ and OH–

c. Mg2+ and PO43–

d. NH4+ and SO4

2–

6. Classify each of the following elements as a metal, a metalloid, or a nonmetal: S, Sr, Co, N, Ga, I, Ar, Si, Li, V.


Symbol P ? H Mg2+ ?



Atomic number ? 11 ? ? 29Protons ? ? ? ? ?Neutrons ? 12 2 ? ?Mass number 31 ? ? 24 64Electrons ? 10 ? ? 29

8. Write the formula for each of the following substances.

a. sodium sulfite heptahydrate

b. iron(III) chloride

c. calcium fluoride

d. sulfuric acid

e. silicon tetrabromide

f. barium hydrogen sulfite

g. perchloric acid

h. dinitrogen tetroxide

9. Name each of the following compounds. Where appropriate, give both the Stock system name and the common name.

a. SrBr2 ∙ 6H2O

b. MnCl2

c. HBrO3

d. NO2

e. LiH

f. BCl3

10. For each of the following, write the symbol and name for the corresponding oxoacid or oxoanion.

a. HClO, hypochlorous acid

b. PO43–, phosphate ion

c. HNO3, nitric acid

d. SO32–, sulfite ion



11. Determine which of the following statements is (are) incorrect with respect to the following chemical equation:

H2 + Cl2 2HCl

a. Two atoms of hydrogen occur on both sides of the equation.

b. The chemical reaction described by this equation consists of a rearrangement of the atoms of hydrogen and chlorine to give hydrogen chloride (HCl).

c. Since one molecule of H2 and one molecule of Cl2 react, two molecules of product must form.

d. This equation describes a chemical change.

12. Balance each of the following equations.

a. ___ Na + ___ O2 ___ Na2O

b. ___ As + ___ H2 ___ AsH3

c. ___ Ba(OH)2 + ___ HNO3 ___ Ba(NO3)2 + ___ H2O

d. ___ Al + ___ H2SO4 ___ Al2(SO4)3 + ___ H2

e. ___ C3H8 + ___ O2 ___ CO2 + H2O

13. A solution of lead sulfate reacts with a solution of sodium chloride to form solid lead chloride and aqueous sodium sulfate. Write the balanced equation for this reaction, including state (phase) labels.

ANSWERS TO ADDITIONAL PROBLEMSIf you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.

1.

Symbol C S2– Al3+

Atomic number 6 16Number of protons 6 16 13Number of neutrons 8 16 14Number of electrons 6 18 10Mass number 14 32 27

(2.3)

2. Isotopes are atoms of the same element that have the same number of protons and the same number of electrons but different numbers of neutrons. Therefore, isotopes have different masses. (2.3)

3. a.

b. (2.3, PS Sk. 1)



4. The atomic mass is the weighted average of isotopic masses.

19.99244 amu 0.9092 = 18.1771 amu

20.99395 amu 0.00257 = 0.0540 amu

21.99138 amu 0.0882 = 1.9396 amu

Atomic mass of neon = 20.1707 = 20.17 amu (2.4, PS Sk. 2)

5.

a. CaH2

b. Al(OH)3

c. Mg3(PO4)2

d. (NH4)2SO4 (2.6, PS Sk. 3)

6. The metals are Sr, Co, Ga, Li, and V. Si is a metalloid. The nonmetals are S, N, I, and Ar. (2.5)

7.

Symbol P Na+ H Mg2+ CuAtomic number 15 11 1 12 29Protons 15 11 1 12 29Neutrons 16 12 2 12 35Mass number 31 23 3 24 64Electrons 15 10 1 10 29

(2.3)

8.

a. Na2SO3 ∙ 7H2O

b. FeCl3

c. CaF2

d. H2SO4

e. SiBr4

f. Ba(HSO3)2

g. HClO4

h. N2O4 (2.8, PS Sk. 4)

9.

a. strontium bromide hexahydrate

b. manganese(II) chloride, manganous chloride

c. bromic acid

d. nitrogen dioxide

e. lithium hydride

f. boron trichloride (2.8, PS Sk. 4)



10.

a. ClO–, hypochlorite ion

b. H3PO4, phosphoric acid

c. NO3–, nitrate ion

d. H2SO3, sulfurous acid (2.8, PS Sk. 6)

11. c (2.10)

12.

a. 4Na + O2 2Na2O

b. 2As + 3H2 2AsH3

c. Ba(OH)2 + 2HNO3 Ba(NO3)2 + 2H2O

d. 2Al + 3H2SO4 Al2(SO4)3 + 3H2

e. C3H8 + 5O2 3CO2 + 4H2O (2.10, PS Sk. 7)

13. PbSO4(aq) + 2NaCl(aq) PbCl2(s) + Na2SO4(aq) (2.6, 2.8, 2.9, 2.10, PS Sk. 6, 7)

CHAPTER POST-TEST1. Indicate whether each of the following statements is true or false. If a statement is false, change it

so that it is true.

a. According to Dalton’s atomic theory, compounds are kinds of matter composed of atoms of two or more elements. True/False: ______________________________________________

__________________________________________________________________________

b. A chemical symbol is used to designate elements, and formulas are used to designate formula units. True/False: _____________________________________________________

__________________________________________________________________________

c. A molecular formula does not indicate the arrangement of atoms in a molecule. True/False:_________________________________________________________________

__________________________________________________________________________

d. KClO3, Na2S, and BF3 are ionic compounds. True/False: ____________________________

__________________________________________________________________________

2. If you have one dozen formula units of Mg(OH)2, how many hydrogen atoms do you have?



3. The structure of sulfuric acid, the most widely used chemical, is

Write the molecular formula for sulfuric acid. (Hint: The order of atoms is H, S, and then O.)

4. Which one of the following statements correctly describes the difference between a structural and a molecular formula?

a. The structural formula indicates the composition of the molecule and the spatial arrangement of the atoms, and the molecular formula shows how the different atoms bond with each other.

b. The structural formula represents the simplest composition of a molecule, and the molecular formula represents the actual composition.

c. The structural formula shows how the atoms are bonded together in a molecule, and the molecular formula shows the atomic composition of the molecule.

d. None of the above are correct statements.

5. Write the symbol and name for the corresponding oxoacid or oxoanion.

a. BrO2–, bromite ion

b. HNO2, nitrous acid

c. IO3–, iodate ion

d. C2H3O2–, acetate ion


Symbol Cr ? Br ? ?Atomic number ? 26 ? ? 24Protons ? ? ? 92 ?Neutrons 28 ? ? 143 29Mass number ? 56 80 ? ?Electrons ? 24 ? 92 24

7. Classify each of the following elements as a metal, a nonmetal, or a metalloid: Mg, Cu, Pb, As, Cl.

8. Which of the following is an incorrect statement? Rewrite that statement so that it is correct.

a. Na and Cs are in the same group in the periodic table.

b. The nonmetallic elements are on the far right side of the periodic table.

c. Most of the known chemical elements are classified as metals.

d. Elements in the same period of the periodic table have similar properties.

e. Fluorine is classified as a nonmetallic element.



9. Write the formula for each of the following.

a. barium peroxide

b. silver chlorite

c. strontium oxalate monohydrate

d. ammonium iodate

e. stannous nitrite

f. phosphoric acid

10. Write the name of each of the following. Where appropriate, give the Stock system name and the common name.

a. KClO4

b. Hg2Br2

c. Mg(C2H3O2)2 ∙ 4H2O

d. CaH2

e. NaMnO4

f. HNO3

11. Fill in the blanks, referring to the following:

2ZnS(s) + 3O2(g) 2SO2(g) + 2ZnO

This equation is a symbolic representation of the (a) _____ between two (b) _____ of (c) _____ in the (d) _____ state and three (e) _____ of oxygen (f) _____ to produce two (g) _____ of gaseous (h) _____ and two formula units of solid (i) _____. The triangle over the arrow indicates that the reactants are (j) _____ to produce the products.

12. Balance the following equations.

a. ___ P + ___ O2 ___ P4O10

b. ___ Ag + ___ NiSO4 ___ Ag2SO4 + ___ Ni

c. ___ C6H14 + ___ O2 ___ CO2 + ___ H2O

d. ___ HF + ___ SiO2 ___ SiF4 + ___ H2O

e. ___ NaHCO3 ___ Na2CO3 + ___ H2O + ___ CO2



ANSWERS TO CHAPTER POST-TESTIf you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.

1.

a. True. (2.1)

b. True. (2.1, 2.6)

c. True. (2.6)

d. False. KClO3 and Na2S are ionic compounds; BF3 is a molecular compound. (2.8)

2. two dozen (2.1)

3. H2SO4 (2.6)

4. c (2.6)

5.

a. HBrO2, bromous acid

b. NO2–, nitrite ion

c. HIO3, iodic acid

d. HC2H3O2, acetic acid (2.8, PS Sk. 5)

6.

Symbol Cr Fe2+ Br U CrAtomic number 24 26 35 92 24Protons 24 26 35 92 24Neutrons 28 30 45 143 29Mass number 52 56 80 235 53Electrons 24 24 36 92 24

(2.3)

7. Mg, Cu, and Pb are metals, Cl is a nonmetal, and As is a metalloid. (2.5)

8. d. Elements in the same group of the periodic table have similar properties. (2.5)

9.

a. BaO2

b. AgClO2

c. SrC2O4 ∙ H2O

d. NH4IO3

e. Sn(NO2)2

f. H3PO4 (2.8, PS Sk. 4)



10.

a. potassium perchlorate

b. mercury(I) bromide or mercurous bromide

c. magnesium acetate tetrahydrate

d. calcium hydride

e. sodium permanganate

f. nitric acid (2.8, PS Sk. 4)

11.

a. chemical reaction

b. formula units

c. zinc sulfide

d. solid (2.9, 2.10)

e. molecules

f. gas

g. molecules

h. sulfur dioxide

i. zinc oxide

j. heated

12.

a. 4P + 5O2 P4O10

b. 2Ag + NiSO4 Ag2SO4 + Ni

c. 2C6H14 + 19O2 12CO2 + 14H2O

d. 4HF + SiO2 SiF4 + 2H2O

e. 2NaHCO3 Na2CO3 + H2O + CO2 (2.10, PS Sk. 6)


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