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9 ACIDS, BASES, AND SALTS

9 ACIDS, BASES, AND SALTS

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to yield 2 H+ per H2SO4 molecule; the second of the 2 hydrogen ions comes off

much less readily than the first. Some acids have hydrogen atoms that do not

produce H + ion in water. For example, each molecule of acetic acid produces only

one H+ ion when it dissociates in water; the other 3 H atoms stay covalently bonded

to the acetate ion:

H O

H O

(4.9.4)

H C C OH

H C C O + H+

H



H



Some of the most widely produced industrial chemicals are acids. Sulfuric acid

ranks first among all chemicals produced in the United States with production of

almost 40 million metric tons per year. Its greatest single use is in the production of

phosphate fertilizers, and it has applications in many other areas including petroleum

refining, alcohol synthesis, iron and steel pickling (corrosion removal), and storage

battery manufacture. Nitric acid ranks about 10th among U.S. chemicals with annual

production of 7–8 million metric tons, and hydrochloric acid is about 25th at about 3

million metric tons (annual production and rank vary from year to year).

The naming of acids is addressed in more detail in Chapter 6. Briefly, acids that

contain only H and another atom are “hydro-ic” acids, such as hydrochloric acid,

HCl. For acids that contain two different amounts of oxygen in the anion part, the

one with more oxygen is an “-ic” acid and the one with less is an “-ous” acid. This is

illustrated by nitric acid, HNO3, and nitrous acid, HNO2. An even greater amount of

oxygen is denoted by a “per-” prefix and less by the “hypo-ous” name. The guidelines discussed above are illustrated for acids formed by chlorine ranging from 0 to 4

oxygen atoms per acid molecule as follows: HCl, hydrochloric acid; HClO, hypochlorous acid; HClO2, chlorous acid; HClO3, chloric acid; HClO4, perchloric acid.



Bases

A base is a substance that contains hydroxide ion OH-, or produces it when

dissolved in water. Most of the best known inorganic bases have a formula unit

composed of a metal cation and 1 or more hydroxide ions. Typical of these are

sodium hydroxide, NaOH, and calcium hydroxide, Ca(OH)2, which dissolve in water

to yield hydroxide ion and their respective metal ions. Other bases such as ammonia,

NH3, do not contain hydroxide ion, but react with water,

NH3 + H2O → NH4+ + OH-



(4.9.5)



to produce hydroxide ion (this reaction proceeds only to a limited extent; most of the

NH3 is in solution as the NH3 molecule).

Bases are named for the cation in them plus “hydroxide.” Therefore, KOH is

potassium hydroxide.



Salts

A salt is an ionic compound consisting of a cation other than H+ and an anion

other than OH-. A salt is produced by a chemical reaction between an acid and a



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base. The other product of such a reaction is always water. Typical salt-producing

reactions are given below:

KOH

base



+



Ca(OH)2 + H2SO4

base







HCl

acid



acid



KCl



+



a salt, potassium chloride







CaSO4



+



a salt, calcium sulfate



H2O



(4.9.6)



water



2H 2O



(4.9.7)



water



Except for those compounds in which the cation is H+ (acids) or the anion is OH(bases), the compounds that consist of a cation and an anion are salts. Therefore, the

rules of nomenclature discussed for ionic compounds in Section 4.8 are those of

salts. The salt product of Reaction 4.9.6, above, consists of K+ cation and Cl- anion,

so the salt is called potassium chloride. The salt product of Reaction 4.9.7, above, is

made up of Ca 2+ cation and SO42- anion and is called calcium sulfate. The reaction

product of LiOH base with H2SO4 acid is composed of Li+ ions and SO42- ions. It

takes 2 singly charged Li+ ions to compensate for the 2- charge of the SO42- anion,

so the formula of the salt is Li2SO4. It is called simply lithium sulfate. It is not

necessary to call it dilithium sulfate because the charges on the ions denote the

relative numbers of ions in the formula.



CHAPTER SUMMARY

The chapter summary below is presented in a programmed format to review the

main points covered in this chapter. It is used most effectively by filling in the

blanks, referring back to the chapter as necessary. The correct answers are given at

the end of the summary.

Chemical bonds are normally formed by the transfer or sharing of 1

, which are those in the 2

. An

especially stable group of electrons attained by many atoms in chemical compounds

is an 3

. An ion consists of 4

. A cation has 5



and an anion has

. An ionic compound is one that contains 7

and is held together by 8

. Both F - and Mg2+

have the electron configuration 9

identical to that of the neutral

. In visualizing a neutral metal atom reacting with a

atom 10

neutral nonmetal atom to produce an ionic compound, the three major energy factors

are 11

. The energy required to separate all

of the ions in a crystalline ionic compound and remove them a sufficient distance

from each other so that there is no interaction between them is called the 12

. The ion formed from Ca is 13

, the ion

14

formed from Cl is

, and the formula of the compound formed from these

ions is 15

. A covalent bond may be described as 16

6



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. In the figure,

H



..



H :N:

..

H

each pair of dots between N and H represents 17

the arrow points to 18

and the circle outlines 19

. A central atom refers to 20

. A triple bond

21



consists of

and is represented in a structural chemical formula as 22

.

With increasing bond order (single
and

. Electronegativity refers to 25

bond strength 24

. A polar covalent bond is one in

which 26

. A coordinate covalent bond is 27

. Three major exceptions to the octet rule are 28



.

Resonance structures are those for which 29

.

Chemical formulas consist of 30

that tell the following about a compound 31

. The percentage elemental composition of a chemical compound is calculated by 32



. The empirical formula of a chemical compound is

calculated by computing the masses of each consituent element in 33

,

dividing each of the resulting values by 34

, dividing each value by 35

,

36

and rounding to

.

The prefixes for numbers 1-10 used to denote relative numbers of atoms of each kind

of atom in a chemical formula are 37

.



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The compound N2O5 is called 38

compound AlCl3 is called 39

contain a “tri-” because 40

a



. The ionic

, which does not

. Ammonium ion, NH4+, is classified as

meaning that it consists of 42

. An acid is 43

.



41



A base is 44

. A salt is 45

duced by 46

The other product of such a reaction is always 47



and is pro.

.



Answers to Chapter Summary

1.

2.

3.

4.

5.

6.

7.

8.

9.

10.

11.

12.

13.

14.

15.

16.

17.

18.

19.

20.

21.

22.

23.

24.

25.

26.

27.



valence electrons

outermost shell of the atom

octet of outer electrons

an atom or group of atoms having an unequal number of electrons and protons

and, therefore, a net electrical charge

a positive charge

a negative charge

cations and anions

ionic bonds

1s22s22p6

neon

ionization energy, electron affinity, and lattice energy

lattice energy

Ca2+

ClCaCl2

one that joins 2 atoms through the sharing of 1 or more pairs of electrons

between them

a pair of electrons shared in a covalent bond

an unshared pair of electrons

a stable octet of electrons around the N atom

an atom to which several other atoms are bonded

3 pair (total of 6) electrons shared in a covalent bond



..



or : :

decreases

increases

the ability of a bonded atom to attract electrons to itself

the electrons involved are not shared equally

one in which only 1 of the 2 atoms contributes the two electrons in the bond



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28. molecules with an uneven number of valence electrons, molecules in which an

atom capable of forming an octet has fewer than 8 outer electrons, and

molecules in which an atom has more than 8 outer electrons

29. it is possible to draw two or more equivalent arrangements of electrons

30. atomic symbols, subscripts, and sometimes parentheses and charges

31. elements in it, relative numbers of each kind of atom, charge, if an ion

32. dividing the mass of a mole of a compound into the mass of each of the constituent elements in a mole

33. 100 g of the element

34. the atomic mass of the element

35. the smallest value

36. the smallest whole number for each element

37. 1-mono, 2-di,3-tri, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octa, 9-nona, 10-deca

38. dinitrogen pentoxide

39. aluminum chloride

40. the charges on ions are used to deduce chemical formulas

41. polyatomic ion

42. 2 or more atoms per ion

43. a substance that dissolves in water to produce hydrogen ion, H+(aq)

44. a substance that contains hydroxide ion OH-, or produces it when dissolved in

water

45. an ionic compound consisting of a cation other than H+ and an anion other than

OH-.

46. a chemical reaction between an acid and a base

47. water



QUESTIONS AND PROBLEMS

Section 4.1. Chemical Bonds and Compound Formation

1. Chlorofluorocarbons (Freons) are composed of molecules in which Cl and F

atoms are bonded to 1 or 2 C atoms. These compounds do not break down well in

the atmosphere until they drift high into the stratosphere, where very short wavelength ultraviolet electromagnetic radiation from the sun is present, leading to the

production of free Cl atoms that react to deplete the stratospheric ozone layer.

Recalling what has been covered so far about the energy of electromagnetic radiation as a function of wavelength, what does this say about the strength of C-Cl

and C-F bonds?

Section 4.2. Chemical Bonding and the Octet Rule

2. Illustrate the octet rule with examples of (a) a cation, (b) an anion, (c) a diatomic

elemental gas, and (d) a covalently bound chemical compound.

Section 4.3. Ionic Bonding

3. When elements with atomic numbers 6 through 9 are covalently or ionically

bound, or when Na, Mg, or A1 have formed ions, which single element do their



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outer electron configurations most closely resemble? Explain your answer in

terms of filled orbitals.

4. Ionic bonds exist because of (a)

.



between (b)



5. In the crystal structure of NaCl, what are the number and type of ions that are

nearest neighbors to each Na+ ion?

6. What major aspect of ions in crystals is not shown in Figure 4.5?

7. What are five energy factors that should be considered in the formation of ionic

NaCl from solid Na and gaseous Cl2?

8. What are two major factors that increase the lattice energy of ions in an ionic

compound?

9. Is energy released or is it absorbed when gaseous ions come together to form an

ionic crystal?

10. What is incomplete about the statement that “the energy change from lattice

energy for NaCl is 785 kilojoules of energy released?

11. How do the sizes of anions and cations compare with their parent atoms?

12. How do the sizes of monatomic (one-atom) cations and anions compare in the

same period?

Section 4.4. Fundamentals of Covalent Bonding

13. Define covalent bond.

14. Explain the energy minimum in the diagram illustrating the H-H covalent bond

in Figure 4.9.

Section 4.5. Covalent Bonds in Compounds

15. What may be said about the liklihood of H atoms being involved in double

covalent bonds?

16. What is represented by a dashed line, –, in a chemical formula?

17. What is represented by the two dots in the formula of phosphine, PH3, below:

H

H P:

H

Section 4.6. Some other Aspects of Covalent Bonding

18. What are multiple bonds?

multiple bonds?



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Which three elements are most likely to form



19. What can be said about the nature of covalent bonds between (a) 2 atoms with

almost identical electronegativity values and (b) 2 atoms with substantially

different electronegativity values?

20. What symbols are used to show bond polarity?

21. What is the “ultimate” in polar bonds?

22. A molecule of NH3 will combine with one of BF3. Describe the kind of bond

formed in the resulting compound.

23. What is required for an atom to form a compound with more than 8 electrons in

the central atom’s outer shell?

24. What are resonance structures?

25. How many total valence electrons are in the nitrate ion, NO3 ? What are the

resonance structures of this ion?

Section 4.7. Chemical Formulas of Compounds

26. Summarize the steps involved in calculating the percentage composition of a

compound from its formula.

27. Phosgene, COCl2, is a poisonous gas that was used for warfare in World War I.

What is its percentage composition?

28. The molecular formula of acetylsalicylic acid (aspirin) is C9H8O4. What is its

percentage composition?

29. Hydrates are compounds in which each formula unit is associated with a definite

number of water molecules. A typical hydrate is copper (II) sulfate

pentahydrate,

CuSO4•5H2O. The water of hydration can be driven off by heating, leaving the

anhydrous compound. Answer the following pertaining to CuSO4•5H2O: (a)

Mass of 1 mole of the compound, (b) mass of H2O in 1 mole of CuSO4•5H2O,

(c) percentage of H2O in CuSO4•5H2O.

30. What are the percentages of oxygen in (a) perchloric acid, (b) chloric acid, (c)

chlorous acid, and (d) hypochlorous acid (these acids were discussed in this

chapter).

31. What is the simplest (empirical) formula of sodium oxalate, Na2C2O4?

32. What is the formula of dichlorine heptoxide? What is its percentage composition?

33. Chlorine dioxide, ClO2, is used as a substitute for chlorine gas in the disinfection

of drinking water. What is the percentage composition of ClO2?

34. Summarize in steps the calculation of empirical formula from the percentage

composition of a compound.



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35. If the empirical formula and the actual formula mass (molecular mass) of a

compound are known, how is the true formula calculated?

36. A chlorofluorocarbon gas is 9.93% C, 31.43% F, and 58.64% Cl. What is its

empirical formula?

37. A 100.0 g portion of the chlorofluorocarbon gas (preceding question) was found

to occupy 25.3 L at 100°C and 1.000 atm pressure. Assuming that the gas

behaved ideally and using the ideal gas equation, how many moles of the gas are

in 100.0 of Freon? What is its molar mass? What is its molecular formula? (To

answer this question it may be necessary to refer back to a discussion of the gas

laws in Section 2.6).

38. A compound is 5.88% H and 94.12% O. What is its empirical formula?

39. The compound from the preceding problem has a molecular mass of 34.0. What

is its molecular formula?

40. A pure liquid compound with an overpowering vinegar odor is 40.0% C, 6.67%

H, and 53.3% O. Its molecular mass is 60.0. What is its empirical formula?

What is its molecular formula?

41. A compound is 29.1% Na, 40.5% S, and 30.4% O. It has a formula mass of

158.0. Fill out the table below pertaining to the compound, give its true formula,

and give its actual formula. It is an ionic compound. What is the anion?

Grams of element

Element in 100 g compound



Mol element

Mol element in

100 g compound mol element with fewest moles



Na

S

O



(d)

(e)

(f)



(a)

(b)

(c)



(g)

(h)

(i)



42. An ionic compound is 41.7% Mg, 54.9% O, and 3.4% H. What is its empirical

formula? Considering the ions in Table 3.6, what is the actual formula?

43. Ethylenediamine is 40.0% C, 13.4% H, and 46.6% N; its formula mass is 60.1.

What are its empirical and molecular formulas?

44. The empirical formula of butane is C2H5 and its molecular mass is 58.14. What

is the molecular formula?

General Questions

45. Suppose that you were asked to give the Lewis formula of formic acid, H2CO2,

where the atomic number of H is 1, that of C is 6, and that of O is 8. The total

number of valence (outer shell) electrons that would have to be placed correctly

in the structure is

A. 6

C. 20

E. 16

B. 18

D. 7



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46. Given the Lewis symbols of hypothetical elements X and Z,



X



Z



The compound most likely formed by these two elements would be

A.

B.

C.

D.

E.



Covalently bound compound formula X2Z 3.

Covalently bound compound formula XZ.

Ionic compound formula X2Z.

Ionic compound XZ 2.

Ionic compound XZ.



47. Of the following, the statement that is untrue regarding ionic NaCl and its

formation from gaseous Na and Cl atoms is

A. Energy called electron repulsion is consumed in putting an electron on a Cl

atom to produce a Cl- ion.

B. Energy called ionization energy is required to remove an electron from a Na

atom to produce a Na+ ion.

C. A relatively large amount of energy (lattice energy) is released when the Na+

and Cl- ions come together to form crystalline NaCl.

D. A particular Cl- ion in the crystal of NaCl has 6 Na+ ions as its nearest

neighbors.

E. Every ion in the crystal of NaCl is closest to ions of opposite charge,

resulting in forces of attraction that account for the stability of ionic bonds.

48. Consider the single, double, and triple bonds connecting the 2 carbon atoms in

the 3 compounds C2H6, C 2H4, and C2H2. Of the following pertaining to these

bonds, the untrue statement is

A. The C=C bond is shorter than the C–C bond.

B. The C≡C bond is stronger than the C=C bond.

C. Because it must accomodate a total of 6 electrons rather than 2, the C≡C

bond is longer than the C–C bond.

D. As bond multiplicity increases, the bond strength increases and the bond

length decreases.

E. The bonds can act like springs in that the atoms connected by the bonds

vibrate when exposed to the right wavelength of infrared radiation.

49. Remember that O and S atoms all have the same number of valence electrons.

Considering the structures (Lewis formulas),



O



S



O



O



S



O



the true statement is

A. Only one of these structures can be correct.

B. Both of these structures are incorrect.

C. The structure on the left shows the incorrect number of outer shell electrons

for the O atom on the left.

D. The structures are equivalent resonance structures.

E. The text showed the sulfur atom on one end, not in the middle.



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Manahan, Stanley E. "CHEMICAL REACTIONS, EQUATIONS, AND STOICHIOMETRY"

Fundamentals of Environmental Chemistry

Boca Raton: CRC Press LLC,2001



5 CHEMICAL REACTIONS, EQUATIONS,

AND STOICHIOMETRY



__________________________

5.1 THE SENTENCES OF CHEMISTRY



As noted earlier, chemistry is a language. Success in the study of chemistry

depends upon how well chemical language is learned. This chapter presents the last

of the most basic parts of the chemical language. When it has been learned, the

reader will have the essential tools needed to speak and write chemistry and to apply

it in environmental and other areas.

Recall that the discussion of chemical language began by learning about the

elements, the atoms composing the elements, and the symbols used to designate

these elements and their atoms. Atoms of the elements bond together in various

combinations to produce chemical compounds. These are designated by chemical

formulas consisting of symbols for the kinds of atoms in the compound and

subscripts indicating the relative numbers of atoms of each kind in the compound. In

chemical language, the symbols of the elements are the letters of the chemical

alphabet and the formulas are the words of chemistry.



Chemical Reactions and Equations: The Sentences of the Chemical

Language

The formation of chemical compounds, their decomposition, and their

interactions with one another fall under the category of chemical reactions.

Chemical reactions are involved in the annual production of millions of kilograms of

industrial chemicals, bacterially mediated degradation of water pollutants, the

chemical analysis of the kinds and quantities of components of a sample, and

practically any other operation involving chemicals. To a very large extent,

chemistry is the study of chemical reactions expressed on paper as chemical

equations. A chemical equation is a sentence of chemistry, made up of words

consisting of chemical formulas. A sentence should be put together according to

rules understood by all those literate in the language. The rules of the chemical



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