1. Chapter 4
The Major Classes of Chemical Reactions

2. Major Classes of Chemical Reactions
4.1 The Role of Water as Solvent
4.2 Writing Equations for Aqueous Ionic Reactions
4.3 Precipitation Reactions
4.4 Acid-Base Reactions
4.5 Oxidation-Reduction (Redox) Reactions
4.6 Elemental Substances in Redox Reactions
4.7 Reversible Reactions: An Introduction to Chemical
Equilibrium

3. Electrical conductivity of ionic solutions
Figure 4.1

4. Definition: An electrolyte
A substance that conducts a current when dissolved in water
Soluble ionic compounds dissociate completely into ions
in water and create a large current; they are called strong
electrolytes.
H2O
KBr(s) K+(aq) + Br-(aq)

5. Definition: A non-electrolyte
Substances that dissolve in water but do not dissociate
into ions; they remain as intact molecules; their solutions do
not conduct an electric current; they are mainly polar covalent
compounds.
Examples include sucrose, ethanol, ethylene glycol

6. Determining Moles of Ions in Aqueous Ionic Solutions
Sample Problem 4.1
PROBLEM:
How many moles of each ion are present in the following solutions?
(a) 5.0 mol of ammonium sulfate dissolved in water
(b) g of cesium bromide dissolved in water
(c) x 1022 formula units of copper(II) nitrate dissolved in water
(d) 35 mL of 0.84 M zinc chloride
PLAN:
Consider molecular formulas and the ionization of salts in water.
SOLUTION:
(a) (NH4)2SO4(s) NH4+(aq) + SO42-(aq)
2 mol NH4+
1 mol (NH4)2SO4
5.0 mol (NH4)2SO4 x
= 10. mol NH4+
5.0 mol SO42-

7. Sample Problem 4.1
(continued)
(b) CsBr(s) Cs+(aq) + Br-(aq)
= mol Cs+
mol CsBr
212.8 g CsBr
78.5 g CsBr x
= mol CsBr
= mol Br-
(c) Cu(NO3)2(s) Cu2+(aq) + 2NO3-(aq)
mol Cu(NO3)2
6.022 x 1023 formula units
7.42 x 1022 formula units Cu(NO3)2 x
= mol Cu(NO3)2
= mol Cu2+
(d) ZnCl2(s) Zn2+(aq) + 2Cl-(aq)
= mol NO3-
1 L
103 mL
0.84 mol ZnCl2 L
35 mL ZnCl2 x x
= 2.9 x 10-2 mol ZnCl2
= 2.9 x 10-2 mol Zn2+
= 5.8 x 10-2 mol Cl-

8. Electron distribution in H2 and H2O
Figure 4.2
Electron distribution in H2 and H2O
Water is an example of a polar molecule.

9. Dissolution of an ionic compound in H2O solvent
Figure 4.3

10. The hydrated proton
Figure 4.4

11. Determining the Molarity of H+ Ions in Aqueous Solutions of Acids
Sample Problem 4.2
PROBLEM:
Nitric acid is a major chemical in the fertilizer and explosives industries. In aqueous solution, each molecule dissociates and the H becomes a solvated H+ ion. What is the molarity of H+(aq) in 1.4 M nitric acid?
PLAN:
Use a balanced chemical equation and the definition of molarity (moles/liter) to find the answer.
SOLUTION:
The nitrate polyatomic ion is NO3-. Nitric acid is HNO3.
HNO3(l) H+(aq) + NO3-(aq)
1.4 M HNO3(aq) contains 1.4 mole of H+(aq)/L and is therefore 1.4 M in H+(aq)

12. A Precipitation Reaction
A reaction between Pb(NO3)2
and NaI, giving insoluble
yellow PbI2

13. Equations for Aqueous Ionic Reactions
Molecular equation: shows all reactants and products as if they
they were intact, undissociated compounds.
Total ionic equation: shows all the soluble ionic substances
dissociated into ions.
Net ionic equation: eliminates the spectator ions* and shows
the actual chemical change that occurs.
(*spectator ions: ions in the total ionic equation that are not
involved in the actual chemical change)

14. A precipitation reaction and its equation
Figure 4.5
A precipitation reaction and its equation

15. Table 4.1 Solubility Rules For Ionic Compounds in Water
Soluble Ionic Compounds
1. All common compounds of Group 1A ions (Li+, Na+, K+, etc.) and the ammonium ion (NH4+) are soluble.
2. All common nitrates (NO3-), acetates (CH3COO-) and most perchlorates (ClO4-) are soluble.
3. All common chlorides (Cl-), bromides (Br-) and iodides (I-) are soluble, except those of Ag+, Pb2+, Cu+, and Hg22+.
4. All common sulfates (SO42-) are soluble except those of Ca2+, Sr2+, Ba2+
and Pb2+.
Insoluble Ionic Compounds
1. All common metal hydroxides are insoluble, except those of Group 1A and the larger members of Group 2A beginning with Ca2+.
2. All common carbonates (CO32-) and phosphates (PO43-) are insoluble, except those of Group 1A and NH4+.
3. All common sulfides are insoluble except those of Group 1A, Group 2A and NH4+.

16. Predicting precipitation reactions and writing ionic equations
Sample Problem 4.3
PROBLEM:
Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction occurs, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions.
(a) sodium sulfate(aq) + strontium nitrate(aq)
(b) ammonium perchlorate(aq) + sodium bromide(aq)
SOLUTION:
PLAN:
(a) Na2SO4(aq) + Sr(NO3)2 (aq) NaNO3(aq) + SrSO4(s)
write ions
2Na+(aq) + SO42-(aq) + Sr2+(aq) + 2NO3-(aq)
2Na+(aq) + 2NO3-(aq) + SrSO4(s)
combine anions and cations
SO42-(aq)+ Sr2+(aq) SrSO4(s)
check for insolubility
(Table 4.1)
(b) NH4ClO4(aq) + NaBr (aq) NH4Br (aq) + NaClO4(aq)
eliminate spectator ions from net ionic equation
All reactants / products are soluble - no reaction occurs.

17. (Bronsted Acids and Bases)
Acid-Base Chemistry
(Bronsted Acids and Bases)
Definitions
An acid is a substance that produces H+ ions when dissolved
in water.
H2O
HX H+(aq) + X-(aq)
A base is a substance that produces OH- ions when
dissolved in water.
H2O
MOH M+(aq) + OH-(aq)

18. Types of Acids and Bases: Strong and Weak
Strong acids and strong bases dissociate completely into ions
when dissolved in water. The ionization reaction is represented
with a single one-way arrow in the forward direction.
Examples: HCl and NaOH
Weak acids and weak bases dissociate partly into ions when
dissolved in water, leaving a large fraction of the acid or base
intact. The ionization reaction is represented with a two-way arrow
indicating an equilibrium between associated and dissociated
(ionized) forms of the acid or base.
Examples: CH3COOH and NH3

19. Table 4.2 Selected Acids and Bases
Strong
Strong
hydrochloric acid, HCl
sodium hydroxide, NaOH
hydrobromic acid, HBr
potassium hydroxide, KOH
hydroiodic acid, HI
calcium hydroxide, Ca(OH)2
nitric acid, HNO3
strontium hydroxide, Sr(OH)2
sulfuric acid, H2SO4
barium hydroxide, Ba(OH)2
perchloric acid, HClO4
Weak
Weak
hydrofluoric acid, HF
ammonia, NH3
phosphoric acid, H3PO4
acetic acid, CH3COOH

20. An Acid-Base Reaction = Neutralization Reaction
One Definition: Occurs when an acid reacts with a base
in aqueous solution, yielding a salt solution and water as
products.
HX(aq) + MOH(aq) MX(aq) + H2O(l)
acid
base
salt
water

21. Sample Problem 4.4
Writing Ionic Equations for Acid-Base Reactions
PROBLEM:
Write balanced molecular, total ionic, and net ionic equations for the following two acid-base reactions and identify the spectator ions.
(a) strontium hydroxide(aq) + perchloric acid(aq)
(b) barium hydroxide(aq) + sulfuric acid(aq)
SOLUTION:
PLAN:
(a) Sr(OH)2(aq) + 2HClO4(aq) H2O(l) + Sr(ClO4)2(aq)
Reactants are strong acids and bases and are thus completely ionized in water.
Sr2+(aq) + 2OH-(aq) + 2H+(aq) + 2ClO4-(aq)
2H2O(l) + Sr2+(aq) + 2ClO4-(aq)
2OH-(aq)+ 2H+(aq) H2O(l)
Products are:
(b) Ba(OH)2(aq) + H2SO4(aq) H2O(l) + BaSO4(aq)
1. water
Ba2+(aq) + 2OH-(aq) + 2H+(aq) + SO42-(aq)
2H2O(l) + Ba2+(aq) + SO42-(aq)
2. spectator ions
2OH-(aq)+ 2H+(aq) H2O(l)

22. An Acid-Base Titration
A Titration: One solution of known concentration is
used to determine the concentration of another solution
through a monitored reaction.
Acid-Base Titration: A standardized solution of base (one
whose concentration is known) is added slowly to an acid
solution of unknown concentration. An indicator is used to
monitor the reaction; its color is different in acid and base.

23. An acid-base titration
Figure 4.7

24. Acid-Base Titrations: Definitions and Assumptions
Equivalence point: occurs when all of the moles of H+ ion
present in the original volume of acid solution have reacted
with an equivalent number of moles of OH- ions added
from the buret.
End point: occurs when a tiny excess of OH- ions changes
the indicator permanently to its color in base.
We assume that the amount of base needed to reach the
end point is equivalent to the amount needed to reach the
equivalence point.

25. Finding the Concentration of Acid from an Acid-Base Titration
Sample Problem 4.5
PROBLEM:
You perform an acid-base titration to standardize an HCl solution by placing mL of HCl in a flask with a few drops of indicator solution. You put M NaOH into a buret, and the initial reading is 0.55 mL. At the end-point, the buret reads mL. What is the concentration of the HCl solution?
PLAN:
SOLUTION:
volume(L) of base
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)
multiply by M of base
( ) mL x
1 L
103 mL
= L
mol of base
molar ratio L
x M
= x 10-3 mol NaOH
mol of acid
Molar ratio is 1:1
divide by L of acid
5.078 x 10-3 mol HCl
0.050 L
= M HCl
M of acid

26. An aqueous strong acid-strong base reaction on the atomic scale
Figure 4.8

27. Oxidation-Reduction (Redox) Reactions
Net movement of electrons from one reactant to another
Driving force: The movement of electrons occurs in a specific
direction, that is, from the reactant (or atom in the reactant) with
less attraction for electrons to the reactant with more attraction
for electrons.

28. The redox process in compound formation
Figure 4.10

29. Table 4.3 Rules for Assigning an Oxidation Number (O.N.)
General rules
1. For an atom in its elemental form (Na, O2, Cl2, etc.): O.N. = 0
2. For a monatomic ion: O.N. = ion charge
3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ion’s charge.
Rules for specific atoms or periodic table groups
1. For Group 1A:
O.N. = +1 in all compounds
2. For Group 2A:
O.N. = +2 in all compounds
3. For hydrogen:
O.N. = +1 in combination with nonmetals
4. For fluorine:
O.N. = -1 in combination with metals and boron
5. For oxygen:
O.N. = -1 in peroxides
O.N. = -2 in all other compounds(except with F)
6. For Group 7A:
O.N. = -1 in combination with metals, nonmetals (except O), and other halogens lower in the group

30. Sample Problem 4.6
Determining the Oxidation Number of an Element
PROBLEM:
Determine the oxidation number (O.N.) of each element in the following compounds:
(a) zinc chloride
(b) sulfur trioxide
(c) nitric acid
PLAN:
O.N.s of the ions in a polyatomic ion add up to the charge of the ion; O.N.s of the ions in the compound add up to zero.
SOLUTION:
(a) ZnCl2. The O.N. for zinc is +2 and that for chloride is -1.
(b) SO3. Each oxygen is an oxide with an O.N. of -2. Therefore, the O.N. of sulfur must be +6.
(c) HNO3. H has an O.N. of +1 and each oxygen is -2. Therefore, N must have an O.N. of +5.

31. Highest and lowest oxidation numbers of reactive main-group elements
Figure 4.11

32. Summary of terminology for oxidation-reduction (redox) reactions
Figure 4.12 e- X
transfer or shift of electrons Y
X loses electron(s)
Y gains electron(s)
X is oxidized
Y is reduced
X is the reducing agent
Y is the oxidizing agent
X increases its oxidation number
Y decreases its oxidation number

33. Sample Problem 4.7
Recognizing Oxidizing and Reducing Agents
PROBLEM:
Identify the oxidizing agent and reducing agent in each of the following:
(a) 2Al(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g)
(b) PbO(s) + CO(g) Pb(s) + CO2(g)
(c) 2H2(g) + O2(g) H2O(g)
PLAN:
Assign an O.N. to each atom and determine which gained and which lost electrons in going from reactants to products. An increase in O.N. is associated with the oxidized species (the reducing agent) and a decrease in O.N. is associated with the reduced species (oxidizing agent).
SOLUTION: +1 +6 -2 +3 +6 -2
(a) 2Al(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g)
The O.N. of Al increases; it is oxidized; it is the reducing agent.
The O.N. of H decreases; it is reduced; it is the oxidizing agent.

34. Sample Problem 4.7
(continued) +2 -2 +2 -2 +4 -2
(b) PbO(s) + CO(g) Pb(s) + CO2(g)
The O.N. of C increases; it is oxidized; it is the reducing agent.
The O.N. of Pb decreases; it is reduced; it is the oxidizing agent. +1 -2
(c) 2H2(g) + O2(g) H2O(g)
The O.N. of H increases; it is oxidized; it is the reducing agent.
The O.N. of O decreases; it is reduced; it is the oxidizing agent.

35. Balancing Redox Equations by the Oxidation Number Method
Sample Problem 4.8
PROBLEM:
Use the oxidation number method to balance the following equations:
(a) Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l)
(b) PbS(s) + O2(g) PbO(s) + SO2(g)
SOLUTION: +1 +5 -2 +2 +5 -2 +4 -2 +1 -2
(a) Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l)
O.N. of Cu increases (it loses 2e-); it is oxidized and is the reducing agent.
O.N. of N decreases (it gains 1e-); it is reduced and is the oxidizing agent.
loses 2e-
balance other ions
Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l)
gains 1e-
x2 to balance e-
balance unchanged polyatomic ions

36. Cu(s) + 2HNO3(aq) Cu(NO3)2(aq) + 2NO2(g) + H2O(l)
Sample Problem 4.8
(continued)
Cu(s) + 2HNO3(aq) Cu(NO3)2(aq) NO2(g) H2O(l)
Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) NO2(g) H2O(l)
Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) NO2(g) H2O(l) +2 -2 +2 -2 +4 -2
(b) PbS(s) + O2(g) PbO(s) + SO2(g)
loses 6e-
PbS(s) O2(g) PbO(s) + SO2(g)
3/2
gains 2e- per O; need 3/2 O2 to make 3O2-
Multiply by 2 to give whole number coefficients.
2PbS(s) + 3O2(g) PbO(s) + 2SO2(g)

37. A redox titration
Figure 4.13

38. Finding an Unknown Concentration by a Redox Titration
Sample Problem 4.9
PROBLEM:
Calcium ion (Ca2+) is required for blood to clot and for many other cell processes. An abnormal Ca2+ concentration is indicative of disease. To measure the Ca2+ concentration, 1.00 mL of human blood was treated with Na2C2O4 solution. The resulting CaC2O4 precipitate was filtered and dissolved in dilute H2SO4. This solution required 2.05 mL of 4.88 x 10-4M KMnO4 to reach the end-point. The unbalanced equation is
KMnO4(aq) + CaC2O4(s) + H2SO4(aq)
MnSO4(aq) + K2SO4(aq) + CaSO4(s) + CO2(g) + H2O(l)
(a) Calculate the amount (in moles) of Ca2+.
(b) Calculate the amount (in moles) of Ca2+ ion concentration expressed in units of mg Ca2+/100 mL blood.
PLAN:
volume of KMnO4 soln
mol of Ca2+
multiply by M
ratio of elements in formula
mol of KMnO4
mol of CaC2O4
molar ratio

39. Sample Problem 4.9
(continued)
SOLUTION: L
103 mL
4.88 x 10-4 mol KMnO4 L
2.05 mL soln x x
= 1.00 x 10-6 mol KMnO4
5 mol CaC2O4
2 mol KMnO4
1.00 x 10-6 mol KMnO4 x
= 2.50 x 10-6 mol CaC2O4
1 mol Ca2+
1 mol CaC2O4
2.50 x 10-6 mol CaC2O4 x
= 2.50 x 10-6 mol Ca2+
2.50 x 10-6 mol Ca2+
1mL blood
2.50 x 10-4 mol Ca2+
100 mL blood
x 100 =
40.08 g Ca2+
mol Ca2+ mg
10-3g
2.50 x 10-4 mol Ca2+
100 mL blood x x
= 10.0 mg Ca2+/100 mL blood

40. Elemental Substances in Redox Reactions
Three types of reactions:
(1) Combination reactions: X + Y Z
(2) Decomposition reactions: Z X + Y
(3) Displacement reactions: X + YZ XZ + Y

41. Combining elements to form an ionic compound
Figure 4.14

42. Decomposing a compound to its elements
Figure 4.15

43. Displacing one metal with another
Figure 4.18

44. cannot displace H2 from any source
strength as reducing agents
can displace H2
from water Li K Ba Ca Na Mg Al Mn An Cr Fe Cd Co Ni Sn Pb H2 Cu
Hg Ag Au
can displace H2
from steam
The activity series
of the metals
can displace H2
from acid
cannot displace H2 from any source
Figure 4.19

45. Identifying the Type of Redox Reaction
Sample Problem 4.10
Identifying the Type of Redox Reaction
PROBLEM:
Classify each of the following redox reactions as a combination, decomposition, or displacement reaction, write a balanced molecular equation for each, as well as total and net ionic equations for part (c), and identify the oxidizing and reducing agents:
(a) magnesium(s) + nitrogen(g) magnesium nitride (aq)
(b) hydrogen peroxide(l) water(l) + oxygen gas
(c) aluminum(s) + lead(II) nitrate(aq) aluminum nitrate(aq) + lead(s)
PLAN:
Combination reactions produce fewer products than reactants.
Decomposition reactions produce more products than reactants.
Displacement reactions have the same number of products and reactants.

46. Sample Problem 4.10
(continued) +2 -3
(a) Combination
Mg(s) N2(g) Mg3N2 (aq) 3
Mg is the reducing agent; N2 is the oxidizing agent. +1 -1 +1 -2
(b) Decomposition
H2O2(l) H2O(l) O2(g)
1/2 or
2 H2O2(l) H2O(l) O2(g)
O is the oxidizing and reducing agent.
(c) Displacement +2 +5 -2 +3 +5 -2
Al(s) + Pb(NO3)2(aq) Al(NO3)3(aq) + Pb(s)
2Al(s) + 3Pb(NO3)2(aq) Al(NO3)3(aq) + 3Pb(s)
Pb is the oxidizing and Al is the reducing agent.

47. The equilibrium state
Figure 4.20