1. How can you differentiate among reactions in aqueous solutions?
An aqueous solution consists of water that contains dissolved substances. Precipitation reactions and acid-base reactions are two examples of the types of reactions that occur in aqueous solution.
Precipitation Reactions When a cation and an anion in solution form a compound that is insoluble in water, the product becomes visible as a solid, or precipitate. For example, in the double-replacement reaction below, potassium iodide and lead(II) nitrate form insoluble lead(II) iodide:
2KI(aq) + Pb(NO3)2(aq)  PbI2(s) + 2KNO3(aq)
Use Table 1, which follows on Slide 2, to determine the solubility of many compounds.
(contd.)

2. (contd.) Table 1 Compounds Solubility Exceptions
Salts of alkali metals and ammonia
Soluble
Some lithium compounds
Nitrate salts and chlorate salts
Few exceptions
Sulfate salts
Compounds of Pb, Ag, Hg, Ba, Sr, and Ca
Chloride salts
Compounds of Ag and some compounds of Hg and Pb
Carbonates, phosphates, chromates, sulfides, and hydroxides
Most are insoluble
Compounds of the alkali metals and of ammonia
(contd.)

3. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Acid-Base Reactions In general, an acid-base reaction in aqueous solution produces a salt and water. The salt forms from the cation of the base and the anion of the acid. For example,
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Many salts are soluble in water, so the reaction between an acid and a base may or may not produce a precipitate. In the above reaction, sodium chloride (NaCl) does not form a precipitate.
1. Differentiate Which of the following reactions is an acid-base reaction? Which is a precipitation reaction? Explain your answers.
a. HCl(aq) + AgNO3(aq)  HNO3(aq) + AgCl(s)
b. 2HCl(aq) + Ca(OH)2(aq)  CaCl2(aq) + 2H2O(l)

4. What are oxidation reduction reactions?
Oxidation-reduction reactions involve the oxidation of a reactant and the reduction of another reactant.
Early chemists viewed oxidation as the gain of oxygen by an element; reduction was viewed as the loss of oxygen by a compound:
Oxidation: A gain of oxygen
Reduction: A loss of oxygen
Modern chemists define oxidation-reduction reactions in terms of a shift of electrons between reactants:
Oxidation: A gain of oxygen OR a loss of electrons
Reduction: A loss of oxygen OR a gain of electrons
Oxidation-reduction reactions are also called redox reactions.
(contd.)

5. Magnesium metal can react with sulfur, forming magnesium sulfide
Magnesium metal can react with sulfur, forming magnesium sulfide. In this reaction, magnesium atoms lose two electrons and sulfur atoms gain two electrons:
Magnesium is oxidized, and sulfur is reduced.
(contd.)

6. Oxidizing and Reducing Agents A substance that undergoes oxidation loses electrons. It is also the reducing agent. A substance that undergoes reduction gains electrons. It is also the oxidizing agent.
In the reaction between magnesium and sulfur:
Magnesium is oxidized and is the reducing agent.
Sulfur is reduced and is the oxidizing agent.
In another example, silver nitrate reacts with copper to form copper(II) nitrate and silver. The equation may be written in ionic form:
2Ag++ 2NO3– + Cu  Cu2+ + 2NO3– + 2Ag
Copper is oxidized and silver is reduced.
Cu  Cu2+ + 2e– (loss of electrons)
2Ag++ 2e– 2Ag (gain of electrons)
Copper is the reducing agent and silver is the oxidizing agent.
(contd.)

7. Oxidation and Reduction With Covalent Compounds In reactions involving covalent compounds, a complete electron transfer does not occur.
Instead, there is a shift of electrons toward the more electronegative atom. Some atoms have a partial gain of electrons; others have a partial loss of electrons.
For carbon compounds, the addition of oxygen or the removal of hydrogen is always oxidation.
2. Identify In the following reaction, identify the substance that is oxidized and the substance that is reduced: H2(g) + O2(g)  H2O(l)
3. Identify Name the oxidizing agent and the reducing agent in the reaction in Question 2.

8. How are oxidation numbers assigned?
An oxidation number is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction.
In general, a bonded atom’s oxidation number is the charge that it would have if the electrons in the bond were assigned to the atom of the more electronegative element.
Oxidation numbers change during redox reactions.
You can use changes in oxidation number to balance equations of complex redox reactions.
Table 2, which follows on Slide 9, summarizes how to assign oxidation numbers.
(contd.)

9. Rules for Assigning Oxidation Numbers
Table 2
Rules for Assigning Oxidation Numbers
1. The oxidation number of a monatomic ion is equal in magnitude and sign to its ionic charge. For example, the oxidation number of (Br1–) is –1.
2. The oxidation number of hydrogen in a compound is +1, except in metal hydrides, such as NaH, where it is –1.
3. The oxidation number of oxygen in a compound is –2, except in peroxides—such as H2O2—where it is –1, and in compounds with the more electronegative fluorine, where it is positive.
4. The oxidation number of an atom in uncombined (elemental) form is 0.
5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0.
6. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.
(contd.)

10. Sample Problem: Assigning Oxidation Numbers to Atoms
What is the oxidation number of each atom in SO2?
Analyze Identify the relevant concepts. Use the set of rules in Table 2 to assign and calculate oxidation numbers.
Solve Apply concepts to this situation. In SO2, there are two oxygen atoms, and each one has an oxidation number of –2 (Rule 3). The sum of the oxidation numbers for the neutral compound must be 0 (Rule 5). Therefore, the oxidation number of sulfur is +4, because: +4 + (2 x (–2) = 0.
Oxidation numbers are often written above the chemical symbols in a formula. For SO2, the oxidation numbers would be written as shown to the right.

11. How do oxidation numbers change in chemical reactions?
In any redox reaction, the oxidation numbers of reacting species change. One reactant is oxidized and another is reduced. For example, look again at the equation for the reaction of silver nitrate and copper:
In this reaction, the oxidation number of Ag decreases from +1 to 0. It gains an electron and is reduced to silver metal.
The oxidation number of copper increases from 0 to +2. It loses two electrons and is oxidized.
Oxidation is an increase in oxidation number.
Reduction is a decrease in oxidation number.
(contd.)

12. Processes Leading to Oxidation and Reduction
Table 3 summarizes the processes of oxidation and reduction.
Table 3
Processes Leading to Oxidation and Reduction
Oxidation
Reduction
Complete loss of electrons (ionic reactions)
Complete gain of electrons (ionic reactions)
Shift of electrons away from an atom in a covalent bond
Shift of electrons toward an atom in a covalent bond
Gain of oxygen
Loss of oxygen
Loss of hydrogen by a covalent compound
Gain of hydrogen by a covalent compound
Increase in oxidation number
Decrease in oxidation number
4. Review How does oxidation compare with reduction in terms of electrons? In terms of oxidation number?

13. How do you identify oxidation-reduction reactions?
Oxidation-reduction reactions always involve the transfer of electrons from one reacting species to another. Examples of oxidation-reduction reactions include:
Single-replacement reactions
Combination reactions
Decomposition reactions
Combustion reactions
Double-replacement reactions and acid-base reactions are not redox reactions.
(contd.)

14. Sample Problem: Identifying Redox Reactions
Use the change in oxidation numbers to identify whether the following reaction is a redox reaction or a reaction of some other type. If the reaction is a redox reaction, identify the element reduced, the element oxidized, the reducing agent, and the oxidizing agent.
Cl2(g) + 2NaBr(aq)  2NaCl + Br2(g)
Analyze Identify the relevant concepts. If changes in oxidation number occur, the reaction is a redox reaction. The element whose oxidation number increases is oxidized and is the reducing agent. The element whose oxidation number decreases is reduced and is the oxidizing agent.
Solve Apply concepts to this situation. Assign oxidation numbers. Then interpret the change (or lack of change) in oxidation numbers to identify if the reaction is a redox reaction.
(contd.)

15. 3KOH(aq) + H3PO4(aq)  K3PO4(s) + 3H2O(l)
The oxidation number of chlorine changes from 0 to –1. Chlorine is reduced. The oxidation number of bromine changes from –1 to 0. The bromide ion is oxidized. Chlorine is the oxidizing agent; the bromide ion is the reducing agent.
5. Interpret Data Is the following reaction a redox reaction? Why or why not?
3KOH(aq) + H3PO4(aq)  K3PO4(s) + 3H2O(l)

16. How do you balance redox reactions?
One way to balance redox reactions is to use the oxidation-number-change method. The following steps are used in this method:
Write the skeleton equation for the reaction.
Assign oxidation numbers to all the atoms in the equation.
Identify which atoms are oxidized and which are reduced.
Use a bracketed line to connect oxidized atoms that are oxidized and another line to connect atoms that are reduced. Write the oxidation-number change on the line. For example, the equation for the process used to obtain metallic iron from iron ore in a blast furnace would appear as follows:
(contd.)

17. Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
Make the total increase in oxidation number equal to the total decrease by using coefficients.
Balance the rest of the equation by inspection (if necessary).
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
6. Differentiate Differentiate each of the following unbalanced equations as an acid-base reaction, a precipitation reaction, or an oxidation-reduction reaction. (Some reactions may have more than one answer.) Balance each equation.
a. Na2CO3(aq) + MgCl2(aq)  NaCl(aq) + MgCO3(s)
b. HNO2(aq) + HI(aq)  NO(g) + I2(s) + H2O(l)
c. KOH(aq) + HF(aq)  KF(aq) + H2O(l)