Aqueous solutions

Electrolytes are substances which conduct electricity when dissolved in water (in aqueous solution).

· To conduct electricity, these substances must produce ions in aqueous solutions.

· Substances can be classified as non-electrolytes, weak electrolytes, and strong electrolytes

· Non-electrolytes do not produce ions and do not conduct electricity in solution

· Strong electrolytes completely dissociate or ionize in solution, and conduct electricity well.

H₂O

HCl (g) à H⁺ (aq) + Cl^- (aq)

· Weak electrolytes partially ionize in solution and conduct electricity poorly.

H₂O

CH₃COOH (aq) _ß^à H⁺ (aq) + CH₃COO^- (aq)

Strong electrolytes consist of three major classes:

1. Strong acids

2. Strong soluble bases

3. Soluble salts

Acids, bases, and salts:

· Acids are substances that produce H⁺ ions in aqueous solution.

· Bases are substances that produce OH^- ions in aqueous solution

· Salts are ionic compounds, with cations other than H⁺ and anions other than OH^- or O^2-.

Strong acids completely ionize in solution

· Common strong acids:

Hydrochloric acid	HCl
Hydrobromic acid	HBr
Hydroiodic acid	HI
Sulfuric acid	H2SO4
Nitric acid	HNO3
Perchloric acid	HClO4
Chloric acid	HClO3

Weak acids only partially ionize in solution

· Represented by the reversible reaction with the double arrow.

Strong bases may or may not completely dissolve in water

· Soluble strong bases consist of the hydroxides of group IA and calcium, strontium, and barium from group IIA

Weak bases only partially ionize in solution, a common example is ammonia

NH₃ (aq) + H₂O (l) _ß^à NH₄⁺ (aq) + OH^- (aq)

Salts may be soluble or insoluble in solution – see solubility rules in Table 4.1 on page 152.

· Which of the following would be soluble? Calcium nitrate, calcium carbonate, calcium chloride, silver nitrate, silver carbonate, silver chloride

Net ionic equations

Solubility rules and knowledge of strong acids and strong soluble bases are needed to write net ionic equations.

· In a net ionic equation, any substance which fully dissociates/ionizes in aqueous solution is written as ions.

These include strong acids, strong soluble bases, and soluble salts.

· Common ions on the reactant and product sides cancel. They are referred to as spectator ions because they do not participate in the reaction.

Ex. H₂SO₄ + 2 NaOH à 2 H₂O + Na₂SO₄

Molecular equation:

H₂SO₄ (aq) + 2 NaOH (aq) à 2 H₂O (l) + Na₂SO₄ (aq)

Complete ionic equation:

2 H⁺ (aq) + SO₄^2- (aq) + 2 Na⁺ (aq) + 2 OH^- (aq) à 2 H₂O (l) + 2 Na⁺ (aq) + SO₄^2- (aq)

Net ionic equation:

2 H⁺ (aq) + 2 OH^- (aq) à 2 H₂O (l)

Precipitation reactions

Precipitation Reactions: A reaction between two salt solutions which results in the formation of an insoluble ionic compound.

· Precipitation reactions are also classified as double replacement/double displacement reactions.

· The cations or anions from the reactants switch places.

· The driving force is the fact that one product is insoluble, so these cations and anions are removed from solution when they precipitate.

Ex. Na₂SO₄ + BaCl₂ à BaSO₄ + 2 NaCl

Complete ionic equation:

2 Na⁺ _{(aq) +} SO₄^-2_(aq) + Ba⁺²_(aq) + 2 Cl^-_(aq) à 2 Na⁺_(aq) + 2 Cl^-_(aq) + BaSO_4(s)

Net ionic equation:

SO₄^-2_(aq) + Ba⁺²_(aq) à BaSO_4(s)

Complete the following precipitation reaction and write a net ionic equation:

· Aqueous calcium chloride reacts with aqueous potassium phosphate.

Acid Base Reactions

Neutralization reactions occur when an acid reacts with a base to produce water and a salt

· H⁺ ions from the acid react with OH^- ions from the base to produce the water

· The salt is produced from the anions from the acid and the cations from the base – most salts produced are soluble in water.

2 HCl(aq) + Ca(OH)₂ (aq) à 2 H₂O (l) + CaCl₂ (aq)

Net ionic equation:

Note: the net ionic equation for all neutralization reactions involving strong acids and strong soluble bases is the one shown above.

A weak acid with a strong soluble base: CH₃COOH (aq) + NaOH (aq) à NaCH₃COO (aq) + H₂O (l)

Net ionic equation:

Complete the following neutralization reactions and write net ionic equations:

Sulfuric acid reacts with sodium hydroxide

Carbonic acid reacts with potassium hydroxide

Carbonate and bicarbonate salts can also act as bases.

Write the molecular equation and the net-ionic equation for the reaction which would occur between sulfuric acid and calcium carbonate.

Molecular equation:

H₂SO₄(aq) + CaCO₃(s) à CaSO₄(aq) + H₂O(l) + CO₂(g)

Note: H₂CO₃ breaks down into H₂O and CO₂ when produced in an exothermic reaction.

This reaction is also referred to as a gas-formation reaction because the gas which is produced is quickly removed from solution, acting as a driving force for the reaction.

Net ionic equation:

Redox Reactions

Oxidation-reduction (redox) reactions involve a transfer of electrons.

· Oxidation numbers/states are used to keep track of electron transfers.

· Oxidation numbers are used to describe the number of electrons an atom has lost, gained, or shared when bonding in a compound

Rules to Determine Oxidation Numbers:

1. The oxidation number of any atom in its elemental state is zero (0). Ex. H₂, O₂, Fe, P₄

2. The oxidation number of a monatomic ion is equal to its charge. Ex. Fe²⁺, O^2-

3. F is assigned an oxidation number of –1 when combined with other elements.

4. Cl, Br, and I are assigned oxidation numbers of –1 except when combined with O or F.

5. H is usually assigned an oxidation number of +1 when combined with other elements. Exception: for hydrides (H combined with a metal) H is assigned –1. Ex. HCl, NaH

6. O is assigned an oxidation number of –2 when combined with other elements. Except in OF₂, peroxides (O₂^2- ion), and superoxides (O₂^- ion). Ex. H₂O, CO₂, OF₂

7. The sum of all oxidation numbers within a species is equal to the overall charge on the species. Ex. NO₃^-

Determine the oxidation number of each atom/ion in the following species.

CO

H₂SO₄

H₂O₂

Ag⁺

CuCl₂

Oxidation numbers are used to describe the number of electrons an atom has lost, gained, or shared when bonding in a compound

· When losing electrons, an atom is assigned a positive oxidation number

· When gaining electrons, an atom is assigned a negative oxidation number

· When sharing electrons, the more electronegative atom is thought of as gaining through sharing and the less electronegative atom is thought of as losing through sharing

· If an atom is sharing with an identical atom, no net gain or loss occurs

· Oxidation numbers take into account all bonds an atom is involved in

Compare this idea to the rules for oxidation #’s:

2 Mg + O₂ à 2 MgO

2 Al + 3 CuCl₂ à 2 AlCl₃ + 3 Cu

2 CH₃OH + 3 O₂ à 2 CO₂ + 4 H₂O

In reactions (comparing reactant atoms to product atoms):

· Oxidation occurs when there is an increase in oxidation number (an apparent loss of electrons)

· Reduction occurs when there is a decrease in oxidation number (an apparent gain of electrons)

Oxidizing and Reducing Agents

Oxidation is always accompanied by reduction.

· Substances in redox reactions that undergo oxidation are called reducing agents. As they lose electrons, they give them to something else – thus accomplishing a reduction of some other atom/ion.

· Substances that become reduced are called oxidizing agents. As they gain electrons, they take them away from something else – thus accomplishing an oxidation of some other atom/ion.

Identify the oxidizing and reducing agents in these reactions.

2 Mg + O₂ à 2 MgO

2 Al + 3 CuCl₂ à 2 AlCl₃ + 3 Cu

2 CH₃OH + 3 O₂ à 2 CO₂ + 4 H₂O

Writing a net ionic equation for some redox reactions helps to focus on the transfer of electrons.

2 Al + 3 CuCl₂ à 2 AlCl₃ + 3 Cu

Note the obvious transfer of electrons as aluminum atoms are transformed into aluminum ions, and as copper ions are transformed into copper atoms.

What observations would be notable as this reaction occurred?

Single displacement reactions -- reactions in which one element displaces another from a compound.

· These are always redox reactions.

· Active metals displace less active metals or hydrogen from compounds in aqueous solution

· The result is an oxidized form of the more active metal and the reduced form (elemental state) of the less active metal or hydrogen

· Relative activities are shown in the activity series on page 140.

Examine these reactions in light of the activity series:

2 Al + 3 CuCl₂ à 2 AlCl₃ + 3 Cu

2 AlCl₃ + 3 Cu à 2 Al + 3 CuCl₂

Would both occur in aqueous solution? Why or why not?

Since hydrogen is included in the activity series, we can also consider how non-oxidizing acids (HCl, H₂SO₄) react with more active metals.

Ex. Zn (s) + H₂SO₄ (aq) à

Very active metals can displace hydrogen from water. Ex. Na (s) + H₂O (l) à NaOH (aq) + H₂ (g)

· These reactions are also exothermic enough that they result in ignition of the flammable hydrogen gas.

Non-metals can also displace other non-metals.

· The activity series for the halogens is:

I₂ < Br₂ < Cl₂ < F₂

Complete the following reactions and write net ionic equations:

Cl₂ (g) + NaI (aq) à

I₂ (g) + NaF (aq) à

Complete the following reactions in aqueous solution and write net-ionic equations for each. Determine whether each is precipitation, acid-base, or redox.

CuSO₄ + Na₂CO₃ à

HClO₃ + Ca(OH)₂ à

H₂SO₄ + Ni à

CH₃COOH + Ba(OH)₂ à

Naming ternary acids and salts

Ternary acids (oxoacids) are compounds of hydrogen, oxygen and a non-metal.

· Non-metals that exhibit more than one oxidation state can form more than one ternary acid, each with a different number of oxygen atoms.

· Suffixes –ous and –ic are used in the names.

· One oxidation state is (arbitrarily) assigned to have the –ic ending. (Ex. HNO₃ = nitric acid, H₂SO₄ = sulfuric acid, HClO₃ = chloric acid.)

· The –ous ending is used for the acid with one less oxygen than the –ic acid. (Ex. HNO₂ = nitrous acid.) Note: the oxidation number has decreased by 2.

· Prefixes hypo- and per- can also be used if more than two oxidation states are possible.

· Hypo- is used with –ous for an ox. state lower than –ous.

· Per- is used with –ic for an ox. state higher than –ic.

Ex. HClO₄ = perchloric acid

HClO₃ = chloric acid

HClO₂ = chlorous acid

HClO = hypochlorous acid

Ternary salts are compounds in which the hydrogen of a ternary acid has been replaced by a cation.

· Names for ternary salts involve polyatomic ions, with the names for the polyatomic ions based on the ternary acids from which they are derived.

· -ic endings are changed to –ate, and –ous endings are changed to –ite.

Ex. ClO₄^- = perchlorate

ClO₃^- = chlorate

ClO₂^- = chlorite

ClO^- = hypochorite

· In some cases, one or more acidic hydrogens may be retained by the anion. In these cases the word hydrogen or dihydrogen is added to the name of the polyatomic ion.

Ex. H₂SO₄ = sulfuric acid

HSO₄^- = hydrogen sulfate

SO₄^2- = sulfate

H₃PO₄ = phosphoric acid

H₂PO₄^- = dihydrogen phosphate

HPO₄^2- = hydrogen phosphate

PO₄^3- = phosphate

Name the following:

HNO₃ Ca(NO₃)₂

H₂SO₄ H₂SO₃

NaHSO₄ Na₂SO₄

HClO₄ HClO

NaClO₄ NaClO