Chapter 8

ACIDS, BASIS, AND IONIC OMPOUNDS

Definitions

• Electrolyte is a "medical/scientific" term for salts, specifically ions. The term electrolyte means that this ion is electrically-charged and moves to either a negative (cathode) or positive (anode) electrode

Defintions

• Ions that move to the cathode (cations) are positively charged

• Ions that move to the anode (anions) are negatively charged

• Electrolyses is the passage of electricity through a solution holding dissolved ions

• Electrolyte is a solute that enables a solution to conduct electricity

• Electrodes are the plates or wires that dip into the solution.

• For example, your body fluids blood, plasma, interstitial fluid (fluid between cells) -- are like seawater and have a high concentration of sodium chloride (table salt, or NaCl). The electrolytes in sodium chloride are:

• sodium ion (Na+) - cation • chloride ion (Cl-) - anion

• Ionization is the gain or loss of electrons. The loss of electrons, which is the more common process in astrophysical environments, converts an atom into a positively charged ion, while the gain of electrons converts an atom into a negatively charged ion

• The ionization energy of an atom measures how strongly an atom holds its electrons. The ionization energy is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom.

• As electrons are removed, the positive charge from the nucleus remains unchanged, however, there is less repulsion between the remaining electrons

• Zeff increases with removal of electrons • Greater energy is needed to remove remaining electrons (i.e. the ionization energy

is higher for each subsequent electron)

Notation for Degrees of Ionization

Suffix

Ionization Examples

Chemist's

Notatio

n

INot ionized

(neutral)H I, He I H, He

II Singly ionizedH II, He

IIH+, He+

III Doubly ionizedHe III, O

IIIHe++, O+

+

What is the process of solutes when dissolved in water?

• Through ionization, solutes release ions in water

• Molecules mix readily because both types of molecules engage in hydrogen bonding. Since the intermolecular attractions are roughly equal, the molecules can break away from each other and form new solute

Ammonia Dissolves in Water solute (NH3), solvent (H2O) hydrogen bonds

What is the process of solutes when dissolved in water?

• Alcohol Dissolves in Water: The -OH group on alcohol is polar and mixes with the polar water through the formation of hydrogen bonds. A wide variety of solutions are in this category such as sugar in water, alcohol in water, acetic and hydrochloric acids.

Do ions carry electricity in water?

• Yes, ions in water do carry electricity.

• Redox reactions primarily involve the transfer of electrons between two chemical species.

• The compound that loses an electron is said to be oxidized, the one that gains an electron is said to be reduced.

• A compound that is oxidized is referred to as a reducing agent, while a compound that is reduced is referred to as the oxidizing agent.

How?

Redox Reactions (Oxidation – Reductions)

• Redox reactions are the transfer of electrons from one reactant to another...

• Sn2+ donated electrons to the Fe3+ (an electron transfer took place).

• In these reactions, the valency (oxidation number) of the reactants change.

• The iron (iii) + tin (ii) have reacted to give iron (ii) + tin (iv) of course, this reaction is carried out in the presence of HCl (Hydrochloric Acid), but the oxidation reduction reaction is only between the iron (iii) and tin (ii).

• Now, a redox reaction is the release and uptake of electrons.• So, the Fe3+ is reduced to Fe2+, and the Sn2+ is oxidized to Sn4+.

For example: 2Fe3+ + Sn2+ -> 2Fe2+ + Sn4+ (8+ each side of the equation)

Oxidation-Reduction

• When there is oxidation, there is also reduction.

• The substance which loses electrons is oxidized.

• The substance which gains electrons is reduced.

• For example: Fe (metal) + Cu2+ -> Fe2+ + Cu (metal)

• Fe donates two electrons to the Cu2+ to form Cu (metal). The Fe lost 2 electrons, so is oxidized.

• The Cu2+ gained 2 electrons, so is reduced (in its valency).

Fe + Cu2+

->Fe2+ + Cu

OxidisedReducingAgent

ReducedOxidisingAgent

Redox Reactions involving acid and bases solutions

• Not only are there an exchange of electrons in these reactions, but also an exchange of protons (hydronium ions), as in any base system

• CuS + HNO3 -> Cu SO4 + NO (g) + H2O (equation not balanced).

• 3CuS + 8HNO3 -> 3 CuSO4 + 8NO(g) + 4H2O (equation balanced)

• 3CuS2+ + 3S2- + 8H+ + 8NO3- -> 3Cu2+ + 3SO42- + 8NO(g) + 4H2O

Redox Reactions Split into 2 separate steps.

• 2Fe3+ + 2e- -> 2Fe2+ (reduction)• (6+) + (2-) -> (4+) (balanced for

charges)• Sn2+ -> Sn4+ + 2e- (oxidation)• (2+) -> (4+) + (2-)

• Add the two half equations: 2Fe3+ + 2e- + Sn2+ -> 2Fe2+ + Sn4+ + 2e-The electrons cancel each other out, so equation is: 2Fe3+ + Sn2+ -> 2F2+ + Sn4+By breaking down the equation into half cells, the oxidation or reduction of each chemical can be determined. The atom which gains electrons reduces its valency, therefore is reduced  and is called the oxidizing agent.

• The atom which loses electrons, increases its oxidation number, therefore is oxidized, and is called the reducing agent.

Another Example• At the cathode: Cu2+(aq) + 2e- ----------) Cu(s) = Reduction• At the anode: 2Br-(aq) ----------) Br2(l) + 2e- = Oxidation• Sum: Cu2+(aq) + 2Br-(aq) ----------) Cu(s) + Br2(l) = Electrolysis

• Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases

• Bases feel slippery, change litmus blue, and become less basic when mixed with acids

[H+] pH Example

Acids

1 X 100 0 HCl 

1 x 10-1 1 Stomach acid

1 x 10-2 2  Lemon juice

1 x 10-3 3  Vinegar

1 x 10-4 4 Soda

1 x 10-5 5 Rainwater

1 x 10-6 6 Milk

Neutral 1 x 10-7 7 Pure water

Bases

1 x 10-8 8 Egg whites

1 x 10-9 9 Baking Soda

1 x 10-10 10 Tums® antacid

1 x 10-11 11 Ammonia

1 x 10-12 12 Mineral Lime - Ca(OH)2

1 x 10-13 13  Drano®

1 x 10-14 14 NaOH

Water as a base

• When an acid reacts with water, the water behaves as a proton acceptor to form the hydronium ion.

• When a base (like ammonia) reacts with water, a proton is transferred from water to the ammonia molecule to form the ammonium ion. Therefore, water is behaving as a proton donor

Salts

• Ionic solids (or salts) contain positive and negative ions, which are held together by the strong force of attraction between particles with opposite charges. When one of these solids dissolves in water, the ions that form the solid are released into solution, where they become associated with the polar solvent molecules.

• H2ONaCl(s)Na+(aq) + Cl-(aq)• We can generally assume that salts dissociate into their ions when they

dissolve in water. Ionic compounds dissolve in water if the energy given off when the ions interact with water molecules compensates for the energy needed to break the ionic bonds in the solid and the energy required to separate the water molecules so that the ions can be inserted into solution

Solubility Rules

1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble.

2. Salts containing nitrate ion (NO3-) are generally soluble.

3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.

4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually anything else is insoluble.

5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4, and CaSO4.

6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.

7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.

8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some other insoluble carbonates include FeCO3, PbCO3. Carbonates become soluble in acid solution.

9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4

10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag2PO4

11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2.

Solubility example

• Take for example the reaction of lead(II) nitrate with sodium chloride to form lead(II) chloride and sodium nitrate, shown below:

• Pb(NO3)2(aq) + 2 NaCl(aq)   PbCl2(s) + 2 NaNO3(aq)• This complete equation may be rewritten in ionic form by using the

solubility rules, lead(II) nitrate is soluble and therefore dissociated, same about NaCl. As products, sodium nitrate is predicted to be soluble and will be dissociated.

• The lead(II) chloride, however, is insoluble. The above equation written in dissociated form is:

• Pb2+(aq) +  2 NO3-(aq)  +  2 Na+(aq)  +  2 Cl-(aq)    PbCl2(s)  +  2 Na+(aq)  +  2 NO3-(aq)

• At this point, one may cancel out those ions which have not participated in the reaction. Notice how the nitrate ions and sodium ions remain unchanged on both sides of the reaction.

• Pb2+(aq)  +  2 NO3-(aq)  +  2 Na+(aq)  +  2 Cl-(aq)    PbCl2(s)  +  2 Na+(aq)  +  2 NO3-(aq)

• What remains is the net ionic equation, showing only those chemical species participating in a chemical process:

• Pb2+(aq) + 2 Cl-(aq)   PbCl2(s)