ACIDS AND BASES …for it cannot be But I am pigeon-liver’d and lack gall To make oppression bitter… Hamlet

Learning objectives Describe properties of acids and bases Define acid and base using Arrhenius and Bronsted definitions Identify Bronsted acids and bases in solution equilibria Distinguish between strength and concentration of acids and bases Estimate pH of common substances Describe phenomenon of acid rain

Properties of acids Acids are sour Acids attack metals Acids react with bases and form salts Acids turn litmus red

Properties of bases Bases taste bitter Bases are slippery Bases react with acids to form salts Bases turn litmus blue

Neutralization The mixing of an acid with a base: ACID + BASE = SALT + WATER The reaction of carbonic acid (CO 2 in H 2 O) to give limestone: H 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2H 2 O

Arrhenius: it’s about water The meaning of acid and base has changed over the years Arrhenius acid is one that generates protons when dissolved in water Arrhenius base is one that generates hydroxide ions when dissolved in water

Hydronium ion is the active ingredient of an acid in water Protons do not exist in solution CH 3 CO 2 H + H 2 O = H 3 O + + CH 3 CO 2 - Vinegar in water produces hydronium ions

Hydroxide ion is the active ingredient of a base in water NH 3 + H 2 O = NH OH - Ammonia, a base, dissolves in water and produces hydroxide ions

The essence of neutralization Elimination of the components of acid and base by combination to give H 2 O H + + OH -  H 2 O ACID BASE

Brønsted and Lowry: All about protons Broader definition of acids and bases Reaction NH 3 + HCl = NH 4 Cl has all elements of acid-base neutralization but no H 2 O Brønsted acid donates a proton Brønsted base accepts a proton

Brønsted acid HCl + H 2 O = H 3 O + + Cl -

Brønsted base NH 3 + H 2 O = NH OH - water NH 3 + HCl = NH 4 + Cl - No water

Substances can be both acids and bases – depends on environment Note that in one instance H 2 O behaves like a base – accepting protons, and in another, behaves like an acid – donating protons HCl + H 2 O = H 3 O + + Cl - In presence of an acid H 2 O is a base NH 3 + H 2 O = NH OH - In presence of a base H 2 O is an acid

Salts Products of acid-base neutralization Contain metal cation and nonmetal anion Acid + base = salt + water HCl + NaOH = NaCl + H 2 O HCl + KOH = KCl + H 2 O HNO 3 + KOH = KNO 3 + H 2 O 2HCl + Ca(OH) 2 = CaCl 2 + 2H 2 O HCN + NaOH = NaCN + H 2 O

Strong coffee (or concentrated?) Not all acids completely donate protons to water molecules HA + H 2 O  A - + H 3 O + Strength: Degree of ionization Concentration: Number of moles per unit volume

Strong and weak Strong acid (HCl) –Fully ionized All H + and Cl - –Corrosive Weak acid (Acetic) –Weakly ionized Mostly CH 3 COOH –Edible

Changing concentration does not change strength Strength refers to degree of ionization: –Strong is completely ionized (100 %) –Weak is partly ionized (1 % - 1:10 6 ) Concentration refers to number of moles per unit volume An acid (or base) can be strong and concentrated, weak and concentrated, strong and dilute, weak and dilute

Ionization of water Even in pure water, molecules are ionized and concentrations of OH - and H 3 O + are equal H 2 O + H 2 O = H 3 O + + OH - [H 3 O + ] = [OH - ] Concentration

In all aqueous solutions, product of concentrations is a constant [H 3 O + ][OH - ] = constant Increasing [H 3 O + ] decreases [OH - ] (acidic conditions) Increasing [OH - ] decreases [H 3 O + ] (basic conditions)

The pH scale – reduces large range of numbers to small In water [H 3 O + ][OH - ] = pH = - log 10 [H 3 O + ] Range of [H 3 O + ] 10 M – M Range of pH -1 to +14 Low pH = acid; high pH = basic pH = 7 = neutral

Relating pH to [H 3 O + ] For pH, take exponent of [H 3 O + ], change sign –10 -1 M (0.1 M) HCl has pH = 1 –Pure water has [H 3 O + ] = M, pH = 7 –Ammonia has [H 3 O + ] = M, pH = 11 Note: change of 1 unit in pH is factor of ten

Estimating pH Estimating pH is often more useful than doing exact calculations Smaller pH value means larger H + concentration Estimating pH

Acidity and the environment Rain is naturally weakly acidic because of CO 2 Alkaline rocks – limestone – neutralize the acid

Acid Rain Acid rain is polluted by acid in the atmosphere. Two common pollutants acidify rain: sulphur dioxide (SO 2 ) and nitrogen oxides (NO X ) Following information from The Green Lane TM, Environment Canada's World Wide Web site -

What’s the big deal? Damage to aquatic life Damage to buildings Damage to forests Damage to air quality

Source of the problem Sulphur dioxide (SO 2 ) byproduct of industrial processes and burning fossil fuels. –Ore smelting –coal-fired power generators –natural gas processing

Where do NO X emissions come from? Main source of NO X is combustion of fuels in motor vehicles, residential and commercial furnaces, industrial and electrical-utility boilers and engines. NO X emissions were 2.5 million tonnes in U.S. NO X emissions for 2000 were 21 million tonnes.

Legislative success with acid rain Eastern Canada Acid Rain program committed Canada to cap SO 2 emissions at 2.3 million tonnes by % reduction from 1980 levels Targets achieved or exceeded By 2001, emissions were 63% reduction from 1980 levels.

Would acid rain remain a problem without further controls? Yes. That is why The Canada-Wide Acid Rain Strategy for Post-2000 calls for further emission reductions in both Canada and the United States. In total, without further controls, almost 800,000 km 2 in southeastern Canada-an area the size of France and the United Kingdom combined- would receive harmful levels of acid rain; that is, levels well above critical load limits for aquatic systems.