Acid and Base Equilibria Chapter 8

Acid and base definitions based on experiments. Acids Tastes sour Conducts electricity Changes the colour of litmus paper from blue to red Turns neutral (green) Bromthymol blue to yellow Reacts with active metals such as zinc and magnesium, liberating hydrogen gas (H2). Reacts with carbonates, releasing carbon dioxide gas (CO2)

Bases Tastes bitter Conducts electricity Changes the colour of litmus paper from red to blue Turns neutral (green) Bromthymol blue to blue Turns colourless phenolphthalein to pink Reacts with an acid to destroy its properties Feels slippery

Arrhenius Theory Review Acids are solutes that produce hydrogen ions/protons in aqueous solutions (increases the concentration of H+) or hydronium ions (H3O+). Ex: HCl H+(aq) + Cl-(aq)  OR H2O(l) + HCl(g) H3O+ (aq) + Cl-(aq) Bases produce hydroxide ions when dissolved in water. NaOH(s)  Na+(aq) + OH-(aq) However, this model does not account for basic properties of compounds that do not contain hydroxide ions, such as ammonia –NH3(aq) {NH3(g) + H2O(l) NH4 +(aq) + OH-­(aq)}

Bronsted-Lowry Theory According to this theory, an acid is a proton donor, and a base is a proton acceptor. A substance can only be classified as one or the other for a particular reaction (as it can change from one reaction to another).  focus on proton transfer Ex. of acid: HCl; when hydrogen chloride reacts with water, a proton is transferred from HCl to H2O H2O + HCl(g) H3O+ (aq) + Cl-(aq) base acid conj. acid conj. base. Ex. of base: when ammonia reacts with water, water now acts as an acid because it donates a proton to ammonia, which is the Bronsted-Lowry base NH3(g) + H2O(l) NH4 +(aq) + OH-­(aq) base acid conj. acid conj. base.

Amphoteric As you can see in the above reactions, water can act as either a base or as an acid. A substance that can act as either a Bronsted-Lowry acid or B-L base given the reaction is called amphoteric (amphiprotic); it can donate or accept a proton. {amphoteric: may act as an acid or base; amphiprotic: may accept or donate protons; for Bronsted-Lowry acids and bases, amphiprotic is always amphoteric, but not in more general definitions}; equilibrium conditrion; like water or bicarbonate ion in baking soda

Amphoteric – Water Another example of water as amphoteric: HCO3-(aq) + H2O(l) H2CO3(aq) + OH-(l) base acid HCO3-(aq) + H2O(l) CO32-(aq) + H3O+(aq) acid base (autoionzation)

Neutralization benefit of B-L rather than Arrhenius is that it defines acids and bases in terms of chemical reactions, so that neutralizations do not have to produce water and salt, as according to Arrhenius Arrhenius neutralization: acid-base neutralization produces water and salt as in NaOH(aq) + HCl(aq)  H2O(l) + NaCl(aq) NH4OH(aq) + HCl(aq)  H20(l) + NH4Cl Acid-base neutralization without hydronium ions, hydroxide ions, or water: NH3(g) + HCl(g)  NH4Cl (proton transferred from Cl atom to N atom)

Reversible Acid-Base Reactions Prompt: Which is the acid? How do we know that? Where is the proton transfer? Reversible Acid-Base Reactions (Equilibrium implied in Bronsted-Lowry reaction) In each proton transfer reaction at equilibrium, both forward and reverse reactions involve Bronsted-Lowry acids and bases. On each side of the reaction are acids and bases, which are called conjugate acid-base pairs. For ex.: HC2H3O2(aq) + H2O(l)  C2H3O2-(aq) + H3O+(aq) acid base conj base conj acid In any acid-base equilibrium, there will always be two acids and two bases. The base on the right (product) is formed by the removal of the proton from the acid on the left. The acid on the right (product) is formed by the addition of a proton to the base on the left. A conjugate acid-base pair is a pair of substances whose molecular formula differ by a single H+ ion (proton). ( the acid has one more proton than the base) The acid in the forward reaction is a proton donor and the base is the proton acceptor. In the reverse reaction, the conjugate acid is the proton donor and the conj. base is the acceptor.

Competition for protons (p.530) View acid-base reactions as a competition for protons between two bases. Strong acids in reactions go to almost 100% ionization or percent reaction. (the proton transfer is almost a complete forward reaction – and almost no reverse proton transfer occurs); therefore, an equilibrium is not established in this case Weak acids have a lower % ionization, so their equilibrium position favours the reactants rather than the products. The stronger an acid, the weaker its conjugate base – the acid has a weaker affinity for the proton, and easily loses/transfers the proton to its conjugate base. The weaker an acid, the stronger its conjugate base. The stronger the base, the stronger the attraction for protons.

Strong and weak acids Bronsted-Lowry explanation of strong and weak acids: HA(aq) is used as general symbol for acid, and A-(aq) as its conjugate base. Ionization reaction: HA(aq) + H2O(l) A-(aq) + H30+(aq) The strength of the acid HA is determined by the extent of the proton transfer. The ionization equation of an acid in water is often abbreviated: HA(aq) + H2O(l)  A-(aq) + H30+(aq) to HA(aq)  A-(aq) + H+(aq)

continued For ex, if we go back to our previous example, HC2H3O2(aq) + H2O(l) C2H302-(aq) + H30+(aq) acid base conj base conj acid then it reduces to HC2H3O2(aq)  C2H302-(aq) + H+(aq) Although this reduced equation shows the change that takes place, it does not demonstrate the important role that water plays in causing the acid to ionize, or that the proton most likely exists as a hydronium ion (H3O+).

Strong Acids (p.534 ) an acid that is assumed to ionize quantitatively (completely) in aqueous solution (100 % ionization is > .99% - however, we assume it’s 100% in calculations) hydrochloric acid HCl(aq), hydrobromic acid HBr(aq), sulfuric acid H2SO4(aq), nitric acid HNO(aq), phosphoric acid H3PO4(aq) Monoprotic acid: an acid that possesses only one ionizable (acidic) proton; ex. HCl Diprotic acid: an acid that possesses two ionizable (acidic) protons; ex. H2SO4 Triprotic acid: possesses three ionizable hydrogen atoms; ex. H3PO4 Note: Dissociation: Chemistry a. The process by which the action of a solvent or a change in physical condition, as in pressure or temperature, causes a molecule to split into simpler groups of atoms, single atoms, or ions. b. The separation of an electrolyte into ions of opposite charge.

Strong Bases (p.537) an ionic substance that (according to Arrhenius) dissociates completely in water to release hydroxide ions dissolve in water (dissociate completely) LiOH(s), NaOH(s), KOH(s), RbOH(s), CsOH(s) Mg(OH)2(s), Ca(OH)2(s), Ba(OH)2(s), Sr(OH)2(s) For every mole of metal hydroxide that is dissolved, one mole of hydroxide ion is produced. NaOH(s)  Na+(aq) + OH-(aq) Group 2 elements – dissolve in water and form 2 hydroxide ions Ba(OH)2(s)  Ba2+(aq) + 2OH-(aq) Hydroxide ions react with hydrogen ions – shifts equilibrium to right, causing undissolved salts to dissolve and produce higher hydroxide ion concentrations

Hydrogen Ion Concentration and pH (p. 540) There is a wide range of concentration of hydrogen ions and hydroxide ions in acidic and basic solutions. The pH scale was developed to measure hydrogen ion concentration. pH = –log[H+(aq)] Example: Calculate the pH of a solution with a hydrogen ion concentration of 5.3 x 10-9. pH = –log [H+(aq)] pH = –log (5.3 x 10-9) (two sig digits) pH = 8.28 The solution has a pH of 8.28.

(p. 541) pH of pure (neutral) water and any neutral solution at SATP is 7.00 as the hydrogen and hydroxide ions are equal.