Ib Hl Chemistry Assessment Statements Topics 8 And 18

Andrew Voyles IB Chemistry Assessment Statements: Topic 8: Acids and Bases Topic 18: Acids and Bases 8.1.1: According to the Bronsted-Lowry theory, acids are defined as proton (H+ ion) donors and bases are defined as proton acceptors. 8.1.2: For a compound to act as a Bronsted-Lowry (BL) acid, it must have a hydrogen atom in it, which it is capable of losing while remaining fairly stable. A BL base must be capable of accepting a hydrogen ion while remaining relatively stable (or reacting to form stable compounds, i.e. a water and a salt). Some compounds (such as water) may act as both a BL acid and a BL base i.e. (H2O-> OH- or H3O+) 8.1.3: The conjugate base will always have one less H atom that the acid and the acid one more than the base. In compounds where there are many hydrogen atoms, the one which is held the weakest is generally the one which is lost, and this must be reflected in the writing of the compound, as in the CH3COOH example below. CH3COOH (Acid)/ CH3COO- (Conjugate Base) NH3 (Base) / NH4+ (Conjugate Acid)

8.2.1: Acids are electrolytes, and have a sour taste, although one should never taste them for identification purposes. They undergo single replacement reactions with metals to form a salt and hydrogen gas, and neutralization reactions with bases to form salts and water. A strip of litmus paper dipped in an acid will change from blue to red, and an acid will turn phenolphthalein colorless. Bases are also electrolytes, and have a bitter taste, although one should never taste them for identification purposes. A strip of litmus paper dipped in a base will change from red to blue, and a base will turn phenolphthalein red. They are also slippery to the touch and undergo neutralization reactions with acids to form salts and water.

8.3.1: Strong and weak acids and bases are defined by their ease of losing (or donating) a proton or hydroxide ion, respectively. A strong acid or base, when placed in water, will almost fully ionize/dissociate, producing H3O+ or hydroxide ions from water. A weak acid or base will, however, only partially do this, leaving some unreacted acid or base

remaining. This is set up as an equilibrium, and so when some of the H3O+ ions produced by a weak acid are reacted, Le Chatlier’s means that more of the acid will react to form H3O- ions. This means that, given an equal number of mols of acid, they will be neutralized by the same amount of strong base, but their solutions will have different pH values. A weak base is the same as this, only it accepts protons and so produces OH- ions from water rather than H3O-. Any solution’s ability to conduct electricity is defined by is charges ions in it. A strong acid will produce more charged ions than a weak one, and so it’s solution will be a better electrical conductor than a weak acid. The same goes for strong/weak bases. 8.3.2: The strong acids are: HCl (hydrochloric), H2SO4 (sulfuric), HNO3 (nitric), and HClO4 (perchloric). All other acids are weak acids. The strong bases are the hydroxides of the group 1 metals, and Ba(OH)2. All other bases are weak. 8.3.3: Strong acids and bases are those which disassociate completely or almost completely in aqueous solution. The strength of an acid or base can be measured with a universal indicator or a pH meter. Also the rate of reaction measured by hydrogen production with metals or CO2 with CaCO3 will reveal the strength of an acid.

8.4.1: pH vales range up and down from 7 (7 being the neutral value of pure water at 20 ºC and 1 atm). Lower pH values are acidic, higher values are basic. 8.4.2: For two acidic solutions, the solution with the lower pH is more acidic. For two basic solutions, the solution with the higher pH is more basic. 8.4.3: A change of 1 in the pH scale represents a 10 times change in the acidity or alkalinity of the solution, due to the fact that the scale is logarithmic, of base 10. 8.4.4: The factor by which the pH changes is equal to 10ΔpH. If the pH increases by two units, then the solution is more alkaline by a factor of 102, 100 (or 1/100 as acidic). If the pH decreases by three units, then the solution is more acidic by a factor of 103, 1000 (or 1/1000 as alkaline).

18.1.1: Kw=[H+][OH-]. The value of Kw is 1 x 10-14 at 25c, but varies with temperature. 18.1.4: n general HA(aq) <=> H+(aq) + A-(aq) or B(aq) + H2O(l) <=> BH+ + OH-(aq). Therefore Ka = [H+][A-]/[HA] and Kb = [BH+][OH-] / [B] 18.1.6: The larger an acid’s Ka or the lower its pKa, the stronger it is. The larger a base’s Kb or the lower its pKb, the stronger it is.

18.2.1: A buffer solution is composed of a weak acid/base and it's conjugate base/acid. A solution of weak acid is made, and this forms a equilibrium with the water: HA + H2O <=> A- + H3O+. To this solution, some of the acid's conjugate base (A-) is added, resulting in an increase in the concentration of A-. Some of this reacts with the H3O+. The result of this is, when equilibrium is re-established, there is a considerable amount of both HA and A- present in the solution, in an dynamic equilibrium. If some other acid is added, this will react with the A-, but this causes the equilibrium to shift to the right, almost completely counteracting any pH change. The addition of a base, which reacts with the HA, cause the equilibrium to be shifted to the left, again resulting in very little pH change. This continues until one of the two components, either HA or A- are completely used up, at which time the pH then changes normally. The reverse of this is true for bases. 18.2.2: The pH of a buffer solution can be found via the following expression Ka = [H+][A-] / [HA] ). This expression can first be rearranged to [H+] = Ka x [HA] /[A-]. Given the concentration of both the Acid and its conjugate base, and the Ka value of the acid, the concentration of H+ can be calculated and this can be converted into a value for pH.

18.3.1: The salt of a strong acid and a strong base will be neutral when dissolved in aqueous solution. The salt of a weak acid and a strong base will be basic when dissolved in aqueous solution. The salt of a strong acid and a weak base will be acidic when dissolved in aqueous solution. The pH of a salt of a weak acid and a weak base will depend on the relative values of Ka and Kb.

18.4.1:

18.5.1: HIn H+ + InHIn and In are different colors, so based on whether the solution containing the indicator solution is acidic or basic, one color will be present. 18.5.2: The pH range of the indicator falls around its pKa value, so to be useful, the pKa must fall within the vertical asymptote at the equivalence point of the titration curve. 18.5.3: Methyl orange: 3.2–4.4 Bromocresol green:3.8–5.4 Methyl red: 4.8–6.0 Bromothymol blue: 6.0–7.6 Phenol red: 6.6–8.0 Phenolphthalein: 8.2–10.0