Acids, Bases, and Acid-Base Equilibria

Acid-Base Theories and Relative Strengths Arrhenius Theory of acids and bases acid – produces H + ions base – produces OH - ions Strong acids and bases ionize completely Problems with this theory: It’s restricted to aqueous solutions and it doesn’t include bases like NH 3 which don’t directly ionize to yield OH -. Does predict formation of salt and water.

Bronsted-Lowry Theory of Acids and Bases Acids are proton (H + ) donors (A Brønsted–Lowry acid must have at least one removable (acidic) proton (H + ) to donate) Bases are proton (H + ) acceptors (A Brønsted–Lowry base must have at least one nonbonding pair of electrons to accept a proton (H + ))

Conjugate acid: formed when a base accepts a proton Conjugate base: formed when an acid donates a proton Conjugate acid-base pair: an acid and a base that differ by only one H + On each side of the equation, there’s an acid and a base

Amphiprotic – a substance that can act as an acid or a base Ex. H 2 O As a base → As an acid →

The relative strengths of acids and bases: – The stronger an acid is, the weaker its conjugate base is – The stronger a base is, the weaker its conjugate acid is In every acid-base reaction, equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base. See Zumdahl sect. 14.2

The Self-ionization of Water K w (ion product constant of water) =  H 3 O +  x  OH -  = 1 x at 25 o C  H 3 O +  =  OH -  = 1 x for pure water (at 25°C)

If a solution is neutral, [H + ] = [OH – ]. If a solution is acidic, [H + ] > [OH – ]. If a solution is basic, [H + ] < [OH – ]. To calculate the concentration of H + or OH - when you only know one of them, use the equation  H 3 O +  x  OH -  = 1 x

The pH Scale pH: it’s a logarithmic scale: pH = -log  H 3 O +  pOH = -log  OH -  pK w = -log K w = 14 (at 25°C) pK w = pH + pOH = 14 Neutral solutions have a pH = 7 Acidic solutions have a pH<7 Basic solutions have a pH>7 Note: when the pH changes by 1, the  H 3 O +  changes 10 fold.

Measuring pH: 2 main ways of measuring pH – 1)with a pH meter – has electrodes that indicate small changes in voltage to detect pH. 2)With acid-base indicators (substances that can exist in either an acid or a base form, and are different colors at different pHs) sample rxn: HIn + H 2 O H 3 O + + In - Works according to LeChatelier’s Principle (more H+ shifts equilibrium to left, less shifts to right; shift Indicated by color change)

Strong Acids and Bases Strong Acids Remember that the seven strong acids are HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 3, and HClO 4. These are strong electrolytes and exist totally as ions in aqueous solution; e.g., HA + H 2 O → H 3 O + + A – So, for the monoprotic strong acids, [H 3 O + ] = [acid] Just use acid concentration for H+ concentration.

Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal (Ca 2+, Sr 2+, and Ba 2+ ) hydroxides. Again, these substances dissociate completely in aqueous solution; e.g., MOH(aq) → M + (aq) + OH – (aq) or M(OH) 2 (aq) → M 2+ (aq) + 2 OH – (aq) For the alkali metal hyroxides: pOH = -log[OH - ] = -log[base] pH = 14 – pOH For the alkaline earth metal hydroxides: pOH = -log[OH - ] = -log{2x[base]}

Weak Acids For a weak acid, the equation for its dissociation is HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A – (aq) Since it is an equilibrium, there is an equilibrium constant related to it, called the acid-dissociation constant, K a : The greater the K a, the stronger the acid.

Calculating pH from k a and initial concentration: Compare Ka to Kw (Ka should be larger, otherwise pH = 7) Ka = [H+][A-]/[HA]. Ka is given. H+ A- HA Init 0 0 Given (G) Change +x +x G-x End x x Assume G (small dissociation, < 5%) Solve (Ka(HA)) 1/2 = x ( Hope M acid /K a > 100, othewise use quadratic eqn to find x ) -log [x] = pH Calculating Percent Ionization: Percent ionization is a measure of acid strength (the stronger the acid, the greater the % ionization) Percent ionization =  100

Polyprotic Acids: Monoprotic acid – 1 ionizable H Polyprotic acid – more than 1 ionizable H. The ionizations occur separately. Each step has its own K a. The first K a is the largest. If the K a values for the first and second dissociation differ by a factor of 10 3 or more, the pH generally depends only on the first dissociation.

Weak Bases Like weak acids, weak bases have an equilibrium constant called the base dissociation constant (K b ). Equilibrium calculations work the same as for acids, using the base dissociation constant instead. Since “x” in this case is the [OH - ] at equilibrium, taking the -log of it gives you the pOH.

Relationship between K a and K b For a conjugate acid–base pair, K a and K b are related in this way: K a × K b = K w Therefore, if you know one of them, you can calculate the other. Also, pK a + pK b = pK w = 14 for a conjugate acid-base pair

Acid-Base Properties of Salt Solutions Combined effect of cation and anion – possible combinations: 1)Salts containing ions from strong acids and strong bases form neutral solutions. (ex. NaCl, KNO 3, BaI 2 ) 2)Salts containing ions from weak acids and strong bases form basic solutions. (ex. Na 2 CO 3, KNO 2, NaCH 3 COO) 3)Salts containing ions from strong acids and weak bases form acidic solns. (ex. NH 4 Cl, NH 4 NO 3 ) 4)Salts containing ions from weak acids and weak bases (or containing ions like Fe 3+ ) –depends on the relative acid and base strength (ex. NH 4 CN, NH 4 NO 2, CrF 3 ). Look for larger Ka or Kb.

Acid-Base Behavior and Chemical Structure Factors that Affect Acid Strength: 1)The H—A bond must be polar with δ + on the H atom and δ – on the A atom 2)Bond strength: Weaker bonds can be broken more easily, making the acid stronger. 3)Stability of A – : More stable anion means stronger acid. The strength of an acid is often a combination of all three factors.

Binary Acids: Binary acids consist of H and one other element. Within a group, H—A bond strength is generally the most important factor (acid strength increases going down a group). Within a period, bond polarity is the most important factor to determine acid strength (acid strength increases going across a period)

Oxyacids: acids in which OH groups and possible additional oxygen atoms are bound to a central nonmetal atom. As the electronegativity of the nonmetal increases, the O-H bond gets weaker, and the acid gets stronger Also, as additional atoms high in electronegativity bond to the center atom, the acid strength increases:

Carboxylic Acids: have a –COOH (carboxyl) group. The electronegativity of the R group attached to the carboxyl group determines the strength. The greater the electronegativity, the stronger the acid. K a = 1.6 x K a = 2.3 x 10 -1