Acids and Bases

Arrhenius Definition Acids produce hydrogen ions in aqueous solution. Bases produce hydroxide ions when dissolved in water. Limits to aqueous solutions. Only one kind of base. NH 3 ammonia could not be an Arrhenius base.

Bronsted-Lowry Definitions And acid is an proton (H + ) donor and a base is a proton acceptor. Acids and bases always come in pairs. HCl is an acid.. When it dissolves in water it gives its proton to water. Water is a base makes hydronium ion

Pairs General equation HA(aq) + H 2 O(l)  H 3 O + (aq) + A - (aq) Acid + Base  Conjugate acid + Conjugate Base This is an equilibrium.

Competition for H+ H+ between H 2 O and A-A- The stronger base controls direction. If H 2 O is a stronger base it takes the H+H+ Equilibrium moves to right.

Acid dissociation constant K a

We can write the expression for any acid. Strong acids dissociate completely. Equilibrium far to right. Conjugate base must be weak. This is an equilibrium expression for the dissociation of a weak acid.

Back to Pairs Strong acids K a is large [H + ] is equal to [HA] A - is a weaker base than water Weak acids K a is small [H + ] <<< [HA] A - is a stronger base than water

Types of Acids Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic). Oxyacids - Proton is attached to the oxygen of an ion. Organic acids contain the Carboxyl group -COOH with the H attached to O Generally very weak.

Amphoteric Behave as both an acid and a base. Water autoionizes 2H 2 O(l)  H 3 O + (aq) + OH - (aq) K W = [H 3 O + ][OH - ]=[H + ][OH - ] At 25ºC, K W = 1.0 x10 -14

In EVERY aqueous solution. Neutral solution[H + ] = [OH - ] = 1.0 x10 -7 Acidic solution [H + ] > [OH - ] Basic solution [H + ] < [OH - ]

pH pH= -log[H + ] Used because [H + ] is usually very small As pH decreases, [H + ] increases exponentially pOH= -log[OH - ] pKa = -log K

Relationships K W = [H + ][OH - ] -log K W = -log([H + ][OH - ]) -log K W = -log[H + ]+ - log[OH - ] pK W = pH + pOH

K W = 1.0 x = pH + pOH [H + ],[OH - ],pH and pOH Given any one of these we can find the other three.

BasicAcidicNeutral [H + ] pH Basic [OH - ] pOH

Calculating pH of Solutions Always write down the major ions in solution. Remember these are equilibria. Remember the chemistry. Don’t try to memorize there is no one way to do this.

Strong Acids HBr, HI, HCl, HNO 3, H 2 SO 4, HClO 4 ALWAYS WRITE THE MAJOR SPECIES Completely dissociated

Strong Acids [H + ] = [HA] [OH - ] is going to be small because of equilibrium = [H + ][OH - ] If [HA]< water contributes H +

Weak Acids Ka will be small. ALWAYS WRITE THE MAJOR SPECIES. It will be an equilibrium problem from the start. Determine whether most of the H + will come from the acid or the water. Compare Ka or Kw

Example Calculate the pH of 2.0 M acetic acid HC 2 H 3 O 2 with a Ka 1.8 x10 -5 Calculate pOH, [OH - ], [H + ]

A mixture of Weak Acids The process is the same. Determine the major species. The stronger will predominate. Bigger Ka if concentrations are comparable Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x ) and 3.4 M HOC 6 H 5 (Ka = 1.6 x )

Percent dissociation = amount dissociated x 100 initial concentration For a weak acid percent dissociation increases as acid becomes more dilute. Calculate the % dissociation of 1.00 M and M Acetic acid (Ka = 1.8 x As [HA] 0 decreases [H + ] decreases but % dissociation increases. Le Chatelier

The other way What is the Ka of a weak acid that is 8.1 % dissociated as M solution?

Bases The OH - is a strong base. Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved.

Bases The hydroxides of alkaline earths Ca(OH) 2 etc. are strong dibasic bases, but they don’t dissolve well in water. Used as antacids because [OH - ] can’t build up.

Bases without OH - Bases are proton acceptors. NH 3 + H 2 O  NH OH - It is the lone pair on nitrogen that accepts the proton. Many weak bases contain N B(aq) + H 2 O(l)  BH + (aq) + OH - (aq) K b is the base dissociation constant.

K b Expression

Strength of Bases Hydroxides are strong. Others are weak. The smaller K b, the weaker the base. Calculate the pH of a solution of 4.0 M pyridine (Kb = 1.7 x )

Polyprotic acids Always dissociate stepwise. The first H + comes of much easier than the second. Ka for the first step is much bigger than Ka for the second. Denoted Ka 1, Ka 2, Ka 3

Polyprotic acid H 2 CO 3 H + + HCO 3 - Ka 1 = 4.3 x HCO 3 - H + + CO 3 -2 Ka 2 = 4.3 x Base in first step is acid in second. In calculations we can normally ignore the second dissociation.

Calculate the Concentration Of all the ions in a solution of 1.00 M Arsenic acid H 3 AsO 4 n Ka 1 = 5.0 x n Ka 2 = 8.0 x n Ka 3 = 6.0 x

Sulfuric acid is special In first step it is a strong acid. Ka 2 = 1.2 x Calculate the concentrations in a 2.0 M solution of H 2 SO 4 Calculate the concentrations in a 2.0 x M solution of H 2 SO 4

Salts as acids an bases Salts are ionic compounds. Salts of the cation of strong bases and the anion of strong acids are neutral. for example NaCl, KNO 3 There is no equilibrium for strong acids and bases. We ignore the reverse reaction.

Basic Salts If the anion of a salt is the conjugate base of a weak acid - basic solution. In an aqueous solution of NaF The major species are Na +, F -, and H 2 O F - + H 2 O HF + OH - K b =[HF][OH - ] [F - ] but Ka = [H + ][F - ] [HF]

Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b =[OH - ] [H + ] K a x K b = K W

K a tells us K b The anion of a weak acid is a weak base. Calculate the pH of a solution of 1.00 M NaCN. Ka of HCN is 6.2 x The CN - ion competes with OH - for the H +

Acidic salts A salt with the cation of a weak base and the anion of a strong acid will be basic. The same development as bases leads to K a x K b = K W Calculate the pH of a solution of 0.40 M NH 4 Cl (the K b of NH x ). Other acidic salts are those of highly charged metal ions. More on this later.

Anion of weak acid, cation of weak base K a > K b acidic K a < K b basic K a = K b Neutral

Structure and Acid-Base Properties Any molecule with an H in it is a potential acid. The stronger the X-H bond the less acidic (compare bond dissociation energies). The more polar the X-H bond the stronger the acid (use electronegativities). The more polar H-O-X bond -stronger acid.

Binary Acids The strength of binary acids increases from left to right across a period. Electronegativity causes increase in polarity.

The strength of binary acids increase from top to bottom within the same group. As bond strength increases acidity decreases. Thus, HI is stronger than HF.

Strength of Oxoacids The more oxygen hooked to the central atom, the more acidic the hydrogen. HClO 4 > HClO 3 > HClO 2 > HClO Remember that the H is attached to an oxygen atom. The oxygens are electronegative Pull electrons away from hydrogen

Electronegativity of Central Atom As electronegativity of central atom increase acid becomes a better proton donor. Electron density is drawn away from O-H bond.

Strength of oxyacids Electron Density ClOH

Strength of oxyacids Electron Density ClOHO

Strength of oxyacids ClOH O O Electron Density

Strength of oxyacids ClOH O O O Electron Density

When central atoms of oxoacids hold the same number of O atoms, acid strength increase from bottom to top in a group and from left to right across a period.

Hydrated metals Highly charged metal ions pull the electrons of surrounding water molecules toward them. Make it easier for H + to come off. Al +3 O H H

Acid-Base Properties of Oxides Non-metal oxides dissolved in water can make acids. SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) Ionic oxides dissolve in water to produce bases. CaO(s) + H 2 O(l) Ca(OH) 2 (aq)

Lewis Acids and Bases Most general definition. Acids are electron pair acceptors. Bases are electron pair donors. BF F F :N:N H H H

Lewis Acids and Bases Boron triflouride wants more electrons. BF F F :N:N H H H

Lewis Acids and Bases Boron triflouride wants more electrons. BF 3 is Lewis base NH 3 is a Lewis Acid. B F F F N H H H

Lewis Acids and Bases Al +3 ( ) H H O Al ( ) 6 H H O