Acids and Bases Acid-Base chemistry –important in our everyday lives acidity of our blood is carefully controlled making sulfuric acid is an important.

Acids and Bases

• Acid-Base chemistry – important in our everyday lives

• acidity of our blood is carefully controlled

• making sulfuric acid is an important industry– sulfuric acid is needed to make fertilizers, polymers,

steel, etc.

• environmental impact of acid rain

Acids and Bases• Properties of acids

– taste sour• e.g. vinegar, citric acid in soda and candies

– react with many metals (to form the metal ion and hydrogen gas

• 2H+ + Zn --> H2 + Zn+2

– react with carbonates (to form CO2)

• 2H+ + CaCO3 --> Ca+2 + H2O + CO2

Acids and Bases

• Properties of bases– bitter taste– slippery feel

Acids and Bases

• Definitions of Acids

• Arrhenius - an acid produces H+ in aqueous solutions

• Bronsted-Lowry - an acid is an H+ donor

• Lewis - e- pair acceptor

Acids and Bases

• Definitions of Bases– Arrhenius - a base produces hydroxide ions in

solution– Bronsted-Lowry - a base is an H+ acceptor – Lewis - a base is an e- pair donor

Acids and Bases

• Acid-Base Definitions– Arrhenius definition is limited to aqueous

solutions, and only allows for one kind of base, those with hydroxide ions.

– Bronsted-Lowry is a more general definitions, water can now act as a base.

– Lewis - the most general definition

Acids and Bases• Terms to know:

– proton - H+ ion

– hydronium ion - H3O+; results from water reacting with H+

• H2O + H+ --> H3O+

– conjugate base - whatever is left from an acid after a proton has been donated

– conjugate acid - whatever has been formed when a proton has been accepted by a base

Acids and Bases

• Acid Dissociation Constant– an equilibrium constant for the ionization or

dissociation of an acid

– Ka

– for HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq)

• Ka = [H3O+][A-] / [HA] which is equivalent to

• Ka = [H+][A-] / [HA]

• remember, pure liquids, like water, are not included in the equilibrium expression

Acids and Bases• Acid Strength

– defined by the equilibrium position of the dissociation reaction

– Strong Acid - equilibrium lies far to the right, i.e. the acid is 100% ionized

– the conjugate base of a strong acid is a much weaker base than water, i.e. the conjugate base will not accept the H+

– the stronger the acid the weaker the conjugate base

Acids and Bases• Weak acid

– not 100% ionized– equilibrium lies to the left– very little HA is ionized– the conjugate base of a weak acid is a stronger

base than water, the conjugate base is more likely to accept an H+ than water

– the weaker the acid, the stronger the conjugate base

Acids and Bases

• The Six Strong Acids– HCl

– H2SO4

– HNO3

– HBr– HI

– HClO4

Acids and Bases

• Oxyacids– Most acids are oxyacids– The acidic proton is attached to an oxygen atom

• Organic acids– Generally weak acids– Contain the -COOH (carboxyl) group

– ex: CH3COOH - acetic acid

Acids and Bases• Water

– The most common amphoteric substance• water can act as both an acid and a base

• water can autoionize:– H2O + H2O <==> H3O+ + OH-

– one water molecule acts as an acid (H+ donor), the other acts as an acid (an H+ acceptor)

– Kw = [H3O+][OH-] = 1.00 x 10-14 @ 25oC• dissociation or ion product constant for water

Acids and Bases

• Kw

– In any aqueous solution at 25oC, the product of [H+] and [OH-] will be 1.0 x 10-14

– So if you know the [H+], you can figure out the [OH-] and vice versa

– If [H+] = [OH-], the solution is neutral– If [H+] > [OH-], the solution is acidic– If [H+] < [OH-], the solution is basic

Acids and Bases• pH scale

– Because the [H+] in any solution is generally quite small, it is easier to use the pH scale to represent a solution’s acidity.

– pH comes from the Danish…potenz or strength of the H+ ion

– pH = - log[H+]– pOH = - log [OH-]

Acids and Bases• pH is a log scale

– when the pH changes by one, the [H+] concentration changes by a power of 10.

• A solution with a pH of 3 has 10 times more H+ than a solution with a pH of 4, and 100 times more H+ than a solution with a pH of 5.

– As pH decreases, the [H+] increases.

– Rule for significant figures for logarithms - the number of places after the decimal point is equal to the number of significant figures in the original number

• pH = - log 1.0 x 10-9 M (2 significant figures in 1.0 x 10-9)

• pH = 9.00 ( 2 places after the decimal point for significant figures)

Acids and Bases• Acid-Base Equilibria…or doing acid-base

problems– the aqueous solutions contain many

components– you must be able to determine which

components are most significant and which can be ignored

– you must be able to determine which reaction is most important of all the possible reactions

Acids and Bases• Ex: Calculate the pH of 1.0 M HCl

– First: Determine the major species

• HCl is a strong acid, 100% ionized

• Major species then will be H+, Cl- and H2O– Since this is an acidic solution with “lots” of H+, the

[OH-] will be insignificant.

– What about H2O <==> H+ + OH- ? Will any H+ come from this reaction?

– Le Chatelier’s principle tells us that the reverse reaction will be favored because of the high concentration of H+ from the HCl, so we can ignore this reaction as a source of H+.

• Thus the pH = - log (1.0) = 0.00

Acids and Bases

• pH of Weak Acid Solutions

• Calculate the pH of 1.0 M HF– Determine the major species: HF (because it is

a weak acid, very little of it ionizes), and H2O

– Which of these major species will provide the H+ ions?

• Consider the Ka and Kw

Acids and Bases

• HF <==> H+ + F- Ka = 7.2 x 10-4

• H2O <==> H+ + OH- Kw = 1.0 x 10-14

• Because the Ka of HF is greater than the Kw, HF is a stronger acid than H2O, and will be the primary contributor of H+ in solution.

• Use the Ka for HF to determine the [H+] at equilibrium, and then determine the pH.

Acids and Bases• Simplifying acid-base equilibrium

calculations– If Ka is small, then “x” is small relative to the

original concentrations. – 1.00 - x becomes essentially 1.00 in these

calculations– Is this a valid assumption? Please check at the

end!

– Compare x to [HAo], if x < 5% of [HAo], then the assumption is valid

Acids and Bases• The pH of a mixture of weak acids

– Calculate the pH of a solution that contains 1.00 M HCN (Ka = 6.2 x 10-10), and 5.00 M HNO2 (Ka = 4.0 x 10-4).

– Determine the major species: HCN, HNO2, and H2O

– Compare Ka’s and Kw.

• The Ka for HNO2 is larger than Kw and the Ka for HCN

• HNO2 <==> H+ + NO2- is the reaction of interest.

Acids and Bases

• % dissociation or % ionization– gives us an idea of the amount of weak acid that

has dissociated

– % dissociation = ([H+] /[HAo]) x 100

– For a weak acid, the % ionization increases as the concentration gets more dilute.

Acids and Bases

• Bases

• Strong Bases– 100% ionized

• Group I hydroxides (e.g. NaOH, KOH, etc)

• heavy Group II hydroxides (Ca(OH)2, Ba(OH)2, Sr(OH)2)

• A base does not have to contain OH- ion to be a base

Acids and Bases• Bases

– Ammonia: NH3 + H2O <==> NH4+ + OH-

– The [OH-] increases due to the reaction of ammonia with water, so ammonia is a base, specifically, a Bronsted-Lowry base, because it is a proton acceptor.

– Many bases are like ammonia with a lone pair of electrons on the nitrogen that can accept H+. Consider these bases as substituted ammonia molecules

– e.g. CH3NH2, (CH3)2NH, (CH3)3N, C2H5NH2

Acids and Bases

• General reaction between a base and water– B + H2O <==> BH+ + OH-

– Kb = [BH+][OH-]/[B]

– In equilibrium problems involving weak bases, if Kb < Kw, then the base will be the primary source of OH-.

Acids and Bases• Polyprotic Acids (many proton acids)

– Some acids can donate more than one H+

– e.g. H2SO4 and H3PO4

– Polyprotic acids lose 1 H+ at a time.

– H2SO4--> H+ + HSO4- Ka1 >> 0

– HSO4- <==> H+ + SO4

-2 Ka2 = 1.2 x 10-2

– The decreasing Ka as protons are lost indicate it is less favorable (as the negative charge on the acid increases) to lose the second (or third) proton

Acids and Bases

• Acid-Base Properties of Salts

• Salt = an ionic compound– salts dissolve in water to form ions– sometimes these ions can react with water to

form weak acids or weak bases

Acids and Bases• Neutral Salts

• When these salts dissolve in water, the pH does not change.

• There is no reaction of the ions from the salt with water.

• The anion and cation are derived from a strong acid and a strong base, respectively.

Acids and Bases

• Ex. NaCl is a neutral salt. Na+, the cation, can be considered as coming from NaOH, a strong base. Cl-

, , the anion, can be considered as coming from HCl, a strong acid. Remember the conjugates of strong acids and strong bases are very weak, weaker than water, and so are unlikely to react with water to reform the acid or base.

Acids and Bases

• Na+ + H2O cannot form NaOH + H+…strong bases do not re-form in solution.

• Cl- + H2O cannot HCl + OH- …strong acids do not re-form in solution.

Acids and Bases• Basic Salts

– The anion from this salt will react with water to form a weak acid and OH-.

– The anion must have come from a weak acid originally.

– Ex: NaCH3COO. The Na+ comes from a strong base, NaOH. CH3COO- comes from a weak acid, CH3COOH.

– Na+ + H2O cannot form NaOH + H+

– CH3COO- + H2O <==> CH3COOH + OH-

Acids and Bases• Acidic Salts

– The cation from an acidic salt will react with water to form a weak base and H+.

– The cation must have come from a weak base.

– Ex: NH4Cl. The NH4+ comes from a weak base,

and will react with water to re-form the weak base and H+ (H3O+). The Cl- comes from a strong acid, and cannot react with water to re-form HCl.

• NH4+ + H2O <==> NH3 + H3O+

Acids and Bases

• How do we know the Kb for the conjugate base of a weak acid? How do we know the Ka for the conjugate acid of a weak base?

• CH3COOH + H2O <==> CH3COO- + OH- Ka = 1.8 x 10-5

• CH3COO- + H2O <==> CH3COOH + OH-

• Kb = [CH3COOH][OH-]/[CH3COO-] = ?

• Ka. Kb = Kw

Acids and Bases

• Salts that are derived from both a weak acid and a weak base– We can predict whether the solution will be

acidic or basic based on the Ka of the acidic ion and the Kb of the basic ion.

– Ka > Kb pH < 7, acidic solution

– Kb > Ka pH > 7, basic solution

– Kb = Ka pH = 7, neutral solution

Acids and Bases• The effect of structure on acid-base properties

– Any molecule containing hydrogen could theoretically act as an acid. However, most of these molecules don’t.

– Organic molecules with lots of C-H bonds are not acidic because the C-H is both strong, and relatively nonpolar, so there is no tendency to lose H+

– However, while the H-Cl bond is stronger than the C-H bond, the H-Cl bond is very polar, and there is a strong tendency to lose H+.

Acids and Bases• Two Factors that determine the acidity of a molecule

containing X-H

– strength of the X-H bond

– polarity of the X-H bond

• Consider polarity of the X-H bond

– H-F > H-Cl > HBr > HI in terms of polarity

– H-F > H-Cl > H-Br > HI in terms of bond strength

• HF is a weak acid, HCl, HBr, and HI are strong acids

• even though the H-F bond is very polar, the H-F bond is very strong, so HF is a weak acid

Acids and Bases

• Oxyacids– contains the grouping H-O-X

– HClO4 Ka >>> 0

– HClO3 Ka = 1

– HClO2 Ka = 1.2 x 10-2

– HClO Ka = 3.5 x 10-8

– As the number of oxygen atoms increase, the acid strength increases

Acids and Bases

• Oxygen is very electrognegative. With increasing number of oxygen atoms, the oxygen atom draw more and more of the electron density towards themselves, thus polarizing and weakening the O-H bond,

Acids and Bases

• This behavior also occurs with hydrated metal ions.

• Al+3 is strongly attracted to water, weakening the O-H bond. The greater the charge on the metal ion, the stronger the acidity.

Acids and Bases

• The greater the electronegativity of X in the H-O-X grouping, the stronger the acid.– X is able to withdraw electron density from the

H-O bond, thus weakening and polarizing the H-O bond, resulting in a stronger acid

HOC

Acids and Bases• Comparison of Electronegativity of X and Ka

Acid E.N. of X Ka

HOCl 3.5 4 x 10-8

HOBr 2.8 2 x 10-9

HOI 2.5 2 x 10-11

Acids and Bases

• Acid-Base Properties of Oxides• An H-O-X grouping in a molecule may result in

a molecule that behaves as a base or an acid.• The nature of the O-X bond will determine the

acidic or basic behavior of the molecule

Acids and Bases• If X is very electronegative, the O - X bond will

be covalent (not as polar because of similar E.N.’s), and strong. The O - X bond will remain intact in water, while the polar and weak H - O bond will break, resulting in acidic behavior.– e.g. HClO2 or H2SO4

• If X is not very electronegative, the O - X bond will be ionic, and can be broken in water, resulting in basic behavior.– e.g. NaOH or KOH

Acids and Bases• Basically…

– Nonmetal oxides will form acids in water• SO2 + H2O <==> H2SO3

• CO2 + H2O <==> H2CO3

– Metal oxides will form bases in water• Na2O + H2O --> 2 NaOH

• CaO + H2O --> Ca(OH)2

• (the oxide ion has a high affinity for protons and reacts with water to form hydroxide ions)

Acids and Bases

• Lewis acids and bases are electron pair acceptors and electron pair donors, respectively

• A very general model for acids-base reactions…see Lewis structures for explanations...