Acids and BasesThree major ways to define acids and bases introduced by Lewis, Brønsted and Arrhenius.

They differ in the role of water

Arrhenius and Brønsted require water, Lewis does not

BrønstedAcid Donates an H+

Base Accepts an H+

HCl + H2O → H3O+ + Cl- Acid

NaOH + H+ → Na+ + H2O Base

Arrhenius Acid Produces H3O+ when added to water

Base Produces OH- when added to water

Acids and Bases

NH3+ H2O → NH4+ + OH-

HCl + H2O → H3O+ + Cl-

Acids and Bases

LewisAcid Accepts electrons

Donates electronsBase Note: Electrons are not transferred between acids and bases, they are shared.

BH3 + NH3 → BH3NH3

Acid Base

BH H

HN H

H

H

: B N H

H

H

H

H

H

Acids and BasesThe Lewis definition is the most general

Consider a Brønsted Acid It donates a H+ H+ leaves electrons behind

A-H A : _

H+ i.e. A accepts the electrons

A is a Lewis Acid

All Brønsted acids are Lewis Acids

An Arrhenius acid, is a Bronsted Acid, since it produces H3O+ when dissolved in water as it “donates” H+ to H2O.

A strong acid, just like a strong electrolyte, is an acid which dissociates completely when dissolved in water.

Strong Acids and Bases

The concentration of H3O+ is thereby the highest possible, determined exactly by how much acid was added to water

Ex) HCl + H2O → H3O+ + Cl-

Ex) H2SO4 + H2O → H3O+ + HSO4

_

Inorganic acids tend to be strong acids (except HF)

A strong base, just like a strong electrolyte, is dissociates completely when dissolved in water.

The concentration of OH_ is thereby the highest possible, determined exactly by

how much acid was added to water.

Ex) NaOH → Na+ + OH_

The hydroxides of alkali metals are strong bases.

Reactivity

A reaction between an acid and a base produces water and a salt

HCl(aq) + NaOH(aq) → H2O (l) + NaCl (aq)

H3O+(aq) + Cl-(aq) + Na+(aq) + OH_ (aq) → H2O(l) + Na+(aq) + Cl-(aq)

H3O+(aq) + OH_(aq) → H2O(l)

Acid Base Water Salt

Ionic Equation

Net equation

A strong acid will react completely with any base.

A strong base will react completely with any acid.

Weak Acids and BasesA weak acid, just like a weak electrolyte, does not dissociate completely when dissolved in water

The concentration of H3O+ is not the highest possible, since much remains in the undissociated form.

Ex) CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq)

The concentration of H3O+ is determined from the dissociation constant, similar to Ksp, and the amount of acid added.

Organic acid tend to be weak acids

A weak base, just like a weak electrolyte, does not dissociate completely when dissolved in water.

The concentration of OH- is not the highest possible, since much remains in the undissociated form.

The concentration of OH- is determined from the dissociation constant, similar to Ksp, and the amount of acid added.

Ex) NH3+ H2O → NH4+ + OH-

Metal Oxides and nitrogen containing organic compounds tend to be weak bases

(l) (s) (aq)2 4 10 3 46H O +P O 4H PO

(s) (l) (aq)2 2Na O +H O 2NaOH

(s) (aq) (aq) (l)2 3 3 2Al O + 6HCl 2 AlCl + 3H O

(s) (aq) (l) (aq)2 3 2 4Al O + 2NaOH + 3H O 2Na[Al(OH) ]

Non-metal oxides react with water to give oxoacids

Amphoteric oxides usually do not dissolve with water by themselves, but react with both strong acids and strong bases to give soluble products

The oxides are anhydrides

Metal oxides react with water to give hydroxide bases

Therefore metal oxides are anhydrides of bases

Therefore non-metal oxides are anhydrides of acids

Metalloid oxides are amphoteric: they react with either strong acids or strong bases

The oxides are anhydrides

1 2 13 14 15 16 17

Strength of acids and bases is correlated with the positions of the oxides on the PT

Metal Metalliod Non-metal

pHThe acidity (or basicity) of a solution is reported as pH:

pH = -log [H3O+] or [H3O+] = 10-pH

p = power of

= concentration of H3O+

in mol./l = molar (M)

For pH < 7 solution is acidic

For pH > 7 solution is basic

Ex) 0.10 M solution of HCl [H3O+] = 0.10 M

pH = - log [0.10] = -(-1.00) = 1.00 # of sig. figs. Increased from 2 to 3?

For logarithmic quantities only the decimal numbers are significant.

Therefore a pH = 1.00 has only 2 sig. figs,

Note: pH does not have units

pOHBasicity of a solution can be reported as pOH:

pOH = -log [OH-] Where [OH-] = conc. of OH in mol/l

pH and pOH are related by: pH + pOH = 14 at 25 oC

Therefore an acidic solution as pOH > 7, and a basic solution has pOH < 7

Exercise Determine the pH and pOH of:

a) 0.275 M HNO3 solution

[H3O+] = 0.275 M pH = - log (0.275) = -(0.561) = 0.561

pOH = 14.000 – pH = 14.000 -0.561 = 13.439

b) 0.0051 M NaOH solution

[OH-] = 0.0051 M pOH = - log (0.0051) = -(-2.29) = 2.29

pH = 14.000 – pOH = 14.000 -2.29 = 11.71

Strength of an Acid• The quantity we use to measure the

strength of an acid is its pKa

• (corresponds to how easily an acid gives up H+). Acids with low pKa values

• are strong while acids with high pKa values are weak: pKa

H2O (pKa = 14)

HCO3-1 (pKa = 10.3)

CH3CO2H (pKa = 4.7)

H3O+ (pKa = 0) This includes all aqueous solutions of HCl, HBr, HI, HNO3, HClO4 and H2SO4

H3PO4 (pKa = 2.15)

HF (pKa = 3.1)

Concentrated HNO3 (pKa = -1)

Concentrated HCl (pKa = -7)

Concentrated H2SO4 (pKa = -3)

Citric acid (pKa = 3.1)

H2CO3 (pKa = 6.4)

STRONG ACIDS

WEAK ACIDS

Aqua Complexes as Acids

A hydrated proton has a pKa of 0 defining the line between the strong and

weak acids.

How about aqua complexes of metals?

What factors affect how easily the aqua complex gives up H+?

Cation pKa Approximate pH of a 1 M solution

Na(OH2)6+ 14.2 7

Ag(OH2)6+ 12 6.0

Mg(OH2)62+ 11.4 5.7

Al(OH2)63+ 5 2.5

Ti(OH2)64+ -4 0

Aqua Complexes as AcidsIf we plot pKa versus z2/r for a variety of aqua

complexes, we see that there is a correlation.

If we only look at those metals with low

electronegativity values (1.5), we can

approximate:

If we introduce an empirical “fudge factor”, we

get a more accurate – if more complex

formula:

r

zpmpKa

2

16.8814.15

)50.1(0960.016.8814.15 1

2

Paulinga pmr

zpmpK

Z2

r