Acids & Bases

Key Characteristics of Acids & Bases

AcidsTaste sour

Reacts with alkali metals (forms H2 gas)

Litmus paper: Red

Neutralizes Bases

BasesTastes bitter

Slippery feel

Litmus paper: Blue

Neutralizes Acids

Theories of Acids & Bases Arrhenius Theory of Acids & Bases

Properties of acids are due to the presence of H+ ions Example:

HCl H+ + Cl- Properties of bases are due to the

presence of OH- ions Example:

NaOH Na+ + OH-

H+ ions in water H+ ions are bare protons These H+ ions react strongly with the

nonbonding pair of electrons in a water molecule This forms the hydronium ion, H3O+

Oftentimes H+ and H3O+ are used interchangeably

HCl H+ + Cl-HCl(g) + H2O(l) H3O+

(aq) + Cl-(aq)

Problems with Arrhenius Arrhenius theory has limitations:

Only deals with aqueous solutions (solutions in water)

Not all acids and bases produce H+ and OH- ionsNH3 for example is a base

Brønsted and Lowry proposed a definition based on acid base reactions transferring H+ ion from one substance to another

Brønsted-Lowry Theory

Theories of Acids & Bases Brønsted-Lowry Theory

Acids are substances that donate H+ ions Acids are proton donors

Bases are substances that accept H+ ions Bases are proton acceptors

Example: HBr + H2O H3O+ + Br-

A B

Brønsted-Lowry Theory The behavior of NH3 can now be

understood:NH3 (aq) + H2O (l) ↔ NH4

+ (aq) + OH-

(aq)

Since NH3 becomes NH4+, it is a

proton acceptor (or a Brønsted-Lowry base)

H2O becomes OH-, which means it is a proton donor (or a Brønsted-Lowry acid)

Brønsted-Lowry TheoryConjugate Acid-Base Pairs

An acid and a base that differ only in the presence or absence of H+ are called a conjugate acid-base pair.

Every acid has a conjugate base.Every base has a conjugate acid.

HX is the conjugate acid of X-

H2O is the conjugate base of H3O+

Brønsted-Lowry Theory These pairs differ by only one hydrogen ion Example

Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base

NH3 + H2O NH4+ + OH-

B A CA CB NH3 acts as a Brønsted base by accepting a

proton. Water acts as a Brønsted acid by donating

a proton.

Brønsted-Lowry Theory Example

HCl (g) + H2O (l) ↔ H3O+(aq) + Cl- (aq)

HSO4- + HCO3

- ↔ SO4-2 + H2CO3

BA CA CB

A B CACB

Theories of Acids & Bases Lewis Acids & Bases

Acids are electron acceptors Bases are electron donors

Example: H2O + NH3 OH- + NH4

+

Is really: H2O + :NH3 OH- + H:NH3

+Electron

pair donor(NH3)

Electron pair acceptor(H+)

Summary Of Theories

•Acids release H+

•Bases release OH-

•Defines acids & bases in H2O

Arrhenius

•Acids – proton donor•Bases – proton acceptor•Can define acids & bases in

solvents other than H2OBrønsted-Lowry

•Acids – electron acceptor•Bases – electron donor•Defines acids & bases

without a solvent

Lewis

The Self-Ionization of Water Even pure water contains a small

number of ions:H2O (l) ↔ H3O+

(aq) + OH- (aq) In pure water, the concentrations of the

ions (H3O+ and OH-) are equal.

[H3O+]=[OH-]= 1x10-7 M

The Self-ionization of Water Writing the equilibrium expression for the

self-ionization of water gives:

Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 this is referred to as the ion product constant

of water This ion product constant of water is given the

symbol Kw

]][[ 3 OHOHKeq

The Self-ionization of water Example #1

What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M?

Kw = [H3O+][OH-]1x10-14 = [H3O+][3.0x10-4]

114-

14

103.310 x 3.010 x 1.0

x

Example #2 If the hydroxide-ion concentration of an

aqueous solution is 1.0 x 10-3 M, what is the [H3O+] in the solution?   Kw = [H3O+][OH-]

1x10-14 = [H3O+][1.0x10-3]

113

14

3 101100.1

101][

x

xxOH

The pH scale Developed by Søren Sørensen in order to

determine the acidity of ales Used in order to simplify the concept of acids

and bases The pH scale goes from 1 to 14 A change in one pH unit corresponds to a power

of ten change in the concentration of hydronium (H3O+) ions A pH = 2.0 has 10 times the concentration of H3O+

than a pH = 3.0, and 100 times greater than pH = 4

The pH scale

pH < 7• Acid

pH = 7• Neutral

pH > 7• Base

Calculations of pH pH can be expressed using the following

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

Example #1 What is the pH of a solution with 0.00010 M H3O+?

Is this solution an acid or a base?

Acid

)00010.0log(pH

4

Calculations of pH Example #2

What is the pH of a solution with the concentration of hydroxide ions 0.0136 M? Is this an acid or a base?pH = -log [H3O+] Kw = [H3O+]

[OH-]

Base

]0136.0][[101 314 MOHxKw

13

14

3 10353.70136.0101][

xxOH

1.12)10353.7log( 13 xpH

Calculations of pH Practice #1

Practice #2

Calculations of pH Example #1

What is the hydronium ion concentration in fruit juice that has a pH of 3.3?

[H3O+] = 10-pH

43.3 100.510][ 3 xOH

Calculations of pH What are the concentrations of the

hydronium and hydroxide ions in a sample of rain that has a pH of 5.05?

[H3O+] = 10-pH Kw = [H3O+][OH-]

60553 1091.810][ xOH .

]][1091.8[101 614 OHxxKw9

6

14

1012.11091.8

101][

x

xxOH

Calculation of pH Practice #1

Practice #2

Strength of Acids & Bases When a solution is considered strong,

it will completely ionize in a solution Nitric acid is an example of strong acid:

HNO3 (l) + H2O (l) NO3- (aq) + H3O+ (aq)

In a solution of nitric acid, no HNO3 molecules are present

Strength is NOT equivalent to concentration!

Strength of Acids & Bases Knowing the strength of an acid is

important for calculating pH If given concentration of strong acid (such

as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions

Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

Strong Acids & Bases Ionize 100%

ExampleNaOH Na+ + OH-

1 M1 M1 M

Na+Na+

Na+

OH-OH-

OH-

Weak Acids & Bases Ionize X%

ExampleHF H+ + F-

? M? M1 M

H+

F-

F-

F-H+

H+ HFHF

Strength of Acids & Bases

Stronger the acid

Weaker the conj. base

Stronger the base

Weaker the conj. acid

Strength of Acids & Bases

Strong Acids

•Perchloric acid, HClO4•Chloric acid, HClO3•Hydrochloric acid, HCl•Hydrobromic acid, HBr•Hydroiodic acid, HI•Nitric acid, HNO3•Sulfuric acid, H2SO4

Strong

Acids

Must be memorized!

Strong Acids 6 of 7 strong acids are monoprotic (HX)

Exists only as H ions and X ionsHI(aq) H+

(aq) + I-(aq)

2M HI = [H+]= [I-] = 2M Determining pH of Strong Acids

For Strong Acids: pH = -log [H+] For monoprotic strong acids: [H+] = [X]