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 Theory

Conjugate 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

ExampleHCl (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 H2

O

Arrhenius

•Acids – proton donor

•Bases – proton acceptor

•Can define acids & bases in solvents other than H2

O

Brønsted-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 3.0

10 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

x

xOH

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 10353.70136.0

101][

x

xOH

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

x

xOH

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%

Example

HF H+ + F-

? M? M1 M

H+

F-

F-

F-

H+

H+ HF

HF

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 ions

HI(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]