ACIDS AND BASES

Chemistry Ms. Piela

Key Characteristics of Acids & BasesAcids

Taste sour

Reacts with alkali metals

(forms H2 gas)

Forms electrolyte solutions (conducts electricity)

pH paper color: Red

Neutralizes Bases

BasesTastes bitter

Slippery feel

Forms electrolyte solutions (conducts electricity)

pH paper color: Blue

Neutralizes Acids

The 3 Main Theories of Acids/Bases

This course will mainly deal with

BL theory

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-

What is an H+?

H+ ions are bare protons These are so reactive that they do not

exist naturally, but will bond with water to form a hydronium ion, or H3O+

ion Oftentimes H+ and H3O+ are used

interchangeably

HCl H+ + Cl-

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

Problems with the Arrhenius theory

Only deals with aqueous solutions (solutions in water)

Not all acids and bases contain H+ and OH- ionsExample: NH3 is a base

Considered the most incomplete theory of acids and bases

Theories of Acids & Bases

Brønsted-Lowry Theory of Acids & Bases Acids are substances that donate H+

ionsAcids are proton (H+) donors

Bases are substances that accept H+ ionsBases are proton (H+) acceptors

Example:

HBr + H2O H3O+ + Br-

A B

Brønsted-Lowry Theory

The behavior of NH3 can be understood now:

NH3 (aq) + H2O (l) ↔ NH4+

(aq) + OH- (aq)

NH3 becomes NH4+, so NH3 is a

proton acceptor (or a Brønsted-Lowry base)

H2O becomes OH-, so H2O is a proton donor (or a Brønsted-Lowry acid)

Brønsted-Lowry Theory

Brønsted-Lowry Theory

Conjugate Acid-Base Pairs Definition: An acid and a base that

differ only in the presence or absence of H+

Every acid has a conjugate base. Every base has a conjugate acid. These pairs only ever differ by

exactly one hydrogen ion

Brønsted-Lowry Theory

Example Problems Identify the Brønsted-Lowry acid, base,

conjugate acid and conjugate base

NH3 + H2O NH4+ + OH-

AB CA CB

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

Amphoteric – substances that can act as both an acid and a base Examples: H2O, HCO3

-

Summary Of Theories

• Acids release H+

• Bases release OH-

Arrhenius

• Acids – proton donor• Bases – proton acceptor

Brønsted-Lowry

• Acids – electron acceptor• Bases – electron donorLewis

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 for his workers

The pH scale goes from 0 to 14 The acidity/basicity of the

solutions depends on the concentration of H+ (or H3O+)

The pH scale

pH < 7Acidic

pH = 7Neutral

pH > 7Basic

pH scale

Low pH values means a high concentration of H+ (acidic)

High pH values means a low concentration of H+ (basic)

Calculations of pH

The Self Ionization of Water In pure water (pH = 7), the concentrations of the

ions (H3O+ and OH-) are equal.

[H3O+]=[OH-]= 1x10-7 This is because water will spontaneously

dissociate naturally:

H2O (l) ↔ H3O+ (aq) + OH- (aq)

Writing the equilibrium expression for the self-ionization of water gives:

]][[ 3 OHOHKeq

The Self-ionization of Water

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 The ion product constant of water has its

own symbol: Kw

Unlike other equilibrium constants, the Kw will always be the same value

Calculations of H3O+/OH-

Example #1 What is the H3O+ concentration in a solution

with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-]

1 x 10-14 = [H3O+][3.0 x 10-4]

Mx 114-

14

103.310 x 3.0

10 x 1.0

_________________________

3.0 x 10-4 3.0 x 10-4

Calculations of H3O+/OH-

If the hydronium-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the hydroxide ion concentration in the solution?   Kw = [H3O+][OH-]

1 x 10-14 = [1 x 10-3][OH-]

Mxx

xOH 11

3

14

101100.1

101][

_________________________

1.0 x 10-3 1.0 x 10-3

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?)00010.0log(pH4pH Acid!

Calculating pH of a solution

Example #2 What is the pH of a solution where the

concentration of hydroxide ions is 0.0136 M? Is this an acid or a base?

Kw = [H3O+][OH-] pH = -log [H3O+] ]0136.0][[101 3

14 MOHxKw

Mxx

OH 1314

3 10353.70136.0

101][

1.12)10353.7log( 13 xpH

Base!

Calculating pH of a solution

Practice #1

Practice #2

Calculating H3O+/OH- from pH Example #1

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

[H3O+] = 10-pH

MxOH 43.3 100.510][ 3

Calculating H3O+/OH- from 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-]

MxOH . 60553 1091.810][

]][1091.8[101 614 OHxxKw

Mxx

xOH 9

6

14

1012.11091.8

101][

Calculating H3O+/OH- from 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

Naming Bases

Bases are soluble metal hydroxides Follow identical naming rules for ionic

compounds Examples

NaOH

Ba(OH)2

NH3

NH4+

Sodium hydroxide

Barium hydroxide

Ammonia

Ammonium

Naming Acids

Binary Acids (HX) If the acid has an anion that ends

in “-ide” use the following basic format to name the acid:“Hydro – root – ic acid”

ExampleHCl Hydrochloric acid

Naming Acids

Example HBr

Practice HI

H2S

Hydrobromic acid

Hydroiodic acid

Hydrosulfuric acid

Naming Acids

Polyatomic acids (aka oxoacids, HxAyOz) Name depends on the polyatomic

used:If polyatomic ends in “-ite”, replace

with “ous acid” If polyatomic ends in “-ate”, replace

with “ic acid”Trick: “I ate something icky”

Naming Acids

Examples HClO4

HClO2

Sulfuric acid

Perchloric acid

Chlorous acid

H2SO4