ACIDS AND BASES

Chemistry Ms. Piela

Key Characteristics of Acids & Bases

Acids Bases

Taste sour Tastes bitter Reacts with alkali metals (forms H 2 gas) Forms electrolyte solutions (conducts electricity) pH paper color: Red Slippery feel Forms electrolyte solutions (conducts electricity) pH paper color: Blue Neutralizes Bases 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: H Cl  H + + Cl  Properties of bases are due to the presence of OH ions  Example: Na OH  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 H 3 O + ion Oftentimes H + and H 3 O + HCl  are used interchangeably H + + Cl HCl (g) + H 2 O (l)  H 3 O + (aq) + Cl (aq)

Problems with the Arrhenius theory  Only deals with aqueous (solutions in water) solutions  Not all acids and bases contain H+ and OH- ions  Example: NH 3 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 + ions  Acids are proton (H + ) donors  Bases are substances that accept H + ions  Bases are proton (H + ) acceptors  Example: HBr + H 2 O  A B H 3 O + + Br -

Brønsted-Lowry Theory  The behavior of NH 3 NH 3 (aq) + H 2 O (l) can be understood now: ↔ NH 4 +

(aq)

+ OH -

(aq)

 NH 3 becomes NH 4 + , so NH 3 acceptor is a proton (or a Brønsted-Lowry base)  H 2 O becomes OH , so H 2 O is a proton donor (or a Brønsted-Lowry acid)

Brønsted-Lowry Theory

Brønsted-Lowry Theory

Brønsted-Lowry Theory  Example Problems  Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH 3 + H 2 O  NH 4 + + OH -

Brønsted-Lowry Theory  Example HCl

(g)

+ H 2 O

(l)

↔ H 3 O +

(aq)

+ Cl -

(aq)

HSO 4 + HCO 3 ↔ SO 4 -2 + H 2 CO 3

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: H 2 O, HCO 3 -

Summary Of Theories

Arrhenius

• Acids release H + • Bases release OH -

Brønsted Lowry

• Acids – proton donor • Bases – proton acceptor

Lewis

• Acids – electron acceptor • Bases – electron donor

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 H 3 O + )

The pH scale pH < 7 pH = 7 pH > 7

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 (H 3 O + and OH ) are equal .

 [H 3 O + ]=[OH ]= 1x10 -7 This is because water will spontaneously dissociate naturally: H 2 O (l) ↔ H 3 O +

(aq)

+ OH -

(aq)

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

K eq

 [

H

3

O

 ][

OH

 ]

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: K w  Unlike other equilibrium constants, the K w the same value will always be

Calculations of H 3 O + /OH  Example #1  What is the H 3 O + = 3.0 x 10 -4 M ?

concentration in a solution with [OH ] K w = [H 3 O + ][OH ] 1 x 10 -14 = [H 3 O + ][3.0 x 10 -4 ] 3.0 x 10 -4 3.0 x 10 -4  1.0

x 10  14 3.0

x 10 4  3 .

3

x

10  11

M

Calculations of H 3 O + /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? K w = [H 3 O + ][OH ] 1 x 10 -14 = [1 x 10 -3 ][OH ] 1.0 x 10 -3 1.0 x 10 -3 [

OH

 ]  1

x

10  14 1 .

0

x

10  3  1

x

10  11

M

Calculations of pH  pH can be expressed using the following equation:  pH = -log [H 3 O + ] Example #1 or [H 3 O + ] = 10 -pH  What is the pH of a solution with 0.00010 M H 3 O + ? Is this solution an acid or a base?

pH pH

   4 log( 0 .

00010 ) 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?

K w = [H 3 O + ][OH ] pH = -log [H 3 O + ] [

H

3

O

 ]

pH



K w

  1

x

10  14 1

x

10 0 .

0136  14   [

H

7 .

353 3

O x

 10 ][ 0 .

0136

M

 13

M

] Base!

 log( 7 .

353

x

10  13 )  12 .

1

Calculating pH of a solution  Practice #1  Practice #2

Calculating H 3 O + /OH from pH  Example #1  What is the hydronium ion concentration in fruit juice that has a pH of 3.3?

[H 3 O + ] = 10 -pH [

H

3

O

 ]  10  3 .

3  5 .

0

x

10  4

M

Calculating H 3 O + /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?

[H 3 O + ] = 10 -pH K w = [H 3 O + ][OH ] [

H

3

O

 ]  10  5

.

05

K w

 8 .

91

x

10  6

M

 1

x

10  14  [ 8 .

91

x

10  6 ][

OH

 ] [

OH

 ]  1

x

10  14 8 .

91

x

10  6  1 .

12

x

10  9

M

Calculating H 3 O + /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: HNO 3 (l) + H 2 O (l) ⇋ NO 3 -

(aq)

+ H 3 O +

(aq)

In a solution of nitric acid, no HNO 3 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 HNO 3 ) assume it is the same as the concentration of hydronium , H 3 O + , ions Given concentration of a strong base, assume it has the same concentration as the hydroxide , OH , ions

Strong Acids & Bases Ionize 100%  Example 1 M NaOH  1 M Na + 1 M + OH Na + OH Na + OH Na OH +

Weak Acids & Bases Ionize X%  Example 1 M HF  ?

M H + ?

M + F H + F HF F H + HF F H +

Naming Bases   Bases are soluble metal hydroxides  Follow identical naming rules for ionic compounds Examples  NaOH  Ba(OH) 2  NH 3  NH 4 +

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 ”  Example  HCl

Naming Acids  Example  HBr  Practice  HI  H 2 S

Naming Acids  Polyatomic acids (aka oxoacids, H x A y O z )  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  HClO 4  HClO 2  Sulfuric acid