Acid-Base Chemistry

Arrhenius acid: Substance that dissolves in water and provides H+ ions

Arrhenius base: Substance that dissolves in water and provides OH- ions

Examples: HCl H+ and Cl- Acid

NaOH Na+ + OH- Base

Bronsted Acid: Substance that donates proton to another substance

Bronsted base: Substance that accepts proton from another substance

Example: HCl + H2O H3O+ + Cl-

HCl acts as acid; H2O acts as base

In the Reverse Reaction,

H3O+ acts as an acid; Cl- acts as a base

Note: (H3O+ = hydronium ion = H+ = proton)

•A monoprotic acid contains one acidic proton.

HCl

•A diprotic acid contains two acidic protons.

H2SO4

•A triprotic acid contains three acidic protons.

H3PO4

•A Brønsted–Lowry acid may be neutral or it may carry a net positive or negative charge.

HCl, H3O+, HSO4−

Conjugate acid: Species formed after base accepts a proton

Conjugate base: Species remaining after an acid donates its proton

Conjugate acid-base pair: an acid and base on opposite sides of the equation that correspond to each other

Examples: HNO3 + H2O H3O+ + NO3-

acid base acid base

Conjugate pairs: HNO3 and NO3-

H2O and H3O+

Example: HS- + H2O H3O+ + S2-

Conjugate pairs: HS- and S2-

H2O and H3O+

Practice: HClO4 + H2O H3O+ + ClO4-

What are the conjugate pairs?HClO4 and ClO4-

H2O and H3O+

H Br + +

gain of H+

acid base base acid

loss of H+

H2O Br− H3O+

•HBr and Br− are a conjugate acid–base pair.

•H2O and H3O+ are a conjugate acid–base pair.

Note: The net charge must be the same on both sides of the equation.

Water can act as both an acid and a base (amphiprotic)!

HClO4 + H2O H3O+ + ClO4 (base)

NH3 + H2O OH- + NH4+ (acid)

Strengths of Acids and Bases

Strong acids/bases: dissociate completely when dissolved in solution

Weak acids/bases: dissociate only partially when dissolved in solution

Examples:

Strong Acid: HCl H+ + Cl- (100% dissociation)

Strong Base: NaOH Na+ + OH (100% dissociation)

Weak Acid: CH3COOH H+ + CH3COO- (1.3% dissociation)

Weak Base: NH3 + H+ NH4+

Table 9.1

Naming Acids

Binary Acids: hydo + root of anion + ic + “acid”

ex. HCl hydrochloric acid, HBr hydrobromic acid

HI

Polyatomic-based Acids: root of polyatomic ion + ic + “acid”

ex. H2SO4 sulfuric acid, H3PO4 phosphoric acid

H2CO3

HNO3

Hydroiodic acid

carbonic acid

nitric acid

The Self-Ionization of Water

H2O + H2O H3O+ + OH-

Pure water: [H3O+]=[OH- = 10-7 M (at 250C)

Neutral Solution: Any solution in which the concentrations of H3O+ and OH- ions are equal (10-7 M)

Acidic Solution: Solutions having a greater concentration of H3O+

than OH- ions ([H3O+] greater than 10-7 M)

Example: A solution with [H3O+] = 10-5 M

Basic Solution: solution having a greater concentration of OH- than H3O+ions ([H3O+] less than 10-7 M)

Example: A solution with [H3O+] = 10-12 M

The pH Scale

pH is a measure of acidity

Scale ranges from 0-14

pH = 7 Neutral

pH < 7 Acidic

pH > 7 Basic

pH represents the concentration of H+ ions in solution

Pure water: 1 x 10-7 moles H+ per liter and1 x 10-7 moles OH- per liter

Solutions with equal concentrations of and ions are called Neutral

Solutions with more than 1 x 10-7 moles H+ per liter are Acidic

Solutions with less than 1 x 10-7 moles H+ per liter are Basic

Note: [H+] x [OH-] = 10-14 (always!)

pH Scale Summary

pH scale refers to amount of H+ ions in solution

pH 7 is neutral, less than 7 is acidic, greater than 7 is basic

Lower pH = more acidic = more H+ ions

Higher pH = more basic = less H+ ions

Each pH unit represents a 10-fold change in H+ ion concentration!

pH 4 has 10 times more H+ ions than pH 5

pH 9 has 10 times fewer H+ ions than pH 8

Mathematical equation for pH:

pH is the negative log of the H3O+ concentration

pH = -log [H3O+]

pH = -log [H3O+]

Any number can be expressed as 10 raised to some exponent: y = 10x

Examples: 100 = 10 2

1000 = 103

0.10 = 10 -1

The log is that exponent!

100 = 10 2; Log of 100 =2

1000 = 103; Log of 1000 = 3

0.10 = 10 –1; Log of 0.10 = -1

We can also take the log of non-whole numbers, but we use our calculators for this.

Example: Find the log of 2.4 x 10-3

Enter 2.4 x 10-3into calculator

Press the “log” key

0.0024 “log” = -2.62

Therefore, 10-2.62 = 2.4 x 10-3

Calculating pH from [H3O+]

pH = -log [H3O+]

Enter [H3O+] into calculator

Press the “log” key

Change the sign

Example: [H3O+] = 1.0 x 10-7 M

pH = -log [H3O+]

pH = -log [1 x 10-7] = 7

Example: [H3O+] = 1 x 10-11M

pH = -log [H3O+]

pH = -log [1 x 10-11] = 11

Example: [H3O+] = 1 x 10-3 M

pH = -log [H3O+]

pH = -log [1 x 10-3] = 3

Example: [H3O+] = 4.2 x 10-5

pH = -log [H3O+]

Enter [H3O+] into calculator (4.2 x 10-5)

Press the “log” key (-4.3767507)

Change the sign (4.3767507)

pH = 4.3767507 = 4.4

Example: [H3O+] = 8.1 x 10-9

pH = -log [H3O+]

Enter [H3O+] into calculator (8.1 x 10-9)

Press the “log” key (-8.091515)

Change the sign (8.091515)

pH = 8.091515 = 8.1

Reactions Between Acids and Bases

Neutralization: reaction between an acid and a base; always produces salt and water

Example: Write a balanced equation for the reaction of hydrochloric acid with sodium hydroxide.

HCl(aq) + NaOH(aq) H—OH(l) + NaCl(aq)

Example: Write a balanced equation for the reaction of hydrochloric acid with magnesium hydroxide.

HCl + Mg(OH)2 H2O + MgCl2

2 2

Titration

Titration: a technique used to determine the concentration of an acid or base in a solution

•If we want to know the concentration of an acid solution, a base of known concentration is added slowly until the acid is neutralized.

•When the acid is neutralized:

# of moles of acid = # of moles of base

•This is called the end point of the titration.

Acid-Base Titration

• Titration is a laboratory procedure used to determine the molarity of an acid.

• In a titration, a base such as NaOH is added to a specific volume of an acid.

Base (NaOH)

Acid

solution

• A few drops of an indicator is added to the acid in the flask.

• The indicator changes color when the base (NaOH) has neutralized the acid.

• At the end point, the indicator has a permanent color.

• The volume of the base used to reach the end point is measured.

• The molarity of the acid is calculated using the neutralization equation for the reaction.

Determining an unknown molarity from titration datarequires three operations:

Moles ofbase

Moles ofbase

Molarity ofacid solution

Molarity ofacid solution

mole–moleconversion

factor

mole–moleconversion

factor

M (mol/L)conversion

factor

M (mol/L)conversion

factor

Moles ofacid

Moles ofacid

Volume ofbase solution

Volume ofbase solution

M (mol/L)conversion

factor

M (mol/L)conversion

factor[1]

[2]

[3]

HOW TO Determine the Molarity of an Acid Solution from Titration

Example: What is the molarity of an HCl solution if 22.5 mL of a 0.100 M NaOH solution are needed to titrate a 25.0 mL sample of the acid?

volume of base (NaOH)22.5 mL

conc. of base (NaOH)0.100 M

volume of acid (HCl)25.0 mL

conc. of acid (HCl)?

Step [1] Determine the number of moles of base used to neutralize the acid.

Volume ofbase solution

Volume ofbase solution

M (mol/L)conversion factor

M (mol/L)conversion factor

mL–L conversion factor

mL–L conversion factor

22.5 mL NaOH x 1 L1000 mL

x 0.100 mol NaOH1 L

=

0.00225 mol NaOH

Step [2] Determine the number of moles of acid that react from the balanced chemical equation.

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

0.00225 mol NaOH x

mole–moleconversion

factor

mole–moleconversion

factor

1 mol HCl1 mol NaOH

= 0.00225 mol HCl

Step [3] Determine the molarity of the acid from the number of moles and the known volume.

M = molL

= 0.00225 mol HCl25.0 mL solution

x 1000 mL1 L

= 0.0900 M HCl

Answer

mL–L conversion factor

mL–L conversion factor

Buffers

Buffer: a solution whose pH changes very little when acid or base is added.

Most buffers are solutions composed of roughlyequal amounts of

•a weak acid

•the salt of its conjugate base

The buffer resists change in pH because

•added base, −OH, reacts with the weak acid

•added acid, H3O+, reacts with the conjugate base

Buffers contain 2 compounds:

Compound with the ability to react with H+ ions

Compound with the ability to react with OH- ions

Example: HCO3- + H+ H2CO3

If acids (H+) are added, react with HCO3-

H2CO3 + OH- HCO3- + H2O

If OH- ions are added, react with H2CO3

H2CO3 is unstable: H2CO3 H2O + CO2

More Examples of a Buffer

CH3COOH + H2O H3O+ + CH3COO−

weak acid conjugatebase

If an acid is added to the following buffer equilibrium,

Adding moreproduct…

…drives the reaction to the left.

then the excess acid reacts with the conjugate base, so the overall pH does not change much.

CH3COOH + −OH H2O + CH3COO−

weak acid conjugatebase

If a base is added to the following buffer equilibrium,

Adding morereactant…

…drives the reaction to the right.

then the excess base reacts with the weak acid, so the overall pH does not change much.