When water self ionizes what products form

CH1000 Fundament als of Chemistry Module 4 – Chapter 15

Arrhenius Acids

Arrhenius Acid: An acid solution contains an excess of H+ ions.

Common Properties of Acids

1. Sour taste

2. Turns litmus paper pink

3. Reacts with:

Metals to produce H2 gas

Bases to yield water and a salt

Carbonates to give carbon dioxide

Arrhenius Bases

Arrhenius Bases: A basic solution contains an excess of OH– ions.

Common Properties of Bases

1. Bitter/caustic taste

2. Turns litmus paper blue

3. Slippery, soapy texture

4. Neutralizes acids

Brønsted-Lowry Acids and Bases

Lewis Acid-Bases

Summary of the Acid/Base

Theories

Reactions of Acids

Base Reactions

Bases can be amphoteric (act as either Brönsted acids or bases)

In general:

Zn(OH)2 (aq) + 2 HBr (aq) ZnBr2 (aq) + 2 H2O (l)

As a base:

NaOH and KOH can also react with metals.

2 NaOH (aq) + 2 Al (s) + 6 H2O (l) 2 NaAl(OH)4 (aq) + 3 H2 (g)

base + metal + water salt + hydrogen

Zn(OH)2 (aq) + 2 NaOH (aq) Na2Zn(OH)4 (aq)

As an acid:

Salts

Salts: products from acid-base reactions.

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

Salts are ionic compounds.

Salts contain a cation (a metal or ammonium ion) derived from the base and an anion (excluding oxide or hydroxide ions) derived from the acid.

Salts are generally crystalline compounds with high melting and boiling points.

Electrolytes and Nonelectrolytes

Electrolytes: compounds that conduct electricity when dissolved in water.

Nonelectrolytes: substances that do not conduct electricity when dissolved in water.

Ion movement causes conduction of electricity in water.

3 classes of compounds, acids, bases, and salts are electrolytes because they produce ions in water when they dissolve.

Comparing Solution Conductivity

(Sugar solution) (Salt solution)(Distilled water)

Dissociation of Electrolytes

Salts dissociate into their respective cations and anions when dissolved in water.

Hydrated sodium (purple) and chloride (green) ions

The negative end of the water dipole is attracted to the positive Na+ ion.

When NaCl dissolves in water, each ion is surrounded by several water molecules.

The permanent dipoles in the water molecules cause specific alignment around the ions.

NaCl (s) Na + (aq) + Cl- (aq)

Electrolyte Ionization

Ionization: process of ion formation in solution. Ionization results from the chemical reaction between a compound and water.

Acids ionize in water, producing the hydronium ion (H3O+) and a counter anion.

Bases ionize in water, producing the hydroxide ion (OH-) and a counter cation.

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

H3PO4 (aq) + H2O (l) H2PO4 - (aq) + H3O

+ (aq)

NH3 (aq) + H2O (l) OH - (aq) + NH4

+ (aq)

Strong and Weak Electrolytes

Strong electrolytes: undergo complete ionization in water. Example: HCl (strong acid)

Weak electrolytes: undergo incomplete ionization in water. Example: CH3COOH (weak acid)

HCl (left) is 100% ionized. CH3COOH exists primarily in the unionized form.

HF (aq) + H2O (l) F - (aq) + H3O

+ (aq)

Double arrows indicate incomplete ionization

(typically weak electrolytes).

Salts

Salts can dissociate into more than 2 ions, depending upon the compound.

A 1 M solution of NaCl produces a total of 2 M of ions.

A 1 M solution of CaCl2 produces a total of 3 M of ions.

NaCl (s) Na+ (aq) + Cl- (aq)

1M 1M 1M

CaCl2 (s) Ca 2+ (aq) + 2 Cl- (aq)

1M 1M 2M

Colligative Properties of Electrolyte Solutions

Colligative properties: depend only on the number of moles of dissolved particles present.

This must be taken into consideration when calculating freezing point depression or boiling point elevation due to the presence of solute particles.

Example: What is the boiling point elevation of a 1.5 m aqueous solution of CaCl2? (Kb for water is 0.512 ºC/m).

Because CaCl2 is a strong electrolyte, 3 mol of ions (1 mol Ca2+ and 2 mol Cl- ions) will be present in the solution.

ΔTb = 1.5 m CaCl2

= 2.3 ºC× 3 mol ions 1 mol CaCl2

0.512 ºC 1 m

×

Autoionization of Water

Pure water auto(self) ionizes according to the reaction:

Based on the reaction stoichiometry:

Concentration H3O+ = Concentration OH– = 1 x 10–7 M

[H3O+] x [OH–] = (1 x 10–7)2 = 1 x 10–14

When acid or base is present in water, [H3O+] and [OH-] change.

In acidic solutions, [H3O+] > [OH–].

In basic solutions, [H3O+] < [OH–].

H2O (l) + H2O (l) H3O + (aq) + OH– (aq)

Introduction to pH

The pH scale

Increasing acidity Increasing basicityHigh H3O +

Low OH- Low H3O

+

High OH-

In pure water, [H3O +] = 1 x 10-7 M, so

pH = - log(1 x 10-7) = 7

pH = - log[H3O +]

pH Calculations

pH = - log[H3O +]

[H3O +] = 1 x 10-5 M

[H3O +] = 2 x 10-5 M

If exactly 1

Exponent = pH pH = 5

If a number between 1 and 10

The pH is between the exponent and next lowest whole

number pH = 4.7

Generalizations

[H3O +] = 10-pH

pH

Neutralization

General Reaction

Example

Overall Ionic Equation: H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l)

All species are included; soluble compounds shown as ions.

Net Ionic Equation:

Spectator ions (orange) are removed from both sides.

acid + base salt + water

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

H+ (aq) + OH- (aq) H2O (l)

Titration

Titration: experiment where the volume of one reagent (titrant) required to react with a measured amount of another reagent is measured.

Titrations allow the amount of an acid or base present in a sample to be determined.

Indicators are used to signal the endpoint of a titration,

the point when enough titrant is added to react with the acid/base present.

Burets deliver measured amounts of the titrant into a solution of the unknown reagent.

Endpoin t

Net Ionic Equations

Rules for Writing Net Ionic Equations

1. Strong electrolytes are written as the corresponding ions. Example: NaOH (aq) is written as Na+(aq) + OH-(aq)

2. Weak electrolytes and nonelectrolytes are written as molecules. Example: CH3OH(aq), CH3COOH(aq), etc.

3. Solids and gases are written as their molecular forms.

4. The net ionic equation does not include spectator ions.

5. The net ionic equation must balance atoms and charge.

Acid Rain

1. Emission of nitrogen or sulfur oxides.

2. Transportation of these chemicals throughout the atmosphere.

3. Chemical reaction of the oxides with water.

4. This forms sulfuric and nitric acids.

5. Precipitation carries the acids to the ground.

General Process for Acid Rain Formation:

Acid rain: atmospheric precipitation more acidic

than typical.

Reading Review

1. What ion is present in an acid? What ion is present in a base?

2. What products are produced in a reaction between an acid and a metal oxide?

3. What type of reaction produces salts?

4. What is the pH range for acids?

5. What is titration?

Slide 1
Arrhenius Acids
Arrhenius Bases
Brønsted-Lowry Acids and Bases
Summary of the Acid/Base Theories
Reactions of Acids
Base Reactions
Salts
Electrolytes and Nonelectrolytes
Dissociation of Electrolytes
Electrolyte Ionization
Strong and Weak Electrolytes
Salts
Colligative Properties of Electrolyte Solutions
Autoionization of Water
Introduction to pH
pH Calculations
pH
Neutralization
Titration
Net Ionic Equations
Acid Rain
Reading Review