Chapter 14 Acids and Bases

Chapter 14Section 1 – Properties of Acids and BasesSection 2 – Acid Base TheoriesSection 3 – Acid Base Reactions

14.1 Properties of Acids and BasesList five general properties of aqueous acids

and bases.

Name common binary acids and oxyacids, given their chemical formulas.

List five acids commonly used in industry and the laboratory, and give two properties of each.

Define acid and base according to Arrhenius’s theory of ionization.

Explain the differences between strong and weak acids and bases.

Properties of: Acids Bases

1. Sour taste2. Conducts electricity3. Turns litmus paper

red4. Reacts with bases

to produce salts and water

5. Reacts with some metals and releases hydrogen gas

o Can you think of a reaction that this occurs?

1. Bitter taste2. Feels slippery3. Conducts

electric current4. Turns litmus

paper blue5. Reacts with

acids to produce salts and water

Binary Acids

Contains only two different elementsHydrogen and an electronegative element

(usually a halogen)

Nomenclature:hydro - _________ - ic acid

Diatomic Nomenclature

OxyacidContains hydrogen, oxygen, and a third

element(hydrogen with a polyatomic ion)

Nomenclature:

Acid Names

Oxyacids

Common Industrial AcidsSulfuric Acid

Sulfuric acid is the most commonly produced industrial chemical in the world.

Nitric Acid

Phosphoric AcidHydrochloric Acid

Concentrated solutions of hydrochloric acid are commonly referred to as muriatic acid.

Acetic AcidPure acetic acid is a clear, colorless, and

pungent-smelling liquid known as glacial acetic acid.

Arrhenius Acids and BasesArrhenius Acids:

Increases concentration of H+ ions in solution

Arrhenius Bases:Increases concentration of OH- ions in

solution

Arrhenius Acid Base Video

Arrhenius AcidsMolecular compounds with ionizable

hydrogen atomsWater solutions are known as aqueous

acidsElectrolytes

Acid StrengthStrong acid:

Ionizes completely in solution and is an electrolyte

Example: HCl, HClO4, HNO3

Weak acid:Releases few hydrogen ions in solution

Hydronium ions, anions and dissolved acid molecules present

Examples: HCN, Organic acids – HC2H3O2

Aqueous Acids

s aq + aq2H O –NaOH( ) Na ( ) OH ( )

aq + l aq + aq–3 2 4NH ( ) H O( ) NH ( ) OH ( )

Base StrengthStrong bases:

Ionic compounds containing metal cation and hydroxide ion (OH-)Dissociates in water

Weak bases:Molecular compounds do not follow Arrhenius

definition: Ammonia (NH3)Produces hydroxide ions when it reacts with water

molecules

Base Strength

Acid Base Strength Video

Acidic solution has greater [H3O+] Basic solution has greater [OH–]

14.2 Acid Base TheoriesDefine and recognize Brønsted-Lowry

acids and bases.

Define a Lewis acid and a Lewis base.

Name compounds that are acids under the Lewis definition but are not acids under the Brønsted-Lowry definition.

+ –3 4HCl NH NH Cl

l + aq aq + aq–2 3 4H O( ) NH ( ) NH ( ) OH ( )

Bronsted-Lowry Acid

Bronsted-Lowry Acid:Proton (H+) donor

Hydrogen chloride acts as a Bronsted-Lowry acid when it reacts with ammonia.

Water can also act as a Bronsted-Lowry acid

+ –3 4HCl NH NH Cl

Bronsted-Lowry Base

Bronsted-Lowry Base:Proton acceptor

Ammonia accepts a proton from hydrochloric acid.

Hydroxide ions produced in solution act as a Bronsted-Lowry base

+ –3 4HCl NH NH Cl

acid base

Bronsted-Lowry Acid Base Reactions

Protons are transferred from one reactant (the acid) to another (the base)

Brønsted-Lowry Acid Base Video

g + l) aq + aq–2 3HCl( ) H O( H O ( ) Cl ( )

Monoprotic AcidsCan donate only one proton (hydrogen ion)

per moleculeOne ionization step

Monoprotic and Diprotic Acids

l + l aq + aq–2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )

aq + l aq + l– 2–4 2 3 4HSO ( ) H O( ) H O ( ) SO ( )

Polyprotic AcidsDonates more than one proton per moleculesMultiple ionization steps

Diprotic – donates 2 protons Ex:Triprotic – donates 3 protons Ex:

Sulfuric acid solutions contain H3O+, HSO4-, SO4

- ions

1.

2.

aq + aq aq aq3 3 4H ( ) : NH ( ) [H — NH ] ( ) or [NH ] ( )

Lewis AcidLewis acid:

Atom, ion, or molecule that ACCEPTS an ELECTRON PAIR to form a covalent bond

A proton (hydrogen ion) is a Lewis acid

Lewis base:Atom, ion, or molecule that DONATES an

ELECTRON PAIR to form a covalent bond

aq + aq aq 3 3 3 3 2Ag ( ) 2 : NH ( ) [H N — Ag — NH ] ( ) or [Ag(NH ) ]

Lewis AcidA lewis acid might not include hydrogenSilver as a lewis acid:

Lewis Acid Base Video

Acid and Base Definitions

Acid Base Definitions Video

14.3 Acid Base ReactionsDescribe a conjugate acid, a conjugate

base, and an amphoteric compound.

Explain the process of neutralization.

Define acid rain, give examples of compounds that can cause acid rain, and describe effects of acid rain.

aq + l aq + aq–2 3HF( ) H O( ) F ( ) H O ( )

acid conjugate base

Conjugate Acid – BaseConjugate Base:

The species that remains after a Bronsted-Lowry acid has given up a proton

Conjugate Acid:The species that remains after a Bronsted-

Lowry base has accepted a proton

aq + l aq + aq–2 3HF( ) H O( ) F ( ) H O ( )

acid1 base2 base1

acid2

Conjugate Acid Base PairsMatch up the acid-base pairs

(proton donor-acceptor pairs)

g + l aq + aq–2 3HCl( ) H O( ) H O ( ) Cl ( )

strong acid base acid weak base

Strength of Acid Base PairsThe stronger the acid, the weaker the

conjugate baseThe stronger the base, the weaker the

conjugate acid

aq + l aq + aq–4 2 3 4HClO ( ) H O( ) H O ( ) ClO ( )

aq + l aq + aq–3 2 3 3CH COOH( ) H O( ) H O ( ) CH COO ( )

stronger acid stronger base weaker acid weaker base

weaker acid weaker base stronger acid stronger base

Proton transfer favors the production of the weaker acid and base.

Acid Base Strength

aq + l aq + aq–2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )

g + l aq aq–3 2 4NH ( ) H O( ) NH ( ) OH ( )

acid1 base2 acid2 base1

base1 acid2 acid1 base2

AmphotericAny species that can react as either an acid

or baseExample: water

Amphoteric Water Video

Amphoteric CompoundsCovalently bonded –OH group in an acid is

referred to as a hydroxyl groupMolecular compounds with hydroxyl groups

can be acidic or amphotericThe behavior of the compound is affected

by the number of oxygen atoms bonded to the atom connected to the –OH group

Oxyacids of Chlorine

2aq + aq aq lHCl( ) NaOH( ) NaCl( ) H O( )

Neutralization ReactionsWhat does it mean to neutralize

something?

Neutralization reactions: Hydronium and hydroxide ions react to form

water The left over cation and anion in solution

produce a salt (ionic compound)

Neutralization Reactions

Neutralization Reaction Video

g + l aq3 2 2 4SO ( ) H O( ) H SO ( )

Acid RainNO, NO2, CO2, SO2, and SO3 gases from

industrial processes can dissolve in atmospheric water to produce acidic solutions.

Very acidic rain is known as acid rain.Acid rain can erode statues and affect ecosystems.

Chapter 15 Acid Base Titration and pH

Chapter 15Section 1 – Aqueous Solutions and the

Concept of pHSection 2 – Determining pH and Titrations

15.1 Aqueous Solutions and pHDescribe the self-ionization of water.

Define pH, and give the pH of a neutral solution at 25°C.

Explain and use the pH scale.

Given [H3O+] or [OH−], find pH.

Given pH, find [H3O+] or [OH−].

l + l aq + aq–2 2 3H O( ) H O( ) H O ( ) OH ( )

Self Ionization of WaterTwo water molecules produce a hydronium ion

and hydroxide ion by proton transfer

In water at 25°C, [H3O+] = 1.0 ×10−7 M and [OH−] = 1.0 × 10−7 M

The ionization constant of water, Kw

Kw = [H3O+][OH−]

At 25OC

Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14

Kw = 1.0 x 10-14

Kw increases as temperature increases

Ion Concentration[H3O+] = [OH−]

neutral

[H3O+] > [OH−]

acidic

[H3O+] > 1.0 × 10−7 M

[OH−] > [H3O+]

basic

[OH−] > 1.0 × 10−7 M

s aq + aq2H O –NaOH( ) Na ( ) OH ( )

-14 -14

-123 – -2

1.0 10 1.0 10[H O ] 1.0 10 M

[OH ] 1.0 10

Calculating ConcentrationStrong acids and bases are considered

completely ionized or dissociated in weak aqueous solutions.

1 mol 1 mol 1 mol1.0 × 10−2 M NaOH => [OH−] = 1.0 ×

10−2 M[H3O+] is calculated using Kw

Kw = [H3O+][OH−] = 1.0 × 10−14

-14 -14

– -10-4

3

1.0 10 1.0 10[OH ] 5.0 10 M

[H O ] 2.0 10

Example Problem 1Given: [HCl] = 2.0 × 10−4 M

[H3O+] = ______________Unknown: [OH-] = ?

Kw = [H3O+][OH−] = 1.0 × 10−14

Example Problem 2A 1.0 10–4 M solution of nitric acid has

been prepared for a laboratory experiment.

Calculate [H3O+]

Calculate [OH–]

Example Problem 2 SolutionGiven:

Concentration of the solution = 1.0 × 10−4

M HNO3

Unknown: [H3O+]

[OH−]Solution:

Assume HNO3 dissociates completelyl + l aq + aq–

3 2 3 3HNO ( ) H O( ) H O ( ) NO ( )

1 mol 1 mol 1 mol 1 mol

Remember: [H3O+][OH−] = 1.0 × 10−14

Example Problem 2 Solution

pHThe pH of a solution is defined as the negative of

the common logarithm of the hydronium ion concentration, [H3O+].

pH = −log [H3O+]

Example: a neutral solution has a [H3O+] = 1×10−7

pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0

pH Values as Specified [H3O+]

pOHThe pOH of a solution is defined as the negative of

the common logarithm of the hydroxide ion concentration, [OH−].

pOH = −log [OH–]

pH + pOH = 14.0

Example: a neutral solution has a [OH–] = 1×10−7

the pH of this solution is?

pOH Video

The pH Scale

Approximate pH Range of Common Materials

Comparing pH and pOH Video

Significant FiguresThere must be as many significant figures to

the right of the decimal as there are in the number whose logarithm was found.

Example: [H3O+] = 1 × 10−7

one significant figure

pH = 7.0

The Circle of pH

pH

pOH

[ H3O+]

[ OH-]

-log [H3O+]

antilog (-pH)

antilog (-pOH)

-log [OH-]

[ H3O+] [ OH-] = 1.0x10-14pH + pOH = 14

pH of Weak Acids and BasesThe pH of solutions of weak acids and

weak bases must be measured experimentally.

The [H3O+] and [OH−] can then be calculated from the measured pH values.

pH Values of Some Common Materials

15.2 Determining pH and TitrationsDescribe how an acid-base indicator

functions.

Explain how to carry out an acid-base titration.

Calculate the molarity of a solution from titration data.

IndicatorsAcid-base indicators: compounds whose

colors are sensitive to pH.

The pH range over which an indicator changes color is called its transition interval.

pH MeterspH meter determines the pH of a

solution by measuring the voltage between the two electrodes that are placed in the solution.

The voltage changes as the hydronium ion concentration in the solution changes.

Measures pH more precisely than indicators

Color Ranges of Indicators

Color Ranges of Indicators

Color Ranges of Indicators

Antacids Video with Methyl Orange

TitrationNeutralization occurs when hydronium

ions and hydroxide ions are supplied in equal numbers by reactants.

H3O+(aq) + OH−(aq) 2H2O(l)Titration: the controlled addition and

measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration.

Titration Pointsequivalence point: point at which the

two solutions used in a titration are present in chemically equivalent amounts

end point: point in a titration at which an indicator changes color

Which indicator do I choose?pH less than 7

Indicators that change color at pH lower than 7 are used to determine the equivalence point of strong-acid/weak-base titrations.

strong-acid/weak-base titration = acidic.pH at 7

Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong-acid/strong base titrations.

strong acids/strong bases = salt solution with a pH of 7.

Which indicator do I choose?pH greater than 7

Indicators that change color at pH higher than 7 are used to determine the equivalence point of weak-acid/strong-base titrations.

weak-acid/strong-base = basic

Titration Curve Strong Acid and a Strong BaseEquivalence Point:

pH at 7

Titration Curve Weak Acid and a Strong BaseEquivalence Point:

pH higher than 7

Titration Curve Strong Acid and a Weak BaseEquivalence Point:

pH less than 7

Molarity and Titrationstandard solution: solution that contains the

precisely known concentration of a solute

primary standard: highly purified solid compound used to check the concentration of the known solution

The standard solution can be used to determine the molarity of another solution by titration.

Performing a Titration – Set up

Performing a Titration – Set up Acid

Performing a Titration – Starting Amount

Performing a Titration – Set up Base

Performing a Titration - Titrating

Performing a Titration – End Point

How to solve a titration problem:

1. Start with the balanced equation for the neutralization reaction, and determine the chemically equivalent amounts of the acid and base.

2. Determine the moles of acid (or base) from the known solution used during the titration.

3. Determine the moles of solute of the unknown solution used during the titration.

4. Determine the molarity of the unknown solution.

1 mol 1 mol 1 mol 1 mol

Molarity and TitrationDetermine the molarity of an acidic solution, 10

mL HCl, by titration

1. Titrate acid with a standard base solution20.00 mL of 5.0 × 10−3 M NaOH was titrated

2. Write the balanced neutralization reaction equation.

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

-3-45.0 10 mol NaOH 1 L

20 mL 1.0 10 mol NaOH used1 L 1000 mL

-4-21.0 10 mol HCl 1000 mL

1.0 10 M HCl10.0 mL 1 L

Molarity and Titration4. Calculate the number of moles of NaOH used

in the titration. 20.0 mL of 5.0 × 10−3 M NaOH is needed to reach

the end point

5. mol of HCl = mol NaOH = 1.0 × 10−4 mol

6. Calculate the molarity of the HCl solution

Example ProblemIn a titration, 27.4 mL of 0.0154 M

Ba(OH)2 is added to a 20.0 mL sample of HCl solution of unknown concentration until the equivalence point is reached. What is the molarity of the acid solution?

Ba(OH)2 + 2HCl BaCl2 + 2H2O1 mol 2 mol 1 mol 2

mol

Example Problem SolutionGiven: 27.4 mL of 0.0154 M Ba(OH)2

Unknown: ? M HCl of 20.0 mLSolution:

Write balanced equation:

2

2 2

mol Ba(OH) 1 LmL of Ba(OH) solution mol Ba(OH)

1 L 1000 mL

22

-42

0.0154 mol Ba(OH)24.7 mL of Ba(OH) solution

1 L1 L

4.22 10 mol Ba(OH)1000 mL

1. Calculate Moles of Given

–42

2

–4

2 mol HCl4.22 10 mol of Ba(OH)

1 mol Ba(OH)

8.44 10 mol HCl

2. Write a mole ratio: moles of base used to moles of acid produced

22

2 mol HClmol of Ba(OH) in known solution mol HCl

mol Ba(OH)

amount of solute in unknown solution (mol) 1000 mL

volume of unknown solution (mL) 1 L

molarity of unknown solution

3. Calculate Unknown Molarity

-2

-48.44 10 mol HCl 1000 mL

20.0 m4.22 10

L 1M l

LHC