Chapt. 18 – Acids and Bases 18.1Acids and Bases: An Introduction 18.2 Strengths of Acids and Bases 18.3Hydrogen Ions and pH 18.4 Neutralization

Chapt. 18 – Acids and Bases

18.1 Acids and Bases: An Introduction

18.2 Strengths of Acids and Bases

18.3 Hydrogen Ions and pH

18.4 Neutralization

Section 18.1 Introduction to Acids and Bases

• Identify the physical and chemical properties of acids and bases.

• Classify solutions as acidic, basic, or neutral.

• Compare the Arrhenius, Brønsted-Lowry, and Lewis models of acids and bases.

Different models help describe the behavior of acids and bases.

Section 18.1 Introduction to Acids and Bases

Key Concepts

• The concentrations of hydrogen ions and hydroxide ions determine whether an aqueous solution is acidic, basic, or neutral.

• An Arrhenius acid must contain an ionizable hydrogen atom. An Arrhennius base must contain an ionizable hydroxide group.

• A Brønsted-Lowry acid is a hydrogen ion donor. A Brønsted-Lowry base is a hydrogen ion acceptor.

• A Lewis acid accepts an electron pair. A Lewis base donates an electron pair.

Acid/Base Properties

Acids have sharp taste or are sour• Carbonic acid (CO2 in water) - soda

• Phosphoric acid (colas)• Citric acid (lemons, citrus fruits)• Acetic acid (vinegar)

Bases taste bitter, feel slippery

Taste and feel are extremely dangerous ways to classify chemicals!

Industrial Chemicals - US

Acids & bases occupy 7 of top 20 spots in list of most widely produced industrial chemicals

These are commodity chemicals – relatively cheap

Sulfuric has many industrial uses

Typical ReactionsSingle replacement reaction: reaction of metals & acids to form metal salt and hydrogen

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

Double replacement reaction: generation of gas & water

NaHCO3(s) + HCl(aq) NaCl(aq) + H2O(l) + CO2(g)

Base reacting with acid: neutralization

Practice

Reactions of acids

Problems 1(a-b), 2 page 635

Problems 64(a-b) page 672

Problems 1 – 2 page 989

Properties - Litmus

AcidBlue litmus Pink

BaseRed litmus Blue

Properties – Conductivity

Pure water essentially nonconductor

Any electrolyte (material that produces ions when dissolved in water) raises conductivity

Acids and bases are electrolytes – raise conductivity

Relation to Ions in SolutionAll aqueous solutions contain hydrogen ions (H+) and hydroxide ions (OH-)Acid – more H+ than OH-

Base – more OH- than H+

Neutral – equal concentrations

+

-

Water – Self Ionization

Water produces small but equal numbers of H+ and OH- (it is neutral)

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

H3O+ = hydronium ion (water molecule covalently bonded to H+)

Mostly OK to use H+ instead of H3O+

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

Arrhenius Acid/Base Model

Acid: Contains hydrogen & ionizes to produce H+ in aqueous solution

HCl(g) H+(aq) + Cl-(aq)

Base: Contains OH- group & ionizes to produce OH- in aqueous solution

NaOH(s) Na+(aq) + OH-(aq)

Model limitation: NH3(g) when dissolved in water produces OH- (base)

Brønsted-Lowry Acid/Base Model

Broadens Arrhenius model

Acid: Hydrogen ion donor

Base: Hydrogen ion acceptor

Uses concept of conjugate acid-base pairs

Conjugate Acid-Base Pairs

Generalized acid dissociation (n=1)

HX(aq) + H2O(l) H3O+(aq) + X-(aq)

Forward reaction:• HX donates H+ to H2O acid

• H2O is base in this process

Acid Base

Reverse reaction:• H3O+ donates H+ to X conjugate acid

• X- is conjugate base in this process

Conjugate

Acid

Conjugate

Base


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

+++ -

Acid? HF

Base? H2O

Conjugate Acid? H3O+

Conjugate Base? F-



All acids and bases that fit Arrhenius definition also fit Brønsted-Lowry definition

Brønsted-Lowry Base

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

NH3 not Arrenhius base

base conjugate acid

acid conjugatebase

Practice

Brønsted-Lowry Model

Problems 3(a-c), 4 page 640

Problem 9 page 643

Problems 19 (a-b) page 649

Problem 69, page 672


Amphoteric Substances

acid

base


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

H2O(l) acting as either an acid or a base

Compounds with this capability called amphoteric

Amphoteric Substances

Can act as either acid or base

Water most common example

Also common is bicarbonate ion, HCO3-

HCO3-(aq) + OH-(aq) CO3

2-(aq) + H2O(l)

acid base conjugate conjugate base acid

HCO3-(aq) + H3O+(aq) H2CO3(aq) + H2O(l)

base acid conjugate conjugate acid base

Lewis Acid/Base ModelMost general model of acids & bases

Same G.N. Lewis that gave us Lewis electron dot structures and shared electron pair model for covalent bond

Acid: electron pair acceptor ( / share)

Base: electron pair donor ( / share)

Includes all substances classified as Brønsted-Lowry acids & bases & many more; also reactions that do not necessarily occur in aqueous solution

Other Lewis Bases

:Cl:....

-H - N - H

H

..

(HCl) (NH4+)

Lewis Acid/Base Model

Formation of HF molecule

Fluoride ion donates electron pair; empty 1s orbital of H accepts pair (covalent bond formed)

Lewis Acid/Base Model

Reaction of gaseous boron trifluoride and gaseous ammonia

Electron deficient B accepts lone pair from ammonia (formation of coordinate covalent bond)

Lewis Acid/Base ModelGaseous sulfur trioxide reacts with solid magnesium oxide

SO3(g) + MgO(s) MgSO4(s)

Lewis acid/base part of reaction:

Will return to this reaction shortly

3 Acid-Base Models (Table 18.2)

Mono- and Polyprotic Acids

Monoprotic – capable of donating only one hydrogen ion

• HCl

Polyprotic – can donate more than one• H2SO4

Above examples simple - inorganic acids, all hydrogens ionize

More complicated for organic acids

Mono- and Polyprotic Acids - HnX

HnX acids that dissociate into nH+(aq) ions and Xn-(aq) ions

X=Cl hydrochloric acid n=1

X=NO3 nitric acid n=1

X=C2H3O2 acetic acid n=1

X=SO4 sulfuric acid n=2

X=CO3 carbonic acid n=2

X=PO4 phosphoric acid n=3

mon

opro

ticpo

lypr

otic

Step-Wise Ionization

Sulfuric acid

H2SO4(aq) + H2O(l) H3O+(aq) + HSO4-(aq)

HSO4-(aq) + H2O(l) H3O+(aq) + SO4

2-(aq)

Conjugate

BaseAcid

Step-Wise Ionization

Phosphoric acid

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

H2PO4-(aq) + H2O(l) H3O+(aq) + HPO4

2-(aq)

HPO42-(aq) + H2O(l) H3O+(aq) + PO4

3-(aq)

Conjugate

BaseAcid

Amphoteric

H1

H2H3

Monoprotic AcidsAcetic acid: organic acid; HF: inorganic acid; each has one ionizable hydrogen atom (one that is involved in a polar bond)

Organic AcidsOrganic acids contain carboxylic acid group, COOH

R - C - O - H

=O

Formula often written as RCOOHR = (almost) anythingFor acetic acid, R = CH3 (methyl)CH3 hydrogens not ionizable – monoprotic even though has 4 H atoms

Ionizable hydrogen

R - C – O + H+

=O–

Organic Acids

Malonic acid (1,3-Propanedioic acid)

How many ionizable hydrogens?

?

2

Anhydrides

Oxides that become acids or bases when added to water

Nonmetal oxides (C, S, N) – acids• CO2(g) + H2O(l) H2CO3 (Carbonic Ac)

• N2O3(g) + H2O(l) 2HNO2 (Nitrous Ac)

• SO3(g) + H2O(l) H2SO4 (Sulfuric Ac)

Metal oxides – bases• CaO(s) + H2O(l) Ca2+(aq) + 2OH-(aq)

Anhydrides and Acid RainNatural rain water: pH 5.6

• Naturally slightly acidic due mainly to CO2

Acid rain: Rain water with pH < 5.6Due to pollutant gases dissolved in air (nonmetal oxides = acid anhydrides)

• SO2, SO3, NO2

Acid rain linked to damage in ecosystems and structures

Damage from Acid RainReacts with metals and materials that contain carbonates

Damages bridges, cars, and other metallic structures

Damages buildings and other structures made of limestone or cement

Acidifies lakes affecting aquatic life

Dissolves and leaches minerals from soil• Makes it difficult for trees to grow

Practice

Acids and Bases Intro – other topics

(neutrality, Arrenhius model, polyprotic acids, ionizable hydrogens, anhydrides)

Problems 5 – 8, 10 – 11 page 643

Problems 55 - 63, 64 (c-d) page 672






18.4 Neutralization

Section 18.2 Strengths of Acids and Bases

• Relate the strength of an acid or base to its degree of ionization.

• Compare the strength of a weak acid with the strength of its conjugate base.

• Explain the relationship between the strengths of acids and bases and the values of their ionization constants.

In solution, strong acids and bases ionize completely, but weak acids and bases ionize only partially.

Section 18.2 Strengths of Acids and Bases

Key Concepts

• Strong acids and strong bases are completely ionized in a dilute aqueous solution. Weak acids and weak bases are partially ionized in a dilute aqueous solution.

• For weak acids and weak bases, the value of the acid or base ionization constant is a measure of the strength of the acid or base.

Acid Strength and Concentration

...are not the same thing

e.g. 1.0 M HCl and 1.0 M H2CO3

stronger acid

same concentrations!Can have a dilute solution of a strong acid and a concentrated solution of a weak acid

Acid Strength

0.10 M HCl conducts electricity better than 0.10 M HC2H3O2 (acetic acid)

Hyd

roch

loric

Ace

tic

Acid Strength

Strong acids dissociate completely

Limited number of them (see Table 18.3)

HCl hydrochloric

HBr hydrobromic

HI hydroiodic

3 binary halogen acids

except HF

HClO4 perchloric

HNO3 nitric

H2SO4 sulfuric

Relationship Between Strengths of Acids and Their Conjugate Bases

The stronger the acid, the weaker is the attraction of the ionizable H for the rest of the molecule

The better the acid is at donating H+, the worse its conjugate base will be at accepting an H+

Strong acid: HCl + H2O → Cl– + H3O+ Weak conj base

Weak acid: HF + H2O F– + H3O+ Strong conj base

Acid Strength & Brønsted-Lowry Model

Strong acid HA has weak conj. base A-

H2O(l) stronger base than A-

Reaction equilibrium lies far to right

+ A-

conjugate base of HA

HA + H2O(l)

acid base

H3O+(aq)

conjugateacid of B

Strong/Weak vs Concentrated/Dilute

Concentrated/Dilute• Number of acid/base molecules per

volume• 6 M HCl concentrated, 0.01 M HCl dilute

Strong/Weak• Measure of degree of ionization• HCl strong acid• CH3COOH (acetic) weak acid

Acid Strength & Brønsted-Lowry Model

Weak acid HA has strong conj. base A-

A- stronger base than H2O(l) Reaction equilibrium lies far to left

+ A-

conjugate base of HA

HA + H2O(l)

acid base

H3O+(aq)

conjugateacid of B

Acid Ionization ConstantsFocus on equilibrium for weak acidsHCN(aq) + H2O(l) H3O+(aq) + CN-(aq)

Write K following standard rules

Ka = [H3O+(aq)] [CN-(aq)] = 6.2x10-10 [HCN(aq)]

Ka acid ionization constant

Small number for weak acids – major species is undissociated form of acid

Table 18.4, page 647 lists values at 25C

Acid Ionization Constants - pKa

HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)

Ka = [H3O+(aq)] [CN-(aq)] = 6.2x10-10 [HCN(aq)]

Ka acid ionization constant

Many texts will utilize quantity pKa

pKa = log(Ka)

For HCN, pKa = 9.21

[Note: SF after DP for log based on SF of number in front of power of ten]

Weak Acids & Ka Values

Strongest Acids

Weakest Acids

Acid Ka Value Conjugate Base

Weakest Bases

Strongest Bases

HF 6.3 x 10-4 F-

HNO2 5.6 x 10-4 NO2-

HCO2H 1.8 x 10-4 HCO2-

CH3CO2H 1.8 x 10-5 CH3CO2-

HOCl 4.0 x 10-8 OCl-

NH4+ 5.6 x 10-10 NH3

HCN 6.2 x 10-10 CN-

Practice

Strengths of Acids


Problems 65 – 68, 71 page 672

Problems 17, 18, 21 page 649


Dilute/concentrated vs weak/strong

Problem 74 page 672

Strong Bases

Alkali & some alkaline earth metal hydroxides are strong Arrhenius bases

M(OH)n(s) Mn+(aq) + n OH-(aq)

Dissociate completelyLimited number (see table 18.5)

NaOH KOH RbOH CsOH (alkali metals)Ca(OH)2, Sr(OH)2, Ba(OH)2 (alkaline

earths)

Strong BasesPicture complicated by low solubility of some

Ca(OH)2(s) Ca2+(aq) + 2OH-(aq)

Calcium hydroxide Ksp = 6.5x10-6

Base strength based on dissolved formCa(OH)2 strong base – dissolved form

[Ca(OH)2(aq)] completely dissociates

Weak Bases

Weak bases mostly Brønsted-Lowry bases

B(aq) + H2O(l) OH-(aq) + BH+(aq)

Most common weak base either ammonia or amine (organic compound with –NH2)

CH3NH2(aq) + H2O(l) OH-(aq) + CH3NH3+(aq)

Kb = [CH3NH3+(aq)] [OH-(aq)] = 4.3x10-10

[CH3NH2(aq)]

Kb= base ionization constant (see table 18.6)

Weak Base Equilibria

Ammonia NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)

Methyl amine

CH3NH2(aq) + H2O(l) CH3NH3+(aq) +

OH−(aq)

Ethyl amineC2H5NH2(aq) + H2O(l) C2H5NH3

+(aq) + OH−(aq)

BicarbonateHCO3

−(aq) + H2O(l) H2CO3 (aq) + OH−

(aq)

Practice

Strengths of Bases

Problems 15 (a-d), 16, 20 page 649

Problems 70, 72, 73 page 672






18.4 Neutralization

Section 18.3 Hydrogen Ions and pH

• Explain pH and pOH.

• Relate pH and pOH to the ion product constant for water.

• Calculate the pH and pOH of aqueous solutions.

pH and pOH are logarithmic scales that express the concentrations of hydrogen ions and hydroxide ions in aqueous solutions.

Section 18.3 Hydrogen Ions and pH

Key Concepts

• The ion product constant for water, Kw, equals the product of the H+ ion concentration and the OH– ion concentration.

Kw = [H+][OH–]

• The pH of a solution is the negative log of the hydrogen ion concentration. The pOH is the negative log of the hydroxide ion concentration. pH plus pOH equals 14.

pH = –log [H+]

pOH = –log [OH–]

pH + pOH = 14.00• A neutral solution has a pH of 7.0 and a pOH of 7.0 because

the concentrations of hydrogen ions and hydroxide ions are equal.

Ion Product Constant for Water (Kw)

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

or


Keq = [H3O+(aq)] [OH-(aq)]

= [H+(aq)] [OH-(aq)]

Keq given special symbol, Kw

Kw = 1.0 x 10-14 at 25oC

Ion product=Kw in any aqueous solution

Ion Product Constant for Water (Kw)

Kw = [H3O+(aq)] [OH-(aq)] = 1.0 x 10-14

In pure water only source of ions is water dissociation, so at 25oC (298 K)

[H3O+] = [OH-] = 1.0 x 10-7 M

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

+ ++ -

[H+] and [OH-] for General Case


Kw = [H+(aq)] [OH-(aq)] = 1.0 x 10-14

H+ or OH- can come from source other than water (i.e., an acid or a base)

Le Châteliers principle equilibrium shifts to left, depressing either H+ or OH- concentrations

Kw must remain constant

Calculation of [H+] and [OH-]

Example problem 18.1, page 651

Kw = [H+(aq)] [OH-(aq)] = 1.0 x 10-14

Given [H+] = 1.0x10-5 M, [OH-] = ?

[OH-]= Kw / [H+]

[OH-]= 1.0 x 10-14 /1.0x10-5 = 1.0x10-9 M

Practice

Kw and Water Ionization Equilibrium

Problems 22(a-d), 23 page 651

Problems 37, 38 page 658

Problems 77 – 78, 80 page 673


pH

Defined as

pH = – log[H+]

Book says scale 0 to 14

Not strictly true! 10 M HCl pH = -1

10 M NaOH pH = 15

Remember log properties

pH 3 solution has 100x [H+] of pH 5

pH 0 - 7 acid pH 7 - 14 basic

pH – Common MaterialsSimilar to Fig 18.14, page 652

pH – Common Materials Similar to Fig 18.14, page 652

Calculating pH

Example problem 18.2, page 653

pH of neutral solution at 298 K?

Kw = 1.00x10-14 = [H+] [OH-]

Neutral solution, [H+] = [OH-]

[H+] = 1.00x10-7

pH = – log[H+] = 7.00

Practice

Calculate pH from [H+]

Problems 24(a-b), 25 page 653



Problem 13 page 989

pOH

pOH = – log[OH-]

pOH 0 - 7 basic pH 7 - 14 acidic

pH + pOH = 14.00

Calculating pOH & pH

Sample problem 18.3, page 654

Ammonia cleaner [OH-] = 4.0x10-3 M

pOH, pH?

pOH = – log[OH-] = 2.40

pH + pOH = 14.00 pH = 11.60

Practice

Calculate pH or pOH from [H+] or [OH-]

Problems 27(a-d), 28(a-b), 29 page 654

Problems 41, 42(c-d), 43 page 658


Problems 14 - 15 page 989

Calculating Ion Concs from pH[H+] = antilog(-pH)

If pH=5, [H+] = antilog(-5) = 10-5

Problem 18.4, page 655

Blood pH = 7.40 @ 298 K; [H+], [OH-] ?

[H+] = antilog(-7.40) = 4.0x10-8 M

[OH-] = Kw / [H+] = 1.0x10-14 / 4.0x10-8

[OH-] = 2.5 x10-7 M(book uses pH + pOH = 14.00 and antilog(-pOH)

Practice

Calculate [H+] or [OH-] from pH


Problem 40 page 658

Problem 79 page 673 (also other topics)


pH of Strong Acids & Bases

Strong acids and bases ~100% ionizedHCl(aq) H+(aq) + Cl-(aq)

(Also HBr, HI, HClO4, HNO3, H2SO4) [H+] = Acid concentration

NaOH(aq) Na+(aq) + OH-(aq) [OH-] = Base concentration

Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq) [OH-] = 2 x Base concentration

Using pH to Calculate Ka

HF(aq) H+(aq) + F-(aq) (weak acid)

pH of 0.100 M solution = 2.12; Ka?

Ka = [H+(aq)] [F-(aq)] / [HF(aq)]

[H+(aq)] = antilog(-2.12) = 7.6x10-3 M (2 SF)

If ignore contribution of H2O to H+(aq)

[F-(aq)] = [H+(aq)]

[HF(aq)] = 0.100 M – 7.6x10-3 M

Ka = 7.6x10-3 7.6x10-3 / 0.092 = 6.3x10-4

Using pH to Calculate Ka

Sample problem 18.5, page 657

pH 0.100 M formic acid = 2.38; Ka?

HCOOH(aq) H+(aq) + HCOO-(aq)

[H+] = antilog(-2.38) = 4.2x10-3 (2 SF)

[HCOO-(aq)] = [H+(aq)]

[HCOOH(aq)] = 0.100 M - 4.2x10-3 M

Ka = [H+(aq)] [HCOO-(aq)] / [HCOOH(aq)]

Ka = 4.2x10-3 4.2x10-3 / 0.096 = 1.8x10-4

Practice

Use pH or pOH to calcuate Ka

Problems 32(a-b), 33(a-c), 34 page 657

Problem 39 page 658

Problem 84* (avoid) page 673


* Relies on assumption that 1st ionization of chromic acid acts like strong acid (no hydrolysis of HCrO4

-)

Measuring pH

Indicators – quick, cheap but not precise

pH meter – can measure to 0.001 pH units under ideal conditions with proper calibration

Methods for Measuring pH

pH meterpH (indicator) paper





18.4 Neutralization

Section 18.4 Neutralization

• Write chemical equations for neutralization reactions.

• Explain how neutralization reactions are used in acid-base titrations.

• Compare the properties of buffered and unbuffered solutions.

In a neutralization reaction, an acid reacts with a base to produce a salt and water.

Section 18.4 Neutralization

Key Concepts

• In a neutralization reaction, an acid and a base react to form a salt and water.

• The net ionic equation for the neutralization of a strong acid by a strong base is H+(aq) + OH–(aq) → H2O(l).

• Titration is the process in which an acid-base neutralization reaction is used to determine the concentration of a solution.

• Buffered solutions contain mixtures of molecules and ions that resist changes in pH.

Neutralization Reaction

Double replacement reaction of acid and base in aqueous solution to produce a salt and water

• Salt – ionic compound, cation from base and anion from acid

Mg(OH)2(aq) + 2HCl(aq) MgCl2(aq) + 2H2O(l)

Base Acid Salt Water

Neutralization & Net Ionic Equation

Salt formed from neutralization of strong acid and strong base will be a strong electrolyte (fully ionized)HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Complete ionic equation:

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

Net ionic: H+(aq) + OH-(aq) H2O(l)

Practice

Neutralization reactions

Problem 49 page 668



Acid-Base TitrationMethod for determining concentration of an unknown acid (base) solution by reacting it with a base (acid) of known concentration HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Stoichiometry applies - if know concentration & volume of NaOH needed to neutralize & initial volume of HCl, can calculate initial concentration of HCl

TitrationTechnique that uses reaction stoichiometry to determine concentration of unknown solutionTitrant (known solution) added from buretIndicators – color-changing chemicals added to help determine when reaction complete (not required if pH measured)Endpoint of titration occurs when reaction complete

Acid-Base TitrationTitrant (solution in buret) is a baseAs base added to acid, H+

reacts with OH– to form H2O; still have excess acid present, so no color changeAt titration’s endpoint,just enough base has been added to neutralize allacid. At this point,indicator changes color

Acid-Base Titration Steps1. Measure volume of unkown.

Determine initial pH

2. Fill buret with titrating solution of known concentration (standard solution)

3. Slowly add standard solution while recording pH until neutralization reaction reaches stoichiometric point (equivalence point)

At equivalence

point, pH changes

rapidly with small

addition of NaOH

Acid-Base Titration - Indicators

Often convenient to use pH indicator to determine equivalence point (point at which number of moles of base added equals number of moles of acid originally present) [OHadd

-(aq) = H+start(aq) for 1:1 type system)

End point: pH at which indicator changes color

Must choose indicator so pH at end point and pH at equivalence point are ~ same

Acid-Base Titration Equivalence Point

Equivalence point for strong acid-strong base titration: pH 7

For weak acid-strong base, pH > 7

For strong acid-weak base, pH < 7

Different from 7 due to salt hydrolysis (topic discussed later)

Must choose pH indicator end point accordingly

Matching of equivalence

point and pH indicator end point – titration of strong acid with strong

base

Matching of equivalence

point and pH indicator end point – titration of weak acid with strong

base

Acid Base Indicators

are weak acids or bases themselves

Hin(aq) + H2O(l) H3O+(aq) + In-

(aq)

Color A Color B

Acid Base Indicators

Ka =[H3O+] [In-]

[Hin]

[In-][Hin]

=Ka

[H3O+]

i.e. ratio of colors changes with pH

Hin(aq) + H2O(l) H3O+(aq) + In-

(aq)

Indicators – See Fig 18.24, p 662

3 Very Common Indicators

Step 1

Known volume of

unknown in flask with indicator

Known concentration

of titrant in buret

Step 2

Titrant added

Pink color appears

momentarily but then

disappears upon mixing

Step 3

Just enough titrant added so that pink

color permanently

remains

Titration Calculation - Prob. 18.6 p 664

18.28 mL of 0.1000 M NaOH for 25.00 mL of formic acid (HCOOH)

1. Write balanced neutralization eqn.

HCOOH(aq)+ NaOH(aq) HCOONa(aq)+H2O(l)

2. Calculate moles titrant added

18.3 mL NaOH x 1 L NaOH = 0.0183 L NaOH 1000 mL NaOH

0.0183 L NaOH x 0.100 mol NaOH = 0.00183 mol L NaOH NaOH

Acid-Base Titration CalculationHCOOH(aq)+ NaOH(aq) HCOONa(aq)+H2O(l)

3. Use mole ratio: calculate moles unknown1.83x10-3 mol NaOH x 1 mol HCOOH

1 mol NaOH

= 1.83x10-3 mol HCOOH4. Calculate concentration of unknown

using moles and initial volumeMHCOOH = 1.83x10-3 mol HCOOH

0.0250 L HCOOH MHCOOH = 7.31x10-2 M

Practice

Acid-Base Titration



Problems 86 – 89, 92 – 93 page 673


Salt HydrolysisPut salt in pure water

Unless salt product of strong acid/strong base reaction, pH won’t be neutral

Neutral Basic AcidicNaNO3 KF NH4Cl

Salt Hydrolysis

KF(s) K+(aq) + F-(aq)

KOH strong base, HF weak acid (see table 18.4, page 647)

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

F-(aq): Brønsted-Lowry base

Solution basic (OH- generated)

Salt Hydrolysis

NH4Cl(s) NH4+(aq) + Cl-(aq)

NH3 weak base, HCl strong acid

NH4+(aq) + H2O(l) NH3(aq) + H3O+ (aq)

NH4+(aq): Brønsted-Lowry acid

Solution acidic (H3O+ generated)

Practice

Salt Hydrolysis





Buffered Solutions

Buffer: Solution that resists change in pH when limited amounts of acid or base added

Add 0.01 mol HCl to solution• 1 L water – pH change: 7.0 to 2.0• Buffer – typical pH change 0.1 unit

Buffer Composition

Mixture of either:(a) Weak acid & salt of its conjugate

baseHF (weak acid) + KF (F- conj base)

(b) Weak base & salt of its conjugate acid

NH3 (weak base) + NH4Cl (NH4+ conj

acid)

Standard buffers mostly made using (a)Having both acid & base present makes buffer resist pH changes

How Buffers Work

Weak acid in buffer mixture can neutralize added base

Conjugate base in buffer mixture can neutralize added acid

Net result: little to no change in solution pH in either case

H2O

How Buffers Work

HA + H3O+A−A−

AddedH3O+

NewHA

HA

H2O

HAHA + H3O+

A−

AddedHO−

NewA−

A−

How Buffers Work

Equimolar Buffer Systems Table 18.7

HF/KF Buffer0.1 M HF (weak ac) + 0.1 M KF (F- conj base)

HF(aq) H+(aq) + F-(aq)F-(aq) from 2 sources:

• Dissociation of HF(aq)• Dissociation of KF(aq)

Add acid – equilibrium shifts to left so H+ consumed

Add base – H+ neutralized to H2O, equil. shifts to right to generate more H+

Either way, pH change minimized

Acetic Acid / Sodium Acetate Buffer

Mix equal volumes of 1 M HC2H3O2 and 1 M NaC2H3O2: buffer pH = 4.75

Add 10 mL of 0.1 M HCl to: • 1 L of buffer, pH = 4.75• 1 L of DI water, pH = 3.0

Add 10 mL of 0.1 M NaOH to:• 1 L of buffer, pH = 4.75• 1 L of DI water, pH = 11.0

Buffer Capacity

Amount of acid or base buffer can absorb without significant pH change0.1 M HF (weak ac) + 0.1 M KF (F- conj base)

HF(aq) H+(aq) + F-(aq)

Add more H+ than available F-, pH must drop

Remove more H+ than can be replaced by remaining HF, pH must rise

pH of Buffer

Easy to determine for equimolar buffers

Equimolar mix NaH2PO4 & Na2HPO4

H2PO4-(aq) H+(aq) + HPO4

2-(aq)

Acid conj. Base

Ka = [H+(aq)][HPO42-(aq)] / [H2PO4

-(aq)]

Since [HPO42-(aq)] = [H2PO4

-(aq)]

pH = -log[H+] = -log(Ka) = -log(6.2x10-8)

pH = 7.21 (see table 18.7, p 667 for others)

Practice

Buffers


Problems 88, 102, 103 pages 673-4

Documents

Chapt. 18 – Acids and Bases 18.1Acids and Bases: An Introduction 18.2 Strengths of Acids and Bases 18.3Hydrogen Ions and pH 18.4 Neutralization