N Acid/Base Definitions n Types of Acids/bases n Polyprotic Acids n The Ion Product for Water n The...

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Acid/Base DefinitionsAcid/Base Definitions Types of Acids/basesTypes of Acids/bases

Polyprotic AcidsPolyprotic Acids The Ion Product for WaterThe Ion Product for Water

The pH and Other “p” ScalesThe pH and Other “p” Scales Aqueous Solutions of Acids and BasesAqueous Solutions of Acids and Bases

HydrolysisHydrolysis The Common Ion EffectThe Common Ion Effect

Buffer SolutionsBuffer Solutions Indicators and TitrationsIndicators and Titrations

Chapter 15. Acids & Bases

Types of Reactions a) Precipitation Reactions. Ionic compounds or salts b) Acid/base Reactions. Acids and Bases c) Redox Reactions. Oxidizing & Reducing agents

What are Acids &Bases?

Definition?

a) Arrhenius

b) Bronsted-Lowry

c) Lewis

Arrhenius definitions AcidAcid Anything that produces hydrogen

ions in a water solution.» HCl (aq) H+ + Cl-

BaseBase Anything that produces hydroxide ions in a water solution.

» NaOH (aq) Na+ + OH-

Arrhenius definitions are limited to aqueous solutions.

Acid base reactions: HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Brønsted-Lowry definitions

Expands the Arrhenius definitions

AcidAcid Proton donor

BaseBase Proton acceptor This definition explains how substances

like ammonia can act as bases.

Eg. HCl(g) + NH3(g) ------> NH4Cl(s)

HCl (acid), NH3 (base).

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

Lewis Definition Lewis was successful in including acid and

bases without proton or hydroxyl ions.

Lewis Acid: A substance that accepts an

electron pair.

Lewis base: A substance that donates an

electron pair.

E.g. BF3(g) + :NH3(g) F3B:NH3(s)

Types of Acids and Bases• Binary acids • Oxyacid • Organic acids • Acidic oxides• Basic oxides • Amine• Polyprotic acids

Binary Acids Compounds containing acidic protons

bonded to a more electronegative atom.

e.g. HF, HCl, HBr, HI, H2S The acidity of the haloacid (HX; X = Cl, Br, I, F) Series increase in the following order: HF < HCl < HBr < HI

Oxyacids Compounds containing acidic - OH groups

in the molecule. Acidity of H2SO4 is greater than H2SO3

because of the extra O (oxygens) The order of acidity of oxyacids from the a

halogen (Cl, Br, or I) shows a similar trend.

HClO4 > HClO3 > HClO2 >HClO perchloric chloric chlorus hyphochlorus

Acidic Oxides

These are usually oxides of non-metallic elements such as P, S and N.

E.g. NO2, SO2, SO3, CO2 They produce oxyacids when

dissolved in water

Basic Oxides

Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.

e.g. CaO + H2O --> Ca(OH)2

Protic Acids

Monoprotic Acids: The form protic refers to acidity or protons. Monoprotic acids have only one acidic proton. e.g. HCl.

Polyprotic Acids: They have more than one acidic proton.

e.g. H2SO4 - diprotic acid

H3PO4 - triprotic acid.

Amines

Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted or Lewis acid/base definitions.

What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions:

a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)

b)HCl(g) + NH3(g) ---> NH4Cl(s)

c)BF3(g) + NH3(g) ---> F3B:NH3(s)

d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)

Common acids and bases AcidsAcids Formula Molarity* nitric HNO3 16

hydrochloric HCl 12 sulfuric H2SO4 18

acetic HC2H3O2 18

BasesBases ammonia NH3(aq) 15 sodium hydroxide NaOH solid

*undiluted.

Acids and bases

AcidicAcidic BasicBasic

– Citrus fruits Baking soda– Aspirin Detergents– Coca Cola Ammonia

cleaners– Vinegar Tums and Rolaids– Vitamin C Soap

Dissociation Equilibrium,

HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

H2O(l) + H2O(l) H3+O(aq) + OH-(aq)

This dissociation is called autoionization of water.

HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

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

Brønsted-Lowry definitions Conjugate acid-base pairsConjugate acid-base pairs.. Acids and bases that are related by

loss or gain of H+ as H3O+ and H2O. Examples.Examples. Acid Base

H3O+ H2O

HC2H3O2 C2H3O2-

NH4+ NH3

H2SO4 HSO4-

HSO4- SO42-

Bronsted acid/conjugate base and base/conjugate acid pairs inacid/base equilibria

HCl(aq): acid H2O(l): base

H3+O(aq): conjugate acid

Cl-(aq): conjugate base H2O/ H3

+O: base/conjugate acid pair HCl/Cl-: acid/conjugate base pair

Select acid, base, acid/conjugate base pair,base/conjugate acid pair

H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4

-(aq) acid base conjugate acid conjugate base base/conjugate acid pair acid/conjugate base pair

Equilibrium, Constant, Ka & Kb

Ka: Acid dissociation constant for a equilibrium reaction.

Kb: Base dissociation constant for a equilibrium reaction.

Acid: HA + H2O H3+O + A-

Base: BOH + H2O B+ + OH-

[H3+O][ A-] [B+ ][OH-]

Ka = --------------- ; Kb = ----------------- [HA] [BOH]

What is Ka

HCl(aq) + H2O(l) <===> H3+O(aq) + Cl-(aq)

E.g. Ka

+O][Cl-] Ka= ----------------- [HCl]

[H+][Cl-] Ka= ----------------- [HCl]

What is Ka1 and Ka2?

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

What is Kb

E.g. H2SO4(aq) + H2O(l) H3

+O(aq) + HSO4-(aq)

+O(aq) + SO42-(aq)

[H3+O][HSO4

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

H2SO4 ; Ka2 = -------------------

[HSO4-]

HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

[H+][C2H3O2-]

H C2H3O2; Ka= ------------------

[H C2H3O2]

[NH4+][OH-]

NH3; Kb= --------------

[ NH3]

Which is weaker?

a. HNO2 ; Ka= 4.0 x 10-4.

b. HOCl2 ; Ka= 1.2 x 10-2.

c. HOCl ; Ka= 3.5 x 10-8.

d. HCN ; Ka= 4.9 x 10-10.

WEAKER/STRONGER Acids and Bases & Ka and Kb values

A larger value of Ka or Kb indicates an equilibrium favoring product side.

Acidity and basicity increase with increasing Ka or Kb.

pKa = - log Ka and pKb = - log Kb

Acidity and basicity decrease with increasing pKa or pKb.

Autoionization of water AutoionizationAutoionization When water molecules react with one

another to form ions.

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

– (10-7M) (10-7M)

Kw = [ H3O+ ] [ OH- ]

= 1.0 x 10-14 at 25oC

Note:Note: [H2O] is constant and is

included in Kw.

ion productof water

What is Kw?

This dissociation is called autoionization of water.

Autoionization of water: Kw = [H3

+O][OH-]

Kw is called ionic product of water

Kw = 1 x 10-14

Why is water important for acid/base equilibria? Water is the medium/solvent for acids

and bases. Acids and bases alter the dissociation

equilibrium of water based on Le Chaterlier’s principle

Comparing Kw and Ka & Kb

Any compound with a Ka value greater than Kw of water will be a an acid in water.

Any compound with a Kb value greater than Kw of water will be a base in water.

pH and other “p” scales

We need to measure and use acids and bases over a very large concentration range.

pH and pOH are systems to keep track of these very large ranges.– pH = -log[H3O+]

– pOH = -log[OH-]– pH + pOH = 14

pH scale

A logarithmic scale used to keep track of the large changes in [H+].

14 7 0

10-14 M 10-7 M 1 M Very Neutral VeryBasic Acidic

When you add an acid, the pH gets smaller.

When you add a base, the pH gets larger.

pH of somecommon materials

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

Substance pH

1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0

1M NaOH 14.0

What is pH? Kw = [H3

+O][OH-] = 1 x 10-14

[H3+O][OH-] = 10-7 x 10-7

Extreme cases: Basic medium [H3

+O][OH-] = 10-14 x 100 Acidic medium [H3

+O][OH-] = 100 x 10-14 pH value is -log[H+] spans only 0-14 in water.

pH, pKw and pOH The relation of pH, Kw and pOH

Kw = [H+][OH-]

log Kw = log [H+] + log [OH-]

-log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14

since Kw =1 x 10-14 14 = pH + pOH pH = 14 - pOH pOH = 14 - pH

Acid and Base Strength Strong acidsStrong acids Ionize completely in water.

HCl, HBr, HI, HClO3, HNO3, HClO4, H2SO4.

Weak acids Weak acids Partially ionize in water. Most acids are weak.

Strong basesStrong bases Ionize completely in water. Strong bases are metal

hydroxides - NaOH, KOH

Weak basesWeak bases Partially ionize in water.

pH and pOH calculations of acid and base solutions a) Strong acids/bases

dissociation is complete for strong acid such as HNO3 or base NaOH

– [H+] is calculated from molarity (M) of the solution

b) weak acids/bases

needs Ka , Kb or percent(%)dissociation

pH of Strong Acid/bases

HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)

Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.

[HNO3] = [H+] = 0.2 mole/L pH = -log [H+] = -log(0.2) pH = 0.699

pH of 0.5 M H2SO4 Solution

+O(aq) + SO42-(aq)

[H3+O][HSO4

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

H2SO4 ; Ka2 = ------------------- ; Ka2 ignored

[HSO4-]

pH of 0.5 M H2SO4 Solution

the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning.

[H2SO4] = [H+] = 0.5 mole/L

pH = -log [H+]

pH = -log(0.5)

pH = 0.30

1.5 x 10-2 M NaOH.1.5 x 10-2 M NaOH.

NaOH is also a strong base dissociates completely in water.

[NaOH] = [HO- ] = 1.5 x 10-2 mole/L

pOH = -log[HO-]= -log(1.5 x 10-2)

pOH = 1.82

As defined and derived previously: pKw= pH + pOH; pKw= 14

pH = pKw + pOH

pH = 14 - pOH

pH = 14 - 1.82 ; pH = 12.18

N Acid/Base Definitions n Types of Acids/bases n Polyprotic Acids n The Ion Product for Water n The...

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