PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 15 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of...

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PRINCIPLES OF CHEMISTRY II

CHEM 1212

CHAPTER 15

DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences

Clayton state university

CHAPTER 15

SOLUTIONS OF ACIDS AND BASES

ARRHENIUS ACIDS

- Acids are substances that ionize in aqueous solutions to produce hydrogen ions (proton, H+)

HCl, HNO3, H2SO4

- Arrhenius acids are covalent compounds in the pure state

Propertiessour taste, change blue litmus paper to red, corrosive

ARRHENIUS BASES

- Bases are substances that ionize in aqueous solutions to produce hydroxide ions (OH-)

NaOH, KOH, Ca(OH)2

- Arrhenius bases are ionic compounds in the pure state

Propertiesbitter taste, change red litmus paper to blue, slippery to touch

BRONSTED-LOWRY ACIDS

- Acids are proton (H+) donors

- Not restricted to aqueous solutions

HCl, HNO3, H2SO4

- Bases are proton acceptors

- Not restricted to aqueous solutions

NH3, dimethyl sulfoxide (DMSO)

- Proton donation cannot occur unless an acceptor is present

BRONSTED-LOWRY BASES

LEWIS ACIDS

- Acids are electron pair acceptors

- Not restricted to protons or aqueous solutions

BF3, B2H6, Al2Cl6, AlF3, PCl5,

Metal ions Can accept four or six pairs of electrons from Lewis bases

Fe3+ + 6H2O(l) → Fe(H2O)63+

- Bases are electron pair donors

- Not restricted to protons or aqueous solutions

NH3, ethers, ketones, carbon monoxide, sulfoxides

- The product of a Lewis acid-base reaction is known as an adduct

- The base donates an electron pair to form coordinate covalent bond

LEWIS BASES

Monoprotic Acid- Donates one proton per molecule (HNO3, HCl)

Diprotic Acid- Donates two protons per molecule (H2SO4, H2CO3)

Triprotic Acid- Donates three proton per molecule (H3PO4, H3AsO4)

Polyprotic Acid- Donates two or more protons per molecule

CONJUGATE ACID BASE PAIRS

- Most Bronsted-Lowry acid-base reactions do not undergo 100% conversion

- Acid-base equilibrium is established

- Every acid has a conjugate base associated with it (by removing H+)

- Every base has a conjugate acid associated with it (by adding H+)

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

- HX donates a proton to H2O to form X-

HX is the acid and X- is its conjugate base

- H2O accepts a proton from HX H2O acts as a base and H3O+ is its conjugate acid

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

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

HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)

AMPHOTERIC SUBSTANCES

- A substance that can lose or accept a proton

- A substance that can function as either Bronsted-Lowry acid or Bronsted-Lowry base

- H2O is the most common

(refer to previous slide for examples)

REACTIONS OF ACIDS AND BASES

Arrhenius acid + Arrhenius base → salt + water

HCl + NaOH → NaCl + H2O

B-L acid + B-L base → conjugate base + conjugate acid

H3PO4 + H2O → H2PO4- + H3O+

AUTOPROTOLYSIS OF WATER

H2O + H2O H3O+ + OH-

- Autoionization (self-ionization) of water

- Pure water molecules (small percentage) interact with one another to form equal amounts of H3O+ and OH- ions

reduces to

H+ + OH-H2OKw

- The number of H3O+ and OH- ions present in a sample of pure water at any given time is small

- At equilibrium (25 oC)

[H3O+] = [OH-] = 1.00 x 10-7 M

- [H3O+] = hydronium ion concentration

- [OH-] = hydroxide ion concentration

- The ion product constant of water (Kw) = [H3O+] x [OH-]

= (1.00 x 10-7) x (1.00 x 10-7)

= 1.00 x 10-14

- Valid in all solutions (pure water and water with solutes)

Addition of Acidic Solute

- increases [H3O+] - [OH-] decreases by the same factor to make product 1.00 x 10-14

Addition of Basic Solute

- increases [OH-] - [H3O+] decreases by the same factor to make product 1.00 x 10-14

Acidic Solution- An aqueous solution in which [H3O+] is higher than [OH-]

Basic Solution- An aqueous solution in which [OH-] is higher than [H3O+]

Neutral Solution- An aqueous solution in which [H3O+] is equal to [OH-]

THE pH CONCEPT

- Negative logarithm of the hydronium ion concentration [H3O+] in an aqueous solution

pH = - log[H3O+]

[H3O+] = 10-pH

- Commonly expressed to 2 decimal places (2 significant figures)

THE pH CONCEPT

- For [H3O+] coefficient of 1.0 - Expressed in exponential notation

- The pH is the negative of the exponent value

[H3O+] = 1.0 x 10-5 M, then pH = 5.00

[H3O+] = 1.0 x 10-3 M, then pH = 3.00

[H3O+] = 1.0 x 10-11 M, then pH = 11.00

THE pH CONCEPT

- For neutral solutions pH is equal to 7.00

- For acidic solutions pH is less than 7.00

- For basic solutions pH is greater than 7.00

- Increasing [H3O+] lowers the pH

THE pH CONCEPT

- A change of 1 unit in pH corresponds to a tenfold change in [H3O+]

pH = 3.00 implies [H3O+] = 1.0 x 10-3 M = 0.0010 M pH = 2.00 implies [H3O+] = 1.0 x 10-2 M = 0.010 M

which is tenfold

- The pH meter and the litmus paper are used to determine pH values of solutions

THE pH CONCEPT

pKw = -log(Kw) = -log(1.00 x 10-14) = 14

pOH = -log[OH-]

[H3O+][OH-] = Kw

Implies that

pH + pOH = pKw

pH + pOH = 14.00

THE pH CONCEPT

STRENGTH OF ACIDS

Strong Acids - Transfer 100% (or very nearly 100%) of their protons

to H2O in aqueous solution- Completely or nearly completely ionize in aqueous solution

- Strong electrolytes HCl, HBr, HClO4, HNO3, H2SO4

Weak Acids - Transfer only a small percentage (< 5%) of their protons

to H2O in aqueous solution Amino acids, Organic acids: acetic acid, citric acid

- Equilibrium position lies to the far right for strong acids

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

- Equilibrium position lies to the far left for weak acids

- Predominant species are H3O+ and A-

- Predominant species is HA

STRENGTH OF ACIDS

- Equilibrium constant for the reaction of a weak acid with water- Represented by Ka (acid dissociation constant)

- H2O is a pure liquid so not included- Acid strength increases with increasing Ka value

- For polyprotic acids, Ka for each dissociation step is smaller than the previous step (weaker acid)

]][AO[HK 3

STRENGTH OF ACIDS

Strong Bases- Completely or nearly completely ionize in aqueous solution

- Strong electrolytes

Hydroxides of Groups IA and IIA are strong bases LiOH, CsOH, Ba(OH)2, Ca(OH)2

Most common in lab: NaOH and KOH

Weak bases- produce small amounts of OH- ions in aqueous solution

Organic bases, methylamine, cocaine, morphineMost common: NH3

STRENGTH OF BASES

- Weak bases produce small amounts of OH- ions in aqueous solution (NH3)

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

- Equilibrium position lies to the far left

- Small amounts of NH4+ and OH- ions are produced

- The name aqueous ammonia is preferred over ammonium hydroxide

STRENGTH OF BASES

- Equilibrium constant for the reaction of a weak base with water- Represented by Kb (base hydrolysis constant)

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

- H2O is a pure liquid so not included

]][OH[BHK b

STRENGTH OF BASES

]][OH[BHKb

]][AO[HK 3

Ka x Kb = [H3O+][OH-] = Kw = 1.00 x 10-14

- Reaction goes to completion when Ka value is very large

- Weak acids have small Ka values

WEAK ACIDS AND BASES

pKa = - logKa

pKb = - logKb

pKa + pKb = pKw

- The stronger an acid the smaller its pKa

- The stronger the acid the weaker its conjugate base

- The stronger the base the weaker its conjugate acid

pH OF STRONG ACIDS

- Differences in acidities of strong acids cannot be measured since they all ionize completely

- This phenomenon is known as leveling effect

Find the pH of 3.9 x 10-2 M HCl

HCl is a strong acid and ionizes completely

HCl(aq) → H+(aq) + Cl-(aq)

pH = - log(3.9 x 10-2) = 1.41

pH OF STRONG BASES

Find the pH of 3.9 x 10-2 M NaOH

NaOH(aq) → Na+(aq) + OH-(aq)

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

[H3O+][3.9 x 10-2] = 1.0 x 10-14

[H3O+] = 2.6 x 10-13

pH = - log(2.6 x 10-13) = 12.59

Find the pH of 3.9 x 10-2 M NaOH

Alternatively

pOH = - log[OH-]

pOH = - log(3.9 x 10-2) = 1.41

pH + pOH = 14

pH = 14 - 1.41 = 12.59

pH OF STRONG BASES

pH OF STRONG ACIDS AND BASES

- For dilute solutions the contribution of H2O should not be neglected

- Acids and bases suppress water ionization

What concentrations of H+ and OH- are producedby H2O dissociation in 1.0 x 10-3 M HCl?

pH = 3[OH-] = Kw/[H3O+] = 1.0 x 10-11

OH- is produced from the dissociation of H2OImplies H2O dissociation = [OH-] = [H3O+] = 1.0 x 10-11

- For dilute solutions the contribution of H2O should not be neglected

- Acids and bases suppress water ionization

What concentrations of H+ and OH- are producedby H2O dissociation in 1.0 x 10-4 M KOH?

[H3O+] = Kw/[OH-] = 1.0 x 10-10

H3O+ (or H+) is produced from the dissociation of H2OImplies H2O dissociation = [OH-] = [H3O+] = 1.0 x 10-10

pH OF STRONG ACIDS AND BASES

WEAK ACID EQUILIBRIUM

For a weak acid HA

HA A- + H+

cHA = total concentration = analytical concentration

= [HA] + [A-]

[HA]][A

][AonDissociatiofFraction

For a weak acid HA

HA A- + H+

- Fraction of dissociation increases with increasing acid strength

- Fraction of dissociation increases with dilution

For a weak acid HA

HA A- + H+

]][A[H

- Assume [H+] ≈ [A-]- F is the initial (formal) concentration of HA- Initial concentration of H+ and A- is 0 each- Final concentration of H+ and A- is x each

- The iCe table may be used for such problems

]][A[H

- The equation reduces to

]][A[H

- If x ≤ 5% of F

That is F – x ≈ F if x ≤ 0.05F

For a weak base B

B + H2O BH+ + OH-

[B]][B

][BHnAssociatioofFraction

WEAK BASE EQUILIBRIUM

For a weak base B

B + H2O BH+ + OH-

- Assume [BH+] ≈ [OH-]- F is the initial (formal) concentration of B

- Initial concentration of BH+ and OH- is 0 each- Final concentration of BH+ and OH- is x each- The iCe table may be used for such problems

]][OH[BH

- The equation reduces to

- If x ≤ 5% of F

That is F – x ≈ F if x ≤ 0.05F

]][OH[BH

- Salts are ionic compounds

- The positive ion is a metal or polyatomic ion

- The negative ion is a nonmetal or polyatomic ion [exception is the hydroxide ion (OH-)]

- Salts dissociate completely into ions in solution

- A reaction between an acid and a hydroxide base produces salt(cation from the base and anion from the acid)

- Solutions of salts may be acidic, basic, or neutral

- Acidity depends on relative values of Ka of the cation and Kb of the anion

- The conjugate base of a strong acid (anion from a strong acid) has no net effect on the pH of a solution (spectator ion)

Cl- from HCl, NO3- from HNO3

- Cation from a strong base has no net effect on the pH of a solution (spectator ion)Na+ from NaOH, K+ from KOH

- NaCl solution contains Na+ and Cl- ions

- Both ions are spectator ions and do not affect the pH of the solution

- pH is determined by autoionization of water

HYDROLYSIS OF SALTS

- Reaction of salt with water to produce hydronium ion or hydroxide ion or both (do not go to 100% completion)

- Not all salts hydrolyze

- The salt of a strong acid and a strong base does not hydrolyze - Neutral solution is the result

- The salt of a strong acid and a weak base hydrolyzes - Acidic solution is the result

- The salt of a weak acid and a strong base hydrolyzes - Basic solution is the result

- The salt of a weak acid and a weak base hydrolyzes - Slightly acidic, neutral, or basic, depending on relative

weaknesses of acid and base

HYDROLYSIS OF SALTS

Acidic Hydrolysis

positive ionof salt + H2O

Conjugatebase + H3O+

- The hydronium ion makes the solution acidic

NH4+ + H2O → NH3 + H3O+

HYDROLYSIS OF SALTS

Basic Hydrolysis

negative ionof salt + H2O

Conjugateacid

- The hydroxide ion makes the solution basic

F- + H2O → HF + OH-

HYDROLYSIS OF SALTS

- When determining the pH of a mixture of acidsonly the pH of the strongest acid is considered

- Contributions by the weaker acids towards pH are neglected

- A weak acid produces fewer protons in the presence of a strong acid

Similarly- A weak base produces fewer hydroxide ions in the

presence of a strong base

MIXTURES OF ACIDS

- Key factors are the strength of the H – A bond and the stability of the A- ion

Binary Acid (HA)- An acidic compound composed of hydrogen and one

other element (mostly a nonmetal)HCl, HI, HBr, H2S, H2O

FACTORS AFFECTING STRENGTH OF ACIDS

Bond Strength of Binary Acids

- Generally decreases down the groups of the periodic table

- Due to increasing size of the other element

- Acidity increases down the groups of the periodic table

- Due to decreasing bond strength

Example

Bond strength of hydrogen halides

HF > HCl > HBr > HI

Acidity of hydrogen halides

HF < HCl < HBr < HI

Stability of the A- Anion

- Depends on the ability of the A atom to accept additional negative charge

- Electronegativity is the factor

- A more electronegative atom results in a stronger acid

- Acidity of nonmetal hydrides increases across periods of the periodic table

CH4 < NH3 < H2O < HF

- Bond strength and electronegativity sometimes predict opposite trends

- Bond strength dominates down a group

- Electronegativity dominates across a period

Oxyacids- Acids containing hydrogen, oxygen, and a third element

The third element may be a- Nonmetal: HNO3, H2SO4, H3PO4

- A transition metal with high oxidation state: H2CrO4

- Carbon in organic acids: CH3COOH

- Acidity increases with electronegativity of the third element

- Hypohalous acids (H – O – X), X = Cl, Br, I

PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 15 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of...

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