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Transcript of PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 15 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of...
![Page 1: PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 15 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state.](https://reader034.fdocuments.in/reader034/viewer/2022051018/56649e9f5503460f94ba194a/html5/thumbnails/1.jpg)
PRINCIPLES OF CHEMISTRY II
CHEM 1212
CHAPTER 15
DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences
Clayton state university
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CHAPTER 15
SOLUTIONS OF ACIDS AND BASES
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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
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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
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BRONSTED-LOWRY ACIDS
- Acids are proton (H+) donors
- Not restricted to aqueous solutions
HCl, HNO3, H2SO4
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- 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
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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+
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- 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
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ACIDS
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
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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+)
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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
CONJUGATE ACID BASE PAIRS
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NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
HF(aq) + H2O(l) H3O+(aq) + F-(aq)
HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)
CONJUGATE ACID BASE PAIRS
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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)
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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+
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AUTOPROTOLYSIS OF WATER
H2O + H2O H3O+ + OH-
Kw
- 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
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- 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
AUTOPROTOLYSIS OF WATER
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- 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)
AUTOPROTOLYSIS OF WATER
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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
AUTOPROTOLYSIS OF WATER
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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
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pH
- 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
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- 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
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- 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
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- 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
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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
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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
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- 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
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
- Predominant species are H3O+ and A-
- Predominant species is HA
STRENGTH OF ACIDS
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- Equilibrium constant for the reaction of a weak acid with water- Represented by Ka (acid dissociation constant)
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
- 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)
[HA]
]][AO[HK 3
a
STRENGTH OF ACIDS
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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
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- 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
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- 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
[B]
]][OH[BHK b
STRENGTH OF BASES
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[B]
]][OH[BHKb
[HA]
]][AO[HK 3
a
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
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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
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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
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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
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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
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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
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- 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
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WEAK ACID EQUILIBRIUM
For a weak acid HA
HA A- + H+
cHA = total concentration = analytical concentration
= [HA] + [A-]
Ka
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WEAK ACID EQUILIBRIUM
HAC
][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
Ka
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For a weak acid HA
HA A- + H+
a
2
KxF
x
[HA]
]][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
Ka
WEAK ACID EQUILIBRIUM
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a
2
KxF
x
[HA]
]][A[H
- The equation reduces to
WEAK ACID EQUILIBRIUM
a
2
KF
x
[HA]
]][A[H
- If x ≤ 5% of F
That is F – x ≈ F if x ≤ 0.05F
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For a weak base B
B + H2O BH+ + OH-
Kb
[B]][B
][BHnAssociatioofFraction
WEAK BASE EQUILIBRIUM
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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
Kb
b
2
KxF
x
[B]
]][OH[BH
WEAK BASE EQUILIBRIUM
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- The equation reduces to
WEAK BASE EQUILIBRIUM
- If x ≤ 5% of F
That is F – x ≈ F if x ≤ 0.05F
b
2
KxF
x
[B]
]][OH[BH
b
2
KF
x
[B]
]][OH[BH
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SALTS
- 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)
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SALTS
- 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
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SALTS
- 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
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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
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- 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
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Acidic Hydrolysis
positive ionof salt + H2O
Conjugatebase + H3O+
- The hydronium ion makes the solution acidic
NH4+ + H2O → NH3 + H3O+
HYDROLYSIS OF SALTS
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Basic Hydrolysis
negative ionof salt + H2O
Conjugateacid
+ OH-
- The hydroxide ion makes the solution basic
F- + H2O → HF + OH-
HYDROLYSIS OF SALTS
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- 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
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- 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
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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
FACTORS AFFECTING STRENGTH OF ACIDS
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Example
Bond strength of hydrogen halides
HF > HCl > HBr > HI
Acidity of hydrogen halides
HF < HCl < HBr < HI
FACTORS AFFECTING STRENGTH OF ACIDS
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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
FACTORS AFFECTING STRENGTH OF ACIDS
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- Bond strength and electronegativity sometimes predict opposite trends
- Bond strength dominates down a group
- Electronegativity dominates across a period
FACTORS AFFECTING STRENGTH OF ACIDS
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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
FACTORS AFFECTING STRENGTH OF ACIDS