Chemistry 281(01) Winter 2014 CTH 277 10:00-11:15 am...

CTH 277 10:00-11:15 am Instructor: Dr. Upali Siriwardane E-mail: [email protected] Office: 311 Carson Taylor Hall ; Phone: 318-257-4941; Office Hours: MTW 8:00 am - 10:00 am; TR 8:30 - 9:30 am & 1:00-2:00 pm. January 16, 2014 Test 1 (Chapters 1&,2), February 6, 2014 Test 2 (Chapters 3 &4) February 25, 2014, Test 3 (Chapters 4 & 5), Comprehensive Final Make Up Exam: February 27, 2012 9:30-10:45 AM, CTH 311.

Chemistry 281(01) Winter 2014

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What are Acids &Bases?

Definition?

a) Arrhenius

b) Bronsted-Lowry

c) Lewis

CHEM 281 Winter 2009 Chapter 4-2

Acid Anything that produces hydrogen ions in a water solution.

HCl (aq) H+ + Cl-

Base 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)

Arrhenius definitions


Expands the Arrhenius definitions

Acid Proton donor

Base 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-

Brønsted-Lowry definitions


Proton in water


Dissociation Equilibrium

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

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

-(aq)

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

This dissociation is called autoionization of water.

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

-(aq)

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


Bronsted conjugate acid/base pairs in equilibria

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

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


Conjugate acid-base pairs.

Acids and bases that are related by loss or gain of H+ as H3O+ and H2O.

Examples. Acid Base

H3O+ H2O

HC2H3O2 C2H3O2-

NH4+ NH3

H2SO4 HSO4-

HSO4- SO42-

Brønsted-Lowry definitions


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


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


Aqua Acids

Acidic proton is on a water molecule coordinated

to a central metal ion

[Fe(OH2)6]3+,Al(OH2)6

3+, Si(OH)4

Acidity increase with charge

Acidity increase as metal become smaller


Anhydrous oxides The Lux/Flood Definition

Covers things which would become acids or bases if dissolved in water.

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

Acids Formula Molarity* nitric HNO3 16 hydrochloric HCl 12 sulfuric H2SO4 18 acetic HC2H3O2 18 Bases ammonia NH3(aq) 15 sodium hydroxide NaOH solid

*undiluted.


Acids and bases

Acidic Basic

– Citrus fruits Baking soda

– Aspirin Detergents

– Coca Cola Ammonia cleaners

– Vinegar Tums and Rolaids

– Vitamin C Soap


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 HCl(aq) + H2O(l) H3

+O(aq) + Cl-(aq)

[H3+O][Cl-]

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


What is Ka1 and Ka2?

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

-(aq)

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)


What is Kb

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


H2SO4 Dissociation

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

+O(aq) + HSO4-(aq)

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

[H3

+O][HSO4-]

H2SO4 ; Ka1 = ------------------- [H2SO4] [H3

+O][SO42-]

H2SO4 ; Ka2 = ------------------- [HSO4

-]


Ka and Kb

E.g. HC2H3O2(aq) + H2O(l) H3

+O(aq) + C2H3O2-(aq)

[H+][C2H3O2-]

H C2H3O2; Ka= ------------------ [H C2H3O2] NH3 (aq) + H2O(l) NH4

+ + OH-(aq) [NH4

+][OH-] NH3; Kb= -------------- [ NH3]


Acidity/Basicity of HA and F-


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 When water molecules react with one another to form ions.

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

(aq)

(10-7M) (10-7M)

Kw = [ H3O+ ] [ OH- ]

= 1.0 x 10-14 at 25oC

Note: [H2O] is constant and is included in Kw.

ion product

of water

Autoionization of 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.


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 and other “p” scales


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 Very

Basic Acidic

When you add an acid, the pH gets smaller.

When you add a base, the pH gets larger.

pH scale


Substance pH

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

1M NaOH 14.0

pH of some common materials


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 acids Ionize completely in water.

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

• Weak acids Partially ionize in water.

Most acids are weak.

• Strong bases Ionize completely in water.

Strong bases are metal hydroxides - NaOH, KOH

• Weak 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


Titration curves

pH

Equivalence

Point

% titration or ml titrant

Buffer region

Overtitration

Indicator

Transition


Indicators Acid-base indicators are highly colored weak

acids or bases.

HIn In- + H+

color 1 color 2

They may have more than one color transition.

Example. Thymol blue

Red - Yellow - Blue

One of the forms may be colorless - phenolphthalein (colorless to pink)


Selection of an indicator for a titration

a) strong acid/strong base

b) weak acid/strong base

c) strong acid/weak base

d) weak acid/weak base

Calculate the pH of the solution at he equivalence point or end point


Common Ion Effect

Weak acid and salt solutions

E.g. HC2H3O2 and NaC2H3O2

Weak base and salt solutions

E.g. NH3 and NH4Cl.

H2O + C2H3O2- <==> OH- + HC2H3O2

(common ion)

H2O + NH4+ <==> H3

+O + NH3

(common ion)


Solutions that resist pH change when small amounts of acid or base are added.

Two types

weak acid and its salt

weak base and its salt

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

Add OH- Add H3O+

shift to right shift to left

Based on the common ion effect.

Buffers


The pH of a buffer does not depend on the absolute amount of the conjugate acid-base pair. It is based on the ratio of the two.

Henderson-Hasselbalch equation.

Easily derived from the Ka or Kb expression.

Starting with an acid

pH = pKa + log

Starting with a base

pH = 14 - ( pKb + log ) [HA]

[A-]

[A-]

[HA]

Buffers


• Control of blood pH

Oxygen is transported primarily by hemoglobin in the red blood cells.

CO2 is transported both in plasma and the red blood cells.

CO2 (aq) + H2O H2CO3 (aq)

H+(aq) + HCO3-(aq)

The bicarbonate

buffer is essential

for controlling

blood pH

Buffers and blood


Main Group Acid/Bases


Amphoteric Oxides


Strength of oxo-acids by Paulings Rules

For OpE(OH)q, pKa ~ 8 - 5p

The successive pKa values of polyprotic acids (i.e. q >1) increase by 5 units for each successive proton transfer.


pKa Values of Oxy Acids


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)


Lewis Acids/Bases


Hard and soft acids and bases


Solvent leveling

• If the solvent contains ionizable protons it is said to be protonic, and if it is protonic, it will engage in acid-base reactions.

• All acids/bases which are stronger than the H3O+(aq) or OH-(aq) ion will react to produce hydronium/hydroxide ion, and so their strength will be leveled to that of the H3O+(aq) or OH-(aq) ion.

• In aqueous solution, the strongest acid/base which can exist is the H3O+(aq) or OH-(aq)


Acid-Base Discrimination Windows


Levelling effect in other protic liquid

all acids are levelled to the strength of the

ammonium ion, NH4+, and all bases are levelled

to the strength of the amide ion, NH2-.

2 NH3 NH4+ + NH2

-.

2HNO3 H2NO3+ + NO2

-

3HF H2F+ + HF2-

2H2SO4 H3SO4+ + HSO4

-


Polycation Formation

NaAl13O4(OH)24(H2O)12(SO4)4 .x H2O


Polyanion Formation:Zeolites

The general method of zeolite production

involves dissolving an aluminium source

(metal or oxide) into an aqueous solution

of sodium or potassium hydroxide. Once

this solution has cooled the silica source

(and organic template, if required) is added

in The form of an aqueous slurry and the

resulting gel stirred until homogenous.

Na12[(AlO2)12(SiO2)12.27H2O


Chemistry 281(01) Winter 2014 CTH 277 10:00-11:15 am...

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Transcript of Chemistry 281(01) Winter 2014 CTH 277 10:00-11:15 am...