analytical chemistry review

5/28/2018 Analytical Chemistry Review

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Determination of Nickel Nickel ion is precipitated with a very selective organicprecipitating reagent called dimethylglyoxime (DMGH) andthen adding a slight excess of aqueous ammonia solution.


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Determination of Carbon and

Hydrogen

Determinations are based on direct volatilization

procedure and require only 5 to 10 mg of sample

The sample is heated in a steam of oxygen in the

presence of catalysts, causing it to decompose

into CO2and H2O

C10H8+ 12 O2 10 CO2+ 4 H2O


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Volumetric method of analysis


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Titration

process wherein a stoichiometrically equivalent

quantity of a standard solution is systematicallyadded to a known quantity of a sample

Requirements of a Titration:

Stoichiometric

Reaction must be rapid

No side reactions

Presence of marked change in solution when thereaction becomes complete

quantitative


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Equivalence point

point in which the amount of standard titrant

added is chemically equivalent to the amount of

the analyte in the sample

Endpoint

estimation of equivalence point manifested by achange in physical properties of solution

Titration error

Difference between equivalence point andendpoint; minimized by:

Using a suitable indicator

Using an indicator blank


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Standard solution

solution of known concentrationstandardization

Process of determining the

concentration of a solution of unknownconcentration


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Common Acid-Base Indicators

pKln Acid/Base

Colors

Methyl Orange 3.46 Red/yellow

Methyl Red 5.00 Red/yellow

Bromcresol

Green

4.66 Yellow/blue

Bromthymol

Blue

7.10 Yellow/blue

Phenolphthalein 9.0 Colorless/red

Methyl Violet 1.6 Yellow/violet

Most Suitable Indicator: the one with pKlnclosest to the

pH at equivalence point


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Strong Acids Strong Bases

HCl NaOHHNO3 KOH

H2SO4 Ba(OH)2

HClO4 CH3NOH

HBr C2H5NOH

HI


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Dissociation Constant, Ka/Kb

When a weak acid or weak base is

dissolved in water, partial dissociation

occurs

HA + H2O H3O++ A-

HA = weak acid; Ka = acid dissociation

Ka= [H3O+] [A-]/ [HA]

For weak base, Kb = base dissociation

constant

B + H2O BH+

+ OH-

-


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Multiplication of one equilibrium-constant

expression by the other gives

Ka x Kb = [H3O+] [A-]/ [HA] x [HA] [OH-]/ [A-]

Ka x Kb= [H3O+] [OH-]

Ka

x Kb= K

w= 1.0 x 10-14at 25C


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pH Calculation

pH = - log [H3O+]

When an acid HA dissolves in water, both acid

and water are suppliers of H3O+:HA + H2O H3O

+ + A- (a)

H2O + H2O H3O+ + OH- (b)

When a base B dissolves in water, both the

base and water are suppliers of OH-

:B + H2O BH

+ + OH- (c)

H2O + H2O H3O+ + OH- (d)


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For Acids

Case 1. The acid is the major supplier of H3O+

This means that the determination of hydronium ion concentration is based on

the equilibrium expression for reaction (a)

Case 2. Water is the major supplier of H3O+

The hydronium ion concentration can be determined by solving the equilibrium

expression for reaction (b)

Case 3. Both the acid and the water are major suppliers of H3O+

Both reactions must be considered

For Bases

Case 1. The base is the major supplier of OH-

This means that the determination of hydroxide ion concentration is based on

the equilibrium expression for reaction (c) Case 2. Water is the major supplier of OH-

The hydroxide ion concentration can be determined by solving the equilibrium

expression for reaction (d)

Case 3. Both the base and the water are major suppliers of OH-

Both reactions (c) and (d) must be considered


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Case Major Supplier of H3O+ Necessary condition

1

23

Strong acid

WaterBoth

CHX>> 10-7

CHX


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Case Major Supplier of OH- Necessary condition

1

2

3

Strong base

Water

Both

CB>> 10-7

CB


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- For weak acids:Case Major Supplier of H3O

+ Necessary condition

1

2

3

Weak acid

Water

Both

Kax CHA>> 10-7

Kax C

HA> 10-7

Kax CHA


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Calculation of pH of conjugate acid-base pairs

o There are two equilibria involved:HA + H2O H3O

+ + A

- (e)

A- + H2O HA + OH

- (f )

oThe ionization of water is neglected, and [H3O

+

] is taken from reaction (e)Ka= [H3O+][A

-]

[HA]

Where:

[HA] = CHA - [H3O+] + [OH

-]

[A-] = CNaA + [H3O

+] - [OH

-]

Ka=[H3O+]( CNaA + [H3O

+] - [OH

-])

CHA - [H3O+] + [OH-]

Solving for [H3O+],

[H3O+]

3 + (CNaA+ Ka) [H3O

+]

2- (KaCHA + Kw) [H3O

+] - KaKw = 0

o If CHA& CNaA 10-3 and Ka& Kb 10-3, then CHAand CNaAare larger compared to the differencebetween [H3O

+] and [OH

-], therefore

[HA] = CHA[A

-] = CNaA

Kabecomes,

Ka=[H3O+]CNaA

CHA

o Note also that conjugate acid-base do not react with each other


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pH buffers

- a mixture of weak acid or base and its conjugate- resists changes in pH upon dilution or addition of acids or base

o effect of dilution:pH = pKa - log [HA]/[A

-]

pH depends on the ratio of the concentrations of acids and conjugate bases rather thantheir absolute value


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Buffer Capacity, - depends on both the concentration of components and the concentration ratio- Defined as the quantity of strong acid or strong base needed to cause 1.0 L of buffer to undergo a pH change

of 1.0 unit- To determine the useful pH range of a buffer use the previous equation

[HA] / [A-] pH from the equation

1 / 10 pHmax= pKa - log 1/10 = pKa+ 1

10 / 1 pH min= pKa - log 10/1 = pKa - 1Therefore the useful pH range is pKa 1.

- In selecting a buffer for a given application there are two considerations to be considered:1. The desired pH and2. The chemical compatibility of the buffer components with the sample

- Preparing a buffer:o By combining the calculated quantities of an acid-base conjugate pairo In cases one of the conjugate pair is unavailable, combining of excess of the available weak acid or

baser with an appropriate amount of strong base or acid can be done.o Preparation of buffers uses the aid of pH meter


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Polyprotic Acids and Polyequivalent

Bases Case 1: A solution containing H2A. The stepwise ionization ofH2A provides two sources of H3O

+:

H2A + H2O H3O++ HA- Ka1=[H3O

+][HA-]

[H2A]HA-+ H2O H3O

++ A2- Ka2=[H3O+][A2-]

[HA-]

if Ka1

>>Ka2

and Ka1

/Ka2

100, the first ionization is the major

source of H3O+;the second ionization can be neglected for the

purpose of calculating the pH


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Case 2: A solution containing H2A + HA-.

Similar to case 1: if Ka1/Ka2 100, the second ionization can be

neglected and it becomes one of calculating the pH of a weak

acid in the presence of its conjugate base--- a buffer problem

Case 3: A solution containing HA-.

Substances such as HA- exhibit both acidic and basic character.

When a salt NaHA is dissolved in water, it dissociates completely

into Na+ and HA-. The HA- can undergo ionization,HA-+ H2O H3O

++ A2- Ka2=[H3O+] [A2-]

[HA-]

- And base ionization

HA-+ H2O H2A + OH- Kb2=[H2A][OH-]= Kw

[HA-] Ka1


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For case three both reactions should be considered,

[H3O+] = [H3O

+]formed[H3O+]lost

But [H3

O+]formed

= [A2-] and [H3

O+]lost

= [OH-]formed

= [H2

A]

Therefore [H3O+] = [A2-][H2A]

Or [A2-] = [H3O+] + [H2A]

Solving for Ka1,

[A2-] = [H3

O+] + [H3

O+][HA-]/Ka1

We obtain

Ka2= [H3O+]([H3O

+]+[H3O+][HA-]/Ka1)

[HA-]

[H3O+]2= K

a1K

a2[HA]

Ka1+ [HA-]

It is frequent that Ka1


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Calculation of pH of solutions of

Salts

Salts of strong acid and strong base(ex NaCl)

pH = 7.0

salts of weak acid and strong base or salts ofweak base and strong acid (ex.NaCH3COO;NH4Cl)

Reverse given dissociation constant

salts of weak acid and weak base (ex.NH4HCOO)

[H3O+] = KwKa/Kb


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Concept of Equivalence

Equivalents (n)

Number of reacting species per mole of the species

For acids and bases = replaceable H+ or OH- For ions = electrons that can be added or removed

Redox reactions = electrons lost or gained

equation of equivalence

Between two aqueous solution

NV = NV where N=M(n) between an aqueous solution and solid

NV = weight solid/EW; EW = formula mass/n


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DOUBLE-INDICATOR TITRATION

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Titration.docx
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DILUTION FACTOR

Simple dilution

DF = final volume/vol of aliquot added

serial dilution

final dilution factor = DF1*DF2*DF3


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Complexometry A titrimetric determination which involves the formation of a soluble

but slightly dissociated complex or complex ion

Metal ion + ligand complex(analyte) (chelate)

metal ion central atom in the complex; lewis acid electron pair acceptor

Ligand Group attached to the central atom; lewis base; can either be an anion or

neutral molecule

Complex Metal ligand

EDTA Well known complex


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Coordinate number

The number of bonds formed by the central atom

coordinate covalent bond

Type of bonding involved in the metal-ligand

complex


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Titration of metals with polydentate

ligands

Polydentate ligands and metals often react in a singlestep thereby avoiding the complications of stepwisereactions

Indicators for EDTA titrations: Eriochrome black T (EBT)

Used exclusively in the pH range 7-11 where the blueform of the indicator predominates in the absence ofmetal ions

Endpoint: red to blue

Calmagite

Stable in aqueous solution;

From red to orange


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Solubility equilibria


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Solubility EquilibriaDeals with substances whose solubility

are low

Ksp(solubility product constant) Equilibrium constant expressing the

solubility of a precipitate in water

Saturated : Ksp = IP

Unsaturated : Ksp > IP Supersaturates : Ksp < IP


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Factors that influence solubilitites

Common ion effect

A precipitate is less soluble in a solution containing an excess of

one of the ions common to the precipitate than it is in pure

form

Diverse ion effect An increase in solubility occurs when salts that contain no ions

in common with the precipitate are present in the solution

Ksp,actual= Ksp,apparent/ In such case, the actual Kspis derived by dividing apparent

solubility product constant by powers of mean ionic activity

coefficient, , which is the measure of the effectiveness with

which a chemical species influences equilibrium


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Debye-Huckel Limiting Law Expression

-log = 0.512 Z+Z-= 0.5 CiZi2

where: - mean ionic activity coefficient

Zcharge of ions

ionic strength

temperature

Generally, solubility increases with increasing temperature

pH

The lower the pH the higher the solubility

complex formation

Presence of complexing agents such as ammonia increases

solubility


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Solubility rules

All nitrates, acetates and perchlorates are

soluble

All halides (except F-) are soluble, except with

Ag+, Hg2+and Pb2+

All sulfates except Ba2+and Pb2+are soluble, butCa2+, Ag+, Hg2+, and Sr2+are only slightly soluble

All sulfides are insoluble except with IA and IIA

elements and (NH4

)2

S

All other common inorganic compounds are

insoluble except Ba(OH)2and Sr(OH)2 which are

soluble. Ca(OH)2is only slightly soluble


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VOLHARD MOHR FAJAN

Ag+determination Direct Indirect Direct

Cl-determination Indirect Direct Direct

Titrant KSCN AgNO3

Indicator Fe3+ CrO42- DCF-Acidity/basicity Acidic Neutral to basic Acidic

Precipitation rxn Ag++ Cl- Ag++ Cl-

Titration rxn Ag++ SCN- Ag++ Cl-

Indicator rxn Fe3++ SCN- Ag++ CrO42-

Color change Brick red Yellow orange Pink


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PermanganimetryUses KMnO4as oxidizing agent (titrant)

Permanganate processes where KMnO4is

used as titrant are self-indicating titrationi.e. endpoint is a permanent faint pink color

Primary standards for KMnO4:

AsO3 Na2C2O4 Pure Fe metal


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Other REDOX Titration

Dichromate Process ( using K2Cr2O7as titrant) Indicator: sodium/barium diphenylbenzidine sulfonate

Color change: colorless to red violet

iodometry (direct titration with iodine)

I2is the oxidizing agent

I2solution is the titrant

Indicator: starch solution

End point: deep blue solution

iodometry (indirect titration with I2)

Used for determining substance with oxidizing properties

Titrant: Na2S2O3(reducung agent)

Indicator: starch solution

End point: disappearance of blue color


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Electrochemistry


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Electrochemistry

Electrochemical cells

Consists mainly of 2 electrodes which areimmersed either into the same solution

or into 2 different solutions in

electrolytic contact with one another


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Galvanic/Voltaic

Cell

Electrolytic Cell

1. Energy Chemical reaction

furnishes electrical

energy

Chemical reaction

is forced to

proceed by

application ofelectrical energy

2. Electrode rxn Spontaneous Non-spontaneous

3. Anode rxn Oxidation Oxidation

4. Cathode rxn Reduction Reduction

5. Anode polarity - +

6. Cathode polarity + -


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Schematic representation for

electrochemical cellsConvention:

Constituents comprising the cell are listed in the

order in which they would be encountered if we

begin at the anode and traveled through thecell solutions to the cathode

Use conventional symbols for ions, elements,

molecules

/ - boundary between 2 phases // - salt bridge

, - indicates that the species are at the same

phase


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analytical chemistry review

Documents

suppliers of h3o

h3o ohkax

log h3o

major supplier of h3o

ha x ha oh

ah2o h2o h3o oh bwhen

acid ha dissolves

suppliers of oh