Mindmap for Physical Chemistry

Negatively charged fast moving particle Atom is no longer unbreakable

Idea of atom introduced No experimental evidence Atom is the smallest unit of matter and is not breakable With experimental evidence from the study of gas Invention of cathode ray tube Behaviour of cathode ray in magnetic and electric field Invention of the word "electron". Raisin pudding model of atom Gold foil scattering experiment The idea of presence of tiny nucleus

Democritus

Dalton Goldstein

Thomson

Rutherford Chadwick

Study of radioactivity

Discovery of neutron The idea of nucleus is breakable

Vaporization chamber Ionization chamber Accelerating plates Construction

Mass of an atom

Radioactivity

Atomic Structure

Development of atomic model

Mass spectrometer

Relative isotopic mass

Relative atomic mass

Relative molecular mass

Interpretation of mass spectrum

Working principle - response to m/e ratio

Vacuum Pump

Recorder

Detector

Deflecting magnetic field

To determine the ionization energy of an atom Determination of isotopic mass Determination of atomic mass Determination of molecular mass Determination of molecular structure The relative mass of a particular isotope in C-12 scale where C-12 isotope is defined to have exactly 12 units of mass.

The weighted average of the relative isotopic masses for a particular element in C-12 scale according to their natural abundance The relative mass of a molecule where that the value is the same as the sum of relative atomic mass of all atoms present in the molecule.

1. Atomic Structure.mmap - 2005/5/23 -

History of discovery

Nature of radioactivity

Discovery of radioactivity

,

and

First order reaction

Independent of temp.

Nuclear bomb

Co-60 / Cs-137 (tumor)

I-131 (thyroid cancer)

P-32 (metabolism of plant)

I-131 (thyroid disease)

Geiger-Muller counter

Constant Half-life

Meaning of Half-life

re-orientation of the spinning proton

Nature state of atoms after the big bang / nucelus formation

increasing of p:n ration

instablility of neutron when isolated

lowering of p:n ratio

binding force between proton and neutron

replusion among proton

100% speed of light

high energy EM wave

up to 99% speed of light

very fast moving electron

5% of speed of light

fast moving helium nuclei

Study the nature of radiation

Extraction of radium and polonium

Becquerel

Curie

Rutherford

particle

particle

radiation

particle

particle

radiation

Natural decay

Origin of radioactivity

Nuclear Reaction

Bio-tracer

Leak Detection

Artificial transmutation

Use of Radioactive isotope

Radiotherapy

Carbon-14 dating

Nuclear power

Molecular mass is numerically the same as molar mass

Measurement of T

Measurement of m

Measurement of V

Measurement of P

Conditions a real gas behaves virtually ideal

Assuming the vapor behaves ideally

Volatile liquid vaporizes easily by heating No interaction among particles Individual particle occupies no volume

interaction among particles is negligible High temperature

Low pressure Barometer in the laboratory

Increase in volume of the gas syringe

Difference in mass of the hypodermic syringe before and after injection

Thermometer (final reading)

Definition of partial pressure - the pressure of a component in the gaseous mixture if that component occupies the same volume by itself alone.

Principle

Setup of apparatus

Operation

T

1/V

Charles' Law V N

Boyle's Law P Avogadro's Law V

Ideal gas Law / Equation PV = nRT

Dalton's Law of partial pressure

Determination of molecular mass of a volatile liquid

Total pressure of a gaseous mixture is the same of the sum of the partial pressure of all of the component present

Take final reading when all readings are steady

Take all initial readings

Allow the system to reach a thermo-equilibrium

Precaution

Precaution

Precaution

particles are far apart

particles are moving fast

Insert the needle fully into the gas syringe

Hold the end of the syringe only

Hold the hypodermic syringe horizontal

Take reading only when the reading is steady

space occupied by individual particle is negligible

Some liquid may have not vaporized The vapor may be still expanding The liquid may leak out. Volatile liquid usually has low surface tension. Body warmth may heat up the liquid and expel it from the syringe Inject the liquid into the deeper part, usually hotter, of the gas syringe to ensure better vaporization.

Take reading only when the reading is steady

mole fraction (percentage by number)

The system may have not reach thermo-equilibrium yet.

2. Mole Concept.mmap - 21/5/2005 -

Behaviour of Ideal gas

How much is one mole ?

Mole Concept

How many ?

How heavy ?

Definition

1st Law

2nd Law

m

1/n

Q

1 F = 96500 C

Faraday constant

m

Avogadro's law

s.t.p.

r.t.p.

Carbon-12 Scale

Avogadro constant (Experimental measurement required)

Molar mass

Molar volume

Molar volume of a gas

Faraday's Laws of electrolysis

Relationship between Current I, Amount of charge Q and time t i.e Q=It

How large ?

Charges

Charges carried by 1 mole of electron

A solution with accurately known concentration Process of determination of concentration of a non-standard solution

criteria

A substance which can be used to prepare a standard solution directly known composition stable composition readily soluble preferably with higher molecular mass

strong acid vs strong base strong acid vs weak base weak acid vs strong base

Choice of indicator

pH meter (titration curve) thermometric titration

Standard solution

Standardization

Primary standard

Equivalence point

Titration

Acid-base titration

Back titration

Redox titration

End point

Without indicator

Example

KMnO4 vs H2C2O4.2H2O

KMnO4 vs FeSO4

Iodometric titration

conductometric titration

Using indicator

The point when the end of titration is detected

The point at which the 2 reagents just react with each other completely.

methyl orange or phenolphthalein methyl orange phenolphthalein

methods for titration of weak acid vs weak alkali iodine / potassium iodide starch solution sodium thiosulphate Determination of concentration of chlorine bleach Constant heating is required

Determination of percentage by mass of calcium carbonate in chalk Analysis of aspirin

Used if the reaction of direct titration is not instantaneous / if no suitable indicator for direct titration

3. Chemical Equations and Stoichiometry.mmap - 8/1/2005 -

Volumetric Analysis

Formulae

Chemical Equations and Stoichiometry

Formulae of Compounds

Determination of Formulae

Chemical Equations

Empirical formula

Molecular formula

Structural formula

Empirical formula

Molecular formula

Word Equation

Chemical Equation

Simplest whole number ratio

Actual no. of atoms

Connectivity of atoms

By Combustion Analysis

From Composition by mass

From empirical formula and molecular mass

Water of crystallization

Stoichiometry

Volumes of gases

Reacting masses

Stoichiometric coefficients

Calculation based on equations

Electromagnetic Spectrum Continuous Spectrum Line Spectrum

Different models of atom

The Electronic Structure of Atoms

Background knowledge

4. The Electronic Structure of Atoms.mmap - 19/1/2005 -

Bohr's model

Setup of the apparatus

gas discharge tube slit

prism or diffraction grating

Line spectrum

Continuum

Ionization of atom

Flame test

Red shift

Planck's equation

ground state

excited state

principal quantum number

Identification of element at remote star

Discovery of helium

Atomic emission spectroscopy

Unevenly spaced energy levels

Multiple transitions take place at the same time

Presence of numerous energy levels (shells)

Process of excitation and relaxation

Introduction of concept of quantum

Convergence limits at the higher frequency end of each series

Unevenly spaced lines in each series

Different series

Numerous lines

photographic film

Characteristic of the hydrogen spectrum

Interpretation of the hydrogen spectrum

Numerous lines

Unevenly spaced lines

Convergence limit and continuum

As a mean of identification

Neon light with different colours

Evidence of subshell

Evidence of shell

Uniqueness of atomic emission spectra

Successive ionization energies

Definition

Atomic Emission spectrum of hydrogen

Ionization energies

1st ionization energies across periods

Electrons are orbiting around the nucleus in orbits with with different energy

The allowed energy of the electron in an atom is quantumized

Able to explain the emission spectrum of hydrogen spectrum but fail to explain the emission spectrum of those multi-electron system.

Quantum numbers

Subsidiary (l)

Principal (n)

energy in multi-electron system

shape of orbital / subshell

l = 0,1,2...(n-1)

shell no.

size

principal energy

n = 1,2,3 ...

By mathematical modeling

Heisenberg's uncertainty principle

Wave-particle duality of electron

Features of Bohr's model

Wave-mechanical model / Orbital model

m = -l,...,0,...+l

spatial orientation of orbital spin of electron

s = +½ or -½

Magnetic (m) Spin (s)

Nodal plane

Nodal surface

Radial probability distribution

Pauli Exclusion principle Orbital / Electron cloud representation

repulsion with the inner eusually stronger repulsion with the e- in the same shell usually weaker

Gets more endothermic when increase in attraction with e> increase in repulsion among electrons, and vice versa Value reflects the net attraction between the outermost electron and the nucleus

Attraction with the nucleus Relative position of the outermost e(Electronic configuration of the atom) Repulsion among electrons

Energy required to remove 1 mole of electrons from 1 mole of gaseous atom

Depending on

General pattern

Definition

Variation

Periodic Table

Atomic radius

Aufbau Principle

Electrons occupy the orbital with the lowest energy first

more evenly distributed electron cloud

full-filled 3d and half-filled 4s orbital

less repulsion

less repulsion

less repulsion

more evenly distributed electron cloud

no electron pair-up in 3d and 4s

half-filled 3d and 4s orbital

3d electrons shield 4s electron from the nucleus and make 4s electrons more energetic

Electrons will occupy degenerate orbitals singly with parallel spin before pairing up occurs.

Each orbital can accommodate 2 electrons with opposite spin

All electrons behave different in an atom i.e. no 2 electrons in an atom have the same set of quantum no.

Order of energy of different orbitals

Pauli exclusion principle

Copper

Completely filled quantum shell e.g. K, L, M

Noble gas cores e.g. [He], [Ne] etc.

Chromium

The difference gets even smaller when 3d orbitals are being filled up

Energy of 4s orbital is only a little bit lower than that of 3d orbital

Hund's Rule (Rule of Maximum Multiplicity)

Irregularity Exception

by notation by electron in boxes use of abbreviation

the p to e- ratio is increasing

the repulsion/shielding among electron is getting less

more difficult to remove electron from a positive particle

usually weaker

repulsion with the e- in the same shell

usually stronger

repulsion with the inner e-

the electronic configuration of an ion is the same as the electron configuration of the preceding atom

the p to e- ratio is increasing

Proof of the presence of shells

the increases in successive ionization energies are too large to be represented by linear scale

shift to the right

shift upwards

appear in tiers

In log scale

Successive ionization energy is always much higher than its predecessor

Plot of Successive ionization energies of an atom

Plot of successive ionization energies of different atoms

including metallic radius

Nuclear charge

Primary shielding effect

Secondary shielding effect

the increase in primary shielding effect outweighs the increase in nuclear charge

the increase in nuclear charge outweighs the increase in secondary shielding effect.

Relative position of the outermost e(Electronic configuration of the atom)

Attraction with the nucleus

half of the distance between 2 chemically bonded nuclei of the same element

Depending on

Repulsion among electrons

contract across a period

expand down a group

half of the distance between 2 nuclei of the same element which are not chemically bonded together in a crystal lattice

General trend

van der Waals' radius

Covalent radius

Variation of Successive ionization energies

Representation of electronic configuration

Building up of electronic configuration on paper

Electronic configurations and the Periodic table

Variation of First Ionization energy across the first 20 elements

s-block, p-block, d-block and f-block elements

Showing 2-3-3 pattern across period 2 and 3

Main trend

Effective nuclear charge (nuclear charge - shielding effect) experienced by the outermost e-

Energy required to move the outermost electron to infinity and escape from the attraction of the nucleus Nuclear charge Primary shielding effect Secondary shielding effect

Getting less endothermic down a group

Getting more endothermic across a period

Position of the outermost e- (Size of atom) the increase in nuclear charge outweighs the increase in secondary shielding effect the increase in primary shielding effect outweighs the increase in nuclear charge Imply presence of subshell

Unexpected drops at O and S (or peaks at N and P)

Unexpected drops at B and Al (or peaks at Be and Mg)

Imply pairing up of electrons in an orbital Electron goes into a more energetic p-subshell (or full filled s-subsell in Be and Mg - more evenly distributed charges) Electrons starts pairing up in O and S - extra repulsion (or half filled p-subshell in N and P more evenly distributed charges)

5. Electronic configurations and the Periodic table.mmap - 6/1/2005 -

Different forms of the same element / compound in the same physical state e.g. graphite and diamond, ozone and diatomic oxygen

298K 1 atm

Standard conditions

pure, if it is a liquid or solid Standard state

Experimental determination

Examples and definitions

Need of standard state and condition

1M if it is an solution

most stable form at 298K, 1 atm

Presence of allotrope and polymorph

Heat of Combustion

Heat of Formation

Heat of Solution

Heat of Neutralization

Enthalpy change as Heat flow (q) of a system at constant pressure returning the system to the original temperature under a specific condition. Formation of 1 mole of water Neutralization of strong/weak acid / alkali A solution with specific concentration / an infinitely dilute solution is prepared. 1 mole of compound is formed from its constituent elements in their standard states Complete combustion of 1 mole of a specific substance in excess oxygen

Heat of Hydration

bomb calorimetry to determine the internal energy change

Formation of 1 mole of hydrated crystal from anhydrous solid Experimental setup conversion to enthalpy change by calculation

simple calorimetry to determine the enthaply change

Philips Harris calorimeter to determine heat of combustion

Common assumptions

Experimental setup

Experimental setup Polystyrene cup + cover Insulating layer thermometer stirrer Specific heat capacity of solution is the same as that of water The rate of heat loss is uniform / No heat loss to the surrounding. The heat capacity of some components in the system can be ignored. The specific heat capacities of some components are only approximation. Heat evolved will be the same as heat flow in the original definition of enthalpy change

6. Energetics.mmap - 17/1/2005 -

Standard enthalpy changes

Significance

Enthalpy and Internal energy

Definition : heat change of a system at constant volume

Definition : heat change of a system at constant pressure

The link between the microscopic interaction among particles and macroscopic observable change.

Enthalpy change

Internal energy change

Enthalpy change = Internal energy change + Work done on the atmosphere

Heat flows out of the system because of the increase in temperature of the system

Potential energy among the particles is converted to heat energy when there is an overall strengthening of attractions among the particles.

Bond formation > Bond breaking

Heat flows into the system as the temperature of the system decreases

Heat energy is converted to potential energy when there is an overall weakening of attractions among the particles.

Bond formation < Bond breaking

Energy cannot be created or destroyed

Some changes has no direct reaction at all

The direct reaction may be too vigorous and too dangerous to be carried out

The extent (% of conversion) of reaction cannot be controlled.

Formation of side products

Balanced equation according to definitions

Calculation involving heat capacity

Calculation of mole

Unit, usually in kJmol-1 must be included

+ and - sign must be included

3. Constituent atoms

2. Combustion products

1. Constituent elements

Patience and careful manipulation

Some common references in constructing energy cycles

Prerequisite

Need of Hess's Law

Born-Haber cycle as energy cycle used to determine the heat of formation

Energy cycle & energy level diagram

Another version of law of conservation of energy

The total enthalpy change of a chemical reaction is independent of the route by which the reaction takes place

Endothermic process

Exothermic process

Relationship between Enthalpy change and Internal energy change

Exothermic and Endothermic process

Energetics

Hess's Law

Calculation

unit and sign of the answer

Attraction among oppositely charged ions Repulsion among similarly charged ions

Some terms

Different arrangements of lattice points

Balance between non-directional electrostatic attraction and repulsion

face centered cubic (f.c.c.) structure simple cubic structure Actual arrangement is depending on

ionic radius

Lattice point

Unit Cell

Coordination number (C.N.) no. of nearest neighbour

e.g. NaCl

Potential well diagram

C.N. = 6 for both ions C.N. = 8 for both ions e.g. CsCl

Stoichiometry of the compound ionic radius ratio Determination of empirical formula 1. the smallest unit containing all the fundamentals of a crystal without repetition. Definition

Determination of no. of different ions in an unit cell

2. the smallest unit of a crystal from which the original crystal structure can be recreated by repetition in 3-dimensional space. Determination of empirical formula The position of a formula unit of the compound considered. Definition - Distance from the centre of an ion to the point of zero electron density between 2 adjacent ions.

Size of isoelectronic ion

Value depending on

determined by X-ray diffraction

p to e- ratio

no. of atoms present

Electron density map polyatomic ion is bigger than simple ion In general, negative ion is bigger than positive ion

As the no. of electrons and electronic configurations are the same, the size is only depending on the number of protons present.

7. Ionic bonding.mmap - 19/1/2005 -

Energetics of formation

Ionic Bonding

Ionic crystals

Born Haber Cycle

Stoichiometry of ionic compounds

Heat of formation

always positive

Removal of 1 mole of electrons from 1 mole of gaseous atoms / ions

always positive

formation of 1 mole of gaseous atoms

formation of 1 mole of compound

Atomization energy of element

Ionization enthalpy of atom

Addition of 1 mole of electrons to 1 mole of gaseous atoms / ions

negative : attraction between the incoming electron and the nucleus > the repulsion between the incoming electron and the existing electrons

depending on the packing efficiency i.e. crystal arrangement

depending on charge density of ions involved

positive : attraction between the incoming electron and the nucleus < the repulsion between the incoming electron and the existing electrons

always negative

Formation of 1 mole of ionic compound from its constituent gaseous ions

Electron Affinity of atom

Lattice energy

Construction of Born Haber Cycle for the hypothetical ionic compounds e.g. MgCl and MgCl3

Cations and anions are just in contact with each other

All ions are perfect spheres

The lattice arrangement of the hypothetical compound is known

Hypothetical lattice energy calculated from the ionic model

Construction of Born Harber cycle for the naturally occurring ionic compound e.g MgCl2

Indeed, the stability of an ionic compound is mainly arising from the formation of ionic bond among oppositely charged ions.

Positive metal ion + free electron are actually more energetic (less stable) than a gaseous metal atom

a measure of strength of ionic bond

Limitation of using stable electronic configuration to explain

Nature prefers the stoichiometry with the most exothermic heat of formation i.e. highest stability comparing with the constituent elements

Assumptions of ionic model

The bond is purely ionic

The value of heat of formation is mainly determined by the values of ionization energies and lattice energy

Strength of Covalent bond

Attraction between nucleus, bonding electron and another nucleus

Replusion among the electrons

directional attraction

e.g. BF3

Expansion of octet

Can be less than octet

Period 2 elements never excess octet

Formation of more bonds

More overall exothermic change

High overall stability

e.g. PCl5, SO42-

Availability of low lying energy orbital (e.g. 3d orbital comparing with 3p orbital)

an electron acceptor with a vacant p-orbital for dative bond formation

Energetically unfavorable

3s orbital has much higher energy than 2p orbital

# of p - # of e- associated with an atom assuming that all bonding e- are belonging to the more electronegative atom

# of p - # of e- associated with an atom assuming that bonding e- are equally shared

e.g. NH4+, H3O+ NH3BF3, Al2Cl6

Overlapping of a lone pair with a vacant orbital

a covalent bond in which the bonding electrons are supplied by one atom only

Promotion of paired electrons to low lying energy orbital (e.g. 2p comparing with 2s, 3d comparing with 3p) to yield more unpaired electrons for bond formation

Increase in electron density along the inter-nuclear axis

Constructive interference of atomic orbitals

Replusion between the nuclei

Potential well diagram

Cases not obeying octet rule

Cases obeying octet rule

Oxidation number

Formal charge

Dative covalent bond

Pairing up of unpaired electrons (bonding electrons)

Overlapping of atomic orbital

Balance between electrostatic attraction and repulsion

Formation of covalent bond

Drawing of Lewis structure

Covalent bonding

Giant Covalent Structure

Shape of molecules and polyatomic ions

Factors affecting bond strength

(Average) Bond enthalpy/energy E(X-Y)

Bond Dissociation Energy D(X-Y)

In the range of 150-1000 kJmol-1 Energy required to break a particular bond The average energy required to break a type of bond Estimation using enthalpy change of atomization

Bond length

Bond order

Using bond energy to estimate the enthalpy change of a reaction no. of bonding electrons / 2

Estimation of bond length Breakup of additivity of covalent radius

Atomic size (covalent radius)

(Primary) shielding effect affects the nuclear attraction experienced by the bonding electrons

The longer the bond length, the weaker will be the bond

Resonance

Polarity of bond covalent bond with certain ionic character is found to be stronger than purely covalent or purely ionic bond bond polarity

Description of shapes

Valence shell electron pair repulsion theory

repulsion among the lone pairs on adjacent atoms

Electrostatic repulsion is very sensitive to distance

Electrons centres tend to separate from each other as far as possible to minimize mutual repulsion

Electron centres : lone pair, single bond, double bond, triple bond

e.g. weak F-F, O-O and N-N bond

repulsion between lone pair-lone pair > lone pair-bond pair > bond pair-bond pair lone pair always occupies equatorial position in a five electron centres system linear, angular (V-shaped or bent), trigonal planar, tetrahedral, trigonal pyramidal, trigonal bipyramidal, seesaw, T-shaped, octehedral, square planar In the describing of the shape of molecule, the presence of lone pair is not included.

Diamond

Graphite

Quartz (SiO2)

tetrahedrally arranged strong C-C single bond only high hardness, high melting point, electrical insulator, excellent heat conductor arrange in hexagonal layers strong C-C bond, bond order 1 1/3, within layer weak van der Waals force between layers

high hardness, high melting point

low hardness, very high melting point, conducting along the layer, slippery between layers

8. Covalent bonding.mmap - 2005/5/21 -

e.g. Cl2

e.g. HCl

The bond is non-polar

Polar covalent bond

Pure covalent bond

Polarity of a molecule

(Permanent) Dipole moment

The bonding electrons are not shared equally

The bonding electrons are shared equally

atoms have different electronegativity

Definition - product of quantity of charge (q) and distance of charge separation (d) The dipole moment along a bond is pointing from the less electronegativity atom to the more electronegativity atom Lone pair

Depending on the vector sum of all dipole moments

Contribute a dipole moment pointing away from the atom Geometry of the molecule magnitude of individual dipole moment A molecule with non-polar bond only

Non-polar molecule

Polar molecule (Dipole)

Do not / weakly attracted by an electric field

A molecule with all dipole moments canceling out each other Non-polar molecules are also attracted by an electric field as a weak temporary dipole moment will be induced onto them by the electric field A molecule with non-zero net dipole moment May not be able to tell the actual direction of the net dipole moment Attracted by an electric field

9. Intermeidate type of bonding.mmap - 26/1/2005 -

Starting from pure covalent bond

Starting from pure ionic bond

Intermediate Type of Bonding

Assumptions of simple ionic model

The lattice arrangement of the compound is known, if it is hypothetical All ions are perfect spheres

Cations and anions are just in contact with each other and clearly separated from each other

Polarization of anion

Mutual repulsion of electron cloud in an ionic lattice

e.g. NaCl (almost purely ionic)

The ions are not perfectly spherical

Certain covalent character

Polarization of anion

p to e- ratio

secondary shielding effect

primary shielding effect

e.g. AgCl, AgBr, AgI, ZnS

Ions may not clearly separated from each other

ease of distortion of electron cloud

inversely proportional to the nuclear attraction on the electron cloud

electronic configuration

shielding effect

nuclear charge

directly proportional to the charge density of a cation

reflects the attraction on the incoming electron to be added to an atom

reflects the attraction on the outermost electron in an atom

Depending on the effective nuclear charge experienced by the bonding electron

Not applicable to noble gases

Increase across a period and up a group

Polarizing power of a cation

Polarizability of an anion

Polarization - distortion of electron cloud of anion under the influence of cation

A very different value implies failure of the ionic model with strong covalent character

A more agreed value implies a situation similar to the assumptions of ionic model pure ionic bond

The bond is purely ionic. i.e. no sharing of electron or incomplete transfer of electron

The lattice energy calculated based on the ionic model is never the same as the experimental value

Anion is always polarized by the cation

ionization energy

Tendency of an atom to attract the bonding electron in a stable molecule

For all real ionic compounds

Electronegativity Pauling's scale

electron affinity

Fluorine has the highest assigned value : 4.0

The percentage of ionic character of a covalent bond is depending on the difference in electronegativity of the 2 elements bonded together

Nature of metallic bond

Sea of electrons model

lattice of positive ions immersed in a sea of delocalized electrons

Non-directional in nature

agitation of delocalized electrons

delocalized electrons capable to reform the metallic bond even if the positive ions are shifted

delocalized electrons get excited and give out luster

Electrostatic attraction among positive ions and sea of delocalized electrons

high ductility and malleability

high conductivity

metallic radius = covalent radius

unique properties of metal

metallic luster

Charge density of the positive ion e.g. Na+ > K+ > Rb+

abab... pattern

C.N. = 12

In molten state, most metallic bond are not broken

high charge density of the positive ion limited the availability of the valence electrons e.g. Hg

no. of valence electrons e.g. Al > Mg > Na

In the range of 100-500 kJmol-1 Attraction on the delocalized electrons

Availability of delocalized electron Big difference between m.p. and b.p.

hexagonal close packing (h.c.p.)

C.N. = 12

abcabc... pattern

C.N. = 8

(face centred) cubic close packing (f.c.c. or c.c.p.) tetrahedral holes and octahedral holes body centered cubic (b.c.c.) Determination of no. of atoms in an unit cell

open-structure (68%)

close-packed structure (74%)

Strength of metallic bond

Metallic crystal (Giant metallic structure)

Metallic bonding

10. Metallic bonding.mmap - 22/2/2005 -

2 per H2O molecule, 1 per HF, NH3, MeOH molecule

H on F, O and N Attraction between a naked proton and a lone pair

An exceptionally strong directional dipole-dipole interaction hydrogen atom bonded onto a very electronegative atom / entity

overall strength depends on

formed between molecules (usually) gives higher b.p.

intermolecular

intramolecular dimerization of carboxylic acids in vapour state / non-aqueous solvent

unique properties of water

formation of secondary structure in protein and DNA

Examples

intermolecular and intramolecular hydrogen bond

Estimation of strength of hydrogen bond

Strength (1/10 of covalent bond) 5-50 kJmol-1

lone pair on a very electronegative atom

including H on CHCl3 lone pair on F, O and N only difference in electronegatvitiy average no. of hydrogen bond per molecule

Hydrogen bonded molecule has higher boiling point that ordinary polar molecule with comparable size.

Experimental

hydrogen bonded molecules also possesses van der Waals' forces

H-bond breaking

H-bond formation

e.g. HF, H2O, NH3, MeOH comparing with other molecules with comparable size

formation of new intermolecular forces > breaking of old intermolecular forces

Period 2 hydrides comparing with hydrides of period 3 to 5

Mixing ethanol and hexane

Mixing trichoromethane and ethyl ethanoate

a substance would be soluble (usually) gives higher solubility in water

formed within a molecule prevents the formation of intermolecular hydrogen bond

overall polarity of the molecule (actually more important than ability to form H-bond)

ability to form hydrogen bond with water e.g. trans-butenedioic acid e.g. alcohols (usually) gives a lower b.p. forming bond is better than not forming bond

Low density of ice

highest density at 4 °C

forming strong bond is better than forming weak bond

open structure of ice

water expands upon freezing diamond like structure

at 0 °C

from below 0 °C to 0 °C melting of ice

DNA

Protein

Variation of density with temperature (K.E.)

double helical structure

single helical structure

from 4 °C to 100 °C

from 0 °C to 4 °C

rapid collapsing of open structure

thermal expansion of solid structure

six membered rings of O linked by O-H▪▪O linkage

density drops

density increases sharply continuous collapsing of open structure

thermal expansion of liquid water

still possesses certain hexagonal network

density increases continuously

density drops

cytosine - guanine pair thymine - adenine pair

hydrogen bond between heterocyclic bases on 2 separate strands

hydrogen bond between N-H and C=O group among different amide groups along the polypeptide chain

3 H-bonds 2 H-bonds

11. Intermolecular Forces.mmap - 2005/5/23 -

Hydrogen bond

Intermolecular Forces

Type of dipole (polar molecule)

van der Waals forces

Molecular crystal

permanent dipole

permanent separation of centre of positive charge and centre of negative charge

no interaction (attraction and repulsion) among particles

individual particle occupies no volume

possessed by all kinds of molecules including polar molecules

the fluctuation of electron cloud creates an instantaneous dipole

possessed by all polar molecules (dipoles) instantaneous/temporary dipole

induced by a permanent dipole

with instantaneous dipole-induced dipole interaction only

b - volume of individual particle

the opposite ends attracted each other

dispersion force

London force

temporary dipole-temporary dipole interaction

induced dipole-induced dipole interaction

branched isomer of hydrocarbon has lower boiling point

lower b.p. and m.p.

higher b.p. and m.p.

the opposite end of different temporary dipoles attract each other

stronger overall strength for molecules with comparable size

weaker overall strength for molecules with comparable size

polarizability of constituent atoms e.g. telfon/poly(tetrafluoroethene)

increases with increasing molecular size / surface area e.g butane vs 2-methylpropane

with dipole-dipole interaction, dipole-induced dipole interaction and instantaneous dipole-induced dipole interaction

an instantaneous dipole induce a temporary dipole on another molecule

other names

a permanent dipole induce a temporary dipole onto another polar/non-polar molecule

also called permanent dipole induced dipole interaction

attractions among opposite ends of permanent dipole

also called permanent dipole-permanent dipole interaction

a - attraction among particles

assumption of ideal gas law

van der Waals' constants

Real gas does not obey ideal gas law perfectly

induced by instantaneous dipole

dipole-induced dipole interaction

dipole-dipole interaction

Condensation of solid and liquid at low temp.

Real gas equation

induced dipole

Evidence

Origin

Polar molecules

instantaneous dipole-induced dipole interaction

Strength (1/10 of covalent bond) 5-50 kJmol-1

Non-polar molecules

depending on the polarizability of the molecule

half of the internuclear distance between 2 atoms of the same element which are not chemically bonded together in a crystal lattice

face-centred cubic structure (f.c.c)

face-centred cubic structure (f.c.c)

van der Waals' radius

iodine

dry ice

Simple molecular

Macromolecular

non-conducting

Giant metallic structure

Giant ionic structure

Giant covalent structure

Molecular structure

Structure and properties of materials

Giant structure

12. Structures and properties of materials.mmap - 2005/6/5 -

usually solid

In general, like dissolves like

the rigid structure transmits the vibration of the C atoms efficiently

bonding electrons in most covalent bonds are not delocalized

intermolecular forces never stronger than covalent bond

solubility depending on the polarity of the solute and solvent

soft and weak

usually gases, liquids or low melting solids

fat and oil - marginal cases, with just over 100 atoms

e.g. H2O, NO2, S8, P4 under 100 atoms

weaker intermolecular forces

e.g. starch, protein, plastic over thousands of atoms stronger intermolecular forces no delocalized electrons

insoluble in any solvent

high melting point

high density

hard

e.g. diamond, quartz, graphite

strong covalent bond

usually non-conductor of electricity

Exception - diamond is an excellent conductor of heat

presence of delocalized electrons along the hexagonal layers

weak van der Waals' forces between layers

flat hexagonal layers

weak van der Waals' forces between layers

usually poor conductor of heat brittle slippery conducting

repulsion cleaves the crystal

weak van der Waals' forces between layers

strong C-C bond, bond order 1 1/3

large space between layers

high melting, even higher than diamond

lower density than diamond

slight dislocation brings ions of similar charge together

only conducts in molten or aqueous state

increase in disorderness

dipole - ion interaction

the ions are not free to move in solid state

brittle

high density

hard

high m.p. and b.p.

Speical case : Graphite

strong ionic bond

presence of ions

hard

high m.p. and b.p.

If soluble, only soluble in aqueous solvent

sometime, strong metallic bond

high conductivity of heat and electricity

malleable and ductile

high density Presence of sea of delocalized electrons Soluble in mercury

Concentration of reactants Pressure Temperature Surface Area Catalyst Light

Definition of Rate

Measuring of Reaction Rate

Rate of Chemical Reactions

Factors affecting Reaction Rates

13. Rate of Chemical Reactions.mmap - 2005/6/5 -

Change in quantity per unit time

Quantity is usually expressed in moldm-3 Time is usually expressed in s

Depending on the units of quantity and time Usually expressed in moldm-3s-1

Instantaneous rate

Average Rate

Unit of rate of reaction

Different rate Initial Rate (Instantaneous rate at t = 0)

Other possible units : mol, g For gases : Pa, mmHg, atm, cm3, dm3

v

by sudden cooling

by dilution with a large amount of water

by removing the catalyst

Need of a calibration curve

Transmittance and Absorbance

by removing one of the reactant

constant amount of Br2 produced / phenol consumed

starch indicator

constant amount of I2 produced / S2O32- consumed

phenolphthalein indicator

constant amount of methanal reacted / HSO3- consumed

constant amount of S ppt. produced to mask the cross

Measuring of mass of the reaction mixture

Measuring of colour intensity

Measuring of volume of gas

Quenching of reaction before titration is done

Methanal Clock experiment

Disproportionation of Na2S2O3 in acidic medium

By measuring a physical quantity related to the amount of a chemical

By titration

Differential form of rate equation / law

Constant time approach (fix the time interval and determine the amount present)

Constant amount approach (fix the amount to be reacted and determine the time required to complete the reaction) Iodine Clock experiment I- vs H2O2

Reaction between Br- and BrO3- in acidic medium

methyl red indicator

Differential form

Order with respect to individual reactant

Order of reaction

Rate constant k

Overall order of the reaction Unit of rate constant is depending on the overall order of the reaction A value depending on temperature only the rate is independent of the concentration of the reactant

Ordinal form

Zeroth order Different order with respect to a reactant First order

Second order

Usually occurring in catalyzed reaction where the availability of the catalyst surface is the limiting factor of the rate the rate is directly proportional to the concentration of the reactant the rate is directly proportional to the square of the concentration of the reactant

Zeroth order reaction

Rate Equation / Law

half life

Rate Equations and Order of Reactions

The concentration of C-14 in atmosphere is assumed to be constant Integrated form

First order reaction

The percentage of C-14 in a living organism is assumed to be constant. half life

carbon-14 dating

By comparing the percentage/beta emission of C-14 in an archaeological sample and the percentage/beta emission of C-14 in a new sample, the age of the archaeological sample can be estimated. The rate of decay is independent of temperature as it doesn't involve collision among particles

Second order reaction half life

Initial Rate method

From the initial rate with different initial concentrations

If there is more than one reactant involved, keep the concentrations of all reactants other than the one being studied virtually constant by keeping them in large excess By plotting concentration against time

Determination of rate equation / law

Horizontal line By plotting graph

By trial and error

By plotting rate against concentration

zeroth order

Straight inclined line

zeroth order

Straight inclined line parabola

first order

second order

By plotting rate against [A] or [A]2 etc to see which will give a straight inclined line passing through the origin By plotting [A] or ln[A] or 1/[A] against time Plotting log (rate) against log (concentration) 14. Rate Equations and Order of Reactions.mmap - 2005/1/31 -

The slope will be the order of reaction with respect to the reactant

The distribution is temperature dependent The total area under the curve is independent of temperature. i.e. the total no. of particles present

e.g iron in Harber process By breaking up a single step into a multi-steps reaction with lower activation energy negative catalyst

positive catalyst

Providing an alternative pathway with a different activation energy, usually lower.

e.g. H+ in decomposition of H2O2 The catalyst has the same phase as the reactants and product

e.g. acid-catalysed esterification The catalyst has a different phase as the reactants and product

Effect of catalyst

homogeneous catalysis

Simple Collision Theory (Effective Collision)

(A)Arrhenius Constant

Energy larger than activation energy Arrhenius Equation

(Z) Collision Frequency

Depending on the nature of the reaction

Vary only a little when the temperature change

(P) Orientation / geometrical / probability factor

The step with the highest activation energy The step with the lowest rate

can be isolated at low temperature

e.g. carbocation

e.g. a pentavalent transition state

Repeat a reaction at two different temperature In the reaction k can be replaced by rate if initial rate method is used k can be replaced by 1/t if constant amount approach is adopted

An energy barrier the reactants need to overcome for the reaction to proceed. By varying the temperature and measuring the rate of a reaction A plot of ln k versus 1/T A plot of ln rate versus 1/T if initial rate method is used A plot of ln 1/t versus 1/T if constant amount approach is adopted

rate-determining step (r.d.s.)

Reaction intermediate

can never be isolated

Determination of activation energy

Multi-stage Reaction

Single stage reaction

transition state (or activated complex)

Activation Energy

Energy profile

The Effect of Temperature Change and Catalyst on Reaction Rate

Maxwell-Boltzmann distribution of molecular speeds

Catalysis

Other examples

heterogeneous catalysis

Haber process

e.g. Catalytic decomposition of hydrogen peroxide by maganese(IV) oxide Fe(s)

Contact process hydrogenation of oil

Enzymes

Catalytic Converter

Gasification of naphtha in product of town gas

Ni(s), Pt(s) or Pd(s)

Pt(s) or V2O5(s)

chemsorption / adsorption on the surface of the catalysis

Ni(s) and NiO

Rh(s) and Pt(s)

zymase in the fermentation by yeast 15. The Effect of Temperature Change and Catalyst on Reaction Rate.mmap - 21/5/2005 -

including all species Excluding those species having no effect on the equilibrium position e.g. solvent(species in large excess), insoluble solid/liquid Examples

Equilibrium constants

At equilibrium and only at equilibrium, the order of reaction with respect to a species will be the same as the stochiometric coefficient of the species Generic equilibrium constant, K Equilibrium constant in terms of concentration, Kc Equilibrium constant in terms of partial pressure, Kp

Significance of equilibrium constant

Determination of Equilibrium constant

Unit of equilibrium constant

Esterification

The equilibrium concentrations of all species are determined by the initial concentrations of them and the equilibrium concentration of any one of them.

Construction of a calibration curve using colorimeter Formation of [Fe(NCS)]2+(aq) complex

Reaction between Fe2+(aq) and Ag+(aq)

Assuming that the rate of reaction is extremely slow and quenching is not required before any titration is done. Preparation of solutions with known [Fe(NCS)]2+(aq) concentration using excess Fe3+(aq) or NCS-(aq) Determination of the equilibrium concentration of [Fe(NCS)]2+(aq) using colorimeter and the calibration curve prepared Assuming that the rate of reaction is extremely slow and quenching is not required before any titration is done The titration is self-indicating if standard KCNS(aq) is used in the determination of the concentration of Ag+(aq). The solution will turn red at the end point with the formation of blood red [Fe(NCS)]2+(aq) complex.

The value is temperature dependent.

The value is depending on the relative stability of the products comparing with the reactants

The value shows the relative amounts of products and reactants when the system is at equilibrium i.e. extent of reaction at equilibrium.

16. Dynamic Equilibrium.mmap - 6/4/2005 -

Equilibrium Law

Equilibrium

Dynamic Equilibrium

Equilibrium position

A state where no macroscopic change is observed.

All chemical equilibria are dynamic in nature. i.e. forward rate = backward rate

The amounts of reactant and product are usually not equal in an equilibrium.

Examples

Thermal dissociation of calcium carbonate

Esterification

Action of acid and alkali on bromine water

Action of acid and alkaline on potassium dichromate

Action of acid and alkali on bismuth chloride

Definition : a set of reactant and product concentration

Using of equilibrium constant and reaction quotient to predict the direction to which the equilibrium position will be shifted.

Q = K : the system is at equilibrium

Q > K : the equilibrium position will be shifted to the left

Q < K : the equilibrium position will be shifted to the right

Effect of change in pressure

this has no effect on the equilibrium position

The rate of a reaction with a higher activation energy is more sensitive to the change in temperature

Catalyst does not change the relative stability of the products comparing with the reactants

Catalyst has no effect on the equilibrium position. It only shortens the time required for a system to reach a state of equilibrium.

Use of change in distribution of molecular speed to explain the effect of change in temperature on the equilibrium position.

by adding an inert gas

by adding one of the product/reactant

by decreasing/increasing the volume

Effect of change in concentration

The principle states that when a system in equilibrium is subjected to a change, the system will response in a way so that the effect of the change imposed will be minimized.

Reaction Quotient (Q)

Le Chatelier's principle Effect of change in temperature

Effect of catalyst

Ka1 is always much larger than Ka2

Successive ionization of polyprotic (polybasic) acid

A more negative pKa (pKb) value means a larger Ka (Kb) value, thus a stronger acid (base).

HA- is relatively more stable when comparing with H2A than A2- when comparing with HAA2- has a much higher charge density than HA- and has a very high potential energy i.e. more unstable It is more difficult to remove a positive proton from negative HA-(aq) than from electrically neutral H2A(aq) Beside the small value of Ka2, the second step of the ionization is also suppressed by the presence of H3O+ from the first step of dissociation which will tend to shift the equilibrium position of the second step to the left. This is known as common ion effect (of H3O+ in this case). Eventually, the amount of H3O+ from the second step of ionization will be very minimal. Determination of acidity constant Ka by half-way neutralization The value of Ka is only depending on temperature but the % of dissociation of a weak acid is depending on concentration as well.

17. Acid Base Equilibrium I.mmap - 6/4/2005 -

Acidity and Basicity constants

Different definitions

Arrhenius definition Bronsted Lowry definition Lewis defintion

acid : substance that produces H+ in water

Introduction of the concept of conjugate acid-base pair

base : substance that produces OH- ions in water

acid : proton donor

base : proton acceptor

acid : electron acceptor

base : electron donor

An Arrhenius acid (base) must be a Bronsted Lowry acid (base) and a Bronsted Lowry acid must be a Lewis acid (base) but not vice versa.

A neutral solution has a pH 7 only at 25°C.

Water is a very poor conductor as it always possesses very small amount of H3O+(aq) and OH-(aq) ions i.e. 1 pair of ions for every 556 million water molecules

Need of calibration using a buffer solution before use

using pH paper / universal indicator

Every unit increase / decrease in pH scale means a ten-fold decrease / increase in the concentration of hydroxonium (H3O+) ion

Originated from French "pouvoir hydrogene" (power of hydrogen).

Auto-ionization (dissociation) of water

Acid Base Equilibrium I : Basic Concepts

pH scale

Measurement of pH

Using pH meter

phenolphthalein or methyl orange methyl orange phenolphthalein no suitable indicator

strong acid vs strong alkali strong acid vs weak alkali weak acid vs strong alkali weak acid vs weak alkali

Examples

Add phenolphthalein to the mixture, then do the first part of the titration (V1). Thereafter, add methyl orange and do the second part of the titration (V2). A mixture of Na2CO3 (V1+V2) and NaHCO3 (V2) A mixture of NaOH (V1) and Na2CO3 (V1+V2) A mixture of NaOH (V1) and NaHCO3 (V2)

18. Acid-base Equilibrium II.mmap - 21/5/2005 -

Choosing of indicator and pH titration Curves

Double indicator titration

Definition : A solution that resists change in pH when a small amount of acid or base is added to it.

Composition : A mixture of weak acid and weak base i.e. usually a weak acid (base) and its salt i.e. the conjugate base (acid).

In the calculation of the pH of a buffer solution, it is assumed that the inter-conversion between the acid (base) and its salt upon mixing is negligible.

The buffering capacity of a buffer solution is depending on the concentrations of the acid/base and its salt.

Human eyes are only capable to detect a colour change when the concentration of one coloured substance is more than 10 times of another coloured substance.

Acid-base indicator is usually a weak acid (base) which colour is different from its conjugate base (acid).

Buffer solution

Acid-base Equilibrium II : Buffers and Indicators

Acid-base indicators

reducing agent is the reagent which is oxidized in a redox reaction oxidizing agent is the reagent which is reduced in a redox reaction

Remember the common strong reducing agents

Remember the common strong oxidizing agents

Oxidation half equation

Reduction half equation

The change in oxidation no. per atom is the same as the no. of electron being lost or gained by that particular atom Figure out the no. of atom that will be oxidized and the no. of atom that will be reduced by using the change in O.N. Add H2O to either side of the equation to get the number of O atom balanced if necessary. Add H+ to either side of the equation to get the number of H atom balanced if necessary. Check the presence of H+ or OH- on the both sides to ensure it agrees with the acidity/alkalinity of the medium used. For reactions in acidic medium, there should be no OH- ion on the both sides of the equation For reactions in alkaline medium, there should be no H+ ion on both sides of the equation

19. Redox Reactions.mmap - 2005/5/23 -

By Combining two half equations

By change in oxidation no.

Check the medium of reaction

Balancing Redox Equations

Definition of redox reaction

Redox Reactions

Oxidation no.

Oxidation

Reduction

Addition of oxygen atom

Removal of hydrogen atom

Losing of electron

Increase in oxidation no.

Removal of oxygen atom

Addition of hydrogen atom

Gaining of electron

A redox reaction in which an element undergoes oxidation and reduction at the same time.

Decrease in oxidation no.

Disproportionation reaction

the sum of the oxidation no. of all atoms in a species equals the overall charge carried by the species.

no. of proton - no. of electron associated with an atom assuming that all bonding electrons are belonging to the more electronegative atom

Rules of assigning oxidation number

The species on the left of the table are oxidizing agents. The species on the right of the table are reducing agents.

Electrochemical series is constructed by arranging the standard electrode potentials measured in an ascending order. i.e. the most negative value first. By convention, the reduction half equations are tabulated in the ECS. Therefore, standard electrode potential is also known as standard reduction potential.

Use of electrochemical series

Concentration and temperature dependence of electrode potential

Those oxidizing agents at the bottom of the list are strong oxidizing agents while those reducing agents at the top of the list are strong reducing agents.

Dependence on concentration

The values list in an ECS are all measured at standard conditions and all species involved are in their standard states. An increase in concentration of an oxidzing agent will shift the equilibrium position to the right and make the value more positive.

The value is only depending on concentration (including partial pressure in case of gas) and temperature, the size of the electrode has no effect on the value.

An increase in concentration of a reducing agent will shift the equilibrium position to the left and make the value more negative. Nernst Equation (Not required in HKAL syllabus) Q : reaction quotient

To predict the overall e.m.f. of an electrochemical cell

To predict the energetic feasibility of a redox reaction

Compare the strengths of oxidizing agents (or reducing agents) A redox reaction with a positive overall e.m.f. will be energetically feasible.

e.g. disproportionation of Cu+(aq) in water

The feasibility of a reaction is also depending on the kinetic stability (activation energy) of the system which is not related to the overall e.m.f. determined. To ensure the overall e.m.f. of the cell be positive, E(cell) = E(cathode) - E(anode) = E(more positive) - E(less positive)

Primary cell

IUPAC Cell diagram

Electrode potential (of a half-cell)

Half cell

Electrochemical cells

Electrochemical Cells

Electrochemical Series (ECS)

A set of symbols are used to represent the set up of an electrochemical cell.

Symbols used Depending on the emphasis and interpretation of the actual setup, it may be possible to construct more than one IUPAC cell diagram for a given chemical cell.

e.g. Zinc-carbon cell

e.g. Lead acid accumulator

Secondary cell (rechargeable cell)

e.g. Hydrogen-oxygen fuel cell Fuel cell

Some examples

For a given half-cell, no matter it is being used as a left-hand or right-hand electrode, the reduced form should be placed on the left if the reduced form and oxidized form are in the same phase. e.g. Pt(s) | Fe2+(aq) , Fe3+(aq) and Fe2+(aq) , Fe3+(aq) | Pt(s)

20. Electrochemical Cells.mmap - 9/5/2005 -

Definition : The difference between the charge on an electrode and the charge in the solution in which the electrode is immersed in.

Practically, it is not possible to determine the absolute potential of a half cell. Therefore, for the sake of comparison, the electrode potential of standard hydrogen electrode is defined as 0 V.

The value is depending on the activity of the electrons in the electrode

This value is also related to the nature, including concentration, of the electrolyte in which the electrode is immersed in.

By definition, overall e.m.f of a cell = E(right) - E(left). When hydrogen electrode is used as the left-hand electrode, E(right, the unknown electrode) = overall e.m.f. of the cell

The sign of the electrode potential of the unknown half cell is the same as the polarity of the right hand electrode (unknown half cell) in the cell diagram.

In some cases, the solution will be saturated with the salt and with some insoluble salt present. e.g. Ag(s)|AgCl(s)|Cl-(aq)

e.g. Pt(s) | Fe2+(aq) , Fe3+(aq)

e.g. C(graphite) | 2I-(aq) , I2(aq)

Platinized = covered with platinum black = covered with platinum powder

An inert electrode, e.g. Pt, is immersed in a mixture of the reduced and oxidized forms of the reagent concerned. e.g. Fe3+(aq) /Fe2+(aq) , I2(aq)/I-(aq)

The metal concerned will be used as the electrode immersed in a solution containing the salt of the metal. e.g. Cu(s) | CuSO4(aq)

By definition, the electrode potential of an unknown half-cell is the e.m.f. of an electrochemical cell measured at standard condition where a standard hydrogen electrode (SHE) is used as the left-hand electrode and the unknown electrode is used as the right-hand electrode.

Involving metal

Not involving metal

Standard hydrogen electrode

Pt(s) | H2(g) | 2H+(aq)

Use a potentiometer

Use a voltmeter with high impedance

To measure the e.m.f. of a cell, the current flowing through the external circuit should be kept as minimal as possible.

The use of porous partition will create a junction potential across the partition and the voltage measured will be different from the case if a salt bridge is used.

The salt bridge is used to provide cations and anions to balance the negative and positive charge cumulated in the two half cells during the discharge of the cell.

Only KNO3 and NH4Cl are used in the salt bridge since the migration speed of the cation and anion are similar and the voltage reading will not be affected.

The electrodes of the two half cells are connected together by an external wire.

The electrolytes of the two half cell are connected together by using a salt bridge.

In some electrochemical cell, a porous partition is used to separate the two electrolytes instead of using a salt bridge.

Measurement of e.m.f of an electrochemical cells.

Phase

Definition : region with identical properties. Phase is not the same as physical state. Usually, there are only 3 physical states i.e. solid, liquid and gas. e.g. Diamond and graphite are both in solid state but they are two different phases of carbon. Consider an ideal gas obeying ideal gas law PV=nRT, for a given amount of gas, P is a function of V and T. i.e. P(V,T).

By plotting a graph using P as the z-axis, V as the x-axis and T as the y-axis, a PVT surface of an ideal gas can be constructed.

PVT surface of of a real substance e.g. water

The isotherms of an ideal gas shows that no matter how low the temperature is, the ideal gas is still a gas and the volume is always inversely proportional to pressure.

Isotherms

The projection along the temperature axis (Isotherms)

Ideal gas

Isotherms

Real substance e.g. water

Real gas doesn't obey Boyle's law at low temperature. Real gas starts condensing at a temp. lower than the critical temp. of the substance.

Phase Equilibrium I : One-component system Isochors

PVT surface

As it is difficult to study a 3-dimensional surface, the projections along the volume axis and the temperature axis are usually studied separately.

Ideal gas

The isochors of an ideal gas shows that no matter how low the temperature, the ideal gas is still a gas and has zero volume at 0 K.

The diagram is known as phase diagram. critical temp : above this temperature, gas cannot be compressed into liquid without cooling. AB : sublimation curve - conditions that solid and vapour can coexist at equilibrium. BC : vaporization / boiling curve - conditions that water and vapour can coexist at equilibrium. BD : fusion / melting curve - conditions that solid and liquid can coexist at equilibrium. B : triple point - the conditions that solid, liquid and vapor can coexist at equilibrium. BE : supercooling curve - the condition that a liquid exists at a temp. lower than its melting point and the liquid is at a metastable state. This is because freezing also requires activation energy. If a liquid is cooled very calmly, the liquid may become supercooled without solidification. Real gas condenses at low temperature and becomes liquid and solid.

Phase diagram of water

Real substance e.g. water

The projection along the volume axis (isochors)

Special features of the phase diagram of water

Special features of the phase diagram of carbon dioxide

Special features of the phase diagram of carbon

21.Phase Equilibrium I . One-component system.mmap - 2005/5/9 -

The fusion curve of water has a negative slope. This implies that the melting point of water decreases with increasing pressure. When a great pressure is applied to ice, the open structure of ice will collapse and the ice will be turned to water. According to the Le Chatelier's principle, when a pressure is applied to an equilibrium system, the system will react in a way to minimize the effect of the change imposed. For ice, it will turn to water which has a smaller volume to lower the increase in pressure.

Like most substances, the fusion curve of carbon dioxide has a positive slope. This implies that the melting point of carbon dioxide increases with increasing pressure. The triple point pressure is higher than atmospheric pressure. This implies that dry ice will sublime directly to gaseous carbon dioxide without going through the liquid state.

It is possible to convert graphite to diamond by applying pressure to graphite. However, the rate of conversion is very slow as this involves rearrangement of atoms in solid state. If graphite is heated to molten state and allowed to crystallize under high pressure, artificial diamond can be made. Moreover, the crystallization is very fast comparing with the natural process of formation of diamond, the diamond crystal obtained will be small and may not be suitable for making jewelry.

The components cannot be separated by fractional distillation completely.

Positive deviation

Extreme negative deviation (a solution of involatile solute)

Negative deviation

Extreme positive deviation (2 immiscible liquids)

Azeotropic mixture (or constant boiling mixture) - A mixture with a constant composition when it is being boiled continuously. i.e. the composition of the liquid mixture = the composition of the vapour The 2 components vaporize independently and give a total vapour pressure higher than that of either component.

The lowering of vapour pressure will cause depression in melting point and elevation of boiling point of the solvent.

22. Phase Equilibrium II . Two-Component System.mmap - 19/5/2005 -

Deviation from Raoult's Law (non-ideal solutions)

Difference between 1-component and 2-component systems

Phase Equilibrium II : Two-Component System

Raoult's Law

For a 1-component system, the (saturated) vapour pressure (P) of a substance is a constant at a given temperature.

In a 1-component system, the properties of a substance is a function of P and T. This requires a 3-dimensional space to represent. Instead of studying the problem in a 4-dimensional space, the study will be limited to the case in which a liquid and its saturated vapour coexist in equilibrium.

In a 2-component system, the properties of a substance will be depending on P, T and X (mole fraction) which will need a 4-dimensional space to represent.

For a 2-component system, the (saturated) vapour pressure at a given temperature will be a function of X (mole fraction) of the component present.

The boiling point of an ideal solution varies linearly with the composition approximately.

Partition of a solute between 2 phases

Phase Equilibrium III : Three-Component Systems

23. Phase Equilibrium III . Three-Component Systems.mmap - 2005/5/23 -

Solvent Extraction (between 2 immiscible solvents)

Partition of a solute between 2 immiscible solvents

Partition Law - At a given temp. and at a state of equilibrium, the ratio of concentrations of a solute in two immiscible solvents is constant.

At equilibrium, the rates of diffusion of the solute between two immiscible solvents in opposite directions are the same.

As the unit of concentration will be canceled out in the calculation, the unit can be expressed in any unit and partition coefficient will have no unit eventually.

Most common organic solvent are lighter than water except chloroform(CHCl3), tetrachloromethane(CCl4) and 1,1,1-trichloroethane

e.g. The partition law is not applicable to the case where ethanoic acid is partitioned between hexane and water as ethanoic acid exists as dimmer in hexane (non-polar solvent) and as monomer in water (polar solvent)

If the density of the solvents are not known, when we say solute A is partitioned between solvent B and C, it usually means

The law is only application to the case if the solute has the same molecular structure in the two solvents.

Partition of a solution between a stationery phase and a mobile phase

Mobile phase

Stationery phase

e.g. different kind of solvent or mixture of solvents

e.g. silica gel on an aluminium foil

e.g. the microscopic water film on the surface of the chromatography paper

It would be more effective to carry out a solvent extraction in successive extractions, a small portion at a time, for a given amount of solvent.

Paper Chromatography (between a stationery phase and a mobile phase)

At a given temperature, with a given chromatography paper and a given solvent, the solute being analyzed is characterized by its Rf value

Analysis of amino acids - A mixture of colourless amino acids can be separated by paper chromatography. The positions of amino acids spot in the chromatograph can be detected by spraying with ninhydrin (a developer), which reacts with amino acids to yield highly coloured products (usually purple in colour).