chemistry synopsis handouts

BASIC

NSTC-6 PREPARATION PR

� Matter: Everything that has

� Mass: The amount of matte

� Element: A simple substan

divided further into more subs

classified into metals and non

� Compound: Two or more

combined together in a fixed p

� Mixture: A substance that c

more elements or compounds

� Homogenous Mixture

uniform composition through

� Heterogeneous Mixtu

mass.

� Chemical Symbols: A o

element to represent it conven

� Chemical Formulae: i

elements present and subscrip

INTERNATIONAL SY The modernized metric sy

SI. It has seven basic quantities.

S. �o. Property

1. Mass

2. Length

3. Time

4. Electric Current

5. Temperature

6. Amount of Subs

7. Luminous Intens

� Derived Quantities: Th

quantities. Example:

ASIC CONCEPTS PREPARED BY

SYNOPSIS HANOSAMA HASAN

ROGRAMME | AGA KHAN HIGHER SECOND

at has mass and occupies space.

matter present in a body.

ubstance that couldn’t be

re substances. (It’s further

d non-metals)

more elements chemically

fixed proportion.

that contains two or

ounds.

ure: A mixture that has

roughout its mass.

ixture: A mixture that has different composition thr

A one- or two-letter designation derived from the n

onveniently.

indicates the composition of a compound through

bscripts to indicate the relative number of atoms of ea

L SYSTEM OF UNITS (SI): tric system using the decimal system as its base is now

ities.

perty Abbreviation Unit Abbrevia

m Kilogram Kg

L Meter M

t Second S

urrent I Ampere A

T Kelvin K

f Substance n Mole Mol

Intensity candela cd

The quantities obtained by the multiplication or di

“The branch of Phy

which studies the co

structure, propertie

occurring inside th

referred as Chemis

ANDOUT 1

NDARY SCHOOL,

KARACHI.

1

ion throughout its

the name of an

rough symbols of

s of each element.

now referred as

breviation

Kg

M

S

A

K

Mol

cd

n or division of SI

Physical Science

e composition,

erties and changes

e the matter is

mistry.”

BASIC CONCEPTS PREPARED BY

SYNOPSIS HANDOUT 1 OSAMA HASAN

NSTC-6 PREPARATION PROGRAMME | AGA KHAN HIGHER SECONDARY SCHOOL,

KARACHI.

2

Density =Mass

Volume Unit =

kg

m�

Volume = length x

breadth x height

Unit = m x m x m

= m3

Temperature Measurements:

Three different scales are widely used for measuring temperature, namely, Celsius

(centigrade), Fahrenheit and Kelvin. The conversion formulae for them are as follows: °C

5=°F − 32

9

K = °C + 273

UNIT CONVERSIONS: SI units could be derived by combining prefixes with a root unit. Some basic prefixes are

as follows:

S. �o. Prefix Multiple Scientific �otation Abbreviation

1. Mega- 1,000,000 106 M

2. Kilo- 1,000 103 k

3. Hector 100 102 h

4. Deka- 10 101 da

5. Deci- 0.1 10-1 d

6. Centi- 0.01 10-2 c

7. Milli- 0.001 10-3 m

8. Micro 0.000,001 10-6 µ

9. Nano- 0.000,000,001 10-9 n

10. Pico 0.000,000,000,001 10-12 p

SIGNIFICANT FIGURES: “The reliable digits which are known with certainty.”

Rules:

� All non-zero numbers are significant

� Zeros placed in between non-zero numbers or after a decimal point are significant.

� Zero that locates the decimal point on number less that (1) is non-significant.

� Zero that locates the decimal point on number larger than (1) is significant.

CHEMICAL LANGUAGE PREPARED BY



KARACHI.

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LAWS OF CHEMICAL COMBINATION:

� Law of Conservation of Mass (Indestructibility of Matter): “In a

chemical reaction, the total mass of reactants is equal to the total mass of products i.e.

matter can neither be created nor be destroyed.” It was put forward by Lavoiser in 1785.

� Law of Definite Proportion: “When elements combine to form compounds, they

do so in definite proportions by weight. A compound always has the same proportion by

whatever method it is prepared.”It was given by Joseph Proust in 1799.

� Law of Multiple Proportions: "When two elements combine to form more than

one compound, the different weights of one of the elements combining with a fixed

weight of the other, bears a simple ratio to each other." The law was proved by Berzilius

& Stas.

� Law of Reciprocal or Equivalent Proportion: "The weights of two

elements combining separately with a fixed weight of a third element do so with each

other in multiple ratios."

BASIC CONCEPTS:

� Mole: “Atomic mass of an element, molecular mass of compounds or formula mass of

an ionic substance when expressed in grams.”

Mole = ��

� Avogadro’s number: “The constant number (6.02 x 1023) atoms, molecules or ions

contained by one mole of any substance.”

� Atom: the basic building block of an element that retains the properties of that element

and can enter into a chemical reaction.

� Molecule: The smallest particle of a substance that retains the physical and chemical

properties of that substance.

� Empirical Formula: The simplest formula of a compound which represents the

elements present in a compound and the simplest ratio between them.

DETERMINATION OF EMPIRICAL FORMULA: S. No Steps

1. Elements present in the compound

2. Mass of each Element

3. Percentage of Each Element (��

�� × ��)

4. Mole Ratio (�� !��

"��#� �� !��) 5. Atomic Ratio (

$�� %��#�&�� $�� %��#�)




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� Molecular Formula: The formula which represents the elements present and the

exact number of atoms of different elements present in a molecule of that compound.

DETERMINATION OF MOLECULAR FORMULA:

Molecular formula = [Empirical Formula] n

Where n = �� '(� ��

)*��(� ��

� Molar Concentrations: The molar concentration (M) is the number of moles of the

solute contained in one cubic decimeter (dm3) or litre of solution.

Molarity = M = No. of moles of solute1 litre of solution

� Molal Concentrations: A molal concentration (m) is the number of moles of the

solute per kilogram of solvent.

Molality = m = No. of moles of solute1 kilogram of solvent

� % Aqueous Solutions of Compounds: A 10 % aqueous solution of a solute

contains 10 g of solute per 100 g solution i.e. 10 g of solute mixed with 90 g water.

� Percentage Concentrations: The general formula is:

� % Concentratiions = No. of grams of solute

No. of grams of solvent x 100%

� Stoichiometry: the study of relationship between the amount of reactant or product

involved in a chemical equation based on balanced chemical equation.

DIRECTIONS FOR STOICHIOMETRY:

Mass-Mass

Relationship

Mass-Mole Relationship Volume-Volume

Relationship

Mass-Volume

Relationship

E�� $�� × F��#�#��

Formula Mass × CoefHicient

Formula Mass

E�� $�� × F��#�#�� → $�� J�#��

No. of Moles → No. of Moles Required

Formula Mass × CoefHicient → Volume

Mass → Moles→ Moles Required→ Volume




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� Limiting Reactant: “The reactant which is consumed first of all during the

chemical reaction is called Limiting Reactant.”

DIRECTIONS FOR LIMITING REACTANT:

O" + QR + SF → T

Step 1 Step 2 Step 3

O" → T

Mass of A → Mole of A → Mole of D

QR → T

Mass of B → Mole of B → Mole of D

SF → T

Mass of C → Mole of C → Mole of D

The source that yields the least quantity of moles would be the limiting reactant.

SOME FACTS TO REMEMBER:

1. 1 mole of a substanceYelement, compound\= 6.02 × 10_` particles Yatoms, molecules\

2. 1 mole of a gas = 22.4 dm` at S. T. P Y0°C and 1 atm pressure\

3. 1 mole of an ionic compoundYNaCl\= 6.02 × 10_` Na ions and 6.02 × 10_` Cl ions = 2 × 6.02 × 10_` ions

4. 1 mole of covalent compound YHCl\ = 6.02 × 10_` molecules of HCl= 6.02 × 10_` atoms of H and 6.02 × 10_` atoms of Cl= 2 × 6.02 × 10_` atoms.

5. 1 atm = 14.7 psi gpounds per square inch, lb in_i j =760 mm of Hg=29.921 in. of Hg

= 760 torr = 20 tons = 1.01325 bar

= 13.6 mm of H20

CHEMICAL TECHNIQUES PREPARED BY



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SEPARATION TECHNIQUES: The separated products could differ in chemical properties or some physical property, such as

size, or crystal modification or other separation into different components.

“A separation process is used to transform a mixture of substances into two or more distinct

products.”

Separation

Between

# Technique Explanation

Solute Solvent

S

O

L

I

D

S

&

S

O

L

I

D

S

1 Destructive

Distillation

Pyrolysis conducted in a distillation apparatus (retort) to

allow the volatile products to be collected.

2 Melting Alloys can be separated by melting them on their

respective temperatures.

3 Wind winnowing An agricultural method developed by ancient cultures for

separating grain from chaff.

S

O

L

I

D

S

&

L

I

Q

U

I

D

S

1 Crystallization A chemical solid-liquid separation technique, in which

mass transfer of a solute from the liquid solution to a pure

solid crystalline phase occurs.

2 Filtration Mechanical or physical operation which is used for the

separation of solids from liquids by interposing a medium

to fluid, through which it can pass.

3 Decantation Process for the separation of mixtures, carefully pouring a

solution from a container, leaving the precipitate

(sediments) in the bottom of the container

4 Distillation Used for separating solids dissolved in liquid

5 Dissolved Air

Floatation

Water treatment process that clarifies wastewaters (or

other waters) by the removal of suspended matter such as

oil or solids.

6 Solid-phase

extraction (SPE)

A separation process that is used to remove solid or semi-

solid compounds from a mixture of impurities based on

their physical and chemical properties

7 Flocculation Process by which fine particulates are caused to clump

together into floc or flakes. The floc or flakes may then

float to the top of the liquid, settle to the bottom of the

liquid, or can be readily filtered from the liquid

8 Cyclonic

Separation

Removing particulates from water stream, without the use

of filters, through vortex separation.

S

O

L

I

D

&

G

A

S

1 Cyclonic

Separation

Removing particulates from an air or gas, without the use

of filters, through vortex separation.

2 Filtration Mechanical or physical operation which is used for the

separation of solids from gases by interposing a medium

to fluid, through which it can pass.




KARACHI.

2 GENERAL SEPARATION METHODS:

L

Q

U

I

D

S

&

S

O

L

I

D

S

1 Decantation Separating a liquid from a solid sediment

2 Dry Distillation Heating of solid materials to produce liquid or gaseous

products (which may condense into solids).

3 Drying Removing liquid from a solid by vaporizing it

5 Expeller Press The liquids in solids are extracted by applying high

pressures.

6 n-hexane Method Separation of liquids from solids by dissolving it in

hexane and later performing distillation of the solution.

7 Super Critical

Fluid Extraction

By exposing the solid to Super Critical Fluid which

dissolves the liquid and then separating them by a mere

change of pressure or temperature.

L

I

Q

U

I

D

S

&

L

I

Q

U

I

D

S

1 Centrifugation Involves the use of the centrifugal force for the separation

of mixtures.

2 Distillation

Separating mixtures based on differences in their

volatilities in a boiling liquid mixture.

3 Fractional

Distillation

Process in which a fractioning column is used in

distillation apparatus to separate components of a liquid

mixture that have different boiling points.

4 Vacuum

Distillation

distillation whereby the pressure above the liquid mixture

to be distilled is reduced to less than its vapor pressure

(usually less than atmospheric pressure) causing

evaporation of the most volatile liquid(s) (those with the

lowest boiling points).

5 Solvent

Extraction

Method

Method to separate compounds based on their relative

solubilities in two different immiscible liquids.

6 Electrophoresis Organic molecules, such as protein are placed in a gel. A

voltage is applied and the molecules move through the gel

because they are charged. The gel restricts the motion so

that different proteins will make different amounts of

progress in any given time.

7 Fractional

Freezing

Process used in process engineering and chemistry to

separate two liquids with different melting points

8 API oil-water

separator

A gravity separation device designed by using Stokes Law

to define the rise velocity of oil droplets based on their

density and size.

Liquid & Gas 1 Demister(Vapour) Removing liquid droplets from gas streams.

Gas & Solid ??? ???

Gas & Liquid ??? ???

Gas & Gas 1 Elutriation/ air

classification

Process for separating lighter particles from heavier ones

using a vertically-directed stream of gas.




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Chromatography: Involves the separation of different dissolved substances as they travel

through a material. The dissolved substances are separated based on their interaction with the

stationary phase.

Steam Distillation: Special type of distillation (a separation process) for temperature

sensitive materials like natural aromatic compounds.

Evaporation: Process by which molecules in a liquid state (e.g. water) spontaneously

become gaseous (e.g. water vapor).

Sublimation: Conversion of solid molecules to gaseous state.

Precipitation: The formation of a solid in a solution during a chemical reaction

Fusion: Conversion of solid molecules to a liquid state.

Gravity separation: An industrial method of separating two components from a suspension

or any other homogeneous mixture where separating the components with gravity is sufficiently

practical.

Sieving: Separates wanted/desired elements from unwanted material using a tool such as a

mesh, net or other filtration or distillation methods, but it is also used for classification of

powders by particle size

Stripping is a chemical separation process where one or more components are removed from a

liquid stream by a vapor stream.

S. 0o. Technique Respective Property

1. Centrifugation Density difference

2. Cyclonic Separation Rotational Effects/Gravity

3. Elutriation/ air classification Size of the particles

4. API oil-water separator Density & size of oil-droplets

5. Fractional Freezing Melting Point

6. Fractional Distillation Boiling Point

7. Electrophoresis Electric Charge

8. Solvent Extraction Method Relative Solubility

9. Vacuum Distillation Pressure

10. Distillation Volatility

11. Dry Distillation Heating on high temperatures

12. Flocculation Density

13. Expeller Press Pressure

TERMINOLOGY:

Miscible: The liquids that can be dissolved in each other

Immiscible: The liquids that do not dissolved and could be separated.

Alloy: A Solid Solution

Ore: A natural deposit containing a mineral of an element to be extracted

Distillate: The material in a distillation apparatus that is collected in the receiver.




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Distilland: The material in a distillation apparatus that is to be distilled.

SOLUBILITY:

SOME FACTS TO REMEMBER:

Soluble Non-Soluble

� All compounds of 0itrates (0O3-), Acetates

(CH3COO-) and Chlorates (???).

� All normal Carbonates (CO3-), Phosphates

(PO33-), Silicates (???) & Sulfides (???)

� All chlorides. � Chlorides of Silver, Mercury and Lead

� All sulfates � Sulfates of lead, barium, strontium and calcium.

� Hydroxides of Ca, Ba, Sr, 0a, K and 0H4. � All Hydroxides

� Common compounds of 0a, K and 0H4. �

TRENDS OF SOLUBILITY:

State Temperature Effect Pressure Effect

Solid Usually increases with

temperature increase

Little Effect

Gas Usually decreases with

temperature increase

Solubility varies in dirtect

proportion to the pressure

applied to it (Henry’s Law1)

FACTOR’S AFFECTING RATE OF SOLUBILITY:

The following procedures INCREASE the rate of solubility:

Pulverizing Increases surface exposed to solvent

Stirring Brings more solvent that is unsaturated into contact with solute.

Heating Increases molecular action and gives rise to mixing by convection currents.

(This heating affects the solubility as well as the rate of solubility)

1 Henry’s Law: “The solubility of a gas (unless the gas is very soluble) is directly proportional to the pressure

applied to the gas”




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QUALITATIVE TESTS FOR GASES:

Gas Tests Observations

Ammonia

0H3

Smell Cautiously Sharp Odour

Test with Litmus Red litmus turns blue

Expose to HCl fumes White fumes (NH4Cl) formed

Carbon Dioxide

CO2

Pass through lime water

(Ca(OH)2)

White precipitates formed

Carbon Monoxide

CO

Burn it and pass product

through lime water

White precipitates formed

Chlorine

Cl2

Smell Cautiously Irritating Odour

Observe color Yellowish Green

Hydrogen

H2

Allow it to mix with some

air, then ignite

Gas explodes

Burn it- trap product Burns with blue flame- product H2O turns

cobalt chloride paper from blue to pink.

Hydrogen

Chloride

HCl

Smell Cautiously Choking Odour

Exhale over the gas Vapour fumes form

Dissolve in water and test

with litmus

Blue litmus turns red

Add AgNO3 to the solution White precipitate formed

Hydrogen Sulfide

H2S

Smell Cautiously Rotten Eggs’ Odour

Test with moist lead acetate

paper

Turns brown black (PbS)

0itric Oxide

0O

Expose to air

Colourless gas turns reddish brown

0itrous Oxide

02O

Insert glowing splint Bursts into flame

Add Nitric oxide gas Remains colourless

Oxygen

O2

Insert glowing splint Bursts into flame

Add Nitric oxide gas Turns Reddish Brown

Sulfur Dioxide

SO2

Smell Cautiously Choking Odour

Allow it to bubble into purple

potassium permanganate

solution

Solution becomes colourless.




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QUALITATIVE TESTS FOR METALS:

Method Procedure Observations

Flame

Test

Clean a platinum wire by dipping it into dilute HNO3 and

heating in the Bunsen flame. Repeat until the flame is

colorless. Dip a platinum wire into the substance being

tested (either solid or solution), and them hold it in the hot

outer part of the Bunsen flame.

Na+ Golden Yellow

K+ Violet

Ca2+ Orange Red

Sr2+ Bright Red

Ba2+ Yellowish Green

Cu2+ Bluish Green

Li+ Crimson

Hydrogen

Sulfide

Test

Bubble Hydrogen Sulfide gas through the solution of a salt

of the metal being tested. Check the color of the precipitate

formed.

Pb Brown Black

Cu Black

Ag Black

Hg Black

Ni Black

Fe Black

Cd Yellow

As Light Yellow

Sb Orange

Zn White

Bi Brown

Ash Test

Add a few drops of cobalt nitrate to the saturated solution

sample of the salt. Dip a piece of filter paper in the solution.

Dry it and then burn it gently in the Bunsen flame.

Al3+ Blue

Zn2+ Green

Mg2+ Light Pink

THE PERIODIC TABLE PREPARED BY



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PERIODIC DEVELOPMENT:

1. 1817 Germany Law of Triads J. W. Dobereiner

2. 1865 England Law of Octaves A. R. Newland

3. 1864-69 Germany Meyer’s Hypothesis Lothar Meyer

4. 1870 Russia Mendeleev’s Periodic Law Dimitri I. Mendeleev

5. 1911 England Modern Periodic Law Henry Mosely

PERIODIC LAWS:

� Law of Triads: “If three elements are arranged in according to their atomic weights,

such that the atomic weight of the middle element is mean of the other two, then they will

exhibit similar properties.”

� Law of Octaves: “If elements are arranged according to increasing atomic weights,

a given set of properties recurs at every eighth place, by analogy with the musical scales.”

� Meyer’s Hypothesis: “The physical properties of elements are the periodic

function of their atomic weights.”

� Mendeleev’s Periodic Law: “The properties of the elements as well as the

formulae and properties of their compounds depend in a periodic manne on the atomic

weight of elements.”

� Modern Periodic Law: “The physical and chemical properties of elements are

periodic function of their atomic numbers.” OR “The physical and chemical properties of

the elements are periodic function of their electronic configuration, which vary with

increasing atomic numbers in a periodic manner.”

DOBEREINER’S TRIADS: Expected Mass Triads with Atomic Masses

7 + 39

2= 23

Li (7) Na(23) K(39)

35.5 + 127

2= 81.25

Cl (35.5) Br(80) I(127)

NEWLAND’S LAW OF OCTAVES: The properties of every eighth element were

similar to those of first.

Li Be B C � O F

7 9 11 12 14 16 19

�a Mg Al Si P S Cl

23 24 27 28 31 33 35

K Ca

39 40

LOTHAR MEYER’S CLASSIFICATION: Meyer plotted atomic volume (mass + density) against atomic weight and obtained the curves, in

which elements with similar physical properties occupied similar position in the curve. i.e.:




KARACHI.

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• Alkali Metals occupied the peak positions

• Inert gases, halogens and elements giving acidic oxides occur on ascending positions

• Elements giving basic oxides occur on descending positions.

• Elements producing amphoteric oxides occupy the curves.

MODERN PERIODIC TABLE: • So far, 110 elements have been discovered.

• Out of which 92 are naturally occurring and 18 are artificially made elements.

1 2

H He

3 4 5 6 7 8 9 10

Li Be B C � O F �e

11 12 13 14 15 16 17 18

�a Mg Al Si P S Cl Ar

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36

K Ca Sc Ti V Cr Mn Fe Co �i Cu Zn Ga Ge As Se Br Kr

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54

Rb Sr Y Zr �b Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

87 88 89 104 105 106 107 108 109 110

Fr Ra Ac Ku Ha

Lanthanide

Series

58 59 60 61 62 63 64 65 66 67 68 69 70 71

Ce Pr �d Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Actinide

Series

90 91 92 93 94 95 96 97 98 99 100 101 102 103

Th Pa U �p Pu Am Cm Bk Cf Es Fm Md �o Lr

PERIODIC TRENDS:

� Atomic Radius: The half of the distance between

two nuclei present in a homo-nuclear diatomic molecule.

� It’s measured in Angstrom A°.

� Atomic Radius ∝ �

�� !��"�

� Atomic Radius ∝ Electron Shell

� Ionic Radius: The distance between the nucleus of an




KARACHI.

3

ion and the point upto which nucleus has influenced

electron is called Ionic Radius. Or the radius of a

spherical ion. � It’s also measured in Angstrom A°.

� Cations have smaller radii than neutral atom.

� Anions have larger radii than the neutral atom.

� Ionic Radius ∝ Electron Cloud

� Ionic Radius ∝ �

�� !��"�

In f-block, atomic radius decreases

with the increase in atomic number

vertically because in this case,

electrons are added to the inner d or

f sub-shells.

� Ionization Potential: The amount of energy

required to remove most loosely bonded electron from

the outer most shell of an atom in its gaseous state.

� Its unit is Kilo-joules per mole (KJ/mol) and is also

measures in electron volts (eV)

� Ionization Potential ∝ �

./01�2�� 34 ��1�32 4�35 2��0

� Ionization Potential ∝ �

�� !��"�

� The removal of electron becomes easier towards

f-shell > d-shell > p-shell

� Electron Affinity: The negative of the energy

change that occurs when an electron is accepted by an

atom in the gaseous state to form an anion.

� A + e6 → A6 + (Energy released) � It’s measured in Kilo-joules per mole (KJ/mol).

� Electron Af=inity ∝ �

>135/� ?�@/�0

� Electron Af=inity ∝ �

�� !��"�

� Elements with stable electronic configuration have zero

electron affinity.

� The large values of electron affinity show that the

elements are very strong oxidizing agents.

� Chlorine has the highest value of electron affinity.

� Electronegativity: The force of attraction by which

an atom attracts a shared pair of electrons.

� It’s measured in electron volts (eV)

� Electronegativity of Fluorine is fixed as 4, which is the

highest value.

� Electronegativity ∝ �

>135/� B/C�

� Electronegativity ∝ Valence Electrons

� Electronegativity ∝ Ionization Potential

� Electronegativity ∝ Electron Af=inity

STATES


KINETIC MOLECULAR T“Kinetic molecular theory explain

different states o

To be noted that:

� The molecules of a gas are in s

seem to resist the force of grav

� The speed and direction of mo

are at all times unpredictable

FOR GASES: � All gases consist of tiny partic

� The molecular size is nearly ne

� The gas molecules are in conti

� Molecules collide with one ano

� Gas pressure is the collision of

� In an ideal gas, no attractive or

� Average kinetic energy of mol

FOR LIQUIDS: � In a liquid, molecules are fairly

� They consist of clusters in whi

� Molecules constitute 70 % of t

� They have a definite volume b

FOR SOLIDS: � In a solid, molecules are closel

� The forces of attraction betwee

� They are arranged in a fixed pa

� They have a definite shape.

GASES: GAS LAWS:

Boyle’s Law: “The volume of

exerted on it provided that tempera

� Mathematically, V � �

� � V

� On plotting volume of a gas ag

� While, on plotting volume of a

obtained.

Gases

TES OF MATTER PREPARED BY



THEORY plains the physical properties of

tes of matter.”

re in such rapid motion that they

gravity.

of motion of any given molecule

ble.

particles called molecules except inert gases which co

arly negligible as compared to the distances between t

continuous motion in straight paths but in random di

ne another and with the walls of container, elastically

ion of gas molecules with the walls of container.

tive or repulsive forces exist in between molecules.

of molecules depends upon the absolute temperature. (

e fairly randomly arranged.

in which they are very close together.

% of the total volume.

ume but do not have any definite shape.

closely packed.

between them are very strong.

ixed pattern, forming a lattice of vibrating masses.

me of a given mass of a gas is inversely proportional t

emperature is kept constant.”

V � k �

�� PV � k � ��

gas against the total pressure, a parabolic curve is obt

e of a gas against the reciprocal of total pressure, a st

Matter

Liquids Solids Pla

In different st

particles of any

differ only in

which they ar

are not diffe

HANDOUT 5

NDARY SCHOOL,

KARACHI.

1

ich consist of atoms.

ween the molecules.

om directions,

tically.

ture. (K.E)avg � T

ional to the pressure

is obtained.

re, a straight line is

Plasma

rent states of matter,

f any given substances

nly in the manner in

ey are arranged and

t different in kind.

STATES OF MATTER PREPARED BY



KARACHI.

2

Charles’s Law: “The volume of a given mass of a gas is directly proportional to its absolute

Temperature, provided that pressure is kept constant.”

� Mathematically, V � T � V � k T � ��

� k � �

��

��

� When the volume is plotted against the temperature, a straight line is obtained.

� At absolute zero the volume of gases is considered to become zero and all the motion ceases

to exist.

� For each Celsius degree rise in temperature, the volume of a gas expands by 1/273 of it’s

volume at 0°C

Avogadro’s Law: “The volume of a gas is directly proportional to the number of molecules

of the gases at constant pressure and temperature.”

� Mathematically, � � � (at constant temperature and pressure)

Graham’s Law of Diffusion: “The rates of diffusion of gases are inversely proportional

to the square root of their densities under same conditions of temperature and pressure.”

� Mathematically, r � �

√�� r � �

√� � ��

��

��

��

��

� For Effusion, � !"#$% &#'(�

� !"#$% &#'(�� !"#$% )*&(�

� !"#$% )*&( �

Dalton’s Law of Partial Pressure: “The total pressure of a mixture of non-reacting

gases is the sum of the partial pressures of various gases in the mixture.”

� Mathematically, P = +� + +- + +. + ⋯ + +0

� �*�&#*1 ��(""!�( $ * 2*"

�$&*1 ��(""!�(� �$1(" $ &3( 4*"

�$&*1 �$1("

Gay-Lussac’s Law (Law of Pressures or Law of Combining Volumes):

“At constant volume, the pressure of a given mass of a gas is directly proportional to its absolute

temperature.”

� Mathematically, P � T � P � kT � ��

� k � ��

��

��

Ideal Gas Equation or Equation of

State: PV = nRT.

This equation is called the equation of state because

when the four variables – pressure, volume,

temperature and number of moles are specified, we

can define the state of gas.

BEHAVIOUR OF GASES: � Real gases deviate from ideal behavior at low

temperature and high pressure.

� Real gases obey the ideal behavior at high temperature and low pressure.

GAS DENSITY:

� Gases are less dense than liquids and solids.

� Gas Density � Pressure

� Gas Density � Molar Masses � Gas Density � 1 TC

Units of Universal Gas Constant

Value of R When

0.082058 L.atm.mol-1.K

-1 P is in atm

62.364 L.torr. mol-1.K

-1 P is in torr

8.3145 J. mol-1.K

-1 P is in

Pa;V is m3




KARACHI.

3

PRESSURE: Pressure is measured by a device named “Manometer”. ∆P,

the difference in barometric pressure Pbar and the pressure of

gas Pgas is given by the difference in mercury levels in open

and closed arms of the device.∆P = ∆h = hopen - hclosed

If Pgas is greater Pbar, Pgas = Pbar + ∆h

If Pgas is less Pbar, Pgas = Pbar – ∆h

P = g . d . h

Where g = 908m/s2;d= density, h= height of column

LIQUIDS: VISCOSITY (η): The force of internal friction which opposes the displacement of different

layers relative to one another.

� The S.I unit of viscosity is (N. sm-C )

� Viscosity is an intensive property.

� Viscosity � Molecular Size

� Viscosity � Intermolecular Attraction

� Viscosity � 1TemperatureC

� Viscosity � Molecualr Shape � The grading of motor oil is done on the basis of viscosity, e.g. grade 30 oil is less viscous

than the grade 40 oil.

SURFACE TENSION (γ): The force per unit length or energy per unit area of the surface

of a liquid.

� Its units are N/m, J/m2, Dynes/cm, erg/cm

2

� Surface Tension is an intensive property.

� Surface Tension � Intermolecular forces

� Surface Tension � Hydrogen Bonding

� Surface Tension � 1TemperatureC

VAPOUR PRESSURE: The pressure exerted by the vapours when they are in equilibrium

with the liquid phase.

� It’s constant at constant temperature.

� Vapour pressure depends upon the nature of liquids.

� Vapour Pressure � Temperature

� Vapour Pressure � 1Intermolecular ForcesC

� The vapour pressure of an aqueous solution is always lowered by addition of more solute.

� On lowering of vapour pressure, the freezing point is also lowered, while the boiling point is

raised.

STATES


SOLIDS:

TYPES OF CRYSTALS:

Crystal Structural

Particles

Pri

Fo

Ionic

Crystals

Cations and

Anions

Elect

Attr

Covalent

Crystals

Atoms Cov

Metallic

Crystals

Cations and

delocalized

electrons

Metal

Molecular

Crystal

(#on-polar)

Atoms or non-

polar molecules

Disp

Fo

Molecular

Crystal

(Polar)

Polar

Molecules

Disp

for

dipol

attra

Hydrogen

Bonded

Molecules with

H bonded to N,

O or F

Hyd

Bo

Crystalline Solids

Characteristic Geometrica

Sharp Melting Point

Different phy.prop. in dif. di

Posses Symmetry

TES OF MATTER PREPARED BY



Principal

Forces

Properties

Electrostatic

Attractions

� Hard, Brittle

� High M.P

� Nonconductors as solids but good

conductors as liquids

� Indefinite growth of crystal.

� Many are soluble in polar solvents.

Covalent

Bond

� Very hard

� High Melting Point or Sublime

� Low Density

� High Refractive Index

� Most are non-conductors of

electricity

Metallic Bond � Lustre, Malleable, Ductile

� High M.P

� Conductors of Electricity and Heat

Dispersion

Forces

� M.P (Extremely low to moderate,

varying with mol. Wt.)

� Readily sublime in some cases

� Soft

� Soluble in non-polar solvents

Dispersion

forces &

dipole dipole

attractions

� Low Melting Point

� Non-conduction of heat and

electricity

� Soluble in other polar and some

non-polar solvents

Hydrogen

Bonds

� M.P (low to moderate)

� Soluble in other hydrogen bonded

liquids and some polar liquids.

Solids

olids

etrical Shape

Point

dif. directions

etry

Amorphous Solids

Non repeitive 3-D struct

Melting Point varies

Same phy. prop. in all direc

Non-symetrical

HANDOUT 5

NDARY SCHOOL,

KARACHI.

4

Example

ood

ents.

NaCl, CAF2,

K2S, MgO

Graphite,

Diamond, Silica,

Carborundum,

SiC, AlN, SiO2

Heat

Metallic

Crystals

Ice, I2, CO2,

HCl, H2S,

CHCl2,

(CH3)2O,

(CH3)2CO

ded

H2O, HF, NH3,

CH3OH,

CH3COOH

lids

tructure

aries

l directions




KARACHI.

5

ISOMORPHISM: The phenomenon in which two substances have the same crystal structure.

E.g. CaCO3 & NaNO3 exist in trigonal crystals|ZnSO4 & NiSO4 exist in Orthorhombic Crystal

Properties:

� They have different physical and chemical properties.

� They have same empirical formula.

� Their solutions on mixing form mixed crystals.

POLYMORPHISM: The phenomenon in which the substance can exist in more than one

crystalline form. E.g. CaCO3 occurs in two crystalline forms (i) Calcite (trigonal) (ii) Aragonite

(orthorhombic)

UNIT CELL: The basic structural unity which when separated in three dimensions generates

the crystal structure.

CELL DIMENSIONS: #ames Sides Angles Examples

� Cubic V � W � X Y � Z � [ � 90 ° NaCl, ZnS

� Tetragonal V � W ≠ X Y � Z � [ � 90 ° SnO2

� Orthorhombic V ≠ W ≠ X Y � Z � [ � 90 ° FeSO4.7H2O

� Trigonal V � W � X Y � Z � [ ≠ 90 ° CaCO3, NaNO3

� Hexagonal V � W ≠ X Y � Z � 90 °; [ � 120° Graphite, snow flakes

� Monoclinic V ≠ W ≠ X Y � [ � 90 °; Z ≠ 90° CuSO4.5H2O

� Triclinic V ≠ W ≠ X Y ≠ Z ≠ [ ≠ 90 ° K2Cr2O7

CHANGE OF STATE:

Critical Pressure: The minimum pressure to liquefy a gas at its critical temperature.

Triple Point: The only temperature and pressure at which three phase a pure substance can

exist in equilibrium with one another in a system containing only the pure substance.

Boiling Point: the temperature at which the vapour pressure of the liquid equals the outside

pressure.

Boiling Point � Pressure

Effect of Concentration on Boiling Point and Freezing Point:

Type Concentrations Boiling

Point

Freezing

Point

Examples

� Molecular #on-ionizing 1 100.51°C -1.86°C Sugar, Urea

� Molecular Ionizing 2 101.02°C -3.72°C HCl

� Ionic dissociated 2 101.02°C -3.72°C NaCl

� Ionic dissociated 3 101.53°C -5.58°C CaCl2, Cu(NO3)2

This indicates that in a 1 m solution of molecules that do not dissociate, boiling and

freezing point differs with a change of (0.51°C) and (-1.86°C) respectively.




KARACHI.

6

Melting Point: the temperature at which there is equilibrium between solid and liquid phases.

Effect of Pressure on Melting Point:

Condition Situation

Substance Expands on Melting bcdef�g +hf�e � +icjjkic

Substance Contracts on Melting bcdef�g +hf�e � 1+icjjkicC

LATENT HEAT VALUES:

Latent Heat of Fusion: The amount of heat required to transform 1 kg of ice into water

completely at O°C.

� Its value is 80 cal/g or 3.34 x 102 J/g.

Latent Heat of Vaporization: The amount of heat required to transform 1 kg of water

into gas (steam) completely at 100°C.

� Its value is 540 cal/g or 2.26 x 103 J/g

ATOMIC STRUCTURE PREPARED BY



KARACHI

1

ATOMIC DEVELOPMENT:

1. England Michael Faraday Indicated the existence of electrons

2. America Robert Millikan Measured the charge on an electron

3. Germany William Crookes Showed the existence of electrons in atoms.

4. Goldstein Showed the existence of protons in atoms.

5. Henry Becquerel Through Radioactivity confirmed the presence of electrons

and protons

6. England James Chadwick Showed the presence of neutrons

PROPERTIES OF CATHODE RAYS: � Produce sharp shadows � Carry a negative charge

� Emerge normally from cathode � Easily deflectable by an electrostatic field

� Penetrate small thicknesses of matter

without any perforations

� Easily deflectable by a magnetic field

� Can exert mechanical pressure

RADIOACTIVITY: “The spontaneous breakdown of an atom by emission of particles and radiations”

Types:

� α- particles (doubly charged He2+

nuclei) � β- particles

� γ- particles (very short wavelength

electromagnetic rays)

� Positron emission

� Electron Capture

PROPERTIES OF:

Properties α- Rays β- Rays γ- Rays

1. Nature 2He4 Nucleus Fast moving electrons very short wavelength

electromagnetic rays

2. Charge Positive Negative Neutral

3. Velocity 3 x 106 to 3 x 10

7 m/s 9 x 10

7 to 27 x 10

7 m/s 3 x 10

8 m/s

4. Penetration Power Very less Intermediate Very high.100 times of β

5. Ionization Energy Very high Very Small Poor ionizers

6. Fluorescence Few substance only Many substance Not possible

X-RAYS: “The high energetic radiation emitted by metals and glass on bombardment with cathode rays,

that cause fluorescence in a variety of substance is called X-rays or Rontgen Rays ( amed after

the discoverer W. Rontgen).”

Characteristics:

� Short Wave Radiations � Extremely high penetration power

� High Ionization Energy � Electromagnetic in nature

� X rays ∝

�

�

� Arise from anode as a result of cathode

rays’ bombardment.




KARACHI

2

SPECTROSCOPY: “The branch of chemistry which deals with the study of absorption or emission of radiation is

called Spectroscopy.”

Planck’s Quantum Theory:

� The emission of radiation is due to the vibrations of charged particles (electrons) in the body.

� The emission is not continuous but in discrete bundles/packets of energy called “Quanta”.

This emitted radiation propagates in the form of waves.

� The energy associated with each quantum for a particular radiation frequency ν is given by:

E = hν

Where h = Planck’s constant = 6.625 x 10-35

J.sec = 6.625 x 10-37

erg.sec

� A body can emit or absorb either one quantum (hν) of energy or some whole number multiple

of it.

Spectrum: A band of rays of different wavelengths obtained from the decomposition of

radiations is called Spectrum.

� Emission Spectra � Absorption Spectra

Emission Spectrum:

When the radiation (absorbed by the object from an electric arc) is passed through a prism in a

spectroscope, it is decomposed into component wavelengths to form an image called an

Emission Spectrum.

� Continuous Spectrum � Line Spectrum

Continuous Spectrum: When a dispersed light is allowed to fall on a photographic plate,

the colors from violet to red (VIBGYOR) are seen without any line of demarcation. Such a

spectrum is called Continuous Spectrum.

Line Spectrum: If dispersed light falling on the photographic plate is sharp, distinct and

well defined by lines of demarcation then such a spectrum is called Line Spectrum.

Hydrogen Spectrum:

Wavelength Series n1 n2 Spectral Region

Less than 4000°A Lyman Series 1 2, 3, 4, … Ultra violet Region

B/w 4000°A to 7000°A Balmer Series 2 3, 4, 5, … Visible Region

More than 7000°A Paschen Series 3 4, 5, 6, … Near Infra Red Region

More than 7000°A Brackett Series 4 5, 6, 7, … Far Infra Red Region

More than 7000°A Pfund Series 5 6, 7, 8, … Radio Waves Region

QUANTUM NUMBERS: “The numbers that describe the distribution of electrons in atoms are called Quantum

umbers.”

Quantum Numbers are derived from the mathematical solution of Schrodinger’s equation

for the Hydrogen atom. It has following four types:

� Principal Quantum Number � Magnetic Quantum Number

� Azimuthal Quantum Number � Spin Quantum Number



NSTC-6 PREPARATION PROGRAMME | AGA KHAN HIGHER SECONDARY SCHOOL, KARACHI

3

QUANTUM

NO.

FORMULA VALUES EXPLAINS ABOUT MISCELLANEOUS

Principal

n

All integral

values

1, 2, 3, 4, …

� Principal Energy State

� Size of Orbit

� Maximum Number of Electrons in

orbit

� Period of the element in Periodic

Table

� It was introduced by Bohr.

� It has all natural no. values.

� Its values are also designated by

alphabets K, L, M. N etc.

n=1 1st energy

level

K shell

n=2 2nd

energy

level

L shell

n=3 3rd

energy

level

M shell

� No. of electrons in an orbit

could be calculated by the

formula 2n2

.

Azimuthal/

Orbital/

Angular

Momentum/

Subsidiary/

l

l = zero

> (n-

1)

0 1 2 3 4 � Splitting of energy levels

� Shapes of orbitals

� Magnitude of angular momentum

of orbiting electron.

� It was introduced by

Sommerfeld.

s p d f g

XXX

XXX

Magnetic

m

M = -l

> 0

>

+l

Max. Values

(2l + 1)

� Orientation (axis of electrons)

� Sub-sub levels inside sub-levels

called orbitals

� It has only whole number

values.

� Maximum values is equals to (2l

+ 1) e.g. l=3, m can have -3, -2,

-1, 0, 1, 2, 3 (total seven values).

Spin

ms or s

Spin up & spin

down

+½ or -½

� Spinning direction of electron

� Rotation of electrons besides

revolution.

+½ ↑ Clock-wise Spin up

Spin Down - ½ ↓ Anti-clockwise




KARACHI

4

ELECTRONIC CONFIGURATION: “The arrangement of electrons in an atom is referred as Electronic Configuration”

Electronic configuration of elements is the fundamental property by which they are

placed in the periodic table and exhibit periodic properties.

Electrons are arranged in different energy levels (K, L, M, N) and sub levels (s, p, d, f) according

to their energy state.

n 1 2 3 4

Energy Level K L M N

Sub Level s(2)

s(2)

, p(6)

s(2)

, p(6)

, d(10)

s(2)

, p(6)

, d(10)

, f(14)

Maximum

Electron

2 8 18 32

Pauli Exclusion Principle: “No two electrons in the same atom can have all the set of

four quantum numbers identical” Or “Only two electrons may occupy the same atomic

orbital and these must have opposite spins.”

Auf Bau or Building-Up Principle: “For any given atom, the electrons are filled to the

orbitals of lowest energy in sequence, two electrons to each orbital”

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

(n+l) Rule: “In building up the electronic configuration of the elements, the orbitals with the

lowest value of (n+l) fills first; when two orbitals have the same value of (n+l), the orbital with

the lower value of n fills first.”

Hund’s Rule: “When several orbitals of the same type are available, the electrons first fill all

the orbitals of same energy singly with parallel spins, before pairing in any one orbital takes

place.”

CHEMICA


CHEMICAL BOND: “The attr

known as Chemical Bond.”

A chemical bond:

� Holds two identical or differen

uncombined atoms.

� Is formed with the creation of

� Is formed due to the interaction

� Makes atoms attain greatest sta

(Octet Rule).

IONIC OR ELECTRO

“The chemical bond formed as a

“The electrostatic

In an ionic bond:

� If the difference of Electronega

greater than 1.7 between two e

bond formed is ionic.

� Only a low electronegative ele

held along with a high electron

(non-metal) and vice versa.

� The atom that loses electrons b

charged (called Cations) and th

electrons becomes negatively c

Anions).

1 A molecule is defined as the smallest pa

original substance.

Primary Bonds

Ionic Bond Covale

Polar Covalent Bond

�onCovalen

ICAL BONDING PREPARED BY



e attractive force that holds atoms together in a com

ifferent atoms as a molecule1, which possesses lower e

on of attractive forces which result in the release of en

raction of valence electrons of the combining atoms.

test stability when they acquire an inert gas electronic

ROVALENT BOND:

as a result of complete transference of valence elec

atom to the other.”

OR

tatic attraction between positive and negative ions

ronegativity ∆E is

two elements, then the

ive element (metal) is

lectronegative element

trons becomes positively

and the atom that gains

ively charged (called

llest particle of an element or a compound that retains the chara

Chemical Bond

onds

ovalent Bond

on-Polar ovalent Bond

Coordinate Covalent Bond

Secondary Bonds (Van der Waal forces)

Hydrogen Bond

Dipole-Dipole Attraction

ANDOUT 7

NDARY SCHOOL,

KARACHI.

1

compound is

ower energy than the

e of energy.

oms.

tronic configuration

electrons from one

ons”

characteristics of the

ds orces)

London Forces

CHEMICAL BONDING PREPARED BY



KARACHI.

2

� The charged species is called ions, which do not

retain the properties of the original atoms.

� The ions do not form an individual molecule in

the liquid or solid phase, but are arranged into a

crystal lattice or giant ion-molecule containing

many of these ions.

� The attractive forces in the crystal lattice greatly

reduce the energy of the system, making it stable.

� Formation of an ionic bond occurs in a number of

steps, details of these steps are as follows:

PATHWAY FOR IONIC BOND FORMATION:

� Metal (s) ───> Metal (g) + ∆H1 Enthalpy of Sublimation

� Gas Molecule ───> Gas atom + ∆H2 Bond Dissociation Energy

� Metal (g) ───> Metal Ion + ∆H3 First Ionization Potential

� Gas Atom ───> Gas Ion+ ∆H4 Electron Affinity

� Metal Ion + Gas Ion ───> Ionic Compound+ ∆H5 Lattice Energy2

CONDITION FOR IONIC BOND FORMATION: For solubility, solvation energy3 must be

greater than the lattice energy. (∆H5+ ∆H4) > (∆H3+ ∆H2+ ∆H1)

CHARACTERISTICS OF IONIC COMPOUNDS:

� Each ion is surrounded by a fixed number of oppositely charged ions, so that the strong

electrostatic forces between ions act in all directions through the crystal.

� Ionic compound are soluble in water and similar polar solvents because of the strong

electrostatic attractions between the ions and polar molecules of the solvent.

� Ionic compounds are insoluble in the organic solvents because there’s no attraction between

the ions of ionic compound and the molecules of non-polar solvents.

� In the presence of a polar solvent, such as water,

the inter-ionic forces are so weakened that the ions

are separated and the free ions are able to move

under the influence of electric current i.e. ionic

compounds are electrolytes in molten and aqueous

state.

� Ionic compounds have low volatilities and low

vapor pressure or they do not vaporize readily at

room temperature.

� Ionic compounds are good electricity conductors in

molten state but their conductivity is smaller than

that of metallic substances.

2 “The energy released when one mole of gaseous ions arrange themselves in definite pattern from crystal lattice is

referred to as the lattice energy.” 3 “The surrounding of the ions by the solvent molecules for solvation releases the energy which is known as the

solvation energy.”




KARACHI.

3

COVALENT BONDING:

“The chemical bond in which two atoms, who tend to gain electrons may combine with each

other by sharing one or more pairs of electrons.”

In a covalent bond:

� If the electronegativity difference between two or more atoms

is zero or not greater than 0.5, then a Non-Polar Covalent

Bond is formed.

� If the electronegativity difference is between 0.5 and 1.7,

there will not be an equal sharing of electrons between the

involved atoms. Then a Polar Covalent Bond is formed.

� The electrons that are shared between the two atoms are the

bonding electrons. � Bonding electrons spend much of their time

between the nuclei, resulting in the attractive

forces between negative charges of electrons and

positive charges of the two nuclei.

� In a polar covalent bonded molecule, the shared

pair of electrons is attracted more towards more

electronegative atom, i.e. a polar bond makes one

part of molecules partially negative(δ-

) and the

other partially positive (δ+

).

� Due to polar bonding, the molecule is observed to

enjoy greater stability than expected.

� In polar molecules, electron cloud shifts towards

the more electronegative atom.

Characteristics of a Covalent Compound:

� Exist as separate covalent molecules, electrically neutral and have little attractive forces for

each other called Van der Waal Forces.

� Mostly low volatile liquids or gases or low melting solids.

� Non-electrolytes

� Insoluble in water and similar polar solvents but soluble in organic solvents.

COORDINATE OR DATIVE OR CO-IONIC COVALENT BOND:

“The covalent bond in which the bonding electron pair is provided by a single atom”

� The group VA, VIA or VIIA form such type of bond.

� Once formed, a co-ordinate covalent bond becomes indistinguishable from a covalent bond.

CHEMICA


MULTIPLE BONDS:

In multiple covalent bonds

much stronger than the single cova

making molecules much more stab

� Bond energy is higher in multi

� Bond lengths in multiple coval

� Multiple bonded compounds a

� Bond stability ∝ Bond energy

� Bond stability ∝ Electron Den

� Bond stability ∝ 1 Bond length

METALLIC BOND: Atoms in metallic crystals

“The strong force of attraction

The valence electrons in a metal ar

strongly because metals have smal

valence electrons. The forces in m

are electrostatic in origin. They ca

released to the common pool due t

binding force. The atomic orbitals

valence electrons are assumed to m

produce a delocalized orbital syste

throughout, on which the electrons

move.

Characteristics of a Meta

� Ability to act between identica

different metallic atoms.

� Lack of direction, as shown by

saturation to permit large numb

� Attractive force varies inversel

� Ability to permit electron trans

� Equilibrium repulsive force, W




bonds, the attraction force is

le covalent bonded molecules

re stable. Due to this fact:

multiple covalent bonds and

covalent bonds are shorter

unds are much stable than single bonded compounds.

nergy

n Density

length

ystals are bonded together by electron cloud.

ction between these differently charged particles fo

Bond.”

etal are not held

e small number of

s in metallic bonds

hey can be easily

l due to weak

rbitals, holding the

ed to merge to

l system

ctrons are free to

etallic Bond:

entical metallic atoms, and at the same time between

wn by the positive retention of properties in the liquid

e number of close neighbors.

versely as some high power of inter atomic distance.

n transfer from one atom to another atom.

Which is atomic in nature.

ANDOUT 7

NDARY SCHOOL,

KARACHI.

4

unds.

es forms Metallic

ween widely

liquid state, and of

tance.

CHEMICA


PROTONIC BRIDGE

“An electrostatic attraction betw

one molecule attracts the negativ

which involves H

Hydrogen bonding is the st

normal covalent bond.

Hydrogen bonding:

� Affects the physical properties

� Increases the melting and boilin

compound.

� Reduces the vapor pressure.

� Increases the heat of vaporizatio

Water has higher boiling p

fluoride. The oxygen atoms have t

pairs of electrons and there are two

atoms (Hδ+

) present in three dimen

Conditions:

� Hydrogen should be bonded to

electronegative element such a

� Hydrogen bonding ∝ ionic cha

� Hydrogen bonding ∝ ∆EN

DIPOLE MOMENT (µ

“A dipole (polar molecule) tend

tendency

Where,

e = magnitude of charge

d = distance between the charges.

The S.I Unit is coulomb meter (Cm

� Dipole moment ∝ ∆EN

� Dipole moment ∝ geometr

�

NATUR

� Diatomic non-

� Diatomic Polar

� Tri-atomic Line

� Tri-atomic Ang

� Tetra Atomic

� Penta Atomic




GE OR HYDROGEN BONDING:

between neighboring molecules is setup when the p

gative pole of neighboring molecule. This type of a

ves Hydrogen is referred as Hydrogen Bonding.”

the strongest of the secondary bonds but is still weak

erties of the compound.

boiling point of the

rization.

iling point than hydrogen

have two non-bonded

are two polar hydrogen

dimensional bonding.

ded to a highly

such as N, O, F etc

ic character

µ):

tends to become oriented in an electrical field. The

ency is referred to as the dipole moment.”

µ = e x d

arges.

ter (Cm)

ometry of molecule

URE DIPOLE MOMENT EXAMPLE

-polar Zero H2, Cl2, Br

Polar Greater than Zero HF, HCl

c Linear Zero CO2, CS

c Angular Greater than Zero H2O, H2S, Ca

Appreciable NH3

Zero CH4, CCl

ANDOUT 7

NDARY SCHOOL,

KARACHI.

5

the positive pole of

of attractive force

”

l weaker than a

The extent of this

PLES

, Br2

, HCl

, CS2

S, CaC2

, CCl4




KARACHI.

6

BOND ENERGY:

“The energy required to break a bond between two atoms in a diatomic molecules is known as

the Bond Energy.”

H – H ───> 2 H ∆H = 435 KJ/moles

Hydrogen Molecule Hydrogen Atom

O = O ───> 2O ∆H = 435 KJ/moles

Oxygen Molecule Oxygen Atom

OR

“The energy released in forming a bond from the free atoms (not from the atoms in their

standard states.)”

H (g) + H (g) ───> H - H ∆H = - 435 KJ/moles

O (g) + O (g) ───> O = O ∆H = - -498 KJ/moles

Reminder: Breaking of Bonds is endothermic and making of bonds is exothermic.

� Bond Energy ∝ Type of Bond

� Bond Energy ∝ Ionic Character

� Bond Energy ∝ Stability

� Bond Energy ∝ 1 Bond Distance

Bond Energy Trend:

Ionic Bond > Polar Covalent Bond > �on-Polar Covalent Bond

BONDING ORBITALS:

“A covalent bond is formed as a result of overlapping of atomic orbitals of combining atoms

giving molecular orbitals.”

The overlapping of atomic orbitals is of two types:

� Sigma(σ) Bonding � Pi(π) Bonding

Sigma Bond:

“A bond which is formed by the axial overlap of two orbitals belonging to different atoms”

It is formed by the axial overlapping of s-s, p-p, s-p orbitals.

Pi Bond:

“A bond which is formed by side ways (parallel) overlapping of atomic orbitals”




KARACHI.

7

Two different theories determine the behavior of electrons in a bonded atom. It is as follows

Valence Bond

Theory

Suggests

The bonding electrons occupy the atomic orbitals4 of the

bonded atoms.

Molecular

Orbital Theory

The bonding electrons occupy molecular orbitals5 which

belong to the whole molecule.

According to molecular orbital theory, linear combination of atomic orbitals (σ) gives

two molecular orbitals:

� σ-Bonding: A molecular orbital with high electron density in the region between two

nuclei having lower energy than either of the parent atomic orbitals from which molecular

orbital is derived.

� σ-Anti Bonding: The molecular orbital having high energy and low electron density

between the nuclei, is less stable than either of the parent atomic orbitals from which it is

derived.

HYBRIDIZATION:

“The mixing of various atomic orbitals to produce the same number of equivalent orbitals

having same shape and energy”

Types of Hybridization:

Mixing Orbitals Type of

Hybridization

Type of Hybrid

Orbitals

Number of

Hybrid Orbitals

One-s & three-p sp3 sp

3 orbitals Four

One-s & two-p sp2 sp

2 orbitals Three

One-s & one-p sp sp orbitals Two

sp3 (Tetrahedral) Hybridization:

� Possesses the character of both s- and p- orbitals in the ration of 1:3.

� sp3 orbitals are directed towards the four corners of a regular tetrahedron in which each angle

is 109°.28’.

� Forms the strongest bonds among hybrid orbitals.

� Examples: CH4, SiCL4, SnCl4 etc

sp2 (Trigonal) Hybridization:

� Possess one-third character of s-orbitals and two third characters of p-orbitals.

4 An orbital in which, the electron is influenced by one nucleus (mono-centric)

5 An orbital in which, the electron is influenced by more than one nucleus (poly-centric)




KARACHI.

8

� These orbitals are co-planar and directed towards the corners of an equilateral triangle

(trigonal) at angles of 120° to each other.

� Examples: C2H4, C2H6 etc

sp (Diagonal) Hybridization:

� Possesses the character of both s- and p- orbitals in the ration of 1:1.

� These orbitals are co-linear at an angle of 180°.

� Examples: C2H2, BeCl2 etc

Hybrid Orbital Model:

Hybridization Lone Pairs Bond Angle

sp Zero 180 °

sp2 Zero 120°

sp3 Zero 109°.28’

sp3 One 107°

sp3 Two 104.5°

MOLECULAR GEOMETRY: � The geometrical shape of a molecule depends upon the repulsion factor among the electrons

pair of an atom.

� There are two types of electron pairs surrounding the central atom:

Bond Pair: These are the result of sharing of unpaired electrons of central atom with unpaired

electrons of surrounding atoms.

Lone Pair: These are the paired electrons, which have not taken part in sharing. They are also

called “Non Bonding Pairs”.

The repulsion factor of these active sets of electrons behaves as follows:

Lone Pair – Lone Pair repulsion > Lone pair – Bond Pair repulsion > Bond – Bond Pair repulsion

� The shape of an atom could be predicted by the number of electrons groups6 in a molecule. A

general observation is as follows:

� 2 Electron Groups: Linear

� 3 Electron Groups: Trigonal Planar

� 4 Electron Groups: Tetrahedral

� 5 Electron Groups: Trigonal Bipyramidal

� 6 Electron Groups: Octahedral

6 An electron group is any collection of valence electrons localized in a region around a central atom that exerts

repulsion on other groups of valence electrons

THERMO


Thermochemistry: The bran

evolved or absorbed during a chem

Thermodynamics: The study

THERMOCHEMICALaccompanied by energy changes w

EXOTHERMIC REACTI

� The chemical reactions accomp

emission of heat energy with th

of product.

� The chemical reactions which

energy released as compared to

energy.

� The chemical reactions in whic

energy is converted to thermal

� Reactants ───> Products + H

H1 > H2

� ∆H = H2-H1 = [a negative valu

� Examples: Combustion, Oxida

Neutralization etc

En

tha

lpy H

──

─>

Exothermic Reac

Reactants

∆H < 0 [negative

Extent of Reactio

Exothermic

MOCHEMISTRY PREPARED BY



branch of chemistry which deals with the measurem

a chemical reaction.

e study based on the principles of conservation of ener

AL REACTIONS: The chemical reactions w

nges with the material changes.

CTIONS: ENDOTHERMIC REAC

ccompanied by the

with the formation

hich have greater

ared to activation

n which chemical

ermal energy.

� The chemical reactions acco

absorption of heat during the

product.

� The chemical reactions whic

activation energy than the en

� The chemical reactions in wh

energy is converted to chemi

ts + Heat

� Reactants + Heat ───> Pro

H1 < H2

value] � ∆H = H2-H1 = [a positive va

Oxidation, � Examples: Reduction, Electr

eaction

tive]

Products

tion───>

En

tha

lpy H

──

─>

Endothermic Re

∆H > 0 [positiv

Reactants

Extent of Reacti

Thermochemical Reactions

ermic Reaction Endothermic Reaction

HANDOUT 8

NDARY SCHOOL,

KARACHI.

1

surement of heat

f energy.

ions which are

EACTIONS:

s accompanied by the

ing the formation of

s which have greater

the energy released.

s in which thermal

chemical energy.

Product

value]

Electrolysis etc

Reaction

Products

sitive]

action───>

ction

THERMOCHEMISTRY PREPARED BY



KARACHI.

2

THERMODYNAMIC PROCESSES:

� Isothermal Process: “The process in which temperature remains constant.” i.e. dT = 0

� Isochoric Process: “The process in which volume remains constant.” i.e. dV = 0

� Isobaric Process: “The process in which pressure remains constant.” i.e. dP = 0

� Adiabatic Process: “The process in which heat remains constant.” i.e. dH = 0

FIRST LAW OF THERMODYNAMICS: This law was enunciated by

Helmholtz.

“Energy can neither be created nor be destroyed, although I may change from one form to

another” OR

“The total energy of a system and its surrounding must remain constant”.

Mathematically, q = (E2-E1) + W → q = ∆E + W → ∆E = q – W→ W = q – ∆E

Facts to remember:

� Heat is supplied to the system: q = [positive] � Work is done by the system: W = [positive]

� Heat is evolved by the system: q = [negative] � Work is done on the system: W = [negative]

Units of Heat Energy:

1 Calorie = 4.2 J 1 J = 0.239 Calorie 1 BTU = 1055 J 1 J = 107 erg

HESS’S LAW OF CONSTANT HEAT SUMMATION: “The heat evolved or absorbed in a given reaction must be independent of the particular manner

in which the reaction takes place”

Let’s consider, for a certain reaction:

Pathway 1: A ──────> D + ∆H

Pathway 2: A ──────> B + ∆H1

B ──────> C + ∆H2

C ──────> D + ∆H3

According to the law, ∆H = ∆H1 + ∆H2 + ∆H3

MEASUREMENT OF HEAT:

Heat Capacity: “The quantity of heat required to change the temperature of system by 1°C

or 1K.”

� = �

ΔT

Q = C ∆T

Molar Heat Capacity: “The heat capacity of one mole of a substance.”

Specific Heat: “The heat capacity of a one-gram sample.”

�� = ��

��=

�

�




KARACHI.

3

�� =�

� × ��

TERMINOLOGY:

� System: “Any specified real or imaginary part of the universe that is under consideration.”

Types of System:

� Open System: “The system that interacts readily with the surroundings, exchanging

matter and energy.”

� Closed System: “The system that exchanges energy but not matter with

surroundings.”

� Isolated System: “The system that exchanges neither matter nor energy with the

surroundings.”

� Homogenous System: “The system which is completely uniform throughout i.e.

consisting only one phase.”

� Heterogeneous System: “The system which is not uniform throughout i.e.

consisting of two or more phases.”

� Surroundings: “The environment of the system or rest of the universe which may act on

the system is known as Surrounding.”

� Interactions: “The exchange of energy and/or matter between a system and

surroundings.”

� Heat of Reaction: “The quantity of heat exchanged between a system and its

surroundings for a reaction at constant temperature.”

� State: “A system is said to be in a certain state when all its properties are fixed.”

� State Function: “The macroscopic properties of a system which has some definite value

for each state and independent of path in which the state is reached.”

Example: Enthalpy, Temperature, etc

� Enthalpy: “The sum of internal energy (E) and the pressure-volume product of a system”

� Enthalpy is an extensive property.

� Enthalpy is a state function.

� Enthalpy changes have unique values.

� Entropy: “The measure of the disorder of the system.”

� Disorder of system increases: d (entropy) = [positive]

� Disorder of system decreases: d (entropy) = [negative]

� Spontaneous changes always occur with an increase of entropy of the universe.

� Entropy of a system as a whole is an irreversible process.

� Entropy is expressed in cal/°C or J/K (SI Unit).

� Gibb’s Free Energy (G): The criterion for spontaneity

� ∆G = [negative] → Reaction is spontaneous and feasible




KARACHI.

4

� ∆G = [positive] → Reaction is impossible

� ∆G = 0 → Reaction is in equilibrium

� Internal Energy: “Sum of energies possessed by a system i.e. translational, vibrational,

rotational, chemical, bond energy, electronic energy, nuclear energy of constituent atoms and

potential energy.”

E = Et + Er + Ee + En + EP.E + Ey

� Macroscopic Properties: “The properties of s system easily measurable for the

entire bulk of the molecules.”

It is divided into following two groups:

� Intensive Properties: “The properties of a system those are independent of the

amount of material concerned.”

Examples: Density, Pressure, Temperature, Refractive index, Viscosity, Surface Tension,

Melting Point, Boiling Point etc

� Extensive Properties: “The properties of a system those are dependent of the amount

of material concerned.”

Examples: Mass, Volume, Mole, Enthalpy, Entropy, Internal Energy, Gibb’s Free Energy

etc

� Heat of Formation: “The change in enthalpy, when 1 mole of a substance is produced

from its elements in the natural state.”

HYDROGEN AND WATER PREPARED BY



KARACHI.

1

HYDROGEN GAS: � Hydrogen atom is the simplest of all atoms with 1 proton, 1 electron and zero neutrons. The

element has following specifications:

I�VESTIGATIO�S OBSERVATIO�S

1. State Gas (colorless, odourless) 2. Electronic configuration 1s1 3. Atomic Radius 0.37 A°Å 4. Ionic Radius (H-) 1.54 5. Ionization Potential (1st) 1312 KJ mol-1 6. Electron Affinity -37 KJ mol-1 7. Electronegativity 2.1 8. Liquefying Point 20.39 K 9. Solidifying Point 13.98 K

ISOTOPES OF HYDROGEN: NAMES FORMULA PROTON ELECTRON NEUTRON ATOMIC MASS %AGE � Protium 1

1H One One Zero 1 99.98 %

� Deuterium 12H or 1

2D One One One 2 0.02 %

� Tritium* 13H One One Two 3 4x10-15 %

*Tritium is a radioactive isotope with the half life of 12.4 years.

POSITION IN PERIODIC TABLE: Scientists are still unable to predict the most appropriate position for hydrogen in the

Periodic Table. This is mostly because of the fact that hydrogen is the simplest element of all and resembles more than one families of the periodic table. A summarized account of its placement in different groups is as follows:

Group I-A(Alkali Metals) Group IV-A(Carbon Family) Group VII-A (Halogens)

Similarities Dissimilarities Similarities Dissimilarities Similarities Dissimilarities

Outermost

Electrons

Orbit

Completion

Half Filled

Valence Shells

Valency Orbit

Completion

Outermost

Electrons

Electronic

Configuration

State Thermodynamic

Properties

Valence Orbit Valency Valence Orbit

Oxidation

State

Ion

Formation

Chemical Bond

Formation

State Diatomic

Molecule

Oxidation

State

Reducing

�ature

Bond

Formation

Reducing

�ature

�ature Oxide

Formation

OCCURRENCE: � By mass, hydrogen makes up only 0.9% of Earth’s crust. � In number of atoms in Earth’s crust, hydrogen ranks third (15.1%) after Oxygen

(53.3%) and Silicon (15.9%). � Hydrogen makes up 89% atoms on the sun and 85 to 95% of the atoms in the atmosphere of

outer planets (Jupiter, Uranus, Saturn and Neptune). � In universe, as a whole, 90% of the atoms are Hydrogen, and the rest are mainly Helium.

HYDROGEN AND WATER

NSTC-6 PREPARATION PROGRAMME

PREPARATION:

ATOMIC HYDROGEN: “The product obtained as the result of dissociation of molecular hydrogen.”

Preparation:

H

Properties:

� It’s more reactive than molecular hydrogen.

BINARY COMPOUNDS OF

hydrogen with metals and non

Ionic Hydride

Covalent Hydride

� By the action of metal on hydrides

� By electrolysis of water

� By Steam and Hydrocarbon Process

� By Action of Steam on Coke

� By Steam Methanol Process

� Thermal Decomposition of Hydrocarbons

� Catalytic Reforming of Alkane Hydrocarbons

� Thermal Decomposition of Ammonia


SYNOPSIS HANDOUT OSAMA HASAN

ROGRAMME | AGA KHAN HIGHER SECONDARY

“The product obtained as the result of dissociation of molecular hydrogen.”

H2 + 104 K. cal

��°�> 2 H

H2 ��

��.� �� > 2 H

than molecular hydrogen. Some of its reactions are as follows:P + 3 H ──────> PH3

O2 + 2 H ──────> H2O2

S + 2 H ──────> H2S

Cl2 + 2H ──────> 2 HCl

CuO + 2 H ──────> Cu + H2O

AgCl + H ──────> Ag + HCl

OMPOUNDS OF HYDROGEN:

"The bi-element compounds of hydrogen with metals and non-metals

are termed as hydrides."

Complex Hydride

Metallic Hydride

Polymeric Hydride

By the action of metal on hydrides : Zn + 2 H+

> Zn

: 2 H2O ��

�.� ��/ ��> 2 H

By Steam and Hydrocarbon Process : CH4(g) + H2O(g) ��

��°�> 3 H

: C + H2O(g)

��°�> 3 H

: CH3OH(g) + H2O(g)

��°�> CO

: CH4

��°�> C + 2 H

Catalytic Reforming of Alkane : C6H14 ��

> C6

Thermal Decomposition of Ammonia : 2 �H3 ��

��°�> �2

SYNOPSIS HANDOUT 13

ECONDARY SCHOOL,

KARACHI.

2

“The product obtained as the result of dissociation of molecular hydrogen.”

reactions are as follows:

Borderline Hydride

> Zn2+

+ H2

> 2 H2 + O2

> 3 H2 + CO

> 3 H2 + CO

> CO2 + 3 H2

> C + 2 H2

6H6 + 4 H2

2 + 3 H2



NSTC-6 PREPARATION PROGRAMME | AGA KHAN HIGHER SECONDARY SCHOOL, KARACHI.

3

Ionic/Saline

Hydrides

Covalent

Hydrides

Complex

Hydrides

Metallic

Hydrides

Polymeric

Hydrides

Borderline

Hydrides

Combination s-block metals except Be and Mg

p-block elements B and Al with alkali metals

Transition elements

Beryllium & Magnesium

I-B, II-B & III-A (In, Tl)

Bonding Ionic Bond Covalent Bond Complex Bonding --------- ???? ????

General

Formula

M+H- & M2+H-2 --------- A+B3+H4 nonstoichiometric M2+H-

2 ????

Preparation By passing H2 over metals at high temperature

Direct and indirect methods.

Alkali metal hydrides + III-A

hydrides

By heating H2 with metal under high pressure

???? ????

Properties � Colorless solids � Non-volatile � High M.P � Insoluble in organic solvents

� Soluble in water � Conduct electricity in molten state

� Colorless gases � Volatile liquids � Low B.P � III-A and IV-A are neutral

� V-A are basic � VI-A and VII-A are acidic.

� Salt like white solids

� Stable upto 300°C

� Soluble in water

� Hard Solids with metallic luster

� Electricity conductors

� Possess magnetic prop.

� Solid, Volatile, white in color

� Properties are intermediate between ionic and covalent hydrides.

� Properties are intermediate between those of metallic and covalent hydrides.

Reactions XH + H2O ──>

XOH + H2

XH + HCl ──>

XCl + H2

XH+C2H5OH──>

C2H5X + H2

???? ???? ???? ???? ????

Examples NaH, LiH, KH, RbH, CsH, CaH2, SrH2 and BaH2.

B2H6, CH4, NH3, H2O, HF, HCl, HI, HBr, H2S

LiAlH4, NaBH4 etc

???? BeH2 and MgH2.

????

Uses � Reducing Agents � Dehydrating Agents

� Source of H2

???? � Reducing Agent � Catalyst

� Reducing agents � Source of Atomic Hydrogen

??? ????




KARACHI.

4

USES OF HYDROGEN: Hydrogen is used: � In the manufacture of NH3, in turn, is used in the manufacture of fertilizers, plastics and

explosives. � To convert Benzene to Cyclohexane, a cyclic hydrocarbon used as an intermediate in the

production of nylon.

C6H6 + 3 H2 ��

> C6H12

� In the synthesis of methyl alcohol (methanol), an industrial solvent and raw material for making other organic compounds.

� To extract pure metals from metal oxides, i.e as Reducing Agent.

WO3 + 3 H2 �� °�

> W + 3 H2O

� Liquid Hydrogen is used a rocket fuel. The fuel tank holds 1.5 x 106 L of liquid hydrogen which powers the shuttle for about 8.5 min at the rate of consumption of 3000 L/s.

� In developing oxyhydrogen welding torch that readily cuts through steel and can be used to melt tungsten (W), which has a melting point of 3400°C.

WATER:

HARD WATER AND WATER SOFTENING:

HARD WATER: The ground water that contains significant concentrations of ions from natural sources, principally Ca2+, Mg2+ and sometimes Fe2+ along with associated anions.

� Temporary Hard Water: The hard water with HCO3- as primary anion.

� Permanent Hard Water: The hard water with Cl- and SO42-as primary anions.

WATER SOFTENING:

� Temporary Hard Water: Temporary hard water can be softened by boiling. A better way is to treat the water with a base, usually Ca(OH)2 that converts HCO3

- to CO32-.

� Permanent Hard Water: Permanent hard water cannot be softened by boiling. Addition of washing soda (Na2CO3) softens permanent hard water. But soluble salts, such as NaCl and Na2SO4 remain in solution.

HEAVY WATER: The water molecule having the Deuterium isotope ‘D’ attached with Oxygen atom in place of Hydrogen (Protium) ‘H’ The formula for Heavy Water is D2O and its molecular mass is 20 a.m.u.

Uses:

� Heavy water is used in nuclear reactors as moderators which absorb the neutrons and slow down the process.

� Heavy water is used in analytical experiments for determining the mechanism of reactions.

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