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Ch i t IGCSE Ed l 2011 2012

Principles of chemistry

States of matter

State Solid Liquid Gas

Arrangement of

particles

Regular

arrangement

(fixed pattern)

Random

arrangement

(no fixed pattern)

Random

arrangement

Proximity Closely packed Still close together

but not as close as

in solid

Far apart

movement Vibrate in their

positions

Slide past each

other

Move everywhere

rapidly

Shape Definite Indefinite Indefinite

Kinetic particle theory

1- All matter is made of tiny invisible particles (atoms, ions or

molecules)

2- Particles move randomly all the time

3- Lighter particles move faster than heavier ones

Changing of state

Sublimation

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Melting a solid

When a solid is heated, particles gain energy, they vibrate faster and the

solid expand, at the melting point the particles vibrate more enough to

overcome the forces between them and the solid changes into liquid.

Boiling a liquid

When a liquid is heated, particles gain more energy; they move faster, this

makes the liquid expand. At the boiling point, the particles get enough energy

to overcome the forces between them and escape in the form of gas.

Evaporation:

When a liquid is left open to the air, some particles of the liquid escape into

the gas state even if the liquid is below the boiling point.

And the rate of evaporation increases as

The temperature increase.

The surface area increase.

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Note

Evaporation occurs at any temperature, but boiling occurs at a certain

temperature which is the boiling point.

Condensing a gas

If a gas is cooled, particles eventually move slowly enough that attractions

between them hold them as a liquid. The gas condenses

Freezing a liquid

If a liquid is cooled, liquid particles will move around more and more slowly.

Eventually, they are moving slowly enough that the forces of attraction

between them will hold them into a solid, the liquid freezes.

Changing between solid and gas [Sublimation]

Small number of substances changes directly from solid to gas without passing

through liquid state.

Examples: Ammonium chloride, Carbon dioxide and iodine

Heating ammonium chloride Dry ice (solid carbon dioxide) subliming Iodine sublimation

Melting point: It is the temperature at which a solid melts, i.e. changes into a liquid

Boiling point: It is the temperature at which a liquid boils, i.e. changes into a gas

Freezing point: It is the temperature at which a liquid freezes, i.e. Changes into

a solid

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Diffusion

It is the random movement of particles so they get mixed up

Or : It is the spreading of particles from regions of high concentration to

regions of lower concentration.

Diffusion in gases

a) Diffusion of bromine

Bromine is a brown liquid which vaporizes

easily at room temperature.

The lower gas jar contains bromine vapour;

the top one contains air. When the lid

between the two jars is removed, bromine

vapor spread to mix with air in the upper jar

b)Diffusion of ammonia and hydrogen chloride gases to form ammonium chloride

A cotton wool soaked in concentrated ammonia solution (as a source of

ammonia gas) is placed at one end of a long glass tube.

A cotton wool soaked in concentrated hydrochloric acid (as a source of

hydrogen chloride gas) is placed at the other end of the tube

When ammonia gas (NH3) is mixed with hydrogen chloride gas (HCl),

white fumes of ammonium chloride (NH4Cl) are formed.

NH3(g) +HCl(g) NH4Cl(s)

ammonia hydrogen White fumesgas chloride gas of ammonium chloride

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The white fumes of ammonium chloride, NH4Cl(s), appear

nearby the source of hydrogen chloride gas since this HCl(g) is

denser and diffuses slower than ammonia gas NH3(g).

Rate of diffusion of gases depends on the molecular mass (Mr),

the smaller the molecular mass the faster the rate of diffusion.

c) If a purple crystal of potassium manganate (VII) is dropped into a

beaker containing water, the purple color will spread throughout

the water

Potassium manganate (VII) crystal inwater

The crystal sinksThe crystal dissolves and becomes smaller

Particles of potassium ion and manganate ionsmove randomly

Color spread everywhere

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Chemistry IGCSE Edexcel 2011 - 2012

Laboratory Glassware

Beaker Conical flask Measuring

cylinder

Pipette Burette Stand Clamp

Test tube Test tube holder Mortar and

pestle

Tripod Gauze Bunsen burner Tongs

Evaporating

dish

Funnel Crucible and lid

Thermometer Separating

funnel

Spatula Top pan

Balance

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Filtration Crystallization Simple distillation Fractional distillation

Used to separate an

isoluble solid from a

solution

Used to separate a solid

crystals from a solution

To separate a solvent from a

solution

To separate two or more miscible

liquids of different boiling points.

Example: sand and

water

Example: Magnesium

sulphate from magnesium

sulphate solution

Example: distilled water from sea

water

Example:Ethanol and water

Solid particles are left

on the filter paper asa residue while the

liquid passes through

the filter paper as a

filterate.

The solution is heated to

crystallization point, thenleft to cool to room

temperature, crystals will

form and can be filtered out

and left to dry in a warm

place or dried between twofilter papers

When the solution is heated, the

solvent changes to vapour, thevapour passes through a condenser

where it is converted back to liquid

and is collected in a conical flask as

the distillate

Mixture of miscible liquids is heated,

the liquid with lower boiling pointreaches the top of the column and

distils over and is collected first

Fractionating column is a glass

colunmn packed with glass beadsallows multiple condensations and

distillations and produces betterseparation between liquids

Methods of separation and purification

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Paper chromatography

Used to separate and identify a mixture of substances present in a

concentrated solution

Example:

1- Separating dyes in ink

2- Colours and flavors in food

A small concentrated spot of the solution containing the mixture

is placed on a base line drawn by a pencil at one end of the

chromatography paper

The paper is dipped in a suitable solvent (e.g. water or ethanol )

taking care that the solvent surface is below the base line and

placed in a sealed container

The solvent gradually moves up the paper.

As the solvent rises through the paper it meets and dissolves the sample

mixture, which will then travel up the paper with the solvent.

Different components of the mixture travel at different rates and

the mixtures are separated into different coloured spots.

The pattern you get is called a chromatogram

The number of spots represents the number of constituents of

the mixture

A single pure substance will produce only one spot

If the constituents of the mixture is insoluble in the solvent

used, it will remain on the baseline Chromatography can be used also to separate colorless

substances, but in this case, the paper must be sprayed by

a locating agent, so that the position of the spots can be

seen.

Each spot could be identified from its R f ratio. This

ratio is calculated from the following formula:-

R f =solventbytraveleddistance

dyethebytraveleddistance

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Example:

The mixture (m) consist of the dyes d1,d3 & d4 because:

They have same color as spots in the mixture

They travelled same distance on the paper

The dye d2 has the same color as one of the spots in the mixture, yet it has

travelled a different distance and so it must be a different compound

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Atom and molecule

Atom: Smallest particle of an element which takes part in a chemicalreaction

Structure of atoms:

1. Nucleus: Contains positively charged protons and neutralneutrons

2. Electrons are negatively charged and moves around the nucleus

in energy levels (energy shells)

Atoms are neutral because the number of positively chargedprotons (p+) are equal to the number of negatively charged

electrons (e-) Mass of electron is very small and can be neglected compared to

mass of nucleus

Relative charge and mass of particles

Particle Symbol Relativecharge

Relative mass

Proton p +1 1Neutron n 0 1Electron e- -1 1/1836

Atomic number and mass number (nucleon number)

Atomic number = number of protons

Mass number = number of protons + number of neutrons

Number of neutrons = Mass number – Atomic number

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Electronic configuration

Electrons are arranged in energy levels or shells around the

nucleus

Each energy level can hold a certain number of electrons

Lower energy levels are filled first before you go to higher ones

The first energy level holds only 2 electrons, the second 8

electrons, the third appears full with 8 electrons but can expand

to a total of 18

Outer shell is called the valence shell and the electrons of the

outer shell are called valence electrons

Outermost shell (Valence shell) should not contain more than 8

electrons

Electronic configuration of the first 20 elements in the periodictable

Elements of the periodic table are arranged according to their

atomic number

Elements of same group of the periodic table contains same

number of valence electrons

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Group 1 (Lithium, Sodium, Potassium,…) has one valence electron

Group 2 (Beryllium, Magnesium, Calcium,…) has two valence

electrons

Group 7 (Fluorine, Chlorine, Bromine,…) has 7 valence electrons

Number of valence electrons gives group number

Number of shells filled with electrons gives period number

Nobel gases (Inert gases)

Group 0 of the periodic table (group 0 ) are very unreactive

because their valence level is completely filled by 8 electrons

except helium that has 2 electrons only

Noble gases are mono-atomic

Other elements tends to react to reach the stable electronic

configuration of the nearest noble gases

Isotopes

Different atoms of same element having same atomic number, but

different mass numbers

Isotopes have same number of protons, different number of

neutrons

Chemical properties of isotopes are similar because isotopes

have same number of valence electrons

Isotopes of carbon

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Atomic mass unit (a.m.u)

It is 1/12 the mass of carbon-12 isotope

Relative atomic mass (Ar)

It is the average mass of the isotopes of an element

Example:

Chlorine consists of two isotopes Cl & Cl , their relative abandence

is 3:1, the relative atomic mass can be calculated as follows:

Ar (Cl) =1+3

37)x(1+35)x(3=

4

(37)+(105)= 35.5 a.m.u.

Molecule: group of atoms (similar or different) combined together

Examples: H2, N2, O2, HCl, H2O, NH3, H2SO4

Diatomic molecueles of similar atoms

35

17

37

17

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Element, compound and mixture

Element: Pure simple substance which can't be split into simpler

forms and consist of same type of atoms

Compound: Two or more elements combined chemically in fixed proportions

Mixture: two or more elements or compounds uncombined chemicallymixed

together in any proportions

Elements can be classified into metals, non-metals and metalloids

Property Metals Non-metals

State at roomtemperature

Solid except mercury(liquid)

Solid, gases and onlyliquid bromine

Melting and

boiling points

High except group 1

(alkali metals)

Low except carbon and

silicon

Appearance Shiny DullElectric and

thermalconductivity

Good Poor or don’t conduct

electricity, exceptgraphite

Effrect of hammering

Malleable and ductile Brittle

Mixture

Compound

Element

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×

××

× × ×

××

×

+ - ×

Chemical bonding

There are many types of bonding between atoms, two of which are:

1) Covalent bond

2) Ionic bond

Ionic bond

Formed by transferring one electron or more from a metal atom

(Forming a positive ion) to a non-metal atom (forming a negative ion)

Ionic bond can be defined as " The electrostatic force of attraction

between oppositely charged ions"

Formation of sodium chloride (table salt)

Sodium atom loses its valence electron to form Na+ ion, the electron is

transferred to chlorine atom to form Cl- ion (chloride ion)

Electrostatic attraction force takes place between oppositely charged ions

The arrangement of valence electrons in sodium chloride can berepresented as

Na + Cl Na Cl

(2,8,1) (2,8,7) (2,8) (2,8,8)

Sodium atom Chlorine atom Sodium ion Chloride ion

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×

× ××

2+

×

× ××

2-

×

××

× ×

×

2+ --

××

× ×

Magnesium oxide (MgO)

Magnesium atom (12Mg) loses two electrons and changes into apositive magnesium ion (Mg+2).

Oxygen atom (8O) gains two electrons and changes into a negativeoxide ion (O2-).

Opposite charged ions are attracted together by the strong "ionicbonds".

Mg + O Mg O2,8,2 2,6 2,8 2,8

Magnesium Oxygen Magnesium Oxide

atom atom ion ion

Same for the formation of calcium oxide

Formation of Calcium Chloride (CaCl2)

Cl

Ca + Cl Ca Cl

Cl

N.B: An ion is an atom or group of atoms carrying positive or negativecharges

×

× × ××

×

× × ××

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Covalent bond

A bond formed by sharing a lone pair of electrons or more between two non-

metal atoms

A lone pair of electrons is attracted to the nuclei of both shared atoms

Single covalent bond

One pair of electrons is shared between two non-metal atoms

A) Hydrogen molecule (H2) H‒H

Each hydrogen atom share one electron to reach the

stable electronic configuration of noble gas helium

(2He).

B) Chlorine molecule (Cl2) Cl‒Cl

Each chlorine atom shares one electron to reach

the stable electronic configuration of noble gas

(18Ar).

Each chlorine atom still has three lone pair of

electrons

C) Water molecule (H2O) H‒O‒H

Each hydrogen atoms shares one electron to

reach the stable electronic configuration of noble

gas helium (2He)

Oxygen atom shares two electrons to reach the

stable electronic configuration of noble gas neon

(10Ne)

Oxygen atom still has two lone pair of electrons

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D) Ammonia molecule (NH3)

Each hydrogen atoms shares one electron to

reach the stable electronic configuration of noble

gas helium (2He)

Nitrogen atom shares three electrons to reach the stable electronic

configuration of noble gas neon (10Ne)

Nitrogen atom still has a lone pair of electrons.

E) Methane molecule

Caron atom shares four electrons with four hydrogen

atoms to reach the stable electronic configuration of

nearest noble gas neon (10Ne)

Each hydrogen atom shares one electron with the

carbon atom to reach the stable electronic

configuration of noble gas helium (2He)

Double covalent bond

Two pairs of electrons are shared between two non-metal atoms

A) Oxygen molecule (O2) O ═ O

Each oxygen atom shares two electrons to

reach the stable electronic configuration of

noble gas neon (10

Ne)

B) Carbon dioxide (CO2) O ═ C ═ O

Each oxygen atom shares two electrons

to reach the stable electronic

configuration of noble gas neon (10Ne)

Carbon atom shares four electrons to

reach the stable electronic configuration

of noble gas (10Ne)

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C) Ethene (C2H4)

Each carbon atom shares four electrons to

reach the stable electronic configuration of

noble gas neon (10Ne)

Each hydrogen atom share one electron to reach the

stable electronic configuration of noble gas helium

(2He).

Triple covalent bond

Three pairs of electrons are shared between two non-metal atoms

A) Nitrogen (N2) N≡N

Each nitrogen atom shares three electrons to

reach the stable configuration of noble gas neon

(10Ne)

B) Ethyne-Acetylene

Each carbon atom shares four

electrons to reach the stable

electronic configuration of noble gas

neon (10Ne)

Each hydrogen atom share one electron to reach the stable electronic

configuration of noble gas helium (2He).

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Covalent compounds Ionic compounds

Most are gases and liquids, feware solids with low melting points

They are solids with high meltingand boiling points

They don’t conduct electricity solids conduct electricity onlywhen molten or dissolved in water

They are usually less soluble inwater

They are usually soluble in water

They have low melting and boilingpoints

They have high melting andboiling points

Explain why?

Ionic compounds have high melting point and boiling point

strong attraction force between oppositely charged ions

need alarge amount of energy to overcome

Magnesium oxide has higher melting point than sodium chloride

Increased charges on magnesium and oxide ions

Greater attraction between ions needs a larger amount of energy to overcome

Metallic bond

Valence electrons of metals like sodium, magnesium, aluminum and iron

are so weakly bounded to metal atoms and are free to move throughoutthe whole metal leaving ions with positive charges

Metallic bond is a lattice of positive metal ions in a sea of

delocalized electrons.

As number of valence electrons in metal increases, hardness,

electric conductivity and melting point of metal also increases.

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Giant structures

1- Giant Ionic structure

2- Giant covalent structure

3- Giant Metallic structure

1) Giant ionic structure

Ionic compounds consist of a regular arrangement of +ve and –ve ions

held together by strong electrostatic attraction force

Structure of sodium chloride

2) Giant covalent structure

Diamond

Diamond is a pure form of carbon

Each carbon atom bonds strongly to four

other carbon atoms in a tetrahedral

arrangement

Graphite

Graphite is another form of carbon Graphite is a layer structure

Within each layer, carbon atom bonds

strongly to three other carbon atoms forming

hexagonal rings

These layers are held together by weak

attraction forces

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Buckminsterfullerene (Simple molecular Structure)

A new discovered form of carbon

A spherical molecule with the formula c60. It resembles a football made of 20 hexagons and 12

pentagons.

Each carbon atom in the structure is bonded covalently

with 3 others

Property Diamond Graphite

Appearance Colorless transparentcrystal which sparklesin light

Dark Grey, shinysolid

Electric conductivity Doesn’t conduct

electricity

Conducts electricity

Hardness A very hardsubstance

A soft material with aslippery feel

Use Drilling & cuttingJewelery

LubricantPencils

Structure of silicon (IV) Oxide (Silicon dioxide)

Silicon dioxide (SiO2) is a giant covalent compound in

which each silicon atom is bonded to four oxygen

atom and each oxygen atom is bonded to two silicon

atoms in a tetrahedral arrangement

Like diamond, silicon dioxide is hard, high melting

point and doesn’t conduct electricity

Explain why? Giant covalent structures have high melting point

Many strong covalent bonds

Needs a large amount of energy to overcome

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Giant metallic structure

A metal structure consists of a regular arrangement (lattice) of

positive ions in a sea of electrons

The metal is held together by the strong attraction between metal

positive ions and the delocalized electrons. Metals are good conductors of electricity because their electrons are

free to move throughout the metal structure.

Metals are malleable (easily shaped) and ductile (easily pulled into

wires) because Layers of metal ions (atoms) can slide over each other

Simple molecular structures

Simple covalent molecules like HCl, Cl2, Br2, I2, NH3,CO2 and CH4 have the

following properties

1- They are gases, liquids or solids with low melting points

Due to the weak attraction force between their molecules

Little amount of energy is needed to overcome

2- They are insoluble in water, unless they react with it.

3- They are soluble in organic solvents

4- They don’t conduct electricity

Bonding Ionic Covalent Metallic

Structure Giant ionic Simple

molecular

Giant

covalent

Giant

metallicMelting

point

High Low Very high Usually high

Electrical

conductivity

Conducts

electricity

whenmolten oraqueous

No No

Except

graphite

Yes

Examples NaCl,MgO,CaCl2

HCl, F2,Cl2, Br2,

I2, NH3,CO2

CH4,H2O,S8,P4,

C60,C2H4,C2H5OH

Diamond,graphiteand siO2

Na, Mg,Ca,Al,Fe,Cu,Zn

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Stoichiometry (Formula, equations and calculations)

Chemical formula and naming of compounds

A) Ionic compounds

Write the positive ion first (Metal ion or ammonium ion) followed by

the negative ion (Non-metal ion or polyatomic ion)

Positive ions Negative ions

H+ Hydrogen ionLi+ Lithium ionNa+ Sodium ionK+ Potassium ionAg+ Silver ionNH4

+ Ammonium ion

H- Hydride ionF- Fluoride ionCl- Chloride ionBr- Bromide ionI- Iodide ionOH- Hydroxide ionNO2

- Nitrite ionNO3

- Nitrate ionCa2+ Calcium ionMg2+ Magnesium ionBa2+ Barium ionPb2+ Lead (II) ionCu2+ Copper (II) ionZn2+ Zinc ionFe2+ Iron (II) ion

O2- Oxide ionS2- Sulphide ionSO3

2- Sulhpite ionSO4

2- Sulphate ionS2O3

2- Thiosulphate ionCO3

2- Carbonate ionSiO3

2- Silicate ionFe3+ Iron (III) ionAl3+ Aluminum ion

N3- Nitride ionP3- Phosphide ionPO4

3- Phosphate ion

The number of positive charge and negative charges must be

balanced to make the total molecule neutral

Aluminum oxide Al3+

O2-

Al2O3

Magnesium oxide Mg2+

O2-

MgO

Sodium Carbonate Na+

CO3

2-Na2CO3

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Examples:

Formula Name Formula Name

Na2CO3 Sodium

carbonate

KI Potassium

iodide

MgSO4 Magnesiumsulphate NaNO3 SodiumnitrateCu(OH)2 Copper(II)

hydroxide

FeSO4 Iron(II)

sulphateCa3(PO4)2 Calcium

phosphate

ZnCO3 Zinc

carbonateNH4Cl Ammonium

chlorideBaSO4 Barium

sulphate

Na2S2O3 Sodiumthiosuphate

Pb(NO3)2 Lead (II)nitrate

AgNO3 Silver Nitrate CaSiO3 Calcium

silicate

B) Covalent compound

Formula of Compound Name

NH3 AmmoniaH2O WaterCH4 MethaneCO Carbon monoxideCO2 Carbon dioxideNO Nitrogen monoxideNO2 Nitrogen dioxideSO2 Sulphur dioxideSO3 Sulphur trioxideHCl Hydrogen chloride

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Chemical equation

During a chemical reaction, reactants are converted into products

Steps for writing a chemical equation

1- A word equation is written, reactants on the left and products on

the right

2- The formulae of reactants and products are written in symbols

3- The equation is balanced

4- State symbol solid (s), liquid (l), gas (g) and aqueous (aq.) is used

to indicate the physical state of substance.

Example:

When hydrogen gas is burnt in air, it forms water vapour.

Hydrogen + Oxygen Water

H2 (g) + O2 (g) H2O (g)

2H2 (g) + O2 (g) 2H2O (l)

When sodium reacts with chlorine gas, sodium chloride is formed

Sodium + Chlorine Sodium chloride

Na (s) + Cl2 (g) NaCl (s)

2Na (s) + Cl2 (g) 2NaCl (s)

Magnesium burns in air to form magnesium oxide

Magnesium + Oxygen Magnesium oxide

Mg (s) + O2 (g) MgO (s)

2Mg (s) + O2 (g) 2MgO (s)

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Chemical calculations

Calculation of relative formula (Molecular) mass Mr

Using the relative atomic mass (Ar) from the periodic table we cancalculate the relative formula mass (Mr)

Example:

Mr of cl2= (2x35.5) = 71

Mr of NaCl = 23 + 35.5 = 58.5

Mr of CaCl2 = 40 + (35.5x2) = 111

Mr of Na2CO3 = (23x2) + 12 + (16x3) = 106

Mr of CO2 = 12 + (16x2) = 44

Mole (unit for measuring amount of a substance)

It is the relative atomic mass (Ar) or relative molecular mass (Mr) in grams

Example:

1 mole of Na = 23 gm

1 mole of Ca = 40 gm

1 mole of H2O = (1x2) + 16 = 18 gm

1 mole of CaCO3 = 40 + 12 + (3x16) = 100 gm

Examples

Calculate the mass of 0.1 mole of water

Mass = Number of moles x Mr = 0.1 x 18 = 1.8 gm

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Calculate the number of moles of sodium chloride whose mass is

117 gm.

N

Avogadro's number (NA) [6.02 x 1023]

It is the number of particles (atoms, molecules, ions and electrons) in one

mole of a substance

Example:

Calculate the number of atoms in 0.5 mole of sodium

Number of atoms = number of moles x NA = 0.5 x 6.02 x1023

= 3.01x1023 atom

Calculate the number of moles of calcium carbonate which contains 12.04

x 1023 molecule.

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Example

Calculate the number of water molecules in 36 gram

Molar volume of gases

1 mole of any gas occupies a volume of 24 dm3 (24000 cm3 ) at room

temperature and pressure (r.t.p)

Example

Calculate the volume of 0.5 mole of ammonia

Volume = number of moles x 24 = 0.5 x 24 = 12 dm3

Calculate the number of moles of nitrogen in 48 dm3

Determining the formula of simple chemical compounds experimentally

Magnesium oxide

When a magnesium ribbon is heated strongly in air, it

burns very brightly to form white powder magnesium

oxide.


The following data was obtained from experiment

Mass of crucible 14.63 g

Mass of crucible and magnesium 14.87 g

Mass of crucible and magnesiumoxide

15.03 g

Mass of magnesium used 0.24 g

Mass of oxygen which has reacted

with magnesium

0.16 g

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Find the formula of magnesium oxide from the experimental data

Mg O

Mass reacting (g) 0.24 0.16Number of moles

=0.01 =0.01

Mole ratio 1 1

Formula MgO

This is the formula which gives the simplest ratio of atoms present; it is

called the Empirical formula

Determining the empirical formula of a hydrate

Many salts crystallize from their solutions with a number of water

molecules bonded to the salt; this is called water of crystallization.

Examples:

Hydrated Magnesium sulphate MgSO4.7H2O

Hydrated Sodium carbonate Na2CO3.10H2O

Finding n in BaCl2.nH2O

A sample of hydrated barium chloride is heated in a crucible until no further

decrease in mass.

The following data was obtained from experiment

Mass of crucible 30.00 g

Mass of crucible + barium chloridecrystals, BaCl2.nH2O

32.44 g

Mass of crucible + anhydrousbarium chloride, BaCl2

32.08 g

From these experimental data, we can find the value of n

Mass of BaCl2 = 32.08 – 30.00 = 2.08 g

Mass of water = 32.44 – 32.08 = 0.36 g

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BaCl2 H2OCombining mass (g) 2.08 0.36

Number of moles

=0.01 =0.02

Mole ratio 1 2Empirical formula BaCl2.2H2O

Determining Molecular formula from empirical formula

If the formula obtained from calculation doesn’t exist in nature, then the

actual molecular formula will be a multiple of the empirical formula.

To calculate the n: number of empirical formula units

Example

An unknown organic compound was found to contain 0.12 g of carbon and

0.02 g of hydrogen and the relative molecular mass (Mr) of the compound is

56. Find the molecular formula of this compound.

C Hmasses (g) 0.12 0.02

Number of moles

=0.01 =0.02

Mole ratio 1 2

Empirical formula CH2

Empirical formula mass = 12 + (1 x 2) = 14

Molecular formula of this organic compound is C4H8

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Moles and chemical equations

When we write a balanced chemical equation, we are indicating the number

of moles of reactants and products involved in the reaction.

Example


2Mg(s) + O2 (g) 2MgO (s)

2 moles 1 mole 2moles

(2 x 24) (2 x 16) 2(24 + 16)

48 g 32g 80 g

We can use the ratio between moles to calculate the mass of products

formed and the reactants involved in the reaction.

Example

Calculate the amount of calcium carbonate produced by thermal

decomposition of 50 gm of calcium carbonate.

CaCO3 CaO + CO2

1 mole 1 mole 1 mole

Number of moles of calcium oxide = 0.5 mole

Mass of CaO formed = number of moles x Mr = 0.5 x 56 = 28 g

Percentage yield

Many chemical reactions don’t produce the same amount of product as

that calculated from chemical equations.

This is caused because these chemical reactions are reversible.

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The actual yield is the measured mass of product obtained in the experiment.

The theoretical yield is the mass calculated from the equation of the reaction.

Example:

When 1000 g of sulphur dioxide is reacted with excess oxygen, 1225 g of

sulphur trioxide is produced.

2SO2 + O2 2SO3

Calculate the percentage yield.

Ratio of SO2:SO3 is 2:2 = 1:1

Number of moles of sulphur trioxide formed = 15.63 mole

Theoretical mass of sulphur trioxide produced = 15.63 x 80 = 1250 g

Moles and solutions

Concentration of a solution is measured in moles per cubic decimeter

(mol dm-3) or grams per cubic decimeter (g dm-3).

Example

Calculate the concentration (in mole dm-3 and g dm3) of a solution of

sodium hydroxide, NaOH, made by dissolving 10 g of solid sodium

hydroxide in water and diluting it to 250 cm3.

250 cm3 =

,

Concentration = 40 g dm-3

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Example

Calculate the mass of potassium hydroxide, KOH, that is needed to

prepare 500 cm3 of 2 mol dm-3 solution in water.

Number of = concentration X Volume of solution

Moles (in mole dm-3) (in dm3)

Number of moles of KOH = 2 x

= 1 mole

Mass of 1 mole of KOH = (1 x 39) + (1 x 16) + (1 x 1) = 56 g

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Electrolysis

Electrolysis is the breaking down of a compound using electric

current.

Electrolysis involves the formation of new substances.

Electric current is the flow of electrons or ions.

Covalent compounds don’t conduct electricity because they don’t

contain ions.

Ionic compounds conduct electricity only when they are molten

(melted) or when dissolved in water (aqueous solution).

Electrolytes are compounds that conduct electricity when molten or

aqueous and breaks down (undergo chemical change) during

electrolysis.

Non-electrolytes are compounds that don’t conduct electricity.

Describe an experiment to distinguish between electrolyte and non-

electrolyte.

Setup the following apparatus

Electrodes

Electric source

Electrolyte solution

Lamp

Active

Inert

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If the lamp glows, then the solution is electrolyte, if it doesn’t glow,

then the solution is a non-electrolyte.

Electrodes are rods that pass the electric current into and out of the

electrolyte.

Positive electrode is called the anode, negative

electrode is called the cathode.

Positive ions (cations) moves to cathode,

negative ions (anions) moves to anode.

Explain why ionic compounds conduct electricity only when molten or

in solution

Because in solid ionic compounds, the ions are not free to move, when they

are melted or dissolved in water, ions liberate from crystal and are free to

move to conduct electricity.

Electrolysis of molten lead (II) bromide [PbBr2] using inert

electrodes

Pb2+(l) + 2e

-Pb(l) (Cathode reaction)

2Br-(l) Br2(g) + 2e

-(Anode reaction)

PbBr2(l) Pb(l) + Br2(g)

Explain why lead (II) bromide has to be melted for electrolysis to

take place

Because lead (II) bromide has a giant ionic structure of lead (II) ions and

bromide ions packed regularly in a crystal lattice, thus their ions are not free

to move.

When solid lead (II) bromide melts, ions become free to move, and thus theelectrons flow in the external circuit.

When a molten ionic compound is electrolyzed, a metal is formed at

the negative electrode and a non-metal is formed at the positive

electrode

Silver/grey Red/Brown

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Electrolysis of aqueous solutions

1) Electrolysis of concentrated sodium chloride solution (Brine)

Water is a weak electrolyte which ionizes according to the following equation

H2O H+ + OH-

So in aqueous solution of sodium chloride, we have the ions Na+, H+, Cl-, OH-.

For positive ions: Na+ Mg2+ Al3+ H+ Cu2+

More likely to be reduced

For negative ions: SO42- NO3

- Cl- Br- I- OH-

More likely to be oxidized

At anode, chloride ions lose electrons and are discharged as chlorine gas

2Cl-(aq) Cl2(g) + 2e-

At cathode, hydrogen ions accepts electrons and

discharge rather than sodium ion.

2H+ (aq) + 2e-

H2(g)

Na+ and OH- remains in the solution, so an aqueous solution of sodium

hydroxide is formed.

Electrolysis of brine (concentrated aqueous solution of sodium

chloride) in a diaphragm cell is used to produce chlorine,

hydrogen and sodium hydroxide on a large scale.

Uses of Chlorine

1) Making bleaches

2) Sterilizing water supplies

3) Making hydrochloric acid

Uses of hydrogen

1) fuel

2) Making ammonia

Uses of sodium hydroxide

1) Making bleaches 2) Making soap 3) Making Paper

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Electrolysis of copper (II) sulphate using inert electrodes (graphite

or platinum electrodes)

The ions present in an aqueous solution of copper (II) sulphate are Cu2+ (aq),

SO42- (aq), H

+(aq) and OH

-(aq) ions.

At Anode

The anode cannot lose electrons because it is inert.

Hydroxide ions (OH-) rather than sulphate ions (SO4

2-) are discharged.

4OH-(aq) O2(g) + 2H2O(l) + 4e

-

At cathode

Copper ions (Cu2+) rather than hydrogen ions (H+) are discharged because

copper is below hydrogen in reactivity series.

Cu2+(aq) + 2e

-Cu(s)

The electrolyte changes from blue to colorless because copper (II) sulphate

solution turned to dilute sulphuric acid.

Electrolysis of copper (II) sulphate using copper electrodes

At anode

Because the anode is not inert, it loses electrons and copper ions go into the

solution. Anode gets smaller

Cu(s) Cu2+(aq) + 2e-

At cathode

Copper ions rather than hydrogen ions are discharged because they are lower in

activity series.

Cu2+(aq) + 2e- Cu(s)

The electrolyte remains the same blue color. This is because copper ions removedfrom the cathode are replaced in solution by copper ions formed at the anode.

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Electrolysis of dilute sulphuric acid using carbon electrodes

At cathode

Only H+

ion (from acid and water) discharge into hydrogen gas

2H+(aq) + 2e- H2(g)

At anode

Hydroxide ion OH- (from water) discharge easier than

sulphate ion SO42- from sulphuric acid.

4OH-(aq) O2(g) + 2H2O(l) + 4e

-

For every four moles of electrons that flow around the circuit, one mole

of oxygen and two moles of hydrogen are produced

Electrolysis calculations

The quantity of electricity flowing through an electrolysis cell is measured in

coulombs (C).

If one ampere (A) is passed for one second, the quantity of electricity is

said to be 1 coulomb.

(Coulombs) (Amperes) (Seconds)

Faraday is the amount of electricity carried by one mole of electrons (6.02

x 1023 electron) and is equal to 96500 coulomb.

1mole of electron = 96500 coulombs

Example

Na+(l) + e- Na(l)

1 mole of sodium, Na, is produced by the flow of 1 mole of electrons (1

Faraday).

Cu2+(aq) + 2e

-Cu(s)

1 mole of copper, Cu, is produced by the flow of 2 moles of electrons (2

Faraday).

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2Cl- (l) Cl2(g) + 2e-

1 mole of chlorine, Cl2, is produced by the flow of 2 moles of electrons (2

Faraday).

Example:

Calculate the number of moles of electrons required to deposit 10 g of silver on

the surface of a fork during an electroplating process.The cathodic reaction is:

Ag+(aq) + e

-Ag(s)

1 mole of electrons 1mole of Ag

1 mole of electrons 108 g of Ag

? 10 g of Ag

Example

What is the mass of copper deposited on the cathode during the electrolysis

of copper (II) sulphate solution if 0.15 amps flow for the 10 minutes?

Electrode equation is: Cu2+(aq) + 2e

-Cu(s)

2 moles of electrons 1mole of copper

2 x 96,000 coulomb 65 g of copper

192,000 coulombs 65 g of copper

90 coulomb ?

90 coulombs give

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Example

During the electrolysis of dilute sulphuric acid using platinum electrodes,

hydrogen is released at the cathode and oxygen at the anode. Calculate the

volumes of hydrogen and oxygen produced (measured at rtp) if 1.0 amp

flows for 20 minutes.

The electrode equations are

2H+(aq) + 2e- H2(g)

4OH-(aq) 2H2O(l) + O2(g) + 4e

-

2 moles of electrons 1 mole of H2

2 x 96500 24,000 cm3 at rtp

192,000 coulombs 24,000 cm3 at rtp

1200 coulomb ?

1200 Coulomb will give

4 moles of electrons 1 mole of O2

4 x 96,000 coulombs 24,000 cm3 of O2

384,000 Coulombs 24,000 cm3 of O2

1200 coulombs ?

1200 coulombs will produce

Therefore, 150 cm3 of hydrogen and 75cm3 of oxygen are produced.

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Example

How long would it take to deposit 0.500 g of silver on the cathode during the

electrolysis of silver (I) nitrate solution using a current of 0.25 amp? The

cathode equation is:

Ag+(aq) + e

-Ag(s)

1 mole of electrons 1 mole of Ag

96,000 Coulombs 108 g of Ag

? 0.500 g of Ag

0.500 g of Ag will need

The time needed to deposit 0.500 g of silver is 1780 seconds.

part 1 (2011)

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