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IGCSE CHEMISTRY REVISION NOTES

1.The particulate nature of matter

 PROPERTIES OF SOLIDS/LIQUIDS AND GASES:

 Solids:

  •  The particles in a solid are arranged in a fixed pattern
  • The particles of a solid are very close to each other
  • The particles can only vibrate about their mean position

Liquids:

  • The particles in a liquid are not arranged in any fixed pattern.
  • The particles in a liquid are very close to each other.
  • The liquid particles can slide past over each other.

Gases:

  • The particles in gases are arranged in a random manner.
  • The particles of a gas are further apart from each other.
  • The gas particles are free to move everywhere rapidly

COMPRESSING GASES 

The gas particles are far apart,while the liquid particles are close to each other. When equal force is applied on the syringe plungers, the liquid does not get squeezed so the plunger does not move. But the gas gets compressed as the plunger is moved inwards as the gas particles come closer. Hence the volume of the gas in the syringe decreases.

 

CHANGE OF STATE:

Some substances directly change state from solid to gas or from gas to a solid . This process is called as sublimation

  Examples of substances that sublime are :

  • Dry ice (solid carbon dioxide)
  • Iodine
  • Arsenic
  • Naphthalein ( The stuff the   camphor balls are made of)

BROWNIAN MOTION:

 Define Brownian motion 

Brownian motion is the random movement of large molecules due to their collision with the faster moving smaller molecules.

DIFFUSION

Define diffusion:Diffusion is the random movement of particles from a region of their high concentration to a region of their low concentration down the concentration gradient.

Speed of diffusion depends upon the Mr( relative molecular mass):

This means that the speed of diffusion of a gas depends on how heavy it's molecules are. Molecules with a lighter mass diffuse faster than those with a heavier mass.

Why do the gases diffuse?- Explanation based on kinetic theory:

Gases diffuse because their particles move in random motion. These particles then collide .This diffusion is from a region of high concentration to a region of low concentration.

Diffusion happens faster in warmer temperatures than in cooler temperatures.

When a liquid spills on a floor and can be smelt far away it means that it has first evaporated then diffused

When a precipitate  is  formed during chemical reactions, the particles, diffuse, collide and then react

   Particles spread out evenly as a result of diffusion

   Generally diffusion happens in liquids and gases as   the particles are free to move. Their particles are constantly moving colliding and changing directions

    Diffusion in gases is faster than diffusion in liquids because the gas particles move rapidly. They are able to move freely because kinetic theory assumes that there are no            forces of attraction between the gas particles while there are weak forces of attraction between liquid particles.

    Diffusion does not happen in solids because the particles are tightly packed and they can only vibrate in their mean positions and not move about.

    Diffusion can occur in liquids which are miscible

    Diffusion is also possible in solids that dissolve in liquids.

    At the same temperature, the molecules that have the lower mass diffuse faster than the heavier molecules. If the lighter and heavier molecules have the same amount of            energy when they collide, then, the lighter ones will bounce off the heavier ones at a faster rate. So, lighter molecules diffuse faster than the heavier molecules.

 

 

2 Experimental techniques

     MEASUREMENT 

    Know your lab apparatus well

  • Volume is measured using a burette, measuring cylinder and pipette.
  • Burette can measure volumes upto 50ml.
  • Volumetric pipettes generally come in 10cm3 and   25cm3 sizes.
  • Note 1ml=1cm3 

      PURITY

Characteristics of pure substances:

  • Pure substances have   sharp melting and boiling points.
  • Example: Boiling point of pure water is 1000C while melting point of pure ice is 0oC.

Effect of impurities on "pure substances":

  • Due to the presence of impurities, the melting and the boiling points will not be sharp any more. Substances will melt and boil over a range of temperatures.
  • The boiling point will be increased further due to the presence of impurities. Example: Impure water will boil above 1000C.
  • The greater the impurity, the greater will be the increase in the boiling point.
  • The melting point is decreased by impurities. So impure ice will not melt at 00C but will melt at a lower temperature
  • Define:

  • Melting point: It is the temperature at which a solid changes into a liquid.
  • Boiling point: It is the temperature at which a liquid changes into a gas.
  • Naming the test for purity of substances:

    The following are the tests to check the purity of substances:

  • Testing for the melting points and boiling points of substances.
  • Performing chromatography and checking the purity of substances 

 

     PURIFICATION TECHNIQUE - FILTRATION

Filtration is insoluble.

  • Solution: A solution contains a solid dissolved in a solvent.
  • Solute: The dissolved solid is called as the solute.
  • Solvent: The liquid that dissolves the solid is the solvent.
  • Filtration: Filtration is a method for separating an insoluble solid from a liquid.
  • Residue: The solid that stays on the filter paper is the residue.
  • If a substance does not dissolve in a solvent, we say that it is insoluble. For example, sand does not dissolve in water – it
  • When a mixture of sand and water is filtered:
  • the sand stays behind in the filter paper (it becomes the residue)
  • the water passes through the filter paper (it becomes the filtrate 

 

 PURIFICATION TECHNIQUE-CRYSTALISATION

Definition:Crystallisation is a separation technique that is used to separate a solid that has dissolved in a liquid and made a solution

To obtain pure crystals of hydrated salt from a metal carbonate or metal oxide.

Example: To make crystals of hydrated Cobalt(II) chloride( CoCl2.6H2O)  Step1: Pour the acid into the beaker and warm it gently. [Remember that warming the acid speeds up the reaction.]

Step2:Add a measured amount of the carbonate to the acid with a spatula and stir it with a glass rod. [It is stirred with a glass rod and not the spatula because spatula which is made of metal might react whereas the glass does not react.]

Step3:The step 2 is repeated till no more cobalt carbonate reacted.{The student knows when no more cobalt carbonate will react when solid cobalt chloride is visible or when there is no more fizzing or no more gas formed [Note you cannot mention colour change as a reason]}

  Step4:The mixture is then filtered or decanted to remove the (unreacted) excess cobalt carbonate.

Step5:The filtrate is then heated(evaporated) until the crystallisation point  is reached. [You come to know when the crystallisation point has reached when crystals          start  forming    on   the  edge      of  the  glass    rod].It    is          then left      to    cool      in           an   evaporating dish.

  Step 6: When crystals form , filter off the crystals. Dry the crystals by pressing them between filter papers or drying them in oven at low temperatures.

Note

  1. If cobalt chloride crystals are heated then the water is lost  (crystals get dehydrated) and cobalt chloride becomes anhydrous and turns blue.[ Rejected answer: To      write that the crystals will break or a powder will form if crystals are heated]
  1. If instead of cobalt carbonate, magnesium carbonate would have been used, then warming the acid is not needed as magnesium carbonate reacts quickly at room temperature and no heat is needed.
  2. Note that the acid should never be in excess as this will make the solution acidic and not neutral and thus the salt will be impure.
  3. Excess oxide/carbonate    is added to ensure that all the acid is neutralised ( used up)
  4. When a metal carbonate is formed, water + carbon dioxide are the byproducts.Carbon dioxide causes bubbling during the reaction, and can be detected using limewater.
  5. Crystals are dried using filter paper (and not heat) to prevent the breakdown of the crystals

 

If you are asked in general  how to obtain salt crystals quickly, then the obvious answer will be to heat/evaporate till the crystallisation point is reached or till it is saturated

Saturated solution: It is a solution in which no more of the solute dissolves.

 

 PURIFICATION TECHNIQUE - SIMPLE DISTILLATION 

Simple distilaltion

Use : Used to separate a dissolved solid ( solute) from a solution. Method:

Simple distillation is a method for separating the solvent from a solution. For example, water can be separated from salt solution by simple distillation. This method works because water has a much lower boiling point than salt. When the solution is heated, the water evaporates. It is then cooled and condensed into a separate container. The salt does not evaporate and so it stays behindEvery pure substance has its own particular melting point and boiling point. One way to check the purity of the separated liquid is to measure its boiling point. For example, pure water boils at 100°C. If it contains any dissolved solids, its boiling point will be higher than this.

 

PURIFICATION TECHNIQUE - FRACTIONAL  DISTILLATION 

Fractional distillation : It is a method of separating two miscible  liquids with a difference in their boiling points.

Fraction: Each distillate that is collected is called as fraction.

The apparatus consists of:

A-Thermometer

B- Beaker C- Tripod

D-Condenser

E- Round bottom flask

F-Fractionating Column

 

   Principle of operation:Two miscible liquids can be separated if they have a difference in their boiling points

Use of fractional distillation: 

  1. To separate ethanol from fermented mixture
  2. To separate a mixture of two liquids with a difference in their boiling points.

    Purpose of the: 

  1. Condenser: The condenser    condenses the    vapour coming out of the fractionating column.
  2. Fractionating column containing beads: The glass beads provide a large surface area for the vaporisation and condensation of the liquid mixture. Small glass beads are preferred over large glass beads as they   provide a large surface area
  3. Thermometer: To check if the complete separation of the substances has occured or no.Example:If you have to separate  ethanol ( Boiling point 780C) and butanol (boiling point 1180C).As temperature increases and till it is below 780C,  ethanol and butanol both evaporate and condense.

There is a range of temperatures in the fractionating  column.  The greatest temperature is lower down while the temperature decreases as you go up i

When temperature reaches 780C, the ethanol starts evaporating and at this temperature it boils off and passes to the condenser as ethanol vapour.The ethanol vapour condenses and changes into liquid ethanol. This liquid ethanol  then gets colled as the distillate

The butanol which has evaporated condenses upon reaching the glass beads and falls back in the flask. Thus it is prevented from reaching the condenser The temperature stays constant till all the ethanol has evaporated. The temperature will only begin to rise when all the ethanol has evaporated.

A possible hazard in this experiment is that the alcohols are flammable due to the heat of the bunsen burner.

Method to check the purity of the liquid obtained:

We may measure the boiling point of the liquid obtained to check its purity.

 

  PURIFICATION TECHNIQUE - CHROMATOGRAPHY

Paper chromatography:

The method of separating pigments(colored substances) using filter paper is paper chromatography.

Key points about chromatography:

  1. The colours separate if
  • The pigments have different solubilities in the solvent.
  • The pigments have different degrees of attraction to   the filter paper

 

3. Atoms , Elements and Compounds 

 KEY DEFINITIONS AND STRUCTURE OF AN ATOM    

       Define:   

  • Proton number ( Atomic number): It is the number of protons in the nucleus of an atom.
  • Nucleon number ( Mass number): It is the total number of protons and neutrons in the nucleus of an atom.
  • Isotope: Atoms of the same element having the same number of protons but different number of neutrons are  called  as isotopes.
  • Elements :Elements contain only one type of an atom
  • Mixture: A mixture contains two or more elements that are not chemically bonded together. They do not have a fixed composition.
  • Compound: A compound is a substance made up of two or more different types of elements joined together by chemical bonds.
  • Covalent     bond:   A   bond   formed     by   sharing   of    electrons   is called a covalent bond.
  • Ionic bond: The bonding between oppositely charged ions is called ionic bonding.
  • Macromolecule: Macromolecules are giant covalent structures
  • Metallic bonding: The bonding that exists between the positive ions and the -vely charged electrons is called as metallic bonding,

   Structure of an atom

An atom is the smallest uncharged particle that can take part in a chemical change.

  • An atom contains a centrally located nucleus.
  • The nucleus contains positively charged   protons   and      neutral neutrons.[Protons +  neutrons= nucleon number or mass number]
  • The electrons revolve around the nucleus in fixed orbits called electronshells or energy levels.
  • An atom is electrically neutral as the number of protons (+vely charged) are equal to the number of electrons (-vely charged).

         

        COMPARING METALS AND NON METALS 

 Comparing elements/compounds and mixtures       

 Element- Cannot be broken to anything simpler by chemical means  Or made up of only one type of an atom

 Compound-Two or more elements chemically bond together to form a compound

 Mixture-Two or more substances that are not chemically bonded together are termed as a mixture

 

Comparing metals and non-metal                   

Physical property

Metals

Non-metals

Electricity

Conduct

Do not conduct

Heat

Conduct

Do not conduct-exception is graphite

Malleable

yes

No. They are brittle.

Ductile

yes

No. They are brittle.

Lustrous

yes

No They are dull-exceptions iodine and graphite

Sonorous

yes

No. They are non-sonorous

                          

 ISOTOPES

  • Atoms of the same element having same atomic number but different atomic mass number are called as isotopes.
  • Isotopes have the same properties because they have the same number of electrons in their outermost shell (valence shell).

There are two types of isotopes:

  • Radioisotopes
  • Non-radioactive isotopes

​Uses of Radioactive isotopes:

Medicinal uses:

  • Cancer treatment: An isotope of cobalt is used for this .The radiations given out by the radioisotope is used to kill cancer cells.
  • Treatment of overactive thyroid glands
  • Locating tumors in the body.  Industrial use:
  • Used to check the leaks in oil and gas pipelines. A radioisotope is added to the oil or the gas in the pipeline and a radiation detector called as the Geiger-Muller counter is used which gives a high reading at the area of the leak.

 ALLOYS

  •  Alloys are mixtures of two metals or one or more metal with a non metal
  •  When a metal is alloyed, the different sized atoms of the second metal make the lattice   arrangement of the first metal irregular. This prevents the metal layers from                   sliding past over each other when a force is applied. Thus making the alloy stronger than the original metal.

      Alloys Properties and Uses 

 

Alloys

Properties

Uses

Brass (Copper+ Zinc)

Stronger than copper but still malleable

Musical instruments. ornaments

Bronze (Copper + Tin)

Very hard

Some moving parts of statues,machines, bells

Stainless steel (Iron+ Chromium+ Nickel)

Does not rust like iron

Car parts, cutlery, surgical instruments

 

  COMPARING IONIC AND COVALENT  BONDS

   Ionic and covalent compounds- A comparison

 

Covalent

Ionic

Formed between

Non -metals

Metals and non-metals

Melting points and boiling points

Low melting and boiling

points because the intermolecular attractive forces are very weak.

Exceptions are: SiO2 with a high melting point

High melting and boiling points because of strong electrostatic forces                                          between the ions in the giant lattice

Solubility

Insoluble in water

( Exceptions: sugar and amino acids-water soluble)

Soluble in water because the water molecules are able to separate the ions from one another and keep them in the solution.

Electrical

Do not conduct electricity

They conduct electricity in

conductivity

because they have no ions.

the molten or aqueous form

Hydrogen chloride gas, a

due to the presence of

covalent compound     reacts

mobile ions.

with water to form HCl acid

which splits up into ions.

 

MACROMOLECULES

Giant covalent structures of Diamond/Graphite/Silicon(IV) oxide

  • Macromolecules are giant covalent structures
  • They have a rigid three dimensional network of strong covalent bonds throughout the crystal.
  • It takes a lot of energy to break the bond
  • They have very high boiling and melting points.
  • Example of giant covalent structures are: Graphite, diamond and [Silicon(IV) oxide-also called as silicon dioxide]

          Graphite

  • Graphite is a black shiny solid.
  • It's carbon atoms are arranged in layers.
  • Every layer consists of carbon atoms in hexagonal ring.
  • Each carbon atom is covalently bonded to 3 other carbon atoms and one electron is set free. Thus in the entire lattice there is a sea of delocalised electrons. Hence graphite conducts electricity because the delocalised electrons can drift along the layers when voltage is applied. As graphite conducts electricity , it is used to make electrodes for electrolysis.
  • Graphite has a slippery feel as the bonding between the layers in graphite is weak. hence the layers can slide past each other. Hence it has a slippery feel.
  • The layers of graphite can flake off because of this weak bonding. hence it is used as a lubricant and in pencil leads.

 

       METALLIC BONDING AND PROPERTIES OF METALS

  • Metallic bonding is a third type of giant structure.
  • The metal atoms are closely packed together in a regular arrangement As they are very close to each other, the valence electrons tend to drift away from the atoms. Thus a sea of delocalised ( mobile) electrons is formed surrounding the positive metal ions. The positively charged metal ions are held together by their strong attraction to the mobile electrons that keep moving between the ions. This is metallic bonding. The electrostatic attraction between the metal ions and the electrons exists in all directions.

          Properties of metals

  • Most metals have a high melting and boiling point: Reason: It takes a lot of energy to weaken the strong forces of attraction between the metal ions and the delocalised electrons in the lattice. These attractive forces can be overcome only when the temperature is too high.
  • Metals are good conductors of heat and electricity. Reason: When a voltage is applied, the delocalised electrons move through the metal lattice towards the positive pole of the cell or power pack. But if the vibrations of atoms becomes faster due to high temperature , The electrons will not be able to move easily through the lattice. That is why the electrical conductivity of a lattice decreases with an increase in temperature. Conduction of heat is due to vibrations of the atoms passing on the energy from one atom to the next. The metallic structure allows the atoms to vibratemore freely. The delocalised electrons can also carry energy through the metal lattice structure quickly.
  • Metals are malleable and ductile.

      The positive ions in a metal are arranged regularly in layers. When a force is applied,  the layers can slide over each other. In a  metallic bond, the attractive forces                  between the metallic ions and the electrons exist in all directions. So when the layers slide, new bonds can easily form. This leaves the metal with a different shape.

 3.  PERIODIC TABLE 

    BASICS:

  1. Periodic table is a method of  classifying elements and it is  used to predict properties of elements.
  2. Elements in a periodic table are arranged in the increasing  order of their atomic numbers( proton numbers).
  3. Elements with similar properties fall under the same vertical columns called groups.
  4. Horizontal rows are called periods.
  5. Metals are on the left hand side of the periodic table.
  6. Non-metals are to the right hand side of the periodic table.
  7. Boron (B),Silicon (Si),Germanium ( Ge),Arsenic (As), Tellurium (Te) and Polonium ( Po) are the metalloids that separate the metals from the non metals

 

Metallic to non-metallic character across a period

  • Metallic character refers to the level of reactivity of a metal to donate electrons during a chemical reaction.
  • Metallic character decreases across a period from left to right.Nonmetallic character relates to the tendency of aelement to accept electrons during chemical reactions.
  • This occurs as atoms more readily accept electrons to fill a valence shell than lose them to remove the unfilled shell.

Group

I

II

III

IV

V

VI

VII

VIII

Number of outer shell electrons

1

2

3

4

5

6

7

8

  1. Group 1    to 3 are metals (except Boron). Their atoms form positive ions ( cations) by losing electrons.
  2. Group 4       have non metals at the top (C and Si) and non-metals below(Ge/Sn/Pb).
  3. Group 5   has non metals like ( N, P and As) and metals like   Sb and Bi.
  4. Groups6( except Po) and 7 are non metals. Their atoms form negative ions ( anions) by gaining electrons.

 

         Group1 Element

Group properties

Group 1 elements are : Li, Na, K, Rb, Cs and Fr.  Group 1 elements are known as soft metals.

Electronic configuration:[ 1st three elements only]  Li -(2,1)

Na-(2, 8,1)

K-( 2.8.8,1)

They show trends in melting point, density and reactivity with water as follows:

Reaction of group 1 metals with water:

General reaction:

Group 1 metals + water -----> Metal hydroxide   + hydrogen Example:

2Li (s) + 2H2O (l) -----> 2LiOH (aq) + H2(g)

All Group I metals dissolve in water to form alkalis.

 

 

 

 

 

 

 

5 Electricity and chemistry

6 Chemical energetics

7 Chemical reactions

8 Acids, bases and salts

9 The Periodic Table

10 Metals

11 Air and water

12 Sulfur

13 Carbonates

14 Organic chemistry