Chapter 4: Chemical Change

Physical Change- changes in the physical state but the substance remains chemically the same and it is often easy to reverse like the changes in physical states of water (H₂O remains H₂O !)

Chemical Change- In this change there is a change in the chemical formula of the substance. So new substances are formed, difficult to reverse.

Exothermic reaction- Energy is given out (fall in temperature of substance)

Endothermic reaction- Energy is taken in (rise in temperature of substance)

Total mass of reactants=Total mass of products always in a reaction

Alkali metal+water à  Metal hydroxide+hydrogen

2Na+2H₂O à  2NaOH + H₂

2H₂ + O_{2 -->} 2H₂O is an highly exothermic reaction. This is an important reaction and this reaction is even used to power rockets!

Types of Reaction

• Synthesis reaction- Where 2 or more substances react together to form just one product. 

2Mg + O_{2 -->} 2MgO The burning of magnesium for example is a synthesis reaction. 

Synthesis reactions usually need heat to start the reaction but are exothermic reactions.

Photosynthesis however is an endothermic reaction. It is a photochemical reactions meaning that it's rate of reaction is affected by the light. More light means more photosynthesis. The green pigment chlorophyll is needed for this reaction as well to trap the sunlight form the Sun.

6CO₂ +6H₂O à C₆H₁₂O₆ +6O₂  

• Decomposition- When one reactant breaks down to give two or more simpler products.

HgO à Hg + O₂  is an example of a decomposition reaction.

Thermal decomposition is the decomposition caused by heat. 

CaCO₃ à CaO + CO₂ is bought by heating limestone (CaCO₃ ) is thermal decomposition

Decomposition can also occur due to light energy.

2AgCl -->2Ag + Cl₂ is an example. This reaction is the base of photography. AgCl is white and Ag is grey ( Ag is sliver!)

• Neutralisation- a chemical reaction between an acid and base to produce a salt and water only.

HCl + NaOH à NaCl + H₂O

•  Precipitation- sudden formation of of a solid by mixing two solutions or by bubbling a gas in a solution 

Like the bubbling carbon dioxide in limewater, forming white precipitate of limestone, this is the test of presence of carbon dioxide.

CO₂+ Ca(OH)₂ --> CaCO₃(S)+ H₂O (limewater test)

This reaction is used to produce insoluble salts and in analytic tests of ions.

Pb(NO₃)₂ + 2KI -->PbI₂(S)+ 2KNO₃  (formation of lead (II) iodide which is yellow in colour, used to identify presence of iodide ions)

• Displacement reactions- It occurs because a more reactive element will displace less reactive one from a solution of one of its compounds.

Zn+CuSO_{4 -->}ZnSO₄ +Cu (colour changes from blue to colourless, since Copper sulphate is blue) Zinc is more reactive than copper therefore displaces it form copper sulphate solution.

Mg+2HCl-->MgCl₂ +H₂ (Magnesium is more reactive than hydrogen)

2K+2H₂0-->2KOH+H₂ (Some metals are so reactive that they displace hydrogen from water)

Cl₂+2KI-->2KCL +I₂  (Chlorine more reactive than iodine, displacement!) The colour changes from colourless to brown because of the formation of iodine.

• Combustion- Involves reaction with oxygen and releases energy, exothermic.

CH₄+2O_2--> CO₂ + 2H₂O is combustion of methane (natural gas) which is used as a fuel.

Respiration- C₆H₁₂O₆ + 6O₂ -->6CO₂ + 6H₂O

• Oxidation- Gain of oxygen atom-loss of electrons-increase in oxidation number.
• Reduction- Loss of oxygen atom- gain of electrons- decrease in oxidation number.
• Reducing agent- A substance that undergoes oxidation to help other substance to go under reduction (by providing oxygen/electrons by oxidising) 
• Oxidising agent-A substance that undergoes reduction to help other substance to go under oxidation (by taking away oxygen/electrons by reduction).
• Important reaction- 2Cu + O₂ --> 2CuO- colour changes from pink (copper) to black (copper (II) oxide)

Ionic reactions- 

• Tests for oxidising agents- adding potassium iodide KI, which is colourless turns to brown, which is the colour of iodine which forms if an oxidising agent is present- oxidises iodide ion I- to Iodine I2.
• Tests for reducing agents- adding potassium magnate (VII) which is purple turns colourless if there is a reducing agent is present.

Chapter 3: Strucutre of materials

Giant metallic lattices- a lattice of positive ions in the sea of electrons.
Giant ionic lattices- a lattice of alternating positive and negative ions. (NaCl)
Giant molecular lattices- a giant molecule making the lattice (Diamond)
Simple molecular lattice- simple molecules in a lattice held together by weak forces. (Chlorine gas)

Metal Crytals

• High density due to dense packing.
• Metals in the lower parts of the Periodic Table have the highest densities since they have higher atomic masses.
• Malleable and ductile- Metals are malleable since the positive ions are arranged in layers which can slide over each other without breaking the structure. Bonds are strong but not rigid.
• Conducts electricity and heat well- mobility of delocalisied electrons helps to conduct electricity and heat.
• Have crystalline structure.

Alloys

• Alloys are formed by mixing molten metals together and allowing them to cool.
• Alloys are stronger than individual metals. This is because the presence of an impurity atom of a different size prevents slipping between the layers making them more stronger.
• Important alloys- Brass- made of zinc and copper and Stainless steel- chromium, nickel and iron.

Ionic Crystals

  
• Ionic compounds form lattices consisting of positive and negative ions.
• There are equal numbers of both types of the ions so the charges balance.
• Ionic compounds are therefore neutral.
• They are hard
• More brittle than metallic crystals. The layers of ions cannot be slid over each other or the lattice breaks due to bringing same charges together.
• Dissolving the ionic lattice in water however breaks up the lattice and keeps ions apart.
• Ions in solution or when ionic compounds are molten (melted) means that the solution can carry a current since the ions are free to move hence can carry current.
• Ionic compounds that don't dissolve in water means that the ionic bonds must be really strong!

GIANT molecular crystals

Graphite, Diamond and Silicon (IV) oxide.

• Giant molecular crystals have strong covalent bonds.
• elements- carbon - diamond and graphite and in compounds- silicon oxide are examples of giant molecular crystals.
• Diamond and silicon oxide have the similar structure.
• Giant molecular crystals have high melting points and are very hard.

Diamond (Carbon form)

• Tetrahedral structure 
• Strong covalent bonds (high bp)
• Rigid and brittle
• No free electrons- doesn't conduct electricity.

Graphite (Carbon form)

• Has free electrons- conducts electricity.
• Carbon atoms arranged in layers which can slip over each other since the layers have weak forces between them (but the carbon atoms have strong covalent bonds)

IMPORTANT

	Diamond		Graphite
	Property	Use	Property	Use
Appearance	Colourless, transparent crystals that sparkle in light	In jewellery and ornamental objects	Dark grey, shiny solid	-
Hardness	The hardest natural substance!	In drill bits, diamond saws, and glass cutters.	Soft- the layers can slip over each other, has slippery feel	In pencils, as a lubricant.
Electrical Conductivity	Doesn't conduct electricity	-	Conduct electricity	Used as electrodes and carbon brushes in electric motors (physics ain't it :P)

Allotropy- When an element can exist in more than one structural form in the same physical state. Like carbon exists as diamond and graphite and both are solids at room temp.

Allotropes-different forms of the same element.

Chapter 3: Bonding

Diatomic Molecules-a molecule containing two atoms like Cl2 for example.
Bonding- Force of attraction between two atoms.
Ions- Charged particles made from an atom by the gain or loss of electrons to gain stability like the sodium atom will lose one electron to achieve a noble gas arrangement 2,8 (before 2,8,1). This will have a sign Na+ (+ sign on top)
Lattice- a regular three-dimensional arrangement of atoms, molecules or ions in a crystalline solid.
Metallic Bonding
This happens only in metals. Here the atoms in the metals lose the electrons from their outershell. These electrons become delocalized and the atoms become positively charged ions. This hence forms an electrostatic force between the regular array of +ve metal ions and the sea of delocalized electrons within a metal solid.
Metallic bonding is the strong attraction between closely packed positive metal ions and a 'sea' of delocalized electrons.

• The delocalized electrons can move freely therefore metals can conduct electricity.
• This electrostatic attraction is so strong that metals have high mp and bp.
• The +ve ions are arranged in layers which can slip over each other making metals malleable and ductile.
• Metallic bonding results in giant metallic lattices.

Non-metals + Non-metals= Covalent Bonding

• Non-metals combine together by sharing pairs of electrons. This is known as a covalent bond. This holds the atoms together. 
• Groups of atoms bonded together in this way is called molecules 
• They don't have any free electrons or ions so they don't conduct electricity.
• This covalent bond is weak and can be easily broken, therefore covalent compounds normally occur as liquids or gases at rtp like chlorine, hydrogen and water.
• Group IV oxides (except carbon dioxide) makes giant molecular lattices like silicon (IV) oxide. These lattices have many strong covalent bonds which make make their melting points very high since a lot of energy is needed to break all of these bonds.
• Group V,VI, VII oxides and CO2 form simple molecular lattice which is simple molecules. These have low melting points since they have less bonds to break.

Non-metals + Metals= Ionic Bonding

 (The two chlorine ions can be written as 2 then the chlorine ion structure as shown, these type of diagrams are always asked in the exam so better practice them like the one above, just show the outershell in these diagrams.)

• An ionic bond is formed between non-metals and metals.
• The metal loses electrons to become a positive ion to attain noble gas arrangement and the non-metal becomes a negative ion by gaining those electrons. 
• So there is a strong electrostatic force between the positive and negative ions. This is the ionic bond.
• Ionic bonds needs a lot of energy to break, therefore ionic compounds have high melting points. They are usually solid at rtp like NaCl (table salt).
• Ionic compounds conduct electricity only when they are molten or dissolved in water, not as a solid. In solids the ions cannot move therefore cannot conduct electricity. Yes, ions conduct electricity in ionic compounds.
• Ionic bonding results in ionic lattices a regular array of alternating +ve and -ve ions like in NaCl.

Formulas of poly atomic ions

Really important to learn these formulas of ions in writing formulas of substances.

Valency	Simple Metal Ions	Simple non-metallic ions		Polyatomic ions
	+ve	+ve	-ve	+ve	-ve
1	Sodium Na+ Potassium K+ Silver Ag+ Copper (I) Cu+	Hydrogen H+	Hydride H- Chloride Cl- Bromide Br- Iodide I-	Ammonium NH4+	Hydroxide OH- Nitrate NO3- Hydrogencarbonate HCO3-
2	Magnesium Mg2+ Calcium Ca2+ Zinc Zn2+ Iron (II) Fe2+ Copper (II) Cu2+		Oxide O2- Sulphide S2-		Sulphate SO42- Carbonate CO32-
3	Aluminium Al 3+ Iron (III) Fe3+		Nitride N3-		Phosphate PO43-

 Summary

More is on the way... keep checking :)

Chapter 3: The Periodic Table

(This is the best of what I could fit. I suggest to use your own periodic table instead :P)

• -Main-group elements- Groups I to 0
• -Alkali Metals- Group I
• -Halogens- Group VII (non-metals)
• -Noble gases- Group 0 (very unreactive)
• -Transition elements- The block of elements between Group II and III
• -The broadest distinction in the table is metals and non-metals. Non-metals are on the right of the thick line while the metals on the left side of the thick line.
• -A metal is an element that does conduct electricity, is malleable and ductile.
• -A non-metal is an element that doesn't conduct electricity and isn't malleable or ductile.
• -A metalloid is an element that have some features of a metal and some features of a non-metal.
• The difference between Metals and non-metals.

Metals	Non-Metals
They are usually solids at rtp except mercury which is liquid at rtp. They have high mp and bp usually	They are usually solids or gases at rtp except bromine which is a liquid at rtp. They have low mp and bp often.
They are usually hard and have high density	Non-metals are softer than metals usually. They have low densities usually.
All metals are good conductors of electricity.	Poor conductors of electricity
Malleable and ductile	Brittle
Grey in colour except gold and sliver, can be polished	Vary in colour, dull (when solid)
Sonorous	Not sonorous

Trends in the Periodic Table

• -Elements in the same group (vertical columns of elements) have the similar chemical properties and physical properties.
• -Elements in the same group have the same number of electrons in their outershell.
• -For main-group elements, the number of the group is the number of electrons in the outershell. For e.g chlorine is in Group VII so it has 7 electrons in its outershell.
• -The period (rows of elements) number tell us the number of shells in the element. Hydrogen in period 1 has 1 shell for e.g.
• -The atomic size of an atom increases down the group (as the number of shells increase) but decreases across a group since the number of electrons in the last orbit increases hence increasing the attractive force between electrons and protons (-ve and +ve) decreasing the size of the atom as a result.
• -Elements become more metallic (ability to lose electrons) down a group and less metallic across a period since the electrons in the last orbit increases so gaining electrons is much easier than losing them all.
• -In metal groups (like group I and group II) the reactivity of the elements increases as you go down the group.
• The most reactive metal is Caesium since Francium is radioactive (metals in Group I in more reactive than the metals in Group II)
• In the group of non-metals, the reactivity of the non-metals increase up the group
• The most reactive non-metal is Fluorine in Group VII. 

Chapter 2 Atoms and molecules

Important definitions-

Atom- The smallest particle of an element that can take part in a chemical reaction. All atoms of the same element have the same number of protons in the nucleus.

Pure substances can be of two types-

1. Element- Substances that cannot be chemically broken down into simpler substances. An element cannot be decomposed by passing electric current through it.
2. Compound- Pure substances made from two or more elements chemically combined together in fixed proportions. 

Mixtures can be of two types-

1. Homogeneous 
2. Heterogeneous (explained in the previous blog)

Difference between Mixtures and compounds

It is really important to understand the difference since most students don't understand the difference.

No	Mixtures	Compounds
1 2  3 4	Substances simply mixed together, no reaction takes place Composition of mixture can be varied like in air the carbon dioxide level could rise and fall, no fixed value Properties of the substances present remain the same. Substances in the mixture can be separated by physical means like filtration and distillation	Substances chemically react together to form a new compound Composition of new compound remains the same. Properties of the new compound formed are different form the elements it is made of. Compound cannot be easily separated into its elements

Kinetic Theory (the Dalton theory not required by syllabus,

• All matter is made up of small particles
• The particles are moving all the time. The higher the temperature the higher the average energy and speed of particles.
• Heavier particles move more slowly than lighter particles at the same temperature. Then take a look at http://igcsechemistryrevision.blogspot.in/2013/03/chapter-2-states-of-matter-and.html for application of kinetic theory in changes in physical states of matter.

Diffusion

• This is the process by which different substances mix as a result of the random motion of their particles. This is the movement of particles from a region where they are in a higher concentration to region of lower concentration
• This eventually leads to the particles spreading out. to make the concentration same throughout.
• Diffusion only takes place in fluids (liquids and gases)
• The lower the molecular mass the faster the particle will diffuse. EG- Hydrogen gas has molecular mass of 2 (H2 =1*2=2) is more likely to diffuse faster than chlorine gas with molecular mass of 71  (Cl2 =35.5*2=71) since it is lighter.

Structure of Atom

• Atoms have three sub-atomic particles- proton, neutron and electron.

Sub atomic Particle	Relative Mass	Relative Charge	Location in atom
Proton	1	+1	In nucleus
Neutron	1	0	In nucleus
Electron	1/1840	-1	Outside nucleus in orbits

• Atomic number = Number of Proton
• Atomic Mass = Number of protons+ Number of neutrons=nucleons
• Number of Electrons = Number of protons = Atomic Number
• Number of neutrons = Mass number-Atomic number=nucleons-protons
• Isotopes- Atoms of the same element with different mass number. They have the same number of protons and electrons but different number of neutrons.They have the same chemical properties since they have the same number of electrons but only physical properties differ. Some isotopes maybe radioactive known as radio-isotopes.(Extremely important to remember this definition and properties of isotopes) 

Radioactivity

• This is the spontaneous decay of unstable radio-isotopes. It is unaffected by temperature or whether the isotope is part of a compound or present as the free element, it is a totally random process.
• There are three types radioactivity

1. Alpha Radiation
2. Beta Radiation
3. Gamma Radiation (no mass)

Radiation	Isotopes involved	Nature	Distance travelled in air	Penetration Paper        Thin     Thick                     Al         Lead
Alpha Radiation	Atomic Number >83	Alpha Particles Helium He42 (double positive charge, no electrons)	A few Centimeters	No	No	No
Beta Radiation	Atomic Number ≤ 83	Beta Particles (electrons single     –ve charge)	A few meters	Yes	No	No
Gamma Radiation	Only occurs with one of the other forms of radiation	Electromagnetic Radiation	Many Kilometers	Yes	Yes	No

Alpha and beta radiation knocks out electrons out of atoms (ionization) to produce positive ions.

Uses of Radioactivity (Important to learn these uses, can be asked )

Industrial

• Radioactive Dating which measures age of an object. 
• Monitors the level of filling in containers
• Checks thickness of sheets of metal or paper
• Detects leaks in gas or oil pipes

Medical

• To kill cancer cells using gamma radiation from cobalt-60
• To sterilize medical instruments, dressings and syringes by gamma radiation to kill bacteria.
• Food treatment to kill bacteria, mould and yeast by gamma radiation 

Electron arrangements in atoms

• Electrons in atoms are arranged in an organised way in different shells
• These electron shells have different electron shells are at different distances from the nucleus of the atom
• Electrons are placed in the shells closest to the nucleus first,and each shell has a maximum number of electrons it can contain. This could be calculated using the formula 2n^{2                     where n is the shell number. For example shell 1 can hold 2 electrons at maximum and shell 2 can hold maximum 8 electrons.}

• Noble gas arrangement is when the outershell is filled with maximum electrons therefore is unreactive since its already stable. These elements are known as noble gases like helium has 2 electrons in outershell and the first shell can only hold 2 electrons not more than that therefore is stable.

• The number of electrons in the outershell play a very important role in deciding the chemistry of the element.
• Atoms gain or lose electrons or share electrons to attain noble gas arrangement. Only the outershell electrons are involved in bonding.Will be discussed later. But remember this, commonly asked question why atoms do this? To attain noble gas arrangement to be stable.

Chapter 2: Separating and purifying substances

Separating heterogeneous mixtures

• Decantation- This is the process of removing a liquid from a solid (Like sand and water) which has settled (sedimented) or form an immiscible heavier liquid (like oil and water) by carefully pouring. It is important you learn this definition, asked in paper 6.
• Filtration- This is where an insoluble material is collected on filter paper- this is the residue. The liquid phase obtained is called the filtrate. It is useful since both phases can be obtained.
• Centrifugation- The separation of an insoluble solid from a liquid in a test tube by rapid spinning during which the solid collects at the bottom of the sample. The liquid then can be decanted carefully.   
• Separating Funnel- This is used to separate immiscible liquids like oil and water. The tap is opened to let the water out first. Then close tap when water is finished and change beaker and empty oil. Hence both liquids are separated.

• Magnetic Properties- Suppose we have a mixture of iron and plastic. We can hence separate iron from the plastic by using a magnet to attract all the iron pieces from the plastic.
• Solubility- If we have a mixture of salt and sand. We can take water which dissolves the salt and not the sand therefore separating the mixture.
• Sublimation-  If we have ammonium chloride and sodium chloride (Salt), ammonium chloride sublimes while salt doesn't therefore the mixture is separated.

Separation of Homogeneous Mixtures

• Evaporation and Crystallization- Used to separated solid in liquid mixtures (like salt and water) Evaporation is used if you want the powder of solid like salt. Evaporation is done by heated the solution on a evaporating dish, letting the liquid evaporate and solid is left. In crystallization the solution is concentrated by evaporating only some of the liquid in a water bath. When small crystals form on glass rod when dipped, it goes for cooling and crystals of solid is formed. The crystals are then filtered off and dried in oven or using filter paper (like tissue paper) 

     Evaporation

• Distillation- Used to obtain liquid from solution. Liquid is first evaporated and then condensed in the condenser. The liquid obtained is known as distillate.

• Fractional Distillation- This is often asked in the exam so pay attention. This is used to separate a mixture of liquids. The liquids have different boiling points hence can be separated. Suppose we have ethanol (alcohol) and water. Ethanol has a lower boiling point (78C) than water (100C) therefore gets boiled first. The water vapor keeps condenses in the fractional column since its below its boiling point.

• Chromatography- This is used to separate a mixture of solids dissolved in a liquid like coloured dyes  by differences in their solubility in a solvent. Please learn the procedure for chromatography since it is frequently asked in most of the papers it is extremely important (bold) 

1.  First draw a line on the chromatography paper (about 1cm form the bottom) with pencil. This is called the base line.
2. Place a drop of the solution on the base line.
3. Place the chromatography paper in a solvent and the solvent shouldn't cross the base line.
4. As solvent moves up the line the dyes with it start to separate into different colours.
5. If the solution is invisible like amino acids, use locating agent to locate the spots and warm the paper in the oven then spots will reveal.
6. Calculate Rf value using the formula = Distance moved by substance/ Distance moved by solvent front. The solvent front is where the solvent stops. The identity of a substance in the dye can be checked by comparing its Rf value to that of a sample we know is pure. Suppose Rf value of an amino acid is 0.31 (random) and the Rf value of a spot on the chromatography paper is 0.31 therefore the solution has an amino acid.
7. One spot means that the solution is pure.

Solubility

• Solution is made of solute (solid dissolved) and solvent ( the liquid in which the solid dissolves in).
• Solubility of solids in liquids increase with temperature. At a fixed temperature there will come a point where no more sold will dissolve in the solvent. This means that the solution is saturated. Temperature has to increase then to dissolve more of the solid.
• Solubility of gases in liquid decreases with an increase in temperature.
• Experiment to find solubility of solid in liquid (asked in p6 once so I recommend you to learn the procedure)

1. Take known mass of the solid using balance and take 100 cm3 of water using burette in beaker.
2. Heat the water ti the required temperature like 30C
3. Add the solid to the liquid and stir and keep adding till no more solid will dissolve in liquid and solution becomes saturated.
4. Filter out excess solid that didn't dissolve in liquid.
5. Then evaporate the liquid off and solid that dissolved remains.
6. Measure the mass of solid that dissolved in the liquid obtained in the evaporation.
7. Calculate solubility of solid at the fixed temperature using mass of solid dissolved (say in grams) divided by volume of liquid used (100cm3 in this example). Units will be g/100cm3 then.

Chapter 2 States of matter and substances

(Chapter 1 is not so important for IGCSE so we can skip that, it can be a little bit long since I added other necessary things to the notes which are required by the syllabus.)

States of Matter
Matter is defined as anything that occupies space and has mass.
There are three states of matter

• Solid
• Liquid
• Gas

Its really important to remember this table, the ones in bold are the key ones to remember and write in the exam to get marks.

Remember liquid and gas are known as fluids.

	Solid	Liquid	Gas
Description	Fixed volume, own shape	Fixed volume, takes shape of container	Any volume, takes shape of container
Arrangement of Particles	In a regular pattern called a lattice	Random	Random
Separation of Particles	Close together, touching	Still close together, just slightly further apart than in the solid pace	Separated, far apart.
Movement of Particles	Vibration about a fixed point	Slow movement in a random way from place to place, sliding other each other	Fast random movement
Attractive forces between particles	Stronger than in the liquid pace	Slightly weaker than the solid pace	No attractive forces between particles

All states of matter respond to a change in temperature.

• An increase in temperature will cause the particles in a solid to vibrate more in their fixed positions and take up more space causing expansion. 
• The reverse is contraction which happens when solid is cooled and therefore less energy in the particles, less vibration therefore less space taken up by the particles.
• When the temperature reaches to melting point then the particle have enough energy to move past each other and change positions and the solid changes into a liquid. This is known as melting.
• The reverse of melting is solidification where liquid is cooled and becomes solid. This happens at freezing point which is the same as melting point. Like in water  water melts or solidifies at 0.C

Melting point is the definite temperature where solid changes to a liquid.

• When a liquid is heated, the particle move faster and take up more space causing expansion. 
• When a liquid is cooled, the particles have less energy and move less faster therefore causing a contraction. 
• When the liquid is heated up to its boiling point, the particle in the liquid have enough energy to break intermolecular forces and so the particles separate and become gas or vapor. This is known as boiling. 
• Vapor is a gas which can be compressed into a liquid without cooling.
• The reverse of boiling is condensation where gas is cooled to become a liquid.

Boiling point is the definite temperature where liquid changes to a gas at a constant temperature. If the surrounding pressure decreases the boiling point also decreases.

 Evaporation though happens at all temperatures.

• For solids like dry ice (solid carbon dioxide) they don't turn into a liquid, they turn directly into a gas when heated, This is known as sublimation. 

Pressure

Only fluids (gases and liquids) are compressible, not solids. Liquids are only slightly compressible while gases can be easily compressed. 

In a sealed container the pressure of gases increases when temperature increases. This is because particles of gas will have more energy and will more faster and therefore will hit the walls of container more frequently and harder leading to an increase in pressure. 

This is a frequently asked question IGCSE, why does this happen, therefore it will be good if you understand the answer properly.

Pure Substances

A pure substance consists of only one substance and boils and melts at definite temperatures. Boiling and melting points can be checked to test the purity of a substance. Water that boils at 100C is pure while water that boils at 110C isn't pure.

Impurities in a substance cause-

• The melting point to decrease
• The boiling point to increase.

Mixtures

Mixtures are a system of two or more substances that can be separated by physical means.

It can be of two types-

• Homogeneous- This means the substances are totally mixed up and are indistinguishable like salt and water. Alcohol and water are miscible meaning they make a solution
• Heterogeneous- This means that the substances remain separate and one substance is spread throughout the other as small particle, droplets or bubbles like sand and water. Water and oil are immiscible and separate into different layers.