Chemistry revision

Unit 1 : Periodic table 

Periods : The rows on a periodic table , Groups : Columns on a periodic table.

Elements: substances that cant be broken down

Compounds : elements that are chemically bonded together 

Molecules : Two or more atoms that are bonded together 

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Chemistry revision

Unit 1 : Periodic table 

Periods : The rows on a periodic table , Groups : Columns on a periodic table.

Elements: substances that cant be broken down

Compounds : elements that are chemically bonded together 

Molecules : Two or more atoms that are bonded together 

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Atomic structure

Atoms are made up of protons , electrons & neutrons.

Protons are positively charged 

Electrons are negatively charged 

Neutrons have no charge 

The Neutrons and the protons make up the nucleus of an atom while the electrons orbit round the atom. 

Mass number : The number of protons & neutrons of an element

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Isotopes

Isotopes are atoms of the same element with the same number of protons but different number of neutrons .

E.g Carbon-12 , Carbon-13 and Carbon-14 

Calculating Relative Atomic Mass (R.A.M.)

Lets start this off with an example!

Example: Naturally occurring silver is 51.84% silver-107 and 48.16% silver-109. Calculate the relative atomic mass of silver.

r.a.m. (Ag) = (51.84/100 x 107) + (48.16/100 x 109)

= 55.469 + 52.494

=107.96

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Metals & Non-metals

Metals are normally shiny. Most of them have high melting and boiling points because of powerful attractions. Metals conduct heat and electricity because delocalized electrons are free to move throughout the structure. Metals are usually easy to shape due to their regular packed molecules. Metals react with water to form bases, and their oxides are also bases. They are good reducing agents because they lose electron

Non metals are normally brittle . Poor conductors of heat and electricity and the form acidic oxides. They are also good oxydising agents .

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Group 1 : Alkali Metals

They are the most reactive metals on the periodic table and their reactivity increases as you go down the group so Lithium being the least reactive and Francium being the most reactive.

Metal

Melting Point (0C)

Boiling Point (0C)

Density (g/cm3)

Lithium

181

1342

0.53

Sodium

98

883

0.97

Potassium

63

760

0.86

Rubidium

39

686

1.53

Francium

29

669

1.88

You can see that as reactivity increases, the melting and boiling points decreases; however, density increases. These points are very low for metals. Remember that potassium, sodium and lithium would float on water due to their densities. But why are they so reactive? Well they only have one electron to lose!

All these metals are extremely reactive. Anyways the metals will quickly react with air to form oxides, and react between rapidly and violently with water to form strongly alkaline solutions of  metal hydroxides.

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Alkali Metals : reactions with water

 Metal + Water à Metal Hydroxide + Hydrogen . All the hydroxides are bases and turn pH paper purple.

 With Sodium:  The sodium floats because it is less dense than water. It melts because its melting point is low and a lot of heat is produced by the reaction. Observations would be that the sodium would turn into a ball and whiz around the surface of the water. It may form a white trail which is sodium hydroxide. This dissolves to make a strongly alkaline solution with the water. When lit, it produces a yellow flame.

 With Lithium: The reaction is very similar to sodium’s reaction, except it is slower. The lithium does not melt due to its higher melting point. When lit, it produces a red flame.

 With Potassium:Potassium’s reaction is faster than sodium’s. Enough heat is produced to ignite the hydrogen, which burns with a lilac flame. The reaction often ends with the potassium spitting around.

 With Rubidium and Caesium – The Two Baddies: The reaction is so violent it can be explosive. When lit, Rubidium forms a red flame and Caesium forms a blue flame

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Alkali metals : Quick notes

§  Group One so +1 charge

§  One electron on outer shell

§  Reactivity increases downwards 

§  Density increases downwards

§  Melting and Boiling points both decrease downwards

§  Very soft and tarnish quickly in air

§  Li, Na and K are stored under oil, whilst Rb and Cs are stored in sealed glass tubes

§  Reacts with air to form oxides

§  Reacts with water to form alkaline hydroxides, which turns pH paper purple §  Positive ions are formed and they are colourless.§ The  Flame Colours: lithium, red; sodium, yellow; potassium , lilac; rubidium, red; caesium, blue

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Group 2 : Alkali earth metals

§  Harder than group one metals

§  Two electrons on outer shell (2+ charge)

§  Form white oxides

§  Forms hydroxides and hydrogen when reacting with water. Reaction is less vigorous than that of group one

§  Reaction increases downwards

§  Silvery-Grey

§  Flame Colours: calcium, brick red; strontium, crimson; barium, apple-green.

The metals are : Berrylium ,  magnesium, calcium, strontium, barium and radium

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Group 7 : Halogens

Halogens: Quick Notes

 

§  Diatomic molecules

§  Seven electrons on outer shell

§  Highly reactive – only need one electron to fill outer shell

§  Form hydrogen halides when reacting with hydrogen

§  Reaction increases as you go up the group

§  Halogens can displace each other

§  Volasil turns iodine pink

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Oxygen & Oxides

Composition of air : Nitrogen - 78% , Oxygen-21% , Argon-0.9% , Carbon-dioxide - 0.04%

Showing air contains 21% of oxygen : The apparatus originally contains 100cm3 of air. This is pushed backwards and forwards of the heated copper, which turns black as copper(II) oxide is formed. This uses up the oxygen. On cooling, around 79cm3 of gas is left in the syringes – 21% has been used up. Therefore, the air contains 21% of oxygen so copper reacted with all the oxygen in the air .

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Methods of seperation

Filtration: Separating a liquid from a solid 

Chromatography: separating two liquids 

Crystallization :Mainly used for purifying substances by forming crystals from a precipitating solution. Crystallization refers to the forming of solid crystals from a homogenous (solution) mixture.

 Distillation : Distillation is good from separating a liquid from a solution. e.g : water from a salt solution

Fractional distillation : Fractional Distillation is used to separate two liquids based on their boiling points.

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Structure and Bonding : Ionic bonding

Ionic bonding is then transfer of electrons from one atom to the other to produce ions. 

One that gains electrons form negative ions(anions)

One that loses electrons form positie ions (cations)

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Ionic bonding : Boiling and melting points

Ionic compounds have high boiling and melting points because they have strong intermolecular forces between the atoms. Tis is because one of them is negative and the other is positive so they attract which creates a strong bonds 

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Structure of ionic compounds

Ionic bonds always produce giant structures. Ions form closely packed regular lattice arrangement. They have high melting/boiling points. The crystals tend to be brittle. Compounds tend to be soluble in water and insoluble in organic solvents.

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Covalent bonding

Covalent bonding is formed when two atoms share electrons to form a full outer shell . The reactants are normally non-metals.

<strong>      These are gases, liquids or solids with low melting points. Examples include water, chlorine, oxygen…etc

</strong>      The covalent bonds between the atoms in a molecule are strong.

<strong>      However, the forces of attraction between these molecules (inter-molecular forces) are weak.

</strong>      They tend to be insoluble in water.

*      They do not conduct electricity because the molecules have no overall charge and there are no electrons mobile enough to move from molecule to molecule.

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Giant covalent structures

<strong>      There are no charged ions.

</strong>      ALL the atoms are joined up to their adjacent atom by extremely strong covalent bonds and packed into giant regular lattices.

<strong>      They have very high melting points, since a lot of heat is needed to provide the energy to break apart the many strong covalent bonds.

</strong>      They tend to be insoluble in water.

They do not conduct electricity


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Giant covalent structures : Diamond

The diamond is the hardest natural substance. It is a form of pure carbon. Each carbon atom forms four covalent bonds to the other carbon atoms. They are arranged in a tetrahedral arrangement. Diamond has a very high melting point, obviously due to very strong carbon-carbon bonds. It does not conduct electricity because all the electrons in the outer levels of the carbon atoms are tightly bonded between the atoms. None of them are free to move around. Diamond is insoluble – like, to both water and other solvents.

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Graphite

Graphite is arranged differently – it has a layer structure. Each graphite layer is strong, but it is easy to separate individual graphite layers. Each carbon atom only forms three covalent bonds. Graphite conducts electricity because the fourth electron is free to move around

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Metallic crystals

Metals are giant structures which consist of a regular array of positive ions in a sea of delocalized electrons. When metal atoms bond together to form solid, visible metal, their outer electrons are no longer attached to particular electrons and are free to move around the whole structure.

 

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Hydrocarbons

ü  Hydrocarbon – compounds that contain carbon and hydrogen only.

ü  Homologous series family of compounds with similar properties because they have similar bonding. They show a graduation in physical properties (mpt/bpt) and similar chemical properties such as the general formula. Alkanes are the simplest.

ü  Saturated – when carbon cannot take anymore bonds – single carbon-carbon bonds.

ü  Unsaturated – presence of a carbon-carbon double bond.

ü  General formula – The formula of different homologous series of carbons.

ü  Isomers – molecules with the same molecular formula but different structural formulae.

 

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Combustion

If there is enough oxygen, alkanes will burn in oxygen completely to give carbon dioxide and water. The general equation for combustion:

 

Hydrocarbon + Oxygen à Carbon dioxide + Water

 

The combustion of methane would be: CH4(g) + 202(g) à C02(g) + 2H20(l)

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Reaction with bromine

Alkanes react with bromine under the presence of ultra-violet light. One hydrogen from the hydrocarbon would be replaced by a bromine atom. This is known as a substitution reaction. Bromine can be used as an indicator for alkanes and alkenes without UV light. Adding bromine water to alkanes produces no colour change. Reacting bromine water to alkenes make it turn from brown to colourless.

 

However, if the mixture of bromine and methane is reacted under UV light, it loses its colour, a mixture of bromomethane and hydrogen bromide gases is formed.

 

CH4(g) + Br2(g) à CH3Br (g) + HBr(g)

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Alkenes

Alkenes have double bonds, making them unsaturated hydrocarbons.

 Alkenes have the general formula of CnH2n – the first four being ethene, propene, butene and pentene.

 Combustion

 Like alkanes, alkenes burn in oxygen or air to give carbon dioxide and water.

 Reaction with Bromine

Alkenes undergo addition reactions, in which part of the double bond breaks and is used to join other atoms onto the two carbon atoms. When added to alkenes, and the test tube is shook, the brown of the bromine would be decolourised, making it suitable as a test for alkenes.

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Fractionating column

Crude oil is separated in fractionating column. This process is fractional distillation, and splits crude oil into various fractions depending on their boiling points and size.

 

Refinery gases

<strong>      A mixture of methane, ethane, propane and butane.

</strong>      Commonly used for domestic heating and cooking.

Gasoline

<strong>      Cars

Kerosene

</strong>      Used as fuel for jet aircraft.

<strong>      As domestic heating oil.

</strong>      As ‘paraffin’ for small heaters and lamps.

Diesel oil

ü  For buses, lorries, some cars and railway engines.

ü  Some is cracked to produce more petrol.

Fuel oil

ü  For ships

ü  Industrial heating

Bitumen

ü  Residue from the bottom which can be used for roads.

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Cracking

Cracking is a useful process in which large hydrocarbon molecules are broken into smaller ones. Most of the hydrocarbons found in crude oil are long-chain alkanes. Cracking can convert these into alkenes and shorter alkanes. It is an example of thermal decomposition.

 How it Works

The fraction is heated to give a gas and is passed over a catalyst of silica or alumina with a temperature of 600-700oC.

 Long alkane à alkene + alkane

 

ü  Sometimes you may get more than one type of alkene/alkane.

Make sure the numbers of carbon and hydrogen are balanced

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Calculations

Calculating Relative Atomic Mass

 Chlorine has two isotopes: chlorine-35 and chlorine-37. A typical sample will be 75% chlorine-35 and 25% chlorine-37.

 The RAM = (0.75 x 35) + (0.25 x 37) = 35.5g

 

Calculating Relative Formula Mass

 

H2

H x 2

2 x 1 = 2

Ca(OH)2

(1 x Ca) + (2 x O) + (2 x H)

40 + (2 x 16) + (2 x 1) = 74

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Calculations using moles

 a)     Firstly, convert the grams of Na into moles:

2.3 / 23g = 0.1 mol

The equation says that 2 moles of Na and 1 mole of Cl (1 mole of a diatomic molecule is always X2) is needed to produce 2 moles of NaCl, so if 0.1 mol of Na is used, then half of that is the amount of chlorine used in the reaction in moles.

 So moles of Cl used = 0.1 / 2 = 0.05 mol

 b)     One mole of any gas has a volume of 24 dm3 (24000cm3) at room temperature and pressure. This is also called the molar volume.

 Cl2 is a gas and the moles used in the reaction = 0.05 mol ,So the volume of Cl2 gas used = 0.05 x 24000 = 1200cm3

 c)     The moles of NaCl produced is 0.1 mol (if 2 moles of Na gives 2 moles of NaCl, then 0.1 mole of Na will give 0.1 mole of NaCl). So all you do is:

i)              Find the RFM of NaCl (58.5) ii)             Multiply that by 0.1 (5.85g)

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Calculations using moles : question

Calculations using moles: The equation for sodium chloride is:

2Na + Cl2 à 2NaCl If 2.3g of Na was used:

a) Find out how many moles of Chlorine was used

b) Find out the volume of Chlorine used in the reaction

c) Find out the mass of sodium chloride produced

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Emperical formula

Mass formula -->

Percentage formula -down

 

Mg

O

Combining Masses

2.4

1.6

Number of moles

2.4/24

1.6/16

=

0.10

0.10

Ratio of Moles

1:1

Empirical Formula

MgO

 

C

H

Percentage

87.5

14.3

Combining Masses

87.5

14.3

Number of moles

85.7/12

14.3/1

=

7.14

14.3

Ratio of Moles

1:2

Empirical Formula

CH2

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Molecular formula

To find the molecular formulae, you need to know the relative formula mass of the compound. Suppose it was 56g for the above question.

 

Firstly, find out the RFM of CH2 = 12 + 2 = 14g

Find out how many times 14 goes into 56, so 56/14 = 4 times

Which means the molecular formula is C4H8!

 

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