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CHEMISTRY


SPECIFICATION EXPLAINED

words and lyrics by Black Jerome


with a lot of help from http://hannahhelpchemistry.blogspot.co.uk/


and BBC GCSE bitesize



i just wanted the notes on flash cards :D


8=====D () 8===(=)=D

Chapter 1 Kinetic Theory and Diffusion


1.1 understand the arrangement, movement and energy of the particles in each of the three states of matter: solid, liquid and gas


1.2 understand how the interconversion of solids, liquids and gases are achieved and recall the names used for these interconversions


1.3 explain the changes in arrangement, movement and energy of particles during these interconversions

1.4 describe and explain experiments to investigate the small size of particles and their movement including: i dilution of coloured solutions ii diffusion experiments

i


Dilution: a substance is put in a solvent to reduce its concentration e.g. Orange squash mmmmmm

ii diffusion experiments

Diffusion: The movement of particles from an area of high concentration to an area of low concentration e.g Diffusion of ammonia and hydrochloric acid in air to form ammonium chloride. Put at either ends of a tube and a ring forms in the tube where the two meet - Ammonium Chloride. Ring forms closer to HCl because ammonia is lighter and moves quicker.

Chapter 2 Atomic Structure


1.5 understand the terms atom and molecule


1.9 Understand that atoms consist of a central nucleus, composed of protons and neutrons, surrounded by electrons, orbiting in shells

Atom is made up of a nucleus (protons and electrons) and orbital electrons in shells.
Molecule is 2 or more atoms bonded together. e.g H2O

Atom is made up of a nucleus (protons and electrons) and orbital electrons in shells.


Molecule is 2 or more atoms bonded together. e.g H2O

1.10 recall the relative mass and relative charge of a proton, neutron and electron

Relative mass of electron is 1/1840 roughly equal to 0

Relative mass of electron is 1/1840 roughly equal to 0

1.11 understand the terms atomic number, mass number, isotopes and relative atomic mass

Atomic number: number of protons (the same as number of electrons.)
Atomic mass: number of protons + neutrons.
Isotopes: Atoms of an element with different numbers of neutrons.
Relative atomic mass (Ar): the mass of one atom of an element in proportion to carbon-12

1.13 understand that the Periodic Table is the arrangement of elements in order of atomic number

Just look m8. Clearly arranged in atomic number

Just look m8. Clearly arranged in atomic number

1.14 deduce the electronic configurations of the first twenty elements from their positions in the Periodic Table


1.15 deduce the number of outer electrons in a main group element from its position in the periodic table

Groups are columns in the periodic table; the group number represents the number of electrons on the outer shell.
Periods are rows on the periodic table; the row corresponds with number of shell.


e.g. Aluminium - Group 3, 3 electrons in outer shell Period 3, 3 shells.


Electronic Configuration: 2, 8, 3


Chapter 3 Bonding


1.5 understand the terms atom (Chapter 2) and molecule

Molecule is 2 or more of the same type of atom  COVALENTLY bonded together. e.g S8. H2O has H2 and is also a molecule as well as a compound.

Molecule is 2 or more of the same type of atom COVALENTLY bonded together. e.g S8. H2O has H2 and is also a molecule as well as a compound.

1.28 describe the formation of ions by the gain or loss of electrons

Atoms with 1,2,3 electrons in their outer shell will lose them to become cations - pussytive.


Atoms with 5,6,7 electrons in their outer shell will gain move electrons to become anions - negative.


Atoms need to have a complete outer (valence) shell to achieve electronic stability.


Any charged particle is an ion

1.31 deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed

Aluminium is in group 3 and has an electronic configuration of 2,8,3.


Al will lose its 3 electrons when it becomes an ion.


Al is a cation and has a positive charge of 3+ because it has 3 more protons than electrons.


Al3+

1.32 explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combination of elements from Groups 1,2,3 and 5,6,7



IONIC BONDING DOT AND CROSS DIAGRAM

Mg lost 2 electrons to complete its outer shell.
Oxygen gained 2 electrons to complete its outer shell.
Mg2+ and O2- are ions in MgO

Mg lost 2 electrons to complete its outer shell.


Oxygen gained 2 electrons to complete its outer shell.


Mg2+ and O2- are ions in MgO

1.33 understand ionic bonding as a strong electrostatic attraction between oppositely charged ions


2 ions are bonded together ionic bonding. The opposite charges (due to the loss or gain of electrons between the two ions) creates a strong electrostatic attraction.

1.38 describe the formation of a covalent bond by the sharing of a pair of electrons

Covalent bond is formed by the sharing pairs of electrons (one from each atom) and the strong attraction between the nuclei.

1.39 understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond

The bonding pair of electrons are negatively charged and so are attracted to the nuclei which are positively charged.

1.40 explain, using dot and cross diagrams, the formation of covalent compounds by electron sharing for the following substances: hydrogen, chlorine, hydrogen chloride, water, methane, ammonia, oxygen, nitrogen, carbon dioxide, ethane, ethene

Covalent bonding is drawn like this. Notice: NH3  has 3 hydrogens. :D. Hydrogen or H2 would have 2 hydrogen atoms with one shared pair of electrons. Oxygen or O2 would have 2 O atoms with two shared pairs of electrons.

Covalent bonding is drawn like this. Notice: NH3 has 3 hydrogens. :D. Hydrogen or H2 would have 2 hydrogen atoms with one shared pair of electrons. Oxygen or O2 would have 2 O atoms with two shared pairs of electrons.

1.46 understand that a metal can be described as a giant structure of positive ions surrounded by a sea of de-localised electrons


Electrons become de-localised and free to move around the stationary now positive ions.

Electrons become de-localised and free to move around the stationary now positive ions.

Chapter 4 Structure -this chapter looks like effort.


1.7 understand the differences between elements, compounds and mixtures



Element - Substance made from only one type of atom (every atom has the same amount of protons).


Compound - Substance made from two or more different elements that have chemically bonded together.


Mixture - Contains 2 or more substances (elements or compounds) which are not chemically bonded together.


1.38 understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions

When melting or boiling anything, heat is needed to break bonds. Stronger bonds means more heat is needed. Ionic compounds have strong bonds due to the strong electrostatic forces between oppositely charged ions. Therefore, more heat is needed to break these bonds, so, ionic compounds have high melting and boiling points.

1.39 understand the relationship between ionic charge and the melting point and boiling point of an ionic compound.

Greater charges have stronger forces of attraction. e.g. + and - ions has a weaker bond than +3 and -3 ions. As stated the stronger the bonds, the more heat is needed to break them and so the greater the ionic charges, the higher the melting and boiling points of the compound.

1.40 describe an ionic crystal as a giant three-dimensional lattice structure held together by the attraction between oppositely charged ions



1.41 draw a diagram to represent the positions of the ions in a crystal of sodium chloride

The oppositely charged ions are held together by the electrostatic forces of attraction. Structure extends in all directions in a regular arrangement forming a 3D lattice. It is giant because a large number of ions are present

The oppositely charged ions are held together by the electrostatic forces of attraction. Structure extends in all directions in a regular arrangement forming a 3D lattice. It is giant because a large number of ions are present

1.46 understand that substances with simple molecular structures are gases or liquids, or solids with low melting points



1.47 explain why substances with simple molecular structures have low melting and boiling points in terms of the relatively weak forces between the molecules

Simple molecular structures are a few atoms with strong covalent bonds e.g. CO2. They have low melting and boiling points because the weak intermolecular forces break down easily when heated.

1.48 explain the high melting and boiling points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds

There are many bonds because there are many atoms in a giant covalent structure. A lot of strong covalent bonds have to be broken to melt or boil a giant covalent structure. Therefore, these structures must have very high melting and boiling points.

1.49 draw diagrams representing the positions of the atoms in diamond and graphite

Diamond, each carbon atom is bonded to four more carbon atoms.  Only made of carbon.The structure continues in all directions. 3D.

Diamond, each carbon atom is bonded to four more carbon atoms. Only made of carbon.The structure continues in all directions. 3D.

1.49 draw diagrams representing the positions of the atoms in diamond and graphite

Carbon atom arranged in layers of hexagons. Each carbon is bonded to three more carbons. Only made of carbon. 

Carbon atom arranged in layers of hexagons. Each carbon is bonded to three more carbons. Only made of carbon.

1.50 explain how the uses of diamond and graphite depend on their structures, limited to graphite as a lubricant and diamond in cutting

Graphite has layers that can slide over each other easily so can be used as lubricants.



Diamond has lots of strong covalent bonds with each carbon atom bonding to four more. Therefore, diamond is very hard and ideal for cutting.

1.48 explain the electrical conductivity and malleability of a metal in terms of its structure and bonding.

Metals have delocalised electrons that are free to move. The electrons carry electricity and so charge passes easily through a metal.


A metal has rows of atoms on top of one another, in pure metals, all the atoms will be the same size and so the metals slide easily over one another.

Chapter 5 Formulae and Equations


1.21 write word equations and balanced chemical equations to represent the reactions studied in this specification

Word equations use the names of the products and reactants in the reaction.


Balanced chemical equations use the chemical symbols for the products and reactants including their quantities. If the quantities aren't equal then they have to be balanced.


e.g. 2H2 + O2 ----> 2H2O


1.22 use the state symbols (s), (l), (g) and (aq) in chemical equations to represent solids, liquids, gases and aqueous solutions respectively

(s) solid


(l) liquid


(g) gas


(aq) aqueous solution-dissolved in water


These go after the element or compound to show what state it's in.

1.30 recall the charges of common ions in this specification.

K+ Na+ Li+ H+ Mg 2+ Ca 2+ Al 3+ Cl- Br- I- F-


These can be told from the periodic table.
OH- NO3- SO4 2- CO3 2-


Easier if you learn these.

Chapter 6 Rates of Reaction


4.17 describe experiments to investigate the effects of changes in surface area of a solid, concentration of solutions, temperature and the use of a catalyst on the rate of a reaction



Surface Area

* Put a set mass of magnesium in hydrochloric acid
* Time the reaction
* Change the from of magnesium keeping the mass the same (powder, wire, strips)

The more surface area (the smaller the pieces of magnesium) the faster the reaction

Concentration of solutions

* Put a set mass of marble chips into dilute hydrochloric acid
* Time the reaction
* Change the ratio of water to hydrochloric acid

The more concentrated the hydrochloric acid (the lower the ratio of water) the faster the reaction

Temperature

Temperature

* Put a set mass of magnesium powder into a set mass of hydrochloric acid
* time the reaction
* Carry out this reaction at different temperaturesThe higher the temperature the faster the rate of reaction

Use of a catalyst

* If you have hydrogen peroxide it will not decompose
* If you put it with manganese dioxide it will decompose into water and oxygen
* The manganese dioxide will be unaltered by the reaction

The more of the catalyst the faster the reaction

4.18 describe the effects of changes in surface area of a solid, concentration of solutions, pressure of gases, temperature and the use of a catalyst on the rate of reaction

Higher temperature, more surface area, higher concentration, higher pressure and use of a catalyst all make a reaction faster.


And vice versa.

4.19 understand the term activation energy and represent it on a reaction profile

Activation energy is the amount of energy required for a reaction start happening.

Activation energy is the amount of energy required for a reaction start happening.

4.20 explain the effects of changes in surface area of a solid, concentration of solutions, pressure of gases and temperature on the rate of a reaction in terms of particle collision theory



Surface area

Collision theory says that to react particles must: Collide with enough energy to react Collide in the right orientation to react (the more frequent the collisions, the more likely this is)


Surface area: Particles collide more frequently if there is more surface area, as there is more contact between the reactants. Faster rate of reaction.



Concentration of solution and Pressure of Gases

* There is more chance of particles colliding at a higher concentration/pressure, so they react more often. Faster rate of reaction.

Temperature

* Particles move about more and will collide more frequently the higher the temperature; react more often. Increases the rate of reaction.

Catalyst

* Provides an alternative pathway for the reaction to start which requires a lower activation energy.

4.21 explain that a catalyst speeds up a reaction by providing an alternative pathway with lower activation energy

A catalyst provides an alternative route for the reaction to start, this route requires less energy to start the reaction.

Chapter 7 Oxygen and Oxides


2.16 recall the gases present in air and their approximate percentage by volume

Nitrogen - 78%


Oxygen - 21%


Argon - 0.9%


CO2 0.037%


And some others that you don't need to know

2.17 describe how experiments involving the reactions elements such as copper, iron and phosphorus with air can be used to determine the percentage by volume of oxygen in air



Experiment on back works for phosphorus and iron as well because they react with oxygen too.


2Cu + O2 ---> 2CuO

By passing 100cm^3 of air through a gas syringe with copper in it, all the oxygen in the air reacts with the copper to form copper oxide. Measure the new volume of gas and subtract that volume from 100cm^3 to get the % of oxygen in air.

By passing 100cm^3 of air through a gas syringe with copper in it, all the oxygen in the air reacts with the copper to form copper oxide. Measure the new volume of gas and subtract that volume from 100cm^3 to get the % of oxygen in air.

2.18 describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese (IV) oxide as a catalyst

Manganese (IV) oxide is unchanged during this reaction. When hydrogen peroxide is added, oxygen and water form. Oxygen is collected via downward displacement.
2H2O2 -->2H2O + O2

Manganese (IV) oxide is unchanged during this reaction. When hydrogen peroxide is added, oxygen and water form. Oxygen is collected via downward displacement.


2H2O2 -->2H2O + O2

2.19 describe the reactions of magnesium, carbon and sulphur with oxygen in air, and the acid-base character of the oxides produced

Non metals: Carbon and sulphur react with oxygen to form a non-metal oxide. These are acids. Acid rain is produced by sulphur dioxide. The reaction gives out heat and light.



Metals: Magnesium will burn in air to form a metal oxide. This is basic or alkali. Burns with a bright, white flame.

2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid

calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxideCaCO3 + 2HCl → CaCl2 + H2O + CO2
 

calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
CaCO3 + 2HCl → CaCl2 + H2O + CO2


2.22 describe the properties of carbon dioxide, limited to its solubility and density

Denser than air


Soluble in water at high pressure

2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density

Carbon dioxide is dissolved into drinks at a high pressure, this makes CO2 bubbles in fizzy drinks.


Some fire extinguishers have CO2 in, because it is denser than air it will fall over the fire creating a barrier between the air and fire. This starves the fire of oxygen because CO2 doesn't burn and so the fire is put out.

2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change

Carbon dioxide prevents heat leaving the earth's atmosphere in rays that the earth emits.
Significant amounts of green house gases will warm up the earth, changing the climate.

5.11 recall that, in car engines, the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming nitrogen oxides

In car engines there is a high enough temperature to cause a reaction between oxygen and nitrogen in the air.
This makes NO or Nitrogen Oxide.

5.12 understand that nitrogen oxides and sulphur dioxide are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain

NO and SO2 are given off into the atmosphere by some industrial processes. They are pollutant gases.
When they are in the atmosphere they react with rain water to create H+ ions. H+ ions are acidic.
When the rain falls the acid can corrode rocks and buildings. Acid can also alter the PH in soil or rivers which can effect an ecosystem.
Also acid rain corrodes limestone, which damages buildings and stuff.


Chapter 8 The Reactivity Series


1.29 understand oxidation as the loss of electrons and reduction as the gain of electrons

OILRIG


Oxidation Is Loss


Reduction Is Gain



2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron

acid + metal > salt + hydrogen
For example:
magnesium + hydrochloric acid > magnesium CHLORIDE + Hydrogen
Mg + 2HCl > MgCl2 + H2



zinc + sulphuric acid > zinc sulphate + hydrogen


Zn + H2SO4 > ZnSO4 + H2

2.29 understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds :potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold

Potassium, sodium, lithium, calcium, magnesium, aluminium, CARBON, zinc, iron, tin, lead, HYDROGEN, copper, mercury, silver and gold, platinum.



Private Second-class MacZitl he can make some gun powder


Has some unneeded elements but i like the anagram.



Potassium is the most reactive and gold is the least reactive.

2.30 describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper


Potassium, sodium, lithium and calcium all react with water and acids
magnesium, zinc and iron all react with acids (and very slowly with water.)
copper doesn't react with either.

The more vigorous the reaction the more reactive the metal. The more things a metal will react with, the more reactive the metal
.

2.31 deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions

With a metal oxide or a metal salt in an aqueous solution:


If a more reactive metal is added, it will displace the current one.


If a less reactive metal is added, there will be no displacement.


From this you can see which metals are more and less reactive.

2.32 understand oxidation and reduction as the addition and removal of oxygen respectively

Oxidation is the gain of oxygen.


Reduction is the loss of oxygen.

2.33 understand the terms :redox, oxidising agent and reducing agent

In a redox reaction, a more reactive metal gains an oxygen from a less reactive metal which loses it. Both reduction and oxidation take place.

The reducing agent is the more reactive metal which reduces the other metal.
The oxidising agent is the less reactive metal which allows the other metal to be oxidised.

Chapter 9 Acids


2.11 understand the difference between hydrogen chloride gas and hydrochloric acid

Hydrogen chloride gas is not acidic because it hasn't split its ions.


Hydrochloric acid is the gas dissolved in water. When this happens the HCl dissociates into its H+ and Cl- ions. H+ is acidic and so hydrochloric acid is acidic.

2.12 explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene

HCl dissociates in water to give H+ and Cl- ions. The H+ ions are acidic making hydrochloric acid..



HCl does not dissociate in methylbenzene, so it doesn't split into its atoms. There are no H+ ions and so it isn't acidic.

2.26 describe the combustion of hydrogen

The combustion of hydrogen is its reaction with oxygen.
Water is created. and a lot of energy.
2H2 + O2 > 2H2O

4.1 describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions


Red litmus paper turns blue in the presence of alkali. (ReBAL)
Blue litmus paper turns red in the presence of acid. (BaRACk Obama)
Phenolphthalein goes pink in alkalis.
Methyl orange is yellow for alkali, red for acid.

4.2 understand how the pH scale, from 0-14, can be used to classify solutions as strongly acidic, wealky acidic, neutral, weakly alkaline or strongly alkaline


4.3 describe the use of universal indicator to measure the approximate pH value of a solution

Acids are from 0-6.
The more red the more acidic.
Alkalis are from 8-14.
The more purple the more acidic.
7 is neutral.
These colours are from universal indicator.

Acids are from 0-6.


The more red the more acidic.


Alkalis are from 8-14.


The more purple the more acidic.


7 is neutral.


These colours are from universal indicator.

4.4 define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions, OH-

Essentially if something is acidic it contains H+ ions, if something is alkaline it contains OH- ions.


4.5 predict the products of reactions between dilute hydrochloric, nitric and sulfuric acids: and metal oxides and metal carbonates (excluding the reactions between nitric acid and metals)

Metal oxides
HCl + metal oxide > metal chloride salt + water


Nitric acid + metal oxide > metal nitrate salt + water


Sulphuric acid + metal oxide > metal sulphate + water


Metal carbonates


HCl + metal carbonate > metal chloride salt + water + CO2


Nitric acid + metal carbonate > metal nitrate salt + water + CO2


Sulphuric acid + metal carbonate > metal sulphate + water + CO2

Chapter 10 Making Salts


4.6 understand the general rules for predicting the solubility of salts in water

i) all common sodium, potassium and ammonium salts are soluble


ii) all nitrates are soluble


iii) common chlorides are soluble, except silver chloride


iv) common sulfates are soluble, except those of barium and calcium


v) common carbonates are insoluble, except those of sodium, potassium and ammonium


4.7 describe experiments to prepare soluble salts from acids

Dilute sulphuric acid is added to an excess of magnesium.
Mg + H2SO4 > MgSO4 + H2
The left over magnesium is filtered off and the mixture boiled down slowly to concentrate.
When it is cooled, crystals will form, these can be blotted dry with a piece of paper.


This can be done with other metals and acids:


HNO3, H2SO4 and HCl make soluble salts with most metals.


Ammonium, potassium and sodium make soluble salts with acids

4.8 describe experiments to prepare the salts using precipitation reactions

Silver nitrate and sodium chloride are added together, the product, silver chloride is made, this salt is insoluble and so will form a white precipitate in the solution. AgNO3 + NaCl > AgCl + NaNO3


Method


Mix two solutions together
Filter the mixture
Wash with distilled water
then set aside to dry and crystallize

4.9 describe experiments to carry out acid-alkali titrations

Put a known volume of acid in a beaker with methyl orange, red.


Set up a burette with alkali in it. Open the tap very slightly, so that it drips very slowly into the acid, stirring the beaker with each drop.


The more alkali that is added the more neutral orange it gets.


When the solution is neutral it will be completely orange, at this point close the tap on the burette.


The level in the burette will have dropped, showing the volume of alkali used. This volume is needed to neutralise the acid.

Chapter 11 Separating and Analysing


1.7 describe experimental techniques for the separation of mixtures, including simple distillation, fractional distillation, filtration, crystallisation and paper chromatography



Distillation

One substance is evaporated off.e.g salt water is heated to 100 degrees; water evaporates off (it rises and the goes down into the condenser where it is cooled back into water), the salt is left in the original flask. See diagram.

One substance is evaporated off.
e.g salt water is heated to 100 degrees; water evaporates off (it rises and the goes down into the condenser where it is cooled back into water), the salt is left in the original flask. See diagram.

Filtration

This consists of a barrier which one component of a mixture can pass through but the other is caught by.
e.g water goes through filter paper, rocks are caught by it.


Fractional Distillation

Works like distillation but separates substances with different boiling points because they will condense at different heights, splitting the substances into their different fractions

Works like distillation but separates substances with different boiling points because they will condense at different heights, splitting the substances into their different fractions

Crystallisation

A solution is warmed allowing the solvent to evaporate, the solution is now left to cool and will form crystals.

Chromatography

chromatography paper is placed in a substance, the different components of the substance will travel at different speeds (due to the size of their particles.)

1.8 explain how information from chromatograms can be used to dientify the composition of a mixture

As the solvent moves up the paper, it carries the mixtures with it.  Different components of the mixture will move at different rates, separating the mixture out.

As the solvent moves up the paper, it carries the mixtures with it. Different components of the mixture will move at different rates, separating the mixture out.

2.27 describe the use of anhydrous copper (II) sulfate in the chemical test for water

Anhydrous copper sulphate will become hydrous copper sulphate when it is reacted with water.
So if anhydrous copper sulphate goes from white to blue in the presence of a liquid it will be water. ACWB sounds like aquifer

2.28 describe a physical test to show whether water is pure

If water is pure it will boil at exactly 100° and freeze at exactly 0°

2.37 describe tests for the cations


i) Li+, Na+, K+, Ca2+ using flame tests


* Lithum: red
* Sodium: orange (so strong can mask other colours)
* Potassium: lilac


Calcium: brick red

ii) NH4+, using sodium hydroxide solution and identifying the ammonia evolved

* NH4 + OH > NH3 + H2O
* ammonium ions + hydroxide ions > ammonia + water


ammonia (pungent smelling gas) turns red litmus paper blue


(ARB)

iii) Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution

Copper (II) : blue precipitate
Iron (II): green precipitate (2 ee's in green)
Iron (III): brown precipitate

2.38 describe tests for the anions


i) Cl-, Br- and I-, using dilute nitric acid and silver nitrate solution


Chloride ions + nitric acid + silver nitrate > white precipitate (silver chloride) (white anyway)
* Bromide ions + nitric acid + silver nitrate > cream precipitate (silver bromide) (creaMMM broMMMiNNNEEE)


Iodide ions + nitric acid + silver nitrate > yellow precipitate (silver iodide) (io is a yellow moon)

ii) SO4 2- (sulphate ions) using dilute hydrochloric acid and barium chloride solution




iii) CO3 2-, using dilute hydrochloric acid and identifying the carbon dioxide evolved

SO4(2-) + HCl + Ba(2+) > white precipitate (barium sulphate) (like a polar bear)



Carbonate + acid > salt + water + carbon dioxide
Carbon dioxide produced will turn lime water cloudy

2.39 describe tests for the gases


i hydrogen


ii oxygen


iii carbon dioxide


iv ammonia


v chlorine

Hydrogen- burns with a 'squeaky pop' sound


Oxygen- Will relight with a glowing splint


Carbon dioxide- Turns lime water cloudy


Ammonia- Damp red litmus paper blue (ARB)


Damp universal indicator purple (Alkali)


Chlorine- bleaches damp litmus paper white (chlorine is white anyway)



Chapter 12 The Periodic Table


2.1 understand the terms group and period

Groups are columns in the periodic table, the number of a group represents the number of electrons on an atom's outer shell.
Periods are rows on the periodic table, the row represents the number of orbitals (electron shells) that a atom has.

2.2 recall the positions of metals and non-metals in the Periodic Table

Metals on the left.
Non-metals on the right.

Metals on the left.


Non-metals on the right.

2.3 explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides

All metals are conductors. Metals form metal-oxides which are alkaline.

Non-metals don't conduct. They form non-metal-oxides which are acidic.

2.4 understand why elements in the same group of the Periodic Table have similar chemical properties

Elements in the same group have the same number of electrons on their outer shell.
This means they will behave in a similar way; they will react and bond similarly.
E.g. in group one, all the elements form +1 ions as they each lose one electron to become stable (to have a full outer shell).

2.5 understand that the noble gases (Group 0 or 8) are a family of inert gases and explain their lack or reactivity in terms of their electronic configurations

Nobel gases are inert, this means they do not react. The reason for this is because they are stable: meaning they have a full outer shell, so they do not need to lose or gain electrons.

Group 1 elements - lithium, sodium and potassium


2.6 describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements

The group one elements- lithium, sodium, potassium- are easily identifiable as the same group due to the fact that they all react vigorously with water. (Same number of electrons on the outer shell).



The reactions become more vigorous as you go down the group. Hydrogen gas is produced as well as metal hydroxide.

Group 1 elements - lithium, sodium and potassium


2.7 describe the relative reactivities of the elements in Group 1

Group one elements get more reactive the further down the group.

Group 1 elements - lithium, sodium and potassium


2.8 explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus

Group one elements need to lose one electron on the outer shell to react. Electrons are held to an atom by the protons in the nucleus. If an electron is close to the nucleus the force holding it in will be very strong, if it is further away it will be weaker.


Group 1 is more reactive down the group
So bigger atoms (towards the bottom of the group) with the outer orbital far from the nucleus will lose their electron more easily: this means they react more easily/quickly/more/vigorously.

Group 7 elements - chlorine, bromine and iodine


2.9 recall the colours and physical states of the elements at room temperature

Fluorine.............. Gas........... yellow (Effigy)(FGY)
Chlorine............. Gas........... yellow-green (Cool GuY)
Bromine............. Liquid........ red-brown (BLuRB)
Iodine................ Solid.......... purple (PISsss) (Purple Iodine Solid)
Astatine............. Solid.......... black (ASS Black)


i just need a way to remember dont judge

2.10 make predictions about the properties of other halogens in this group

We would expect the colour to keep getting darker and the melting and boiling points to keep getting higher further down the group.

2.13 describe the relative reactivities of the elements in Group 7

Group 7 elements become less reactive as you go down the group.



At the top, the positive charge of the proton in the nucleus is close to the outside orbital this makes it easy for them to pull in the one electron they need to become stable, meaning they are very reactive.



The further away the outside orbital from the nucleus, the weaker the reaction.

2.14 describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts

A more reactive halogen will displace a less reactive one that is bonded as a salt. This will only happen if the salt is dissolved in water or a gas.

2.15 understand these displacement reactions as redox reactions

When a more reactive halogen displaces a less reactive one this is a redox reaction.


This means that one element has been reduced (gained electrons) and one has been oxidised (lost electrons.)


OILRIG


Chapter 13 Electrolysis


1.48 understand that an electric current is a flow of electrons or ions

An electric current is a flow of electrons, although it can also be a flow of ions (as they have a charge.)

1.49 understand why covalent compounds do not conduct electricity

In covalent compounds there are no electrons free to move, this means there can be no transfer of electricity through a covalent compound

1.50 understand why ionic compounds conduct electricity only when molten or in solution

When ionic compounds are molten or in solution, the positive and negative ions separate this means that there are ions free to flow, and so they can conduct electricity.

1.51 describe simple experiments to distinguish between electrolytes and non-electrolytes

Set up an electric circuit with an LED and a break in the wire, put both ends of wire into a solution/molten substance.


If the LED lights up then there is a current flowing, this will only be able to happen if the solution is conducting: so it must be an electrolyte.


But if the LED does not light up then there is no current flowing, and so the solution has not conducted electricity meaning it must be a nonelectrolyte.

1.52 understand that electrolysis involves the formation of new substances when ionic compounds conduct electricity

In electrolysis ionic compounds conduct electricity.


Positively charged ions move to one end, negatively to the other, these then lose their charge and turn into atoms.


Therefore, new substances are formed.

1.53 describe experiments to investigate electrolysis, using inert electrodes, of molten slats such as lead (II) bromide and predict the products

Inert electrodes are ones that don't react with any other substances, but only play a role in the transfer of electrons.
Lead bromide will make lead and bromine, you can use chemical tests to see is you got these products from electrolysis.

Inert electrodes are ones that don't react with any other substances, but only play a role in the transfer of electrons.


Lead bromide will make lead and bromine, you can use chemical tests to see is you got these products from electrolysis.

1.54 describe experiments to investigate electrolysis, using inert electrodes, of aqueous solutions of sodium chloride, copper(II) sulphate and dilute sulphuric acid and predict the products

Place inert electrodes into an aqueous solution. (H2O)
At the cathode a cation will form an atom. At the anode the anion will form an atom. The other ions remain in the solution.
NaCl: H2(g) at the cathode; Cl(g) at the anode
Cu(II)SO4: Cu(s) at the cathode; O2(g) + H2O(l) at the anode
H2SO4: H2(g) at the cathode; O2(g)
+ H2O(l) at the anode


If metal in solution is more reactive than hydrogen, hydrogen from water will be a product, as the metal will bond with the oxygen.

1.55 write ionic half-equations representing the reactions at the electrodes during electrolysis

At the positive electrode, electrons will be lost: to show this we write the lost electrons as products:
2Br- > Br2 + 2e-
Make sure the charges are equal on both sides: 1- > 1-.
At the negative electrode, electrons will be gained so we write them as reactants:
2H+ + 2e- > H2
And to make sure the charges are the same on both sides: 0 > 0.

Chapter 14 Introducing Energy Changes in Reactions


2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper (II) carbonate

When metal carbonates are heated they become carbon dioxide and a metal.
For example:
copper carbonate > copper oxide + carbon dioxide
CuCO3 > CuO + CO2

4.10 understand that chemical reactions in which heat energy is given out are described as exothermic and those in which heat energy is taken in are endothermic.

In an exothermic reaction heat is given out. Because bonds are made which gives out energy.
Think exit; exo.

In a endothermic reaction heat is taken in. Because bonds are Broken which requires energy.


Think into: endo

4.13 understand the use of ΔH to represent enthalpy change for exothermic and endothermic reactions

ΔH is the symbol that represents the amount of energy lost or gained in a reaction.

+ΔH is endothermic (because it gains heat)
-ΔH is exothermic (because it loses heat)

4.14 represent exothermic and endothermic reactions on a simple energy level diagram

On the left is an exothermic reaction. There is a lower energy level at the end.
An endothermic reaction would have a higher energy level at the end.

On the left is an exothermic reaction. There is a lower energy level at the end.


An endothermic reaction would have a higher energy level at the end.

Chapter 15 Introducing Reversible Reactions


4.22 understand that some reactions are reversible and are indicated by the symbol ⇌ in equations

Some reactions can happen both ways: the reactants can make the products and the products can make the reactants.

This symbol shows that ⇌.

4.23 describe reversible reactions such as the dehydration of hydrated copper (II) sulfate and the effect of heat on ammonium chloride

If you add water to copper sulphate you can make hydrated copper sulphate. If you remover the water from hydrated copper sulphate you can make copper sulphate.

When heated, ammonium chloride splits into to hydrogen chloride and ammonia. Hydrogen chloride and ammonia can be reacted to make ammonium chloride.


4.24 understand the concept of dynamic equilibrium

Dynamic equilibrium is when a reversible reaction is happening both ways at the same time, at the same rate.

4.25 predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions,

Increasing pressure favours side with least molecules. (Makes more)
Decreasing pressure favours side with most molecules. (Makes more)

If you increase the temperature, there will be more products that are produced by an endothermic reaction.
This is because the reaction is trying to use up the extra heat, and it does so by putting the energy into making bonds.

Chapter 16 Manufacturing Chemicals


5.22 understand that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia

Ammonia is made by reacting nitrogen from the air and hydrogen. Hydrogen comes as a natural gas or from cracking hydrocarbons.

5.23 describe the manufacture of ammonia by the Haber process, including the essential conditions.


5.24 understand how the cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated

Conditions: 
i) a temperature of about 450°C
ii) a pressure of about 200 atm
iii) an iron catalyst

Conditions:


i) a temperature of about 450°C


ii) a pressure of about 200 atm


iii) an iron catalyst

5.25 describe the use of ammonia in the manufacture of nitric acid and fertlisers

Ammonia is put into fertilisers (NPK) as it contains nitrate ions plants need to make amino acids ---> proteins for plants to grow.

Ammonia is also reacted with oxygen to produce nitric acid:
4NH3 + 8O2 > 4HNO3 + 4H2O

5.26 recall the raw materials used in the manufacture of sulfuric acid


5.27 describe the manufacture of sulfuric acid by the contact process, including the essential conditions

Sulphur (sulphur is found in rocks and some natural gasses), oxygen from the air and water.
(Stage 1) S + 2O > SO2 (Stage 2) SO2 + O > SO3

Sulphur (sulphur is found in rocks and some natural gasses), oxygen from the air and water.


(Stage 1) S + 2O > SO2
(Stage 2) SO2 + O > SO3

5.27 describe the manufacture of sulfuric acid by the contact process, including the essential conditions

(Stage 3) H2SO4(l) +SO3(g) > H2S2O7 (l) (oleum)


H2S2O7(l) + H2O(l) > 2H2SO4(l)




You can react SO3 with H2O but this produces an uncontrollable fog of concentrated sulfuric acid. Using oleum is safer.

5.28 describe the use of sulfuric acid in the manufacture of detergents, fertlisers and paints

Detergents: used to 'sulphonate' products (apperently.)
Fertilisers: reacted to make phosphates soluble to plants: reacted to make ammonia easier to handle.
Paints: reacted with titanium ore to make a main pigment in paint.

5.29 describe the manufacture of sodium hydroxide and chlorine by the electrolysis of concentrated sodium chloride solution brine) in a diaphragm cell

Brine is NaCl solution.
In electrolysis, chlorine is created at the anode from Cl-, hydrogen is created at the cathode from H+, NaOH is left in the solution.

2NaCl + 2H2O > 2NaOH + H2 + Cl2

This happens in a diaphragm cell.

5.30 write ionic half-equations for the reactions at the electrodes in the diaphragm cell

2Cl- ---> Cl2 + 2e-
2H+ + 2e- ---> H2

5.31 describe important uses of sodium hydroxide, including the manufacture of bleach, paper and soap; and of chlorine, including sterilising water supplies and in the manufacture of bleach and hydrochloric acid

Sodium Hydroxide: Bleach; paper; soap


Chlorine: sterilising water; bleach; hydrochloric acid

Chapter 17 Metals


2.34 describe the conditions under which iron rusts

Iron rusts in the presence of oxygen and water. Only the oxygen is really needed for the oxidation of iron to occur. Rusting is accelerated in the presence of electrolytes such as salt.



Note: Many metals corrode, but it is only the corrosion of iron that is referred to as rusting.

2.35 describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising

Grease, oil, paint and plastic prevent air and/or water from coming into contact with iron. This means the reaction that rusts iron can't occur.

Galvanising is coating in zinc. The zinc reacts in the air to form ZnCO3 which prevents air and/or water from coming into contact with the iron.

2.36 understand the sacrifical protection of iron in terms of the reactivity series

Sacrificial is covering a metal with a more reactive metal. What this means is water and/or air will react with the more reactive metal instead of the one underneath.

5.1 explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series

Anything below carbon can be displaced from its ore by carbon.
 
Anything above carbon can't so is extracted by electrolysis.

Anything below carbon can be displaced from its ore by carbon.



Anything above carbon can't so is extracted by electrolysis.

5.2 describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis

Bauxite is purified into aluminium oxide.
This is then dissolved in molten cryolite to bring down the boiling point
The walls of the tank are the negative electrode; here aluminium is made.

Bauxite is purified into aluminium oxide.


This is then dissolved in molten cryolite to bring down the boiling point


The walls of the tank are the negative electrode; here aluminium is made.

5.2 describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis CONTINUED

The aluminium sinks to the bottom and is tapped off.


Oxygen is formed at the anode.


The oxygen formed reacts with the graphite anode to from carbon dioxide; so the anode has to be replaced


It is very expensive to supply the electricity needed for this electrolysis.

5.3 write ionic half-equations for the reactions at the electrodes in aluminium extraction

Al3+ + 3e- →Al


2O2- → O2 + 4e-

5.4 describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace

Iron is displaced from its ore by carbon (from coke):2Fe2O3 + 3C > 4Fe + 3CO2It's also displaced by COFe2O3 + 3CO > 2Fe + 3CO2Limestone reacts with impurities to from 'slag' which is tapped off.Oxygen allows burning of coke

Iron is displaced from its ore by carbon (from coke):
2Fe2O3 + 3C > 4Fe + 3CO2
It's also displaced by CO
Fe2O3 + 3CO > 2Fe + 3CO2
Limestone reacts with impurities to from 'slag' which is tapped off.
Oxygen allows burning of coke

5.5 explain the uses of aluminium and iron, terms of their properties

Aluminium is light and does not corrode so is used for the bodies of planes.

Iron is used in electromagnets as it is a soft magnetic material and used in ships and bridges because it's strong.

Chapter 18 Introducing Organic Chemistry


3.1 explain the terms of homologous series, hydrocarbon, saturated, unsaturated, general formula and isomerism

Compounds in the same homologous series have the same general formula and similar chemical properties.
A hydrocarbon is made up only of hydrogen and carbon, chemically bonded together.
Saturated means something has bonded as many time as possible. Unsaturated means that more bonds can be made.
General formula is the formula to work out the molecular formula.
Isomers have the same molecular formula but different structures.

3.3 draw displayed formulae for alkanes with up to five carbon atoms in a molecule and name the straight-chain isomers

Pentane has 5 carbon and 12 hydrogen atoms.

Pentane has 5 carbon and 12 hydrogen atoms.

3.7 draw displayed formulae for alkenes with up to four carbon atoms in a molecule, and name the straight-chain isomers

Can only have double bond between two carbons. Therefore there isn't a methene.

Can only have double bond between two carbons. Therefore there isn't a methene.

Chapter 19 Alkanes, Alkenes and Alcohols


3.1 explain the terms homologous series, hydrocarbon, saturated, unsaturated, general formula and isomerism

Compounds in the same homologous series have the same general formula and similar chemical properties.
A hydrocarbon is made up only of hydrogen and carbon, chemically bonded together.
Saturated means something has bonded as many time as possible. Unsaturated means that more bonds can be made.
General formula is the formula to work out the molecular formula.
Isomers have the same molecular formula but different structures.

3.2 recall that alkanes have the general formula CnH2n+2

Alkanes is a homologous series with the formula CnH2n +2.

3.4 recall the prodcuts of the complete and incomplete combustion of alkanes

Complete combustion gives carbon dioxide and water.

Incomplete combustion gives carbon monoxide and water.

3.5 describe the substitution reaction of methane with bromine to form bromomethane in the presence of UV light

In UV light bromine and methane will form bromomethane:

CH4 + Br2CH3Br + HBr



What has happened in this reaction is a bromine has taken the place of a hydrogen (substitution).

3.6 recall that alkenes have the general formula CnH2n

Alkenes is a homologous series and has the general formula CnH2n

3.8 describe the addition reaction of alkenes with bromine, including the decolourisation of bromine water as a test for alkenes

An alkene will make its double bond into a single bond, to bond to two bromines. Bromine is added to the molecule. The product made is colourless. When alkenes are put in bromine water it turns from brown to colourless (a good way of testing for alkenes.)
For example:

C2H4(g) + Br2 (aq) → C2H4Br2 (aq)

3.9 describe the manufacture of ethanol by passing ethene and steam over a phosphoric acid catalyst at a temperature of about 300°c and a pressure of about 60-70 atm

C2H4 (ethene) + H2O (steam) > C2H5OH (ethanol)


The conditions are: temperature of about 300°c and a pressure of about 60-70 atm


Catalyst: phosphoric acid

3.10 describe the manufacture of ethanol by the fermentation of sugars, for example glucose, at a temperature of about 30°C

Ethanol can be made by the anaerobic respiration of microorganisms.

Glucose > ethanol + carbon dioxide

This happens at 30 degrees.

3.11 evaluate the factors relevant to the choice of method used in the manufacture of ethanol, for example the relative availability sugar cane and crude oil

Fermentation


Cane sugar widely avalible/ cheap/ renewable


Slow process, done in batches


Impurities in the product
Hydrating (ethene and steam)


Crude oil (cracked to make ethene) expensive/ non-renewable


Fast process, continuous


Pure product

3.12 describe the dehydration of ethanol to ethene, using aluminium oxide

C2H5OH > C2H4 + H2O

ethanol > ethene + water

aluminium oxide is the catalyst for this reaction.

5.10 understand that incomplete combustion of fuels may produce carbon monoxide and explain that carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen

Hydrocarbons (from crude oil) + oxygen > carbon dioxide + water

However, if there is not enough oxygen around. Then fuels will combust incompletely:
Hydrocarbons + oxygen > carbon monoxide + carbon + water
Carbon monoxide combines with heamaglobin in red blood cells, meaning they can't carry oxygen around the body

Chapter 20 Useful products from crude oil


5.6 understand that crude oil is a mixture of hydrocarbons

Crude oil contains different molecules made only of hydrogen and carbon. They are hydrocarbons.

5.7 describe and explain how the industrial process of fractional distillation separates crude oil into fractions

Crude oil is heated until it boils. As a gas it floats upwards. 
The different compounds have different condensing points. When they condense, they can be collected.

Crude oil is heated until it boils. As a gas it floats upwards.


The different compounds have different condensing points. When they condense, they can be collected.

5.8 recall the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen

Refinery Gases- Bottled Gas


Gasoline- Petrol for vehicles


Kerosene- Jet fuel


Diesel- Diesel fuels for trucks and shiz


Fuel Oil- Fuel for ships and factories.


Bitumen- Roads and Roofing

5.9 describe the trend in boiling point and viscosity of the main fractions

Fractions with low boiling points are less viscous.

Fractions with high boiling points are more viscous.

5.10 understand that incomplete combustion of fuels may produce carbon monoxide and explain that carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen

Hydrocarbons (from crude oil) + oxygen > carbon dioxide + water

However, if there is not enough oxygen around. Then fuels will combust incompletely:
Hydrocarbons + oxygen > carbon monoxide + carbon + water
Carbon monoxide combines with heamaglobin in red blood cells, meaning they can't carry oxygen around the body

5.13 understand that fractional distillation of crude oil produces more long-chain hydrocarbons that can be used directly and fewer short-chain hydrocarbons than required and explain why this makes cracking necessary

Long chain hydrocarbons are less flammable and more viscous.
Short chain hydrocarbons burn well and flow well.
Therefore, people want short chain hydrocarbons more. Supply and demand.
Long chain hydrocarbons can be cracked which breaks them up into short chain ones.

5.14 describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes by catalytic cracking, using silica or alumina as the catalyst and a temperature in the range of 600-700°c

Long chain hydrocarbons are passed over a hot catalyst (silica or alumina at 600-700 degrees) this causes them to break down into smaller molecules.

As some atoms are lost from molecules, they become unsaturated and can therefore form a double bond. This is how you get alkenes from the process as well as shorter chain alkanes.

Chapter 21 Polymers


5.15 understand that an addition polymer is formed by joining up many small molecules called monomers

Monomers are alkenes with a double bond. If this bond is broken there can be other things bonded, if a carbon from another monomer is bonded to it then you can create a chain; do this many times and you have a polymer.


This is an addition reaction.

5.16 draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and poly(chloroethene)

Draw them in this format. Ignore PTFE.
There is no double bond.

Draw them in this format. Ignore PTFE.


There is no double bond.

5.17 deduce the structure of a monomer from the repeat unit of an addition polymer

A monomer that is repeated in a polymer looks much like the repeat unit; apart from, instead of having an empty bond either end, it has a double bond in the middle.

5.18 describe some uses for polymers, including poly(ethene), poly(propene) and poly(chloroethene)

Polyethene: plastic carrier bags; plastic bottles
Polypropene: crates; ropes
Polychloroethene: piping; cable insulation.

5.19 explain that addition polymers are hard to dispose of as their inertness means that they do not easily biodegrade

Polymers are saturated so they don't react, they are inert. This means they don't decompose easily.

5.20 understand that some polymers, such as nylon, form by a different process called condensation polymerisation

Condensation polymerisation occurs if two monomers combine but a small molecule is lost.


Nylon-6,6 is made from 2 monomers each containing 6 carbon atoms. A dicarboxylic acid -COOH and a diamine -NH2 join together with the loss of a water molecule each time.


Industrially making nylon is done at 350 degrees, but can be done a room temp. with a few modifications.

5.21 understand that condensation polymerisation produces a small molecule, such as water, as well as the polymer

Two monomers come together by losing a molecule. Atoms from each monomer join together to make the molecule: commonly a H atom from one and a OH molecule from another form water. The two monomers then join together, making a polymer.

Chapter 22 RAMs and Moles


1.11 understand the terms atomic number, mass number, isotopes and relative atomic mass.

Atomic number: number of protons (the same as number of electrons.)
Atomic mass: number of protons + neutrons.
Isotopes: Atoms of an element with different numbers of neutrons.
Relative atomic mass (Ar): the mass of one atom of an element in proportion to carbon-12

1.12 calculate the relative atomic mass of an element from the relative abundances of its isotopes

Relative atomic mass is the average weight of an atom of an element. There is variation in the weight due to the fact that there are different isotopes with in an element. (Different isotopes have different weights due to less or more neutrons.)



It is calculated like this:


(% of isotope x its mass) + (% of isotope x its mass) / 100


(75x35)+(25x37)=35.5

1.16 calculate relative formula masses (Mr) from relative atomic masses (Ar)

Add up the Ar of all the atoms in the molecule to give Mr.

1.17 understand the use of the term mole to present the amount of substance


1.18 understand the term mole as the Avogadro number of particles (atoms, molecules, formulae, ions or electrons) in a substance

Having a mole of something is having 6.022x10^23 of it.


6.022x10^23 is Avogadros constant.

1.19 carry out mole calculations using relative atomic mass (Ar) and relative formula mass (Mr)

Mr MASSive Mol

Mr MASSive Mol

1.23 understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Weigh your compound, remove one element through a reaction, then weigh again.


The first weight is AB the second weight is A so doing AB-A= B.


Now that you have the weight of A and B you can work out the formulae by doing the weight divided by the Ar. Mr Massive Mol


e.g A= 22g with an Ar of 11 A= 22/11= 2


B=72g with an Ar of 18 B=72/18= 4


A2B4 simplified to AB2

1.24 calculate empirical and molecular formulae from experimental data

Find the masses of the elements in the compound. Make these masses into percentages, divide each percentage by the Ar of that element and you will have the number of atoms of it in a molecule.
e.g Carbon= % mass 80, Ar 12 C= 80/12= 6.7
Hydrogen= % mass 20, Ar 1 H= 20/1= 20 =C6.7H20
To find the empirical formula divide the molecular formula by the smallest number.
e.g. C6.7H20 /6.7=CH3

Chapter 23 Calculations from equations


1.20 understand the term molar volume of a gas and use its values (24dm^3 and 24000 cm^3) at room temperature and pressure (rtp) in calculations

At standard temperature and pressure, one mole of any gas will occupy 24000 cm3; also known as 24dm3.

1.23 understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Weigh your compound, remove one element through a reaction, then weigh again.


The first weight is AB the second weight is A so doing AB-A= B.


Now that you have the weight of A and B you can work out the formulae by doing the weight divided by the Ar. Mr Massive Mol


e.g A= 22g with an Ar of 11 A= 22/11= 2


B=72g with an Ar of 18 B=72/18= 4


A2B4 simplified to AB2

1.25 calculate reacting masses using experimental data and chemical equations

This means knowing that the mass of A and the mass of B will equal the mass of AB.
Bearing in mind that the mass of one mole of an element is its atomic mass.
If we have 2 moles of A(24) and react it with 5 moles of B(25)
then we have (2x24)+(5x25) and can calculate that A2B5 will weigh 173g

1.26 calculate percentage yield

Percentage yield is: (actual yield / theoretical yield) x 100
Theoretical yield is what you expect to get, for example if a reaction doesn't finish you may end up with a lower yield. (Work this out using Mr Mass Mol and the ratio)

Chapter 24 Electrolysis calculations


1.56 recall that one faraday represents one mole of electrons

One Faraday is 96500 coulombs. That is the amount of coulombs in one mole of electrons.

1.57 calculate the amounts of the products of the electrolysis of molten salts and aqueous solutions

Haven't done yet

Chapter 25 Energy Calculations


4.11 describe simple calorimetry experiments for reactions, such as combustion, displacement, dissolving and neutralisation in which heat energy changes can be calculated from measured temperature changes

Haven't done yet

4.12 calculate molar enthalpy change from heat energy change

Haven't done yet

4.15 understand that the breaking of bonds is endothermic and the making of bonds is exothermic

Haven't done yet

4.16 use average bond energies to calculate the enthalpy change during a simple chemical reaction

Haven't done yet

Chapter 26 Titration calculations


1.27 carry out mole calculations using volumes and molar concentrations

Moles / volume = concentration moldm^-3


Mass / volume = concentration gdm^-3

4.9 describe experiments to carry out acid-alkali titrations

Put a known volume of acid in a beaker with methyl orange, red.


Set up a burette with alkali in it. Open the tap very slightly, so that it drips very slowly into the acid, stirring the beaker with each drop.


The more alkali that is added the more neutral orange it gets.


When the solution is neutral it will be completely orange, at this point close the tap on the burette.


The level in the burette will have dropped, showing the volume of alkali used. This volume is needed to neutralise the acid.