IGCSE Chemistry
From the Edexcel IGCSE 2009 Syllabus including triple science statements
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Unit 1: The Periodic Table
The periodic table contains about a hundred or so elements that have been currently discovered. The rows are known as periods and elements of the same period have the same number of electron shells. The columns are known as groups and elements of the same group have the same number of electrons on their outer shell. Group one has one outer electron,
An element is a substance that cannot be broken down into anything simpler. KCl for example (potassium chloride) is NOT an element because it can be broken down into K (potassium) and Cl (chlorine). The potassium and chlorine are the elements.
A compound is two or more elements chemically bonded together. An example would be KCl (potassium chloride), which consists of the elements potassium and chloride chemically bonded together.
Atoms are the building blocks of substances.
Molecules are two or more atoms bonded together. It doesn’t have to be a compound. Elements such as O2 and Br2 are diatomic molecules – they exist in pairs.
Atomic Structure
Atoms are made up of protons, neutrons and electrons. Protons are positively charged. Electrons are negatively charged. Neutrons don’t have a charge.
An atom consists of a nucleus, which contains protons and neutrons; and some electron shells which surround the nucleus and contain electrons. The neutrons however, are different. The number of protons and the number of neutrons add up to make the mass number of an element.
Understanding the Lack of Reactivity in Noble Gases (Group 0)
Noble gases have eight electrons on their outer shell, therefore, there is no need for them to gain or lose electrons. Basically they have a full outer shell so they don’t need to react. This is what makes them so unreactive.
How to Read Each Square on The Periodic Table
You probably already know that the periodic table is made up of lots and lots of squares, each containing an element and information about it.
Anyways we already know what the atomic mass number is (the number of protons + neutrons). It says 12.011 here but this is probably because this picture came from some super complicated periodic table. In IGCSE level however, the atomic mass should read 12. Anyways, the atomic number is the number of protons (and electrons), so to find the number of neutrons, if asked to, simply subtract the atomic mass by the atomic number.
Example: Calculate the number of neutrons Carbon has.
The answer: 12 – 6 = 6 neutrons
The Arrangement of Electrons
Atoms are surrounded by electron shells which contain electrons. But the arrangement is the same for ALL the elements, not matter how different they are.
Each shell can only hold a certain number of electrons. The very first shell can hold only two electrons. The second shell can hold eight. The third sometimes appear full with eight but can expand to a total of eighteen. However, this is beyond GCSE level, and for now, the shells only hold eight.
So how do you find the electron configuration? Well let’s use potassium (K) as an example.
Look up the atomic number of potassium. It should say 19. This tells you the number of protons, which is equal to the number of electrons so we can use that.
Arrange the electrons in shells, always filling up the inner shell before you go to the outer one. Remember the first, innermost shell can only take 2 electrons, the second one can take 8, and the third one, 8. You will find that you have one electron left. That goes on the fourth shell.
Your electron configuration should look like this: 2, 8, 8, 1.
Example: Work out the electron configuration of chlorine.
Chlorine has an atomic number of 17 – so 17 electrons.
17 – 2 (as the innermost shell only holds two electrons) = 15
15 – 8 (as the second shell only holds eight electrons) = 7 (This number is the number of electrons Chlorine has on its outer shell).
7 electrons does not fill up the third shell so we are left with the configuration: 2, 8, 7.
Isotopes
The number of neutrons in an atom can vary slightly. For example, there are three kinds of carbon atom, called carbon-12, carbon-13 and carbon-14. They all have the same number of protons, but the number of neutrons vary. These different atoms of carbon are called isotopes. Isotopes are atoms that have the same atomic number, but different mass numbers. They have the same number of protons, but different numbers of neutrons. The fact that they have varying numbers of neutrons makes no difference whatsoever to their chemical reactions. The chemical properties are governed by the number and arrangement of the electrons.
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
Now what did we do there? Well I simply calculated 51.84% of 107 (of silver) and 48.16% of 109 (of silver), and added the two answers! What we end up with is 107.96. Round that up to a whole number and the average relative atomic mass of silver is 108.
Calculating the Abundance (percentage) of an Isotope
Example: Copper consists of two isotopes, copper-63 and copper-65. Its relative atomic mass is 63.62. Find the abundance of each isotope.
Let y/100 = abundance of copper-63
Let (100-y)/100 = abundance of copper-65
63.62 = (y/100 x 63) + [(100-y)/100 x 65]
63.62 = 63y +6500 – 65y
-2y = -135
y = 69
Abundance of copper-63 = 69%
Abundance of copper-65 = 100 – 69 = 31%
About Metals and Non-Metals
The IGCSE spec. states you have to recall the positions of metals and non-metals on the periodic table. That’s easy! Its on page two. Have a look. Its colour-coded.
Anyways, this section covers 2.2, 2.3 and 2.5.
Metals
Metals tend to be shiny. They tend to 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
Non-metals tend to be brittle. They are poor conductors of heat and electricity. They form acidic oxides and are good oxydising agents because they gain electrons.
Aluminium Oxide
Aluminium oxide is amphoteric. It can neutralize both an acid and a base.
Reaction with acids
Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. That means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride solution.
In this (and similar reactions with other acids), aluminium oxide is showing the basic side of its amphoteric nature.
Reaction with bases
Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodium hydroxide solution. Various aluminates are formed - compounds where the aluminium is found in the negative ion. This is possible because aluminium has the ability to form covalent bonds with oxygen.
Group 1: The Alkali Metals
Alkali metals are metals that are part of group one. They are extremely reactive metals, and reactivity increases DOWNWARDS – in other words, lithium is the least reactive and francium.
Some Basic Physical Properties
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!
The metals are also very soft and easy to cut, becoming softer as you go down the group. They are shiny and silver when cut, but tarnish within seconds on exposure to air.
Storage and Handling
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.
To stop them reacting with oxygen or water vapour in the air, lithium, sodium and potassium are stored under oil. Rubidium and caesium are so reactive that they have to be stored in sealed glass tubes to stop any possibility of oxygen getting at them.
Great care must be taken not to touch any of these metals with bare fingers. There could be enough sweat on your skin to give a reaction producing lots of heat and a very corrosive metal hydroxide.
Reactions with Water
All these metals react with water to produce a metal hydroxide and hydrogen.
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.
Explaining the Increase in Reactivity
The differences between reactions depend in part on how easily the outer electron of the metal is lost in each case. That depends on how strongly it is attracted to the nucleus. The more electron shells an atom has, the less powerful the attraction forces are. For example, Lithium is a lot less reactive than Potassium. This is because there are less shells which shield the full attraction of the nucleus from the This makes the electron harder to lose. However, potassium has a lot more electron shells which shield the outer electron from the nucleus. This weakens the attraction in compared to lithium, and therefore, the electron is easier to lose.
Compounds of Alkali Metals
All group one metal ions are colourless. That means that their compounds will be colourless or white unless they are combined with a coloured negative ion (remember metals would become positive ions because they lose electrons, whereas, most non-metals gain electrons). Potassium dichromate is orange, for example, because the dichromate ion is orange. Group one compounds are typical ionic solids and are mostly soluble in water.
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
§ Flame Colours: lithium, red; sodium, yellow; potassium, lilac; rubidium, red; caesium, blue.
§ Forget about Francium you don’t need to know much about it.
Group 2: Alkali Earth Metals
Alkali earth metals belong to Group two. They are beryllium, magnesium, calcium, strontium, barium and radium. These metals are harder than those in group one. They are silvery grey in colour. They tarnish quickly, however they don’t just disappear into thin air because the oxides the metals form when reacting with air would form an outer coat that protects the metal from the air. They are good conductors of heat and electricity. They burn in oxygen to form white oxides. They react with water to form hydroxides and hydrogen, but the reaction is a lot less than that of group one. Also, reactivity increases down the group.
Flame Colours
-Calcium –brick red
-Strontium -crimson
-Barium -apple-green
Well that’s it for Group two!
Alkali Earth Metals: Quick Notes
§ 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.
Element / State / ColoursFlourine / Gas / Yellow
Chlorine / Gas / Green
Bromine / Liquid / Orange – Brown vapour
Iodine / Solid / Dark Grey – Purple vapour
Group 7: The Halogens
Halogens are group seven elements. Their elements are diatomic molecules. They exist in pairs, such as F2 and Cl2. These two elements are gases, bromine is a liquid and iodine is a solid. Astatine is radioactive.