This document contains summary notes needed for the AQA GCSE Chemistry Exam (HIGHER/SEPERATE SCIENCE) that allowed me and others to achieve a grade 9, and so can you!
Chemistry
Paper 1
Atomic Structure and the Periodic table:
● An atom is the smallest part of an element that can exist
● An element is a substance that contains only one type of atom
● A compound contains atoms of two or more elements, which are chemically combined in fixed proportions
● In reactions, the total mass of the products is always equal to the total mass of the reactants - because no atoms are
lost or made
● Separating mixtures:
○ Mixtures consist of two or more elements or compounds, which are not chemically combined
○ Filtration - separates soluble solids from insoluble solids, such as salt (soluble) and sand (insoluble), by
dissolving the salt in water and filtering the mixture.
○ Crystallisation - is used to obtain a soluble solid from a solution, such as salt crystals(1. The mixture is
gently warmed, 2. The water evaporates leaving crystals of pure salt)
○ Simple distillation - used to obtain a solvent from a solution.
○ Fractional distillation - used to separate mixtures in which the components have different boiling points,
crude oil and hydrocarbons
○ Chromatography - is used to separate the different soluble, coloured components of a mixture, such as the
coloured parts of a fizzy drink
● Scientific models of the atoms:
○ In early models, atoms were thought to be tiny spheres that could not be divided into smaller particles
○ In 1898, Thomson discovered electrons. Overall an atom is neutral. Thomson thought atoms contained tiny
negative electrons and were surrounded by a sea of positive charge - this led to the plum pudding model.
○ Later, Geiger and Marsden carried out the alpha particle experiment. They bombarded a thin sheet of gold
with alpha particles. Most positive particles went straight through the atoms, some atoms were deflected.
Rutherford concluded that there was a positive charge in the middle of an atom.
○ This was called the nucleus, and the nuclear atom model replaced the plum pudding model
○ Bohr deduced that electrons orbit the nucleus at specific distances, otherwise they would spiral inwards
○ 20 years later, the nucleus became more accepted as an idea. Chadwick then provided evidence of
neutrons in the nucleus.
○ Later evidence said that the positive charge in the nucleus could be divided into protons.
○
, ●
● Isotopes of an element have the same number of protons but a different number of neutrons
● Metal atoms lose electrons to form positive ions
● Non-metal atoms gain electrons to form negative ions
● The electrons in an atom occupy the lowest available shell or energy level
● Development of the Periodic table:
○ John Newlands put together a table in 1863, with only 63 known elements - in order of atomic weight
○ He noticed patterns but the missing elements caused problems
○ Dimitri Mendeleev realised that there were undiscovered elements. In his table, 1869, he left gaps and
reordered elements.
○ Each element was placed in a vertical column with elements that had similar properties.
○ Mendeleev used his table to predict other elements. When subatomic particles were later discovered, it
was revealed that Mendeleev organised the elements in order of increasing atomic number (number of
protons)
● Group 0 = Noble gases. They have a full outer shell of electrons. They have a stable configuration, so they are
unreactive non-metals. The boiling points increase down the group
● Group 1 = Alkali Metals. They have: 1 electron in the outer shell, have low melting and boiling points that decrease
down the group and become more reactive down the group (because the outer electrons get further away from the
nucleus) They are stored under oil because they react very vigorously with oxygen and water. When they react with
water, a metal hydroxide is formed and hydrogen gas is given off. They have a low density; lithium, sodium and
potassium float on water because they are less dense than it. When a metal hydroxide is dissolved in water, an
alkaline solution is produced. Alkali metals react with non-metals to form ionic compounds - the metal atom loses one
electron to form a metal ion with a +1 charge.
● Group 7 = Non-metals and are called halogens. They have 7 electrons in the outer shell. The halogens are molecules
made of pairs of atoms. Reactivity decreases down the group and the further down group 7, the higher the relative
molecular mass, melting point and boiling point. Halogens react with metals to produce ionic salts, the halogen atom
gains 1 electron to form a halide ion with a negative charge. A more reactive halogen will displace a less reactive
halogen from an aqueous solution of its salt; chlorine would displace bromine from potassium bromide
● Transition Metals = Are in the centre of the periodic table, between groups 2 and 3. They form coloured compounds
and have ions with different charges. They can be used as catalysts and are good conductors of heat and electricity.
They are also malleable.
Bonding, Structure and the properties of matter:
● Chemical bonds = There are three types of strong chemical bonds, ionic bonding, covalent bonding and metallic
bonding.
● Ionic bonding:
○ Occur between positive and negative ions, they are formed when atoms lose or gain electrons, giving them
an overall charge
○ Ions have a complete outer shell of electrons
○ Ionic bonding involves the transfer of electrons from metal atoms to non-metal atoms. The metal atom
loses electrons to form a positive ion and the non-metal atom gains electrons to form a negative ion.
○ The ionic bond is a strong electrostatic force of attraction between the positive and negative ion
○ Ionic compounds have high melting and boiling points because ionic bonds are very strong and require lots
of energy to overcome them.
● Ionic compound properties:
, ○ Ionic compounds are giant structures of ions
○ They are held together by strong electrostatic forces of attraction that act in all directions between
oppositely charged ions.
○ They have high melting and boiling points
○ They do not conduct the electricity when solid, because the ions cannot move
○ They conduct electricity when molten or in solution because the ions can move and carry their charge.
○
● Metallic bonding:
○ Occurs in metallic elements (like iron and copper), and alloys (like stainless steel)
○ Metals have a giant structure in which electrons in the outer shell are delocalised (not bound to one atom)
○ This produces a regular arrangement (lattice) of positive ions held together by electrostatic attraction to the
delocalised electrons.
○ A metallic bond is an attraction between the positive ions and the delocalised negatively charged electrons
○
● Covalent bonding:
○ A covalent bond is a shared pair of electrons between atoms.
○ They occur in non-metallic elements like oxygen, and compounds of non-metals like sulphur dioxide
○ Covalent bonds are very strong.
○ Simple structure covalent bonds are:
■ Water = H-O-H
■ Hydrogen = H-H
■ Hydrogen chloride = H-Cl
■ Methane =
■ Oxygen = O-O
■ Nitrogen =
○ Others have giant covalent structures, like diamond and silicon dioxide
● Simple molecules:
○ Contain a relatively small number of non-metal atoms joined together by covalent bonds.
○ The molecules have no overall charge, so they cannot conduct electricity
○ They are usually liquids and gases, that have relatively low melting and boiling points
○ They have weak intermolecular forces, compared to the strength of the covalent bonds in the molecules
themselves.
○ The larger the molecules are, the stronger the intermolecular forces. This means they have higher melting
and boiling points
○ Going down group 7, the molecules get larger. Fluorine and chlorine are gases at room temperature, but
bromine is a liquid and iodine is a solid.
● Giant covalent structures:
○ All atoms are linked by strong covalent bonds
○ The bonds must be broken for the substance to melt or boil, this means if they have a giant covalent
structure - they have very high melting and boiling points.
○ Diamond:
■ Is a form of carbon
■ It has a giant, rigid covalent structure (lattice)
■ Each carbon atom forms four strong covalent bonds with other carbon atoms
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