We offer Easter Revision, May Half Term Revision and Christmas/New Year Revision courses to work in cohesion with the build up towards the GCSE and IGCSE examinations.
Taking place in our tuition centres in High Wycombe, St Albans and Harrow, our tutor-led GCSE revision courses are designed to provide in-depth, focused preparation for the GCSE exams. We aim to work on exam technique, confidence and attainment to make sure you have all bases covered before the exam.
1. Animal Cells (multicellular organism)
Store carbohydrates as glycogen Heterotrophic – they consume other organisms as food.
A – Nucleus – for nervous coordination, controls functions within the cell.
B – Cytoplasm – Where most chemical reactions inside the cell take place. >
C – Cell Membrane – partially permeable membrane that controls what enters and exits the cell
D – Mitochondria – Where aerobic respiration occurs
NaOH + HCl → NaCl + H2Oh) A base will readily accept an H+ ion from an acid, to form a new compound, e.g. OH- forms H2O and NH3 forms NH4+. i) **Hydrated refers to a crystalline compound containing water molecules. Anhydrous refers to a substance that contains no water molecules. Water of Crystallisation refers to water molecules that form an essential part of the crystalline structure of a compound.** j) Using percentage composition or mass composition, you can find the formula of a salt by finding the ratio of water molecules to molecules of the anhydrous salt. This, along with stoichiometry, allows the formula to be found. This can also be obtained using experimental data. **k) Perform acid-base titrations and carry out structured titrations** – this is a practical element, but can be described in an exam as: - Use a burette with either acid or alkali in the top, and drop it down into a conical containing the other solution. - Do a rough run using an indicator such as methyl orange, and then use this value to repeat the titration more exactly. - If a salt needs to be obtained, repeat the titre without the indicator, filter and evaporate to give pure crystals. **1.1.4** **a) An oxidation number is a measure of the number of electrons that an atom uses to bond with atoms of other elements.** They are also known as oxidation states. Oxidation numbers are derived from a set of rules: * Elements have an oxidation number of 0 (e.g. Na = 0, Cl2 = 0) * Neutral molecules have an overall oxidation number of 0 (e.g. NaCl = 0, where Na = +1 and Cl = -1) * Charged species have an overall oxidation number the same as the charge (e.g. [SO4]2- = -2) * Metal ions have an oxidation number equal to their group number (1-3) because they have a charge that is equal to their group number (e.g. Na+ = +1, Mg2+ = +2) * Combined hydrogen has an oxidation number of +1, except in metal hydrides where it has an oxidation number of -1. * Combined oxygen has an oxidation number of -2, except in peroxides, where it is -1 or when combined with fluorine, which is more electronegative than oxygen. * Combined halogens usually have an oxidation number of -1. b) **OIL RIG Oxidation is the loss of electrons or an increase in oxidation number of a species. Reduction is the gain of electrons or a decrease in oxidation number of a species. ** c) Roman numerals can be used to indicate the magnitude of the oxidation state of an element when the name may be ambiguous, e.g. nitrate(III) and nitrate(V), which are both nitrate groups, but in nitrate(III) the nitrogen has an oxidation number of +3, whereas in nitrate(V) the nitrogen has an oxidation number of +5. d) You can write formulae using oxidation numbers – if the overall compound is neutral, then the overall sum of the oxidation numbers must be zero. If the compound has an overall charge then the sum of the oxidation numbers must match that charge. E.g. SO42- has a charge of 2- and so its oxidation state must be -2. Oxygen has an oxidation number of -2, which means that in this case, sulphur must be +6, in order to balance the compound with a 2- charge. e) (i) Metals generally form ions by losing electrons, which increases their oxidation number. They form positive ions. Most metals have 1-3 electrons in their outer shell, and so to gain stability, they need to lose electrons. (ii) Non-metals generally gain electrons to form ions, which decreases their oxidation number. They form negative ions. Non-metals have 4-8 electrons in their outer shell and so in order to gain stability, it is easier to gain electrons. The noble gases, which have 8 outer shell electrons, are inert and do not easily form ions. f) The reaction of a metal with dilute acid is a redox reaction. The generic equation is metal + acid → salt + hydrogen. E.g. Mg + 2HCl → MgCl2 + H2 E.g. Ca + H2SO4 → CaSO4 + H2 g) **INTERPRET AND MAKE PREDICTIONS FORM REDOX EQUATIONS IN TERMS OF OXIDATION NUMBERS AND ELECTRON LOSS OR GAI.** **MODULE 2** **1.2.1** a) **The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of unipositive gaseous ions. Successive ionisation energies are a measure of the energy required to remove each electron in turn. For example, the second ionisation energy of an element is the energy required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of 2+ gaseous ions.** b) Ionisation energies are affected by nuclear charge, electron shielding and distance of the outermost electron from the nucleus: * Nuclear charge – the more protons there are in the nucleus, the greater the positive charge, which means that the electrons are more tightly held. As you go across the period, the electrons are held in the same space but held by a stronger nuclear charge, so the ionisation energies increase. * Electron shielding – ionisation energies decrease as you go down a group as there are more shells of electrons between the nucleus and the outer shell electrons, which reduces the effective nuclear charge on the electrons and makes it easier to remove them. * Distance of outermost electrons from the radius – as you go down a group, the atomic radius increases due to more shells of electrons, which means there is a greater distance between the nucleus and the outermost electrons, which causes the effective nuclear charge to be weaker. c) **PREDICT SUCCESSIVE IONISATION ENERGIES** d) **A shell is a group of atomic orbitals with the same principal quantum number, n. This is also known as a main energy level. Principal Quantum Number, n, is a number representing the relative overall energy of each orbital which increases with distance from the nucleus. The sets of orbitals with the same n value are referred to as electron shells or energy levels. ** The first energy level can hold only two electrons, in an s orbital. The second can hold 8, in an s orbital and a p orbital. The 3rd and 4th can both hold 18, in s, p and d orbitals. e) **An orbital is a region that can hold up to two electrons with opposite spins. ** f) An s orbital is a roughly spherical shaped orbital, which holds 2 electrons. A p orbital is propeller shaped and can also hold 2 electrons. g) An s subshell is made up from one s-orbital and can hold one pair of electrons. A p subshell is made up from three p-orbitals, perpendicular to each other, and can hold 6 electrons in total, two in each p orbital. A d subshell is made up from 5 d-orbitals, each of which can hold two electrons, so the d-subshell can hold 10 electrons in total. h) An s orbital has less energy than a p orbital, which has less energy than a d orbital. The first energy level is the lowest energy, and then the second energy level etc. The orbitals fill in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p... i) **DEDUCE THE ELECTRON CONFIGURATIONS OF ATOMS AND IONS** ii) **CLASSIFY THE ELEMENTS INTO S P AND D BLOCKS** **1.2.2** a) **Ionic bonding is the electrostatic attraction between oppositely charged ions. ** b) **DOT AND CROSS DIAGRAMS IONIC** c) You can predict the charge of ions from their position on the ionic table, due to the electron movement that will make them the most stable. Groups 1, 2 and 3 lose 1, 2 and 3 electrons respectively to form their ions, therefore usually have a charge of 1+, 2+ or 3+. Group 7 usually gains one electron, group 6 gains 2 and group 5 gains 3, therefore these groups usually have a charge of -1, -2, and -3 respectively. d) Some common radical groups are ions: Nitrate: NO3- Carbonate: CO32- Sulphate: SO42- Ammonium: NH4+ e) **A covalent bond is a shared pair of electrons** f) **DOT AND CROSS COVALENT** g) The shape of a simple molecule is determined by the repulsion between electron pairs surrounding the central atom to give maximum stability. h) Lone pairs of electrons repel more than bonded pairs of electrons, and so the angle created by lone pairs is greater than the angle created by bond pairs. i)** EXPLAIN SHAPES OF AND BOND ANGLES IN MOLECULES AND IONS SURROUNDING A CENTRAL ATOM ** j)** PREDICT THE SHAPES OF AND BOND ANGLES IN MOECULES ANALAGOUS TO THOSE STUDIED** k) **Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond.** l) When covalently bonded atoms have different electronegativity, the electrons in the bond will sit closer to the more electronegative species, which causes a permanent dipole within the bond, which causes a polar bond, as the less electronegative end will be ᵟ+ and the more electronegative end will be ᵟ-. m) There are three types of intermolecular forces – Van der Waals, dipole-dipole and hydrogen bonds. * Van der Waals (e.g. noble gases) – Electrons in a molecule are constantly in flux, and at different times there may be more electrons in one end of the molecule than the other, which creates an instantaneous dipole on the molecule. When this comes up against another molecule, it can create a temporary induced dipole, where the negative end of that molecule is electrostatically attracted to the positive end of the other molecule. This attraction is what causes Van der Waals. * Permanent Dipoles (e.g. Hydrogen chloride) – Some molecules have a polar bond, due to the electronegativity of the species in the molecule, and if all the polar bonds act in the same direction, the molecule is polar, i.e. it has a ᵟ+ and ᵟ- end. This can be attracted to another molecule that has opposite orientation. n) **Hydrogen bonding is an attraction between an electron deficient hydrogen atom, and a lone pair of electrons on a highly electronegative atom on a different molecules.** There are only a few species in which hydrogen bonding can occur: nitrogen, oxygen and fluorine. Some examples of molecules which have hydrogen bonding are H2O and NH3 o) Water has some anomalous properties for its simple molecular structure, which can be explained by hydrogen bonding: * Water is denser than ice – this is because ice is an open lattice with hydrogen bonds holding the water molecules apart. When ice melts, the hydrogen bonds collapse, allowing the water molecules to move closer together, which increases the density of water. * Water has a relatively high freezing and boiling point – The hydrogen bonds are much stronger intermolecular forces than van der Waals, which are found in most simple molecules. As water has hydrogen bonds, more energy is required to overcome these than if there were just Van der Waals. p) **Metallic bonding is the attraction of positive metal ions to delocalised electrons. ** q) r) s) **MODULE 3** a) The periodic table is arranged in order of increasing atomic number, which denotes the number of protons in the nucleus of an atom of a given element. It is arranged in groups and periods, with each period having repeating trends in physical and chemical properties and each group having similar physical and chemical properties. b) **Periodicity is the regular periodic variation of properties of element s with atomic number and position in the periodic table.** c) In each group, elements have the same number of outer shell electrons, as well as having the same kind of orbitals in the outer shell, which results in them having similar properties. The repeating pattern of similarity is caused by the repeating pattern of similar electronic configuration. d) **DESCRIBE AND EXPLAIN THE VARIATION OF THE FIRST IONISATION ENERGIES OF ELEMENTS** e) In periods two and three, as you go across the period, the number of protons in the nucleus of the atom increases. As the atoms are neutral, this means that the number of electrons surrounding the nucleus also increases.