# OCR Chemistry A F321

Mindmap on all of F321

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• Created by: Miss T
• Created on: 23-12-12 15:02
• F321
• Atomic structure
• The nucleus contains protons (+ive charge) and neutrons (no charge)
• The atomic number is the number of protons in an atom's nucleus
• The mass number is the total number of protons and neutrons in the nucleus
• Ions do not have the same number of electrons as protons and so have an overall charge
• Isotopes and Relative masses
• isotopes are atoms having the same number of protons but different numbers of neutrons
• The relative atomic mass is the weighted mean mass of an atom relative to 1/12th of the mass of carbon- 12
• The average relative mass is equal to the sum of each isotope's  mass for an element, x, it's relative abundance
• The relative mass of a compound is equal to the sum of the individual relative atomic masses
• The Mole
• A mole is the SI unit for amount of substance and has units of mol
• One mole of substance is simply the relative formula mass for an element in grams
• The Emperical Formula is the simplest whole number ratio of atoms in each element present in a compound
• The molecular formula is the actual number of atoms of each element in a molecule
• Calculations using the mole
• Mass calculations : n = m/M
• Gas volumes: n = v(in dm^3)/ 24.0   OR n = v(in cm^3) /24000
• 1 mole of any gas occupies 24dm^3 @ RTP
• Gas volumes: n = v(in dm^3)/ 24.0   OR n = v(in cm^3) /24000
• solution calculations: n= v(in dm^3) x c OR n = v(in cm^3)/1000 x v
• A dilute solution consists of a small amount of dissolved solute, A concentrated solution consists of a large amount of solute.
• Calculations using the mole
• Mass calculations : n = m/M
• 1 mole of any gas occupies 24dm^3 @ RTP
• solution calculations: n= v(in dm^3) x c OR n = v(in cm^3)/1000 x v
• A dilute solution consists of a small amount of dissolved solute, A concentrated solution consists of a large amount of solute.
• Acids and bases
•   An acid is a hydrogen ion (H+) or proton donor in solution, whereas a base is a hydrogen ion or proton acceptor in solution.
• Hydrochloric acid (HCl), sulfuric acid (H2SO4) and nitric acid (HNO3) are common acids.
• Bases include metal oxides (e.g. MgO), metal hydroxides (e.g. NaOH) and ammonia (NH3).
•  Alkalis are soluble bases and form hydroxide ions, OH-, in solution.
• Reactions of acids and bases
• Acids and bases
•   An acid is a hydrogen ion (H+) or proton donor in solution, whereas a base is a hydrogen ion or proton acceptor in solution.
• Hydrochloric acid (HCl), sulfuric acid (H2SO4) and nitric acid (HNO3) are common acids.
• Bases include metal oxides (e.g. MgO), metal hydroxides (e.g. NaOH) and ammonia (NH3).
•  Alkalis are soluble bases and form hydroxide ions, OH-, in solution.
• Salts are formed when a hydrogen ion from the acid is replaced by a metal ion, or an ammonium ion.
• Acids react with bases to form a salt and water only; they react with metal carbonates to form a salt, water and carbon dioxide gas.
• Metals react with acids to form a salt and hydrogen gas.
• Salts may chemically combine with water as water of crystallisation in hydrated salts. (Without water in anhydrous salts.)
• Oxidation numbers
• An oxidation number indicates the formal charge of a chemically combined particle in a compound.
• The oxidation number of metals usually equals the group number (as a positive value) and minus (8 – group number) for non-metals.
• An element has been oxidised if the oxidation number increases, and reduced if the oxidation number decreases.
• When they react, metals are normally oxidised (they lose electrons), whereas non-metals gain electrons and are reduced.
• Oxidation Is Loss Reduction Is Gain
• An element has been oxidised if the oxidation number increases, and reduced if the oxidation number decreases.
• Electronic Structure
• Electrons occupy energy levels around the nucleus of the atom, where each shell has a principal quantum number.
• For principal quantum number, n = 1, the number of electrons is 2; for n = 2, the number is 8; then 18; then 32 electrons for n = 4.
• Main energy levels are sub-divided into sub-shells and these consist of orbitals called s, p and d-orbitals
• Elements have an electronic configuration that can be shown in s, p or d notation, for example, sodium is 1s2, 2s2, 2p6, 3s1.
• Chemical Bonding
• Ionic bonding takes place when positive ions and negative ions are attracted in a giant ionic structure
• Covalent bonding is the sharing of electron pair(s) between nuclei of atoms
• The covalent bond and ionic bond are both very strong chemical bonds.
• A dative covalent bond is one formed in which both electrons are donated from the same atom.
• Molecular shapes
•  The shape of a molecule is determined by the repulsion between bonded electrons and non-bonded electrons (lone pairs).
• Lone electron pairs repel more than bonded pairs of electrons and give rise to distorted shapes
• By deducing the number of bonded electron pairs and lone pairs of electrons, the shape of a molecule may be predicted
• BF3 is trigonal planar; CH4 and NH4+ are tetrahedral; SF6 is octahedral; H2O is non-linear (V-shaped/bent); CO2 is linear and ammonia, NH3, as pyramidal
• Intermolecular forces
• Hydrogen bonding arises in molecules in which a hydrogen atom is bonded to either an N or O atom
• Electronegativity is the ability of an atom in a covalent bond to attract a bonded pair of electrons towards itself.
• Water molecules, and other substances consisting of hydrogen bonding, have anomalous properties as a result
• An intermolecular force exists between molecules and may include hydrogen bonding, dipole-dipole or van der Waals’ forces
• Bonding and physical properties
• Giant structures have high melting and boiling points due to strong chemical bonds acting throughout the structure
• Metals are very good electrical conductors as a result of having mobile electrons
• Giant ionic structures conduct electricity when molten, and when dissolved in water due to mobile ions, not electrons
• Metals consist of a close-packed arrangement of positive ions, through which delocalised electrons move
• Periodicity
• In the Periodic Table, ionisation energies increase moving across a period from left to right, and decrease moving down a group
• Electron structures, atomic radii, melting points and boiling points all show periodicity
• Elements in the same group have similar chemical and physical properties
• When the elements are arranged in order of their atomic number, there is a regular repetition of physical and chemical properties
• Group 2 elements – the alkaline earth metals
• Metal hydroxides are weak alkalis and typically have a pH between 8 and 11
• Reactivity increases on descending the group because the outer two electrons are further from the nucleus and are less shielded
• Group 2 carbonates decompose with greater difficulty as the group is descended, to form the metal oxide and carbon dioxide gas
• These elements all react with water to form a solution of the hydroxide and hydrogen gas. These elements react with oxygen to form the oxide
• Group 7 elements – the halogens
• Halogen atoms gain one electron to form halide ions, X-, and this ability becomes easier on moving up the group
• Halogens dissolve in organic solvents, like hexane, to form characteristic colours, for example, iodine forms a purple solution
• Halogen atoms become larger on descending the group, so a gained electron is only weakly attracted due to greater shielding
• All halogens exist as diatomic molecules in which van der Waals’ intermolecular forces act between the molecules
• Reactions of chlorine and halide ions
• Chlorine is used to kill germs in water supplies, but is also toxic to humans at higher doses
• Chlorine also disproportionates in cold, aqueous sodium hydroxide to form sodium chloride, sodium chlorate(I) and water
• Halide ions are detected with silver(I) nitrate solution and the subsequent reaction with ammonia solution
• Chlorine disproportionates in water to form hydrochloric acid and chloric(I) acid, the latter being an oxidising agent