OCR Chemistry A F321

Mindmap on all of F321 

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  • Created by: Tulsi
  • Created on: 23-12-12 15:02
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  • 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

      Comments

      Kathryn2


      It's really informative, but you forgot to mention the f subshell, and that it goes 4s then 3d in order of energy levels.

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