CHEM TEST 15/11/16

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  • Chem test 15/11/16
    • Definitions
      • Electron orbital
        • Represents a region of space around the nucleus that holds 2 electrons.
      • Ionisation energy
        • Amount of energy needed to remove one electron form each atom in a mole of atoms of an element in gaseous state.
      • Dative covalent bond
        • A shared pair of electrons in which the bonded pair has been provided by one of the atoms only.
      • Electronegativity
        • The attraction of a bonded atom in a covalent bond.
      • Ionic bonding
        • Electrostatic attraction between positive and negative ions.
      • Covalent bonding
        • Strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
      • Intermolecular forces
        • Weak interactions between dipoles of diffent molecules.
    • Electron structure
      • Shells
        • 2 e- N=1
        • 8 e- N= 2
        • 32 e- N=4
        • 18 e- N=3
        • 3rd shell - 3 subshells
          • 5 orbitals in d 10 e-
        • 1st shell - 1 subshell
          • 1 orbital in s - 2e-
        • 4th shell -4 subshells
          • 7 orbitals in f 14 e-
        • 2nd shell - 2 subshells
          • 3 orbitals in p- 6e-
        • S subshell - spherical shape
        • P subshell has 3 different dumbbell shapes
      • Quantum theory
        • Energy is transferred in fixed amounts.
        • A packet of energy is a quantum.
        • Packets of energy is a quanta.
      • Stability of sub shells
        • A full subshell makes the element more stable compared to half filled.
        • Cr and Cu are different.
          • Cr is actually (Ar) 3d 5, 4s 1
          • Cu is actually (Ar) 3d 18 4s 1
        • 2p orbitals have slightly higher energy than 2s orbitals.
        • Why a drop between group 2 and 3
          • Boron has a 2p orbital. beryllium has a 2s orbital. Boron therefore has a lower energy than beryllium.
      • Dative covalent bonds
        • Carbon monoxide: C=O Oxygen gives two e- to C.
    • Ionisation energy
      • Factors affecting ionisation energy
        • Increase In nuclear charge
          • As electrons are ripped off, you have more positive charge.
            • This is because there will be more protons than neutrons.
        • Distance from the nucleus.
        • Shielding
          • e- repel each other. The e- in inner shells repel the outer shells and this reduces the positive nuclear charge.
            • The greater the shielding, the less ionisation energy is needed to remove it.
      • Flame tests
        • Sodium - yellow
        • Calcium - brick red
        • Strontium - pale pink
        • Barium - apple green
        • Copper- green/blue
        • Potassium - lilac
        • Lithium - red
      • Most/least reactive
        • Alkali metals - most reactive, low IE
        • Noble gases,Least reactive, high IE
      • Trends
        • IE increase across a period.
    • Charges on ions
      • Carbonate CO3 2-
      • Nitride N3-
      • Sulphate SO4 2-
    • Structure
      • Shapes and angles
        • Shapes
          • 4 bonded pairs, tetrahedral, 109.5 eg CH4
          • 3 bonded pairs, 1 lone pair, pyramidal, 107 eg NH3
          • 2 bonded pairs, two lone pairs, non-linear, 104.6 eg h2O
          • 2 electron pairs, linear, 180 eg CO2
          • 3 electron pairs, trigonal planar, 120, eg BF3
          • 6 electron pairs, octanedral, 90, eg SF6
    • Electronegativity and polarity
      • Bonded electron pair
        • It is shared evenly unless nuclear charges are different, atoms are different sizes, shared electrons may be closer to one nucleus than the other.
      • Measuring it
        • Pauling electronegativity values. Across the periodic table, the nuclear charge increases and the atomic radius decreases.
        • Ionic or covalent? If the electronegatitivty difference is 0 it is covalent, 0 to 1.8 is polar covalent and anything over 1.8 is ionic.
      • bond polarity
        • Non- polar bonds
          • The bonded electron pair is shared equally between the bonded atoms.
          • A bond will be non-polar when the bonded atoms are the same or the bonded atoms have the same or similar electronegativity
          • When a bond is formed between H, O or Cl, the bonded elements are the same so this forms a pure covalent biond
        • Polar bonds
          • The bonded electron pair is shared unequally
          • A bond is polar when the atoms are different and have different electronegativity values.
            • This then forms a polar covalent bond,
        • Polar molecules
          • When a permanent dipole is upon a compound
    • Bonding
      • Ionic
        • Anions - cl- No3- eg
        • Cations - Na+, Ca2+ eg
        • Giant ionic lattice
        • Solid at RT- not enough energy to overcome the dstrong electrostatic forces of attraction.
        • High melting and boiling points
        • Dissolve in polar solvents (like water)
          • To be soluble, the ionic lattice must be broken down and the water molecules must attract and surround the ions.
        • solids don't conduct electricity. Once molten or dissolved it does.
      • Covalent
        • It is localised - attraction only between the two atoms
        • Happens between elements like H2, H2O and NH4+
    • Intermolecular forces
      • Induced dipole-dipole interactions (London forces)
        • Very weak but are everywhere
        • Movement of electrons produces a changing dipole in a molecule. An instantaneous dipole will appear and induces a dipole on a neighbouring molecule.
        • The more electrons in each molecule: The larger the instantaneous and induced dipoles., the greater the induced dipole-dipole, the stronger the attractive forces.
      • Simple molecular substances
        • Made up of simple molecules. In a solid state, simple molecules form a regular structure called a simple molecular lattice,
          • Molecules are held by weak intermolecular forces. The atoms within each molecule are bonded strongly by covalent bonds.
        • Properties
          • Low melting and boiling point. Simple molecular lattices are  broken down very easily.
          • When a simple molecular lattice is broken apart, only weak intermolecular forces  break. The covalent bonds don't break.
          • Solubility
            • Non- polar: When a simple molecular is added to a non-polar substance, intermolecular forces form between the molecules and the solvent. The interactions weaken the intermolecular forces  and this makes the intermolecular forces break and the compound dissolves.
            • Polar substances:  Polar covalent substances may dissolve in polar solvents as the polar solute molecules and the polar solvent molecules can attract each other.
          • There are no mobile charged particles so it cannot conduct electricity.
      • Hydrogen bonding
        • A special type of permanent dipole-dipole interaction with elements containing O,N,F
        • Strongest type of intermolecular attraction.
        • Solid (ice)m is less dense than the liquid (water). Hydrogen bonds hold water molecules apart in an open lattice structure.. The water molecules in ice are further apart than in water.  Solid ice is less dense than liquid water so it floats.
        • Water has a high M/BP this is because H+ forces are extra on top of London forces. A lot of energy is needed to break the hydrogen bonds.
        • Water has relatively high surface tension and viscosity.

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