Edexcel Chemistry - Topic 2: Bonding and Structure

  • Created by: Ryan C-S
  • Created on: 15-03-18 23:40

Ionic Bonding

  • The electrostatic force of attraction between oppositely charged ions formed by electron transfer.
  • Metal atoms lose electrons to form +ve ions
  • Non-metal atoms gain electrons to form -ve ions
  • Ionic Crystals have a giant lattice structure
  • Ionic bonding is stronger when there are smaller ions and larger charges involved
  • Physical properties include:
    High Melting Points (Strong attractive forces between ions)
    Non-conductor of Electricity when Solid (Ions held tightly together and cannot move)
    Conductor of Electricity when in Solution or Molten (Ions free to move)
    Brittle - Moving the ions along will cause ions to be next to similar ions. Like charges cause repulsion so they push they layers apart
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Evidence for Ionic Bonding

  • X-ray diffraction shows the likelihood of finding electrons in a region. The contours are lines of equal electron density.
  • Maps for an ionic compound show discrete or seperate circles as electron density falls to zero between the ions and the ions are arranged in a regular pattern.
  • Electron density maps don't show the edge of an ion so it isn't possible to measure an ionic radius
  • Attaching electrodes to a U-tube containing CuCrO4 solution results in a blue colour of (Cu)2+ accumulating aroung the -ve electrode and a yellow colour of (CrO4)2- accumulating around the +ve electrode
  • Placing a drop of Potassium Manganate solution on a piece of filter paper on a microscope slide and attaching it to a 24V DC power supply will cause the purple colour of the (MnO4)- ion will move towards the +ve electrode
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Covalent Bonding

  • The Electrostatic force of attraction between the shared pair of electrons and two nuclei
  • Giant covalent molecules have high melting points as they contain several strong covalent bonds in a macromolecular structure so require a lot of energy to break the multiple bonds
  • X-ray diffractions for a covalent bond have a significant electron density between the atoms
  • Double and Triple Bonds have a greater electron density between them so they have a greater force of attraction between the nuclei and electrons resulting in a shorter bond length and greater bond strength

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Dative Covalent Bonding

  • Forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms - also referred to as a co-ordinate bond
  • Common examples include (NH4)+ and (H3O)+
  • Dative Covalent Bonds are represented by an arrow going from the atom that is providing the lone pair of electrons
  • A dimer is formed when two molecules covalently bond together with dative covalent bonds e.g. Al2Cl6

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VSEPR Theory

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory is used to determine the shape and structure of covalent molecules
  • When explaining VSEPR Theory, the structure is:

1) State number of bonding pairs of electrons
2) State number of lone pairs of electrons
3) State that electron pairs repel to reach a position of minimal repulsion
4) If there are no lone pairs, the electron pairs repel equally
5) If there are lone pairs, the lone pairs repel more than bonding pairs
6) State shape and bond angle

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Shapes of Molecules

Linear (2bp; 0lp) -                      180 degrees e.g. CO2

Trigonal Planar (3bp; 0lp) -        120 degrees e.g. BF3

Tetrahedral (4bp; 0lp) -              109.5 degrees e.g. CH4

Trigonal Pyramidal (3bp; 1lp) -   107 degrees e.g. NH3

Bent (2bp; 2lp) -                          104.5 degrees e.g. H2O

Trigonal Bipyramidal (5bp; 0lp) - 120 and 90 degrees e.g. PCl5

Octahedral (6bp; 0lp) -                 90 degrees e.g. SF6

Square Planar (4bp; 2lp) -           90 degrees e.g. XeF4

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Electronegativity

  • Electronegativity is the Relative Tendency of an atom to attract a bonding pair of electrons in a covalent bond towards itself.
  • Electronegativity is measured on the Pauling Scale (ranges from 0.0 to 4.0)
  • F, O, N and Cl are the most electronegative elements
  • Electronegativity increases across a period as there are more protons so the nuclear charge is stronger and the atomic radius is shorter so the bonding electrons are attracted more.
  • Electronegativity decreases down a group as there are more shells of electrons which increases shielding which weakens the nuclear charge so the bonding electrons are attracted less.
  • The polarity of a bond depends on the difference in electronegativities of the bonding atoms
  • 0.0 to 0.3 = Nonpolar Covalent bond
  • 0.3 to 1.7 = Polar Covalent bond
  • 1.7 to 4.0 = Ionic bond
  • If a molecule is symmetrical but contains polar bonds it will be a non-polar molecule overall as the polarities cancel out
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Van der Waals Forces

  • Van der Waals Forces (or London Forces) form in all molecular substances due to the random movement of electrons within atoms creating temporary differences in charges as electron densities fluctuate resulting in a slightly positive and slightly negative side to each atom.
  • The temporary charges can cause dipoles to form in neighbouring molecules. These are called induced dipoles.
  • Stronger Van der Waals forces occur when there are more electrons as there is a higher chance that temporary dipoles will form.
  • This means more energy is required to break them apart so boiling points get higher - this can explain the trend in boiling points down the Group 7 series and in alkanes.
  • Straight chain molecules also have a higher boiling point because there is a larger surface area of contact for Van der Waals force to form on compared to branched molecules.
  • Ionic compounds don't form Van der Waals forces
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Permanent Dipole-Dipole Forces

  • Permanent dipole-dipole forces occur between polar molecules
  • Stronger than Van der Waals forces so compunds have higher boiling points
  • Polar molecules have a permanent dipole e.g. C-Cl, C-F, C-Br or H-Cl bonds
  • Polar molecules are asymmetrical and have a bond where there is a significant differenence in electronegativity between the atoms.
  • Permanent dipole forces occur in addition to Van der Waals forcesSee the source image
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Hydrogen Bonding

  • Occurs in compounds that have a hydrogen atom attached to either Nitrogen, Oxygen or Fluorine (which must have a lone pair of electrons)
  • The is a large difference in electronegativity difference between the H and the O/N/F
  • Hydrogen bonds have a bond angle of 180 degrees around the H atom because there are two pairs of electrons around the H atom involved in the bond. These electrons repel to a position of minimal repulsion.
  • Alcohols, carboxylic acids, proteins, amides and amines can all form hydrogen bonds
  • Hydrogen bonds are strong intermolecular forces so have higher boiling points
  • Water can form multiple hydrogen bonds so requires more energy to break the bonds and therefore means the boiling point of water is a lot higher.
  • Hydrogen bonds cause water molecules to be held further apart when in a solid state which reduces the density of ice allowing it to float on liquid water.
  • H2O, NH3 and HF have anomalously high boiling points due to hydrogen bonding.
  • The general trend down a group for boiling points is that of an increase because there are more electrons so therefore there are stronger Van der Waals forces.
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Solvents and Solubility

  • Solubility of a solute in a solvent requires a balance of energy to break bonds in the solute and solvent against energy released when making new bonds between the solute and solvent
  • When an ionic lattice dissolves in water the negative ions are attracted to the δ+ hydrogens on the polar water and the positive ions are attracted to the δ- oxygen.
  • The higher the charge density, the greater the hydration enthalpy (smaller ions/large charges) as the ions attract the water molecules more strongly.
  • Solutes with similar intermolecular forces to those in the solvent will generally dissolve e.g. non-polar solutes will dissolve in non-polar solvents.
  • Propanone is a useful solvent as it has polar and non polar characteristics so can dissolve a wide range of solutes
  • Smaller alcohols are soluble in water because they form hydrogen bonds with water. Longer chains cause Van der Waals force to take precedent and become less soluble in water.
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Metallic Bonding

  • The electrostatic force of attraction between positive metal ions and delocalised electrons
  • Metallic bonds are stronger if:
    There are large numbers of protons (stronger nuclear attraction)
    There are large numbers of delocalised electrons per atom
    The ion is smaller
  • Metals have high melting points because the strong electrostatic forces between metal ions and the sea of delocalised electrons require a lot of energy to break
  • Magnesium has stronger metallic bonding than sodium because it is smaller and has more protons and electrons so its melting point is higher
  • Metals can conduct electricity well because electrons can flow through the structure
  • Metals are malleable and ductile because the positive ion in the lattice are identical so planes of ions can easily slide over one another.
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Macromolecular Structures

  • Giant lattices (Macromolecular structures) are present in all three types of bond
  • Ionic: NaCl - alternating positive and negative ions
  • Giant Covalent: Diamond - tetrahedral arrangement of carbon atoms
  • Metallic: Magnesium - closely packed structure of repeating positive ions
  • Molecular: Iodine - molecules held together by weak Van der Waals forces in a solid structure
  • Ice - tetrahedral arrangement of water molecules held together by hydrogen bonds

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Carbon Allotropes

  • An allotrope is a different arrangement of the same atoms of a chemical element to form structures with different properties
  • Diamond is a tetrahedral arrangement of Carbon atoms. It can't conduct electricity because all the electrons are localised and cannot move. 
  • Graphite is a planar arrangement of carbon atoms with 3 covalent bonds per atom in each layer. The fourth valence electron is delocalised. Graphite can therefore conduct electricity well as electrons are free to move between the layers. One layer is called graphene.
  • Carbon nanotubes have high tensile strengths due to the covalent bonds, can conduct electricity due to the delocalised electrons

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Structure and Properties

See the source image (http://www.4college.co.uk/a/aa/bonding2.gif)

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great resource

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