2.2 Electrons, Bonding and Structure

2.2.1 Shells and Orbitals
1. s-orbitals
  1. Spherical shape
  2. From n=1, each shell contains 1 s orbital
  3. Giving 2s electrons in each shell
2. p-orbitals
  1. has a 3D dumb-bell shape
  2. Present from n=2 upwards
  3. Each shell has 3 p-orbitals at right angles to each other
  4. Each p-orbital can hold 2 electrons, giving a possible total of 6p electrons
3. d-orbitals and f-orbitals
  1. From n=3 upwards, each shell contains 5 d-orbitals, giving a possiblesum of 10d electrons
  2. From n=4 upwards each shell contains 7 f-orbitals, giving a possible total of 14f electrons
2.2.2 Sub-shells and energy levels
1. Sub-shells
  1. Have the same principle quantum number
  2. There are s,p,d and f subshells as each is only made of one type of atomic orbital
2. Filling shells and sub-shells
  1. Lowest energy filled first
    1. An energy level must be full before a higher starts to fill
  2. Sub-shells are made of several orbitals, each with the same energy level
    1. Each orbital is filled singly before pairing starts
    2. The 4s is at a slightly lower energy level than 3d, so 4s fills first
3. Electron configuration
  1. Written in the form nx^y
    1. n=shell number
      1. x=type of orbital
      2. y=the number of electrons in the orbitals making up the sub-shell
  2. Electron configuration of ions
    1. Ions are formed by gaining or loosing electrons
    2. If they form a positive ion then the electrons found in the highest energy levels are lost first
2.2.3 Chemical bonding
1. Chemical reactions are accompanied by electron transfers and often bem=come more stable by transferring electrons or combining electrons with other atoms
  1. The most stable and unreactive elements are the noble gases as they have a full outer shell, other elements react to try and achieve the same configuration
    1. The 8 outer shell electrons are made from 2 in the s and 6 in the p
2. 3 main types of chemical bonding
  1. Ionic bonding
    1. Metal and a non-metal
    2. Electrons are transferred from the metal to the non-metal, to form oppositely charged ions that attract each other
  2. Metallic bonding
    1. Occurs in metals
    2. Electrons are shared between all the atoms
  3. Covalent bonding
    1. 2 non-metals
    2. Electrons shared between the atoms and are attracted to the nuclei of both bonded atoms
2.2.4 Ionic bonding
1. Ionic bonds
  1. Electrons are transferred
  2. Oppositely charged ions are formed
  3. Metal ion is positive
  4. Non-metal ion is positive
2. Giant ionic lattices
  1. Each ion is surrounded by oppositely charged ions
  2. Ions attract from all directions, forming a 3D lattice
2.2.5 Structures of ionic compounds
1. Due to the huge amounts of electrostatic attraction, they have uniques properties
2. High melting and boiling points
  1. Solids at room temperature as a large amount of energy is needed to break the electrostatic bonds that hold the opposingly charged ions together in the lattice
  2. The greater the charge, the stronger the electrostatic forces between the ions, meaning more energy is required to break up the ionic lattice during melting
3. Electrical conductivity
  1. When solid they don't conduct electricity as the ions are held in fixed positions and so cannot move and therefore cannot conduct electricity
  2. When dissolved or melted the solid lattice breaks down and the ions are free to move, so it can now conduct electricity
4. Solubility
  1. An ionic lattice dissolves in polar solvents such as water, polar substances contain substances that have polar bonds
  2. Polar water molecules break down an ionic lattice by surrounding each ion to form a solution
    1. The slight charges of the polar substances are able to attract the charged ions in the giant ionic lattice, meaning the lattice is disrupted and ions are pulled out of it
2.2.6 Covalent bonding
1. An electron occupies the space between the 2 atoms' nuclei, this shared pair of electrons forms a covalent bond
  1. The attraction overcomes the repulsion between the two positively charged nuclei
2. Single covalent bonding
  1. If atoms are bonded by one shared pair of electrons then it is known as a single bond
    1. Ionic bonds act in all directions, covalent only in one
3. Lone pairs
  1. A pair of outer-shell electrons that aren't involved in chemical bonding
4. Average bond enthalpy
  1. Not all covalent bonds are the same strength
2.2.7 Dative covalent bonding
1. One atom supplies both the shared electrons in a covalent bond
2. Examples
  1. The ammonium ion, NH4+
    1. 3 covalent bonds and 1 dative
    2. One N has a lone pair and bonds to an H+
    3. Overall ion has a 1+ charge
  2. The oxonium ion, H3O+
    1. Acid+water forms oxonium ions
      1. One of the lone pairs from the O in H2Oforms a dative covalent bond
2.2.8 Structures of covalent compounnds
1. 2 possible structures
  1. Giant covalent
    1. High melting and boiling point as high temperatures are needed to break the strong covalent bonds within the lattice
      1. They cannot conduct electricity as they have no free charged particles (except for graphite)
        They are insoluble in both polar and non-polar solvents as the covalent bonds are too strong to be broken by either polar or non-polar solvents
  2. Simple molecular
    1. Low melting and boiling points are the intermolecular forces are weak and so relatively little energy is needed to break them
      1. Not conductors of electricity as there are no free charged particles
        Generally soluble in non-polar solvents such as hexane. Weak London forces are able to form between covalent molecules and the solvents, helping the molecular lattice to break down and the substance dissolves
2.2.9 Shapes of molecules and ions
1. Electron repulsion theory
  1. The shape of a molecule or ion is determined by the number of electron pairs in the outer shell surrounding the central atom. These pairs can be bonding or lone pairs
    1. All electrons have a negative charge, each electron pair repels other electron pairs
      1. The shape adopted will be the shape that allows all the pairs of electrons to be as far apart as possible
2. Bonded pairs: 1 Name of shape: linear Angle:
  1. Bonded pairs: 2 Name of shape: linear Angle:180
    1. Bonded pairs: 3 Name of shape: trigonal planar Angle: 120
      1. Bonded pairs: 4 Name of shape: tetrahedral Angle: 109.5
        Bonded pairs: 5 Name of shape: trigonal bipyramidial Angle: 90 and 120
        Bonded pairs:6 Name of shape: octahedral Angle: 90
        Bonded pairs: 3 Lone pairs: 1 Name of shape: pyramidal Angle: 107
        Bonded pairs: 2 Lone pairs: 2 Name of shape: non-linear Angle: 1045
2.2.10 Electronegativity and bond polarity
1. Electronegativity measures the attraction of a bonded atom for the pair of electrons in a covalent bond
  1. Electronegativity increases towards the top right corner of the periodic table, fluorine being the most electronegative element
    1. If atoms in a covalent bond are different, one of the atoms is likely to be more strongly attracted to the bonding electrons. The bonding atom with a greater attraction for the electron pair is said to be more electronegative than the other atom
      1. This creates a small charge difference across the bond. This difference is always present and is called a permanent dipole. this is now a polar covalent bond
2. Polar molecules have polar bonds
  1. If a molecule is non-symmetrical it can create a charge difference across the whole molecule. In symmetrical molecules the dipoles of the bonds cancel out
    1. Oxygen is one of the most electronegative elements and bonds containing oxygen will be polar, but this doesn't mean the molecule will be polar
2.2.11 Intermolecular forces
1. Intermolecular forces occur due to constant random movements of the electrons within the shells of the atoms in molecules
  1. 2 main types of intermolecular force are hydrogen bonds and van der Waals forces
2. Van der Waals forces:
  1. London (dispersion) forces
    1. They are caused by random movements of electrons in atoms ' shells. This movement unbalances the distribution of charge within the electron shells
      1. At any moment there will be an instantaneous dipole across the molecule
        The instantaneous dipole induces a dipole in neighbouring molecules, which can further induce dipoles in neighbouring molecules
        Increases with the number of electrons
    2. The effect of London forces on boiling points;
      1. As the number of electrons increases, so does the strength of the London forces
  2. Permanent dipole- permanent dipole interactions
    1. Molecules with permanent dipoles will be attracted to other molecules with permanent dipoles. E.g. like a magnet
  3. Permanent dipole-induced dipole interactions
    1. Permanent dipoles can come from polar bonds being present. when this is near to a non-polar bond it can cause electrons in the shells of the nearby molecules to shift slightly, causing the non-polar molecule to become slightly polar and the attraction then occurs
2.2.12 Hydrogen bonding
1. A strong permenant dipole-permenant dipole interaction between an electron deficient hydrogen and the lone pair of a N,F or O molecule
2. Allows water to be denser than ice as the hydrogen bonds become fixed in an open structure in ice
3. Water has a high boiling point as the hydrogen bonds are much stronger than other intermolecular forces
4. Can explain high surface tension and viscosity of water

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2.2 Electrons, Bonding and Structure

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