Intermolecular forces (A level)

Kirsten Rowland
Slide Set by , created almost 3 years ago

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Kirsten Rowland
Created by Kirsten Rowland almost 3 years ago
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Slide 1

    Intermolecular Forces
    These: are weak electrostatic forces of attraction between neighbouring molecules are much weaker than covalent, ionic or metallic bonds only influence the physical properties of materials GIANT STRUCTURESCovalent (e.g. diamond) or ionic (e.g. NaCl) have high melting and boiling points. Imf (intermolecular forces) are not applicable because no molecules exist.SIMPLE MOLECULES(e.g. H2O, H2, CH4) have much lower melting points and boiling points because of imf.
    Imf influence physical properties: melting and boiling points solubility in water (and other solvents) 3D shapes of complex molecules e.g. DNA viscosity of liquids density And so on...Boiling points are good indicators of imf. Strong bonds within molecules are not broken when molecular substances are vapourised. Weak imf between molecules are broken when molecular substances are vapourised.Boiling point increase = Imf strength increaseBoiling point variations suggest 3 types of imf: Van der Waal forces Dipole-dipole forces Hydrogen bonds

Slide 2

    Hydrogen Bonding:
    You need two components for hydrogen bonding to occur: A VERY electronegative atom with a lone pair of electrons (only fluorine, oxygen and nitrogen meet this criteria) A H atom bonded directly to the electronegative atom The bonds are formed through the electronegative atom drawing hydrogen's electron away (de-shielding the hydrogen) making it slightly positive.A HYDROGEN BOND = the attraction between a lone pair of electrons on a N, O or F atom and a de-shielded H atom in a H-F, H-O or H-N bond.
    Caption: : Hydrogen bonding in water - the delta symbol by each atom means it is either slightly more positive or negative

Slide 3

Slide 4

    Caption: : NH3 has approximately one H bond per mole
    The strength of individual hydrogen bonds decreases with the electronegativity of each element BUT water forms two bonds per molecule so order of boiling point is: H2O >> HF > NH3 not HF > H2O > NH3.

Slide 5

    Because of it strong H bonds water has an unusually high surface tension (why if you make drops with a pipette they stay as little globules). This creates a "skin-like" effect on the surface of the water, allowing insects like the water skater to "walk" on water.Because of H bonds ice has an unusually low density compared to water. For ice: the maximum number of H bonds are operative molecules are held apart in a tetrahedral arrangement of covalent an H bonds there is much empty space between molecules it is a larger volume than he same mass of water it has a lower density than water Expansion during freezing of water can burst pipes. Also ice floats on water which is why fish (and actually every species on the planet) don't die.

Slide 6

    Dipole-Dipole Forces
    Dipoles in molecules arise from uneven electron distribution caused by differences in the electronegativity of atoms in the molecule.e.g.1 Hydrogen ChlorideH < Cl in electronegativity the bond electron is drawn towards the chlorine there is a permanent dipole in the molecule; delta + one end delta - the other the dipole-dipole force occurs between the delta + of one mole and the delta - of the neighbouring mole
    Caption: : Hydrogen chloride dipole

Slide 7

    e.g.2 Hydrogen Sulphide H < S in electronegativity this is a permanent dipole the dipole-dipole force is the same as in the chloride, except that the delta + is formed by a sort of electron cloud produced by the hydrogens in a neighbouring molecule It is important to note that weaker van der Waal forces also occur wherever dipole-dipole forces occur.
    Caption: : The free sulphur electrons are attracted to the slightly positive electron cloud created by the hydrogens in a different molecule of hydrogen sulphide

Slide 8

    Not dipoles?
    Caption: : BF3 molecule
    Compounds like CCl4, BF3 and BeCl2 do not contain dipole-dipole forces. With these, the individual bonds are polar (e.g. delta + Be-delta - Cl) but the molecules are not because they are symmetrical, meaning: bond dipoles cancel each other out they are non-polar molecules

Slide 9

    Van der Waal Forces
    Ne, I2 and CH4 are some examples of non-polar molecules. This means that the electron cloud of the molecule is in constant random motion, leading to momentary electron density imbalance which leads to a temporary dipole in the neighbouring molecule.THIS LEADS TO MOMENTARY ATTRACTION BETWEEN TEMPORARY DIPOLES WHICH IS THE VAN DER WAALS FORCE.Van der Waal forces get stronger as the number of electrons in the molecule decreases. For example, the boiling points of noble gases increase down the group as electron shells are added.Similarly, the boiling point of group IV hydrides (CH4, SiH4 etc.) increases down the group as Mr increases.

Slide 10

    PENTANE Formula = C5H12 Boiling point = 36 degrees Celsius Mr = 84 Straight chain molecule: CH3-CH2-CH2-CH2-CH3This means that: moles can get close there are strong Van der Waal forces it has a high boiling boint
    2,2, DIMETHYLPROPANE Formula = C5H12 Boiling point = 10 degrees Celsius Mr = 84 Branched chain molecule (imagine lines between the centre and the top and bottom molecules):                                                       CH3                                               CH3-C-CH3                                                      CH3This means that: moles cannot get close there are weak Van der Waal forces it has a low melting point