2. Bonding & Molecular Structure Physical Science K Warne

3. Atoms and molecules (simple and giant) Carry out all of the following tasks with respect to the following elements & compounds: oxygen, water, petrol, sugar, carbon (diamond & graphite), table salt and copper metal.  Find out which of these (or any other) substances found in atomic form.  Define a molecule.  Write formula for each of these substances using: o Molecular formula for simple molecules o Empirical formulae for giant molecules  Draw these molecules showing both space filling and ball and stick representations.  Group these substances into: o Simple molecules (relatively few atoms per molecule): o Giant molecules (millions of atoms per molecule): Describe the difference between inter and intra molecular forces. State which of these forces are involved in phase changes. Explain the boiling point of water in relation to the other hydrides of group six..  Link the macroscopic properties of these substances to their microscopic (particle) structure.

4. Two different types of bonds occur in water. Intramolecular –Between hydrogen atoms and oxygen atoms inside the molecules Covalent bonds Intermolecular –Between two different water molecules Hydrogen bonds Intermolecular forces O ++ -- H H O ++ --HH O ++ -- H H

5. Bond Polarity in Water The oxygen atom has greater electronegativity so it is surrounded by greater electron density than the hydrogen atoms. O H H -- ++ ++ ++ -- The water molecule is a DIPOLE - it has two oppositely charged “poles”. OHH ++++ ----

6. The oppositely charged poles on the water molecules attract each other. This electrostatic attraction constitutes Hydrogen bonding. It is the strongest form of intermolecular attraction. Hydrogen bonding exists only between molecules in which hydrogen is bonded to a very electronegative atom (H-O-X, H- NX 2 or H-F). Result in abnormally high boiling points. Hydrogen Bonding O ++ -- H H O ++ --HH O ++ -- H H Example: Water

7. Molecular solids Iodine - I 2 Strong covalent bonds Weaker intermolecular bonds Covalently bonded molecules held together by weaker intermolecular bonding. PROPERTIES Low melting points Brittle Soluble in non-polar solvents Non - conducting

8. IMF vs Mp & Bp Halogen X 2 (diatomic) Molecular Mass (M r g.mol -1 ) Mp / Bp ( o C) Flourine F 2 pale yellow gas 19 x 2 = 38 -220 / -188 Chlorine Cl 2 pale green gas 35.5 x 2 = 71 -101/ -35 Bromine Br 2 red volatile liquid 80 x 2 = 160 -7 / 59 Iodine I 2 purple solid - sublimes 127 x 2 = 254 114 / 184 Mp & Bp increases with SIZE and molecule MASS.. : X..... : X.....

9. Density Density = mass/volume (g.cm -3 ) High density – solids Many particles per cm 3 Low density – gases Few particles per cm3 1cm x 1cm x 1cm = 1cm 3

10. Effect of Temp average kinetic energyTemperature is a measure of average kinetic energy. As temperature rises more particles have high energy. Maxwell-Boltzman Curve Average E k Increases

11. Effect of Temp Temperature is a measure of average kinetic energy. As temperature rises – particles move faster – and further apart – substances expand – used in a thermometer. Maxwell-Boltzman Curve Average E k

12. Viscosity Viscosity is a measure of how thick (viscous) and sticky a liquid is.Viscosity is a measure of how thick (viscous) and sticky a liquid is. Viscosity reduces the ability of a liquid to flow.Viscosity reduces the ability of a liquid to flow. Liquids that flow readily (water) have a low viscosity.Liquids that flow readily (water) have a low viscosity. Viscosity is a function of (depends on) the attractive forces of the molecules of the liquid.Viscosity is a function of (depends on) the attractive forces of the molecules of the liquid. Strong forces – high viscosityStrong forces – high viscosity Temperature also greatly affects viscosity: as temperature increases, viscosity decreases.Temperature also greatly affects viscosity: as temperature increases, viscosity decreases. Kinetic energy enables particles to overcome forces.

13. Bonding & Molecular Structure Atoms and molecules (simple and giant)   State that the only substances found in atomic form are the noble gases  Describe a molecule as a group of two or more atoms that are attracted to each other by relatively strong forces or bonds  Give examples of molecules based on the above description e.g. o Small molecules (relatively few atoms per molecule): oxygen, water, petrol, sugar, o Giant molecules (millions of atoms per molecule): diamond, a sodium chloride crystal, a metal crystal like a piece of copper, …  Recognize molecules from models (space filling, ball and stick, …)  Draw diagrams to represent molecules using circles to represent atoms  Represent molecules using o Molecular formula for small molecules, e.g. O 2, H 2 O, C 8 H 18, C 12 H 22 O 11, … o Empirical formulae for giant molecules, e.g. C, NaCℓ, Cu,  Give the formula of a molecule from a diagram of the molecule and vice versa Describe the difference between inter and intra molecular forces.  Linking macroscopic properties of materials to micro (particle) structure