1. 8.4 Water

2. Density of water Density = mass/volume
The density of liquid water is calculated using the formula:
Density = mass/volume
where: D = density (g/mL or g/m-3)
m = mass (g)
V = volume (mL)

3. Temperature and density
Generally as the temperature of a
substance increases it expands and its
volume also increases. This in turn results
in a decrease in density.
BUT:
As liquid water cools its density increases
until 4oC, however, below this temperature
density decreases. As a result the density
of ice is lower than the density of liquid
water

4. Melting and boiling points of H2O
Melting point: at 0OC ice and liquid water can exist in equilibrium with one another.
Boiling point: at 100OC liquid water and water vapour exist in equilibrium with one another.

5. The effect of solutes The presence of solutes can effect the
freezing and boiling points of water.
Generally the freezing point of water
decreases as the concentration of solute
increases. For example, salt water freezes
at a lower temperature that pure water.

6. The effect of solutes Normal freezing point of water 0OC
Normal boiling point of water
100OC
Freezing point decreases on adding solutes
Boiling point rises on adding solutes

7. Water molecules
A water molecule has an oxygen atom covalently bonded to two hydrogen atoms
There are 2 bonding pairs of electrons
There are two pairs of non-bonding electrons
Remember: covalent bond = shared electrons

8. Valence shell electron pair repulsion
The shape of a molecule depends on the arrangement of electron pairs around the central atom in the molecule
As a general rule: the electron pairs in the valence energy level of an atom repel each other and so are arranged as far apart as possible.

9. Electronegativity
Question: Do metals tend to attract electrons to obtain a stable valence shell?
Answer: No, they tend to give up electrons
Question: Do non-metallic elements tend to attract electrons to obtain a stable valence shell?
Answer: Yes

10. Electronegativity
The electron attracting ability of atoms is called electronegativity.
Because non-metals have a strong attraction for electrons they tend to have higher electronegativity than metals.
Electronegativities increase across a period to Group VII and decrease down the groups.

11. Electronegativity and polar bonds
Question: What type of bond is likely to form between two atoms that differ greatly in electronegativity (eg: Li & F) – ionic or covalent?
Answer: ionic, because of the strong tendency for an electron to be transferred completely from one atom to the other

12. Polar bonds
When a bond is formed between two identical atoms such as two Cl atoms in Cl2 the electrons are shared equally, and the bond is non-polar covalent.
When different types of atoms form a covalent bond, the electrons are shared unequally (eg: they spend more time near one nucleus than the other) and the bond is polar covalent

13. Polar bonds δ +H δ- Cl Example HCl: The electrons spend more
time near the Cl atom than the H atom,
because Cl is more electronegative. As a
result the Cl end of the molecule has a slight
negative charge while the H end is slightly
positive. This is written as:
δ +H δ- Cl
Note: δ (delta) means ‘small amount of’

14. Dipole
In the case of HCl the positive charge is equal and opposite to the negative charge.
δ +H δ- Cl
A pair of equal and opposite charges separated by some fixed distance is called a dipole.

15. Non- polar molecules Molecules can also be classified as polar
or non-polar. The shape of a molecule will
determine whether a molecule will be polar
or non-polar.

16. Non-polar molecules
If the shape of a molecule is such that the individual dipoles cancel out (eg: there is no net dipole) then the molecule is non-polar.
For example, BeF2:
BeF2 is linear. The Fl atoms are more electronegative and so attract electrons more strongly. The Be-F2 bonds are therefore polar.
The dipoles of the two bonds have the same magnitude but opposite directions. The sum of the dipoles is 0 and
the molecule is
non-polar.

17. Polar molecules
Molecules with an uneven charge distribution are called polar molecules.
Molecules are polar if the sum of their bond dipoles, does not equal zero. For example, H2O:
Uneven charge distribution or asymmetrical charge distribution results in a polar molecule

18. Summary When deciding whether a molecule is non-polar or polar:
Use electronegativities to decide whether it has any polar bonds.
Use the shape of the molecule to decide whether the polar bonds cancel out (non-polar) or produce a net dipole (polar).
If a molecule is symmetrical it will be non-polar, assymmetrical molecules will be polar.
Note: to have a polar bond there has to be a difference in electronegativity between the atoms

19. Intermolecular forces
Covalent molecular substances can have the
following types of intermolecular forces.
Dipole – dipole forces: form when the positive end of one polar molecule attracts
the negative end of
another.

20. Intermolecular forces
Dispersion forces: form when a temporary dipole (resulting from the uneven distribution of electrons around an atom or molecule) induces a dipole in a neighbouring atom. The two species are then attracted to each other.

21. Intermolecular forces
Hydrogen bonding: involves a hydrogen atom bonded to a O, N or F atom in one molecule becoming attached to O, N or F atom in a different molecule.

22. Intermolecular forces
These forces in order of increasing strength
are:
Dispersion < dipole-dipole < hydrogen bonds
The strength of these intermolecular forces
impacts on the melting and boiling point of
substances. As a general rule the weaker the
intermolecular forces the lower the melting
and boiling point.

23. Solutions For a solution to form:
The intermolecular forces between the solute molecules must be overcome
Intermolecular forces between some of the solvent molecules must be overcome
Intermolecular forces must form between the solute and solvent molecules

24. Solutions Generally a solute will dissolve in a solvent if:
the molecular forces within the solute and within the solvent are similar to those forming BETWEEN the solute and solvent molecules.
in other words ‘like dissolves like’ (eg: polar solvents tend to dissolve polar solutes)

25. Solubility of ionic compounds
Most ionic substances (eg: NaCl) are soluble in water because the polar water molecules
have a strong
attraction for
the charged
ions.
positively charge Na+ (cation) is attracted to the negatively charged end of the water molecule and the negatively charge Cl- (anion) is attracted to the positively charged end of the water molecule.
Aqueous solutions of ionic compounds produce ions in solution (eg: the ions become dispersed throughout the solution) this process is called dissociation.

26. Solubility of covalent molecular substances
Most soluble
Polar substances that are able to form hydrogen bonds with water are very soluble (eg: ammonia, glucose).
Covalent molecular substances that form dipole-dipole of dispersion forces with water are partially soluble.
Large covalent molecules do not dissolve in water due to the presence of a large non-polar portion in the molecule.
Note: Covalent molecular compounds may be non-polar or polar.
Non polar compounds eg: butane and carbon tetrachloride form dispersion forces with water and have low solubility
Polar molecular compounds such as chloroform form dipole-dipole forces with water and have slightly greater solubility.
Least soluble

27. Precipitation reactions
When solutions of two ionic substances
are mixed a precipitate (solid) will form if
one or more of the cation-anion
combinations produces an insoluble
substance.

28. Precipitation reactions
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
In the above reaction a white solid – silver
chloride – is produced. Note that this is
essentially a reaction between Ag +(aq) and
Cl-(aq). Therefore the net ionic equation for this
reaction is :
Ag +(aq) + Cl-(aq). AgCl(s)