8.4 Water

1. 8.4 Water

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

7. There are two pairs of non-bonding electrons There are 2 bonding pairs of electrons Water molecules • A water molecule has an oxygen atom covalently bonded to two hydrogen atoms

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 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:

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:

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.

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.

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. 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)