Bonding & Molecular Theory

1. Bonding & Molecular Theory Everything written in black has to go into your notebook Everything written in blue should already be in there

2. H H O H H Molecule of hydrogen (H2) 2 hydrogen atoms bonded together Hydrogen is an element Molecule of water (H2O) 2 hydrogen atoms and 1 oxygen atom bonded together Water is a compound

3. Ion • An ion is an atom or group of atoms who have an electric charge • An ion is an atom which has lost or gained electrons • A positive charge means it has lost electrons • A negative charge means it has gained electrons HL OL

4. Electropositive elements: tend to lose electrons when they form a bond Electronegative elements: tend to gain electrons when they form a bond Eg. sodium chloride (NaCl) Sodium loses an electron and chlorine gains an electron Na+ Cl-

5. Ionic Bond • An ionic bond is formed when one or more electrons are transferred from one atom to another

6. Cl- Na+ 11 P 12 N 17 P 19 N

7. O2- Mg 2+ 12 P 12 N 16 P 16 N

8. Properties of Ionic compounds • Usually crystalline solids • High melting and boiling points • Conduct electricity when molten or dissolved in water (because its ions are free to move) • Usually dissolve in water

9. Covalent Bond • A covalent bond is formed when atoms share a pair of electrons

10. E.g. chlorine (Cl2) 7 electrons on outer energy level (shell) Cl Cl

11. E.g. hydrogen gas (H2) 1 electron on outer energy level (shell) H H

12. If the electrons are shared equally we get a pure covalent bond • E.g. H2, O2, Cl2 • If they are shared unequally we get a polar covalent bond • E.g. H2O, HCl, CH4 • A polar covalent bond involves the unequal sharing of electrons

13. Properties of Covalent Compounds • Usually liquids or gases • Have low melting and boiling points • Do not conduct electricity (no ions) • Insoluble in water (but soluble in organic solvents e.g. benzene)

14. Writing formulae for ionic compounds • Example: Write the chemical formula for potassium chloride

15. Chlorine • Group 7 • 7 electrons in outer shell • Valency of 1 • Potassium: • Group 1 • 1 electron in outer shell • Valency of 1 K Cl KCl

16. Example 2: Write the chemical formula for magnesium bromide

17. Bromine • Group 7 • 7 electrons in outer shell • Valency of 1 • Magnesium: • Group 2 • 2 electrons in outer shell • Valency of 2 Br Br Mg MgBr2

18. Group Ions

19. Metallic Bond • The metallic bond is the attraction between the positive ions and the cloud of electrons • In a metal, atoms readily lose electrons to form positive ions • This leaves electrons which are free to roam throughout the positive ions

21. This is why… • Metals are good conductors of electricity because electrons are free to move through the metal • They are malleable and ductile because layers of ions can slide over one another • They have high melting and boiling points because of strong bonds

22. Dative bonds • This is like a covalent bond, but the pair of electrons only come from one atom • i.e. Rather than each atom contributing one electron, both electrons come from the same atom

23. Electronegativity • The electronegativity of an element is a measure of the attraction of an atom of that element for a shared pair of electrons • Electronegativity is the attraction for electrons HL OL

24. Trends in electronegativity • Values decrease as you go down a group because: • the atomic radius increases so the bonding pair of electrons are further from the nucleus and are less strongly attracted • Values increase as you go across a period because • the atomic radius decreases due to the increasing nuclear charge (the electrons are attracted more strongly to the nucleus)

25. Sharing of electrons between two atoms • Same electronegativity value: pure covalent bond • E.g. Cl2 • Difference in values is less than 1.7: polar covalent bond • E.g. HCl H Cl • Difference in values is greater than 1.7: ionic bond • E.g. NaCl Na+ Cl- Δ+ Δ-

26. The Hydrogen Bond • If a H atom is bonded with a very electronegative element, it only gets a small share of the electrons. • A polar covalent bond is formed, where H has a slightly positive charge and the other atom has a slightly negative charge

27. Eg. The water molecule Δ- Δ- O Δ+ Δ+ H H

28. The hydrogen in one molecule is attracted to the oxygen in an adjoining molecule • This is called the hydrogen bond • They arise in water (H2O), ammonia (NH3) and hydrogen chloride (HCl) • Hydrogen bonds hold the molecules together in a crystal of ice

29. Hydrogen bonds explain why • Water has a high freezing and boiling point • Water has ‘surface tension’ • Water is a good solvent for ionic compounds because they are polar (Polar means they have a positive part and a negative part)

30. Water doesn’t dissolve covalent substances because they are non-polar • NB: Because water is polar, it will dissolve polar compounds • “LIKE DISSOLVES LIKE”

31. Van Der Waal’s Forces • Weak forces of attraction that exist between molecules, because of the movement of their electrons

33. Dipole Moment • A molecule has dipole moment if its centre of positive charge does not coincide with its centre of negative charge (Dipole moment is really another way of saying if a molecule is polar or not)

34. Examples • H2 (hydrogen molecule) H H -The bond is non-polar -H2 has no dipole moment

35. HCl (hydrogen chloride) Δ+ Δ- Cl H -The bond is polar -NaCl has dipole moment

36. CO2 (carbon dioxide) Δ+ Δ- Δ+ Δ- O O C -Each bond is polar -But CO2 has no dipole moment (because the centre of positive charge coincides with the centre of negative charge)

37. H2O (water) Δ- Δ- O Δ+ Δ+ H H -Each bond is polar -H2O has dipole moment (because of its v – shape), the centre of positive charge does not coincide with the centre of negative charge)

38. Crystals • A regular, solid, repeating structure of atoms, ions or molecules arranged in a pattern called a lattice

39. 1. Ionic Crystals • Positive and negative ions at each point of the lattice structure held together by ionic bonds • E.g. Sodium chloride, (NaCl) • Potassium iodide, (KI)

40. NaCl (sodium chloride) Cl- Na+ Cl- Na+ Na+ Cl- Cl- Na+

41. It is a 3 dimensional cube • Every Na+ ion is surrounded by 6 Cl- ions • Every Cl- ion is surrounded by 6 Na+ ions

42. 2. Covalent (atomic) Crystals • Atoms (not ions) at each point of the lattice structure held together by covalent bonds • E.g. Carbon (in the form of diamond or graphite)

43. DIAMOND • Tetrahedral structure is repeated over and over

44. Diamond is very hard and unreactive because there are 4 strong, covalent bonds • Diamond does not conduct electricity because there are no electrons free to move

45. GRAPHITE This pattern is repeated over and over

47. Graphite is soft because it has 3 covalent bonds and 1 weak van der Waals bond • Graphite can conduct electricity because it has electrons which are free to move • The layers in graphite are able to slide over each other • Graphite is used as a lubricant

48. 3. Molecular Crystals • Molecules at each lattice point held together by Van Der Waal’s forces • Iodine, I2 • Ice

49. 4. Metallic Crystals • Positive metal ions held together by the metallic bond between the positive ions and the cloud of electrons

50. Allotropes are different forms of the same element. • E.g. diamond and graphite are allotropes of carbon