1. BONDING

2. OVERVIEW

3. IONIC BONDING

4. LESSON 1: IONIC BONDING Objectives: Reflect on prior knowledge of bonding Refresh knowledge and understanding of ionic bonding

5. REFLECTING ON BONDING Brainstorm everything you already know about bonding. You have 2 minutes

6. RECAPPING IONIC BONDING An ionic bond is: The electrostatic attraction between two oppositely charged ions sodium fluoride lithium oxide Ionic bonds typically form between a metal and a non-metal Ionically bonded compounds are often referred to as salt Li + O 2- Li + Na + F-F-

7. FORMATION OF SIMPLE IONS Positive ions (cations) Positive ions are formed when metals lose their outer shell electrons Group 1: Li  Li + + e - Group 2: Ca  Ca 2+ + 2e - Group 3: Al  Al 3+ + 3e - Transition metals – form multiple different ions Fe  Fe 2+ + 2e - Fe  Fe 3+ + 3e - Negative ions (anions) Negative ions are formed when non-metals gain enough electrons to complete their outer shells Group 5: N + 3e -  N 3- Group 6: O + 2e -  O 2- Group 7: F + e-  F -

8. POLYATOMIC IONS Many ions are made of multiple atoms with an overall negative charge The negative ones are mostly acids that have lost their hydrogens You need to know about:  Sulphate, SO 4 2-  Phosphate, PO 4 3-  Nitrate, NO 3 -  Carbonate, CO 3 2-  Hydrogen carbonate, HCO 3 -  Ethanoate (acetate), CH 3 CO 2 -  Hydroxide, OH -  Ammonium, NH 4 +

9. THE FORMULA OF IONIC COMPOUNDS Ionic compounds are always neutral, so the charges must balance Example 1: calcium reacting with fluorine: Calcium forms Ca 2+, fluorine forms F - The formula is CaF 2 so two F - charges cancel the one Ca 2+ Example 2: iron (II) reacting with phosphate Iron (II) is the Fe 2+ ion, phosphate is PO 4 3- The formula is Fe 3 (PO 4 ) 2 The 6 + charges from iron (2 + x 3) balance the 6 - charges (3 - x 2) from phosphate Look for the lowest common multiple Ionic compounds do not form molecules so these are always empirical formulae

10. THE NAMES OF IONIC COMPOUNDS The cation gives the first part of the name Normally a metal except in the case of ammonium (NH 4 + ) In the case of transition metals, Roman numerals tell you the charge on the metal ion The anion gives the second part of the name Simple ions: ‘-ide’…e.g. chloride, fluoride, nitride etc Complex ions: just their name: sulphate, phosphate etc Note: the ‘-ate’ ending usually refers to polyatomic ions containing oxygen, which provides the negativity…more on this in the redox unit Examples: CaF 2 : calcium fluoride Fe 3 (PO 4 ) 2 : iron (II) phosphate

11. Deduce the formulae and names of the ionic compounds formed between: 1. Lithium and fluorine 2. Magnesium and iodine 3. Aluminium and oxygen 4. Iron (II) and sulphur 5. Calcium and nitrogen 6. Sodium and sulphate ions 7. Chloride and ammonium ions 8. Iron (III) and sulphate ions 9. Iron (II) and nitrate ions 10. Potassium and carbonate ions 11. Work through the simulation here: http://www.learner.org/interactives/periodic/groups_interactive.html http://www.learner.org/interactives/periodic/groups_interactive.html

12. KEY POINTS Ionic bonds are the attraction between two oppositely charged ions Ionic bonds form between metals and non metals Metals lose their outer shell Non-metals complete their outer shell The number of each ion in the formula is determined by the lowest common multiple of their charges

13. COVALENT BONDING

14. A covalent bond is the attraction of two atoms to a shared pair of electrons watercarbon dioxide Each O has two single bondseach C has two double bonds Atoms aim for complete outer-shells, and each covalent bonds gives them one electron Atoms form as many bonds as they have gaps in their outer-shells Covalent bonds typically form between two non-metals H H O O O C

15. HOW MANY BONDS? Atoms (usually) form bonds according to the ‘octet’ rule This means they try to get a full outer shell of 8 electrons (except hydrogen which is full at 2) Atoms form as many bonds as they have ‘gaps’ in their outer shells, with each bond gaining them one electron: Group 7: 7 electrons, 1 gap  1 bond Group 6: 6 electrons, 2 gaps  2 bonds Group 5: 5 electrons, 3 gaps  3 bonds Group 0/8: 8 electrons, 0 gaps  0 bonds Covalent bonds can be: Single: one shared electron pair, X-X Double: two shared electron pairs, X=X Triple: three shared electron pairs, X  X

16. COVALENT STRUCTURES Molecular As in water and methane Giant lattice As in silicon dioxide More on these later in the unit

17. LEWIS STRUCTURES Show the position of outer-shell electrons in a covalent compound Various types: all show the same thing, any is fine. Blue Circles: These are the bonding pairs of electrons – the ones involved in the bonds. Red Circles: These are non-bonding or lone pairs of electrons. They are very important, but students often forget about them!

18. WORKING OUT A LEWIS STRUCTURE Example: diazene, N 2 H 2 Step 1: Write the number of electrons in each atom and the number of bonds each atom can form Nitrogen: 5 electrons, 3 bonds Hydrogen: 1 electron, 1 bond Step 2: Draw the structure using lines for bonds There will be 2 N-H bonds and 1 N=N bond Step 3: Add in the lone pairsThe N started with 5 electrons, and 3 are in bonds, so that leaves 2 remaining…each N will have one lone pair Don’t worry about the shape…more on that later!

19. TIME TO PRACTICE…AGAIN Draw Lewis structures for the following, bearing in mind the previous two slides 1. H 2 2. O 2 3. N 2 4. H 2 O 5. HCl 6. NH 3 7. CO 2 8. HCN 9. C 2 H 4 10. C 2 H 2

20. THE DATIVE-BOND Sometimes an atom will contribute both of the electrons in a covalent bond, this is called a dative (covalent) bond E.g. In this example, the lone pair from a water molecule has formed a dative bond to a hydrogen ion (H + ) You can show dative bonds with an arrow to say where the electrons came from…but do not have to

21. THE EXPANDED OCTET In this example, the Lewis structure of SO 3 shows it with 12 electrons in the outer shell This is because sulphur can make use of its empty d-orbitals (the 3d ones) This is called an expanded octet Period 2 elements can’t do this as they have no d- orbitals

22. TIME TO PRACTICE…YET AGAIN Draw Lewis structures for the following, bearing in mind the previous two slides 1. NH 4 + 2. SO 2 3. B 2 F 6 4. Al 2 Cl 6 (yes covalent!) 5. SF 6 6. PCl 5 7. CO 8. XeF 4

23. KEY POINTS Atoms (generally) form covalent bonds according to the octet rule. Each covalent bond gives an atom one extra electron In dative bonds, both the electrons in the bond come from the same atom Period 3 (and above) elements can break the octet rule by using empty d- orbitals and might have 12 or more electrons in their outer shell

24. HOMEWORK Check tables 9 and 10 of the data booklet Draw a graph of bond length vs. bond enthalpy (strength) The easiest way is to enter the data into Excel and get it to draw it Identify and explain the relationship between bond length and bond strength. Identify any significant exceptions to this trend and explain why they occur.