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Bonding. Mr Field. Using this slide show. The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to! Slide shows should be shared with students (preferable electronic to save paper) and they should add their own notes as they go along.

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Bonding


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    1. Bonding Mr Field

    2. Using this slide show • The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to! • Slide shows should be shared with students (preferable electronic to save paper) and they should add their own notes as they go along. • A good tip for students to improve understanding of the calculations is to get them to highlight numbers in the question and through the maths in different colours so they can see where numbers are coming from and going to. • The slide show is designed for my teaching style, and contains only the bare minimum of explanation, which I will elaborate on as I present it. Please adapt it to your teaching style, and add any notes that you feel necessary.

    3. Menu: • Lesson 1 – Ionic Bonding • Lesson 2 – Covalent Bonding • Lesson 3 – Structures • Lesson 4 – Physical Properties • Lesson 5 – Molecular Shapes • Lesson 6 – Intermolecular Properties • Lesson 7-9 – Internal Assessment • Lesson 10 – HL – Sigma and pi bonds • Lesson 11 – HL – Hybridisation • Lesson 12 – HL – Delocalisation

    4. Lesson 1 Ionic Bonding

    5. Overview • Copy this onto an A4 page. You should add to it as a regular review throughout the unit.

    6. Assessment • This unit will be assessed by: • An internal assessment (24%) at the end of the unit • A joint test along with Periodicity at the end of that unit

    7. We Are Here

    8. Lesson 1: Ionic Bonding • Objectives: • Reflect on prior knowledge of bonding • Refresh knowledge and understanding of ionic bonding

    9. Reflecting on bonding • Brainstorm everything you already know about bonding. • You have one minute

    10. 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 salts Na+ Li+ Li+ O2- F-

    11. Ionic Structures • Don’t worry about this now, you will be looking at in more in Lesson 3

    12. 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  Ca2+ + 2e- • Group 3: Al  Al3+ + 3e- • Transition metals – form multiple different ions • Fe  Fe2+ + 2e- • Fe  Fe3+ + 3e- • Negative ions (anions) • Negative ions are formed when non-metals gain enough electrons to complete their outer shells • Group 5: N + 3e- N3- • Group 6: O + 2e-  O2- • Group 7: F + e-  F-

    13. 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, SO42- • Phosphate, PO43- • Nitrate, NO3- • Carbonate, CO32- • Hydrogen carbonate, HCO3- • Ethanoate (acetate), CH3CO2- • Hydroxide, OH- • Ammonium, NH4+

    14. The formula of ionic compounds • Ionic compounds are always neutral, so the charges must balance • Example 1: calcium reacting with fluorine: • Calcium forms Ca2+, fluorine forms F- • The formula is CaF2 so two F- charges cancel the one Ca2+ • Example 2: iron (II) reacting with phosphate • Iron (II) is the Fe2+ ion, phosphate is PO43- • The formula is Fe3(PO4)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

    15. The names of ionic compounds • The cation gives the first part of the name • Normally a metal except in the case of ammonium (NH4+) • 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: • CaF2: calcium fluoride • Fe3(PO4)2: iron (II) phosphate

    16. Your turn • Deduce the formulae and names of the ionic compounds formed between: • Lithium and fluorine • Magnesium and iodine • Aluminium and oxygen • Iron (II) and sulphur • Calcium and nitrogen • Sodium and sulphate ions • Chloride and ammonium ions • Iron (III) and sulphate ions • Iron (II) and nitrate ions • Potassium and carbonate ions • Work through the simulation here: http://www.learner.org/interactives/periodic/groups_interactive.html

    17. 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

    18. Homework • Research and make notes on metallic bonding. Including: • Description of the nature of the metallic bond • Factors affecting the strength of metallic bonds • Explanation of the malleability of metals • Explanation of the electrical conductivity of metals • Factors affecting the conductivity of metals

    19. Lesson 2 Covalent Bonding

    20. Refresh • Predict and explain which of the following compounds are ionic: • NaCl • BF3 • CaCl2 • N2O • P4O6 • FeS • CBr4.

    21. We Are Here

    22. Lesson 2: Covalent Bonding • Objectives: • Refresh knowledge and understanding of covalent bonding • Learn how to draw Lewis structures • Identify examples of dative bonding • Identify instances of expanded octets

    23. 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 salts Na+ Li+ Li+ O2- F-

    24. Covalent bonding • A covalent bond is the attraction of two atoms to a shared pair of electrons watercarbon dioxide Each O has two single bonds each 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 O C O H H O

    25. 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

    26. Covalent Structures • Molecular • As in water and methane • Giant lattice • As in silicon dioxide • More on these later in the unit

    27. Lewis structures • Show the position of outer-shell electrons in a covalent compound • Various types: all show the same thing, any is fine dots and crosses crosses only dots only lines 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!

    28. Working out a Lewis structure Don’t worry about the shape…more on that later! • Example: diazene, N2H2

    29. Time to practice…again • Draw Lewis structures for the following, bearing in mind the previous two slides • H2 • O2 • N2 • H2O • HCl • NH3 • CO2 • HCN • C2H4 • C2H2

    30. 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

    31. The expanded octet • In this example, the Lewis structure of SO3 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

    32. Time to practice…again • Draw Lewis structures for the following, bearing in mind the previous two slides • NH4+ • SO2 • B2F6 • Al2Cl6 (yes covalent!) • SF6 • PCl5 • CO • XeF4

    33. 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.

    34. 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

    35. Lesson 3 Structures

    36. Refresh • How many lone pairs and bonding pairs of electrons surround xenon in the XeF4 molecule? Lone pairs Bonding pairs • 4 8 • 0 8 • 0 4 • 2 4

    37. We Are Here

    38. Lesson 3: Structures • Objectives: • Describe and compare the structures and properties of: • Allotropes of carbon • Silicon and silicon dioxide • Ionic compounds

    39. Marketplace – in four groups • Each group needs to produce a learning resource to teach the other students about their chosen topic. • Once the resources are completed, one person should remain with the resource whilst the remaining members circulate and learn from the other stations….you should manage your time, taking turns manning your station to make sure everyone makes it round class. • There will be a test at the end. • Groups should look at the structure, bonding, properties and uses of: • Group 1: Diamond and graphite • Group 2: Buckminster Fullerenes and carbon nanotubes • Group 3: Silicon and silicon dioxide • Group 4: Ionic compounds (less focus on uses)

    40. Test Time • Complete the test here • You have 10 minutes

    41. Alternative to Marketplace Activity • Produce a table summarising the differences in properties of three allotropes of carbon: diamond (this should include silicon and silicon dioxide which have similar structures), graphite and buckminster fullerene. Look at structure, properties (explained), uses (and how these are related to the properties) • Draw a mind-map summarising the three main types of structure: giant ionic, giant covalent, molecular including a drawing of each, three example compounds, and an explanation of their properties. • Produce a Venn Diagram to summarise then similarities and differences between the three main types of bonding • Produce a flow chart that can be followed to determine the type of bonding present in an element/compound and (for ionic/covalent) produce Lewis diagrams to describe it.

    42. Homework • Graphene is a recently discovered ‘super-material’ • Research its structure and properties and as many potential uses for it as possible • Make sure you have

    43. Key Points • Carbon (diamond, graphite and fullerenes), silicon and silicon dioxide exhibit giant covalent (macromolecular) structures • For example diamond: each carbon is bonded to exactly four others and so on • Ionic compounds form giant ionic lattices • NaCl: every Na+ ion surrounded by 6 Cl- ions, every Cl- ion surrounded by 6 Na+ ions

    44. Lesson 4 Physical Properties

    45. Refresh • Unlike many covalently bonded substances, graphite is an excellent conductor. Describe the structure and bonding in graphite and explain why it is such a good conductor.

    46. We Are Here

    47. Lesson 4: Testing Properties • Objectives: • Design and conduct an experiment to use the physical properties of compounds to differentiate between them

    48. Design Challenge • You will be provided with samples of 6 unknown substances: • Lead shot, iodine granules, graphite, sodium chloride, sugar, lead bromide • You need to design and conduct a series of experiments to let you determine the nature of the bonding in each, using their physical properties. Focus on: • Solubility in polar solvents (such as water) • Solubility in non-polar solvents (such as hexane) • Melting/boiling point • Electrical conductivity (when solid, molten or in solution) • Volatility (evaporatingness…not a word but you know what I mean!)

    49. Homework • Bring laptop to next lesson and make sure you have installed ACD/Labs ChemSketch (Google it!) • Identify which substance is which out of: • Lead • Graphite • Iodine • Sugar • Sodium chloride • Lead bromide • Fully explain your reasoning

    50. Key Points • The structure and bonding of substances has a very significant effect on their properties