3. Reaction Rates During the course of a chemical reaction, reactants are being converted into products. Measurement of the rate of reaction involves measuring the ‘change in the amount’ of a reactant or product in a certain time. The rate of reaction changes as it progresses, being relatively fast at the start and slowing towards the end. What is being measured is the average rate over the time interval chosen. Reactions can be followed by measuring changes in concentration, mass and volume.

4. Where property = mass/volume/concentration The above is used when there is no change in mass/volume/concentration measured, for example during a colour change reaction.

5. Collision Theory A chemical reaction can only occur if there is a successful collision between reactant molecules. From national 5 we know that we can speed up a chemical reaction by; 1.Decreasing particle size (increasing surface area) 2.Increasing concentration (of reactant) 3.Increasing temperature 4.Adding a catalyst

6. Collision Theory – Particle Size The smaller the particle size, the higher the surface area. The higher the surface area, the greater the number of collisions that can occur at any one time. The greater the number of collisions, the faster the reaction. Therefore the smaller the particle size, the faster the reaction rate.

7. Collision Theory – Concentration The higher the concentration, the higher the number of particles. The higher the number of particles, the greater the chance of collisions that can occur. The greater the number of collisions, the faster the reaction. Therefore the higher the concentration, the faster the reaction rate.

8. Collision Theory – Temperature The higher the temperature, the higher the energy the particles have. The higher the energy, the faster the particles move. The faster the particles move, the greater the chance that they can collide with sufficient energy (activation energy) to be successful. The greater the number of collisions, the faster the reaction. Therefore the higher the temperature, the faster the reaction rate.

9. Collision Theory – Catalyst A catalyst speeds up a chemical reaction by lowering the activation energy. (i.e. catalyst provides another ‘easier’ route) The lower the activation energy, the greater the chance of successful collisions. The more collisions in a period of time, the faster the reaction rate.

10. Catalysts A catalyst is a substance which speeds up a chemical reaction without getting used up or changed itself. There are two main categories of catalyst; a) Heterogeneous + b) Homogenous.

11. Heterogeneous Catalysts Heterogeneous catalysts have active sites on their surface. Reactant molecules form weak bonds with the surface in a process called adsorption. At the same time bonds with the adsorbed reactant molecules are weakened. The reactant molecules are also held at a favourable angle for a collision with another reactant molecule to occur. The product molecules then leave the active site in a stage called desorption. The active site is then available again.

12. surface Active site Desorption – product molecules formed.

13. Unfortunately unwanted substances can often be adsorbed onto the active sites thus making them unavailable for the normal reactants. (Example; lead in petrol.) When this happens the catalyst is said to be poisoned. Sometimes it is not possible to regenerate a poisoned catalyst and it must be replaced/renewed. This adds to industry’s costs so every effort is made to remove any impurities from reactants that might poison a catalyst.

14. You’ll need 3 colours when drawing this diagram (underneath - if possible – the previous catalyst note)

15. Homogeneous Catalysts A catalyst that is in the same state as the reactants is said to be a homogeneous catalyst. The catalyst forms an intermediate compound with one of the reactants. (this intermediate compound later decomposes to reform the catalyst.) For example (using Reactants A and B) Reactant A + Catalyst Intermediate Intermediate + Reactant B Product + Catalyst

16. Potential Energy Diagrams Labels include; Exothermic or Endothermic Activated complex Enthalpy change Reaction pathway Potential Energy (KJ) Activation Energy

18. Summary of Bonding types in first 20 elements.

19. Inter – means in between. In other words an INTERmolecular bonds means bonds in between the molecules. Intra – means within. In other words an INTRAmolecular bond means bonds within the molecule.

20. Types of Bonding in elements There are 3 types; 1.Metallic Bonding – (intramolecular) 2.Covalent Bonding – (intramolecular) 3.Van der Waal’s/London Forces – (intermolecular)

21. Metallic Bonding Metallic bonding – unsurprisingly – only appears in metal elements. Metallic bonding occurs between (positively charged) metal ions and delocalised outer shell electrons. ‘delocalised’ means the electrons are common to all of the ions (i.e. they move from one to another)

22. The movement of delocalised electrons allow metal elements to conduct electricity

23. Covalent Bonding Covalent bonding occurs between two non metal atoms. Covalent bonds are held together through the attraction between the positively charged nucleus of one atom and the negatively charged outer electrons of the other atom. Outer electrons are shared in covalent bonding.

24. Van der Waals’ Bonding Van der Waals bonding is weak bonding which occurs BETWEEN molecules. London forces are temporary dipole to temporary dipole attractions. Temporary dipoles occur when electrons lie slightly closer to one atom than the other. This means for a short time one of the atoms is slightly negative and the other is slightly positive (i.e. electrons not shared equally) London forces (and other intermolecular forces) are useful when explaining patterns in the periodic table e.g. melting/boiling points.

25. Bonding in Specific Groups Groups 1, 2 and 3 All elements in groups 1, 2 and 3 have strong metallic bonds holding them together. Metallic bonds allows metals to be shaped (i.e. malleable and ductile) Metals have high melting/boiling pts due to strong metallic bonds Boron is the only exception as it has very complex bonding. B 12 is almost as hard as diamond. This suggests a covalent network structure.

26. The 3 Structures of Carbon (group 4) Each atom covalently bonds to 4 other atoms. This means covalent bonds must be broken to melt/ boil = very high m.pt/b.pt values. No free electrons = no conduction. Tunnels between atoms allow light through = transparent structure.

27. Each atom forms 3 covalent bonds and its last valence electron becomes delocalised. As the delocalised electrons are only held weakly they can flow i.e. graphite conducts electricity. The delocalised orbitals sit between the layers - as a result there are 3 strong covalent bonds WITHIN the layers but only weak interaction BETWEEN the layers. Due to these weak interactions, graphite is flaky as the layers can be easily separated. Graphite layers are offset (i.e. not above each other) - light can’t travel through it meaning it is not transparent.

28. 3. Buckminsterfullerene (aka ‘Bucky Ball’) The fullerenes, despite being large molecules, are discrete covalent molecules. The smallest of fullerenes is a molecule known as Buckminsterfullerene (C 60 ). This is a spherical molecule containing 5 and 6 membered carbon rings. The properties are still being researched so the full applications are still unknown.

29. Group 5 (nitrogen and phosphorus) Nitrogen atoms form diatomic molecules with a triple covalent bond. This means that nitrogen only has London forces between the molecules. London forces are easily broken and as a result nitrogen has a low boiling pt. This is why nitrogen is a gas at room temperature. Phosphorus forms tetrahedral P 4 molecules which are larger than N 2 molecules and as a result it has stronger London forces between its molecules. This stronger attraction means phosphorus has a higher boiling pt and is a solid at room temp.

30. Group 6 (oxygen and sulphur) Oxygen atoms form diatomic molecules with a double covalent bond. This means that oxygen only has London forces interaction between the molecules. London forces are easily broken and as a result oxygen has a low boiling pt. Hence oxygen is gas @ room temp. Sulphur forms 8 membered rings. The London forces between the molecules are strong enough in sulphur to make it solid at room temperature.

31. Group 7 (the halogens) The halogens form diatomic molecules (i.e. they bond with themselves) As with oxygen and nitrogen this results in the halogens only having London forces with each other molecule. Fluorine and chlorine are very volatile (and therefore reactive) gases due to these weak intermolecular forces. Group 8 (the noble gases) As the noble gases have a stable outer electron shell they do not form bonds. As a result they remain monoatomic.

33. Explaining the Melting and Boiling pts Trend In small discrete covalent molecules the melting and boiling points are low. This is because only weak intermolecular London forces have to be overcome when boiling or melting. The strong covalent bonds are left unaffected. In the covalent network solids (carbon, silicon and boron) strong covalent bonds MUST be broken when melting or boiling. Breaking these bonds requires a lot more energy and therefore we get very high values. In the metal groups (1, 2, 3) strong metallic bonds MUST be overcome thus they have high melting/boiling points.

34. Summary of Bonding types in first 20 elements. Diatomic molecules Other small molecules Discrete covalent molecules Monatomic elements Covalent network Metallic bonding Except Buckminsterfullerenes !

36. Types of Bonding in compounds There are 4 main types; 1.Ionic Bonding – intramolecular 2.Covalent Network Bonding – intramolecular 3.Polar/Non Polar Covalent Bonding - intramolecular 4.Intermolecular (Van der Waals) 1.Hydrogen bonding (present in H 2 O, NH 3 and HF) 2.Permanent dipole – Permanent dipole interactions 3.London forces – Temporary dipole interactions Strongest to Weakest

37. Ionic Bonding Ionic bonding is an electrostatic attraction between the positive ions and negative ions. Ionic bonding is related to the electronegativities of elements. The greater the difference in ‘e.n’ the less likely the elements are to share outer electrons. (electronegativity definition and trends are found in later section of jotter.) Instead the element with the higher ‘e.n’ value will gain the electrons to form a negative ion and the element with lower ‘e.n’ value will lose the electrons to form a positive ion. Due to the trends of electronegativity, the elements that are far apart from one another in the periodic table form ionic bonds. (normally metal and non metal.) Caesium fluoride is the compound with the greatest ionic character.

38. Structure Ionic compounds do not form molecules. Instead the positive and negative ions come together to form lattice structures. When the lattice forms, energy is released. This is known as lattice energy or enthalpy. The overall charge of the lattice must be zero and therefore this affects the number of each ions we have present. In sodium chloride (NaCl) there is an equal number of Na + and Cl - ions. In calcium fluoride (CaF 2 ) there are twice as many F - ions than Ca 2+

39. Covalent Compounds There are 3 types of covalent bonding in compounds (all involving combinations of non metals) ; 1. Covalent network structures. 2. Polar covalent molecules 3. Non Polar covalent molecules

40. Covalent Network

41. These covalent network compounds have the same properties as covalent network elements. Both SiC and SiO 2 have very high melting pts. as melting requires breaking strong covalent bonds. Silicon carbide (structurally similar to diamond) has many uses due to it’s strength, durability and low cost. Silicon carbide is often referred to as ‘carborundum’

42. Most covalent compounds are made from atoms with slightly different electronegativity (e.n.) This difference is not significant enough for one of the atoms to fully remove an electron from the other. (Approx difference of between 0.5 and 1.6) As a result the atom (element) with the higher e.n. holds the electron slightly closer to itself and therefore becomes slightly negative ( ∂- ) The atom with the lesser e.n. is therefore slightly positive ( ∂+ ) as the electron is sitting further away from it. Covalent bonds with unequal electron sharing are called polar covalent bonds. Polar Covalent Bonding

43. Non Polar Covalent Bonding Non polar (or pure) covalent bonding normally occurs when; 1.electrons are equally shared between the two different atoms. i.e. equal electronegativity. E.g. Phosphorus Hydride 2.the compound structure is symmetrical and therefore charges are overall balanced. E.g. CH 4 (methane) and CO 2 (carbon dioxide)

44. Both polar and non polar? It’s possible for non polar covalent molecules to have individual polar bonds. For example; Carbon Tetrachloride (CCl 4 )

45. ∂ ∂

46. Summary of electronegativity values + bonding In general; 1.If the electronegativity difference (usually called ΔEN) is less than 0.5, then the bond is non polar (pure) covalent. 2.If the ΔEN is between 0.5 and 1.6, the bond is considered polar covalent 3.If the ΔEN is greater than 2.0, then the bond is ionic.

47. Boiling Points Polar covalent molecules have higher boiling points than non polar covalent molecules with a similar mass. This is because the intermolecular forces are stronger (changing from London forces to permanent – permanent dipole interactions.) Permanent dipole to dipole interaction is caused via the constant attraction between the ∂+ atoms and ∂- atoms of neighbouring molecules. Properties of Polar/ Non Polar Covalent Bonds

48. Solvent Action ‘like substances dissolve in like substances’ This means that polar molecules will dissolve in polar solutions but not in non polar solutions and vice versa. This is due to the attraction between ∂+ and ∂- atoms of the water and the polar substance. Ionic compounds dissolve in polar solutions in a similar way due to the interaction between the ions and the ∂+ and ∂- atoms. Ions surrounded by a layer of water molecules – held by electrostatic attraction – are said to be hydrated.

49. Behaviour in electric field Copy figure 4.8 on page 49 Viscosity (thickness/ability to pour) Summarise textbook notes on page 54 (2/3 sentences max) Miscibility (ability to mix) Summarise textbook notes on page 57 (2/3 sentences max)

50. Physical properties of some hydrides Boiling Points

51. The above bonds are very polar due to the large difference in the electronegativity values. This interaction is called hydrogen bonding. Hydrogen bonding occurs in any molecule that contains any of the above bonds but mainly in hydrogen fluoride (HF), water (H 2 O) and ammonia (NH 3 ). These intermolecular forces affect the properties of these compounds. Hydrogen bonding is stronger than both London forces and permanent dipole to dipole interactions but weaker than covalent bonding.

52. Water is unusual as it’s solid form (ice) is less dense than it’s liquid form (water). This means that ice floats on water whereas most other solids sink in their own liquid forms. This phenomenon is due to the structure and bonding which takes place between the water molecules. As water molecules cool, they contract. However, at 4 o C they begin to expand. This is because of hydrogen bonds between the water molecules. This decreases the density of ice (greater volume/same mass) compared to that of the liquid water. Ice floating is vital to ‘real life’ – i.e. fish/marine life surviving under frozen lakes etc (https://www.youtube.com/watch?v=T4GCShGvw-M) Density of Water

53. Trends in the Periodic Table

54. Density (measured in g/cm 3 ) Across a period (starting from group 1) the density increases towards the centre (group 4) and then decreases. Density tends to increase down a group (as atomic number increases.)

55. Atomic size Atomic size (or covalent atomic radius) is half the distance between the nuclei of two bonded atoms. Single bond lengths between atoms of different elements can be found by adding their individual covalent radii. e.g. the covalent radii of hydrogen and chlorine are 37 and 99 pm so the bond length in HCl is 37 + 99 pm = 136 pm There are two clear trends in the periodic table; 1.Going across a period covalent radii decreases. This is because… 2.Going down a group covalent radii increases. This is because… the nuclear charge increases but the number of electron shells stays the same – i.e. the outer electrons are held more tightly, making the atom smaller. the number of electron shells increases and therefore the inner (full) electron shells shield the outer electrons from the nuclear charge. (known as ‘shielding effect’) i.e. the outer electron are held less tightly, making the atom bigger.

56. First Ionisation Energy Definition; First ionisation energy is the energy required to remove one mole of electrons from an element in gaseous state. (example at top of page 11 in data book – no excuses!) Trends; Across a period… increases due to an increased nuclear charge holding the outer electrons more closely (smaller atoms.) This means you need MORE energy to remove a mole of electrons. Down a group… decreases due to the shielding effect (bigger atoms.) This means the outer electrons are further away from the nucleus and therefore this attraction is less, thus it is easier (less energy) to remove one mole of electrons.

57. 2.The first ionisation energy of lithium = 520kJmol -1 but its second ionisation energy value = 7298kJmol -1. Why is there such a big difference between the two values? Lithium achieves a stable outer electron shell (octet) when it loses one electron. Therefore first ionisation energy is small. The second electron would therefore be removed from an stable octet which is unfavourable and requires a lot of energy. Other types of ionisation questions; 1. Calculation (whiteboard examples)

58. Electronegativity Electronegativity is a measure of the tendency of an atom to attract electrons. (think pulling power) Electronegativity is measured on the Pauling scale. Trends; Electronegativity increases across a period. Electronegativity decreases down a group. (Hint – Fluorine is the highest value)

59. Melting and Boiling Points Generally, the stronger the bond between atoms, the higher the energy required to break that bond. Melting/boiling points are varied and don't generally form a trend across a period however; Metals generally possess a high melting point. Most non-metals possess low melting points. Group 4 have the highest values. Metal group example; In group 1 (alkali metals), the melting/boiling pts decrease as the atomic number increases. This is because there is an decrease in the attraction between the particles. (refer to bonding) Non metal group example; In group 7 (halogens) the melting/boiling pts increase as the atomic number increases. This is because there is an increase in the attraction between the particles. (refer to bonding)