Unit 1 - Chemical Changes and Structure

1. Unit 1 - Chemical Changes and Structure

2. Reaction Rates

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 averagerate over the time interval chosen. • Reactions can be followed by measuring changes in concentration, mass and volume.

4. Nat 5 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. Higher

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; • Decreasing particle size (increasing surface area) • Increasing concentration (of reactant) • Increasing temperature • 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. • 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) • The greater the number of collisions, the faster the reaction. • Therefore the higher the temperature, the faster the reaction rate.

9. Potential Energy Diagrams See jotter for labelled diagrams Labels include; • Exothermic or Endothermic • Activated complex • Enthalpy change • Reaction pathway • Potential Energy (KJ) • Activation Energy

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 siteson their surface. Reactant molecules form weak bonds with the surface in a process called adsorption. At the same time bonds within 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. 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 notpossible to regenerate a poisonedcatalyst and it must be replaced/renewed. This adds to industry costs so every effort is made to remove any impurities from reactants that might poison a catalyst.

14. 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

15. 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.

16. Trends in the Periodic Table

17. Density (measured in g/cm3) 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.)

18. Good to know but not essential... 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. 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 H-Cl is 37 + 99 pm = 136 pm There are two clear trends in the periodic table; • Going across a period covalent radii decreases. This is because… • Going down a group covalent radii increases. This is because…

19. 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.

20. 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; • Calculation (whiteboard examples)

21. The first compounds of the noble gases were formed using Xenon (Xe). Tetrafluoroammonium octafluoroxenate (NF4)2XeF8) is a strong oxidiser and can be used in propelling rockets. Suggest why xenon compounds were the first to be made. (Hint it’s related to atomic size…) Xenon has the largest atoms of any noble gas. This means that the electron removed in ionisation to create Xe+ is the furthest from the nucleus. Hence, the first ionisation energy of Xenon is the lowest of any of the noble gases.

22. 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)

23. refer to bonding 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)

25. O 2 Bonding, Structure and Properties of Elements He S Fe C

26. Summary of Bonding types in first 20 elements.

27. 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.

28. Types of Bonding in elements There are 3 types; Metallic Bonding – (intramolecular) Covalent Bonding – (intramolecular) Van der Waal’s a) London Forces– (intermolecular) Splits into 3 sub categories

29. Metallic Bonding • Metallic bonding 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)

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

31. Covalent Molecular Bonding whiteboard example • Covalent bonding (usually) 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.

32. Van der Waals’ Forces • Van der Waals forces are weak bonding which occurs BETWEEN molecules. • London forces are one of three types of Van der Waals’ forces. • London forcesare 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) • Van der Waals forces are useful when explaining patterns in the periodic table e.g. melting/boiling points.

33. Bonding in Specific Groups Groups 1, 2 and 3 All elements in groups 1, 2 and 3 have strong metallic bonds holding them together in a giant lattice structure. Metallic bonds allows metals to be shaped (i.e. malleable and ductile) Metals have high melting/boiling pts due to strong metallic bonds Boron isthe only exceptionas it has very complex bonding. B12 is almost as hard as diamond. This suggests a covalent networkstructure.

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

35. Graphite (covalent network) Each atom forms 3 covalent bonds and its last valence electron becomes delocalised. As the delocalised electronsare 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.

36. The fullerenes, despite being large molecules, are discrete covalent molecules. • The smallest of fullerenes is a molecule known as Buckminsterfullerene (C60). • 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. 'Bucky Ball' (covalent molecule) 3. Buckminsterfullerene (aka ‘Bucky Ball’)

37. Nitrogen atoms form diatomic molecules with a triple covalent bond. • This means that nitrogen only has London forcesbetween 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 P4 molecules which are larger than N2 molecules. As a result it has stronger London forces between its molecules. This stronger attraction means phosphorus has a higher boiling ptand is asolidat room temp. Group 5 (nitrogen and phosphorus)

38. 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. Group 6 (oxygen and sulphur)

39. 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 between 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.

41. 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 brokenwhen 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.

42. Summary of Bonding types in first 20 elements. Diatomic molecules Other small molecules Metallic bonding Monatomic elements Covalent network Discrete covalent molecules Except Buckminsterfullerenes ! H He Li Be B C N O F Ne Mg Al Si P S Cl Ar Na Ca K

43. Bonding, Structure and Properties of Compounds

44. Types of Bonding in compounds There are 4 main types; Ionic Bonding – intramolecular Covalent Network Bonding – intramolecular Polar/Non Polar Covalent Bonding - intramolecular Intermolecular (Van der Waals) Hydrogen bonding (present in H2O, NH3 and HF) Permanent dipole – Permanent dipole interactions London forces – Temporary dipole interactions Strongest to Weakest

45. 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 previous 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.

46. Structure diagram • 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 zeroand 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 (CaF2) there are twice as many F- ions than Ca2+

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

48. Covalent Network

49. These covalent network compounds have the same properties as covalent network elements. • Both SiC and SiO2 have very high melting pts. as melting requires breaking strong covalent bonds. • Silicon carbide – known as ‘carborundum’ - is structurally similar to diamond and has many uses due to it’s strength, durability and low cost.

50. Polar Covalent Bonding 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 theatom (element) with the higher e.n. holds the electron slightly closerto itself and therefore becomes slightly negative (∂-) The atom with the lesser e.n. is therefore slightly positive (∂+) as the electron is sitting further awayfrom it. Covalent bonds with unequal electron sharing are called polar covalent bonds. Molecules are attracted via permanent to permanent dipole interactions