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Using Covalent Structures

Using Covalent Structures. Separating Immiscible Liquids. “Immiscible” means “two liquids that can’t be dissolved”, e.g. oil and water:. Separating these liquids is fairly easy – you simply allow them to settle and then “tap off” the heavier liquid at the bottom using a separating funnel.

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Using Covalent Structures

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  1. Using Covalent Structures

  2. Separating Immiscible Liquids “Immiscible” means “two liquids that can’t be dissolved”, e.g. oil and water: Separating these liquids is fairly easy – you simply allow them to settle and then “tap off” the heavier liquid at the bottom using a separating funnel.

  3. Separating miscible liquids Miscible liquids are liquids that have dissolved together, so separating them is much harder. Here’s an example – distillation: This apparatus can be used to separate water and ethanol because they have different _____ ______. The ______ will evaporate first, turn back into a _______ in the condenser and collect in the _______. The water remains in the round flask, as long as the _______ does not exceed 100OC. Words – temperature, boiling points, ethanol, beaker, liquid

  4. Fractional Distillation of Air Gaseous nitrogen out at -190OC Liquefied air at -200OC Liquid oxygen out at -185OC Air can be distilled in the same way – you simply have to get it cold enough first:

  5. Chromatography R G B X 1 2 3 Z Chromatography can be used to separate a mixture of different inks. Some example questions… 1) Ink X contains two different colours. What are they? 2) Which ink is ink Z made out of?

  6. Rf value This line marks the distance travelled by the solvent Distance travelled by substance Distance travelled by solvent Rf value = R G B The Rf value is a way of measuring how far a substance has moved:

  7. Example questions 10cm 8cm 5cm 2cm R G B Calculate the Rf values of the following:

  8. Using Chromatography Dye 1 Dye 2 Dye 3 Chromatography can be used to test which foods contain which ingredients. For example, consider the dye Sudan 1, which was found in 450 foods in 2005. Which dye contains Sudan 1? Sudan 1

  9. Topic 4 – Groups in the Periodic Table

  10. The Periodic Table Noble gases (Group 0) Alkali metals (Group 1) These elements are called the “transition metals” Halogens (Group 7) This line divides metals from non-metals

  11. A closer look at metals Metals are defined as elements that readily lose electrons to form positive ions. The electrons in the highest shells are delocalised and surround positive ions. There are a number of ways of drawing this: + + + + + + + + + + + + + + + + + + + + + + + + + + + + - - + + + - + + - + + - + + + + + - + - - Delocalised electrons

  12. Properties of metals 03/01/2020 Metals have very high melting points (which means that they are usually _____) whereas non-metals will melt at lower ___________ All metals conduct heat and __________ very well, whereas non-metals don’t (usually) Metals are strong and ______ but bendable. Non-metals are usually _____ or they will snap. Metals will _____ when freshly cut or scratched, whereas non-metals are usually dull. Metals have higher _______ than non-metals (i.e. they weigh more) Metals can be used to make ______ (a mixture of different metals) Words - alloys, electricity, solids, weak, densities, temperatures, tough, shine

  13. The Transition Metals Some facts… 1) They are all ______ and solid (except _________) 2) They are ____ reactive than the alkali metals 3) They can form __________ compounds, usually _______ 4) They can be used as a ______ (a chemical that speeds up a reaction) Words – hard, coloured, mercury, less, catalyst, insoluble

  14. Different forms of elements and compounds + + + + + + + + + + + + + + + + + + + + + + + + + H Cl- Na+ Cl- Na+ Cl- Na+ H C H Cl- Na+ Cl- Na+ Cl- Na+ H Elements and compounds can form many different structures, including: 1) Ionic, like sodium chloride: 2) Giant covalent structures, like graphite: 3) Metallic, like iron: 4) Simple covalent molecules, like methane:

  15. Different properties Using information from previous lessons, complete this table:

  16. Group 1 – The alkali metals 03/01/2020 Watch video of these metals reacting with water (from Sky One’s Brainiac)

  17. Group 1 – The alkali metals Potassium + water potassium hydroxide + hydrogen 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) Some facts… 1) These metals all have ___ electron in their outer shell. 2) Density increases as you go down the group, while melting point ________ 2) Reactivity increases as you go _______ the group. This is because the electrons are further away from the _______ every time a _____ is added, so they are given up more easily. 3) They all react with water to form an alkali (hence their name) and __________, e.g: Words – down, one, shell, hydrogen, nucleus, decreases

  18. Trends in Group 1 Take away one of the electrons Take away one of the electrons + + Consider a sodium atom: Sodium ion Now consider a potassium atom: Potassium ion Potassium loses its electron more easily because its further away – potassium is MORE REACTIVE

  19. Group 7 – The halogens

  20. The Halogens - Chlorine Each molecule has a strong force holding the atoms together, but the forces between molecules are very weak so chlorine is a gas at room temperature and is pale yellow.

  21. The Halogens - Bromine The forces between molecules are slightly stronger so bromine is a liquid at room temperature. It is reddish-brown in colour.

  22. The Halogens - Iodine Iodine is a solid at room temperature but with gentle heating it will melt. The atoms will remain in pairs. In solid form iodine is grey like metal but gaseous iodine is purple.

  23. The halogens – some reactions H H Cl Cl Halogen + metal metal halide Na + - This can be dissolved in water to form an acid Cl Na Cl Halogen + hydrogen a hydrogen halide 1) Halogen + metal: + 2) Halogen + hydrogen: +

  24. Trends in Group 7 Add an electron Add an electron - - Consider a fluorine atom: Flouride ion Now consider a chlorine atom: Chloride ion Chlorine doesn’t gain an electron as easily as fluorine so it is LESS REACTIVE

  25. Displacement reactions Decreasing reactivity To put it simply, a MORE reactive halogen will displace a LESS reactive halogen from a solution of its salt.

  26. Group 0 – The Noble gases 03/01/2020 • Questions: • How many electrons do these elements have in their outer shell? • How does this affect their reactivity?

  27. How the Noble Gases were discovered Sir William Ramsay, 1852-1916 Ramsay’s blue plaque outside a house in Notting Hill A while ago we discussed the idea that Mendeleev left gaps in the periodic table to account for undiscovered elements: • I discovered most of the Noble Gases around the year 1900. I did this by: • Noticing the density of nitrogen made in certain reactions differed from the density of nitrogen in air • I then developed a hypothesis about the existence of undiscovered elements • I then tested this hypothesis with further experiments.

  28. Group 0 – The Noble gases Some facts… • All of the noble gases have a full outer shell, so they are very ______ • They all have _____ melting and boiling points and are inflammable 3) They exist as single atoms rather then _________ molecules • Helium is ________ then air and is used in balloons and airships (as well as for talking in a silly voice) • Argon is used in light bulbs (because it is so unreactive) and argon , krypton and ____ are used in fancy lights Words – neon, stable, low, diatomic, lighter

  29. Properties of the Noble Gases What numbers would you expect Xenon to have?

  30. Topic 5 – Chemical Reactions

  31. Endothermic and exothermic reactions 03/01/2020 Energy Energy Step 1: Energy must be SUPPLIED to break bonds: Step 2: Energy is RELEASED when new bonds are made: A reaction is EXOTHERMIC if more energy is RELEASED then SUPPLIED. If more energy is SUPPLIED then is RELEASED then the reaction is ENDOTHERMIC

  32. Common examples of these reactions Burning Photosynthesis Cooling packs Hand warmer packs Are these reactions exothermic or endothermic?

  33. Example reactions 03/01/2020

  34. Energy level diagrams 03/01/2020 Energy required to break bonds (endothermic) Energy given out making bonds (exothermic) Energy level Reaction progress

  35. Exothermic vs endothermic: 03/01/2020 EXOTHERMIC – more energy is given out than is taken in (e.g. burning, respiration) ENDOTHERMIC – energy is taken in but not necessarily given out (e.g. photosynthesis)

  36. Rates of Reaction Oh no! Here comes another one and it’s got more energy… No effect! It didn’t have enough energy! Here comes another one. Look at how slow it’s going… Hi. I’m Mike Marble. I’m about to have some acid poured onto me. Let’s see what happens… It missed! Here comes an acid particle…

  37. Measuring the Rate of Reaction Two common methods:

  38. Rates of Reaction Chemical reactions occur when different atoms or molecules _____ with each other but they HAVE to collide with enough _______. • Basically, the more collisions we get and the more energetic they are the _______ the reaction goes. The rate at which the reaction happens depends on four things: • The _______ of the reactants, • Their concentration • Their surface area • The ______ the reactants are under Words – energy, quicker, pressure, temperature, collide

  39. Catalysts Summary Catalysts are used to ____ __ a reaction to increase the rate at which a product is made or to make a process ________. They are not normally ___ __ in a reaction. Cars use catalytic converters to remove unwanted gases. They take gases like carbon _______ and react them with water to form carbon dioxide. They have a large _____ ___ and work best at high temperatures. Carbon monoxide + oxygen carbon dioxide Words – surface area, speed up, used up, cheaper, monoxide

  40. Topic 6 – Quantitative Chemistry

  41. Mass and atomic number revision 03/01/2020 MASS NUMBER = number of protons + number of neutrons 4 He SYMBOL 2 PROTON NUMBER = number of protons (obviously)

  42. Relative formula mass, Mr Relative atomic mass of O = 16 Relative atomic mass of H = 1 The relative formula mass of a compound is the relative atomic masses of all the elements in the compound added together. E.g. water H2O: Therefore Mr for water = 16 + (2x1) = 18 Work out Mr for the following compounds: • HCl • NaOH • MgCl2 • H2SO4 • K2CO3 H=1, Cl=35 so Mr = 36 Na=23, O=16, H=1 so Mr = 40 Mg=24, Cl=35 so Mr = 24+(2x35) = 94 H=1, S=32, O=16 so Mr = (2x1)+32+(4x16) = 98 K=39, C=12, O=16 so Mr = (2x39)+12+(3x16) = 138

  43. Empirical formulae Empirical formulae is simply a way of showing how many atoms are in a molecule (like a chemical formula). For example, CaO, CaCO3, H20 and KMnO4 are all empirical formulae. Here’s how to work them out: A classic exam question: Find the simplest formula of 2.24g of iron reacting with 0.96g of oxygen. Step 1: Divide both masses by the relative atomic mass: For iron 2.24/56 = 0.04 For oxygen 0.96/16 = 0.06 Step 2: Write this as a ratio and simplify: 0.04:0.06 is equivalent to 2:3 Step 3: Write the formula: 2 iron atoms for 3 oxygen atoms means the formula is Fe2O3

  44. Example questions • Find the empirical formula of magnesium oxide which contains 48g of magnesium and 32g of oxygen. • Find the empirical formula of a compound that contains 42g of nitrogen and 9g of hydrogen. • Find the empirical formula of a compound containing 20g of calcium, 6g of carbon and 24g of oxygen. MgO NH3 CaCO3

  45. Calculating percentage mass Mass of element Ar x100% Percentage mass (%) = Relative formula mass Mr If you can work out Mr then this bit is easy… Calculate the percentage mass of magnesium in magnesium oxide, MgO: Ar for magnesium = 24 Ar for oxygen = 16 Mr for magnesium oxide = 24 + 16 = 40 Therefore percentage mass = 24/40 x 100% = 60% • Calculate the percentage mass of the following: • Hydrogen in hydrochloric acid, HCl • Potassium in potassium chloride, KCl • Calcium in calcium chloride, CaCl2 • Oxygen in water, H2O 3% 52% 36% 89%

  46. Recap questions 03/01/2020 • Work out the relative formula mass of: • Carbon dioxide CO2 • Calcium oxide CaO • Methane CH4 44 56 16 • Work out the percentage mass of: • Carbon in carbon dioxide CO2 • Calcium in calcium oxide CaO • Hydrogen in methane CH4 27% 71% 25%

  47. Balancing equations Sodium + water sodium hydroxide + hydrogen O Na Na H H H H H O Consider the following reaction: + + This equation doesn’t balance – there are 2 hydrogen atoms on the left hand side (the “reactants” and 3 on the right hand side (the “products”)

  48. Balancing equations Sodium + water sodium hydroxide + hydrogen O O Na Na Na Na H H H H H H H H O O 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) We need to balance the equation: + + Now the equation is balanced, and we can write it as:

  49. Some examples Mg + O2 Zn + HCl Fe + Cl2 NaOH + HCl CH4 + O2 Ca + H2O NaOH + H2SO4 CH3OH + O2 MgO ZnCl2 + H2 FeCl3 NaCl + H2O CO2 + H2O Ca(OH)2 + H2 Na2SO4 + H2O CO2 + H2O 2 2 2 3 2 2 2 2 3 2 2 2 2 2 4

  50. Calculating the mass of a product IGNORE the oxygen in step 2 – the question doesn’t ask for it Step 1: READ the equation: 2Mg + O2 2MgO E.g. what mass of magnesium oxide is produced when 60g of magnesium is burned in air? Step 2: WORK OUT the relative formula masses (Mr): 2Mg = 2 x 24 = 48 2MgO = 2 x (24+16) = 80 • Step 3: LEARN and APPLY the following 3 points: • 48g of Mg makes 80g of MgO • 1g of Mg makes 80/48 = 1.66g of MgO • 60g of Mg makes 1.66 x 60 = 100g of MgO

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