Atomic Structure & Periodic Table

Transcription

1 Year Topic 1 Atomic Structure & Periodic Table 2 Bonding 3 10 Quantitative Chemistry 4 Chemical Changes 5 Energy Changes 6 Rate of Reaction 7 Organic Chemistry 8 11 Chemical Analysis 9 Chemistry of the Atmosphere 10 Using Resources 1

2 1. An element is a pure substance made of one type of atom. 2. There are about 100 different elements. Elements are shown in the periodic table. 3. An atom is the smallest part of an element that can exist. 4. Atoms of each element are represented by a chemical symbol, e.g.represents an atom of oxygen, 5. Na represents an atom of sodium. 6. Compounds are formed from elements by chemical reactions. 7. Chemical reactions always involve the formation of one or more new substances, and often involve a detectable energy change. 8. Compounds contain two or more elements chemically combined in fixed proportions and can be represented by formulae using the symbols of the atoms from which they were formed. e.g. NaCl 9. Compounds can only be separated into elements by chemical reactions. 10. Chemical reactions can be represented by 11. word equations e.g. 12. sodium + chlorine sodium chloride 13. or 14. symbol equations e.g Na (s) + Cl2 (g) 2NaCl (s) 16. In chemical equations, the three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions. 17. A mixture consists of two or more elements or compounds not chemically combined together. The chemical properties of each substance in the mixture are unchanged. 18. Mixtures can be separated by physical processes that do not involve chemical reactions and no new substances are made. 19. See different KMaps (*)for details of the following processes 20. filtration (insoluble solid from liquid e.g. sand & water) 21. crystallisation (soluble solid from liquid e.g. salt & water) Atoms 40. chromatography (soluble dyes, additives e.g. ink colours) 41. simple distillation (soluble solid or liquid from another liquid e.g. alcohol & water) 22. All substances are made of atoms. 23. Atoms are very small, having a radius of about 0.1 nm (1 x m). 24. The radius of a nucleus is less than 1/ of that of the atom (about 1 x m). 25. Almost all of the mass of an atom is in the nucleus. 26. Atoms can be represented as shown in this example: 27. The number of protons in an atom of an element is its atomic number. 28. All atoms of a particular element have the same number of protons. 29. Atoms of different elements have different numbers of protons. 30. The sum of the protons and neutrons in an atom is its mass number. 31. Atoms have no overall electrical charge because the number of electrons is equal to the number of protons in the nucleus 32. The numbers of protons, neutrons and electrons in an atom or ion, can be calculated from its atomic number and mass number. 33. e.g. Na as shown above(picture at the top of this column) 34. PROTONS = ELECTRONS = Atomic Number so..na has 11 protons & 11 electrons 35. NEUTRONS = Mass Number Atomic Number so..na has = 12 neutrons 36. Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element The relative atomic mass of an element is an average value that takes account of the abundance (natural occurrence) of the isotopes of the element. (mass of isotope x % abundance of isotope) e.g. Chlorine has 2 isotopes Cl 35 75% abundance; Cl 37 25% abundance Relative atomic mass = (35 x 75) + (37 x 25) = fractional distillation (complex mixture of liquids or gases e.g. crude oil or air) in nucleus 2

3 Development of the Atomic Model 43. New experimental evidence may lead to a scientific model being changed or replaced. 44. Before the discovery of the electron, atoms were thought to be tiny spheres that could not be divided. 45. Electrons (1906) were discovered before protons (1909) and neutrons (1932) The discovery of the electron led to the plum pudding model of the atom. 48. The plum pudding model suggested that the atom is a ball of positive charge with negative electrons embedded in it. 50. The results from the alpha particle scattering experiment (see below) led to the conclusion that the mass of an atom was concentrated at the centre(nucleus) and that the nucleus was charged. This nuclear model replaced the plum pudding model. 51. Niels Bohr adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances. The theoretical calculations of Bohr agreed with experimental observations. 52. Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge. The name proton was given to these particles The experimental work of James Chadwick provided the evidence to show the existence of neutrons within the nucleus. This was about 20 years after the nucleus became an accepted scientific idea. 55. Alpha Particle Scattering Experiments 56. No deflection for most no collision (lots of empty space) 57. Deflection collision/charges interfere (electrons) 58. Rebound Collison with dense solid structure, repulsion of like charges (NUCLEUS positive) 59. Deflection is related to the mass and charge of the particles 68. The OUTER electrons dictate the chemical reactions of an element. 69. Elements with FULL OUTER SHELLS are stable and are therefore INERT they don t react e.g. Helium 70. Other elements react to form molecules and compounds in order to obtain full outer shells and increase their stability. 71. Group 1 elements readily lose their 1 outer electron which makes them very reactive. 72. Group 7 elements only need 1 more electron to become stable. They can gain or share to fill their outer shell. See Periodic Table and Bonding Topics Electronic Structure 60. The electrons in an atom occupy the lowest available energy levels(innermost available shells). 61. The electronic structure of an atom can be represented by numbers or by a diagram. 62. For example, the electronic structure of sodium is 2,8,1 or (see diagram) 63. The first lowest energy (inner) level can have a maximum of two electrons 64. The second and third energy levels can contain up to eight electrons 65. The number of electrons in each level corresponds to the number of elements in each period of the Periodic Table] 66. The last number / number of electrons in the outermost energy level is equal to the group number

4 MODERN PERIODIC TABLE 1. The elements in the periodic table are arranged in order of atomic (proton) number 2. The table is called a periodic table because similar properties occur at regular intervals. 3. Elements with similar properties are in columns, known as groups. 4. Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and this gives them similar chemical properties. 5. Elements that react to form positive ions are metals. 6. Elements that do not form positive ions are non-metals. 7. The majority of elements are metals. 8. Metals are found to the left and towards the bottom of the periodic table. 9. Non-metals are found towards the right and top of the periodic table. Development of the Periodic Table 10. Before the discovery of protons, neutrons and electrons, scientists attempted to classify the elements by arranging them in order of their atomic weights. 11. The early periodic tables were incomplete and some elements were placed in inappropriate groups if the strict order of atomic weights. 12. Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not been discovered and in some places changed the order based on atomic weights. 13. Elements with properties predicted by Mendeleev were discovered and filled the gaps e.g. Germanium, Ge 14. Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct e.g. Ar and K 4

5 GROUP 0 Noble Gases 15. The elements in Group 0 of the periodic table are called the noble gases. 16. They are unreactive and do not easily form molecules because their atoms have stable arrangements of electrons. 17. The noble gases have eight electrons in their outer shell, except for helium, which has only two electrons. 18. The atomic radius increases down the group as there are more shells 19. The boiling points of the noble gases increase with increasing relative atomic mass (going down the group). 20. Noble gases were used in light bulbs to stop the filament reacting with oxygen 21. They can also produce different coloured lights. Transition Metals TRIPLE CHEMISTRY ONLY 22. The transition elements are metals with similar properties which are different from those of the elements in Group Melting points & densities are much higher than Group 1 metals, which makes transition metals much stronger and harder. They make good building materials e.g. iron/steel 24. Transition metals are less reactive with oxygen, water and halogens than Group 1 and 2 metals. 25. Iron rusts very slowly in the presence of air(oxygen) and water. 26. Copper doesn t corrode very easily (low reactivity) which means that it stays shiny. 27. Many transition elements have ions with different charges, form coloured compounds and are useful as catalysts. 28. Iron is used as a catalyst in the making of ammonia in the Haber process; Nickel is the catalyst margarine production 29. Iron has two possible ions, Fe 2+ and Fe Copper sulfate is blue, Cobalt Chloride is pink; both compounds can be used to test for water as their colour changes when water is removed (Cu - white, Co - blue) see reversible reactions 5

6 GROUP 7: Halogens 1. The elements in Group 1 of the periodic table are known as the alkali metals and have characteristic properties because of the single electron in their outer shell. 2. The reactivity increases going down Group 1 as the outer electron is further from the positive nucleus. There is less attraction and so the electron is more easily lost. 3. Group 1 metals react very quickly with oxygen which causes the surface of the metal to become dull. They are only shiny when freshly cut. 4. Group 1 metals are very soft and are easily cut with a knife. 5. Alkali metals have low density and float on water. 6. Alkali metals react violently with water. The reaction produces hydrogen and the metal hydroxide (alkali). 7. e.g. Sodium + water sodium hydroxide + hydrogen 8. 2Na(s) + 2H2O(l) 2NaOH(aq) + H2 (g) 9. Observations during the experiment: Float on surface of water (lower density) Bubbles (gas produced pop test = hydrogen) Universal Indicator changes from green to purple (alkali produced metal hydroxide) Metal melts into a ball (heat produced - exothermic, low melting point) Potassium produces an immediate lilac flame; sodium is yellow and lithium red, only when lit. 10. Group 1 metals are stored in oil to remove oxygen and water. 11. Group 1 metals form salts, called chlorides, when they react with chlorine. The reaction is very violent. 12. sodium + chlorine sodium chloride 2Na(s) + Cl2(g) 2NaCl(s) 13. Sodium chloride is table salt, NaCl The elements in Group 7 of the periodic table are known as the halogens and have similar reactions because they all have seven electrons in their outer shell. 16. The halogens are non-metals and consist of molecules made of pairs of atoms (diatomic molecules) e.g. Cl2, Br2 17. In Group 7, the further down the group an element is the higher its relative molecular mass, melting point and boiling point. 18. Reactivity decreases going down Group 7 because the outer electrons are further away from the positive nucleus. Less attraction means it I harder to attract the extra electron to fill the outer shell. 19. A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt e.g. 20. chlorine + Sodium bromine + Sodium bromide chloride Cl2 + 2NaBr(aq) Br2 + 2NaCl(aq) Halogens form covalent bonds with nonmetals. The halogen shares 1 electron to fill its outer shell, forming molecules. 23. Halogens form ionic compounds with metals. The halogen accepts 1 electron from the metal to fill its outer shell. This forms halide ions with a -1 charge e.g. chloride, Cl Element name ends in INE 25. Compound name ends in IDE 6

7 1. States of Matter The three states of matter are solid, liquid and gas. 2. The three states of matter can be represented by a simple model. In this model, particles are represented by small solid spheres. 3. Particle theory can help to explain melting, boiling, freezing and condensing. 4. Melting and freezing take place at the melting point (M.pt). 5. Boiling and condensing take place at the boiling point (B.pt). M.pt. B.pt. 6. The amount of energy needed to change state from solid to liquid and from liquid to gas depends on the strength of the forces between the particles of the substance. 7. The nature of the particles involved depends on the type of bonding and the structure of the substance. 8. The stronger the forces between the particles the higher the melting point and boiling point of the substance. 9. (HT only) Limitations of the simple model above include that in the model there are no forces, that all particles are represented as spheres and that the spheres are solid. 10. BONDING There are three types of strong chemical bonds: ionic, covalent and metallic. 11. METALLIC bonding occurs in metallic elements and alloys. 13. For metallic bonding the particles are atoms which share delocalised electrons. 14. Metals consist of giant structures of atoms arranged in a regular pattern. 15. The bonding in metals may be represented in the following form: 16. The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure. 17. The sharing of delocalised (free) electrons gives rise to strong metallic bonds. 18. Metals have giant structures of atoms with strong metallic bonding. 19. This means that most metals have high melting and boiling points. 20. In pure metals, atoms are arranged in layers which can slide over each other to allow metals to be bent and shaped (malleable, ductile). 21. Pure metals are too soft for many uses and so are mixed with other metals to make alloys. 22. Alloys are harder than pure metals because the distortion of the layers of atoms in the structure stops the layers from sliding over each other. 23. Metals are good electrical conductors because the delocalised electrons in the metal carry electrical charge through the metal. 24. Metals are good thermal conductors because energy is transferred by the delocalised electrons 12. IONIC bonding occurs in compounds formed from metals combined with non-metals. 25. For ionic bonding the particles are oppositely charged ions. 26. When a metal atom reacts with a non-metal atom electrons in the outer shell of the metal atom are transferred. 27. Metal atoms lose electrons to become positively charged ions. 28. Non-metal atoms gain electrons to become negatively charged ions. 29. The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (full outer shell). 30. The charge on the ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 relates to the group number of the element in the periodic table. 31. Grou 32. Grou 33. Grou 34. Grou p 1 p 2 p 6 p The electron transfer during the formation of an ionic compound can be represented by a dot and cross diagram, e.g. for sodium chloride. 40. An ionic compound is a giant structure of ions held together by strong electrostatic forces of attraction between the oppositely charged ions. These forces act in all directions in the lattice and this is called ionic bonding. 41. The structure of sodium chloride can be represented in the following forms: 42. There are limitations of the dot and cross, ball and stick, two and three-dimensional diagrams to represent a giant ionic structure e.g. 43. Atoms are not solid spheres (balls) 44. The forces between atoms are not solid structures (sticks) 45. The forces act in all directions 46. The empirical formula (ratio of positive and negative ions) of an ionic compound can be deduced from a model or diagram that shows the ions in the structure. 55. Ionic compounds are high melting solids because a lot of energy is needed to overcome the strong , Na Cl , Mg , Na2 O , Mg O electrostatic attraction between the ions in the giant lattice structure. 56. Ionic compounds are soluble in water and will conduct electricity when molten or dissolved as the ions become free to move. Cl2 7

8 57. COVALENT bonding occurs in most non-metallic elements and in compounds of non-metals. 58. For covalent bonding the particles are atoms which share pairs of electrons. 59. When atoms share pairs of electrons, they form covalent bonds. 60. These bonds between atoms are strong. 61. Covalently bonded substances may consist of small molecules. 62. e.g. hydrogen, H2; chlorine,cl2; oxygen, O2; 63. nitrogen, N2; hydrogen chloride, HCl; water, H2O; ammonia, NH3; methane, CH4 64. The covalent bonds in molecules can be represented in the following forms: 65. Substances that consist of small molecules are usually gases or liquids that have relatively low melting points and boiling points. 66. These substances have only weak forces between the molecules (intermolecular forces). 67. When intermolecular forces are overcome, not the covalent bonds, the substance melts or boils. 68. The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points. 69. These substances do not conduct electricity because the molecules do not have an overall electric charge. 70. Drawing dot & cross diagrams: e.g. Cl2 71. How many outer electrons? How many are needed for full shell? How many atoms are sharing with? 1/1 in overlap make a pair 74. Subtract sharing e from outer e 7-1 =6 in outside ring 75. Some covalently bonded substances have very large molecules, such as polymers. 76. The atoms in the polymer molecules are linked to other atoms by strong covalent bonds. 77. The intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature. 78. Polymers can be represented in the form: where n is a large number. 79. Some covalently bonded substances have giant covalent structures, such as diamond and silicon dioxide (silica). 80. Substances that consist of giant covalent structures are solids with very high melting points. 81. All of the atoms in these structures are linked to other atoms by strong covalent bonds which have to be overcome to melt or boil these substances. 82. In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity (no delocalised electrons). 83. In graphite, each carbon atom forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers. 84. In graphite, one electron from each carbon atom is delocalised therefore graphite is similar to metals and will conduct electricity. 85. Graphene is a single layer of graphite and has properties that make it useful in electronics and composites. 86. Fullerenes are molecules of carbon atoms with hollow shapes. 87. The structure of fullerenes is based on hexagonal rings of carbon atoms but they may also contain rings with five or seven carbon atoms. The first fullerene to be discovered was Buckminsterfullerene (C60) which has a spherical shape. 88. Carbon nanotubes are cylindrical fullerenes with very high length to diameter ratios. Their properties make them useful for nanotechnology, electronics and materials. 89. TRIPLE CHEMISTRY ONLY 90. Nanoscience refers to structures that are nanometers (nm) in size, of the order of a few hundred atoms 91. 1,000,000,000nm = 1m; 1nm = 1x10-9 m 92. Surface area = w x h x6 sides; Volume = w x h x d 93. Nanoparticles, are smaller than fine particles (PM2.5), which have diameters between 100 and 2500nm (1 x 10-7 m and 2.5 x 10-6 m). 94. Coarse particles (PM10) have diameters between 1 x 10-5 m and 2.5 x 10-6 m. Coarse particles are often referred to as dust. 95. Nanoparticles may have properties different from those for the same materials in bulk because of their high surface area to volume ratio. 96. As the side of cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of It may also mean that smaller quantities are needed to be effective than for materials with normal particle sizes. 98. Nanoparticles have many applications in medicine, in electronics, in cosmetics and sun creams, as deodorants, and as catalysts. 99. New applications for nanoparticulate materials are an important area of research, many of the potential risks are still unknown. 8

9 QUANTITATIVE CHEMISTRY & THE LAW OF CONSERVATION OF MASS 1. The relative formula mass (Mr) of a compound is the sum of the relative atomic masses of the atoms in the numbers shown in the formula. e.g. CaCO x3 = The law of conservation of mass states that no atoms are lost or made during a chemical reaction so the mass of the products equals the mass of the reactants. 3. This means that chemical reactions can be represented by symbol equations which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation. CH 4 + 2O 2 2H 2O + CO atom of carbon, 4 atoms of hydrogen and 4 atoms of oxygen on both sides 5. Large numbers before symbols (2O2)multiplies the number of atoms or molecules (not chemically joined) 6. Small numbers following symbols multiplies the previous atom only (CH4) indicates the number of that atom within a molecule (chemically joined) 7. In a balanced chemical equation, the sum of the relative formula masses of the reactants in the quantities shown equals the sum of the relative formula masses of the products in the quantities shown. 8. Mr reactants = Mr products CH 4 + 2O 2 2H 2O + CO 2 (12+4) + 2(16x2) 2(1x2+16) + 12+(16x2) = = Some reactions may appear to involve a change in mass but this can usually be explained because a reactant or product is a gas and its mass has not been taken into account e.g. 10. thermal decompositions of metal carbonates - carbon dioxide is produced and escapes into the atmosphere leaving the metal oxide as the only solid product CuCO3 CuO + CO2 11. metal reacts with oxygen - the mass of the oxide produced is greater than the mass of the metal 12. 2Mg + O2 2MgO MOLES (HT only) 13. Chemical amounts are measured in moles. 14. The symbol for the unit mole is mol. 15. The mass of one mole of a substance in grams is numerically equal to its relative formula mass e.g. 1 mole of HF has a mass of = 20g 16. mol = mass/mr How many moles in 5g of HF? 5/20 = 0.4 mol What is the mass of 5 moles of HF? mass = mol x Mr 5x g 17. One mole of a substance contains the same number of particles e.g. atoms, molecules or ions as one mole of any other substance. 18. The number of atoms, molecules or ions in a mole of a given substance is the Avogadro constant. 19. The value of the Avogadro constant is 6.02 x particles per mole. How many molecules in 0.5mol HF? 1 mol contains 6.02 x atoms 0.5 mol = 6.02 x x 0.5 = 3.01 x REACTING MASSES (HT only) 20. The masses of reactants and products can be calculated from balanced symbol equations. 21. Chemical equations can be interpreted in terms of moles. 22. The number of moles reacting in a balanced equation is a fixed ratio Mg (s) + 2HCI (aq) MgCI 2(aq) + H 2(g) Moles Mass 1x24g 2(1+35.5)g 24+(2x35.5)g 1x2g 24g 73g 95g 2g 23. How many grams of Magnesium chloride are produced from 6g of magnesium? Equation Mg (s) MgCI 2(aq) Mole ratio from equation 1 1 Molar mass ratio 24g 95g Reduce to 1g 1g 95/24 Adjust to mass in question 6g 95/24 x 6 = 23.75g 24. The balancing numbers in a symbol equation can be calculated from the masses of reactants and products by 25. converting the masses in grams to amounts in moles (mol = mass/mr) 26. converting the numbers of moles to simple whole number ratios see Magnesium Oxide experiment in chemical change topic. 27. e.g. Write a balanced equation for decomposition of BaCO3 given that 30.6g Barium oxide, BaO are formed from 39.4g of carbonate. BaCO3 BaO + CO2 Mass ratio 39.4g 30.6g No. moles Mole ratio 1 1 Balanced Equation BaCO3 BaO + CO2 9

10 28. Empirical formula is the simplest whole number ratio of atoms in a compound 29. The mass or percentage mass can be used 30. If percentages are given remember to check that they add up to 100% EMPIRICAL FORMULA Elements Ca C O Mass or percentage of each Divide by Ar = Mole ratio 40/40 = 1 12/12 = 1 48/16 = 3 Divide all by smallest No. 1/1 = 1 1/1 = 1 3/1 = 3 Simplest whole number ratio CaCO3 LIMITING REACTANTS (HT only) 31. In a chemical reaction involving two reactants, it is common to use an excess of one of the reactants to ensure that all of the other reactant is used. 32. The reactant that is completely used up is called the limiting reactant because it limits the amount of products. 33. Values that fall outside the mole or mass ratio from the equation indicate limiting reagents and those in excess. 34. Calculate the mole ratio and mass ratio, then compare to values in the question 2Ca + O2 2CaO Mole ratio If there were 10 moles of Ca and 2 moles of oxygen, (2 moles oxygen needs 4 moles Ca) Mass ratio 2x40 2x16 2(40+16) If there were 40g of Ca and 50g of oxygen (40g Ca needs 16g of oxygen) Percentage (%) Yield 35. Even though no atoms are gained or lost in a chemical reaction, it is not always possible to obtain the calculated amount of a product because: the reaction may not go to completion because it is reversible some of the product may be lost when it is separated from the reaction mixture some of the reactants may react in ways different to the expected reaction. 36. The amount of a product obtained is known as the yield. 37. When compared with the maximum theoretical amount as a percentage, it is called the percentage yield. 38. % Yield = Mass of product actually made x 100 Maximum theoretical mass of product 39. (HT only) The theoretical mass of a product can be calculated from a given mass of reactant and the balanced equation for the reaction. Calculate the mass of ammonia that can be made from 12.0g hydrogen 20.3g of ammonia was formed in this reaction. Calculate the percentage yield. N2 + 3H2 2NH3 3 moles 2 moles 3 (2 x 1)g 2 (14 + 3x1)g 6g 34g So 12g 2x34g = 78g % Yield = Mass of ammonia actually made x 100 Maximum theoretical mass of ammonia 20.3 x 100 = 26.0% Atom Economy 41. Atom economy (atom utilisation) is a measure of the amount of starting materials that end up as useful products. 42. Other products are called by-products. 43. It is important for sustainable development and for economic reasons to use reactions with high atom economy. 44. The percentage atom economy of a reaction is calculated using the balanced equation for the reaction as follows: Mr of desired product from equation x 100% Sum of Mr of all reactants from equation (HT only) 45. Atom economy can explain why a particular reaction pathway is chosen to produce a specified product given appropriate data such as yield, rate, equilibrium position and usefulness of by-products. Hydrogen can be made in several ways. Method 1 CO + H2O CO2 + H2 Method 2 CH4 + 2H2O CO2 + 4H2 46. Calculate the atom economy for each method and explain which is more preferable. Method 1 2 x 1 X 100 % = 2 x 100 = 4.34% (12+16) + (2x1) Method 2 4 x (2x1) X 100 % = 8 x 100 = 15.38% 12+(4x1) + 2(2x1) Method 2 is preferable as more of the reagent atoms are converted to the desired product (H2) 10

11 CONCENTRATION by mass 47. Many chemical reactions take place in solutions. 48. The concentration of a solution can be measured in mass per given volume of solution, e.g. grams per dm3 (g/dm3). 49. Dilute solutions have a low concentration less mass per unit volume 50. Concentration = mass volume dm 3 = 1000cm 3 What is the concentration of a solution containing 12g of NaCl in 500cm 3? Conc = /1000 = 24 g/dm Mass = conc x volume (cm 3 ) 1000 What mass of HCl needs to be dissolved in 100cm 3 to produce a standard solution of concentration of 45g/dm 3? 52. Mass = 45 x 100/1000 = 4.5g CONCENTRATION by moles CHEM ONLY HT 53. The concentration of a solution can be measured in mol/dm The amount in moles of solute or the mass in grams of solute in a given volume of solution can be calculated from its concentration in mol/dm Titration Calculation Calculate the number of moles of the substance with known concentration usually 25cm 3 alkali in the burette. Use the equation to calculate the molar ratio of acid to alkali Calculate the number of moles of the other substance using the molar ratio and answer to step 1 Calculate the concentration of the unknown using step 3 and average titre volume from the experiment 56. Example: n1=c x v/1000 c is given in question v= 25/1000 HCl + NaOH = 1 : 1 H2SO4 + 2KOH = 1 : 2 nalkali = nacid 2nalkali = nacid so nacid = nalkali 2 c = n2 v/1000 GASES CHEM ONLY HT 57. Equal amounts in moles of gases occupy the same volume under the same conditions of temperature and pressure. 58. The volume of one mole of any gas at room temperature and pressure (20 o C and 1 atmosphere pressure) is 24 dm 3. Volume = mol x 24 dm The volumes of gaseous reactants and products can be calculated from the balanced equation for the reaction. e.g. What volume of hydrogen is needed to react with 10dm 3 of oxygen? 1mole 2 moles so 10 dm 3 2x10 = 20dm 3 equation O2 (g) + 2H2 (g) 2H2O (g) moles volume 24 dm 3 2 x 24 dm 3 2 x 24 dm 3 Mass of 1 mol 2x16= 32g 2x2x1= 4g 2(2x1+16)= 36g 60. Volumes of a gas at room temperature and pressure can be calculated from its mass and Mr What volume of hydrogen is needed to produce 12g of water? mol H2O = 12/36 so mol H2 = 1/3 so vol = 24 x 1/3 = 8 dm 3 What volume of hydrogen gas reacts with 160cm 3 of oxygen at the same temperature and pressure? 1 mol oxygen reacts with 2 mol hydrogen 160cm 3 reacts with 2 x 160cm 3 =320cm Remember to measure the volume of all solutions in a titration using the meniscus at eye level cm 3 portions of a sodium hydroxide were titrated with a standardised solution of 0.75 mol/dm3 sulfuric acid solution using phenolphthalein indicator. If the average titration was cm 3 of sulfuric acid, what is the molar concentration of the sodium hydroxide? 2NaOH(aq) + H2SO4(aq) Na2SO4 + 2H2O(l) moles H2SO4 in titration = concentration H2SO4 x volume in dm 3 = 0.75 x (17.70/1000) = mol From the balanced equation, for every mole of H2SO4, two moles of NaOH react, therefore moles NaOH = 2 x moles H2SO4 moles NaOH = x 2 = mol Concentration of NaOH = moles NaOH / volume in dm3 = / (25.00/1000) = mol/dm 3 11

12 CHEMICAL CHANGE 1. A chemical reaction involves the rearrangement of atoms to produce new substances. 2. Signs of a chemical change: Colour change Bubbles + gas produced Temperature change Solid forms (precipitate) Smell changes ph changes 3. Metals react with oxygen to produce metal oxides. 4. The reactions are oxidation reactions because the metals gain oxygen. 5. The oxidation of magnesium produces UV light and white magnesium oxide 6. 2Mg + O2 2MgO 7. The oxidation of magnesium can be used to find the empirical formula of magnesium oxide. Reactivity of Metals 1. When metals react with other substances the metal atoms form positive ions. 2. The reactivity of a metal is related to its tendency to form positive ions. 3. Metals can be arranged in order of their reactivity in a reactivity series. 4. The non-metals hydrogen and carbon are often included in the reactivity series. 5. Metals above hydrogen in the reactivity series can displace hydrogen from acids. 8. The masses of Mg and O can be converted to moles to give the mole ratio and hence empirical formula of the compound: : = 1.2 : 1 9. Experimental errors lead to inaccuracy i.e. the Mg is too high because not all the Mg reacted Extraction of Metals 18. Unreactive metals such as gold are found in the Earth as the metal itself (native metals) 19. Most metals are found as compounds that require chemical reactions to extract the metal from rocks called ores. 20. Metals less reactive than carbon can be extracted from their oxides by reduction with carbon. 21. Reduction involves the loss of oxygen. 22. Iron ore contains iron oxide (Fe2O3)which is reduced in a Blast Furnace with coke (carbon) 23. Coke burns in oxygen C + O2 CO2 Oxidation of C 24. Carbon dioxide reacts with more coke CO2 + C 2CO Oxidation of C, reduction of CO2 25. Carbon monoxide reduces the iron ore Fe2O3 + 3CO 2Fe + 3CO2 Reduction of Fe2O3 Oxidation of CO 26. REDOX reactions involve both reduction and oxidation 27. Oxidation and reduction can also be described in terms of electrons 28. Oxidation is the loss of electrons O.I.L 29. Reduction is the gain of electrons R.I.G HT only 30. Ionic equations are used to describe the oxidation and reduction within displacement reactions. 31. Ions that do not change in a reaction are called spectator ions e.g. Cl - in the equation below 32. Overall displacement reaction: CuCl2 + Ca Cu + CaCl2 33. Ionic equation: Cu 2+ + Ca Ca 2+ + Cu 34. Half equations show the oxidation and reduction steps (loss and gain of electrons) 35. Oxidation half equation: Ca Ca e - metals become positive ions 36. Reduction half equation: Cu e - Cu metal ions become metal atoms The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids. 8. Metal + acid metal salt + hydrogen 9. Zinc + hydrochloric acid zinc chloride + hydrogen 10. Observations: bubbles =gas +pop test, metal disappears 11. Metal + water metal hydroxide + hydrogen 12. Calcium + water calcium hydroxide + hydrogen 13. Observations: bubbles =gas pop test 14. Universal Indicator changes from green to purple 15. A more reactive metal can displace a less reactive metal from a compound. 16. Iron + copper sulfate iron sulfate + copper 17. Observations: silver nail becomes coated in orange metal, blue solution fades to pale green 12

13 Acids & Alkalis 1. Acids produce hydrogen ions (H + ) in aqueous solutions. 2. Aqueous solutions (aq) are soluble substances dissolved in water 3. Hydrochloric Acid HCl 4. Sulfuric Acid H2SO4 5. Aqueous solutions of alkalis contain hydroxide ions (OH )., e.g. sodium hydroxide, NaOH 6. The ph scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a ph probe. 20. Acids are neutralised by alkalis (e.g. soluble metal hydroxides) bases (e,g. insoluble metal hydroxides and metal oxides) 21. alkalis and bases react with acids to produce a salts and water 22. metal carbonates to produce salts, water and carbon dioxide. 23. The particular salt produced in any reaction between an acid and a base or alkali depends on the positive ions from the base (usually a metal) the acid: (chloride sulfate nitrate) 29. Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates (see diagrams for method below). 30. The solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt. 31. Salt solutions can be crystallised to produce solid salts A solution with ph 7 is neutral. 9. Aqueous solutions of acids have ph values of less than Strong acids have ph -0-3 and weak acids Aqueous solutions of alkalis have ph values greater than Strong alkalis have ph and weak alkalis ph In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water. 14. H + + OH - H 2O (HT only) 15. A strong acid is completely ionised in aqueous solution. Examples of strong acids are hydrochloric, nitric and sulfuric acids. HCl H + + Cl A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids. HA H + + A dilute and concentrated relates to the amount of substance dissolved in a given volume g/dm 3 or mol/dm 3 See Quantitative Chem Unit for calculation 18. For a given concentration of aqueous solutions, the stronger an acid, the lower the ph. 19. As the ph decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of Sodium hydroxide + hydrochloric acid Sodium chloride + water NaOH + HCl NaCl + H2O 25. The formulae of common ions can be used to deduce the formulae of salts e.g. Na + Mg 2+ Cl - Na Cl Mg Cl2 SO4 2- Na2 SO4 Mg SO4 (HT only) 26. The reaction between metals and acids are REDOX reactions 27. Magnesium, zinc and iron react with hydrochloric and sulfuric acids to produce hydrogen 28. Mg + 2HCl MgCl2 + H2 Oxidation: Mg Mg 2+ +2e - Reduction: 2H + + 2e - H2 TITRATION - CHEM ONLY 33. The volumes of acid and alkali solutions that react with each other can be measured by titration using a suitable indicator. See Quantitative Chem Unit for calculation Add known volume (pipette)of solution Add suitable indicator such as phenolphthalein 36. Record starting volume on burette containing solution 2. Use bottom of meniscus at eye level for greater precision 37. Add volume of liquid from burette until equivalence / end point is reached first colour change 38. Swirl to mix the acid and alkali thoroughly throughout 39. Add slowly at the end to improve accuracy of end point. 40. Note final volume of acid and calculate volume added (titre). 41. Repeat until concordant results (within 0.20cm3 of each other) 42. Calculate average titre from concordant results only 13

14 ELECTROLYSIS 1. Compounds can be split (decomposed) into their component elements by electrolysis. 2. When an ionic compound is melted or dissolved in water, the ions are free to move about within the liquid or solution. 3. These liquids and solutions are able to conduct electricity and are called electrolytes. 4. Passing an electric current through electrolytes causes the ions to move to the electrodes. 5. Electrodes are inert (unreactive) and charged from an external circuit 6. Positively charged ions (cations) move to the negative electrode (the cathode). 7. Positive ions are produced by metals or hydrogen 8. Negatively charged ions (anions) move to the positive electrode (the anode). 9. Negative ions are produced by nonmetal elements 10. Ions are discharged at the electrodes producing elements. MOLTEN COMPOUNDS 11. When a simple ionic compound is electrolysed in the molten state using inert electrodes, the metal is produced at the cathode and the nonmetal is produced at the anode. 12. e.g. Lead bromide Cathode Lead, Pb PbBr2 Anode Bromine, Br Solid lead bromide does not conduct electricity because the ions are not free to move. The solid must be heated to melt it and free the ions. Extraction of metals using Electrolysis 15. Metals can be extracted from molten compounds using electrolysis. 16. Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon. 17. Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current. 18. Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide (bauxite) and cryolite which lowers the meting point and therefore saves energy and lowers costs 19. Molten aluminium metal is produced at the negative electrode 20. Oxygen gas is produced at the positive electrodes 21. The positive carbon(graphite) anodes must be continually replaced because they wear away in the reaction with oxygen. C+O2 CO2 HT Only: 22. OIL RIG Oxidation is Loss of electrons Reduction is gain of electrons 23. Reactions at electrodes can be represented by half equations. 24. At the cathode (negative electrode), positively charged ions gain electrons and so the reactions are reductions. 25. Reduction at the Cathode (-) Al e - Al 26. At the anode (positive electrode), negatively charged ions lose electrons and so the reactions are oxidations. 27. Oxidation at the Anode (+) 2O 2- O2 + 4e - Predicting products of electrolysis 28. If MOLTEN, (-)metal at cathode, (+)non-metal at anode 29. If AQUEOUS SOLUTION, (-) hydrogen unless Cu, Ag, Au (+) oxygen unless group 7 halogen AQUEOUS SOLUTIONS 30. The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved. 31. At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen. 32. At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced. 33. This happens because in the aqueous solution water molecules break down producing hydrogen ions and hydroxide ions that are discharged. H2O (l) H + (aq) + OH - HT ONLY 34. REDUCTION 2H + + 2e - H2 35. OXIDATION 4OH - O2 + 2H2O + 4e - or 4OH - 4e - O2 + 2H2O 36. Brine is a solution of sodium chloride. 37. Positive ions: sodium (Na + ) and hydrogen (H + ) 38. Negative ions: chlorine (Cl - ) and hydroxide (OH - ) 39. Hydrogen is less reactive than sodium, so hydrogen gas (H2) is produced at the negative electrode. 40. Chlorine gas (Cl2) is produced at the positive electrode. 41. Sodium hydroxide is produced from the ions that remain in solution. 42. Sodium hydroxide is used to make soap. 43. Electroplating is coating a metal object with a different metal. 44. Silver ions (Ag + ) move to the negatively charged spoon and become silver atoms 45. Ag + + e - Ag 14

15 ENERGY 46. Energy is conserved in chemical reactions. The amount of energy in the universe at the end of a chemical reaction is the same as before the reaction takes place. 47. If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred. 48. An exothermic reaction is one that transfers energy to the surroundings so the temperature of the surroundings increases. 49. Exothermic reactions include combustion, many oxidation reactions and neutralisation. 50. EX = out (like exit); EN = in (like entrance); Therm = heat 51. Everyday uses of exothermic reactions include self-heating cans and hand warmers. 52. An endothermic reaction is one that takes in energy from the surroundings so the temperature of the surroundings decreases. 53. Endothermic reactions include thermal decompositions and the reaction of citric acid and sodium hydrogen carbonate. Some sports injury packs are based on endothermic reactions. 54. Energy changes can be measured using an insulated cup with a lid to reduce the heat loss and make measurements more accurate. 55. Chemical reactions can occur only when reacting particles collide with each other and with sufficient energy. 56. The minimum amount of energy that particles must have to react is called the activation energy. 57. Reaction profiles can be used to show the relative energies of reactants and products, the activation energy and the overall energy change of a reaction. 58. In an endothermic reaction, the energy of the products is higher than the reactants 59. In an exothermic reaction the products are lower energy than the reactants 60. During a chemical reaction energy must be supplied to break bonds in the reactants energy is released when bonds in the products are formed. 66. The energy transferred in chemical reactions can be calculated using bond energies that will be given in a question. 67. Calculate the overall enthalpy change for the following reaction and state why it is endothermic or exothermic. CH3OCH3 + 3 O2 2 CO2 + 3 H2O The energy needed to break bonds and the energy released when bonds are formed can be calculated from bond energies. 63. The difference between the sum of the energy needed to break bonds in the reactants and the sum of the energy released when bonds in the products are formed is the overall energy (enthalpy) change of the reaction, H. H = bonds broken(energy IN) - bonds made(energy OUT) 64. In an exothermic reaction, the energy released from forming new bonds is greater than the energy needed to break existing bonds. Exothermic OUT (make) > IN (break) 65. In an endothermic reaction, the energy needed to break existing bonds is greater than the energy released from forming new bonds. Endothermic IN (break) > OUT (make) Calculate sum of all bonds broken (left side) 6 x 413(CH) + 2x 358(CO) + 3x 498(OO) =4688 Calculate sum of all bonds made (right side) 4x 745(CO) + 6x 467(OH) =5782 Calculate difference Left (IN) Right (OUT) = kj/mol More energy out than in so temperature will rise EXOthermic reaction. 15

16 Cells & Batteries CHEM ONLY 69. Cells contain chemicals which react to produce electricity. 70. The voltage produced by a cell is dependent upon a number of factors including the type of electrode and electrolyte. 71. A simple cell can be made by connecting two different metals in contact with an electrolyte. Fuel Cells CHEM ONLY 79. Fuel cells are supplied by an external source of fuel (e.g. hydrogen) and oxygen or air. 80. The fuel is oxidised electrochemically within the fuel cell to produce a potential difference (voltage). 72. Batteries consist of two or more cells connected together in series to provide a greater voltage. 73. In non-rechargeable cells and batteries the chemical reactions stop when one of the reactants has been used up. 74. Alkaline batteries are non-rechargeable. 75. Rechargeable cells and batteries can be recharged because the chemical reactions are reversed when an external electrical current is supplied The overall reaction in a hydrogen fuel cell involves the oxidation of hydrogen to produce water. 2H2 + O2 2H2O (HT only) 83. Oxidation H2 2H + + 2e - Reduction O2 + 4e - 2O Hydrogen fuel cells offer a potential alternative to rechargeable cells and batteries. 76. The relative reactivity of different metals used in each half cell affects the voltage. 77. The bigger the difference in reactivity, the larger the voltage. 78. The direction of electron flow (current) affects the sign but not the value of the voltage 85. Hydrogen fuel cells produce less waste and pollution in comparison with rechargeable cells and batteries. There are no metals and potentially toxic chemicals sent to landfill. 86. Hydrogen is a flammable gas which increases fire risk and issues with safe storage of the fuel. Gasses need to be compressed to reduce storage space and there is currently no easy way to distribute the fuel. 87. Hydrogen fuel cells are more efficient and quiet than petrol/diesel engines but they are more expensive. 88. Hydrogen is abundant and readily available, however many of the extraction methods may rely on fossil fuels, which causes pollution and is not sustainable. E.g. using methane as a raw material to produce hydrogen 16

17 RATE OF REACTION 1. The rate of a chemical reaction can be found by measuring the quantity of a reactant used or the quantity of product formed over time: 2. Mean rate of reaction = quantity of reactant used time taken OR Mean rate of reaction = quantity of product formed time taken 3. The quantity of reactant or product can be measured by the mass in grams or by a volume in cm The units of rate of reaction may be given as g/s or cm 3 /s HT only mol/s. 5. Graphs show the quantity of product formed or quantity of reactant used up against time 6. The slope/gradient on these graphs is a measure of the rate of reaction. 7. (HT only) Draw and calculate the gradient of a tangent to the curve on these graphs as a measure of rate of reacting at a specific time. 8. Describing Graphs means say what you see. If numbers are given in the graph, quote them where appropriate. 9. As temperature increases so does rate of reaction. This means that reactions finish faster at higher temperatures, as the graph shows the reactant is used up faster at 100 o C, so it levels off sooner. 10. As temperature increases, rate of reaction increases very quickly. As temperature continues to increase the rate of reaction increases more slowly. Eventually the rate of reaction levels-off as reagents are used up Collision theory explains how various factors affect rates of reactions. 14. Chemical reactions can occur only when reacting particles collide with each other and with sufficient energy. 15. The rate of a reaction depends on two things: the frequency of collisions between particles. The more often particles collide, the more likely they are to react. the energy with which particles collide. If particles collide with less energy than the activation energy, they will not react. 16. The minimum amount of energy that particles must have to react is called the activation energy. 17. Increasing temperature increases the speed of the particles (because they gain kinetic energy) so they collide often and with more energy, so more collisions have the energy greater than the activation energy. This increases the rate of reaction. 18. Increasing the pressure of gases brings the particles closer together so there are more frequent collisions. This increases the rate of reaction. 19. Increasing the concentration of reactants increases the number of particles, so there are more frequent collisions which increases the rate of reaction. 20. Increasing the surface area of a SOLID (you cannot change the surface area of a liquid or gas) increases the frequency of collisions and hence increases the rate of reaction. 21. Larger particles have a smaller surface area 22. To increase rate, use powdered solids. 23. Be prepared to calculate surface area of a model to show this e.g. for a cube w x h x 6 sides (mm 2 ) 17

18 Catalysts 24. Catalysts change the rate of chemical reactions but are not used up during the reaction. 25. Catalysts are not included in the chemical equation for the reaction. 26. Different reactions need different catalysts. 27. Iron is he catalyst for the production of ammonia in the Haber process 28. Nickel is the catalyst for the production of margarine 29. Enzymes act as catalysts in biological systems. 30. Catalysts increase the rate of reaction by providing a different pathway for the reaction that has a lower activation energy. 31. A reaction profile for a catalysed reaction can be drawn in the following form Measuring Rate of Reaction 32. Colour changes can be measured using a disappearing cross, or light sensor. 33. Mass changes can be measured where gases are produced and escape the reaction vessel 34. Gas production can be measured by displacement of water (inverted measuring cylinder or burette) or in a gas syringe. 35. Errors and uncertainty in measurement and the resolution of the measuring apparatus should be considered. 36. Using sensors reduces human errors e.g. when does the cross actually disappear? 37. Using more precise equipment, increases the resolution and reduces uncertainty in measurements e.g. a burette has smaller scale divisions than a measuring cylinder, so uncertainty improves to 0.1cm 3 from 1cm 3 per reading. 38. If a start and end volume or temperature reading are taken remember to double the uncertainty in the overall measurement. i.e. the difference includes the uncertainty in each of the start and end readings

19 Reversible Reactions 1. In some chemical reactions, the products of the reaction can react to produce the original reactants. Such reactions are called reversible reactions and are represented: A + B C + D 2. For example: ammonium chloride decomposes into two gases when heated and returns to a white solid as the vapours cool If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction. The same amount of energy is transferred in each case. For example: The anhydrous copper sulfate is used as a test for water white blue colour change 7. Anhydrous cobalt chloride can also be used blue pink. 8. The direction of reversible reactions can be changed by changing the conditions. 9. When a reversible reaction occurs in apparatus which prevents the escape of reactants and products, equilibrium is reached when the forward and reverse reactions occur at exactly the same rate. HT ONLY 10. The relative amounts of all the reactants and products at equilibrium depend on the conditions of the reaction. 11. The effects of changing conditions on a system at equilibrium can be predicted using Le Chatelier s Principle. 12. If a system is at equilibrium and a change is made to any of the conditions, then the system responds to counteract the change. 13. If the concentration of one of the reactants or products is changed, the system is no longer at equilibrium and the concentrations of all the substances will change until equilibrium is reached again. If the concentration of a reactant is increased, more products will be formed until equilibrium is reached again. If the concentration of a product is decreased, more reactants will react until equilibrium is reached again. 14. If the temperature of a system at equilibrium is increased: the relative amount of products at equilibrium increases for an endothermic reaction the relative amount of products at equilibrium decreases for an exothermic reaction. 15. If the temperature of a system at equilibrium is decreased: the relative amount of products at equilibrium decreases for an endothermic reaction the relative amount of products at equilibrium increases for an exothermic reaction. 16. For gaseous reactions at equilibrium: an increase in pressure causes the equilibrium position to shift towards the side with the smaller number of molecules as shown by the symbol equation for that reaction a decrease in pressure causes the equilibrium position to shift towards the side with the larger number of molecules as shown by the symbol equation for that reaction. 19

20 Organic Chemistry 1. Organic chemistry is the study of carbon based molecules and their reactions. 2. Carbon is a non-metallic element that can form a wide range of molecules which are the basis of life on our planet. 3. Fats, proteins and carbohydrates are all carbon based molecules. 4. Hydrocarbons are made of carbon and hydrogen atoms only. 5. Carbon belongs to group 4 of the Periodic Table and forms 4 covalent bonds. 6. The vast array of natural and synthetic carbon compounds occur due to the ability of carbon atoms to form families of similar compounds. 7. Compounds with the same functional group react in similar ways and belong to the same chemical family or homologous series. 8. Homologous series is a group of compounds with a general formula and properties that change by a regular amount. 9. The functional group dictates its reactions and properties. 10. Some properties of hydrocarbons depend on the size of their molecules, including boiling point, viscosity and flammability. 11. These properties influence how hydrocarbons are used as fuels. 12. Viscosity is low when substances are runny. Fractions containing small molecules have low viscosity. 13. Smaller molecules are more flammable because they are more volatile (evaporate easily). 14. Small molecules have lower boiling points because the forces between the molecules (intermolecular forces) are weak, so less energy is needed to separate the molecules when it boils (physical change - liquid to gas). 26. Crude oil is a finite resource found in rocks which is formed from the 15. Hydrocarbons are named after the number of carbon atoms they contain and the family they belong to: Carbon number prefix remains of an ancient biomass consisting mainly of plankton that was buried in mud. 27. Crude oil is a mixture of a very large number of compounds. Most of the compounds in crude oil are hydrocarbons. 28. The many hydrocarbons in crude oil may be by fractional distillation. 29. The crude oil mixture is heated(350 o C) at the bottom of the fractionating column. Most of the components evaporate and rise. 30. As the components cool, they condense and are collected at different levels within the column which corresponds to their boiling points. 31. The separated fractions each contain molecules with a similar number of carbon atoms and hence boiling point. 32. Small molecules with the lowest boiling points are collected at the top and the largest molecules with the highest boiling points are collected at the bottom. 33. The diagram shows the range of other properties of the fractions which are related to the size of the molecules. 34. The fractions can be processed to produce fuels e.g. petrol, diesel; and feedstock for the petrochemical industry such as solvents, lubricants, polymers, detergents. Homologous Series alkane alkene alcohol C1 meth methane methanol C2 eth ethane ethene ethanol C3 prop propane propene propanol C4 but butane butene butanol Simple Hydrocarbons Alkanes 16. Most of the hydrocarbons in crude oil are hydrocarbons called alkanes. 17. The general formula for the homologous series of alkanes is CnH2n The first four members of the alkanes are methane, ethane, propane and butane. 19. Alkane molecules can be represented by the following formula: molecular C2H6 or structural 20. Alkanes have single covalent bonds between the carbon atoms. 21. Alkanes are saturated hydrocarbons, they cannot form any more bonds (no double bonds). 22. Alkanes are not very reactive, but make very good fuels. 23. Methane (CH4) is called natural gas. It is used as the mains gas supply (fuel) at home and in the laboratory. 24. Methane is produced naturally from the decomposition of plant and animal material (biogas). 25. Methane is a greenhouse gas that increases global warming. 35. Hydrocarbon Fuels store energy in a chemical store which can be released as heat and light when the fuel is burned. 36. Combustion(burning) is a chemical reaction with oxygen (oxidation). Heat, fuel and oxygen are all needed. 37. During combustion, the carbon and hydrogen atoms in the fuel are oxidized (gain oxygen). 38. The complete combustion of a hydrocarbon produces carbon dioxide and water. 39. e.g. C4H O2 4CO2 + 5H20 (CO2 = C in hydrocarbon; H2O = half H s in hydrocarbon; O2= CO2 + half H20) 20

21 40. Incomplete combustion produces carbon monoxide and soot(carbon particles) which are harmful pollutants. 41. Impurities in fuels can increase pollution. Sulfur will produce sulfur dioxide which dissolves in water to form acid rain. Alkenes Chem only 42. Alkenes are hydrocarbons with a double carbon-carbon bond, C=C (functional group). 43. The general formula for the homologous series of alkenes is C nh 2n 44. Alkene molecules can be represented in the following formula: molecular C 3H 6 or structural 45. Alkene molecules are unsaturated because they contain two fewer hydrogen atoms than the alkane with the same number of carbon atoms (have a double bond). 46. The first four members of the homologous series of alkenes are ethene, propene, butene and pentene. 47. Alkenes are used to produce polymers and as starting materials for the production of many other chemicals. Supply & Demand 48. There is a high demand for fuels with small molecules e.g. petrol but the supply from the fractional distillation of crude oil is too low meet the demand. 49. Other larger molecules produced from crude oil have a higher supply than demand. 50. Hydrocarbons can be broken down (cracked) to produce smaller, more useful molecules. 51. Cracking is a thermal decomposition reaction. 52. Large chain alkanes are broken down into smaller, more useful alkanes and another type of hydrocarbon called alkenes. E.g. C12 H26 C8H18 + C4H8 Long alkane shorter alkane alkene 53. The total number of carbon(c) and hydrogen (H) atoms on both sides must be equal. 54. Cracking can be done by various methods including catalytic cracking and steam cracking. 55. Cracking can be done in the lab as shown 56. Catalytic cracking uses a catalyst (aluminum oxide) to speed up the reaction and reduce temperatures ( o C). The catalyst can also be regenerated and reused. This saves money, resources and energy. 57. Catalytic cracking is the best way to produce shorter alkane fuels such as petrol. 58. Steam cracking uses much higher temperatures (up to 900 o C) and high pressures. It produces more alkenes than catalytic cracking. 21

22 Reactions of alkenes 59. Alkenes are more reactive than alkanes. 60. The test for alkenes is the decolourisation of bromine water (see diagram). CHEM ONLY 61. Alkenes react with oxygen in combustion reactions in the same way as other hydrocarbons, but they tend to burn in air with smoky flames because of incomplete combustion. 62. Alkenes react with hydrogen, water and the halogens, by addition. 63. The addition of atoms across the carbon-carbon double bond changes the double bond to a single carbon-carbon bond. 64. Margarine is made by adding hydrogen across the double bonds present in unsaturated vegetable oils 65. The addition of hydrogen requires heating with a catalyst (nickel) e.g. 66. The addition of halogens (Group 7) happens at room temperature e.g. Addition polymerisation 69. Alkenes can be used to make polymers such as the common plastics poly(ethene) and poly(propene). 70. Poly(ethene) is formed from ethene (monomer). 71. Monomers always have the functional group C=C (alkenes) to allow the addition to take place. 72. The double bond breaks allowing new bonds to be formed between the monomers. 73. In addition polymerisation reactions, many small molecules (monomers) join together to form very large molecules (polymer). [mono = one, poly = many] 74. In addition polymers the repeating unit has the same atoms as the monomer because no other molecule is formed in the reaction. 75. n next to the repeating unit indicates that there are many units joined together (any integer in maths). 76. Monomers and repeating units are always drawn in this way (H shape) so that the different polymers can be identified easily. 67. Water can be added across a double bond when the alkene is heated in the presence of an acid (H + ) catalyst e.g. 68. Alcohols are produced from the reactions of alkenes and steam (water)

23 CHEM ONLY 78. Alcohols contain the functional group C OH. 79. Methanol, ethanol, propanol and butanol are the first four members of a homologous series of alcohols. 80. Alcohols can be represented in the following forms: CH3CH2OH or 81. Aqueous solutions of ethanol are produced when sugar solutions are fermented using yeast at C 82. Above this temperature the enzymes from the yeast are denatured and will not work. 83. Fermentation is a batch process that is much slower than the continuous industrial production of alcohol from ethene. 84. Fermentation uses plants as a raw material which is sustainable whereas ethene comes from finte crude oil resources. 85. Alcohols are flammable and burn to produce carbon dioxide and water e.g. CH3CH2OH + 5 O2 2 CO2 + 3 H2O 86. The energy released from burning alcohols is proportional to the number of carbon atoms. 87. CHEM ONLY 88. Carboxylic acids have the functional group COOH. 89. The first four members of a homologous series of carboxylic acids are methanoic acid, ethanoic acid, propanoic acid and butanoic acid. 90. The structures of carboxylic acids can be represented in the following forms: CH3COOH or 91. Carboxylic acids become less soluble in water as the carbon chain length increases. 92. (HT only) Carboxylic acids are weak acids due to partial ionisation 93. Carboxylic acids react with carbonates to produce carbon dioxide 94. Carboxylic acids react with alcohols to produce esters 95. Ethylethanoate is an ester produced from ethanol and ethanoic acid. CHEM ONLY 96. Condensation polymerisation involves monomers with two functional groups. 97. When these types of monomers react they join together, usually losing small molecules such as water, and so the reactions are called condensation reactions. 98. The simplest polymers are produced from two different monomers with two of the same functional groups on each monomer. 99. For example: ethane diol and hexanedioic acid polymerise to produce a polyester: CHEM ONLY 100. Amino acids have two different functional groups in a molecule Amino acids react by condensation polymerisation to produce polypeptides For example: glycine polymerises to produce the polypeptide H2NCH2COOH (-HNCH2COO-)n + n H2O 103. Different amino acids can be combined in the same chain to produce proteins DNA (deoxyribonucleic acid) is a large molecule essential for life DNA encodes genetic instructions for the development and functioning of living organisms and viruses Most DNA molecules are two polymer chains, made from four different monomers called nucleotides, in the form of a double helix Other naturally occurring polymers important for life include proteins, starch and cellulose. Polymer Polypeptides Protein Starch, Cellulose DNA Monomer Same amino acid Different Amino Acids Glucose 4 different Nucleotides 23

24 CHEMICAL ANALYSIS 1. In everyday language, a pure substance can mean a substance that has had nothing added to it, so it is unadulterated and in its natural state, e.g. pure orange juice. 2. In chemistry, a pure substance is a single element or compound, not mixed with any other substance. 3. Pure substances cannot be broken down into other substances by physical or chemical means Pure elements and compounds have specific melting & boiling temperatures. 6. Mixtures (impure substances) melt over a range of temperatures. 7. Melting point and boiling point data can be used to distinguish pure substances from mixtures. 8. Boiling points can be used to separate mixtures using distillation. GAS TESTS 19. Hydrogen: use a burning splint held at the open end of a test tube of the gas. 20. Hydrogen burns rapidly with a pop sound. 21. Oxygen: use a glowing splint inserted into a test tube of the gas. 22. The splint relights in oxygen. 23. Carbon dioxide uses an aqueous solution of calcium hydroxide (lime water). 24. When carbon dioxide is shaken with or bubbled through limewater the limewater turns milky (cloudy). 25. Chlorine uses litmus paper. 26. When damp litmus paper is put into chlorine gas the litmus paper is bleached and turns white. Chromatography 9. Chromatography can be used to separate mixtures and can help identify substances. 10. Chromatography involves a stationary phase and a mobile phase. 11. Separation depends on the distribution of substances between the phases. See required practical sheet. 12. Distribution depends on the solubility in the solvent and its mass. 13. The ratio of the distance moved by a compound (centre of spot from origin) to the distance moved by the solvent can be expressed as its Rf value: 14. Rf = distance moved by substance distance moved by solvent 15. Different compounds have different Rf values in different solvents, which can be used to help identify the compounds. 16. The measurements and answers are given to 2 significant figures 17. The compounds in a mixture may separate into different spots depending on the solvent but a pure compound will produce a single spot in all solvents. 18. Remember that the baseline MUST be drawn in pencil so that it is insoluble in the solvent and does not interfere with the chromatogram. FLAME TESTS - CHEM ONLY 27. Flame tests can be used to identify some metal ions (cations). 28. lithium compounds result in a crimson flame 29. sodium compounds result in a yellow flame 30. potassium compounds result in a lilac flame 31. calcium compounds result in an orange-red flame 32. copper compounds result in a green flame. 33. If a sample containing a mixture of ions is used some flame colours can be masked. PRECIPITATION REACTIONS 34. Sodium hydroxide solution can be used to identify some metal ions(cations). 35. Solutions of aluminium, calcium and magnesium ions form white precipitates Mg 2+ (aq) + OH - (aq) Mg(OH)2 (s) 36. aluminium hydroxide precipitate dissolves in excess sodium hydroxide solution. 37. Copper(II) forms a blue precipitate, iron(ii) a green precipitate and iron(iii) a brown precipitate. 24

25 ANION TESTS 38. Halide ions in solution produce precipitates with silver nitrate solution in the presence of dilute nitric acid. 39. Silver chloride is white, 40. silver bromide is cream 41. silver iodide is yellow. 44. Carbonates react with dilute acids to form carbon dioxide gas. Carbon dioxide can be identified with limewater. 48. Sulfate ions in solution produce a white precipitate with barium chloride solution in the presence of dilute hydrochloric acid BaCl2(aq) + CaSO4(aq) BaSO4 (s) + CaCl2(aq) 42. AgNO3(aq) + NaCl(aq) AgCl (s) + NaNO3(aq) 43. Check the state symbols for the (s) precipitate 46. CaCO3(s) + HCl(aq) CaCl2 + H2O(l) + CO2(g) 47. Bubbles will be observed - check the state symbols for the (g) gas INSTRUMENTAL METHODS - CHEM ONLY 51. Elements and compounds can be detected and identified using instrumental methods. 52. Instrumental methods are accurate, sensitive and rapid. 53. Flame emission spectroscopy is an example of an instrumental method used to analyse metal ions in solutions. 54. The sample is put into a flame and the light given out is passed through a spectroscope. 55. The output is a line spectrum that can be analysed to identify the metal ions in the solution and measure their concentrations. 56. The spectrum produced from each metal is unique, like a fingerprint. 57. The unknown sample has the same pattern of lines as strontium 58. The process can only be used with solutions 59. The machine can detect very small quantities (very sensitive) 60. Automated computer libraries of the line spectra of all elements, makes the identification of unknown samples very fast and accurate. 50. Check the state symbols for the (s) precipitate TESTS FOR CATIONS TESTS FOR ANIONS 61. A formulation is a mixture that has been designed as a useful product (see NPK fertilisers topic) 62. Many products are complex mixtures in which each chemical has a particular purpose. 63. Formulations are made by mixing the components in carefully measured quantities to ensure that the product has the required properties. 64. Formulations include fuels, cleaning agents, paints, medicines, alloys, fertilisers and foods. 25

26 The Current ATMOSPHERE 1. For 200 million years, the proportions of different gases in the atmosphere have been much the same as they are today: The Early Atmosphere 7. Theories about what was in the Earth s early atmosphere and how the atmosphere was formed have changed and developed over time. 8. Evidence for the early atmosphere is limited because of the time scale of 4.6 billion years. 9. At the start of this period the Earth s atmosphere may have been like the atmospheres of Mars and Venus today, consisting of mainly carbon dioxide with little or no oxygen gas. 2. about four-fifths (approximately 78 %) nitrogen 3. about one-fifth (approximately 21 %) oxygen 4. small proportions of various other gases, including carbon dioxide, water vapour and noble gases such as argon. 5. One theory suggests that during the first billion years of the Earth s existence there was intense volcanic activity that released gases that formed the early atmosphere and water vapour that eventually condensed to form the oceans once the Earth had cooled. 6. Volcanoes also produced nitrogen which gradually built up in the atmosphere and there may have been small proportions of methane and ammonia. 10. Algae first produced oxygen about 2.7 billion years ago and soon after this oxygen appeared in the atmosphere. 11. Over the next billion years plants evolved and the percentage of oxygen gradually increased to a level that enabled animals to evolve. 12. Algae and plants produced the oxygen that is now in the atmosphere by photosynthesis, which can be represented by the equation: 16. The combustion of fuels is a major source of atmospheric pollutants. 17. The gases released into the atmosphere when a fuel is burned may include carbon dioxide, water vapour, carbon monoxide, sulfur dioxide and oxides of nitrogen. 18. Most fuels, including coal, contain carbon and/or hydrogen and may also contain some sulfur. 19. Burning (combustion) causes the fuel and impurities to react with oxygen to produce sulfur dioxide. S + O2 SO2 20. Solid particles and unburned hydrocarbons may also be released that form particulates in the atmosphere. 21. Carbon monoxide is a toxic gas. It is colourless and odourless and so is not easily detected. 22. Sulfur dioxide and oxides of nitrogen cause respiratory problems in humans and cause acid rain. 23. Particulates cause global dimming (haze) and health problems for humans. 13. Algae and plants decreased the percentage of carbon dioxide in the atmosphere by photosynthesis. 14. When the oceans formed carbon dioxide dissolved in the water and carbonates were precipitated producing sediments (eventually forming sedimentary rocks, like limestone), reducing the amount of carbon dioxide in the atmosphere. 15. Carbon dioxide was also decreased by the formation of fossil fuels that contain carbon such as oil and natural gas. 16. The Carbon Cycle shows how Carbon dioxide is produced and stored. 24. Complete combustion of carbon fuels produces carbon dioxide 25. CH4 + 2O2 CO2 + 2H2O 26. Incomplete combustion produces carbon monoxide and soot (carbon particles). 27. C2H6 + 2O2 CO + C + 3H2O 26

27 The Greenhouse Effect & Climate Change 28. Greenhouse gases in the atmosphere maintain temperatures on Earth high enough to support life. 29. Water vapour, carbon dioxide and methane are greenhouse gases. 30. The Greenhouse Effect is a result of the interaction of short (UV) and long (IR) wavelength radiation with the Earth and its atmosphere. 32. Some human activities increase the amounts of greenhouse gases in the atmosphere, including carbon dioxide and methane An increase in average global temperature is a major cause of climate change. 39. There is a positive correlation between increasing atmospheric carbon dioxide concentration and increasing surface temperature The increasing use of fossil fuels for energy production is the main reason carbon dioxide levels have risen dramatically. 40. There are several potential effects and implications of global climate change. 41. Increased temperatures melt ice caps that cause increased sea levels and occurrence and flooding. Flooding reduces land available for crops and cattle, hence food process will rise. 42. Increased air pollution causes loss of visibility and increase in respiratory illness. 35. Based on peer-reviewed evidence, many scientists believe that human activities will cause the temperature of the Earth s atmosphere to increase at the surface and that this will result in global climate change. 36. However, it is difficult to model such complex systems as global climate change. 37. This leads to simplified models, speculation and opinions presented in the media that may be based on only parts of the evidence and which may be biased. E.g. fuel companies. 43. The carbon footprint is the total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product, service or event. 44. The carbon footprint can be reduced by reducing emissions of carbon dioxide and methane. 45. Reducing car use - car sharing, using public transport, walk, cycle 46. Reducing heating use lower thermostats for water & heating, use solar heating systems 47. Reduce electricity use turn lights, plugs, chargers off when not in use, use renewables 48. Reducing use of plastics reusable bags, reduce bottled water & plastic cups 49. Recycle and reduce waste (including food waste), 50. Actions to reduce carbon footprints may be limited because: 51. Access to recycling services is limited 52. Access to alternative products is limited or more expensive 27

28 USING RESOURCES 1. Humans use the Earth s resources to provide warmth, shelter, food and transport. 2. Natural resources, supplemented by agriculture, provide food, timber, clothing and fuels. E.g. plant materials such as wood, cotton; animal resources such as meat, wool, leather 3. Finite resources from the Earth, oceans and atmosphere are processed to provide energy and materials e.g. fossil fuel oil is made into plastics, solvents as well as fuels for cars; minerals and ores provide metals and building materials. 4. Finite materials are non-renewable, they will eventually run out. 5. Reducing usage and finding more sustainable alternatives will extend the life of finite resources. 6. Chemistry plays an important role in improving agricultural and industrial processes to provide new products and in sustainable development. 7. Sustainable development is development that meets the needs of current generations without compromising the ability of future generations to meet their own needs. 8. Sustainable development incudes reducing fuel and water consumption, reducing waste, reusing and recycling materials. Waste Water Treatment 24. Urban lifestyles and industrial processes produce large amounts of waste water that require treatment before being released into the environment. 25. Sewage and agricultural waste water require removal of organic matter and harmful microbes. 26. Industrial waste water may require removal of organic matter and harmful chemicals. 27. Sewage treatment includes: 28. screening and grit removal 29. sedimentation to produce sewage sludge and effluent 30. anaerobic digestion of sewage sludge 31. aerobic biological treatment of effluent. Potable Water 9. Water of appropriate quality is essential for life. 10. For humans, drinking water should have sufficiently low levels of dissolved salts and microbes. 11. Water that is safe to drink is called potable water. 12. Potable water is not pure water in the chemical sense because it contains dissolved substances. 13. The methods used to produce potable water depend on available supplies of water and local conditions. 14. In the United Kingdom (UK), rain provides water with low levels of dissolved substances (fresh water) that collects in the ground and in lakes and rivers, and most potable water is produced by: choosing an appropriate source of fresh water passing the water through filter beds sterilising. 15. Sterilising agents used for potable water include chlorine, ozone or ultraviolet light. 16. If supplies of fresh water are limited, desalination of salty water or sea water may be required. 17. Desalination can be done by distillation 18. Pure water boils at 100 o C 19. The salt water is heated; pure water evaporates and passes to into the condenser. As the vapour cools, pure liquid water forms and is collected. The salt and other solid impurities remain in the original flask. 20. Large scale desalination can be achieved by processes that use membranes such as reverse osmosis. 21. Membranes allow water to pass through but not salt 22. These processes require large amounts of energy. 23. Both processes also produce more concentrated salt water (brine) which cannot be returned to the ocean as it is toxic to animal and plant life. 28

29 Alternative Metal Extraction - HT ONLY 32. The Earth s resources of metal ores are limited (finite). 33. If current increased use continues, we will have no iron ore remaining in 50 years Most metals are extracted from high grade ores by smelting (reduction by heating with carbon) or electrolysis, where metals are more reactive than carbon. See chemical reactions sheet. 36. High grade ores are finite resources. There are also significant environmental issues associated with mining ores and extracting metals. 37. Copper ores are becoming scarce and new ways of extracting copper from low-grade ores include phytomining, and bioleaching. 38. These methods avoid traditional mining methods of digging, moving and disposing of large amounts of rock. 39. Phytomining uses plants to absorb metal compounds. The plants are harvested and then burned to produce ash that contains metal compounds. 40. Phytomining only reaches metal sources at the depth of its roots. 41. Bioleaching uses bacteria to produce leachate solutions that contain metal compounds. Alloys & Corrosion of Metals Triple Chemistry Only 46. Most metals in everyday use are alloys as their properties can be made fit for purpose. 47. Alloys can be designed to be stronger, lower density, lower or higher melting, resistant to corrosion etc. 48. Bronze is an alloy of copper and tin. 49. Brass is an alloy of copper and zinc. 50. Gold used as jewellery is usually an alloy with silver, copper and zinc as pure gold is too soft. 51. The proportion of gold in the alloy is measured in carats. 24 carat being 100 % (pure gold), and 18 carat being 75 % gold. 52. Steels are alloys of iron that contain specific amounts of carbon and other metals. 53. High carbon steel ( % C) is strong but brittle. 54. Low carbon steel ( % C)is softer and more easily shaped. More likely to rust than high carbon steel. 55. Steels containing chromium and nickel (stainless steels) are hard and resistant to corrosion. 56. Aluminium alloys are low density and strong, which makes them useful in aircraft manufacture. 57. Corrosion is the destruction of materials by chemical reactions with substances in the environment. 58. Rusting is an example of corrosion specific to iron. 59. Iron + oxygen Iron oxide 4Fe + 3O2 2Fe2O3 60. Both air and water are necessary for iron to rust Where air and or water are excluded, the nail does not rust. 63. Salt and acid increase the rate of rusting.. Corrosion can be prevented by applying a coating that acts as a barrier blocking access of oxygen and water. 65. Greasing, painting or electroplating works in this way, however, if the coating is damaged, the iron underneath starts to rust. 66. Aluminium has a natural oxide coating that protects the metal from further corrosion. It is the reason that the reactivity of aluminium appears to be lower than it should be. 67. This enables aluminium to be used for drink cans, electricity cables. 68. Some coatings are reactive and contain a more reactive metal to provide sacrificial protection, e.g. zinc is used to galvanise iron. 69. Unlike paint, if the galvanized coating is scratched, the iron doesn t rust because the zinc corrodes first, protecting the iron, even when oxygen and water are able to reach the surface The metal compounds can be processed to obtain the metal. For example, copper can be obtained from solutions of copper compounds by displacement using scrap iron or by electrolysis. 44. These methods use less energy; produce less pollution and waste than mining as well as using ores or waste material not suitable for smelting. 45. The methods are slower and produce less metal than large scale smelting of high grade ores

30 Life Cycle Assessment 71. Life cycle assessments (LCAs) are carried out to assess the environmental impact of products in each of these stages: 72. extracting and processing raw materials 73. manufacturing and packaging 74. use and operation during its lifetime 75. disposal at the end of its useful life, including transport and distribution at each stage. 76. Use of water, resources, energy sources and production of some wastes can be fairly easily quantified. 77. Allocating numerical values to pollutant effects is less straightforward and requires value judgements, so LCA is not a purely objective process. 78. Selective or abbreviated LCAs can be devised to evaluate a product but these can be misused to reach pre-determined conclusions, e.g. in support of claims for advertising purposes. 79. When comparing the impact on the environment of the stages in the life of a product, use and quote the provided data to justify choices The reduction in use, reuse and recycling of materials by end users reduces the use of limited resources, use of energy sources, waste and environmental impacts. 82. Metals, glass, building materials, clay ceramics and most plastics are produced from limited raw materials. 83. Much of the energy for the processes comes from limited resources. 84. Obtaining raw materials from the Earth, by quarrying and mining, causes environmental impacts. (noise, dust, air pollution, ugly ) Some products, such as glass bottles, can be reused. Glass bottles can be crushed and melted to make different glass products. 87. Other products cannot be reused and so are recycled for a different use. 88. Metals can be recycled by melting and recasting or reforming into different products. 89. The amount of separation required for recycling depends on the material and the properties required of the final product. This can be expensive and time consuming. 90. For example, some scrap steel can be added to iron from a blast furnace to reduce the amount of iron that needs to be extracted from iron ore. 30

31 Glass, Ceramics & Polymers Triple Chemistry Only 3. Most of the glass we use is soda-lime glass, made by heating a mixture of sand, sodium carbonate and limestone. 4. Borosilicate glass, made from sand and boron trioxide, melts at higher temperatures than soda-lime glass Clay ceramics, including pottery and bricks, are made by shaping wet clay and then heating in a furnace. 6. Ceramics are brittle and break easily. 7. Ceramics have very high heat resistance and are often used in situations where high temperatures are required. 8. Ceramics are poor thermal and electrical conductors. This makes them useful insulators e.g. in electricity pylons. 2. Polymers 9. The properties of polymers depend on what monomers they are made from and the conditions under which they are made. 10. For example, low density (LD) and high density (HD) poly(ethene) are produced from ethene. See Organic Chemistry sheet for formation of polymers. 11. High density polythene molecules can pack more tightly together, producing harder more rigid plastic which are often used for bottles. 12. Low density polythene is made from smaller, branched chains which produces a more flexible lower melting point. They are used for plastic food bags. 13. Thermosoftening polymers melt when they are heated because the forces between the chains are weak and easily overcome Thermosetting polymers do not melt when they are heated. They have cross link bonds between the chains which are not easily broken, so the object holds it shape. 15. Most composites are made of two materials, a matrix or binder surrounding and binding together fibres or fragments of the other material, which is called the reinforcement. 16. Concrete is a composite material. The matrix is cement which is reinforced with aggregate (gravel) and metal bars. 17. Carbon fibre, Kevlar and fibre glass are all composite materials. 18. These materials are used in applications such as sports equipment, bulletproof vests and boat building. 19. Composite materials can be designed to specifically suit a particular use by giving a desired combination of properties. 31

32 Haber Process - CHEM ONLY 20. The Haber process is used to manufacture ammonia, which can be used to produce nitrogen-based fertilisers. 21. The raw materials for the Haber process are nitrogen and hydrogen. 22. Nitrogen comes from air and hydrogen from natural gas, methane. 23. The purified gases are passed over a catalyst of iron at a high temperature (about 450 C) and a high pressure (about 200 atmospheres). 24. Some of the hydrogen and nitrogen reacts to form ammonia. 25. On cooling, the ammonia liquefies and is removed. 26. The remaining hydrogen and nitrogen are recycled. 27. The reaction is reversible so some of the ammonia produced breaks down into nitrogen and hydrogen: 28. nitrogen + hydrogen ammonia HT ONLY 29. N2(g) + 3H2(g) 2NH3(g) 30. The equilibrium reaction requires a compromise between pressure and temperature to increase the yield. 31. The forward reaction is exothermic, so high temperatures favour the reverse reaction. Although high temperatures increase rate, they lower yield. 32. There are fewer particles on the right of the equation (4 moles on the left, 2 moles on the right), high pressures will favour the forward reaction, increase yield and rate of reaction. 33. The iron catalyst speeds up the reaction and does not affect the position of equilibrium or the yield. 34. Ammonia can be used to manufacture ammonium salts 35. NH 3 + HNO 3 NH 4NO 3 Ammonium nitrate 36. NH 3 + H 2SO 4 (NH 4) 2SO 4 Ammonium sulfate 37. and nitric acid. 4NH 3(g) + 5O 2(g) 4NO (g) + 6H 2O (g) 2NO (g) + O 2(g) 2NO 2(g) 3NO 2(g) + H 2O (l) 2HNO 3(aq) + NO (g) NPK Fertilisers 38. Compounds of nitrogen, phosphorus and potassium are used as fertilisers to improve agricultural productivity. 39. NPK fertilisers contain formulations of appropriate percentages of various salt compounds of all three NPK elements. 40. Industrial production of NPK fertilisers can be achieved using a variety of raw materials in several integrated processes. From Haber Process + Nitric acid Ammonia Nitric acid Ammonium nitrate Potassium chloride Calcium nitrate Phosphoric Acid Ammonium phosphate NPK + Sulfuric Acid 41. Potassium chloride, potassium sulfate and phosphate rock are obtained by mining 42. Phosphate rock cannot be used directly as a fertilizer, but the purified potassium salts can because they are soluble in water. 43. Phosphate rock is treated with nitric acid or sulfuric acid to produce soluble salts that can be used as fertilisers. 44. The process involves acidifying phosphate rock with nitric acid to produce a mixture of phosphoric acid and calcium nitrate 45. Calcium nitrate is crystallised out by cooling to 0 o C to produce nitrogen fertilizer. 46. The filtrate is composed mainly of phosphoric acid with some nitric acid and traces of calcium nitrate, and this is neutralized with ammonia to produce, the fertilizer, ammonium nitrate. 47. If sulfuric acid is added to phosphate rock, calcium phosphate and calcium sulfate are produced. 48. Ammonium salts can be made in the laboratory by neutralising ammonia solution with nitric or sulfuric acid by titration. NH3(aq) + HNO3 NH4NO3 Phosphate Rock Calcium phosphate Calcium sulfate 2NH3(aq) + H2SO4 (NH4)2SO4 49. Unlike the industrial process, there is no waste, or byproducts, in this reaction, other than water. 50. Water is removed by crystallisation. Evaporating excess water and then cooling slowly. 32