GCSE chemistry mindmap

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the electron, John Dalton said the solid sphere made up the different elements. JJ Thompson's experiments showed that showed that an atom must contain small negative charges (discovery of electrons). Ernest Rutherford's alpha particle scattering experiment showed that the mass was concentrated at the centre of the atom. Niels Bohr proposed that electrons orbited in fixed shells; this was supported by experimental observations. A beam of alpha particles are directed at a very thin gold foil Provided the evidence to show the existence of neutrons within the nucleus Show chemical reactions - need reactant(s) and product(s) energy always involves and energy change Atoms of the same element with the same number of protons and different numbers of neutrons Uses words to show reaction reactants → products magnesium + oxygen → magnesium oxide Uses symbols to show reaction reactants → products 2Mg + O₂ → 2Mgo C1 Most of the alpha particles passed right through. A few (+) alpha particles were deflected by the positive nucleus. A tiny number of particles reflected back from the nucleus. Law of conservation of mass states the total mass of products = the total mass of reactants. Does not show what is happening to the atoms or the number of atoms. Shows the number of atoms and molecules in the reaction, these need to be balanced. 35Cl (75%) and 37 Cl (25%) Relative abundance = (% isotope 1 x mass isotope 1) + (% isotope 2 x mass isotope 2) + 100 e.g. (25 x 37) + (75x 35) ÷ 100 = 35.5 Metals Non metals Halogens 1 H Li Be B Na Mg Al Si P S Ge As Se Br K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Fr Ra Ac Rf Db Sg Bh Hs Mt ?? ? Tl Pb Bi Po At Rn With metals Alkali metals With hydrogen Transition metals With aqueous solution of a halide salt To the left of the Periodic table To the right of the Periodic table Consist of molecules made of a pair of atoms Melting and boiling points increase down the group (gas →→liquid → solid) Reactivity decreases down the group Forms a metal halide Forms a hydrogen halide A more reactive halogen will displace the less reactive halogen from the salt Form positive ions. Conductors, high melting and boiling points, ductile, malleable. Form negative ions. Insulators, low melting and boiling points. Halogens 3 4 5 Metal + halogen → metal halide e.g. Sodium + chlorine → sodium chloride Hydrogen + halogen →→ hydrogen halide e.g. Hydrogen + bromine → hydrogen bromide 6 Chlorine + potassium bromide → potassium chloride + bromine Noble gases 0 He Ne Cl Ar Kr Group Have seven electrons in their outer shell. Form -1 ions. 7 CNOF Metals to the left of this line, non metals to the right Increasing atomic mass number. Metals and non metals Increasing proton number means an electron is more easily gained e.g. NaCl metal atom loses outer shell electrons and halogen gains an outer shell electron e.g. Cl₂ + H₂ → 2HCI e.g. Cl₂ +2KBr →2KCI + Br₂ The Periodic table Elements arranged in order of atomic number Group 0 Noble gases AQA Chemistry C1: Atomic Structure & The Periodic Table Development of the Periodic table Unreactive, do not form molecules Boiling points increase down the group Group 1 Before discovery of protons, neutrons and electrons Elements with similar properties are in columns called groups Transition metals (Chemistry only) This is due to having full outer shells of electrons. Increasing atomic number. Mendeleev Elements arranged in order of atomic weight Left gaps for elements that hadn't been discovered yet Alkali metals With oxygen With water With chlorine Compared to group 1 Elements in the same group have the same number of outer shell electrons and elements in the same period (row) have the same number of electron shells. Typical properties . Very reactive with oxygen, water and chlorine Reactivity increases down the group Forms a metal oxide Forms a metal hydroxide and hydrogen Forms a metal chloride Early periodic tables were incomplete, some elements were placed in inappropriate groups if the strict order atomic weights was followed. Elements with properties predicted by Mendeleev were discovered and filled in the gaps. Knowledge of isotopes explained why order based on atomic weights was not always correct. Only have one electron in their outer shell. Form +1 ions. Metal + oxygen → metal oxide • Less reactive • Harder • Denser • Higher melting points Negative outer electron is further away from the positive nucleus so is more easily lost. Metal + water → metal hydroxide + hydrogen Metal + chlorine → metal chloride C1 Many have different ion possibilities with different charges • Used as catalysts Form coloured compounds . e.g. 4Na + O₂ → 2Na₂O e.g. 2Na + 2H₂O → 2NaOH + H₂ e.g. 2Na+ Cl₂ → 2NaCl Cu²+ is blue • Ni²+ is pale green, used in the manufacture of margarine Fe²+ is green, used in the Haber process • Fe³+ is reddish-brown • Mn²+ is pale pink vrage Metallic Covalent Ionic High melting and boiling points Dot and cross diagram Particles are oppositely charged ions Do not conduct electricity when solid Giant structure Particles are atoms that share pairs of electrons Do conduct electricity when molten or dissolved Particles are atoms which share delocalised electrons Electrons are transferred so that all atoms have a noble gas configuration (full outer shells). Na (2, 8, 1) (28.7) Nat. Ch Occurs in compounds formed from metals combined with non metals. Occurs in most non metallic elements and in compounds of non metals. Large amounts of energy needed to break the bonds. Occurs in metallic elements and alloys. lons are held in a fixed position in the lattice and cannot move. Lattice breaks apart and the ions are free to move. (2,8) lonic bonding Metal atoms lose electrons and become positively charged ions Non metals atoms gain electrons to become negatively charged ions -[@]*[@] Na (2,8,8) Properties of ionic compounds Structure Solid, liquid, gas bonds Chemical Melting and freezing happen at melting point, boiling and condensing happen at boiling point. The three states of matter lonic compounds Group 1 metals form +1 ions Group 2 metals form +2 ions AQA Chemistry C2: Bonding, Structure & The Properties of Matter Group 6 non metals form -2 ions Group 7 non metals form -1 ions • Held together by strong electrostatic forces of attraction between oppositely charged ions Forces act in all directions in the lattice SOLD Good conductors of electricity Good conductors of thermal energy The amount of energy needed for a state change depends on the strength of forces between particles in the substance. Properties of metals and alloys LIQUID Metallic bonding Metals as conductors Giant structure of atoms arranged in a regular pattern GAS Alloys Delocalised electrons carry electrical charge through the metal. Energy is transferred by the delocalised electrons. (HT only) Limitations of simple model: There are no forces in the model • All particles are shown as spheres Spheres are solid Delocalised electrons Mixture of two or more elements at least one of which is a metal Pure metal High melting and boiling points Pure metals can be bent and shaped Metalions S g C2 Alloy solid liquid gas This is due to the strong metallic bonds. Atoms are arranged in layers that can slide over each other. Harder than pure metals because atoms of different sizes disrupt the layers so they cannot slide over each other. Electrons in the outer shell of metal atoms are delocalised and free to move through the whole structure. This sharing of electrons leads to strong metallic bonds. Grang Autor Mass appears to increase during a reaction Mass appears to decrease during a reaction Conservation of mass M, Balanced symbol equations Avogadro constant The sum of the relative atomic masses of the atoms in the numbers shown in the formula One of the reactants is a gas One of the products is a gas and has escaped No atoms are lost or made during a chemical reaction Represent chemical reactions and have the same number of atoms of each element on both sides of the equation Chemical amounts are measured in moles (mol) The sum of the M, of the reactants in the quantities shown equals the sum of the M, of the products in the quantities shown. Magnesium + oxygen → magnesium oxide Calcium carbonate → carbon dioxide + calcium oxide Mass changes when a reactant or product is a gas Mass of the products equals the mass of the reactants. Number of moles= mass (a) or mass (a) A₁ M, H₂ + Cl₂ → 2HCI Subscript Normal script Subscript numbers show the number of atoms of the element to its left. Normal script numbers show the number of molecules. One mole of any substance will contain the same number of particles, atoms, molecules or ions. Mass of one mole of a substance in grams = relative formula mass → 2MgO 48g + 32g = 80g 80g = 80g 2Mg + 4.7 0.05 mol 984 equations. and balanced symbol Conservation of mass Moles (HT only) mass (M₁) AQA Chemistry C3: Quantitative Chemistrv 6.02 x 10²3 per mole One mole of H₂O will contain 6.02 x 10²3 molecules One mole of NaCl will contain 6.02 x 10²3 Na+ ions Relative One mole of H₂O = 18g (1 +1 +16) One mole of Mg = 24g How many moles of sulfuric acid molecules are there in 4.7g of sulfuric acid (H₂SO₂)? Give your answer to 1 significant figure. (M, of H₂SO₂) formula equations (HT only) substances in Amounts of The reactant that is completely used up (HT only) Limiting reactants Chemical measurements Using moles to balance equations (HT only) Chemical equations show the number of moles reacting and the number of moles made Limits the amount of product that is made Whenever a measurement is taken, there is always some uncertainty about the result obtained Concentration of solutions Measured in mass per given volume of solution (g/dm³) Can determine whether the mean Less moles of product are made. value falls within the range of uncertainty of the result Conc.= mass (g) volume (dm³) The balancing numbers in a symbol equation can be calculated from the masses of reactants and products Mg + 2HCl →MgCl₂ + H₂ One mole of magnesium reacts with two moles of hydrochloric acid to make one mole of magnesium chloride and one mole of hydrogen C3 1. Calculate the mean 2. Calculate the range of the results Estimate of 3. Example: 1. Mean value is 46.5s 2. Range of results is 44s to 49s = 5s 3. Time taken was 46.5s +2.5s uncertainty in mean would be half the range HT only Greater mass = higher concentration. Greater volume = lower concentration. Convert the masses in grams to amounts in moles and convert the number of moles to simple whole number ratios. If you have a 60g of Mg, what mass of HCI do you need to convert it to MgCl₂? A,: Mg =24 so mass of 1 mole of Mg = 24g M, : HCI (1 + 35.5) so mass of 1 mole of HCI = 36.5g So 60g of Mg is 60/24 = 2.5 moles Balanced symbol equation tells us that for every one mole of Mg, you need two moles of HCI to react with it. So you need 2.5x2 = 5 moles of HCI You will need 5 x 36.5g of HCI= 182.5g A measure of the amount of starting materials that end up as useful products Calculate the atom economy for making hydrogen by reacting zinc with hydrochloric acid: Zn + 2HCl → ZnCl₂ + H₂ M, of H₂ = 1+1=2 M, of Zn + 2HCl = 65 +1 +1 + 35.5 + 35.5 = 138 Atom economy = ²/138 × 100 = ²/138 x 100 = 1.45% This method is unlikely to be chosen as it has a low atom economy. CaCO, ) CaO + CO, M, of CaCO3 = 40 + 12 + (16x3) = 100 M, of CaO = 40 + 16 = 56 Atom economy - Relative formula mass of desired product from equation x 100 Sum of relative formula mass of all reactants from equation 100g of CaCO3 would make 56 g of CaO So 200g would make 112g Yield is the amount of product obtained It is not always possible to obtain the calculated amount of a product HT only: 200g of calcium carbonate is heated. It decomposes to make calcium oxide and carbon dioxide. Calculate the theoretical mass of calcium oxide made. Percentage yield is comparing the amount of product obtained as a percentage of the maximum theoretical amount Atom economy Concentration of a solution is the amount of solute per volume of solution Using concentrations of solutions in mol/dm³ (HT only, chemistry only) AQA Chemistry C3: Quantitative Chemistry % Yield = Mass of product made x 100 Max. theoretical mass Percentage yield 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. Concentration = (mol/dm³) Titration amount (mol) volume (dm³) A piece of sodium metal is heated in chlorine gas. A maximum theoretical mass of 10g for sodium chloride was calculated, but the actual yield was only 8g. Calculate the percentage yield. Percentage yield = 8/10 x 100 = 80% If the volumes of two solutions that react completely are known and the concentrations of one solution is known, the concentration of the other solution can be calculated. Equal amounts of moles or gases occupy the same volume under the same conditions of temperature and pressure High atom economy is important or sustainable development and economic reasons Use of amount of substance in relation to volumes of gases (HT only, chemistry only) What is the concentration of a solution that has 35.0g of solute in 0.5dm³ of solution? 2NaOH(aq) + H₂SO₂(aq) → Na₂SO4(aq) + 2H₂O(l) It takes 12.20cm³ of sulfuric acid to neutralise 24.00cm³ of sodium hydroxide solution, which has a concentration of 0.50mol/dm³. Calculate the concentration of the sulfuric acid in mol/dm³: 0.5 mol/dm³ x (24/1000) dm³ = 0.012 mol of NaOH The equation shows that 2 mol of NaOH reacts with 1 mol of H₂SO4, so the number of moles in 12.20cm³ of sulfuric acid is (0.012/2) = 0.006 mol of sulfuric acid Calculate the concentration of sulfuric acid in mol/dm³ 0.006 mol x (1000/12.2) dm³ -0.49mol/dm³ C3 35/0.5 = 70 g/dm³ The volume of one mole of any gas at room temperature and pressure (20°C and 1 atmospheric pressure) is 24 dm³ What is the volume of 11.6 g of butane (C₂H₁) gas at RTP? M.: (4 x 12) + (10 x 1) = 58 11.6/58 0.20 mol Volume = 0.20 x 24 = 4.8 dm³ Calculate the concentration of sulfuric acid in g/dm³: H₂SO4= (2x1) + 32 + (4x16) = 98g 0.49 x 98g = 48.2g/dm³ No. of moles of gas-vol of gas (dm³) 24dm³ 6g of a hydrocarbon gas had a volume of 4.8 dm³. Calculate its molecular mass. 1 mole = 24 dm³, so 4.8/24 = 0.2 mol M, = 6/0.2 = 30 If 6g = 0.2 mol, 1 mol equals 30 g For displacement reactions Acid name Hydrochloric acid Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons) Sulfuric acid Nitric acid Metals and oxygen Reduction Ionic half equations (HT only) Oxidation Ionic half equations show what happens to each of the reactants during reactions Salt name Chloride Sulfate Acids can be Neutralisation neutralised Nitrate sodium hydroxide + hydrochloric acid → sodium chloride + water calcium carbonate + sulfuric acid → calcium sulfate, + carbon dioxide + water by alkalis and bases The half-equation for copper (II) ions is: Cu²+ + 2e →→ Cu For example: The ionic equation for the reaction between iron and copper (II) ions is: Fe + Cu²+ →→ Fe²+ + Cu Metals react with oxygen to form metal oxides The half-equation for iron (II) is: Fe→→ Fe²+ + 2e- This is when oxygen is removed from a compound during a reaction This is when oxygen is gained by a compound during a reaction Oxidation and reduction in terms of electrons (HT ONLY) An alkali is a soluble base e.g. metal hydroxide. A base is a substance that neutralises an acid e.g. a soluble metal hydroxide or a metal oxide. Neutralisation of acids and salt production Metal oxides Reactions with acids e.g. metal oxides reacting with hydrogen, extracting low reactivity metals magnesium + oxygen → magnesium oxide 2Mg + 0₂ → 2MgO HT ONLY: Reactions between metals and acids are redox reactions as the metal donates electrons to the hydrogen ions. This displaces hydrogen as a gas while the metal ions are left in the solution. Acids react with some metals to produce salts and hydrogen. e.g. metals reacting with oxygen, rusting of iron metal + acid → metal salt + hydrogen Reactions of acids and metals Reactions of acids AQA Chemistry C4: Chemical Changes Reactivity of metals The reactivity series Metals form positive ions when they react Carbon and hydrogen Displacement magnesium + hydrochloric acid → magnesium chloride + hydrogen zinc + sulfuric acid → zinc sulfate + hydrogen Extraction using carbon Metals less reactive than carbon can be extracted from their oxides by reduction. Extraction of metals and reduction Group 1 metals Group 2 metals Zinc, iron and copper The reactivity of a metal is related to its tendency to form positive ions Carbon and hydrogen are non-metals but are included in the reactivity series A more reactive metal can displace a less reactive metal from a compound. For example: zinc oxide + carbon → zinc + carbon dioxide Unreactive metals, such as gold, are found in the Earth as the metal itself. They can be mined from the ground. Reactions with water Reactions get more vigorous as you go down the group Do not react with water Do not react with water The reactivity series arranges metals in order of their reactivity (their tendency to form positive ions). These two non-metals are included in the reactivity series as they can be used to extract some metals from their ores, depending on their reactivity. Silver nitrate + Sodium chloride → Sodium nitrate + Silver chloride C4 Reactions with acid Reactions get more vigorous as you go down the group Observable reactions include fizzing and temperature increases Zinc and iron react slowly with acid. Copper does not react with acid. potassium most reactive K sodium Na calcium magnesium aluminium carbon rbon zinc iron tin lead ieau hydrogen copper silver gold Ca Mg Al Zn Fe Sn 31 Pb H Cu Ag Au platinum least roactive Pt At the negative electrode At the positive electrode Strong acids Weak acids Hydrogen ion concentration Soluble salts The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved. Metal will be produced on the electrode if it is less reactive than hydrogen. Hydrogen will be produced if the metal is more reactive than hydrogen. Oxygen is formed at positive electrode. If you have a halide ion (CI, I, Br) then you will get chlorine, bromine or iodine formed at that electrode. Electrolysis of aqueous solutions Production of soluble salts acidic Completely ionised in aqueous solutions e.g. hydrochloric, nitric and sulfuric acids. Only partially ionised in aqueous solutions e.g. ethanoic acid, citric acid. As the pH decreases by one unit (becoming a stronger acid), the hydrogen ion concentration increases by a factor of 10. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Soluble salts can be made from reacting acids with solid insoluble substances (e.g. metals, metal oxides, hydroxides and carbonates). neutral Add the solid to the acid until no more dissolves. Filter off excess solid and then crystallise to produce solid salts. Process of electrolysis alkaline In neutralisation reactions, hydrogen ions react with hydroxide ions to produce water: H* + OH → H₂O Electrode Acids Where do the ions go? Alkalis (HT ONLY) Strong and weak acids You can use universal indicator or a pH probe to measure the acidity or alkalinity of a solution against the pH scale. Soluble salts Splitting up using electricity Anode Cathode Cations Anions When an ionic compound is melted or dissolved in water, the ions are free to move. These are then able to conduct electricity and are called electrolytes. Passing an electric current though electrolytes causes the ions to move to the electrodes. The positive electrode is called the anode. The negative electrode is called the cathode. Cations are positive ions and they move to the negative cathode. neutralisation The pH scale and Anions are negative ions and they move to the positive anode. Electrolysis AQA Chemistry C4: Chemical Changes Reactions of acids Acids produce hydrogen ions (H+) in aqueous solutions. LeadBons Pb. Titrations are used to work out the precise volumes of acid and alkali solutions that react with each other. Aqueous solutions of alkalis contain hydroxide ions (OH). Titrations (Chemistry only) BromidelonsBr Moltenlead bromide Extracting metals using electrolysis Calculating the chemical quantities in titrations involving concentrations in mol/dm³ and in g/dm³ (HT ONLY): C4 Metals can be extracted from molten compounds using electrolysis. 2NaOH(aq) + H₂SO₂(aq) → Na₂SO₂(aq) + 2H₂O(1) This process is used when the metal is too reactive to be extracted by reduction with carbon. 1. Use the pipette to add 25 cm³ of alkali to a conical flask and add a few drops of indicator. The process is expensive due to large amounts of energy needed to produce the electrical current. Example: aluminium is extracted in this way. 2. Fill the burette with acid and note the starting volume. Slowly add the acid from the burette to the alkali in the conical flask, swirling to mix. It takes 12.20cm³ of sulfuric acid to neutralise 24.00cm³ of sodium hydroxide solution, which has a concentration of 0.50mol/dm³. Calculate the concentration of the sulfuric acid in g/dm³ 0.5 mol/dm³ x (24/1000) dm³ = 0.012 mol of NaOH Higher tier: You can display what is happening at each electrode using half-equations: At the cathode: Pb²+ + 2e → Pb At the anode: 2Br →Br₂ +2e 3. Stop adding the acid when the end-point is reached (the appropriate colour change in the indicator happens). Note the final volume reading. Repeat steps 1 to 3 until you get consistent readings. The equation shows that 2 mol of NaOH reacts with 1 mol of H₂SO₂, so the number of moles in 12.20cm³ of sulfuric acid is (0.012/2) = 0.006 mol of sulfuric acid Calculate the concentration of sulfuric acid in mol/dm³ 0.006 mol x (1000/12.2) dm³ -0.49mol/dm³ Calculate the concentration of sulfuric acid in g/dm³ H₂SO₂ = (2x1) + 32 + (4x16) = 98g 0.49 x 98g = 48.2g/dm³ Reaction profiles Endothermic Overall energy change of a reaction Bond energy calculation Exothermic Breaking bonds in reactants Making bonds in products Show the overall energy change of a reaction Exothermic Energy is taken in from the surroundings so the temperature of the surroundings decreases Endothermic Energy is transferred to the surroundings so the temperature of the surroundings increases Endothermic process Exothermic process Energy released making new bonds is greater than the energy taken in breaking existing bonds. Energy needed to break existing bonds is greater than the energy released making new bonds. Calculate the overall energy change for the forward reaction N₂ + 3H₂2NH3 Bond energies (in kJ/mol): H-H 436, H-N 391, NEN 945 Bond breaking: 945 + (3 x 436) = 945 + 1308 = 2253 kJ/mol Bond making: 6 x 391 2346 kJ/mol Overall energy change = 2253-2346= -93kJ/mol Therefore reaction is exothermic overall. Batteries Simple cell Non-rechargeable cells • Thermal decomposition • Sports injury packs The energy change of reactions (HT only) Rechargeable cells . Combustion • Hand warmers • Neutralisation Types of reaction AQA Chemistry C5: Energy Changes Make a simple cell by connecting two different metals in contact with an electrolyte Cells and batteries (Chemistry only) Consist of two or more cells connected together in series to provide a greater voltage. lonic half equations Stop when one of the reactants has been used up Hydrogen fuel cells Can be recharged because the chemical reactions are reversed when an external electrical current is supplied Increase the voltage by increasing the reactivity difference between the two metals. Alkaline batteries Negative electrode: 2H₂(g) + 40H- (aq) → 4H₂O (1) + 4e- Rechargeable batteries Word equation: hydrogen + oxygen →water Advantages: No pollutants produced Can be a range of sizes Reaction profiles Endothermic Exothermic Fuel cells (Chemistry only) Energy Energy Reactants Time Reactants Activation energy Positive electrode: O₂ (g) + 2H₂O (1) + 4e → 40H- (aq) Time Symbol equation: 2H₂ + O₂ → 2H₂O Disadvantages: Hydrogen is highly flammable Hydrogen is difficult to store Chemical reactions only happen when particles collide with sufficient energy Activation /energy Products Activation energy Products C5 The minimum amount of energy that colliding particles must have in order to react is called the activation energy. Products are at a higher energy level than the reactants. As the reactants form products, energy is transferred from the surroundings to the reaction mixture. The temperature of the surroundings decreases because energy is taken in during the reaction. Products are at a lower energy level than the reactants. When the reactants form products, energy is transferred to the surroundings. The temperature of the surroundings increases because energy is released during the reaction. Aspire for Quantity Mass Volume Rate of reaction REKTANTS TIME Reversible reactions Rate of chemical reaction Representing reversible reactions If a catalyst is used in a reaction, it is The direction Grams (g) cm³ not shown in the word equation. Grams per cm³ (g/cm³) HT: moles per second (mol/s) This can be calculated by measuring the quantity of reactant used or product formed in a given time. Unit A+B Catalyst Enzymes How do they work? Volnoon? heat 10 BO In some chemical reactions, the products can react again to re-form the reactants. A+B= The direction of reversible reactions can be changed by changing conditions: These are biological catalysts. C + D A catalyst changes the rate of a chemical reaction but is not used in the reaction. C + D 60 Catalysts provide a different reaction pathway where reactants do not require as much energy to react when they collide. Reversible reactions cool Energy changes and reversible reactions Rate quantity of reactant used time taken Rate quantity of product formed time taken If one direction of a reversible reaction is exothermic, the opposite direction is endothermic. The same amount of energy is transferred in each case. Calculating rates of reactions -0.42 on's Equilibrium Catalysts Equilibrium in reversible reactions Rate of reaction For example: Hydrated copper sulfate AQA Chemistry C6: The Rate and extent of chemical change Factors affecting rates Reversible reactions and dynamic equilibrium Changing conditions and equilibrium (HT) The relative amounts of reactants and products at equilibrium depend on the conditions of the reaction. endothermic 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 exactly at the same rate. exothermic Temperature Anhydrous copper + Water sulfate Concentration Surface area Factors affecting the rate of reaction Pressure (of gases) Collision theory and activation energy Activation energy Chemical reactions can only occur when reacting Collision theory particles collide with each other with sufficient energy. Le Chatelier's Principles Changing concentration Changing pressure (gaseous reactions) The higher the temperature, the quicker the rate of reaction. Changing temperature The higher the concentration, the quicker the rate of reaction. The larger the surface area of a reactant solid, the quicker the rate of reaction. This is the minimum amount of energy colliding particles in a reaction need in order to react. When gases react, the higher the pressure upon them, the quicker the rate of reaction. C6 Increasing the temperature increases the frequency of collisions and makes the collisions more energetic, therefore increasing the rate of reaction. Increasing the concentration, pressure (gases) and surface area (solids) of reactions increases the frequency of collisions, therefore increasing the rate of reaction. States that when a system experiences a disturbance (change in condition), it will respond to restore a new equilibrium state. If the concentration of a reactant is increased, more products will be formed. If the concentration of a product is decreased, more reactants will react. If the temperature of a system at equilibrium is increased: Exothermic reaction = products decrease - Endothermic reaction products increase For a gaseous system at equilibrium: - Pressure increase = equilibrium position shifts to side of equation with smaller number of molecules. Pressure decrease = equilibrium position shifts to side of equation with larger number of molecules. Crude oil Hydrocarbons General formula for alkanes Alkanes to alkenes Alkenes Properties of alkenes Cracking Steam cracking A finite resource These make up the majority of the compounds in crude oil Catalytic cracking C₂H2n+2 Consisting mainly of plankton that was buried in the mud, crude oil is the remains of ancient biomass. Most of these hydrocarbons are called alkanes. For example: Long chain alkanes are cracked into short chain alkenes. Alkenes are hydrocarbons with a double bond (some are formed during the cracking process). Alkenes are more reactive that alkanes and react with bromine water. Bromine water changes from orange to colourless in the presence of alkenes. C₂H6 C6H₂4 The breaking down of long chain hydrocarbons into smaller chains The heavy fraction is heated until vaporised The heavy fraction is heated until vaporised Crude oil, and alkanes hydrocarbons Display formula for first four alkanes H H-C-H H Methane (CH₂) H- The smaller chains are more useful. Cracking can be done by various methods including catalytic cracking and steam cracking. HHH Η Η Η Propane (C₂H₂) -C-H After vaporisation, the vapour is passed over a hot catalyst forming smaller, more useful hydrocarbons. Cracking and alkenes After vaporisation, the vapour is mixed with steam and heated to a very high temperature forming smaller, more useful hydrocarbons. Carbon compounds as fuels and feedstock HH H-C-C-H HH Ethane (C₂H6) Carbon compounds as fuels and feedstock f AQA Chemistry C7: Organic Chemistrv Alkenes and uses HHHH H-C-C-C-C-H HHH as polymers HH Butane (C₂H₂0) Why do we crack long chains? Hydrocarbon chains Combustion in oil The hydrocarbons in Fractions crude oil can be split into fractions Boiling points Using fractions Fractional distillation and petrochemicals Fractions can be processed to produce fuels and feedstock for petrochemical industry C10H22 Decane →→ pentane + propene + ethane C5H12 + C3H6 + C₂H4 Used to produce polymers. They are also used as the starting materials of many other chemicals, such as alcohol, plastics and detergents. Without cracking, many of the long hydrocarbons would be wasted as there is not much demand for these as for the shorter chains. Properties of hydrocarbons Hydrocarbon chains in crude oil come in lots of different lengths. The boiling point of the chain depends on its length. During fractional distillation, they boil and separate at different temperatures due to this. Each fraction contains molecules with a similar number of carbon atoms in them. The process used to do this is called fractional distillation. We depend on many of these fuels; petrol, diesel and kerosene. During the complete combustion of hydrocarbons, the carbon and hydrogen in the fuels are oxidised, releasing carbon dioxide, water and energy. Boiling point (temperature at which liquid boils) Many useful materials are made by the petrochemical industry; solvents, lubricants and polymers. Viscosity (how easily it flows) Flammability (how easily it burns) Crude Oil The oil is heated in a furnace 20 °C 150°C 400 *C C7 Butane & Propane Complete combustion of methane: Methane + oxygen → carbon dioxide + water + energy CH,(g) + 2O,(g) > COz(E) +2_ H,O (I) Petrol 200 °C R Kerosene 300°C RITT Diesel 370 °C R Fuel Oil Lubricating oil, Parrafin Wax Asphalt As the hydrocarbon chain length increases, boiling point increases. As the hydrocarbon chain length increases, viscosity increases. As the hydrocarbon chain length increases, flammability decreases. for Ethene C,H, H C-C-CH H. H Propene C₂H₂ I Butene C₂H₂ Pentene C₂H₂ Functional group Carboxylic acid reactions Strength (HT only) Polymers Displaying polymers Alkenes Unsaturated General formula for alkenes -COOH For example: CH₂COOH Carboxylic acids react with carbonates, water and alcohols. Carboxylic acids are weak acids Hydrocarbons with a double carbon-carbon bond. Alkenes are unsaturated because they contain two fewer hydrogen atoms than their alkane counterparts. C₂H₂n Methanoic acid, ethanoic acid, propanoic acid and butanoic acid are the first four of the homologous series. Carboxylic acids and carbonates: These acids are neutralised by carbonates Carboxylic acids and water: These acids dissolve water. Carboxylic acids and alcohols: The acids react with alcohols to form esters. Structure and formula of alkenes Carboxylic acids only partially ionise in water. Alkenes are used to make polymers by addition polymerisation. An aqueous solution of a weak acid with have a high pH (but still below 7). In addition polymers, the repeating unit has the same atoms as the monomer. Carboxylic acids polymerisation Functional group Addition polymerisation Alkene reactions Many small molecules join together to form polymers (very large molecules). It can be displayed like this: H ethene HJn repeating unit of poly(ethere) Alkenes are hydrocarbons in the functional group C=C. Reactions of alkenes Reactions of alkenes and alcohols Alkenes react with oxygen in the same way as other hydrocarbons, just with a smoky flame due to incomplete combustion. Condensation polymerisation (HT only) AQA Chemistry C7: Organic Chemistry Synthetic and naturally occurring polymers H Amino acids Amino acids have two functional groups in a molecule. They react by condensation polymerisation to produce peptides. H C Glycine Condensation polymerisation O-H Alcohols DNA and naturally occurring polymers The functional group of an organic compound determined their reactions. Alkenes also react with hydrogen, water and the halogens. The C=C bond allows for the addition of other atoms. Functional group Alcohol reactions Fermentation DNA DNA structure -OH For example: CH-CH₂OH Alcohols react with sodium, air and water. Ethanol is produced from fermentation. H Condensation polymerisation involves monomers with two functional groups H-C-O-HH-C-C-O-H H Methanol H-C- HH C7 II HH Ethanol HHHH ## -C-OH H-C-C-C-C-O-H HHHH Propanol Butanol Methanol, ethanol, propanol and butanol are the first four of the homologous series. Alcohols and sodium: bubbling, hydrogen gas given off and salt formed. Alcohols and air: alcohols burn in air releasing carbon dioxide and water. Alcohols and water: alcohols dissolve in water to form a neutral solution. When sugar solutions are fermented using yeast, aqueous solutions of ethanol are produced. The conditions needed for this process include a moderate temperature (25-50°C), water (from sugar solution) and an absence of oxygen. Deoxyribonucleic acid is a large molecule essential for life. DNA gives the genetic instructions to ensure development and functioning of living organisms and viruses. Most DNA molecules are two polymer chains made from four different monomers, called nucleotides. They are in the double helix formation. Natural Other naturally occurring polymers include proteins, starch polymers and cellulose and are all important for life. When these types of monomers react they join together and usually lose small molecules, such as water. This is why they are called condensation reactions. A for Melting point of a pure substance Formulation How are formulations made? Examples of formulations. Pure substances R, Values Chromatography A pure substances is a single element or compound, not mixed with any other substance. Pure substances Melting point of an impure substance By mixing chemicals that have a particular purpose in careful quantities. A formulation is a mixture that has been designed as a useful product. Fuels, cleaning agents, paints, medicines and fertilisers. Can be used to separate mixtures and help identify substances. The ratio of the distance moved by a compound to the distance moved by solvent. The compounds in a mixture separate into different spots. Formulations Pure substances melt and boil at specific temperatures. Heating graphs can be used to distinguish pure substances from impure. Pure substances Mixture Purity, formulations and chromatography Chromatography Position solvent reaches Mixture separated Solvent Involves a mobile phase (e.g. water or ethanol) and a stationary phase (e.g. chromatography paper). R₂ = distance moved by substance distance moved by solvent This depends on the solvent used. A pure substance will produce a single spot in all solvents whereas an impure substance will produce multiple spots. Gas AQA Chemistry C8: Chemical Analysis Hydrogen Identification of common gases Oxygen Element Chlorine Lithium Sodium Potassium Calcium Copper Carbon dioxide Flame tests (chem only) Test Burning splint Glowing splint Litmus paper (damp) Colour flames Limewater Crimson Yellow Lilac Orange-red Green Identification of ions (CHEMISTRY ONLY) Flame emission spectroscopy Positive result 'Pop' sound. Re-lights the splint. Bleaches the paper white. Metal hydroxides (chem only) Goes cloudy (as a solid calcium carbonate forms). Instrumental methods Sodium hydroxide White precipitates Coloured precipitates Carbonates Halide ions Instrumental methods Flame emission spectroscopy Carbonates, halides and sulfates (chem only) Sulfate ions Is added to solutions to identify metal ions. Aluminium, calcium and magnesium ions form this with sodium hydroxide solution. Copper (II)=blue Iron (II) - green Iron (III) = brown React with dilute acids to form carbon dioxide. When in a solution, they produce precipitates with silver nitrate solution in the presence of nitric acid. C8 When in a solutions they produce a white precipitate with barium chloride solutions in the presence of hydrochloric acid. Methods that rely on machines An instrumental method used to analyse metal ions. Can be used to identify elements and compounds. These methods are accurate, sensitive and rapid. The sample solution is put into a flame and the light that is given out is put through a spectroscope. The output line spectrum, can be analysed to identify the metal ions in the solution. It can also be used to measure concentrations. argon nitrogen Volcano activity 1st Billion years Other gases Reducing carbon dioxide in the atmosphere Combustion of fuels oxygen Gases from burning fuels Particulates Gas Nitrogen Oxygen Argon Carbon dioxide Billions of years ago there was intense volcanic activity Released from volcanic eruptions When the oceans formed, carbon dioxide dissolved into it Atmospheric pollutants from fuels Percentage ~80% This released gases (mainly CO₂) that formed to early atmosphere and water vapour that condensed to form the oceans. Source of atmospheric pollutants. Most fuels may also contain some sulfur. -20% Nitrogen was also released, gradually building up in the atmosphere. Small proportions of ammonia and methane also produced. Carbon dioxide, water vapour, carbon monoxide, sulfur dioxide and oxides of nitrogen. 0.93% This formed carbonate precipitates, forming sediments. This reduced the levels of carbon dioxide in the atmosphere. Solid particles and unburned 0.04% hydrocarbons released when burning fuels. Carbon monoxide Sulfur dioxide and oxides of nitrogen atmosphere gases in the Proportions of Particulates The Earth's early atmosphere Algae and plants Oxygen in the atmosphere Properties and effects of atmospheric pollutants How oxygen increased How carbon dioxide decreased AQA Chemistry C9: Chemistry of the atmosphere Composition and evolution of the atmosphere Common atmospheric pollutants These produced the oxygen that is now in the atmosphere, through photosynthesis. Toxic, colourless and odourless gas. Not easily detected, can kill. First produced by algae 2.7 billion years ago. Cause respiratory problems in humans and acid rain which affects the environment. Cause global dimming and health problems in humans. Reducing carbon dioxide in the atmosphere Carbon footprints The total amount of greenhouse gases emitted over the full life cycle of a product/event. This can be reduced by reducing emissions of carbon dioxide and methane. Formation of sedimentary rocks and fossil fuels CO₂ and methane as greenhouse gases Algae and plants carbon dioxide + water →glucose + oxygen 6CO₂ + 6H₂O → C6H₁2O6 + 60₂ These are made out of the remains of biological matter, formed over millions of years change Global climate Over the next billion years plants evolved to gradually produce more oxygen. This gradually increased to a level that enabled animals to evolve. Effects of climate change Rising sea levels Extreme weather events such as severe storms Change in amount and distribution of rainfall Changes to distribution of wildlife species with some becoming extinct Greenhouse gases Carbon dioxide, water vapour and methane These gradually reduced the carbon dioxide levels in the atmosphere by absorbing it for photosynthesis. The greenhouse effect Carbon dioxide Remains of biological matter falls to the bottom of oceans. Over millions of years layers of sediment settled on top of them and the huge pressures turned them into coal, oil, natural gas and sedimentary rocks. The sedimentary rocks contain carbon dioxide from the biological matter. Methane Climate change C9 Human activities and greenhouse gases Examples of greenhouse gases that maintain temperatures on Earth in order to support life Radiation from the Sun enters the Earth's atmosphere and reflects off of the Earth. Some of this radiation is re-radiated back by the atmosphere to the Earth, warming up the global temperature. Human activities that increase carbon dioxide levels include burning fossil fuels and deforestation. Human activities that increase methane levels include raising livestock (for food) and using landfills (the decay of organic matter released methane). There is evidence to suggest that human activities will cause the Earth's atmospheric temperature to increase and cause climate change. Corrosion Preventing corrosion Sacrificial corrosion NPK fertilisers Fertiliser examples Treatment Nitric acid Sulfuric acid Phosphoric acid The destruction of materials by chemical reactions with substances in the environment Coatings can be added to metals to act as a barrier When a more reactive metal is used to coat a less reactive metal These contain nitrogen, phosphorous and potassium Potassium chloride, potassium sulfate and phosphate rock are obtained by mining Phosphate rock An example of this is iron rusting; iron reacts with oxygen from the air to form iron oxide (rust) water needs to be present for iron to rust. Examples of this are greasing, painting and electroplating. Aluminium has an oxide coating that protects the metal from further corrosion. This means that the coating will react with the air and not the underlying metal. An example of this is zinc used to galvanise iron. Formulations of various salts containing appropriate percentages of the elements. Phosphate rock needs to be treated with an acid to produce a soluble salt which is then used as a fertiliser. Ammonia can be used to manufacture ammonium salts and nitric acid. Products The acid is neutralised with ammonia to produce ammonium phosphate, a NPK fertiliser. Calcium phosphate and calcium sulfate (a single superphosphate). Calcium phosphate (a triple superphosphate). Pressure fertilisers Production and uses of NPK Temperature Corrosion and its prevention Alloys are useful materials Gold Alloys Steels 10 Gold jewellery is usually an alloy with silver, copper and zinc. The carat of the jewellery is a measure of the amount of gold in it e.g. 18 carat is 75% gold, 24 carat is 100% gold. Alloys of iron, carbon and other metals. High carbon steel is strong but brittle. Low carbon steel is softer and easily shaped. The Haber process-conditions and equilibrium The reactants side of the equation has more molecules of gas. This means that if pressure is increased, equilibrium shifts towards the production of ammonia (Le Chatelier's principle). The pressure needs to be as high as possible. A mixture of two elements, one of which must be a metal e.g. Bronze is an alloy of copper and tin and Brass is an alloy of copper and zinc. Steel containing chromium and nickel (stainless) are hard and corrosion resistant. Aluminium alloys are low density. Using materials Ceramics, polymers and composites AQA Chemistry C10: Using Resources The forward reaction is exothermic. Decreasing temperature increases ammonia production at equilibrium. The exothermic reaction that occurs releases energy to surrounding, opposing the temperature decreases. Too low though and collisions would be too infrequent to be financially viable. The Haber process and the use of NPK fertilisers The Haber process Composite materials Ceramic materials Polymers The Haber process Raw materials Catalyst Polymers Thermosetting Thermosoftening A mixture of materials put together for a specific purpose e.g. strength Made from clay Many monomers can make polymers Iron polymers that do not melt when they are heated. Soda-lime glass, made by heating sand, sodium carbonate and limestone. polymers that melt when they are heated. Borosilicate glass, made from sand and boron trioxide, melts at higher temperatures than soda-lime glass. MDF wood (woodchips, shavings, sawdust and resin) Used to manufacture ammonia Concrete (cement, sand and gravel) Made by shaping wet clay and then heating in a furnace, common examples include pottery and bricks. Nitrogen from the air while hydrogen from natural gas These factors affect the properties of the polymer. Low density (LD) polymers and high density (HD) polymers are produced from ethene. These are formed under different conditions. Ammonia is used to produce fertilisers Nitrogen + hydrogen ammonia Both of these gases are purified before being passed over an iron catalyst. This is completed under high temperature (about 450°C) and pressure (about 200 atmospheres). The catalyst speeds up both directions of the reaction, therefore not actually increasing the amount of valuable product.