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The structure of atoms C2 chapter 1

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1 The structure of atoms C2 chapter 1
Atoms are made up of protons, neutrons and electrons. Particle Charge Mass Proton +1 1 Neutron Electron -1 Nucleus Electrons are arranged in energy levels (shells). The fourth shell can hold 18 electrons (but you don’t need to go that far). The first shell can hold 2 electrons. All atoms want a full outer shell of electrons, and they will do that by gaining electrons, losing electrons or sharing electrons. The second shell can hold 8 electrons. The third shell can hold 8 electrons. Electron configurations can be written 2,8,8,18.

2 Electrons gained/lost
Ionic bonding A metal atom gives electrons to a non-metal atom, so that both of them have full outer shells. When drawing the electron configuration for an ion, we draw square brackets around it, and put the charge on the outside. Electron configurations can be written Na [2,8]+ and Cl [2,8,8]- Na + Cl - We draw the electrons as different shapes to show which atom they came from (dot and cross diagrams) Charges on ions from different groups Group 1 2 3 4 5 6 7 Outer shell electrons 2/8 Electrons gained/lost Loses 1 Loses Loses 3 N/A Gains 3 Gains 2 Gains 1 Ion charge + 2+ 3+ 3- 2- -

3 Ionic formulae + 2 Mg2+ F- MgF2 - - Mg F F Remember: Swap n’ drop
If the positive and negative ions have different charges, then you will need different numbers of ions to balance out the charge e.g. MgF2. F - Mg + 2 F - Remember: Swap n’ drop Compound ions Swap the ion charge numbers, and drop them down to the bottom right of the element symbol. Get rid of the charge. These ions are on your data sheet, so you don’t need to remember them. If your compound has a compound ion (OH-, NO3-, SO42-, CO32-, NH4+) and there is more than one of them, you need to put it in brackets. E.G Mg(NO3)2, or (NH4)2O, or Al2(SO4)3 Mg2+ F- MgF2

4 simplified dot and cross diagram
Covalent bonding Covalent bonding is a shared pair of electrons. covalent bond Covalent bonding occurs between non metals. Group 4 elements share 4 electrons. Group 5 elements share 3 electrons. Group 6 elements share 2 electrons. Group 7 elements and hydrogen share 1 electron. F F F simplified dot and cross diagram H N H O solid line (one line = 1 pair) F H Cl O C

5 Metallic bonding Metal atoms get rid of their outer shell electrons to get a full outer shell. This means that metals are made of positive ions, surrounded by a ‘sea’ of delocalised (free) electrons, which can move.

6 Properties of ionic compounds
C2 chapter 2 Properties of ionic compounds Ionic compounds have strong attractions between oppositely charged ions. When they are in liquid form or in solution, they can conduct electricity because the ions can move. They can’t conduct electricity when solid because the ions can’t move. Ionic compounds have high melting points because it takes a lot of energy to break the forces between the ions. Ion Opposite charge Giant Lattice

7 Properties of small covalent substances
Small covalent substances are made up of non-metals. They are called molecules. Small covalent substances cannot conduct electricity because there are no free electrons. The covalent bonds (the bits where the atoms share electrons) are strong. The intermolecular forces are weak and require a small amount of energy to break. This is why the melting points of small covalent substances are low. Covalent bond Intermolecular force Share Low melting point

8 Properties of giant covalent structures
The atoms are held in giant lattices by strong covalent bonds. This makes them hard. The covalent bonds are strong and require a lot of energy to break, so they have very high melting points. Diamond Silicon dioxide In diamond, each carbon is bonded to 4 other carbons. In Silicon dioxide, the silicon is bonded to 4 oxygens. Each oxygen is bonded to 2 silicon atoms. Macromolecule Hard High melting point Lattice Strong covalent bonds

9 Properties of graphite
In graphite, each carbon atom is bonded to 3 other carbon atoms. The carbon atoms are arranged in hexagons. This means that there is one extra electron which is free to move between the layers. Since graphite has free electrons, it can conduct electricty. There are weak forces of attraction between the layers in graphite, making it soft and slippery. Layers Free electrons Bonded to 3 carbons Slippery

10 Properties of giant metallic structures
In giant metallic structures, the atoms are held together in closely packed layers, which can slide over each other, making the metal malleable. The metal atoms release their outer shell electrons, meaning that there are delocalised electrons in the structure. The attraction between the positive ions and the negative electrons gives metals a high melting point. The delocalised electrons allow metals to conduct electricity. Metallic bonding Delocalised electrons Conductor High melting point Close packed layers Shape memory alloys are alloys that can return to their original shape when heated.

11 Properties of polymers
Polymers are made up of chains. If the chains are close together, then the polymer is harder and less flexible than if they are far apart (HDPE vs LDPE). Some polymers have cross-links which hold the chains together. These are called thermosetting polymers. These polymers are harder and have higher melting points. Polymers without cross-links are called thermosoftening polymers. They are soft and flexible. Thermosoftening Thermosetting Cross-links Chains Branched chains High density polythene (HDPE) Low density polythene (LDPE)

12 Nanoscience Nanoparticles are used in:
Glass – titanium oxide particles on windows help clean them. Sun screen – titanium oxide particles help block the sun. Cosmetics – nanoparticles are able to travel through the pores in the skin more easily. Risks of nanoparticles Their high surface area means that they could easily catch on fire. Their small size means that they could get around the body and they may have unpredictable effects on peoples’ health. 1-100nm across A few hundred atoms Larger surface area Different properties than the same substance made from larger particles.

13 C 12 6 C2 chapter 3 Particle Charge Mass Proton +1 1 Neutron Electron
No. of protons = atomic number No. of electrons in an atom = atomic number. Number of neutrons = mass number – atomic number 12 Mass number = protons + neutrons C 6 Atomic number= protons Particle Charge Mass Proton +1 1 Neutron Electron -1 Two atoms of the same element are isotopes if they have the same number of protons but a different number of neutrons. Isotopes have different properties, and some are radioactive, but they have the same chemical properties because they have the same number of electrons in the outer shell.

14 Relative atomic mass: The average mass of all the isotopes of an atom compared to 1/12 of a carbon-12 atom. In almost every case, it is the same as the mass number of the most common isotope. Relative formula mass: The total mass of all the elements in a compound. E.g in CO2, the relative formula mass is 12 + (16 x 2) = 44. Mole: A really big number (6.02x1023). One mole of a substance weighs its relative atomic mass in grams.

15 Empirical formula Element name Mass of element (g)
Follow the same steps every time to calculate the empirical formula. Write down the mass of each element (mass can be in grams or a percentage) Divide the mass by the relative atomic mass of the element. Divide numbers by the smallest number to get the ratio of elements. If there is a half in the ratio (for example 1: 2.5), multiply all the numbers by 2. These numbers give the empirical formula. Example A compound has 24 g of carbon and 64 g of oxygen. What is its empirical formula? Element name Mass of element (g) Mass ÷ Relative Atomic Mass Divide by the smaller number Ratio Carbon 24 24 ÷ 12 = 2 2 ÷ 2 = 1 1 Oxygen 64 64 ÷ 16 = 4 4 ÷ 2 = 2 2 The empirical formula of this compound is CO2.

16 Percentage composition of an element in a compound =
(Relative atomic mass of the element ÷ Relative formula mass of the compound) x 100 e.g. the % of magnesium in magnesium oxide (MgO) is (24 ÷ 40) x 100 = 60% Yield and % yield The yield of a chemical reaction is the amount of product made. However, a chemical reaction never produces as much product as it should for several reasons: The reaction may be reversible (so the products may turn back into reactants) Some reactants may react to give unexpected products Some products may be left behind in the apparatus The reactant may not be completely pure Some reactions produce more than one product and it may be difficult to separate the product that we want from the reaction mixture. Percentage yield of a chemical reaction = Mass of product actually made ÷ Mass of product that should be made in theory.

17 + + Balancing equations CH4 + 2O2  CO2 + 2H2O
There are 4 hydrogens here. You multiply the big number by the little number. There are 4 hydrogens here, bonded together. There are 2 molecules of oxygen not bonded together. You can only change the BIG numbers in equations – we cannot change the small numbers or add or take away any reactants or products.

18 If we had 32g of oxygen reacts to form water, what mass of water is produced?
O H  H2O Work out the number of moles of oxygen (O2 has an Mr of 32) 2 x 18 = 36g of water is produced. 32 ÷ 32 = 1 mole of oxygen. Look at the ratio of oxygen to water. Work out the mass of water. Work out the number of moles of water. 1 mole of oxygen produces 2 moles of water. 1 x 2 = 2 2 moles of water are produced.

19 Reversible reactions HLit ⇌ Lit- + H+ CuSO4.5H2O ⇌ CuSO4 + 5H2O
In a reversible reaction, the products turn back into reactants. The symbol for reversible reactions is ⇌. HLit ⇌ Lit- + H+ Litmus is an indicator that turns red in an acid and blue in an alkali. Red Litmus Blue Litmus CuSO4.5H2O ⇌ CuSO4 + 5H2O Hydrated copper Sulphate (blue) Anhydrous copper Sulphate (white) NH4Cl ⇌ NH3 + HCl

20 Chromatography is a way of separating out mixtures of liquids.
A piece of with the ink is dipped in water. As the water creeps up the paper, it carries the inks with it. Different inks travel different distances depending on how well they stick to the paper. You can use chromatography to identify inks by comparing how far the ink has travelled with known inks. Advantages of instrumental methods Highly accurate and sensitive They are quicker They enable very small samples to be analysed. Instrumental methods Other methods can be used that involve instruments. They have advantages and disadvantages. Disadvantages of instrumental methods It is usually very expensive It takes special training to use It gives results that can only be interpreted by comparison with known substances.

21 Gas chromatography The sample is vaporised. The ‘carrier’ gas moves the vapour through the coiled column. The compounds will have different attractions to the particles in the column. Compounds that stay longer have a longer retention time and compounds that get through quickly have a short retention time. Chromatogram Most common substance. Substance with the shortest retention time. Substance with the longest retention time.

22 Mass spectrometry It provides an accurate way of measuring the relative molecular mass of a compound. The sample is turned into + ions and deflected in a magnetic field. The more it deflects, the lighter it is. The far right peak of a mass spectrum is the molecular ion peak. It is the mass of the sample.

23 C2 chapter 4 Rate of reaction = amount of product produced ÷ time
Rate of reaction = amount of reactant used up produced ÷ time This is the amount of Product. If the product is a gas, it can be measured. This is the amount of Reactant. If a gas is produced, you can measure the loss in mass as the reaction goes on.

24 Two particles can only react if: They collide.
Collision theory Two particles can only react if: They collide. They have enough energy to react (this energy is called the activation energy) At the beginning of the reaction, the rate is fast, but then it slows down as there is less reactant as more product is produced, so it is harder for them to collide. If you want to increase the rate of a reaction, you need to do something to increase the number of collisions or do something that increases the number of particles that have the activation energy.

25 The effect of surface area on rate.
Surface area is the amount of a solid that is in contact with the air. If you cut up a solid, you are giving it more surface area. Powders have a greater surface area than lumps. low surface area high surface area 2m Increasing surface area increases the rate of the reaction, because there are more particles exposed to the other reactant, and so there are more successful collisions.

26 The effect of temperature on rate.
Temperature is linked to the amount of energy the particles have. When the temperature is high, the particles move faster and they have more energy. Increasing temperature makes the particles move faster. This means that there are more collisions. Also, since the particles have more energy, then more particles have the activation energy for the reaction.

27 The effect of concentration and pressure on rate.
Concentration is the number of particles in a certain volume of liquid. Pressure is the number of gas particles in a certain volume of gas. The higher the concentration and pressure, the faster the rate. This is because there are more particles in a smaller space, and so there are more successful collisions.

28 The effect of catalyst on rate.
A catalyst works by lowering the activation energy needed to make the reaction happen. It does not give the particles more energy, or increase the number of collisions. A catalyst is a substance that speeds up a reaction without being used up itself. reaction (time) energy (kJ)

29 Catalysts in action Enzymes Enzymes are biological catalysts. They are used in biological washing powders which allow them to break apart food stains. Nanoparticle catalysts Nanoparticles are more effective catalysts because they have a high surface area to volume ratio. Catalytic converters Catalytic converters turn carbon monoxide and nitrogen oxides into nitrogen and carbon dioxide.

30 Exothermic and endothermic reactions
Exothermic reactions give out energy. They feel hot. Endothermic reactions absorb energy from the surroundings. They feel cold. Examples of exothermic reactions: Respiration Combustrion Examples of endothermic reactions: Photosynthesis Thermal decomposition (because the substance is constantly being heated)

31 Energy and reversible reactions
Blue solid White solid CuSO4.5H2O CuSO4 + 5H2O Hydrated copper sulphate Anhydrous copper sulphate The forward reaction is endothermic because it requires constant heating. Anhydrous copper sulphate can be used to test for water because it will turn blue if water is present.

32 Calcium oxide + water → Calcium Hydroxide
Exothermic reaction Uses the reaction: Calcium oxide + water → Calcium Hydroxide Irreversible Endothermic reaction – can be reused by putting in the fridge Exothermic reaction – can be reused by heating in boiling water.

33 C2 5 Acids and alkalis C2 chapter 5.
Acid – a substance that has hydrogen in it. It releases hydrogen ions when it dissolves in water. HCl(g)  H+(aq) + Cl-(aq) Alkali – an alkali is made up of a metal and a hydroxide e.g. potassium hydroxide (KOH), or sodium hydroxide (NaOH). All alkalis are soluble in water. NaOH(s)  Na+(aq) + OH-(aq) State symbols (s) = solid (l) = liquid (g) = gas (aq) = dissolved in water (aqueous solution) Base – A base reacts in the same way as an alkali, but it is not soluble in water. Most metal hydroxides are bases. Examples include iron oxide. Universal indicator is used to measure pH.

34 C2 5 Reactions of acids and salts
Acid + alkali  Salt + water (neutralisation reaction) Acid + base  Salt + water (neutralisation reaction) Acid + metal  Salt + hydrogen Acid + carbonate  Salt + carbon dioxide + water Hydrochloric acid makes chloride salts. Sulphuric acid makes sulphate salts. Nitric acid makes nitrate salts. Precipitate reactions Two soluble salts can be mixed together to to make an insoluble salt and another soluble salt. The insoluble salt is called a precipitate. Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)

35 C2 5 Electrolysis Electrolysis is performed on ionic compounds in order to split them into separate elements. It can only work on ionic compounds that are in liquid state, or in aqueous solution so that the ions can move. Ions cannot move when they are in a solid state. Lead bromide  Lead + bromine PbBr2(l)  Pb(s) + Br2(g) Positive ions head towards the negative electrode and negative ions head towards the positive electrode.

36 Na Br This is a liquid, which means that the ions are not joined together. They are floating about (so always use big numbers for the ions). (l) This ion is negative (you can tell by looking at the ions sheet which you get in the exams) It will lose electrons at the positive electrode to get 0 charge. This ion is positive (you can tell by looking at the ions sheet which you get in exams) It will gain electrons at the negative electrode to get 0 charge. Negative ions always end up in pairs, when they become neutral, so you need 2 negative ions. They make a molecule with 2 atoms joined together (so always use small numbers for the molecule) in them and twice the electrons as the charge on the ion. 2Cl-  Cl2 + 2e- (1- charge, 2 electrons) 2O2-  O2 + 4e- (2- charge, 4 electrons) 2N3-  N2 +6e- (3- charge, 6 electrons) It gains the same number of electrons as the number on its positive charge. Na+ + e-  Na (1+ charge, 1 electron) Ca2+ +2e-  Ca (2+ charge, 2 electrons) Al3+ + 3e-  Al (3+ charge, 3 electrons)

37 C2 5 Electrolysis of solutions
Water contains hydrogen ions and hydroxide ions. If you have a solution of an ionic substance, then you need to decide if the positive ion that goes towards the negative electrode will be hydrogen or the metal or if the negative ion that goes towards the positive electrode will be hydroxide or the other ion. For positive ions, the lower element on the reactivity series will go towards the negative electrode. If hydrogen goes towards the electrode, this is the half equation: 2H+ + 2e-  H2 Hydrogen For negative ions, if the ion is an element ( such as chloride, bromide or iodide), then that will go towards the positive electrode. Otherwise, hydroxide will go towards the positive electrode. The half equation for hydroxide is: 4OH-  2H2O + O2 + 4e-

38 The equation at the anodes is:
2O2-  O2 + 2e- The oxygen produced reacts with the carbon electrodes to make carbon dioxide. The carbon electrodes need to be replaced. Aluminium ore is called bauxite. It contains aluminium oxide. Aluminium oxide is dissolved in molten cryoliote to reduce the melting point. Aluminium is used because it is light and strong. Although it is reactive, it has a layer of aluminium oxide which prevents it from reacting. At the negative cathode, aluminium ions gain electrons to become aluminium atoms. Al3+ + 3e-  Al The overall equation is 2Al2O3(l)  4Al(l) + 3O2(g)

39 C2 5 Electrolysis of brine
Brine is sodium chloride solution – NaCl(aq). It contains hydrogen ions, sodium ions, chloride ions and hydroxide ions. The hydrogen ions go to the electrode to form hydrogen gas. The chloride ions go towards the electrode to form chlorine gas. Sodium hydroxide is left in the solution. Used to make margarine and rocket fuel. Used to make bleach. 2Cl-  Cl2 + 2e- 2H+ + 2e-  H2 Used to make soap.

40 C2 5 Electroplating This occurs when the electrodes are made of the same metal as the positive ions in the solution. The positive electrode will not attract negative ions. What will happen is that the atoms in the positive electrode will turn into ions. This is useful in: Purification of copper. At the anode – Cu  Cu2+ + 2e- At the cathode – Cu2+ + 2e-  Cu Electroplating At the anode – Ag  Ag+ + e- At the cathode – Ag+ + e-  Ag The cathode is a steel spoon that gets coated in silver to make it look silver.


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