Chemistry revision pack (2012 14)


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Chemistry revision pack (2012 14)

  1. 1. CHEMISTRY REVISION GUIDE for CIE IGCSE Coordinated Science (2013 and 2014 Syllabus)This revision guide is designed to help you study for the Some very useful websites to help you further yourchemistry part of the IGCSE Coordinated Science course. understanding include: • - whilst not the prettiestThe guide contains everything that the syllabus says you need site this contains a lot of very useful and nicelyyou need to know, and (almost) nothing extra. explained information. • -The material that is in the supplementary part of the course well presented with many clear diagrams, animations(which can be ignored by core candidates) is marked by two and quizzes. Can occasionally lack signs (++) or highlighted in dashed boxes: • - whilst mostly targeted at A-Levels this site contains very detailed information suitable for those looking to deepen their knowledge and hit the highest grades. Finally, remember that revision is not just reading but shouldWhilst this guide is intended to help with your revision, it be an active process and could involve:should not be your only revision. It is intended as a starting •Making notespoint but only a starting point. You should make sure that you •Condensing class notesalso read your text books and use the internet to supplement •Drawing Mind-mapsyour study in conjunction with your syllabus document. •Practicing past exam questions •Making flashcards The golden rule is that what makes you think makes youWhilst this guide does cover the entire syllabus, it just has the learn (and reading on its own does not do this).bare minimum and is not in itself sufficient for thosecandidates aiming for the highest grades. If that is you, youshould make sure you read around a range of sources to get adeeper knowledge and understanding. Happy studying, Mr Field.
  2. 2. C1: THE PARTICULATE NATURE OF MATTER Solids, Liquids andAtom: The smallest particle An atom:of matter Some atoms: GasesMolecule: A small particle Molecules of an element: Molecules of amade from more than one compound:atom bonded togetherElement: A substance A solid element: A gaseous element: SOLIDS LIQUIDS AND GASESmade of only one type of The particles in solids, liquids and gases are held near to each other by forces ofatom attraction. The strength of these forces determines a substance’s melting and boiling points. In a solid, the forces of attraction are strongest, holding the particles tightly inCompound: A substance A solid compound A gaseous compound: position. As the solid is heated, and the particles vibrate faster, these forces aremade from two or more partially overcome allowing the particles to move freely as a liquid – this is calleddifferent elements bonded melting. As the liquid is heated more, the particles gain so much energy that thetogether forces of attraction break completely allowing particles to ‘fly around’ as a gas – this is called boiling. The reverse of the these processes are condensing and freezing. Under specific conditions, some solids can turn straight to gases – a process called subliming (the reverse is called desubliming). PROPERTIESMixture: A substance A mixture of compounds and elements: Solids Liquids Gasesmade from two or more •Have a fixed shape •Take the shape of their •Take the shape of theirelements or compounds •Can’t be compressed container containermixed but not joined •Particles close together •Can’t be compressed •Can be compressed in a regular pattern •Particles close together •Particles widely spaced •Particles vibrate around but disordered in random order a fixed point •Particles move freely •Particles moving very fast.
  3. 3. PAPER CHROMATOGRAPHY best possible separation of spots. C2: EXPERIMENTAL Paper chromatography is a technique that can be used to separate mixtures of dyes or pigments and TECHNIQUES is used to test the purity of a mixture or to see what it contains. Firstly a very strong solution of the mixture is prepared which is used to build up aFILTRATION small intense spot on a piece of absorbent paper.Used to separate solids This is then placed in a jar of solvent (with a lid). Asfrom liquids. The mixture the solvent soaks up the paper, it dissolves theis poured through a filter mixture-spot, causing it to move up the paper withpaper in a funnel. The the solvent. However since compounds haveliquid can pass through different levels of solubility, they move up thethe small holes in the paper at different speeds causing the individualfilter paper (to become components to separate out. The solvent orthe filtrate) and the solid combination of solvents can be changed to get the gets left behind(called the residue). PURITY FRACTIONAL DISTILLATION It is important for chemists to be able to purify the When the liquids being distilled have similarCRYSTALLISATION compounds they make, this is because the boiling points, normal distillation can’t separateCrystallisation is used to separate mixtures of solid dissolved in impurities could be dangerous or just un-useful. them completely but simply gives a purerliquid and relies on the fact that solids are more soluble at This is especially true for chemists making mixture. In this case a fractionating column ishigher temperatures. A solution containing a solid is cooled compounds that are consumed by people such as used. This provides a large surface area fordown until crystals form in the solution, these can then be drugs or food additives since the impurities may be condensation meaning much purer ‘fractions’ arecollected by filtration. toxic which would be very bad news! produced. The most important use of this is separating crude oil into it’s useful components.The related technique of recrystallisation can be used to WHICH TECHNIQUE? DISTILLING AIRseparate a mixture of two soluble solids by taking advantage of You need to be able to select appropriate methods Fractional distillation can bethe difference in their solubility. The mixture is dissolved in the to separate a given mixture. The key to this is look conducted on very cold liquidsmallest possible amount of hot solvent. As the solution for differences in the properties of the air to separate it intocools, the less soluble compound forms crystals that can be components of the mixture such as their nitrogen, oxygen and so on.collected by filtration whilst the more soluble compound stays state, solubility, melting/boiling point and so on. The idea is the same but thedissolved. Then pick the method that best takes advantage of equipment is a little different.DISTILLATION this difference.In distillation a mixture ofliquids is separated using the MELTING/BOILING POINTSdifferences in their boiling No two substances have the exact same meltingpoints. The mixture is heated and boiling points. We can take advantage of thisuntil the liquid with the lowest to test the purity of a compound we have made. Ifboiling point boils, the vapours we know what the melting or boiling point of thethen condense on the cold pure compound should be, we can then measuresurface of the condenser and the melting or boiling point of a sample we havethe pure(er) liquid is collected. produced and the closer it is to the pure value, the more pure it is likely to be.
  4. 4. C3: ATOMS, ELEMENTS AND A NOBLE MATTER H Group VIII: Noble Gases The Noble Gases (He, Ne, Ar Group VII: Halogens COMPOUNDS – Structures and Non-metals Group I: Alkali Metals Group II: Alkali-Earth etc) have full outer shells containing either 2 or 8 Bonding electrons. This is very stable which is why the Noble gasesELECTRON ARRANGEMENT/CONFIGURATION Transition Metals are so unreactive.Electrons are arranged around atoms in specific shells. The Othermost important shell is the outer one as this controls an Other elements tend to Metalsatom’s chemistry. We call the electrons in the outer shell react in such a way as to‘valence electrons’ because they are used for bonding. The achieve a full outer shell bynumber of electrons in the outer shell is the same an gaining or losing electronselement’s group number. until they achieve this Noble Gas configuration. Lanthanides and Actinides (metals)The number of electrons around an atom is given by theatom’s proton number. They are arranged in shells as follows: •1st Shell – Holds two electrons CHEMICAL VS PHYSICAL STRUCTURE OF THE PERIODIC TABLE ISOTOPES •2nd/3rd/4th Shells – Hold 8 electrons CHANGES Elements arranged in order of increasing Isotopes are atoms•Example 1: Carbon. Proton •Example 2: Chlorine. Proton Chemical changes make new proton number. with the same protonnumber is 6 which means number is 17 which means substances whereas physical Periods: The rows in the periodic table. number but differentthere are 6 electrons: 2 in the there are 17 electrons: 2 in changes do not. •For example Li, C and O are all in period 2. nucleon number.1st shell and 4 in the second the 1st shell, 8 in the second Groups: The columns in the PT. and 7 in the 3rd. For example if you melted some •Use roman numbers: I, II, III, IV, V, VI, VII, For example carbon solid sugar to a liquid, it is still VIII (i.e. 1, 2, 3, 4, 5, 6, 7) has two main sugar, just in a different form. If C •Eg. F, Cl, Br, I are all in different periods isotopes – C-12 and you were foolish enough to taste but the same group (Group VII). C-13. Carbon has a Cl it , it would taste sweet before it •Elements in the same group have similar proton number of 6 burnt your tongue! When carbon properties and react in similar ways: the so they both containChecking Your Answer: To check is burnt with oxygen, it makes a halogens all react in the same way with 6 protons and 6you are right, the period gives the new substance: carbon dioxide. sodium to form sodium fluoride (NaF), electrons but C-12number of shells and the group gives the number of electrons The carbon is a black solid sodium chloride (NaCl), sodium bromide has 6 neutrons and C-in the outer shell. For example chlorine is in Period 3 and whereas the carbon dioxide is a (NaBr) and sodium iodide (NaI) 13 has 7.Group VII so it has 3 shells and 7 electrons in the outer shell. colourless gas.Ions: The configuration of ions is the same as for atoms but ATOMIC STRUCTURE what the element have to take electrons away from positive ions and add Eg 1: Boron has 5 Atoms are made of:extra for negative ions. For example O/O2- Li/Li+ In a square on the periodic table protons, 6 neutrons (i.e. Protons: mass = 1, charge = +1 the smaller number, the proton 11 - 5) and 5 electrons Neutrons: mass = 1, charge = 0 number, gives the number of Electrons: mass = 0, charge = -1 protons or electrons and the Eg 2: Phosphorus has 15 O O2- Li Li+ The numbers of each vary from bigger number, the nucleon protons, 16 neutrons (i.e. element to element but it is the number the number of protons 31 - 16) and 15 electrons number of protons which decides and neutrons together.
  5. 5. C3: ATOMS, ELEMENTS AND Group VIII: Noble Gases IONIC BONDING An ionic bond is the attraction between two oppositely charged ions. Cations (positive) are formed Group VII: Halogens COMPOUNDS – Bonding and Non-metals when atoms (usually metals) lose electrons. Anions (negative) are formed when atoms (usually non- metals) gain electrons. An example is table salt: NaCl, made of positive Na + ions and negative Cl- ions. Structure Atoms will lose or gain electrons until they have a complete outer shell: elements in Groups I, II and IIIMOLECULES When a substance melts, it is will lose 1, 2 and 3 electrons respectively to form 1+, 2+ and 3+ ions. Atoms in Groups V, VI and VIIA molecule is a small particle these weak intermolecular gain 3, 2 and 1 electrons to form 3-, 2- and 1- ions. In an ionic compound the number of positive andmade from (usually) a few forces that break NOT the negative and charges must cancel out to neutral.non-metal atoms bonded strong covalent bonds. Example: NaF, sodium in Group I forms a 1+ ion Example: Li2O, lithium in Group I forms a 1+ iontogether. and fluorine in group VII forms a 1- ion so one but oxygen in Group VI forms a 2- ion so two Li+ Molecular compounds haveThe atoms in a molecule are low melting points and are Na+ is needed to balance out one F- are needed to balance out one O2-joined by strong covalent volatile (evaporate easily) duebonds. In a solid each to the weak intermolecularmolecule is held close to its forces, and insulate electricity F- Na+ Li+ O2- Li+neighbour by weak as all electrons are stuck inintermolecular forces. bonds and so unable to move. COVALENT BONDING GIANT IONIC LATTICES A covalent bond forms between two atoms and is the attraction The positive and negative ions in of two atoms to a shared pair of electrons. Small groups of an ionic compound don’t form covalent bonded atoms can join together to form molecules. molecules but form crystals madeGIANT COVALENT LATTICES Graphite: made of carbon of a repeating pattern of positiveA crystal made of a repeating atoms arranged in hexagonal The atoms share enough electrons to complete their outer and negative ions called a giantpattern of atoms joined with sheets with long weak bonds shells. Example: H O*, hydrogen is Example: CO *, carbon is has ionic lattice. Eg sodium chloride: 2 2covalent bonds that repeats between the sheets. This has one valence electron and four valence electrons somillions of times in all means the sheets can easily needs one more to complete needs four more to completedirections. separate making graphite a the 1st shell, oxygen has six its outer shell, oxygen needs good lubricant: valence electrons electrons so two more. Thus each carbonDiamond is made of carbon needs two more. Thus one will react with twoatoms arranged so that each C oxygen will react with two oxygens, sharing twois bonded in a pyramid hydrogens: electrons with each one. Aarrangement to 4 others. This bond involving two sharedmakes it very hard, ideal for use Properties of Ionic Compounds pairs is a double industrial drills: Silicon (IV) oxide (SiO2) has a H O H When you melt or dissolve an structure with each Si ionic compound it conducts O C O electricity because the ions are joined to 4 O and each O joined to 2 Si. It is free to move towards the positive the main ingredient *Nb: In these diagrams only draw the outer shell and use and negative electrodes. When in glass. different shapes/colours to show where electrons have come solid the ions are stuck in position from. You should be able to draw at least: H2O, CH4, Cl2, HCl, H2, and there are no free electrons so N2, O2, CO2, C2H4 they don’t conduct.
  6. 6. CHEMICAL FORMULAS diagram of a molecule C4: STOICHIOMETRY – Formulas tell you the atoms that make up a for example glucose. compound. By counting you can see Formulas and Equations Eg 1. H2O – two H, one O there are 6 carbons, 12 hydrogens and Eg 2. C2H6O – two C, six H, one O 6 oxygens so the Eg 3*. Mg(OH)2 – one Mg, two O, two HSYMBOL EQUATIONS formula is C6H12O6 . Eg 4*. CH2(CH3)2 – three C, eight H•Show the reactants you start with and the products you When writing a formulamake using symbols not words *In this case everything in brackets is doubled you should put any•Must contain an arrow () NOT an equals sign (=) You may be asked to metal atoms first, and then everything else in•Must be balanced – same number of atoms on each side. write a formula given a alphabetical order.•Balancing is done by placing numbers called coefficients infront of the formulas for the compounds/elements. For WORD EQUATIONS IONIC FORMULASexample, ‘O2‘ means there is one oxygen molecule involved in •These tell you the names of the chemicals involved in You can deduce the formula of an ionica reaction but ‘2O2’ would mean there are two. reaction compound if you know the charges on the Example:. CH4(g) + O2(g)  CO2)g) + H2O(g)* •The left hand side shows you what you start with and ions involved. The total positive chargeThis is unbalanced as there are 4 ‘H’ on the left but only 2 ‘H’ is called the reactants must balance out the total negative chargeon the right. This must be corrected by placing a ‘2’ in front of •The right hand side shows you what you make and is so you must look for the lowest commonthe ‘H2O’ so there are now 2 waters: called the products multiple (LCM) of the charges. CH4 (g) + O2(g)  CO2(g) +2H2O(g) •The left and right are connected by an arrow ( not Eg1. Calcium nitrate is made of Ca2+ ions andNow the ‘H’ balances but there 4 ‘O’ on the right and only 2 ‘=‘) which means ‘makes’ or ‘becomes’ NO3- ions. The LCM of 2 and 1 is 2 whichon the left. This must be balanced by placing a ‘2’ in front of •E.g. :When you react a metal with oxygen to make a means you need 1 Ca2+ ion and 2 NO3- ionsthe ‘O2’ so that there are 2 oxygen molecules: metal oxide, the equation might be: so the formula is Ca(NO3)2 CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) Iron + oxygen  iron oxideNow there is 1 ‘C’, 4 ‘H’ and 4 ‘O’ on each side so it balances. •Many fuels burn in oxygen to produce carbon dioxide Eg2. Aluminium oxide is made of Al3+ ions and water for example: and O2- ions. The LCM of 2 and 3 is 6 whichIn ionic equations, we tend to look only at the ions that Methane + oxygen  carbon dioxide + water means you need 2 Al3+ ions and 3 O2- ions soactually change. For example, when iron reacts with copper the formula is Al2O3.sulphate to form iron sulphate and copper the equation is: Fe(s) + Cu2+(aq) + SO42-(aq)  Fe2+(aq) + SO42-(aq) + Cu(s) CHEMICAL MASSESIn this case, the sulphate ion (SO42-) remains unchanged (we The relative atomic mass (Ar) of an element is the Example 1: Water, H2Ocall it a spectator ion) so it can be left out of the equation to mass of one atom relative to 1/12th the mass of C- The Ar for H and O are 1.01 and 16.00 so:give: 12. It is just a number that allows us to compare Mr(H2O) = 2 x 1.01 + 1 x 16.00 = 18.02 Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s) the mass of atoms of different elements. Ar can beThis allows us to see more clearly the actual chemical changes found on the periodic table as the ‘large’ number Example 2: Magnesium Hydroxide, Mg(OH)2taking place. in each square. For example Ar for carbon is 12.01 The Ar for Mg, O and H are 24.31, 16.00 and 1.01: and for iron is 55.85. Ar has no units since it is only Mr(Mg(OH)2) = 1 x 24.31 + 2 x 16.00 + 2 x 1.01Note: You can’t change the little numbers (ie the 2 in H2O ) as a relative number, allowing us to compare things. = 58.33this changes the compound to something completelydifferent. The relative formula mass (Mr) is the combined Ar Example 3: Decane, CH3(CH2)8CH3 of all the elements in the formula for a substance. The Ar for C and H are 12.01 and 1.01*The state symbols (s), (l), (g) and (aq) are used to indicate Mr also has no units for the same reason as above. Mr(decane) = 10 x 12.01 + 22 x 1.01 = 142.34solid, liquid, gas and ‘aqueous solution’ (dissolved in water).
  7. 7. THE MOLE C4: STOICHIOMETRY – The A mole is 6.02x1023 (this number is called Avogadro’s constant) of something. It is chosen so that a mole of something has the same mass in grams (molar mass, Mm) as its formula mass. E.g.: the Mr of Mole Concept water is 18.02 so the Mm of is 18.02g; the Mr of decane is 142.34 so the Mm is 142.34g. Importantly this means that 18.02 g of water and 142.34g decane contains the same number of molecules.EQUATIONS AND MOLE RATIOS THE MOLES AND MASSESEquations can be used to help us calculate the numbers of moles of substances If you know the mass in grams of substance, you can calculate the number of molesinvolved in a reaction. We can see this by studying the following reaction: as follows: 2C2H6 + 7O2  4CO2 + 6H2O Moles = Mass / Molar massQ1: How many moles of CO2 are produced by burning 1.0 mol of C2H6? We say that Eg 1. How many moles is 27.03 g of H2O?C2H6 is our ‘known’ and CO2 is our ‘unknown’ so: Moles (H2O) = Mass / Molar mass = 27.03 / (2 x 1.01 + 16.00) = 1.50 mol Moles CO2 = moles known/knowns in eqn x unknowns in eqn = 1.0 / 2 x 4 = 2.0 mol Eg 2. What is the mass of 0.05 mol of H2O. This time the equation must beQ2: If 0.01 mol of CO2 is produced, how much H2O must also be produced? This rearranged to give:time CO2 is our known and H2O is our unknown so: Mass (H2O) = Moles x molar mass = 0.05 x (2 x 1.01 + 16.00) = 0.901g Moles H2O = moles known/knowns in eqn x unknowns in eqn Note: Mass must be given in grams – you may need to convert from kg: x1000 = 0.01 / 4 x 6 = 0.015 mol*You must make sure your equation is balanced or your mole ratio will be wrong. THE MOLES AND GASES One mole of any gas has a volume of 24.0 dm3 (remember dm3 is the symbol forCALCULATING REACTING QUANTITIES decimetres cubed, aka litres) at room temperature and pressure. So for a gas:Using what we know about calculating moles, we can now answer questions like: IfI have 100g X, how much Y is made? The key is to convert the known to moles 1st. Moles = Volume / 24.0Example: What volume of H2 gas would be produced by reacting 12.15g Eg 1. How many moles of CO2 are present in 60 dm3?magnesium with excess hydrochloric acid? Moles (CO2) = Volume / 24.0 = 60/24.0 = 2.50 mol1. First we need a balanced equation: Eg 2. What is the volume of 0.20 mol of H2 gas?.This time the equation must be Mg + 2HCl  MgCl2 + H2 rearranged to give:2. Then calculate moles of Mg (our known) we start with: Volume (H2) = Moles x 24.0 = 0.20 x 24.0 = 4.80 dm3 Moles Mg = mass/molar mass = 12.15/24.30 = 0.50 mol *The volume must be in dm3 – to convert from cm3 divide by 10003. Next we work out how many moles of H2 ( our unknown) we expect to produce: Moles H2 = moles known/knowns in eqn x unknowns in eqn = 0.50 / 1 x 1 = 0.50 mol THE MOLE AND SOLUTIONS4. Finally we calculate the volume using our equations for a gas: The concentration (strength) of a solution is measured in mol dm-3 (moles per Volume H2 = moles x 24.0 = 0.50 x 24.0 = 12.0 dm3 decimetre cubed). A 1.0 mol dm-3 solution contains 1 mol of substance dissolved in each litre.LIMITING REACTANTS moles of H2O could you make from 3 mol Moles = Concentration x Volume*This is the reactant that will run out first. of H2 and 3 mol of O2. H2: 3/2 = 1.5, O2: Eg 1. How many moles of NaOH are present in 2.5 dm3 of a 1.5 mol dm-3 solution?It is important as this is the one you 3/1 = 3. This means there is enough O2 to Moles (NaOH) = concentration x volume = 1.5 x 2.5 = 3.75 molshould then use for your calculations. do the reaction 3 times but only enough Eg 2. 0.15 mol NaCl is dissolved in 250 cm3 water. What concentration is this? ThisYou calculate it by dividing the number of H2 for 1.5 times so H2 is the limiting time you must rearrange the equation to:moles of reactant by the number of reactant. Thus, moles H2O = 1.5 x (2/2) = Concentration = moles/volume = 0.15/(250/1000)* = 0.60 mol dm-3times they appear in the equation. For 1.5 mol. *The volume must be in dm3 – to convert from cm3 divide by 1000example 2H2 + O2  2H2O. How many
  8. 8. ELECTROLYSIS In the electrolysis of copper chloride (CuCl2) C5: ELECTRICITY AND Electrolysis is a process in which electricity is (right) positive copper ions move to the cathode used to break compounds down into their and form copper metal. Negative chloride ions CHEMISTRY elements. The mixture being electrolysed is more to the anode and form chlorine gas. called an electrolyte and must be liquid (either melted or dissolved) to allow the ions to move. Bubbles of gas Molten Anode formed Cathode Salt Solution (positive (negative Salt Cations (positive ions – remember they are electrode) electrode) Metal, except with reactive metals (K, ’puss-itive’) move to the cathode (the negative Cl Cl Na, Li Ca, Mg) in which case H2 gas is electrode) where they gain electrons, usually Cl Cl Cu2+Cathode Metal produced plus a solution of metal forming a metal (or H2). Layer of metal hydroxide Cl- Cl- Cu formed Anions (negative ions) move to the anode (the Non Metal, except sulphates in which positive electrode) where they lose CuAnode Non-metal case O2 Cl- Cu2+ electrons, usually forming a non-metal (other Cu Cations move Anions move than H2). Cl-ELECTROLYSIS OF COPPER SULPHATE to anode Cl- Cu2+ Cl- to cathodeWhen copper sulphate is electrolysed using carbonelectrodes, you produce O2 gas at the anode and a layer of Cu EXTRACTING ALUMINIUM Aluminium oxide ( the electrolyte) is dissolved inmetal at the cathode. This can be used to electroplate items Aluminium can’t be extracted by reduction of molten ‘cryolite’ and placed in a large carbonby setting them as the cathode. However, when two copper aluminium oxide (Al2O3) using carbon as carbon lined vessel which acts as the cathode. A largeelectrodes are used, what ends up happening is a transfer of is less reactive than aluminium. Instead anode made of carbon is lowered into thecopper from the anode to the cathode, this is used to purify aluminium is produced by electrolysis. electrolyte. The processes that take place are:copper. At the cathode:When copper is made it contains lots of impurities. The copper Aluminium ions gain electronsis purified by electrolysis. A large lump of impure copper is to make liquid aluminiumused as the anode, the electrolyte is copper sulphate solution Al3+ + 3e-  Al(l)and the cathode is made of pure copper. At the anode:At the anode, instead of anions losing electrons, neutral Oxide ions lose electrons tocopper atoms lose electrons to become copper ions . make oxygen gas Cu(s)  Cu2+(aq) + 2e- O2-  ½ O2(g) + 2e-These then move through the electrolyte to the cathode The oxygen reacts with thewhere they become copper atoms again. carbon anode so it has to be Cu2+(aq) + 2e-  Cu(s) replaced regularly The anode loses mass as copper atoms leave ELECTROLYSIS OF BRINE it and the cathode When sodium chloride solution (brine) is gains mass as copper electrolysed , chlorine gas is produced at the atoms join it. The anode and hydrogen gas at the cathode (because impurities sink to the sodium is too reactive). A solution of sodium bottom as a pile of hydroxide is left behind. sludge.
  9. 9. QUANTIFYING ENERGY C6: ENERGY CHANGES IN Using the ideas you learn in physics about specific heat capacity, you may have to calculate the amount of energy released by one mole of a substance. CHEMICAL REACTIONS Example: When 0.250 mol of Metal X reacts fully with 500 cm3 of 2.0 mol dm-3 HCl solution, the temperature increases by 15.4OC. How much energy is released when 1.0 mol X reacts with HCl?EXOTHERMIC REACTIONS First calculate the heat evolved:Exothermic reactions get hotter – the temperature increases. Heat evolved = m.c.ΔT = 500 x 4.2 x 15.4 = 32340 J*The energy given out can be used to keep the reaction going Then calculate heat released per mole:so that once started, they don’t stop until they have run out of Heat per mole = heat evolved / moles = 32340/0.250 = 129360J = 129.4 kJreactants. *ΔT is the temperature rise, m is the mass of the solution in grams which is assumed to equal itsImportant examples of exothermic reactions include: volume in cm3, c is the specific heat capacity of water which is 4.2 J K-1 g-1 •Combustion of fuels •Acid-base neutralisations •Displacement reactions •Respiration in cellsENDOTHERMIC REACTIONSEndothermic reactions reactions get colder – the temperaturedecreases. Generally endothermic reactions need a constantenergy supply to keep them goingImportant examples of exothermic reactions include: •Dissolving of many (but not all) salts •Thermal decompositions Yes this unit really is this small – in fact you don’t even really need the stuff about quantifying energy, •Photosynthesis I just put it in there as it often proves useful!! •Cooking!!!ENERGY CHANGESIn exothermic reactions, chemical energy stored in thereactants gets converted to heat energy. The products haveless chemical energy than the reactants and the difference isthe amount of heat released.In endothermic reactions, heat energy gets converted tochemical energy. The products have more chemical energythan the reactants and the difference between the two is theenergy that has to be supplied to make the reaction go.
  10. 10. MEASURING REACTION RATES If a reaction produces gas, you can easily measure the reaction rate by collecting the gas (either in an C7: CHEMICAL REACTIONS upturned measuring cylinder full of water or a gas syringe) and recording how much has been collected each second.RATES OF REACTIONThe ‘speed’ of a reaction is called its rate and is simply theamount of new product formed every second.For a chemical reaction to happen, the reacting particles needto collide with enough energy. Anything that increases thenumber of collisions or their energy will increase the rate.TemperatureIncreasing temperature increases the rate of a reaction. MEASURING REACTION RATES INVESTIGATING REACTION RATESThis is because particles are moving faster which means more On a graph showing the change in concentration of reactants To investigate a factor influencingcollisions and higher energy collisions. or products, the gradient of the line tells you the reaction rate: reaction rate, you must change it steeper = faster, whilst keeping the others constant.Concentration flat = stopped For example, investigating theIncreasing the concentration of a solution increases the rate of effect of concentration, you coulda reaction. carry out the reaction at 5 different concentrations whilst making sureThis is because it means there are more particles available to the temperature, particle size andreact which leads to more collisions. presence/absence of a catalyst remains the same.Surface Area/Particle size Increasing the total REDOX REACTIONS together. DANGEROUS RATES surface area of Reduction means a substances Factories that produce particles (by using loses oxygen. Oxidation means a Another way to look at this is to flammable powders (for finer powder) substance gains oxygen. For think of oxidation as the loss of example bread flour) have to increases the rate of a example: electrons and reduction as the be careful about sparks since reaction because it 2Fe2O3 + 3C  4Fe + 3CO2 gain of electrons (OILRIG). Eg: in the very fine powder particles means more particles Fe2O3 is reduced because it loses the electrolysis of molten sodium burn with a VERY high reaction at the surface are oxygen to become Fe. C is bromide. At the anode: rate causing explosions. exposed to collisions. oxidised because it gains oxygen 2Br-  Br2 + 2e- to become CO2. C is called a This is an oxidation because the Similar is true underground inCatalysts reducing agent because it causes bromide ions lose electrons. coal mines where gas can buildCatalysts are substances that speed up a reaction without Fe2O3 to get reduced. Reactions At the cathode: up. Gas can be thought of asgetting used up. Whenever a catalyst is present, the rate of like this are called redox Na+ + e-  Na the finest possible powder soreaction increases. Catalysts DO TAKE PART in reactions, they reactions because an oxidation This is a reduction because the they too react explosively fast.just aren’t changed by them. AND a reduction take place sodium ions gain electrons.
  11. 11. THE pH SCALE Litmus indicator C8: ACIDS, BASES AND SALTS Neutral substances have a pH=7 turns red in acids Acids have a pH of less than 7 and blue in alkalis. – Reactions of Acids Alkalis have a pH greater than 7 Universal indicator pH can be measured with colour has many coloursREACTIONS OF ACIDS changing indicators or digital pH (see chart).You need to memorise these reactions, each one shows the metersgeneral word equation then a specific example with symbols. WHAT IS THE SALT?Acids and Metals To work out which salt is formed during neutralisation reactions you need to know the ions formed by• Acid + Metal  Salt + Hydrogen the acid or alkali when it dissolves.•Hydrochloric acid + lithium  lithium chloride + hydrogen• 2HCl(aq) + 2Li(s)  2LiCl(aq) + H2(g) Substance Cation(s) Formed Anion(s) Formed Working out the name is easy, you just combine theAcids and Base (like alkali but not always soluble) Hydrochloric acid, HCl 1 H+ Cl- , chloride name of the cation from• Acid + Base  Salt + water Nitric acid, HNO3 1 H+ NO3- , nitrate the alkali with the anion•Sulphuric acid + sodium hydroxide sodium sulphate + water Sulphuric acid, H2SO4 2H + SO4 2- , sulphate from the acid.• H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) Phosphoric acid, H3PO4 3 H+ PO43- , phosphate For example potassiumAcids and Carbonates Sodium hydroxide, NaOH Na + , sodium 1 OH- sulphate and sulphuric acid Acid + Carbonate  Salt + Water + Carbon Dioxide Potassium hydroxide, KOH K+ , potassium 1 OH - makes potassium sulphate.•Nitric acid + calcium carbonate calcium nitrate + water + Magnesium hydroxide, Mg(OH)2 Mg2+ , magnesium 2 OH- Magnesium hydroxide and. carbon dioxide Ammonium hydroxide, NH4OH NH4+ , ammonium 1 OH- phosphoric acid makes• 2HNO3(aq) + CaCO3(s)  Ca(NO3)2(aq) + H2O(l) + CO2(g) magnesium phosphatePREPARING SALTS Working out the formula of the salt is a little more complicated, the key is to make sure the positiveTo prepare any given salt, you first need and negative charges on the cancel each other out to zero. to work out which acid and alkali to Eg 1. Potassium nitrate Eg 2. Magnesium phosphatereact together (see right). Then react K+ has one plus charge Mg2+ has two plus chargesthem in appropriate quantities so they ACID SO4 2- has two minus charges PO43- has three minus chargesexactly neutralise each other. You caneither calculate the right amounts You need two K+ to balance out one So you need three Mg2+ to balance out(see Unit C4) or find it experimentally SO42- so the formula is K2SO4 two PO43- so the formula is Mg3(PO4)2from a titration. Finally, to write a balanced equation, you need to get the right number of waters, the simplest way is to remember that each ‘H+’ from an acid makes one water.Once you have done this you can use Eg 1. Potassium hydroxide and sulphuric acid Eg 2. Magnesium phosphatethe appropriate techniques to separate As we have seen it makes K2SO4 which requires As we have seen it makes Mg3(PO4)2 whichthe salt from the rest of the solution one H2SO4 and two KOH. Two H2O are made requires two H3PO4 and three Mg(OH)2. Six H2O(See Unit C2). since the one H2SO4 produces two H+ ions are made since each of the two H3PO4 produces three H+ ions. ALKALI + H2SO4 + 2KOH  K2SO4 + 2H2O 2H3PO4 + 3Mg(OH)2  Mg3(PO4)2 + 6H20 INDICATOR
  12. 12. TESTING FOR IONS: Most of these involve forming insoluble precipitates – they go cloudy. C8: ACIDS, BASES AND SALTS Test for... By.... Positive The reaction – Chemical Testing result Chloride ions, Add acidified silver White precipitate Forms insoluble silver chloride:TESTING GASES Cl- nitrate Cl-(aq) + AgNO3(aq)  AgCl(s) + NO3-(aq)Hydrogen: Sulphate ions, Add acidified White precipitate Insoluble barium sulphate formed:•A test tube of hydrogen produces a ‘squeaky pop’ with a SO42- barium nitrate SO42-(aq) + Ba(NO3)2(aq)  BaSO4(s) + 2NO3-(aq)lighted splint Carbonate Add acid and Rapid gas formation The acid reacts with carbonate to makeOxygen: ions, CO32- bubble the gas which turns carbon dioxide gas:•A test tube of oxygen can re-light a glowing splint. formed in limewater cloudy CO32-(s) + 2H+(aq)  CO2(g) + H2O(l) limewater The CO2 reacts with limewater to makeChlorine: insoluble calcium carbonate.•Bleaches the colour from damp litmus paper. Nitrate ions, Boil with NaOH and Vapours turn red The nitrate gets reduced by aluminiumAmmonia: NO3- aluminium foil. litmus paper blue which is a strong reducing agent and forms•Turns damp red litmus paper blue. Test the gas with ammonia.Carbon dioxide: damp red litmus Ammonia is an alkali so can turn the red•Turns limewater cloudy. paper. litmus paper blue. Copper (II), Add sodium Blue precipitate Insoluble copper (II) hydroxide formed: Cu2+ hydroxide followed that dissolves when Cu2+(aq) + 2NaOH(aq)  Cu(OH)2(s) + 2Na+(aq)OXIDES by ammonia ammonia added When ammonia is added a soluble complexThe oxides of most metals are basic (the opposite of acidic). For solution forms so the precipitate dissolves.example sodium oxide (Na2O) forms the alkali sodium 2+ Iron (II), Fe Add sodium Green precipitate Insoluble iron (II) hydroxide formed:hydroxide when it reacts with water. hydroxide followed insoluble in Fe2+(aq) + 2NaOH(aq)  Fe(OH)2(s) + 2Na+(aq) by ammonia ammonia Ammonia does not react with the iron (II)Most oxides of non-metals are acidic. For example, sulphur solution. hydroxide so it does not dissolve.trioxide (SO3) forms sulphuric acid when it dissolves in water. Iron (III), Fe3+ Add sodium Brown precipitate Insoluble iron (III) hydroxide formed: hydroxide followed insoluble in Fe3+(aq) + 3NaOH(aq)  Fe(OH)3(s) + 3Na+(aq)Some oxides form neutral solutions in water for example by ammonia Yet More Tests ammonia Ammonia does not react with the iron (III)carbon monoxide (CO) and nitrogen monoxide (NO). The other solution. to remember these chemical hydroxide so it does not dissolve. You needmain example is dihydrogen monoxide – better known as Zinc, Zn2+ Add sodium tests: White precipitate Insoluble zinc hydroxide formed:water! hydroxide(see Unit C7) •Oxygen followed soluble in both Zn2+(aq) + 2NaOH(aq)  Zn(OH)2(s) + 2Na+(aq) by ammonia lighting a test-tube more Both ammonia and sodium hydroxide react •Hydrogen – ammonia or of H2ACID ENVIRONMENTS solution or more a sodium hydroxide with the zinc hydroxide to form a soluble with a splint gives squeaky popAcid soils grow poor crops so the acidity is reduced by sodium hydroxide. when bubbled •Carbon dioxide – complex.neutralising it with lime (CaO, calcium oxide) through limewater it turns it turn red Ammonium, Add sodium Vapours cloudy. The NH4+ ion is acidic so the NaOH NH4 + hydroxide solution litmus paper blue. neutralises it producing ammonia:Acidic gases from factory chimneys (like sulphur dioxide) can and warm it. NH4+(aq) + -OH(aq)  NH3(g) + H2O(l)dissolve in the water in clouds to form harmful acid rain.
  13. 13. GROUP I (Li, Na, K....) GROUP VII (F, Cl, Br, I...) The metals of Group I (aka the alkali metals) are The elements of Group VII are better known as C9: THE PERIODIC TABLE soft, silvery grey, reactive metals. Down the group the halogens. As we go down the group they get: they get: •Less reactive •Softer •Higher melting point (Cl2 is gas, Br2 is •Lower melting point liquid, I2 is solid)THE PERIODIC TABLE •More reactive •Darker colour (Cl2 is pale green, Br2 isThe periodic table is arranged in order of increasing proton They all react with water as follows: reddy-brown, I2 is dark brown)number – starting at Hydrogen with a proton number of one Metal + water  metal hydroxide + hydrogen They will react with ions of other halogens (halideand working along the rows. •Lithium + water  lithium hydroxide + hydrogen ions) that are below them in the group. For • Li + H2O  LiOH + H2 example:Periods: •Lithium reacts the slowest, Na reacts faster, K Cl2 + 2Br-  2Cl- + Br2The rows in the periodic table are called periods. Going along a reacts faster still and so on. Because Cl is more reactive than Br. However,period, the elements change from metals to non-metals. Br2 + Cl-  no reactionUsually, one or two elements in the period are called As Br is less reactive than Cl.metalloids – these have some properties of a metal and someproperties of a non-metal. TRANSITION ELEMENTS GROUP 0/VIII (He, Ne, Ar, Kr....)Groups: These are the metals in the long middle block of The gases of Group 0 are called the Noble GasesThese are the columns in the periodic table. Elements in the the periodic table. because they are very unreactive. This is becausesame group share similar properties. they have full outer shells of electrons which is Their important properties include: very stable.Groups I and II are always metals. Groups VII and 0/VIII are •High melting/boiling pointsalways non-metals and elements in groups III, IV, V and VI can •High densities They exist as single atoms rather than metals, metalloids or non-metals depending on the period. •Form strongly coloured compounds •(Often) Act as catalysts – both as They are used whenever an inert (unreactive)The Periodic Table and Atomic Structure: elements and when combined in atmosphere is needed. For example:The periodic table can be used to work out the arrangement of compounds •Light Bulbs – Argon surrounds the coiledelectrons: filament as even when white hot, it won’t react. •Period number = number of shells Helium has a very low density (1/7th that of air) •Group number = electrons in outer shell so is used to make airships and blimps float.For example: Chlorine is in Period 3and Group VII so it has 3 electron shells METALS (really belongs in C10 but didn’t quite fit) •High melting/boiling pointand 7 electrons in the outer shell. Most of the known elements are metals. •Sonorous – ‘ring’ when hit Cl •Ductile – can be pulled into wiresElements with only a few electrons in All metals: Conduct electricity, conduct heat, are Many metals react with:their outer shell tend to be metals, shiny •Acids – to form salt and hydrogenwhereas those with many electrons •Oxygen – to form (basic) oxidestend to be non-metals. Most metals are also: •Sulphur – to form sulphides •Malleable – can be beaten into shape When metals bond to non metals they form ionic •Strong bonds.
  14. 14. REACTIVITY OF METALS REACTIVITY SERIES The reactivity of metals can be seen by the way they react with steam or MOST REACTIVE C10: METALS with acid (see Unit C6 for the reactivity series). Potassium, K Reaction with water (see Unit C2 for details of this reaction): Sodium, Na The most reactive metals (K-Ca) react with cold water, fairly reactive metals Calcium, CaEXTRACTING METALS FROM THEIR ORES (Mg-Fe) will only react with steam whereas the least reactive metals (Sn- Magnesium, MgRocks that contain a significant amount of a metal are called Pt) don’t react at all. Aluminium, Al REACTIVITYores. The metals are present as compounds – often oxides or Reaction with dilute acids (see Unit C9 for details) (Carbon, C)sulphides of the metal. For example lead can be extracted from The reaction of metals with acids shows a similar patter with the most Zinc, Znan ore called galena (PbS, lead sulphide). reactive metals (K-Ca) reacting violently, the fairly reactive metals (Mg-Pb) Iron, FeMetals that are less reactive than carbon can be extracted by reacting gradually more slowly and the least reactive metals (Cu-Pt) not Tin, Snusing carbon as a reducing agent (to steal the oxygen/ reacting at all. Lead, Pbsulphur). More reactive metals are extracted by electrolysis. (Hydrogen, H) Displacement Reactions Copper, CuIron is less reactive than carbon so can be reduced by it. This is The reactivity of metals relates to how easily they form ions, more Silver, Agdone in a blast furnace. Study the diagram then read the reactive metals like K form K+ ions much more easily than less reactive Gold, Aufollowing: metals like Cu can form Cu+ ions. A more reactive metal will reduce a less Platinum, Pt•Step 1: Carbon (coke) reacts with oxygen (from the hot air reactive metal:blast) Eg 1. Reaction with aqueous ions LEAST REACTIVE C (s)+ O2(g)  CO2(g) Zinc + Copper sulphate  Zinc sulphate + copper•Step 2: Carbon dioxide reacts with more carbon to make Zn(s) + Cu2+(aq) + SO42-(aq)  Zn2+(aq) + SO42-(aq) + Cu2+(aq) USES OF METALScarbon monoxide This happens because Zn is more reactive than Cu so is able to reduce it. Metals have many uses CO2(g) + C(s)  2CO(g) The Cu2+ gains electrons to become Cu so is reduced by the Zn. including:•Step 3: Carbon monoxide reduces the iron oxide (iron ore) to Eg 2. Reaction with metal oxides •Aluminium – and itsmake molten liquid iron. Iron oxide + aluminium  aluminium oxide + iron alloys used for aircraft Fe2O3(s) + CO(g)  Fe(l) + CO2(g) This happens since Al is more reactive than Fe so is able to reduce it. as they have low densityThe limestone (CaCO3) reacts with impurities such as silicon to and great strength These are called displacement reactions because the more reactive metalform an easy-to-collect waste called slag (calcium •Aluminium – used for takes the place of the less reactive metal.silicate, CaSiO3): CaCO3 +SiO2  CaSiO 3+ CO2 food containers as the ALLOYS Alloys are often harder than the waterproof oxide layer Alloys are ‘mixtures of metals’ metals they are made from. In pure on its surface prevents (although sometimes they can metals atoms are neatly lined up corrosion which could contain a non-metal) that are made meaning they can slip past each taint the food. by mixing molten metals. easily when hit. In alloys there are •Zinc - used to protect Step 3 happens here atoms of different sizes which don’t steel either by coating it Alloys often have very different (galvanising) or as line up neatly so can’t slip past each properties to the metals they are sacrificial protection – Step 2 happens here other so easily making them harder. made from and by varying their i.e. on a ship’s hull – a Element Alloy Step 1 happens here metals can be tailored to have lump of zinc prevents specific desirable properties – this is rust as it is more called metallurgy. reactive so corrodes instead of the steel hull.
  15. 15. CARBON DIOXIDE, CO2 Thermal decomposition of carbonates e.g.: There are many ways to produce CO2 including: CaCO3  CaO + CO2 C11: AIR AND WATER Burning carbon-containing fuels: As a by product of respiration in living cells: CH4 + 2O2  CO2 + 2H2O C6H12O6 + O2  CO2 + H2OWATER, H2O Drinking Water AIR POLLUTION RUSTINGWater is the most useful Water drawn from rivers can Many of man’s activities pollute the air. Pollutants include: Rust (hydrated iron (III) oxide) affectscompound known to man. In contain pollutants such as most structures made of iron (or Carbon monoxide, COthe home it is used for fertilizers, dissolved organic steel) and causes huge damage: •Formed when fuels burn without enough, cleaning and matter, harmful bacteria and Iron + oxygen + water  hydrated •CO prevents the blood from carrying oxygen leading to deathtransporting waste. In industry industrial waste that make it iron (III) oxide by suffocationit is used for cooling hot unfit to drink. At treatment Rust can be prevented by takingmachinery, cleaning and as a plants, two main processes are Sulphur dioxide, SO2 steps making sure either oxygen orsolvent. Water is useful for used to make water safe: •Formed by burning fossil fuels containing sulphur impurities. water can’t reach the iron. The maincleaning as it is able to dissolve •Dissolves in water in clouds to form sulphurous acid which ways to do this involve covering the Filtration – the water is passed falls as acid rain. metal with: paint (bridges and othermany types of ‘dirt’. through a series of increasingly •Acid rain corrodes buildings and damages ecosystems. structures); oil/grease (movingA simple test for water is that fine filters that trap suspended •Irritates the respiratory system when inhaled. machine parts) or another metal suchit is able to turn cobalt chloride particles. Activated carbon is as zinc (galvanising).paper from blue to pink. used to filter out dissolved Nitrogen Oxides, NOx pollutants. •Formed by burning fuels in engines and power stations. FERTILISERS Chlorination – chlorine is •Dissolves in cloud water to form nitric acid thus acid rain. Fertilisers are chemicals applied to added to the water which •Irritates the respiratory system when inhaled. plants to improve their growth and destroys bacteria. increase the amounts of products NITROGEN AND AMMONIA to make an economical such as fruits, nuts, leaves, roots andAIR man’s activities such as Ammonia (NH3) is a smelly gas. amount of ammonia. To speed flowers that they produce for us.Air is a mixture of gases burning fossil fuels and One way to produce it is to it up, the reaction is done at They work by supplying plants withcomprising: deforestation. This is a concern react ammonium (NH4+) salts high temperature (~450OC) the vital elements they need as CO2 is able to absorb the with an alkali (OH-) eg: with an iron oxide catalyst. including Nitrogen - in the form or infrared radiation (heat) NH4Cl + NaOH  NH3 + H2O High pressure (~200 times nitrate (NO3- containing) salts; radiated by the ground when . + NaCl atmospheric pressure) is used phosphorous – in the form of the sun heats it up (the to increase the proportion of phosphate (PO43- containing) salts greenhouse effect). More CO2 Ammonia is vital to produce NH3 formed. and potassium (K+ containing) salts. means more trapped heat the nitrates used in fertilisers leading to global warming. and explosives. It is produced The nitrogen comes from the CATALYTIC CONVERTERS by the Haber process: air and hydrogen comes fromThe ‘other’ is mostly argon Global warming is a major reacting methane (CH4) gas Fit to a car’s exhaust and use awith CO2, water vapour and N2(g) + 3H2(g) D 2NH3(g) platinum or palladium catalyst to problem because with steam.many trace gases. temperatures are rising faster The reaction is reversible convert harmful gases to safer gases: than nature’s ability to adapt – which means much of the Nitrogen and oxygen can be for example nitrogen oxides areAlthough the proportion of separated from air by cooling it reduced back to nitrogen gas and this makes the future of both product turns back tocarbon dioxide is very small farming and of our ecosystems reactants as soon as it is made, to a liquid and using fractional oxygen gas.(~0.04%) it is increasing due to very uncertain. this means it takes a long time distillation.
  16. 16. THE CONTACT PROCESS Sulphuric acid is produced by the Contact Process. C12: SULPHUR This involves are three chemical reactions. First sulphur is burnt in air to produce sulphur dioxide (SO2): S (s) + O2 (g)  SO2 (g)SULPHURIC ACID, H2SO4Sulphuric acid is a very important compound used in many Secondly SO2 is reacted with further oxygen to make sulphur trioxide (SO3):industrial processes including: •Fertiliser production 2SO2 (g) + O2 (g) D SO3 (g) •Oil refining •Paper making This reaction is reversible, so to maximise the amount of SO3 made, they use a high temperature •Steel making (425OC), medium-high pressure (1-2 times atmospheric pressure) and a catalyst (vanadium (V) oxide,It is also the acid found in car batteries. V2O5).Sulphuric acid is a strong acid which when diluted in water Finally, the sulphur trioxide is produced by first dissolving it in sulphuric acid to make oleum (H 2S2O7)produces two protons and a sulphate ion: which then makes more sulphuric acid on the addition of water: H2SO4(l)  2H+(aq) + SO42-(aq) SO3 (g) + H2SO4 (l)  H2S2O7 (l)It exhibits all the reactions typical of an acid as seen by its H2S2O7 (l) + H2O (l)  2H2SO4 (l)reactions with metals, alkalis, metal oxides and carbonates.(see Unit C8 for details). Note: trying to dissolve SO3 directly in water produces a very fine mist of sulphuric with limited uses. This is another tiny unit with very little to learn.
  17. 17. USES OF CALCIUM CARBONATE Powdered calcium carbonate can be added directly to soils to raise their pH (reduce their acidity). C13: CARBONATES We can also make calcium oxide (CaO, aka ‘quicklime’) by heating powdered calcium carbonate to about 1000OC, producing carbon dioxide as a by-product:CALCIUM CARBONATE, CaCO3 CaCO3(s)  CaO(s) + CO2(g)Calcium carbonate is a very common mineral and makes up thebulk of many common rocks including: This is called a thermal decomposition as heat is used to break down or decompose the calcium •Chalk carbonate. Calcium oxide is one of the key ingredients in cement. •Limestone •Marble Another useful product, calcium hydroxide (Ca(OH)2, ‘slaked lime’) is made by adding water to calcium oxide:Whilst solid limestone is often used in construction, powdered CaO + H2O  Ca(OH)2limestone has many industrial uses. Slaked lime has many uses including: •Raising soil pH quickly (when powdered calcium carbonate might take too long) •Neutralising acidic industrial waste •Sewage treatment – it helps small particles of waste to clump together into easily removed lumps. Another mini-unit with very little in it!