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Reaction Rates Following the rate of a reaction. Reaction rate is the speed of a reaction. The faster the reaction the higher the rate. The progress of a chemical reaction can be followed by examining the reaction rate. There are several methods that can be used to follow a reaction rate. Change in volume of any gases produced Change in mass of reaction mixture if gases are given off or taken in. Other changes which can be monitored are: colour, pH, concentration and many more As these changes happen over a period of time, the reaction rate is expressed as Rate = Change in something Change in time. Changes in Volume The volume of the gas can be measured using the following experiments. gas Measuring cylinder 50cm3 of 1 mol/l HCl 0.1 g of Mg water OR Gas syringe 50cm3 of 1 mol/l HCl 0.1 g of Mg 1 As the experiment proceeds the volume of gas is measured and from these results a graph can be drawn. Time (min) 0 2 4 6 8 10 12 14 16 Volume of gas (g) 0 40 70 90 100 105 105 105 105 The volume increases as more gas is formed. Volume of gas produced (cm3) Time (min) The graph is steeper at the beginning as the rate is higher (i.e. the gas is made more quickly) The graph levels off at 10 minutes when no more gas is being made (i.e. the reaction is finished) In this experiment the rate would be calculated by: Rate = Change in volume of gas (cm3) Change in time (min) And would have the units: cm3/min i.e. in the first 2 minutes between 2 and 6 min Rate = 40 - 0 – 0 (cm3)) 2 – 0 (min) Rate =90 – 40 (cm3) 6 – 2 (min) = 40 (cm3)) / 2 (min) = 50 (cm3) / 4 (min) =20 cm3/min = 12.5 cm3/min 2 Changes in mass The mass change can be followed using the experiment below. 50cm3 of 1 mol/l HCl 0.1 g of Mg Balance 127.300g As the experiment proceeds the mass is measured. Time (min) 0 2 4 6 8 10 12 14 16 Mass of flask (g) 127.3 126.5 125.9 125.5 125.3 125.2 125.2 125.2 125.2 The mass falls as the gas is made and escapes. From these results a graph can be drawn of mass of flask against time. Mass of flask (g) Time (min) Rate = Change in mass of flask (g) Change in time (min) 3 And would have the units: g/min The mass of gas formed can also be measured by taking the mass at any time away from the initial mass at time zero. i.e. Before the reaction starts (time zero) the mass At 2 minutes the mass Therefore the mass of gas formed = 127.3 g – 126.5g = =I = 1277.3 126.5g 0.8g We can therefore work out another row for our table: ‘Mass of gas formed’. Time (min) 0 2 4 6 8 10 12 14 16 Mass of flask (g) 127.3 126.5 125.9 125.5 125.3 125.4 125.4 125.4 125.4 Mass of gas formed (g) 0.0 0.8 1.4 1.8 2.0 2.1 2.1 2.1 2.1 This will give a different shaped graph. From these results a graph can be drawn of mass of flask against time. Mass of gas formed (g) Time (min) Rate = Change in mass of gas formed (g) Change in time (min) would have the units: g/min The rate can be calculated from the graph or table Rate = Change in mass of gas formed (g) Change in time (min) i.e. in the first 2 minutes between 2 and 6 min Rate = 0.8 – 0.0 (g) 2 – 0 (min) Rate =1.8 – 0.8 (g) 6 – 2 (min) = 0.8 (g) / 2 (min) = 1.0 (g) / 4 (min) = 0.4 g/min = 0.25 g min The rate is faster at the start of the reaction, which we can tell from the graph because it is steeper. 4 And Atomic structure Everything is made up of tiny particles called atoms. Atoms are mostly empty space made up of smaller sub-atomic particles. At the centre of the atom is the nucleus. This contains two types of particles, called protons and neutrons. Spinning around the nucleus are very fast moving particles called electrons. They move in different levels, called shells. Protons have a positive electrical charge. Neutrons have no electrical charge and are neutral. Moving around the nucleus are very small negatively charged electrons. Nucleus containing the Protons and Neutrons (positive) (positive) (neutral) Electron Cloud containing the Electrons (negative) (negative) Name of particle Proton Neutron Electron Electric Charge Mass (amu) 1 1 1/1840 (almost zero) 1+ 0 1- Position Nucleus Nucleus Electron Cloud The number of protons in the nucleus of an atom is called the Atomic number. Every element has a different atomic number. The atom itself is neutral because the number of protons and electrons are the same. Therefore the positive charge of the nucleus is equal to the sum of the negative charges of the electrons. Atomic Number = Number of protons (also number of electrons) 5 Every atom has two numbers. Atomic Number = number of protons Mass Number = number of protons + neutrons Electrons are not counted when calculating the mass number as they add so little to the atom’s mass. Using a system called Nuclide notation, scientists can show the symbol, mass number and atomic number of an element. Examples: 7 3 Li MassNumber 19 AtomicNumber 9 F Electrons move around the nucleus in orbits. The electron orbits are called shells. Each shell can hold only a certain number of electrons. The first shell The second and third shells The last or outermost shell can hold a maximum of 2 electrons. can hold a maximum of 8 electrons. can hold a maximum of 8 electrons. Electrons further away from the nucleus have more energy then those close to the nucleus. E.g. an atom of the element carbon has an atomic number of 6. Its electrons are arranged so that each electron in the second shell is in a separate quarter of the shell. electron arrangement 2,4 6 An atom of the element oxygen has an atomic number of 8. Its electron arrangement is 2,6 O Electron arrangement and properties Sodium has an atomic number of 11. So its electron arrangement is 2,8,1 Na If we look at Group 1 (the alkali metals) These metals all react violently with water and form alkalis. They all have similar chemical properties. All of the alkali metals one electron in their outer shell. Lithium Sodium Potassium Rubidium Caesium Li Na K Rb Cs 2,1 2,8,1 2,8,8,1 2,8,18,8,1 2,8,18,18,8,1 If we look at Group 8 (0) (the Noble) All the elements in group 0, the noble gases, are inactive and have virtually no chemistry. The outer shell is Full, (8 outer electrons, except He which is full with 2 outer electrons). Helium Neon Argon Krypton Xenon Radon He Ne Ar Kr Xe Rn 2, 2,8 2,8,8, 2,8,18,8 2,8,18,18,8, 2,8,18,32,18,8, Ne It is the number of electrons in the outer shell, which gives atoms their properties. 7 Isotopes All atoms of an element must contain the same number of protons, but they can contain different numbers of neutrons. We can therefore get atoms of the same element with different masses. These are called isotopes. Isotopes: atoms with: Same number of protons but different number of neutrons but different Mass Number OR Same Atomic Number 35 37 17 17 Cl Cl e.g.35Cl has 17 protons and 18 neutrons 37 Cl has 17 protons and 20 neutrons RELATIVE ATOMIC MASS (RAM) (page 4 of the data book) Most elements contain a mixture of isotopes. e.g. 37 35 exist in the proportions 25% 75% Cl Cl The relative atomic mass is the average mass of all the isotopes an element has. It is rarely a whole number but has been rounded off to the nearest 0.5. Relative Atomic Masses in the data book as always closest to the most abundant isotope. The ram for chlorine (on page 4 of the data book) is 35.5. This is because there is more 35Cl than 37Cl It is possible to get equal proportions isotopes. 107 109 e.g. Ag Ag 50% 50% The ram is therefore 108 (i.e. half way between 107 and 109). 8 Covalent Bonding Non-metal elements Group 0 elements, the noble gases exist as single atoms, the monatomic gases. These elements are very unreactive. The rest of the non-metallic elements exist as molecules which are groups of atoms joined together with covalent bonds. Some elements exist as diatomic molecules (a molecule made of two atoms) There are seven non-metal elements in the periodic table which occur as diatomic molecules. They are : hydrogen (H2) oxygen (O2) O nitrogen (N2) fluorine (F2) chlorine (Cl2) bromine (Br2) iodine (I2) Compounds Covalent compounds Covalent compounds contain only non-metal elements eg hydrogen oxide (pure water). Both hydrogen and oxygen are non-metals, so hydrogen oxide is covalent. 9 Covalent Bonding (non-metal/non-metal) The Noble gases are stable elements as they have a full outer electron shell. Other elements react until their atoms obtain a full outer shell and become stable. Non-metal atoms obtain a full outer shell by sharing their outer electrons with other nonmetal atoms. The sharing of outer electrons is called a COVALENT BOND. The substance formed is then called a covalent molecule. For example chlorine atoms have 7 outer electrons by two atoms sharing one electron each they both obtain a full outer shell. Cl Chlorine atom (7 outer electrons) Cl Cl Chlorine atom Cl Chlorine molecule (Cl 2) (7 outer electrons) (each atom now has 8 outer electrons) One shared pair of electrons is called a covalent bond. A full structural formula for a Cl 2 molecule is drawn as Cl Cl It is possible to form double and triple covalent bonds. Double bonds O Oxygen atom (6 outer electrons) O O Oxygen atom O Oxygen molecule (O 2) (6 outer electrons) (each atom now has 8 outer electrons) The full structural formula for the oxygen molecule O 2 is O O 10 Triple Bonds N N Nitrogen atom Nitrogen atom (5 outer electrons) (5 outer electrons) N N Nitrogen molecule (N 2) (each atom now has 8 outer electrons) The full structural formula for the nitrogen molecule N 2 is N This contain triple covalent bond. 11 N Compounds. We can apply the same rules to covalent compounds. Covalent bond Hydrogen chloride H H Cl + Hydrogen atom 1 outer electron Chlorine atom 7 outer electrons Cl Hydrogen chloride molecule (HCl) Each atom in a hydrogen chloride molecule has a full outer shell. The hydrogen atom has 2 outer electrons like the noble gas helium. The chlorine atom has 8 outer electrons like the noble gas argon. Covalent molecules can be made from more than two atoms. Water (H2O) H O Hydrogen has 1 outer electron It needs to obtain 1 more electron Oxygen has 6 outer electrons It needs to obtain 2 more electrons The solution is for the oxygen atom to share its electrons with 2 hydrogen atoms. O H H The full structural formula for H2O is O H 12 H Covalent Bonds Covalent bonds hold the atoms in a molecule together. Take Hydrogen (H2) as an example. Each hydrogen atom has 1 positive proton in the nucleus and 1 negative electron in its electron shell. Two hydrogen atoms share their outer electron to obtain a full outer shell. + + - The negatively charged electrons are attracted to both positively charged nuclei and this holds the atoms together. 13 Formulae The number of electrons which an atom needs is called its valency. This tells us the number of covalent bonds that it will form. Using this we can work out the full structural formula For example: Water Hydrogen has 1 outer electron; it needs 1 more outer electron and so forms 1 bond. Oxygen has 6 outer electrons; it needs 2 more outer electrons and so forms 2 bonds. We therefore draw the full structural formula O H H As all elements in the same group of the periodic table have the same number of outer electrons we can save time by looking at the group Group Group 4 Group 5 Group 6 Group7 Number of outer electrons 4 5 6 7 Valency (Number of electrons needed ) 4 3 2 1 Number of bonds formed 4 3 2 1 For example Phosphorus chloride Element Phosphorus (P) Chlorine (Cl) Group 5 7 Number of outer electrons 5 7 Valency (Number of electrons needed ) 3 1 Number of bonds formed 3 1 Full structural formula Cl P Cl Cl For example carbon sulphide Element Carbon (C) Sulphur (S) Group 4 6 Number of outer electrons 4 6 Valency (Number of electrons needed ) 4 2 Number of bonds formed 4 2 Full structural formula S C 14 S Working out molecular formulae Molecular formulae use symbols to show the number of each atom in a molecule. We can deduce the molecular formula from the structural formula or work it out using the valencies. Example Phosphorus Chloride Symbols Valency Swap Formula P 3 1 P1 Cl 1 3 Cl3 The formula for phosphorus chloride is PCl3 Sometimes you will need to simplify the formulae. Example Carbon sulphide Symbols Valency Swap Simplify Formula C 4 2 1 C1 S 2 4 2 S2 The formula for carbon sulphide is CS2 As Carbon sulphide would work out as C2S4 This is simplified to CS2 to get the correct number of atoms as show by the full structural formula; S C S 15 Compounds which do not obey the rules. Some covalent compounds do not obey the valency rules. The names of these compounds tell us the molecular formulae. eg In the compound carbon dioxide, the prefix di means two. So the compound contains one carbon and two oxygen atoms. The molecular formula is CO2 Prefixes in compound names are: mono di tri tetra penta = = = = = 1 2 3 4 5 Examples of compounds where the name tells us the formula are: carbon monoxide CO nitrogen dioxide NO2 sulphur trioxide SO3 dinitrogen tetraoxide N2O4 (note in these formulae we do not simplify) 16 Shapes of covalent molecules The electrons in the first electron shell are found in a sphere but the other electron shells actually hold the electrons in four orbitals. Thus hydrogen and oxygen which we represent as; H O Should really be represented as and This would give us the molecule water drawn as: We do not have to use orbitals when we show how molecules form, however it is important when working out the actual shape of molecules Each orbital can hold two electrons giving the 8 outer electrons which atoms need to have a stable outer shell. 17 When molecules form, the orbital’s overlap and gives us one of four possible shapes. The shape depends on the number of bonds formed between the atoms. Number of Bonds Shape Name of shape 1 H F Linear 2 Angular 3 Pyramidal 4 Tetrahedral 18 Covalent Molecules and Covalent Networks Covalent substances can be divided into two groups depending on their structure. Covalent molecules So far we have only considered covalent molecules where the molecular formula tells us the exact number of atoms in a molecule. eg water H2O A molecule of water has 1 oxygen atom joined to two hydrogen atoms. Covalent networks A few covalent substances have their atoms joined into a huge network. e.g. diamond (carbon) (C) and silicon dioxide (SiO2) These substances will contain many millions of atoms. Their formulae do not tell us the number of each atom but the ratio of the atoms. e.g. The formula for silicon dioxide is SiO2. This does not mean that the silicon dioxide network contains 1 silicon atom joined to 2 oxygen atoms. It means the ratio of Si : O is 1:2 i.e. there are twice as many oxygen atoms as silicon atoms. 19 Ions Atoms of the same element always have the number of protons but the number of electrons can change when a compound is formed. This gives the atom a charge and we call it an ion. Metal atoms form positive ions Non-metal atoms form negative ions. Positive and negative ions are found together in some compounds. A positive ion is made when an atom loses electrons A negative ion is made when an atom gains electrons. For example: Atom Symbol Na Electrons Lost or Gained 1 electron lost Ion formed Na Al 3 electrons lost Al Cl 1 electron gained Cl O 2 electrons gained O What information do we get from an ion symbol? Here is a symbol for the magnesium atom. Mass number 25 Atomic number 12 Mg This magnesium atom contains :- 12 protons 13 neutrons 12 electrons Here is a symbol for the magnesium ion. Mass number 25 Atomic number 12 magnesium ion contains :- Mg 2+ Electric charge (for ions only) 12 protons 13 neutrons 10 electrons Note: in an ion the number of protons is not equal to the number of electrons. 20 + 3+ 2- Ionic Bonding (metal / non-metal) In order to become stable elements want the stable electron arrangement as their nearest Noble gas (8 Outer electrons, 2 for Li & Be) Elements can also obtain the stable electron arrangement of the noble gases by forming charged particles called IONS. Metal atoms form positive ions by losing electrons. loses 1 e sodium atom sodium ion Na+ Na Electron arrangement Na+ Na+ Na 2,8,1 + e- 2,8, number of positive charges = number of electrons lost = periodic table group number Non-metal atoms form negative ions by gaining electrons. Cl - Cl gains 1 e e.g. chlorine atom Cl Electron arrangement +e chloride ion - Cl- 2,8,7 2,8,8 number of = number of negative charges electrons gained These types of equations are known as an ion-electron equations and can be written for the formation of any simple ion. 21 For example Magnesium atom loses 2 e to form a Magnesium ion Mg 2,8,8,2 Mg 2 ++ 2,8,8, 2e and Sulphur atom gains 2e to form a Sulphide ion S + 2e 2,6 S22,8, Note : The names of the non-metals change to have the ending ‘ –ide ‘ when an ion is formed. Transition metal ions The charge on the metal ion is shown by a Roman number after the name, e.g. an iron (III) ion is Fe3+ a copper (II) ion is Cu2+ Complex ions Some atoms form groups of atoms which have an overall charge. These are listed on page 4 of the data book. e.g. sulphate SO42nitrate NO3 The names of these ions usually end in ‘ ate ‘ or ‘ ite ‘ and contain oxygen. 22 Ionic Compounds (Metal/ Non-metal) We know: metals become stable by Non-metals become stable by losing electrons gaining electrons Ionic compounds are formed when metals transfer their outer electrons to the non-metals. Eg Sodium wants to lose 1 electron and Chlorine wants to gain 1 electron Na atom 2,8,1 Cl atom 2,8,7 1 electron transferred Na+ ion 2,8, Cl- ion 2,8,8, Sodium chloride Na+ Cl- Eg Calcium wants to lose 2 electrons and Chlorine wants to gain 1 electron Ca atom Cl atom Cl atom 2,8,8,2 2,8,7 2,8,7 1 electron transferred 1 electron transferred Ca2+ ion 2,8,8, Cl- ion 2,8,8, Cl- ion 2,8,8, Calcium chloride Ca2+ (Cl-)2 23 Formulae of ionic compounds e.g. aluminium oxide iron (III) bromide a) Al O Fe Br V S S F 3 2 2 2 3 3 3 1 1 1 3 3 Al2O3 copper (II) nitrate Cu NO3 2 1 1 FeBr3 1 2 2 Cu(NO3)2 Note; i) The valency is obtained from the group in the periodic table; ii) iv) 1 2 3 4 5 6 7 8 Valency 1 2 3 4 3 2 1 0 metals give positive ions. non-metals give negative ions. therefore; iii) Group Group 1 2 3 4 5 6 7 8 Valency 1+ 2+ 3+ 4+/- 3- 2- 1- 0 The valency of transition metals is given in roman numerals. If it is not given Assume: silver All other transition metals Complex ions are obtained from the data book page 4. Complex ions are identified by the name not ending in –ide. The exceptions are hydroxide (OH-) and ammonium (NH4+). 24 valency (I) valency (II) STRUCTURE OF COMPOUNDS Ionic compounds form a neat array oppositely charged ions. e.g. sodium chloride Na+ Cl- This structure is called an Ionic lattice (sometimes a crystal lattice). The attractions between oppositely charged ions are very strong and difficult to break.. An Ionic formula does not tell use the actual number of ions present. It tells us the ratio of ions i.e. Na+ Cl- is 1:1 for every one Na+ ion there is one Cl- ion. 25 Acids and Bases The pH Scale A continuous scale from below 0 to above 14 It is a measure of the hydrogen ion (H+) concentration of a solution Dissociation of Water Water exists as an equilibrium between water molecules and hydrogen and hydroxide ions H2O (l) molecules ⇌H + (aq) + OH – (aq) ions In water equilibrium lies over to the left – mainly exists as In water and other neutral solutions (pH 7) the concentration of H+ = the concentration of OH- 26 pH OH- versus H+ Examples concentration of Hydrochloric acid, Nitric acid, Sulphuric acid pH less than 7 Acidic - H+ OH concentration of Neutral pH = 7 Alkaline pH more than 7 - H+ = OH concentration of - + H < OH Water, Sugar solution, salt solution Sodium hydroxide, Potassium hydroxide Bases/Alkalis A base is a substance that neutralises and acid (brings the pH of the acid up to 7 ). An alkali is a base which is soluble in water. Making Acids or Alkalis Acids Acids are made by dissolving soluble non-metal oxides in water. o These make acidic solutions by raising the H+ concentration of water o CO2 (forms carbonic acid, H2CO3 when dissolved in water) o NO2 (forms nitric acid, HNO3 when dissolved in water) Alkalis Alkalis can be made by dissolving soluble metal oxides in water to make metal hydroxides - (see data book). For example: Na2O + H2O 2 NaOH - o This makes an alkaline solution because it raises the OH concentration of the water o Common bases are metal oxides, metal hydroxides or metal carbonates. For example: Na2O Na2OH Na2CO3 Neutral Insoluble oxides will not affect the pH of water. They will not lower or raise the pH of water from 7 because they do not dissolve into the water. 27 Common acids and alkalis used in the laboratory hydrochloric nitric sulphuric ethanoic Common Laboratory Acids HCl HNO3 H2SO4 CH3COOH Common Laboratory Bases sodium hydroxide NaOH Potassium hydroxide KOH calcium hydroxide Ca(OH)2 ammonia NH3 Dilution of acids and alkalis Acids Acids have a pH below 7, which means that they have more H+ than OH-. Diluting an acid, increases the pH towards 7 (concentration of H+ decreases) Alkalis - Alkalis have above 7, which means that they have more OH than H+. Diluting an alkali, decreases the pH towards 7 (concentration of OH- decreases) 28 Neutralisation (Salt Preparation) Neutralisation is the reaction of acids with bases in which salt and water are produced. Base – a substance that reacts with an acid to neutralise it eg metal oxide, metal hydroxide or metal carbonate Salt – a substance in which the hydrogen ion of an acid has been replaced by a metal or ammonium ion During a neutralisation the pH of the resulting solution is pH 7 This can easily been seen by using an indicator. Neutralisation Reaction Equations Acid + Base Salt + Water Acid + metal oxide salt + H2O Acid + metal hydroxide (alkali) salt + H2O Acid + metal carbonate salt + H2O + CO2 29 Naming of Salts The name metal in the base appears in the name of the salt eg copper (II) carbonate produces a copper (II) salt The name of the negative ion in the acid appears in the name of the salt acid Hydrochloric Nitric Sulphuric negative ion Chloride Nitrate Sulphate name of salt Chloride Nitrate Sulphate Examples: nitric acid + iron (II) oxide HNO3 + FeO hydrochloric + copper (II) hydroxide chloride iron (II) nitrate + water Fe(NO3)2 + H2O copper (II) + water + H2O acid HCl + Cu(OH)2 CuCl2 sulphuric carbonate + calcium sulphate calcium + water + carbon acid dioxide H2SO4 + CaCO3 CaSO4 + H2O + CO2 30 Neutralisation reactions In neutralisation reactions the hydrogen ions from the acid are reacting with ions from the base to form water molecules. This can be seen by omitting the spectator ions, (the ions which do not change during the reaction). For example when hydrochloric acid is neutralised by sodium hydroxide we can see what is happening by following 4 steps. 1. Write a balanced equation including state symbols 2HCl(aq) + NaOH(aq) NaCl(aq) + H2O (l) 2. Rewrite including ionic formulae for underlined substances - - H+Cl (aq) + Na+OH - Na+Cl (aq) + H2O(l) 3. Identify the ions which do not change -Spectator Ions - - - H+Cl (aq) + Na+OH Na+Cl (aq) + H2O(l) (sodium and chloride ions are the spectator ions in this equation) 4. Rewrite the equation missing out the spectator ions H+ (aq) + OH - H2O(l) 31 Titrations A titration is a method of analysis that allows you to determine the precise endpoint of a reaction and therefore the precise quantity of reactant in the titration flask. Neutralisation reactions can be carried out very accurately by using a burette and a pipette in a titration. The volumes obtained can be used to calculate the concentrations of the acid or alkali. Pipette Filler Pipette Conical flask 32 FORMULAE When determining formula we use valency which we can obtain from the group in the periodic table. 1 1 1+ 2 2 2+ 3 3 3+ 4 4 4- H Li Na K Rb Cs Fr Be Mg Ca Sr Ba Ra B Al Ti V Cr Mn Fe Co Ni Cu Zn Ga Zr Nb Mo Tc Ru Rh Pd Ag Cd In Hf Ta W Re Os Ir Pt Au Hg Tl C Si Ge Sn Pb Sc Y La Ac 5 6 5 6 3 - 2- N P As Sb Bi O S Se Te Po 7 7 1- 8/0 Group 0 Valency 0 Charge F Cl Br I At He Ne Ar Kr Xe Rn Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr To determine a formula cross over the valences and simplify e.g. symbol nitrogen fluoride N F valency 3 1 symbol carbon C oxide O valency 4 2 Formula N1F3 Formula C2O4 Simplify NF3 Simplify CO2 Ionic Formulae In ionic formulae we show the charges. Note The valency of transition metals is given in roman numerals. Complex ions are obtained from the data book page 8. e.g. a) b) calcium chloride Ca2+ ClCa2+ (Cl-)2 Formulae: Ca2+(Cl-)2 iron (III) bromide Fe3+ BrFe3+ (Br-)3 Fe3+(Br-)3 33 copper (II) nitrate Cu2+ NO3Cu2+ (NO3-)2 Cu2+(NO3-)2 What does a formula tell you? What a formula tells you depends on the structure of a substance. Covalent molecules The formula gives the number of atoms present in the molecule> e.g. H2O contains 2 H atoms and 1 O atom Covalent Network The formula gives the simplest ratio of atoms in the substance. e.g. SiO2 contains Si and O atoms in the ratio 1:2 Ionic Lattice The formula gives the simplest ratio of ions in the substance. e.g. Fe3+(Br-)3 contains Fe3+ and Br- ions in the ratio 1:3 34 Balanced equations In a reactions atoms / ions are not destroyed they just ‘rearrange’ what other atoms / ions they are attached to. In an equation there must be the same number of each atom/ion on each side of the equation. For example when magnesium reacts with oxygen gas to form magnesium oxide. Mg + O2 MgO There must be two oxygen on the right. For this to be possible we must have 2MgO Mg + O2 2MgO For this to be possible we must have started with 2 Mg 2Mg + O2 2MgO + O2 2MgO Thus the balanced equation is: 2Mg We describe this as 2 moles of magnesium reacting with 1 mole of oxygen forming 2 molesa of magnesium oxide. 2Mg 2 mol + O2 1mol 2MgO 2 mol 35 Mole calculations Using the data booklet page 7 we can work out the gram formula masses (GFM) of elements and compounds. GFM is defined as the mass f 1 mole. Calcium nitrate Ca(NO3)2 6 x O = 6 x 16 = 96 2 x N = 2 x 14 = 28 1 x Ca = 1 x 40 = 40 GFM = 144g = 1 mole Once we know the GFM we can calculate masses and number of moles using the triangle m Number of moles n Mass x GFM What is the mass of 0.1 moles of calcium nitrate? m = n x GFM = 0.1 x 144g = 14.4g How many moles of calcium nitrate are there in 72g of calcium nitrate? n = = = m/GFM 72/144 0.5 moles 36 Calculations from balanced equations To do this use the following technique: Here are the steps to carry out: 1. Put a ? over what you need to calculate. 2. Put a over what you know about. 3. Use the balanced equation to write mole ratio. 4. Calculate the mass of each substance 5. Using what you actually have calculate what you want to know What mass of MgO will be formed when 98g of Mg is burned in excess oxygen? √ 2Mg + Mole ratio 2 mol 2(24.5) 49g Actually have 98g ? 2MgO 2 mol 2(24.5 + 16) 81g 98 x 81 49 = 162g O2 37 Concentration of Solutions Concentrations are expressed as the number of moles of a substance dissolved in 1 litre of water. (mol l-1). We use the triangle Once we know the GFM we can calculate masses and number of moles using the triangle n concentration (mol l-1) number of moles c x V volume (in litres) What is the concentration if 0.1 moles of calcium nitrate is dissolved in 250cm3 of water? c = n/V 250 cm3 = 250/1000 l = 0.25 l = 0.1 / 0.25 = 0.4 mol l-1 What mass of Calcium nitrate is needed to make 100cm3 of a 0.1 mol l-1 solution? m = = = = = n x GFM n x 144g n x 144g 0.01 x 144 1.44g We work out GFM from the formula as shown above. To calculate n we use: n = c x v = 0.1 x v v = 100cm3 = 100/1000 l = 0.1 l = 0.1 x 0.1 = 0.01 mol 38 Calculations for titrations To do this use the following technique: Here are the steps to carry out: 1. Put a ? over what you need to calculate. 2. Put a over what you know about. 3. Use the balanced equation to write mole ratio. 4. Calculate the number of moles you actually have of the substance you know about 5. Use the mole ratio to get the number of moles of the substance asked about. 6. Calculate the answer to the question Remember to turn all volume into litre by dividing cm3 by 1000. i.e. 20cm3 = 20/1000 = 0.020 litres Example If 20cm3 of NaOH is neutralised by 20cm3 of 1 mol l-1 H2SO4, what is the concentration of the NaOH? Mole ratio √ H2SO4 1 mol + ? 2NaOH 2 mol Na2SO4+2H2O Actually have c =1 mol l-1 v = 20cm3 = 0.020 litre n = 1 x 0.02 = 0.02 mol v = 20cm3 = 0.020 litre c=? c = n/v 0.04 mol C = n/v = 0.04 / 0.02 = 2 mol l-1 39 Unit 2 Hydrocarbons. Burning hydrocarbons always produces carbon dioxide and water. Results U-tube Colourless liquid condensed Cobalt chloride paper changes from blue to pink. Water present Conclusion Boiling tube Lime water turns cloudy Carbon dioxide present This tells us that hydrocarbons contain Carbon : because carbon dioxide is made Hydrogen: because water is made It does not tell us that they contain oxygen because the oxygen in carbon dioxide and water can have come from the air on burning. Word Equation : methane + oxygen → carbon dioxide + water Formula equation : CH4 + 2O2 → CO2 40 + 2H2O Measuring the Energy The energy produced when a fuel is burned is calculated using the equation: Eh = cmΔT where Eh = energy gained by the water c = specific heat capacity of water = 4.18 kJ0C-1kg-1 m = mass of water (1 litre = 1kg) ΔT= change in temperature of water For example when 200cm3 of water is heated from 20oC to 50oC using the following apparatus. c = specific heat capacity of water = 4.18 kJ0C-1kg-1 m = 200 / 1000 = 0.2kg ΔT= 50 – 20 = 30oC Eh = 4.18 x 0.2 x 30 = 25.08 kJ As the reaction is exothermic we write: Eh = - 25.08kJ 41 Hydrocarbons Hydrocarbons are molecules containing hydrogen and carbon only Carbon Hydrogen- – - always forms 4 bonds always forms 1bond The hydrocarbons are grouped into families called Homologous series. Homologous Series- A group of hydrocarbons with: Similar chemical properties Same General Formula 42 Alkanes General formula CnH 2n+2 Name No. of Carbon atoms Molecular formula Methane 1 CH4 Ethane 2 C2H6 Propane 3 C3H8 Butane 4 C4H10 Pentane 5 C5H12 Hexane 6 C6H14 Heptane 7 C7H16 Octane 8 C8H18 Structural Formula Methane H H CH4 C H H Ethane H Propane H H H C2H6 C C H H H H H C C C H H H H C3H8 H Note: the bonds between the atoms are covalent single bonds 43 Alkenes General Formula CnH2n Ethene H Propene H H H C2H4 C C H H H C C C H C3H6 H H Note: Alkenes contain a double bond (C=C). Alkene Ethene Propene Butene Pentene Hexene No. of Carbon Atoms Molecular Formula 2 C2H4 3 C3H6 4 C4H8 5 C5H10 6 C6H12 44 Cyclo alkanes General Formula CnH 2n Cyclopropane H H C H C C H H Cyclobutane C3H6 H H H C4H8 H C C H H C C H H H Alkane Number of Carbon atoms Molecular formula Cyclopropane Cyclobutane Cyclopentane Cyclohexane 3 4 5 6 C3H6 C4H8 C5H10 C6H12 45 Saturated or unsaturated hydrocarbons. Alkanes and cycloalkanes are saturated because all the (C-C) bonds are single bonds. Alkenes are unsaturated because they contain a (C=C) double bond.. The double bond is responsible for the quick reaction with bromine water. Bromine can be used to tell the difference between saturated and unsaturated hydrocarbons. Addition Reactions The bromine adds to the alkene across the double bond. The reaction involves the breaking of the carbon-carbon double bond to form a carbon-carbon single bond. C3H6 H + Br2 C3H6Br2 H C C H H H + Br Br H C H C Br Br Changing Alkenes to Alkane Alkenes can react with Hydrogen under certain conditions to form Alkanes. Ethene + Hydrogen Ethane C2H4 + H2 C2H6 Propene + Hydrogen Propane C3H6 + H2 C3H8 46 H Isomers Isomers: same molecular formula but different structural formula. Butane – formula is C4H10 and there are two different structures. H H H H H C C C C H H H H H This is a straight chain molecule. This is a branched chain molecule. For the molecule C4H8 H H H C C H H C C H H H Cyclobutane H H H H C C C C H H H Butene 47 H Naming Alkanes Identify the longest chain. Number the carbon atoms to place the branch on the lowest number possible. Identify the position of the chains. Identify the number of each branch. 2,3,4-trimethylpentane Naming Alkenes This is similar to alkanes but th position of the double bond must be as low as possible . 3 – methybut – 1 – ene 48 Alkanols (alcohols) General formula CnH2n+1OH Alkanols contain the functional group –OH (hydroxyl) Name No. of Carbon atoms Molecular formula Methanol 1 CH3OH Ethanol 2 C2H5OH Propanol 3 C3H7OH Butanol 4 C4H9OH Pentanol 5 C5H11OH Hexanol 6 C6H13OH Heptanol 7 C7H15OH Octanol 8 C8H17OH Structural Formula Methanol H H CH3OH C OH H Ethanol H Propanol H H H C C H H H H C2H5OH OH (CH3CH2OH) H C C C H H H C3H7OH OH (CH3CH2CH2OH) 49 Naming alkanols Alkanols are named like the alkenes but here the number of the hydroxyl group (-OH) must be as low as possible. Butan-1-ol Butan-2-ol 50 Uses of Alkanols (alcohols) Alcoholic Drinks. Ethanol is found in alcoholic drinks. Fuels Alcohols burn very easily with a clean flame and so they can be used as a fuel. The products of combustion are carbon dioxide and water. C2H5OH + 3O2 2CO2 + 3H2O Ethanol can be used as a petrol replacement. Methanol is used as the fuel for drag racers. CH3OH + 1½ O2 CO2 + 2H2O Solvents Alkanols are good solvents for many types of substances. The carbon chain helps alkanols dissolve in covalent substances The hydroxyl group (-OH) helps alkanols dissolve in water. Ethanol is used as a solvent in perfumes, aftershave and mouthwash as it is good at dissolving and evaporates easily as well. 51 Alkanoic acids (carboxylic acids) General formula CnH2n+1COOH Alkanoic acids contain the functional group –COOH (carboxyl) O C OH Name No. of Carbon atoms Molecular formula Methanoic 1 HCOOH Ethanoic 2 CH3COOH Propanoic 3 C2H5COOH Butanoic 4 C3H7COOH Pentanoic 5 C4H9COOH Hexanoic 6 C5H11COOH Heptanoic 7 C6H13COOH Octanoic 8 C7H15COOH Be careful when writing the molecular formula or naming alkanoic acids. The stem will have 1 carbon less than you expect as it also contains the carboxyl group COOH. i.e. C4H9COOH is pentanoic acid as there are 5 carbon atoms in this alkanoic acid. 52 Structural Formula Methanoic acid O H Ethaoic acid H HCOOH C OH H O CH3COOH C C OH H H O C C C H H H Propanoic acid H C2H5COOH OH (CH3CH2COOH) Uses of Alkanoic acids (carboxylic acids) food preservatives (i.e. vinegar is a solution of ethanoic acid used for pickled onions) cleaning products and descaling agents (i.e. vinegar in solutions for washing windows) manufacture of esters and plastics 53 Esters Esters are made by reacting alkanols with alkanoic acids. Uses of Esters Food flavourings Esters are added to foods as their smell influences the flavour, for example pineapple cubes contain methyl butanoate Fragrances. The smells of specific esters allows them to be used in things such as perfumes and air fresheners. Solvents. Many esters are sold to remove stains from clothes and house furnishings. Materials Polyesters are esters used in clothing and fabrics such as nylon. 54 POLYMERS Polymers are macromolecules (big molecules). Monomers are the small molecules that combine to form a polymer. A polymer is the large molecule formed by combining many monomers Polymers are very large molecules containing many repeating units. Plastics are made by a process called polymerisation of which there are two types: Addition polymerisation Condensation polymerisation 55 Addition polymerisation. Addition polymers are made from alkenes The names of polymers come from the name of the monomer Propene makes poly(propene) Chloroethene makes poly(chloroethene) Ethene makes poly(ethene) The monomer for poly(ethene) is ethene. This is an unsaturated molecule containing a carbon to carbon double bond, which makes the ethene molecule reactive. The double bond splits open leaving free bonds to join onto other molecules forming a long chain. H H C C H H Ethene H H H H H H H H C C C C C C C C H H H H H H H H Poly(ethene) Repeating unit H Butene can also be used to make a polymer, called poly(butene.) H H H C C C C H H H To make things easier the molecule is drawn in the shape of a letter H 56 H CH3 H CH3 H CH3 H CH3 H CH3 H C C C C C C C C H CH3 H CH3 H CH3 H Butene Repeating unit C CH3 H C CH3 Poly(butene) Since addition polymers are formed by opening a C=C and joining to give a chain they have nothing but carbon atoms in the main chain. 57 Condensation polymerisation. In condensation polymerisation many small molecules combine to form a polymer. In condensation polymerisation molecules are joined together by removing water. Polyesters are formed by condensation polymerisation from: alcohols with two hydroxyl groups (–OH) and carboxylic acids with two carboxyl groups (–COOH) This means that the polyester molecules can continue to grow in both directions with many ester linkages. In order to obtain the structures of the monomers used to make the polyester: 1. Locate the ester link: 58 O C O 2. Water adds on like this: O C O O H H 3. The carboxyl and hydroxyl groups in the monomers can now be formed: O C O H + H O ….. Natural Polymers Many addition and condensation polymers are made by the chemical industry; however, polymers are found in nature in both plants and animals. Starch is a natural condensation polymer made from glucose in plants. Proteins are also natural condensation polymers. They are the major structural material of plant and animal tissue e.g. muscles, skin, hair and are involved in the maintenance and regulation of life processes. 59 Unit 2 Metals Metals conduct electricity because they have a sea of delocalised electrons. LEO: Loss Electrons Oxidation GER: Gain Electrons Reduction Reactions of metals (OXIDATION) 1. Metals and water. metal e.g. + water metal hydroxide potassium + water + hydrogen potassium hydroxide + hydrogen (formula) 2K + 2H2O 2KOH + H2 (ionic) + 2H2O 2K+ 2OH- + H2 2K + 2. Metals and acid. Metals above hydrogen in the electrochemical series react with acids. metal e.g. + acid magnesium + hydrochloric acid salt + hydrogen magnesium + chloride hydrogen (formula) Mg + 2HCl MgCl2 + H2 (ionic) Mg + 2H+ + 2Cl- Mg2+ + 2Cl- + H2 3. Metals and oxygen e.g. metal + oxygen metal oxide magnesium + oxygen magnesium oxide (formula) 2Mg(s) + O2(g) 2MgO(s) (ionic) 2Mg(s)) + O2(g) 2Mg2+O2- (s) 60 4. The Reactivity Series These reactions give an indication of the reactivity of the metal and are summarised below. This is called the reactivity series. Metal Reaction with Oxygen Potassium Sodium React React Lithium Calcium Water react Aluminium with form forming and hydrogen Magnesium Zinc React forming metal hydroxide to Acid steam salt and Iron Tin metal oxide hydrogen Lead Copper No Mercury No Silver No Gold Reaction Reaction 61 Reaction Extracting metals from compounds (REDUCTION) Metals are found as compounds called ores. The more reactive metals form the most stable ores and so are hardest to obtain. 5. Methods of extraction. Heating metal oxides The least reactive metals can be obtained from their ores simply by heating e.g. Metal oxide metal + oxygen mercury oxide mercury + oxygen 2HgO 2Hg + O2 Heating metal oxides with carbon More reactive metals are extracted using carbon to remove the oxygen. e.g. Metal oxide + carbon metal + carbon dioxide lead oxide + carbon lead + carbon dioxide PbO2 + C Pb + CO2 Electrolysis Electricity is needed to obtain the most active metals from their compounds. e.g. aluminium oxide aluminium + oxygen 4Al + 3O2 (Al3+)2(O2-)3 4Al + 3O2 Al3+ + 3e- Al 2Al2O3 62 Reduction More about Electrolysis For electrolysis to happen the ionic lattice must be broken down by dissolving or melting the compound. This makes the ions free to move and so carry the current. e.g. Electrolysis of molten lead (II) bromide- Pb2+(Br-)2 (l) molten lead (II) bromide Product at the negative electrode lead At the negative electrode : Pb (metal ions gain electrons) – At the positive electrode : 2Br (non-metal ions lose electrons) 2+ (aq) + 2e → (aq) → 6. Percentage of Metal in an Ore. Step 1 Calculate the gram formula mass (gfm) K2CO3 - potassium carbonate 3 x 16 = 48 1 x 12 = 12 2 x 39 = 78 = 138g Step 2 Calculate the total mass of the metal in the formula K 3 x 39 = 78g Step 3 Divide the mass of the metal by the gfm and multiply by 100% % mass of the metal = 78 x 100% 138 = 56.5% 63 Product at the positive electrode bromine Br2 (g) + 2e Pb (s) Making electricity Electricity is a flow of charged particles: flow of electrons through metals, flow of ions through solutions or melts. Electrons always flow from metals high in the electrochemical series through the wires to metals lower in the electrochemical series. The further apart two metals are in the electrochemical series the larger the voltage obtained. . 2.7V Mg 1.1V Cu Zn Cu electrolyte 64 Redox & Displacement When electrons go into an electrode we get REDUCTION When electrons leave an electrode we get OXIDATION Electron flow V Zn Ion Bridge Zn2+SO42- Cu Cu2+SO42- The copper ions in the solution accept these electrons and turn into copper metal. Zn Cu2+ + 2e- Zn2+ Cu + 2e- oxidation reduction The oxidation and reduction equations can be combined to form what is known as a redox equation e.g. Zn + Cu2+ Zn2+ 65 + Cu redox Cells involving non-metals. The electrochemical series in the data book has some reactions that involve non-metals V Zn Ion Bridge Zn2+SO42- C I2 / K+I- If we are told that the following reaction occurs: I2 + 2e- 2I- This is Reduction which is where electrons go into an electrode. So electrons must be going in at the C electrode. This means that electrons must be leaving at the Zn electrode (Oxidation) Zn2+ Zn + 2e- reduction So electrons must be flowing from the Zinc to the carbon. Electron flow V Zn Ion Bridge Zn2+SO42- C I2 / K+I- The diagram shows that electrons flow from zinc metal to the carbon rod and the reactions are shown below: Zn Zn2+ + 2ereduction I2 + 2e 2I oxidation The redox equation will be: I2 + Zn 2I- + 66 Zn2+ redox Displacement Metals can displace ions of a less reactive metal. For example: Zn(s) + Cu2+SO42- Cu(s) + Zn2+SO42- Omitting the sulphate ions (SO42-) and the state symbols makes the process clearer. Zn + Cu2+ Cu + Zn2+ Being the more reactive metal the zinc metal is losing its electrons which it gives to the copper ions.. Zn Zn2+ + Cu2+ + 2e- Cu 2 e- oxidation reduction Adding the oxidation and reduction equations together gives us the redox equation Zn + Cu2+ Cu + Zn2+ redox Metals reacting with acids is an example of a displacement reaction. Zn(s) + (H+)2SO42- Zn2+SO42- + H2 The zinc metal is displacing the hydrogen ions. The metals which do not react with acids, can not displace the hydrogen ions and so these metals must be less reactive than hydrogen. Therefore hydrogen is put into the reactivity series between lead and copper. The process of losing electrons is called oxidation and that of gaining electrons is called reduction. Fuel Cells & Rechargeable Batteries. A fuel cell is a device that converts chemical energy from a fuel such as hydrogen into electrical energy through a chemical reaction with oxygen or some other oxidising agent (a substance that causes oxidation) 67 Fertilisers Fertilisers contain the essential elements (NPK) that plants need. Examples are ammonium salts (such as NH4Cl), nitrates (such as KNO3) potassium salts (such as K2SO4) phosphates (such as K3PO4) Percentage composition Fertilisers need to contain different proportions of N : P : K for different plants. When comparing fertilisers it is useful to know the percentage of an element in that fertiliser. calculate the percentage of nitrogen in ammonium nitrate we first calculate the GFM NH4NO3 3x0 1xN 4xH 1XN = 3 x 16 = 48 = 1 x 14 = 14 =4x1 = 4 = 1 X 14 = 14 GFM = 80 g From the total number of nitrogen atoms in ammonium nitrate calculate the total mass of nitrogen. There are 2 nitrogen atoms in the formula therefore total mass nitrogen = 14 + 14 = 28 Percentage nitrogen = mass of nitrogen x 100 GFM = 28 x 100 80 68 =35% To Making fertilisers Making Fertilisers by Neutralisation Nitrate fertilisers are formed using nitric acid: HNO3 + NaOH NaNO3 + H2O Phosphate fertilisers are formed using phosphoric acid: H3PO4 + 3NH3 (NH4)3PO4 For fertilisers containing potassium, an acid is reacted with a potassium compound: K2CO3 + HNO3 KNO3 + H2O + CO2 Ammonium fertilisers are formed using ammonia and nitric acid: NH3 + HNO3 NH4NO3 The Haber Process To make ammonium nitrate we need ammonia and nitric acid. To make nitric acid we need ammonia so the production of ammonia is essential to making ammonium nitrate. Ammonia is formed by reacting nitrogen (from air) and Hydrogen (from natural gas). Iron is used as a catalyst and he reaction is carried out at 400oC. N2(g) + H2(g) NH3(g) The reaction is reversible i.e. it can go in both directions. To help stop this, the ammonia is cooled down to turn it into a liquid. Any unreacted nitrogen and hydrogen is recycled. Unreacted N2 and H2 recycled Nitrogen Hydrogen Reaction Chamber Fe catalyst 400 0C 200 atm N2/H2/NH3 Cooler liquid NH3 The ammonia is then used to make nitric acid. 69 NUCLEAR CHEMISTRY Types of radiation There are 3 types of radiation, alpha ( ), beta ( an electrical field. ) and gamma ( ). Their properties can be studied using α γ β + α β γ paper 2cm Al 5cm concrete Alpha particles Alpha radiation consists of helium nuclei, Alpha can only travel a few centimetres. Beta particles A beta particle is an electron, . Beta can travel metres. Gamma waves Gamma waves are electronegative waves. They can travel 100’s of metres. 70 Nuclear equations In nuclear equations, the sum of the atomic number and mass number on each side of the equation should balance. Alpha emissions Beta emissions Half-life The half-life of a radioisotope never changes. It is the time taken for the sample's activity to fall by half. Example 1 The mass of a radioisotope falls from 1.6g to 0.1g in 2 hours. What is the half-life of this radioisotope? Answer 1.6 g 0.8g 0.4g 0.2g 0.1g The activity halves 4 times in 2 hours Half life = 2/4 = 0.5 hour Example 2 If a 1g sample of a radioisotope with a half-life of 3 days has an activity of 32 c.p.m., how long would it take for the activity to fall to 8 c.p.m.? Answer If 1g sample is 32 cpm. 32 cpm 16 cpm Activity halves 2 times. Half life = 3 days Time taken = 2 x 3 = 6 days 8 cpm 71 Carbon dating Carbon -14 is present in the atmosphere. Carbon dioxide is responsible for carbon-14 entering the food chain. The levels of carbon dioxide in living things stays constant as it is absorbed through photosynthesis or in food, and it decays by beta emission. When living things die no more carbon-14 is absorbed and as the C-14 continue to decay by beta emission its levels fall. We can use the half-life of C-14 (5,700 years) to calculate the age of the object. Example If a piece of wood has 25 % the radiocativity of a living tree then it has undergone 2 half-lives. 100% 50% 25% Half-life = 5700 years Tree died 2 x 5,700 = 11,400 years ago. Uses of radioisotopes Radioisotopes of elements have a wide variety of uses. Cobalt-60 Used in medicine to treat cancer.. Iodine-133 Used to treat thyroid gland. 72