The purpose of this booklet is to bridge the gap between GCSE Double Award Science and Separate Science Chemistry and to prepare for A level.

Transcription

1 A level CHEMISTRY 2018 The purpose of this booklet is to bridge the gap between GCSE Double Award Science and Separate Science Chemistry and to prepare for A level. By the start of the course you should be confident with the content, much of which should already be familiar. So by the start of term: read through the notes (pages 2 to 9). learn the ionic charges given in the table on page 7. complete all the even-numbered questions in the exercises 1 that follow page 10, (writing your answers in the spaces provided.) If there are any exercises that you are finding difficult then try the odd numbers as well. 1 Taken from UA GCE Chem Moles wkbk Iss3 Hurstpierpoint College 1 PJM/SMC

2 What is it all about? Chemistry is the creative central science, the science of changing matter. What is the matter? Nearly all real stuff / matter is a mixture of substances. A pure substance has characteristic properties such as colour, solubility, melting point (Tm) and boiling point (Tb). Elements are the simplest pure substances. There are about one hundred elements, divided into metals and nonmetals although the division is not always clear-cut. Metals conduct electricity whereas nonmetals do not; graphite is a metalloid as it does conduct but generally behaves as a non-metal. Compounds, of which there are millions, are more complicated pure substances, being composed of two or more elements bonded (not mixed!) together. In chemical changes, chemical reactions, new substances are formed. Elements cannot be broken down, they can only build up (synthesise) compounds. Compounds can be converted into other compounds and may be broken down (decomposed) into simpler compounds or into elements. There are three common states of matter: solid, liquid and gas. It is possible to have one state spread through another as in colloids (such as foam, suspensions, gels). Atomic Theory (formulated by Dalton in 1808, updated with comments in italics). 1. All matter is made of minute indivisible particles, atoms. Atoms are themselves made of smaller particles but an atom is the smallest, electrically neutral particle of an element that can participate in a chemical reaction. 2. Atoms cannot be created or destroyed except in nuclear reactions. 3. The atoms of an element are identical. Atoms of an element have the same atomic number but not necessarily the same mass number (isotopes). 4. Atoms of an element are different from the atoms of all other elements. 5. When elements react to form compounds, the atoms of the elements combine, bond, in simple whole number ratios. In chemical reactions, new substances are formed as atoms change their partners. Hurstpierpoint College 2 PJM/SMC

3 Relative atomic mass scale Hydrogen atoms are the lightest. The relative atomic mass, r.a.m. or A r, of an element is the number of times heavier an atom of the element is than one hydrogen atom. This definition will be revised at the start of the AS course. Since a carbon atom is twelve times heavier than a hydrogen atom A r (C) = 12 For water particles, made of three atoms bonded together, the relative formula mass, M r, is calculated thus M r (H 2O) = 2 * = 18.0 and for hydrated aluminium sulfate, the relative formula mass M r (Al 2(SO 4) 3.16H 2O = 2* ( *16.0) + 16*18.0 = Note that the.16h 2O is the water of crystallization, water that is part of the crystals. Relative formula mass is perhaps a better term for what is more commonly known as relative molecular mass. M r (NaCl) = = 58.5 Note that relative atomic and relative formula (molecular) masses have no units. Model of an atom Atomic Structure subatomic particle mass on A r charge scale proton neutron 1 0 electron almost 0-1 volume occupied by electrons nucleus holding protons & neutrons Numbers atomic number = proton number ( = electron number ) mass number = neutron number + proton number Points Both atomic number and mass number really refer to the nucleus of the atom. The mass number of an atom is effectively the same as its A r, because the mass of the electrons is negligible. Hurstpierpoint College 3 PJM/SMC

4 Electron arrangement Electrons are arranged in shells (energy levels) around the nucleus. The first shell (nearest the nucleus the lowest energy level) can hold only two electrons. The second shell can hold only eight electrons. The third shell can hold eight electrons. group period 1 H 1 2 Li Be Na Mg K Ca B 2.3 Al C 2.4 Si N 2.5 P O 2.6 S F 2.7 Cl He 2 Ne 2.8 Ar Periodic Table Elements arranged in order of atomic number. Period Group Noble gases Atomic size Elements having the same electron shell filled. Elements have the same number of electrons in the outer shell (valence electrons) and have similar chemical properties. Elements have a full (or empty) outer shell and are chemically unreactive. Atoms of the other elements bond to get this arrangement. Across a period the increasing nuclear charge causes the atoms to get smaller. The chemistry of an element is governed by its outer valence electrons. When showing the valence electrons (in a dot & cross diagram), set out in the order north, south, east, west before pairing in the same sequence, e.g SYMBOL Hurstpierpoint College 4 PJM/SMC

5 Covalent bonding Nonmetal atoms can effectively gain electrons by sharing pairs forming molecules. Covalent bonding then results from the attraction of the nuclei for the shared pair of electrons. Using the dot & cross approach: show just one atom of each element before bonding (to reduce clutter); bonding involves the pairing of unpaired electrons. Examples: structural formula molecular formula Points Note Covalent bonds are formed by non-metal atoms sharing pairs of electrons. The number of covalent bonds formed by an atom = number of unpaired valence electrons Particles of atoms bonded covalently are called MOLECULES. Inside molecules the bonds are strong. Between molecules there are weak forces. Use the terms bonds & forces to differentiate the relative order of magnitude of interactions. It is sensible to learn the formulae of the most common covalent substances. elements: hydrogen H 2, oxygen O 2, halogen (i.e. chlorine/bromine/iodine) Hal 2, nitrogen N 2 compounds with trivial names: water H 2O, ammonia NH 3, methane CH 4, ethane C 2H 6, propane C 3H 8, butane C 4H 10. Others have systematic names providing an indication of their formulae: carbon monoxide CO, carbon dioxide CO 2, phosphorus trichloride PCl 3, silicon tetrachloride SiCl 4, phosphorus pentachloride PCl 5. Hurstpierpoint College 5 PJM/SMC

6 Ionic bonding When a metal combines with a nonmetal, electrons are transferred between the atoms forming ions. Ionic bonding then results from attractions between the oppositely charged ions. Using the dot & cross approach: show just one atom of each element before bonding (to reduce clutter); ions are charged as a result of having unequal numbers of protons and electrons. Examples formula Points Ionic bonds are formed when electrons are transferred between atoms. Metal atoms lose electrons, becoming positive IONS, ca+ions. Non-metal atoms gain electrons, becoming negative IONS, anions. Ionic bonding is the strong electrostatic attraction between oppositely charged ions.. Ions are arranged in a giant lattice, with ions of the opposite charge nearest, e.g. sodium chloride structure: Hurstpierpoint College 6 PJM/SMC

7 Ionic Formulae The charges of the most commonly encountered ions are: Cations Anions gp 1 metals M + halides (gp 7) Hal - 1+ hydrogen H + + ammonium NH 4 1- hydrogencarbonate HCO 3- hydroxide OH - silver Ag + manganate(vii) MnO 4 - nitrate NO 3 - nitrite NO 2 - most metals M 2+ carbonate, CO 3 2- chromate(vi) CrO dichromate(vi) Cr 2O 2-7 oxide O 2- sulfide S 2- sulfite SO 3 2- sulfate SO aluminium Al phosphate PO 4 3- chromium(iii) Cr 3+ iron(iii) Fe 3+ Hurstpierpoint College 7 PJM/SMC

8 Rules for naming ionic compounds These will be refined during the course 1. the cation is stated before the anion. 2. binary (two element) compounds end in -ide, e.g. Na 2O sodium oxide, MgCl 2 magnesium chloride, LiH lithium hydride 3. exceptions to rule 2 are hydroxides and cyanides, e.g. Ca(OH) 2 calcium hydroxide and KCN potassium cyanide 4. with metals other than those in groups 1, 2 or 3, e.g. the transition metals which can have more than one charge, the charge must be shown using Roman numerals, e.g. FeCl 2 iron(ii) chloride & FeCl 3 iron(iii) chloride and similarly PbO lead(ii) oxide & PbO 2 lead(iv) oxide. 5. oxyanions end in -ate or ite, the latter indicating a lower oxygen content than the former, e.g. sulfate SO & sulfite SO 3 and nitrate NO 3 & nitrite NO 2-. Method for deducing the formulae of ionic compounds In an ionic compound the charge of the cations must be balanced by the charge of the anions. Knowing the charge of the ions therefore gives a simple way of deducing the chemical formula. Ions Substance formula (see page 7) sodium chloride Na 1+ Cl 1- NaCl sodium oxide Na 1+ O 2- Na 2O magnesium oxide Mg 2+ O 2- MgO calcium chloride Ca 2+ Cl 1- CaCl 2 aluminium oxide Al 3+ O 2- Al 2O 3 sodium hydroxide Na 1+ OH 1- NaOH calcium hydroxide Ca 2+ OH 1- Ca(OH) 2 calcium carbonate Ca 2+ CO 3 2- CaCO 3 calcium hydrogencarbonate Ca 2+ HCO 3 1- Ca(HCO 3) 2 Note that in the formula the figure 1 is not written (but is assumed if no other number present) a bracket must be placed around a radical if it is multiplied by 2 or more numbers are cancelled / simplified if possible. Hurstpierpoint College 8 PJM/SMC

9 Writing equations Tackle in four steps. 1. Write word equation. 2. Put formulae below words. Metal just symbol Covalent stuff (compound & non-metal) you need to know formula Ionic compound use ionic charges 3. Balance. The numbers of atoms on each side of the equation must be the same. Extra numbers are only allowed in front of a formula. 4. Add state symbols. This is an optional extra in most cases but is essential in some situations (e.g. precipitation reactions). (s) solid (l) liquid (g) gas (aq) aqueous, i.e. dissolved in water Examples These have been set out step by step to show the working; in practice this is unnecessary! magnesium + oxygen magnesium oxide Mg + O 2 MgO 2 Mg + O 2 2 MgO 2 Mg (s) + O 2 (g) 2 MgO (s) sodium + water (hydrogen sodium + hydrogen hydroxide) hydroxide Na + H 2O NaOH + H 2 2 Na + 2 H 2O 2 NaOH + H 2 2 Na (s) + 2 H 2O (l) 2 NaOH (aq) + H 2 (g) sulfuric (hydrogen + iron(iii) iron(iii) + water acid sulfate) hydroxide sulfate H 2SO 4 + Fe(OH) 3 Fe 2(SO 4) 3 + H 2O 3 H 2SO Fe(OH) 3 Fe 2(SO 4) H 2O 3 H 2SO 4 (aq) + 2 Fe(OH) 3 (s) Fe 2(SO 4) 3 (aq) + 6 H 2O (l) Hurstpierpoint College 9 PJM/SMC

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11 Exercise 1: Calculating the relative formula mass of compounds See page 3 of the notes. You will find relative atomic masses on the periodic table on page H2O 2 CO2 3 H2SO4 4 CuSO4 5 C2H5OH 6 Na2CO3 7 Pb3O4 8 KMnO4 9 K2Cr2O7 10 CH3COCH3 11 Ca(OH)2 12 Mg(NO3)2 13 Ca(HCO3)2 14 Pb(NO3)2 15 Al(NO3)3 16 Fe2(SO4)3 17 (NH4)2SO4 18 CuSO4.5H2O 19 (COOH)2.2H2O 20 (NH4)2SO4.Fe2(SO4)3.24H2O Hurstpierpoint College 11 PJM/SMC

12 Exercise 2: Writing formulae from their names See pages 5, 7 & 8 of the notes. 1 Carbon Dioxide 2 Carbon Monoxide 3 Sulfur Trioxide 4 Phosphorus Pentachloride 5 Silicon Tetrachloride 6 Dinitrogen tetraoxide 7 Potassium Bromide 8 Magnesium Chloride 9 Silver Chloride 10 Sodium Hydroxide 11 Calcium Carbonate 12 Sodium Carbonate 13 Lithium Sulfate 14 Sodium Phosphate 15 Aluminium Oxide 16 Magnesium Hydroxide 17 Aluminium Hydroxide 18 Barium Nitrate 19 Ammonium Carbonate 20 Aluminium Sulfate Hurstpierpoint College 12 PJM/SMC

13 31 Copper(II) Oxide 32 Copper(II) Sulfate 33 Iron(II) Chloride 34 Copper(I) Chloride 35 Lead(II) Carbonate 36 Lead(IV) Oxide 37 Tin(IV) Chloride 38 Iron(III) Chloride 39 Iron(III) Sulfate 40. Silver(I)Oxide Hurstpierpoint College 13 PJM/SMC

14 Exercise 3: Deducing names from formulae See pages 5, 7 & 8 of the notes. Questions 1-12 Some of these you just need to know, so look them up if you don t. 1 CO2 2 NH3 3 SO3 4 CH4 5 H2SO4 6 HNO3 7 C8H18 8 NaCl 9 Ca(NO3)2 10 (NH4)2CO3 11 KHCO3 12 CsAt Questions Roman numerals will be needed in these 13 FeSO4 14 FeCl3 15 PbO 16 CuCl 17 AgNO3 18 Co(NO3)2 19. PbCl4 20. Cr(OH)3 Hurstpierpoint College 14 PJM/SMC

15 Exercise 4 Balancing equations Balance the following equations. All the formulae are correct. 1 H2 + O2 H2O 2 BaCl2 + NaOH Ba(OH)2 + NaCl 3 H2SO4 + KOH K2SO4 + H2O 4 K2CO3 + HCl KCl + H2O + CO2 5 CaCO3 + HNO3 Ca(NO3)2 + H2O + CO2 6 Ca + H2O Ca(OH)2 + H2 7 Pb(NO3)2 + NaI PbI2 + NaNO3 8 Al2(SO4)3 + NaOH Al(OH)3 + Na2SO4 9 N2 + H2 NH3 10 H3PO4 + NaOH Na3PO4 + H2O 11 NaNO3 NaNO2 + O2 12 CH4 + O2 CO2 + H2O 13 C4H10 + O2 CO2 + H2O 14 H3PO4 + NaOH NaH2PO4 + H2O 15 6NaOH + Cl2 NaClO3 + NaCl + H2O 16 Fe2O3 + CO Fe + CO2 17 C2H5OH + PCl3 C2H5Cl + H3PO3 18 2KMnO4 + HCl MnCl2 + Cl2 + 8H2O + KCl 19 Al(OH)3 + NaOH NaAlO2 + H2O 20 Pb(NO3)2 PbO + NO2 + O2 Hurstpierpoint College 15 PJM/SMC

16 Exercise 5 Writing equations in symbols from equations in words In the following examples you will need to convert the names of the materials into formulae and then balance the resulting equation. 1. Hydrogen gas reacts with oxygen gas to make water. 2 Liquid silicon tetrachloride reacts with water to produce solid silicon dioxide and hydrogen chloride gas. 3. Zinc metal reacts with copper sulfate solution to produce solid copper metal and zinc sulfate solution. 4. When octane (C8H18) vapour is burnt with excess air in a car engine, carbon dioxide and water vapour are produced. 5 When magnesium is added to dilute nitric acid, a solution of magnesium nitrate is produced and bubbles of hydrogen gas. Hurstpierpoint College 16 PJM/SMC

17 6. When lead(ii) nitrate crystals are heated in a dry tube lead(ii) oxide, nitrogen dioxide gas and oxygen are produced. 7 When a solution of calcium hydrogencarbonate is heated, a precipitate of calcium carbonate is produced together with carbon dioxide gas and water. 8. Solid calcium hydroxide reacts with solid ammonium chloride on heating to produce solid calcium chloride, steam and ammonia gas. 9. A solution of ammonium hydroxide will neutralize sulfuric acid to make ammonium sulfate solution and water. 10 When solutions of silver nitrate and calcium chloride are mixed a white precipitate of silver chloride is formed. Hurstpierpoint College 17 PJM/SMC

18 Exercise 6 What s wrong here? The following equations have one or more mistakes. Are the formulae correct? Is the balancing correct? Are the state symbols correct and, most importantly, does the reaction actually happen? Identify the error(s) and then rewrite the equation correctly. 1 Na (s) + H2O (l) NaOH (aq) + H (g) 2 PbNO3 (aq) + NaCl (aq) PbCl (s) + NaNO3 (aq) 3 CaOH2 (aq) + 2 HCl (aq) CaCl2 (aq) + 2H2O (l) 4 C2H4 (g) + 2 O2 (g) 2 CO2 (g) + 2 H2 (g) 5 MgSO4 (aq) + 2 NaOH Ca(OH)2 (s) + Na2SO4 (aq) 6 Cu(NO3)2 (s) + CuO (s) 2 NO (g) + O3 (g) 7 Cu (s) + H2SO4 (aq) CuSO4 (aq) + H2 (g) 8 AlCl2 (s) + 2 KOH (aq) Al(OH)2 (s) + 2 KCl (aq) 9 NaCO3 (s) + 2 HCl (aq) NaCl2 (aq) + CO2 (g) + H2O (l) 10 2 AgNO3 (aq) + MgCl2 (aq) Mg(NO3)2 (s) + 2 AgCl (aq) Hurstpierpoint College 18 PJM/SMC