A-Level OCR Chemistry Revision Pack PDF

THE ULTIMATE A-LEVEL OCR CHEMISTRY CHEATSHEET PACK  Achieve Better Grades with Less Stress Our intelligent platform provides high quality, personalised support that is more effective than private tutoring available 24/7 at a fraction of the cost Bite-size Smart Mini Revisio...

THE ULTIMATE A-LEVEL OCR CHEMISTRY CHEATSHEET PACK  Achieve Better Grades with Less Stress Our intelligent platform provides high quality, personalised support that is more effective than private tutoring available 24/7 at a fraction of the cost Bite-size Smart Mini Revision 24/7 Tutor Personalised Predicted Videos Quizzes Guides Support Learning Paths Exam Packs LEARN MORE How to Use The aim of this pack is simple — we wanted to condense the A-level Chemistry course into a few super condensed pages. Now you have a concise summary of the entire course that focuses on the most important definitions, key terms, diagrams and concepts. We’ve spent weeks working with top designers, academic writers and illustrators to ensure this is the best cheatsheet out there. Our promise to you is you won’t find anything better. The cheatsheet pack has been built off the OCR specification to ensure no important information is missed — below is a table which summarises how our cheatsheets map to the OCR specification. Specification Points Cheatsheet 2.1.1 Atomic Structure & Isotopes 2.1.1 - 2.1.2 Compounds, Formulae, Equations & Moles 2.1.4 - 2.1.5 Acids, Bases & Redox 2.2.2 Electron Configuration & Bonding 3.1.1 Periodicity 3.1.2 - 3.1.3 Group 2 & Group 7 3.2.1 Enthalpy 3.2.2 Reaction Rates 3.2.3 Equilibrium, The Haber Process & Partial Pressure 4.1.1 Introduction to Organic Chemistry 4.1.2 - 4.1.3 Alkanes & Alkenes 4.2.1 - 4.2.2 Alcohols & Haloalkanes 4.2.3 - 4.2.4 Synthesis & Analysis 5.1.1 Rate Equations & Rate Determination 5.1.2 - 5.1.3 Acids, Bases & Buffers 5.2.1 Lattice Enthalpy & Born-Haber Cycles 5.2.2 - 5.2.3 Entropy & Redox 5.3.1 Transition Elements 6.1.1 Benzene & Phenol 6.1.2 - 6.2.1 Carbonyls 6.2.2 Amines & Amino Acids 6.2.3 - 6.2.5 Condensation Polymers & Synthesis 6.3.1 Chromatography & Tests for Ions We hope you enjoy using it and wish you the best of luck in your A-levels. SnapRevise ATOMIC STRUCTURE & ISOTOPES CHEAT SHEET Atomic Structure Relative Masses Electronic Configuration Atoms are the components that make up all elements. Relative isotopic mass is the mass of an atom of an isotope Electrons orbit the central nucleus in shells. Each shell can hold Atoms are made up of three types of sub-atomic particles – compared with 1/12th of the mass of an atom of carbon-12. For 2n2 electrons, where n is the principal quantum number. protons, neutrons, and electrons an isotope, the relative isotopic mass = its mass number. Electron shells are made up of atomic orbitals, which are Protons and neutrons make up the nucleus, where most of the regions in space where electrons may be found. mass is concentrated. Electrons orbit the nucleus in shells Relative atomic mass is the ratio of the average mass of Each shell is composed of one or more orbitals and each orbital an atom of an element to 1/12th of the mass of an atom of can hold one pair of electrons. carbon-12. There are four main types of orbitals: s-, p-, d-, and f- Relative molecular mass is the ratio of the average mass of a molecule of an element or compound to 1/12th of the mass of an atom of carbon-12. Relative formula mass is similar to relative molecular mass but applies to ionic compounds. Mass Spectrometry There are 1 s-orbital, 3 p-orbitals, 5 d-orbitals and 7-p orbitals possible in each subshell. Shell/quantum number n = 1 can have the occupancy 1s2 Mass spectrometry is a form of molecular chemical analysis n = 2 can have the occupancy 2s22p6 Particle Relative Mass Relative Charge that allows the masses of individual molecules or isotopes to be n = 3 can have the occupancy 3s23p63d10 determined. n = 4 can have the occupancy 4s24p64d104f14 Proton 1 +1 Mass spectrometry can be used to provide structural Within each shell, Neutron 1 0 information, identify an unknown compound, or determine the orbitals that are of relative abundance of each isotope of an element. the same energy level Electron 1/2000 -1 are grouped together The mass spectrum gives information about the relative in sub-shells. abundance of isotopes on the y axis and about the relative Sub-shells have isotopic mass on the x axis. different energy Mass Number & Isotopes The mass spectrum can be used to determine the relative levels. Note that 4s is atomic mass (Ar) lower in energy than 3d, so 4s will fill first. Element, X Shells and sub-shells are filled with electrons according to a set of rules: ◦ Atomic orbitals with the same energy fill individually first before pairing A mass spectrum for a molecular sample shows the relative ◦ Aufbau principle – the lowest available energy level is filled molecular mass on the x axis. first Electron configuration is written with n representing principal quantum Mass number, A, is the total number of protons and neutrons in number. X is the type of orbital and y is the nucleus the number of electrons in the orbitals Atomic number, Z, is the number of protons. The number of of the subshell e.g. potassium has 19 positively charged protons is equal to the number of negatively electrons and its electron configuration charged electrons in an atom, making the atom neutrally is written as 1s22s22p63s23p64s1 charged Electronic configurations of ions can be determined by using Mass number = number of protons + number of neutrons the same building up principle. Atomic number = number of protons = number of electrons e.g. K+ is the same as potassium but one electron has been Ions are formed by atoms losing or gaining electrons. removed, therefore it is: 1s22s22p63s23p64s1 ◦ A charge of x- means that the number of electrons in the ion Take care with d-block ions, as 4s fills before 3d, but it also is the atomic number + x empties before 3d when forming ions Electrons have an intrinsic property (spin). For two electrons ◦ A charge of x+ means that the number of electrons in the ion in the same orbital, the spin must be opposite to minimise the is the atomic number - x repulsion. Isotopes are atoms with the same number of protons and different numbers of neutrons. Therefore, they have different mass numbers but the same atomic number. Isotopes of the same element have the same electronic configuration so react in the same way in chemical reactions but have slightly different physical properties. No more than two electrons can fill an atomic orbital snaprevise.co.uk COMPOUNDS, FORMULAE, EQUATIONS & MOLES CHEAT SHEET Ionic Compounds The Mole & Concentration Balanced Equations An ion is an atom or molecule with a net charge formed The mole is the unit used to quantify the amount of a When a chemical reaction occurs, no atoms are created or destroyed. through the gain or loss of electrons substance. It can be applied to any amount of chemical species, The atoms in the reactants rearrange to form the products. including atoms, electrons, molecules and ions A cation is formed from the loss of electrons In a balanced equation, there is the same number of atoms of A mole is the amount of substance that contains the same each element in both the reactants and products. An anion is formed from the gain of electrons number of atoms or particles as 12 g of carbon-12. State symbols are written after every species to indicate the Ionic compounds are composed of oppositely charged ions. The The number of particles in 12g of 12C is the Avogadro constant physical state overall charge is zero as the ionic charges balance. of 6.022 x 1023 mol-1. ◦ Solid (s) e.g. write the formula of sodium carbonate ◦ Liquid (l) Sodium is in Group 1, and so it will lose one electron, Na+ ◦ Gaseous (g) The carbonate ion has the formula CO32- n is the number of moles (mol) ◦ Aqueous (aq) - dissolved in water We need 2 x Na+ to balance the -2 charge of CO32- m is the mass (g) The formula is Na2CO3 Ionic equations can be written for any reaction involving ions M is the molar mass (g mol-1) in solution, where only the reacting ions and the products they The concentration of a solution is the amount of solute present form are included. Empirical & Molecular Formula in a known volume of solution The net ionic equation shows only the ions directly involved in The empirical formula is the simplest whole-number ratio of the reaction (removing spectator ions). atoms of each element present in a compound. e.g. NaCl (aq) + AgNO3 (aq) → AgCl (s) + NaNO3 (aq) c is the concentration (mol dm-3) Na+ (aq) + Cl- (aq) + Ag + (aq) + NO3- (aq) → AgCl (s) + Na+ The empirical formula can be calculated from the composition n is the number of moles in solution (mol) (aq) + NO3- (aq) by mass or percentage by mass. V is the volume (dm3) Net ionic equation: e.g. 6.2 g of P is combined with O2 to form 14.2 g of Cl- (aq) + Ag+ (aq) → AgCl (s) phosphorous oxide. Calculate the empirical formula of the Remember: compound. ◦ 1 dm3 = 1000 cm3 mass of O2: 14.2 g - 6.2 g = 8 g ◦ 1 m3 = 1000 dm3 number of moles of each element: Atom Economy & Percentage Yield Atom economy is a theoretical measure of the proportion of Gas Equations atoms from the reactants that form the desired product. In order to calculate it a balanced chemical equation is required. Divide through by the smallest number of moles to get the One mole of any gas under standard conditions will occupy the whole number ratio: same volume. % atom economy = The molar gas volume is 24 dm-3 mol-1 under standard conditions of 298 K and 100 kPa Maximising atom economy has important economic, ethical and environmental advantages: The number of moles of gas can be calculated using the equation: ◦ More sustainable (uses fewer raw materials) Empirical formula: P2O5 ◦ Minimises chemical waste The molecular formula gives the number and type of atoms of ◦ Maximises efficiency each element in a molecule. It is made up of a whole number of n is the number of moles of gas (mol) empirical units. ◦ Less money is spent on separation processes V is volume (dm-3) The molecular formula can be determined using the empirical In an ideal gas the assumptions are made that: The limiting reagent is the reagent not in excess. It dictates the formula and relative molecular mass of the molecule. ◦ Intermolecular forces between the gas particles are negligible theoretical yield and the amount of product actually formed. e.g. Determine the molecular formula of a compound with ◦ The volume of the particles themselves, relative to the empirical formula CH2 and a relative molecular mass of 224. Percentage yield is a measure of the percentage of reactants volume of their container, is negligible Relative molecular mass of the empirical formula: that have been converted into the desired product. It gives a The ideal gas equation is: measure of the efficiency of a reaction route. C H2 12 + (1 × 2) = 14 The percentage yield is reduced by the formation of unwanted p is pressure (Pa) by-products, any reactant that remains unreacted, or product Divide the relative molecular mass by that of the empirical V is volume (m3) that cannot be extracted from the reaction vessel. n is the number of moles (mol) formula: R is the gas constant (8.314 JK-1) 224/14=16 T is temperature (K) Molecular formula: 16 ×CH2 = C16H32 snaprevise.co.uk ACIDS, BASES & REDOX CHEAT SHEET Acids Oxidation Numbers Redox Titrations Common acids include: Redox titrations can show how much oxidising agent is needed Oxidation number is a number representing the number of ◦ Hydrochloric acid, HCl to react exactly with a reducing agent electrons lost or gained by an atom in a compound. ◦ Sulphuric acid, H2SO4 ◦ Nitric acid, HNO3 Oxidation is a loss of electrons during a reaction or an increase Redox reactions often involve colour changes, making it easy to in oxidation number. see when a reaction has reached completion Acids dissociate in water: HA (aq) ⇌ A– (aq) + H+ (aq) Reduction is a gain of electrons during a reaction or a decrease Manganate(VII) ions are readily reduced to Mn2+ ions under This process of HA separating into A– and H+ ions is called in oxidation number acidic conditions. The MnO4- ions are purple, Mn2+ are dissociation. colourless). The rules for assigning oxidation numbers: ◦ Oxidation: Fe2+ → Fe3+ + e– The strength of an acid describes how much of it dissociates ◦ An uncombined element has an oxidation number of 0 when it dissolves. ◦ Reduction: 8H+ + MnO4- + 5e– → Mn2+ + 4H2O ◦ A simple ion (of a single element) has an oxidation number ◦ Overall: 8H+ + MnO4- + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+ A strong acid is an acid which dissociates almost completely in equal to the charge on the ion water or aqueous solution. HA (aq) ⇌ A– (aq) + H+ (aq) ◦ The sum of oxidation numbers of the elements in a In iodine-thiosulphate titrations, iodine is reduced to form compound is equal to the overall charge of the compound iodide ions and thiosulphate ions are oxidised. Iodine is often A weak acid is an acid which is only partially dissociated in combined with starch, which makes it dark blue. As iodine is water or aqueous solution. HA (aq) ⇌ A– (aq) + H+ (aq) ◦ The charge on a complex ion, e.g. NH4+, is equal to the sum of consumed the solution turns pale yellow at the end point the oxidation numbers Hydrochloric acid is a strong acid: ◦ Oxidation: 2S2O32– (aq) → S4O62– (aq) + 2e– ◦ The most electronegative element in a compound always has HCl (aq) Æ H (aq) + Cl (aq) + – a negative oxidation number ◦ Reduction: I2 (aq) + 2e– → 2I– (aq) ◦ Oxygen is always –2 except in peroxides where its –1 ◦ Overall: 2S2O32– (aq) + I2 (aq) → S4O62– (aq) + 2I– (aq) Bases ◦ Hydrogen is +1 except in metal hydrides where its –1 Common bases include: ◦ Sodium hydroxide, NaOH Oxidation numbers are represented by Roman numerals when Titrations naming compounds ◦ Potassium hydroxide, KOH ◦ Ammonia, NH3 Acid–base titrations can be used to find the concentration of a Oxidation numbers can be used to write formulae for a sample of either an acid or a base. Alkalis also dissociate in water, forming bases: compound XOH Æ X+ + OH- e.g. NaOH Æ Na+ + OH- A known concentration of an acid is gradually added to a known volume of a base of unknown concentration until the Redox Reactions solution is neutralised i.e. the titration reaches the ‘end point’ Reduction: The gain of electrons and decrease in oxidation A burette is used to gradually add the acid, a pipette is used number of an element to add a known volume of the base, and an indicator is used to cause a colour change when the reaction reaches the end point Oxidation: The loss of electrons and increase in oxidation number of an element Calculations are used to determine the concentration of the unknown solution. Neutralisation Redox reactions involve both oxidation and reduction Acids and bases react together in neutralisation reactions: These calculations are those you already know: Oxidising agents cause oxidation of other species, and so are n=cxv m = n x Mr themselves reduced Acid–carbonate reactions: Reducing agents cause reduction of other species, and so are themselves oxidised Acid–metal reactions: In the reaction below, H is reduced, Na is oxidised. 2 HCl + 2 Na Æ 2 NaCl + H2 +1 0 +1 0 In each example H+ reacts with OH- to form water. snaprevise.co.uk ELECTRON CONFIGURATION & BONDING CHEAT SHEET Ionic Bonding Giant Covalent Structures Bond Polarity An ionic bond is an electrostatic attraction between positive Giant covalent lattices consist of billions of atoms held together Electronegativity is the power of an atom to attract the pair of and negative ions, resulting in a giant lattice by a network of strong covalent bonds. electrons in a covalent bond. It is affected by: A lattice is a giant three-dimensional structure, where each ion Carbon forms three types of giant covalent structures: graphite, ◦ Nuclear charge is surrounded by oppositely charged ions ◦ Atomic radius graphene and diamond. They are often formed between ions of metals and non-metals, ◦ Shielding by electrons in inner shells where each atom aims to gain a stable full outer shell. In diamond, each carbon atom is bonded to four other carbon atoms A polar bond will form between two atoms with different Electrons are represented electronegativity values where the electrons in the bond are not by dots and crosses, This makes it very strong, with an incredibly shared equally. This causes a permanent dipole. which helps visualise the high melting point, hardness, insolubility. It is an Partial negative charges (δ-), and partial positive charges (δ+) origin of the electrons electrical insulator because all the outer shell are used to show that a bond is polar. clearly electrons are being used in covalent bonds, so Molecules containing polar bonds are not always polar. The strength of an ionic there are no mobile charge carriers The symmetry of polar bonds can cancel the effect of any bond depends on permanent dipole. ◦ The charge on the ions - Ions Metallic Bonding with higher charges will have a greater electrostatic attraction between them and will Metallic bonding is the strong electrostatic Intermolecular Forces form stronger bonds. attraction between positive metal ions and Intermolecular interactions are the forces of attraction between negative delocalised electrons in a metal lattice molecules that do not involve the transfer of electrons. They are ◦ The distance between the ions - Smaller ions have a smaller Metals have a fixed lattice structure of positive ions. the result of the constant and random movements of electrons internuclear distance so the electrostatic forces of attraction The outer shell of electrons is delocalised, which can carry current. within electron shells will be greater. Smaller ions form stronger ionic bonds in Metals have high melting and boiling points due to the large The greater the strength of the bonds, the higher the melting more closely packed lattices. amounts of energy needed to overcome the metallic bonds. and boiling points. E.g., down a group boiling point increases, Properties of ionic structures The greater the charge of a metal ion, the more electrons are due to greater induced dipole=dipole forces, as electron ◦ High melting and boiling points delocalised. The smaller the size of the metal ions, the closer number increases ◦ Soluble in polar solvents they are to the delocalised electrons. This results in stronger ◦ Electrical conductivity when molten or dissolved in water bonds and a higher melting point Metals are malleable and ductile. The ions can slide and move past each other as there are no bonds holding specific ions Covalent Bonds together Covalent bonding usually occurs between two atoms of non- metals, so that the atoms each gain a stable full outer shell A covalent bond is a strong electrostatic attraction between Simple Molecules Induced dipole-dipole (London) forces are caused by the a shared pair of outer electrons and the nuclei of the bonded Simple molecular structures have low boiling points. In the solid temporary unequal distribution of charge due to the constantly atoms state they are held together by weak intermolecular forces moving electrons. The temporary dipole can induce a Unlike ionic bonds, covalent bonds are localised Covalent bonding pairs and lone pairs of electrons are charge temporary dipole in a neighbouring atom and the two dipoles A single covalent bond involves one shared pair of outer clouds that repel each other will be attracted to each other. This can occur between almost electrons, while a multiple covalent bond involves more than Electron pairs in the outer shell arrange themselves as far apart all molecules. one shared pair of outer electrons as possible to minimise repulsion Permanent dipole-dipole forces are caused by molecules with Bonding pairs repel each other equally. a permanent dipole being attracted to the opposite charge in While lone pairs repel other pairs more other permanent dipoles, if correctly aligned. because they are more electron dense. A normal line represents the bond is in plane of the paper. A dotted wedge is a bond going into the paper and a bold Hydrogen bonding requires: wedge is a bond coming out of the paper. ◦ An electron-deficient hydrogen atom and a more The shape of a molecule is determined by electronegative atom. These atoms are oxygen, fluorine and Two single covalent A double covalent electron-pair repulsion theory nitrogen bonds in H2O bond in O2 ◦ A lone pair of electrons on a highly electronegative atom on A co-ordinate (dative covalent) bond involves one atom donating another molecule both electrons in a covalent bond. Electron Bonding Lone Bond Hydrogen bonding in water allows ice to be less dense than Example water as hydrogen bonds hold the water molecules apart at a A dative covalent bond is represented by an arrow from Pairs Pairs Pairs Angles the atom that is donating both electrons to the atom that is Linear 2 2 0 180° CO2 fixed position, forming an open lattice accepting both electrons Water has high melting and boiling points than expected due to Trigonal Planar 3 3 0 120° BF3 hydrogen bonds being the strongest intermolecular forces ◦ Average bond enthalpy is a measure of the enthalpy required to break one mole of covalent bonds, hence it indicates the Tetrahedral 4 4 0 109.5° CH4 strength of covalent bonds. Pyramidal 4 3 1 107° NH3 ◦ Covalent compounds can have two structures, either: simple Non-linear 4 2 2 104.5° H2O molecular structure OR giant covalent lattice. Octahedral 6 6 0 90° SF6 snaprevise.co.uk PERIODICITY CHEAT SHEET Periodic Table Period 2 & 3 Ionisation Energy In the Periodic Table, elements are arranged in order of Across Period 2, two electrons fill the 2s sub-shell before six Ionisation energy is a measure of the energy required to increasing atomic number electrons will fill the 2p sub-shell completely remove an electron from an atom of an element to form an ion. Elements can be grouped into periods (horizontal rows) and Across Period 3, two electrons will fill the 3s sub-shell before six groups (vertical columns). electrons will fill the 3p sub-shell First ionisation energy is the energy required to remove one electron from each atom in one mole of the gaseous element to Groups contain elements with similar physical and chemical Across Period 2 and Period 3, the trend in meting point relates form one mole of gaseous 1+ ions. properties, as they have the same number of outer shell to the structure of the elements. The melting point increases X(g) → X+(g) + e- electrons from Group 1 to Group 14 because the elements have giant structures (metallic and then covalent). The melting points Successive ionisation energies apply to the removal of electrons Periodicity is a regularly repeating pattern of atomic, physical, decrease from group 14 to 15 because the structure changes after the first ionisation energy. The nth ionisation energy is: and chemical properties with increasing atomic number to simple molecular, which is only held together by weak intermolecular forces X(n-1)+ (g) → Xn+(g) + e- The Periodic Table can be split into s-, p-, d-, and f- blocks. Which is determined by which orbital the highest energy Successive ionisation energies provide evidence for the shell electron is in. structure of atoms. ◦ Within each shell successive ionisation energies increase, as there is less electron repulsion ◦ Between shells, there are big jumps in ionisation energies, as the electric is removed from a shell closer to the nucleus Factors affecting ionisation energies: ◦ Atomic radii - The larger the atomic radius, the further away the outer electrons are held from the nucleus, and the smaller the nuclear attraction. ◦ Nuclear charge - The greater the nuclear charge, the greater the attractive force on the outer electrons. Across the period, the atomic radius decreases as effective ◦ Shielding - electrons repel each other due to their negative nuclear charge increases and there is no increase in shielding charge. The greater the number of inner shells of electrons, the greater the repulsion of the outer shell of electrons. The greater the attraction, the harder it is to remove an electron. Therefore, the ionisation energy will be larger. Atomic radii show periodicity. Across a period, the radius decreases while down a group, the radius increases. Ionisation energy increases across a period, electrons are all added to the same shell resulting in greater attraction. Ionisation energy decreases across a group, as the number of shells increases, so does the atomic radius and shielding, reducing attraction. snaprevise.co.uk GROUP 2 & GROUP 7 CHEAT SHEET Group 2, Alkaline Earth Metals Group 7, The Halogens Use of Chlorine and Chlorate The halogens exist as diatomic (X2) molecules Chlorine can dissolve in water to form hydrochloric acid and All Group 2 elements have 2 electrons in their outer s-subshell. Halogen Appearance chloric(I) acid in a disproportion reaction (chlorine is both oxidised and reduced). Down the group ionisation energy decreases, as the atomic F2 Pale yellow gas Cl2 (g) + H2O (l) → HCl (aq) + HClO (aq) radius and shielding increases, decreasing the attraction of the Cl2 Green gas Oxidation States: 0 -1 +1 electron to the nucleus Br2 Red-brown liquid Chloric(I) acid is an oxidising agent that kills bacteria and Melting point also decrease down the group due to the sanitises water increased atomic radii and shielding, so metallic bonding is I2 Black solid weaker Cl2 (g) + 2NaOH (aq) → NaCl (aq) + NaClO (aq) + H2O (l) Down the group electronegativity decreases, as atomic radius Group 2 elements undergo redox reactions: and shielding increases. Sodium chlorate(I) has a bleaching action which allows it to be ◦ with water to form hydroxides and hydrogen gas: Down the group boiling point increases, as the molecules have an active ingredient in household bleach. Mg (s) + 2H2O(l) → Mg(OH)2 (aq) + H2 (g) a greater surface area and electrons for stronger London forces 0 +1 +2 0 The advantages of using chlorine to treat water is that it Halogen Appearance Boiling point London forces prevents the spread of waterborne disease and sanitises water. ◦ with oxygen to form oxides: F2 Pale yellow gas Lowest Weakest However, it is added without direct customer consent, chlorine Mg (s) + O2 (g) → MgO (s) is a respiratory irritant and it could react to form chlorinated 0 0 +2-2 Cl2 Green gas hydrocarbons, implicated in cancers. ◦ with dilute acids to form salts: Br2 Red-brown liquid Mg (s) + HCl(aq) → MgCl2 (aq) + H2 (g) 0 +1 +2 0 I2 Black solid Highest Strongest The reactions involve the loss of electrons to form 2+ ions, The halogens have seven electrons in their outer shell (s2p5), which requires the input of the first two ionisation energies so they need to gain one electron to obtain a stable octet, forming ions with a charge of 1–, this is shown by the reduction First and second ionisation energy values decrease down group of chlorine: 2 because of increase in atomic radius, increase in shielding, and hence a decrease in attraction between the nucleus and Cl2 + 2 e– Æ 2 Cl– outer electron A more reactive halogen will displace the halide ion of a less reactive halogen from solution As ionisation becomes easier, reactivity increases,therefore reactivity increases down group 2 ◦ If a solution of chlorine (Cl2) is added to a solution of bromide (Br–), the chlorine will displace the bromine Group 2 Group 2 Cl2 + 2 Br– Æ 2 Cl– + Br2 Solubility of Hydroxide Element Hydroxide Halogen Colour in Water Mg Mg(OH)2 Slightly soluble Chlorine Pale green Bromine Orange Ca Ca(OH)2 Sparingly soluble Iodine Brown Precipitation reactions with aqueous silver ions are used to test More soluble than Mg(OH)2 for halide ions. Silver halides are formed when AgNO3 (aq) is Sr Sr(OH)2 and Ca(OH)2 added to a halide solution: Ag+ (aq) + X– (aq) Æ AgX (s) Ba Ba(OH)2 Most soluble Silver Colour of Addition of Dilute Addition of Concentrated Halide Precipitate Ammonia Ammonia Magnesium hydroxide is used in indigestion tablets to neutralise Dissolves to give a Dissolves to give a AgCl White stomach acid colourless solution colourless solution Dissolves to give a AgBr Cream Does not dissolve Calcium hydroxide is used in increase the pH of acidic soils colourless solution AgI Yellow Does not dissolve Does not dissolve snaprevise.co.uk ENTHALPY CHEAT SHEET Enthalpy Coffee Cup Calorimetry Hess’s Law Enthalpy, H, is the thermal energy that is stored in a system A coffee cup calorimeter can be used to calculate enthalpy Hess’s law is that the enthalpy change of a reaction is changes of neutralisation independent of the route taken Enthalpy change is the heat energy change measured under conditions of constant pressure. ◦ The reaction mixture is placed in a Styrofoam cup with a lid to keep it insulated, along with a stirrer and a thermometer. Enthalpy change values are usually given under standard The Styrofoam cup can be held within another Styrofoam conditions (100 kPa, 298K, 1 mol dm-3 & the standard state of Ɵ cup in a beaker for maximum insulation the substance). Standard enthalpy changes are denoted by ΔH. ◦ A measured volume of the first reactant is added, and the temperature is recorded until stable. A measured amount total enthalpy of products (kJ mol–1) of the second reactant is added, and the temperature is measured evert minute whilst constantly stirring total enthalpy of reactants (kJ mol–1) ◦ A graph of temperature against time is plotted, with a line of The standard enthalpy change of reaction (ΔHrϴ) is the enthalpy Enthalpy changes of combustion can be used to find the best fit change that accompanies a reaction in molar quantities shown enthalpy change of a reaction in a chemical equation under standard conditions ◦ q can be calculated using ∆T and converting the volumes of the solutions into masses using their densities (assumed to The standard enthalpy of formation (ΔfHƟ) is the enthalpy be 1 g cm-3) change when one mole of a substance is formed from its constituent elements, with all reactants and products being in their standard states and under standard conditions Bond Enthalpies The standard enthalpy of combustion (ΔcHƟ) is the enthalpy Average bond enthalpy is the energy required to break one change when one mole of a substance is completely burnt in mole of a specified type of bond in a gaseous molecule oxygen, with all reactants and products being in their standard states and under standard conditions The strength of a covalent bond varies according to the environment in which they are found, so an average value is The standard enthalpy change of neutralisation (ΔneutH ) is the ϴ taken Enthalpy changes of formation can be used to find the enthalpy energy change that accompanies the reaction of an acid by a change of a reaction base to form one mole of H2O (l) under standard conditions The actual bond enthalpy is specific to each individual molecule The enthalpy of formation for an element is 0 kJ mol-1. Bond breaking is endothermic, while bond making is Exothermic & Endothermic Reactions exothermic Enthalpy change can be predicted using bond enthalpies Exothermic reactions increase the temperature of the surroundings. ΔH is negative The enthalpy change values calculated from bond enthalpies are approximate and not as accurate as those calculated from Hess’ Law cycles Spirit Burner Calorimetry A spirit burner calorimeter can be used to calculate enthalpy changes of combustion Calorimetry ◦ The substance being heated or cooled, usually water, is Calorimetry is the process Endothermic reactions placed in a beaker with a thermometer. A spirit burner is of measuring the amount decrease the temperature placed underneath the beaker to heat it of heat given off or taken in of the surroundings ΔH is ◦ The spirit burner containing the fuel is weighed and a during a chemical reaction positive. known volume of water is added to the beaker and its initial temperature measured q is the heat change (J) Activation energy is the ◦ The spirit burner is burnt, and the water continuously stirred m is the mass of the minimum energy required for substance (g) a reaction to occur ◦ After a few minutes, the flame is extinguished, and the spirit –1 –1 burger reweighed. The final temperature of the water is c is the specific heat capacity (J g K ) measured ΔT is the temperature change (K or °C) ◦ The measured values are used to calculate q snaprevise.co.uk REACTION RATES CHEAT SHEET Collision Theory Rate of Reaction Calculating the Rate of Reaction Particles in a gas or liquid are constantly moving and colliding The rate of reaction is the change in the concentration of The rate of a reaction can be calculated by recording the with each other. However, not all collisions result in a reaction. a reactant or product in a given time. amount of a reactant or a product at regular time intervals For a reaction to occur, two particles must collide with: during a single reaction (continuous monitoring) or by ◦ Sufficient energy to overcome the activation energy determining the initial rate of several reactions and finding an ◦ The correct orientation Factors affecting the rate of reaction: average. ◦ Temperature- increasing the temperature increases the kinetic energy of the molecules, leading to more frequent There are several methods of continuous monitoring successful collisions and an increase in the rate of reaction ◦ Pressure — increasing the pressure of a gaseous reaction ◦ Monitoring by gas collection – use a syringe to measure the increases the number of gaseous molecules in a given volume, volume of gas evolved so molecules are closer together. This leads to more frequent ◦ Mass loss – set up the system on a balance so the decrease in successful collisions and an increase in the rate of reaction mass of reactants can be measured Maxwell-Boltzmann Distribution ◦ Concentration — increasing the concentration of an aqueous reactant increases the number of molecules in a given volume, ◦ Colorimetry – concentration change is monitored by The Maxwell-Boltzmann curve shows the distribution of so molecules are closer together. This leads to more frequent measuring the amount of light absorbed by a sample molecular kinetic energies in a gas at a constant temperature successful collisions and an increase in the rate of reaction In the Boltzmann distribution: ◦ Catalysts — adding A graph can be plotted using the variable measured against ◦ The area under the curve is equal to the total number of a catalyst provides time: molecules in the system an alternative ◦ The curve starts at the origin — no molecules have zero pathway that has The gradient of the graph is equal to the rate i.e. energy a lower activation ◦ y/x = concentration change/time = rate ◦ Only molecules with an energy greater than the activation energy, leading energy, Ea, can react to more frequent ◦ The curve’s peak successful collisions represents the most and an increase in probable energy that the rate of reaction a molecule will have ◦ The average energy is to the right of, and slightly The Effect of Temperature and greater than, the most probable energy Catalysts on The Maxwell-Boltzmann Distribution Catalysts At higher temperatures, the kinetic energy of the molecules increases, so the molecules move faster A catalyst is a substance that increases the rate of a chemical A greater proportion of molecules have an energy greater than reaction without being used up in the process. the activation energy, so more frequent successful collisions Catalysts work by reacting with the reactants to form an occur, increasing the rate of reaction. intermediate. Catalysts are then regenerated later in the The area under the curve remains the same as the number reaction of molecules in the system Homogeneous catalysts are in the same phase as the reactants remains the same Heterogenous catalysts are in a different phase to the Catalysts lower the activation reactants energy of a reaction by Heterogeneous catalysts work by: providing an alternative reaction route with a lower ◦ Adsorption — the reactants forming weak bonds with the activation energy atoms on the surface of the catalyst, holding the reactants in A greater proportion of the correct position for them to react. molecules have an energy ◦ Desorption — the products detach from the atoms on the greater than the new, lower surface of the catalyst activation energy, and more ◦ As the products detach more reactants can be adsorbed, and frequent successful collisions the process is repeated occur, increasing the rate of The use of catalysts has economic, environmental and social reaction. benefits The addition of the catalyst ◦ The demand for fossil fuels is lowered does not change the ◦ Emission of pollutants are lowered distribution of the molecular ◦ Reduced energy use means lower costs energies. snaprevise.co.uk EQUILIBRIUM, THE HABER PROCESS & PARTIAL PRESSURE CHEAT SHEET Chemical Equilibrium Haber Process Equilibrium Constant Reversible reactions at equilibrium are denoted by the Nitrogen gas reacts with hydrogen gas to form ammonia in the The equilibrium constant, KC, indicates where the equilibrium equilibrium symbol: ⇌ Haber process lies – it is the ratio of the concentration of products and –1 Many chemical reactions are reversible N2 (g) + 3H2 (g) ⇌ 2NH3 (g) ΔH = –92 kj mol reactants in a reversible reaction The optimum conditions for the Haber process are: The concentration, in mol dm–3, of a species X involved in the A chemical system is in dynamic equilibrium when all three of ◦ High pressure — increasing the pressure will cause the expression for KC is represented by [X] the following conditions are met: position of equilibrium to shift to the right, favouring the ◦ The concentration of reactants and products is constant For the general reaction: aA + bB ⇌ dD + eE forwards reaction as there are fewer moles on the right ◦ The rate of the forwards reaction is the same as the rate of (two moles) than on the left (four moles). Therefore, more the backwards reaction ammonia is produced, and the pressure will be reduced, ◦ The reaction is in a closed system opposing the change Le Chatelier’s principle- when a system in dynamic equilibrium ◦ Low temperature — the forward reaction is exothermic, The units of the equilibrium constant will vary depending is subjected to change, the position of the equilibrium will shift so decreasing the temperature will cause the position of on how many species take part in the reaction and the to oppose the change equilibrium to shift to the right, favouring the forwards stoichiometry reaction. Therefore, more ammonia is produced and the If the conditions of a reaction are changed, the position of The magnitude of KC indicates the extent of the reaction: temperature increases, opposing the change equilibrium will shift to favour either the forwards or backwards ◦ KC = 1: equilibrium lies halfway between the reactants and the These optimum conditions are not favourable because: reaction to oppose the change. products ◦ Maintaining high pressures is expensive and unsafe The effect of a change in temperature on the position of ◦ Low temperature would mean a slow rate of reaction ◦ KC > 1: equilibrium lies further to the right and the products equilibrium will depend on the enthalpy change of the reaction. In industry, conditions must be used that strike a balance are favoured If the forward reaction is exothermic and the backwards between obtaining a good yield and being economically feasible ◦ KC < 1: equilibrium lies further to the left and the reactants reaction is endothermic: Compromise conditions must be used in industrial processes to are favoured ◦ Increasing the temperature will cause the position of ensure a good yield whilst considering the following factors: Equilibrium Constant, Kp equilibrium to shift in the endothermic direction to absorb ◦ The rate of reaction the added heat ◦ Cost and risks of equipment ◦ Decreasing the temperature will cause the position of ◦ Side reactions Kp is an equilibrium constant associated with equilibrium equilibrium to shift in the exothermic direction to add more The industrial conditions used today in the Haber process are a reactions, only involving gases. It is written in terms of partial heat compromise and are: pressures rather than concentrations. ◦ 400–500 °C — allows a reasonable rate of reaction and yield Changing the concentration of a reactant or product will For the equilibrium reaction: ◦ 200 atm — allows a high yield without costing too much or cause the position of equilibrium to shift to oppose this aA (g) + bB (g) ⇌ cC (g) + dD (g) posing a safety risk change ◦ Iron catalyst increases the rate of reaction and allows ◦ Increasing the concentration of the reactants will cause equilibrium to be established more quickly at a lower the position of equilibrium to shift to make more of the temperature, saving energy and increasing profits product ◦ Increasing the concentration of the products will cause the position of equilibrium to shift to remove the extra product, making more reactants Partial Pressure The units of Kp are variable and depend on the specific reaction Mole fractions indicate the fraction of a mixture occupied by a that is under consideration Changing the pressure of a system will only change the position particular gas. of the equilibrium if the reaction involves gases. Temperature can affect the position of equilibrium ◦ Increasing the pressure of the system will cause the position ◦ Increasing the temperature will cause the position of of equilibrium to shift to the side with the fewest moles of equilibrium to shift in the endothermic direction. Increasing Kp gas ◦ Decreasing the temperature will cause the position of Partial pressure is the pressure each gas in a mixture would ◦ Decreasing the pressure of the system will cause the equilibrium to shift in the exothermic direction. Decreasing Kp exert on its own. The partial pressure of a gas A is denoted by position of equilibrium to shift to the side with more moles p(A) or PA. Pressure can affect the position of equilibrium of gas The total pressure of a gas mixture is the sum of all partial ◦ Increasing the pressure will shift the position of equilibrium Adding a catalyst does not change the position of pressures from each gas. to the side with less moles of gas equilibrium, but increases the rate at which the equilibrium is The amount of pressure a gas exerts in a fixed volume depends ◦ Decreasing the pressure will shift the position of equilibrium established on how many particles there are – more particles means greater to the side with more moles of gas pressure. Changing pressure will not affect the value of Kp; instead, the p(A) = mole fraction of gas A x total pressure position of equilibrium will shift to keep Kp constant. p(A) partial pressure of gas A (kPa) Catalysts do not affect the value of Kp. Instead, it affects how mole fraction of gas A (mol) quickly equilibrium is reached. total pressure (kPa) snaprevise.co.uk INTRODUCTION TO ORGANIC CHEMISTRY CHEAT SHEET Formulae Nomenclature Reaction Mechanisms General formula — the simplest algebraic formula for Hydrocarbons can be: Carbon Atoms Prefix in alkyl group a homologous series ◦ Aliphatic — carbon atoms form 1 Methyl straight or branched chains 2 Ethyl Structural formula gives the minimum detail on the arrangement of atoms in a molecule, without drawing any bonds ◦ Alicyclic — carbon atoms form a ring 3 Propyl 4 Butyl ◦ Aromatic — carbon atoms form a ring Molecular formula shows the number and types of atoms of 5 Pentyl and have a delocalised electron each element in a compound. However, it does not give any 6 Hexyl system information on how the molecule is bonded together. ◦ Saturated — containing only single bonds Bond fission can be homolytic or heterolytic Skeletal formula is a simplified formula used to represent organic molecules. Lines represent bonds between atoms, Homologous series are compounds with the same functional Homolytic Fission Heterolytic Fission junctions are carbon atoms. Other labels are omitted. group and similar chemical and physical properties. They differ by the number of repeating units they contain When the bond breaks, each When the bond breaks, both Displayed formula shows the relative positioning of atoms and electron in the bond goes to the electrons in the bond go the bonds between them. All atoms and bonds are shown A functional group is the group of atoms responsible for the a different atom. to the same atom. characteristic reactions of a compound. Empirical formula the simplest whole–number ratio of each element present in a compound. To name a compound: This results in the formation of This results in the formation ◦ The stem is the main part of the name derived from the highly reactive free radicals, of a positively charged cation longest carbon chain. each with an unpaired and a negatively charged Isomerism ◦ The suffix after the stem, comes from the most significant electron, represented by anion. Isomers are compounds with the same molecular formula but functional group a dot. a different arrangement of atoms ◦ The prefix before the stem comes from functional groups attached to the main carbon chain Bonds are formed on the collision of: Structural isomers are compounds with the same molecular ◦ Numbers and hyphens indicating the position of functional ◦ Two free radials with unpaired electrons formula but a different structural formula. They may contain the same functional group but on a different carbon atom, or groups on the carbon chain different functional groups e.g. ketones and aldehydes with ◦ Functional groups are prioritised alphabetically the same number of carbon atoms have the same molecular ◦ Oppositely charged ions formula Functional Compound Prefix Suffix Group Bonds can also be formed when ions are attracted to dipoles, Alkanes – –ane e.g. nucleophilic substitution reactions Alkenes – –ene Alcohols Hydroxy– –ol Carboxylic Acids – –oic acid Fluoro– Chloro– Haloalkanes – Bromo– Iodo– Aldehydes – –al Ketones – –one snaprevise.co.uk ALKANES & ALKENES CHEAT SHEET Alkanes Alkenes Addition Reactions of Alkenes Alkanes are a homologous series made up of saturated Alkenes are a homologous series made up of unsaturated An electrophile is an electron pair acceptor attracted to areas of hydrocarbons, containing only carbon and hydrogen atoms hydrocarbons containing at least one C=C double bond, with high electron density joined by sigma bonds the general formula CnH2n. Electrophilic addition with HBr Carbon atoms have four electrons in their outer shell. Therefore, C-H — these are σ-bonds formed due to the direct overlap of each carbon atom can form four covalent bonds, resulting in the electron clouds belonging to each carbon atom tetrahedral geometry C=C — these contain both a σ-bond and two π-bonds. It is an Alkanes are not polar and only weak London area of high electron density, so can react with electrophiles. forces of attraction occur, as carbon and π-bonds — formed due to the overlap of adjacent p-orbitals above hydrogen have similar electronegativities. and below the carbon atoms. It restricts the rotation around the Major and minor products form Boiling point of alkanes increases with chain planar C=C double bond, leading to stereoisomerism in some alkenes when an unsymmetrical alkene length, as there is a greater surface area and Due to electron pair repulsion, alkenes have a trigonal planar reacts. The major product will form number of electrons for stronger London forces. shape with a bond angle of 120°. from the most stable carbocation Shorter carbon chain alkanes and more branched alkanes have intermediate, with the most alkyl a smaller the surface area for contact between molecules and groups attached. weaker London forces, resulting in lower boiling points Stereoisomerism Tertiary > secondary> primary C-C and C-H σ bonds are strong and non-polar therefore Stereoisomers are organic compounds with the same molecular carbocations. This reduces the alkanes are not reactive. The electronegativities of C and H are and structural formulae but a different arrangement of atoms in positive charge, due to inductive so similar that the C-H bonds are considered non-polar space electron donation. E/Z isomerism is a type of stereoisomerism that can arise in Electrophilic addition with Br2. A dipole is induced in the halogen molecule by the high electron Combustion of Alkanes alkenes due to the restricted rotation around the C=C bond If a carbon atom has two of the same substituent attached, it density of the C=C Alkanes can be used as a fuel source will not show E/Z isomerism Bromine water acts as a test for unsaturation. An unsaturated In complete combustion, the alkane burns with a clean Substituents can be assigned priorities based on atomic mass compound will decolourise the bromine water, as a dihaloalkane blue flame. Water vapour and carbon dioxide are formed using Cahn-Ingold-Prelog rules to name E/Z isomers. The will form, using up the bromine (greenhouse gases). E.g. greater the atomic mass, the higher the priority At 150° in the presence of a catalyst (Ni), alkenes undergo C3H8 (g) + 5O2 (g) Æ 3CO2 (g) + 4H2O (g) When the highest priority groups are on different sides of the addition reactions with H2 to form alkanes In incomplete combustion, the alkane burns with a dirty yellow double bond, the isomer is an E-isomer. At high pressure in flame. It can produce carbon, carbon monoxide and unburned When the highest priority groups are on the same side of the the presence of a hydrocarbons as products. double bond, the isomer is a Z-isomer sulfuric acid catalyst, CH4 (g) + O2 (g) Æ C (s) + 2H2O (g) gaseous alkenes CH4 (g) + 1.5O2 (g) Æ CO (g) + 2H2O (g) react with steam to form alcohols As with addition with HBr, a major and minor product forms Chlorination of Alkanes Alkanes react with the halogens, specifically chlorine and Addition Polymers bromine, in the presence of UV light to form haloalkanes Methane reacts with chlorine to form chloromethane and Addition polymers are macromolecules made from small hydrogen chloride: CH4 + Cl2 Æ CH3Cl + HCl repeating units of monomers This reaction is a free radical substitution with the steps: They are formed from alkenes ◦ Initiation — free radicals are formed when exposed to UV Cl2 Cis-trans isomerism is a type of E/Z isomerism. It arises when and substituted alkenes Æ 2Cl the carbon atoms at each end of the double bond are attached Addition polymers are ◦ Propagation — free radicals are used up and created in a to two different groups, but at least one of the groups attached unreactive ss the main carbon chain reaction Cl + CH4 Æ CH3 + HCl to one carbon are identical to a group on the other carbon chain is saturated and non-polar, with only weak London forced CH3 + Cl2 Æ CH3Cl + Cl between the chains ◦ Termination — free radicals are removed. Polyalkanes are non-biodegradable therefore pose a risk to the 2Cl Æ Cl2