Chemistry Edexcel IAS Unit 1 Notes PDF

Chemistry Edexcel IAS Unit 1 Notes Ruaa Topic 1: Formulae, equations and amount of substance Simplifying ionic equations (tips): Write the normal equation replacing the reactants and products with ions, and remove spectator ions If there is a solid product amongst only solutions, simply write the i...

Chemistry Edexcel IAS Unit 1 Notes Ruaa Topic 1: Formulae, equations and amount of substance Simplifying ionic equations (tips): Write the normal equation replacing the reactants and products with ions, and remove spectator ions If there is a solid product amongst only solutions, simply write the ions needed to make the product as the rest will be spectator ions Reactions of acids: o Metal + acid -> salt + hydrogen o Metal oxide + acid -> salt + water o Metal hydroxide + acid -> salt + water o Alkali + acid -> salt + water o Metal carbonate + acid -> salt + water + carbon dioxide o Hydrogencarbonate + acid -> salt + water + carbon dioxide Displacement reactions: Keep in mind that single displacement reactions are redox. Double displacement reactions are not redox (no overall change in oxidation numbers). Precipitation reactions: The test for carbon dioxide: Ca(OH)2 (aq) + CO2 (g) -> CaCO3 (s) + H2O (l) Calcium carbonate is a white precipitate, limewater turns milky/cloudy so CO2 is present. The test for sulfates (two examples) SO42- (aq) + Ba2+ (aq) -> BaSO4 (s) Barium sulfate is a white precipitate which forms, indicating presence of sulfate ions. The test for halides: Cl- (aq) + Ag+ (aq) -> AgCl (s) 1 Relative atomic mass / Ar definition: The weighted average mass of an atom compared to 1/12 of the mass of 12C Ar = mean mass of an atom of an element/ 1/12 the mass of 12C Mole calculations: 1. Mol = mass in g/molar mass N = m/M 2. % yield = actual * 100/ theoretical 3. Atom economy = molar mass of desired product/molar mass of all products 4. V = mol * molar volume (24 at r.t.p) 5. Mass concentration = mass/V 6. Molar concentration = n/V 7. Concentration in ppm = solute mass * 106 / solvent mass 8. Concentration in ppm = V of gas * 106 / V of air 9. pV = nRT p must be in Pa, V must be in m3, n is mol, T must be in K, R is given (8.31 J mol-1K-1) Unit conversions (!!!) kPa = 1,000 Pa m3 = 1,000,000 cm3 m3 = 1000 dm3 0 K = -273 celcius Topic 2: Atomic structure and the Periodic Table Mass spectrometry: How it works (don’t need to know but helpful) 1. Sample is vaporized because it must be in a gaseous state. 2. Vapor is bombarded with high energy electrons which ionize the vapor by removing one electron or more. (It is preferred to only remove one electron because if the charge is +1 then the m/z ratio will be dependent only on the mass.) 3. An electric or magnetic field will cause the ions to accelerate. 4. A larger m/z ratio will mean a smaller deflection. 5. The ions pass through a slit and are collected on a metallic plate, the strength of the magnetic field is what controls the ions which are detected. 2 How to determine Mr of diatomic molecules: The peaks at 35 and 37 correspond to relative isotopic mass and the peaks at 70, 72 and 74 correspond to the possible combinations of atoms. The probability of each combination is what determines the height of those peaks. The ratio of 35:37 in the sample is 3:1 which gives an approximate Ar of 35.5 so getting a 35 and 35 combination (70) is more likely than 37 and 37. (a full explanation is in book page 46) Orbitals and electronic configuration: Quantum shell > subshell > orbital The first quantum shell has 1 subshell (1s) The second quantum shell has two subshells (2s and 2p) 3 The third quantum shell has three subshells (3s 3p 3d) The fourth quantum shell has four subshells (4s 4p 4d 4f) An orbital can only have 2 electrons, no matter what orbital it is. It’s the entire subshell that can hold more. subshell No. of orbitals Max e- no. in subshell s 1 2 p 3 6 d 5 10 f 7 14 You must be able to draw and name these in the exam. Rules of filling electron shells: 1. Aufabu principle: Electrons are added to lowest energy orbitals first. 2. Hund’s rule: electrons will occupy orbitals singly before pairing. 3. Pauli exclusion principle: electrons in the same orbitals have opposite spins. 4. A filled or half-filled orbital is more stable, so you need to make exceptions when filling orbitals. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d…… Notice how the 4s comes before the 3d, this is because it has lower energy (this is the only anomaly you need to know of) Rule 4 example: 24Cr : 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5 Instead of filling the 4s subshell completely, one electron is given to the 3d subshell so they are both half-filled and more stable. 4  This is how to draw out the electronic configurations, start by filling every orbital with one arrow then pair the remaining ones (Hund’s rule) Ionisation energy (definition): A measure of the energy required to remove one electron from one mole of an atom of an element in the gaseous state. Note: it is VERY important that the element be in a gaseous state. First ionisation energy: State symbols are always required for this equation, you will lose X (g) -> X+(g) + e- marks for not putting the state symbol. Second ionisation energy: X+ (g) -> X2+ (g) + e- (and so on) To find logarithm of ionisation energy, use 10 = a and N = ionisation energy 5 Ionisation energy is affected by: 1. Atomic radius: as atomic radius increases, IE decreases because electrons are further away from the nucleus, attraction is less and less energy is required to remove them. 2. Electron-electron repulsion/shielding (or spin pair repulsion): more repulsion means lower IE because there is higher energy in the quantum shell which gives a lower IE. Shielding also exists between electrons in different orbitals. 3. Nuclear charge (number of protons): as nuclear charge increases, IE increases because there is more attraction between the electrons and the nucleus and it is harder to remove the electrons. IE = energy of electron when removed – energy of electron when in orbital 4. Higher energy quantum shell > lower IE because further from nucleus. Note: the increase in nuclear charge often has a stronger effect than shielding which lowers the energy of an electron overall even if there is shielding. There is a large jump in IE energy when you need to break through a quantum shell to reach the next electron. Note: it is clear from this graph that there are 3 quantum shells which have electrons. Common question: explain what can be deduced from the graph about the electronic structure of sodium (or whatever element) Overall increase in IE, jump between electron 1 and 2 and electron 9 and 10. Large jump means start of a new quantum shell. Last 2 electrons are much harder to remove than the pervious 8, hence they are in the first quantum shell. Trends across the periodic table: Periodic properties are sometimes referred to as periodicity and they are the regularly repeating patterns of atomic, physical and chemical properties. 6  1. Atomic radii Atomic radius is the measure of the distance between the nucleus and the boundary of an electron cloud. These boundaries are not well defined so we use the distance between 2 nuclei to determine atomic radius: o Atomic radius decreases across a period. This is because the nuclear charge of an atom increases and there is more attraction between the nucleus and the outer shell electrons, pulling them closer to the nucleus and decreasing the distance between the nucleus and the boundary. o Atomic radius increases down a group. This is because the number of occupied shells increases, making the atom bigger. 2. Melting and boiling points o Elements with large lattice structures have high melting and boiling points. o Elements with simple molecular structures have low melting and boiling points. Mg and Al have higher melting points than Na because they have a higher charge and more delocalized electrons in the lattice Silicon has a higher melting point than phosphorus because it is a giant covalent structure and strong covalent bonds need more energy to break than weak intermolecular forces in P 7 The slight increase in melting point for sulfur is because sulfur is a larger molecule than P and Cl so it has stronger van der Waals forces (unit 2) 3. Ionisation energy o Across a period there is an increase in the first ionisation energy. This is because more protons are being added causing an increase in nuclear charge. There is an increased attraction between the outer shell electrons and the nucleus so they will be more tightly held which makes it more difficult to remove an electron. This increase in nuclear charge is stronger than the increase in electron-electron repulsion. o Down a group there is a decrease in the first ionisation energy. This is because the atomic radius increases so electrons are farther away from the nucleus so the attraction diminishes and it becomes easier to remove electrons. There is also stronger e-e repulsion. (this pattern exists for groups 1-2 and 5-8) ANOMALIES: The first IE of Boron is lower than the first IE of Beryllium. The outer electron of Boron has more energy as it is in a 2p orbital so the IE required to remove it is lower than it is to remove the electron in the 2s orbital of beryllium. Topic 3: Bonding and structure Ionic bonding is the electrostatic attraction between positive and negative ions (or cations and anions) The strength of an ionic bond can be determined by calculating the amount of energy required in 1 mol of a solid to separate the ions to infinity (when they are no longer attracted to one another). The smaller the ion and bigger the charge, the more energy needed to overcome the attraction. Proof ions exist: you can pass a current through Copper(II)Chromate(VI) which is green. Cu2+ ions are blue and will turn the solution near the anode blue. CrO42- ions are yellow and will turn solution near cathode yellow. 8  O2- Mg2 + The atomic radius of an ion with a larger positive charge is smaller because electrons are more strongly attracted to the nucleus Polarisation and polarising power Polarisation is the distortion of an anion’s electron density causing a region where electrons exist in orbital overlap. Polarising power is the ability of a cation to attract the electrons from an anion towards itself. (a cation with a high charge and small size has the highest polarizing power) Polarisation happens when there is a high charge and a small cation or a high charge and a large anion. The slight overlap and sharing of electrons causes a degree of covalent bonding to exist in an ionic substance. Properties of ionic compounds: o Brittle: when stress is applied, layers slide causing same charged ions to be next to each other which repel and cause it to break apart o Conduct electricity when molten (mobile ions can carry current) 9  o High melting and boiling point (lattice structure) o Soluble (energy required to break apart the structure can be provided by hydration because the positive and negative ions are both attracted to water molecules because of their polarity) Covalent bonding is the electrostatic attraction between the shared pair of electrons and the nuclei of two non-metals. End-to-end bond (sigma bonds) σ bonds Side-by-side bonds (pi bonds) π bonds π bonds can only be formed after a σ bond is formed so they only exist in molecules with double or triple bonds. In double bonds, one is σ and the other is π In triple bonds the first is σ and the other 2 are π Bond length is the distance between the nuclei of two covalently bonded atoms Bond strength is the amount of energy needed to break one mole of the molecule in a gaseous state. The shorter the bond length the stronger the bond (because there’s an increased electrostatic attraction) Electronegativity is the ability of an atom to attract a bonding pair of electrons 10 Chlorine has a higher electronegativity than hydrogen so the electrons are attracted to the chlorine making this a polar bond. Chlorine is delta negative *partially negative and hydrogen is delta positive *partially positive If the difference in electronegativity is large enough, the electrons will be transferred and the bond will be ionic. o Electronegativity increases across a period due to increase in nuclear charge o Electronegativity decreases down a group because there are more energy levels which shield the bonded electrons from the nucleus Polar bonds have the same overlap of orbitals as normal covalent structures but the difference in electronegativity causes a slight positive and slight negative charge -> some ionic character. Bonding continuum shows how bonding types range from purely covalent to purely ionic. After a certain difference in electronegativity the bonds can only be between a metal and a non-metal and are considered ionic. The octet rule refers to the tendency of atoms to form compounds that have 8 electrons in their outer shell, but there are some exceptions to this rule: o Beryllium chloride (BeCl2) which has 4 electrons in its outer shell o Boron trichloride (BCl3) which has 6 electrons in its outer shell o Phosphorus (V) chloride (PCl5) which has 10 electrons in its outer shell o Sulfur hexafluoride (SF6) which has 12 electrons in its outer shell 11 Dative bonds form when an empty orbital of an atom overlaps with an orbital containing a non-bonding pair of electrons. The electrons are shared but they come from only one of the atoms. The dative bond is denoted by an arrow from the donor to the acceptor. Bonding in Aluminum chloride AlCl3 (which forms Al2Cl6 when it’s just above its sublimation temperature, a dimer; a molecule consisting of 2 identical molecules linked together) Dative bond Aluminum chloride is considered a covalent bond because the difference in electronegativity between Al and Cl is quite small, so this leads the electrons to be partially shared. More info: the Al ion is highly charge dense, and when it polarises the chlorine anion, it stretches and distorts the electron cloud so much it shows high covalent character. It's an Aluminium thing because it has a 3+ charge while at the same time being small. So it's fundamentally ionic, but the electron cloud is distorted so much in the lattice structure over to the cation that the covalent character is very high. Shapes of molecules Electron pair repulsion theory (EPR or VESPR) o The shape of a molecule or ion is caused by the repulsion between electron pairs – bond and lone pairs – that surround the central atom. 12 o The electron pairs arrange themselves in a way such that the repulsion between them is at a minimum or the space between them is at a maximum. o Lone pair-lone pair repulsion > bond pair-lone pair repulsion > bond pair-bond pair repulsion (both diagrams are in the book pages 82 & 83) 13  Tip: the angles decrease by 1.5 when a lone pair is added. Tetrahedral (no lone pairs): 109.5 Trigonal pyramidal (1 lone pair): 107 V-shaped (2 lone pairs): 104.5 Non-polar and polar molecules It is important to note that just because a molecule has polar bonds does not mean it is a polar molecule overall. This symbol is used to denote a dipole with the arrowhead pointing to the more electronegative atom. Symmetrical molecules are non-polar. 14 All of these molecules are symmetrical and have no overall dipole because the dipoles cancel each other out The dipoles in a symmetrical molecule can reinforce each other and cause a net dipole: The charge is unbalanced because of the hydrogen which causes the bottom part to be delta negative. Metallic bonding is the electrostatic attraction between cations in a sea of dealocalised electrons. Properties of metallic lattices + explanations: 1. High melting temperatures o To melt a metal you must overcome the forces of attraction between the cations and electrons in the lattice to allow the cations to move freely and the energy to overcome these forces is very large. o The melting temperature depends on the number of delocalized electrons per cation. o Metals in group 1 have relatively low melting temperatures, whereas metals in the d-block have high melting temperatures. 2. Electrical conductivity o The delocalized electrons move to the positive terminal of a cell. Flow of electrons -> electrical current 3. Thermal conductivity 15  o Free-moving electrons pass kinetic energy along the metal, the cations also pass kinetic energy to one another as they are closely packed. 4. Malleable and ductile o Metals can be hammered into different shapes (malleable) and drawn into wires (ductile). o When stress is applied to a metal the layers slide over each other. o The electrons move with the cations to prevent them from being next to each other and repelling each other. Giant covalent lattices 1. Diamond o Bonding: each carbon atom forms four σ bonds to four other carbon atoms. o Tetrahedral arrangement. o Extremely hard because the bonds are very strong. o Extremely high melting temperature because there is a huge number of bonds, so a large energy is required to break them. 2. Graphite o Bonding: each carbon atom is bonded to 3 others by σ bonds. o The remaining electron is in a p orbital. The orbitals overlap forming a cloud of delocalized electrons between layers. o The delocalized electrons means graphite can conduct electricity – but electrons can only move parallel to layers. 3. Graphene The same as graphite but is just one layer. 16 Topic 4: Introductory organic chemistry and alkanes. General formulae Alkane: CnH2n+2 Alkene: CnH2n Cycloalkane: CnH2n Cycloalkene: CnH2n-2 Displayed formula shows all the atoms and all the bonds Structural formula shows the bonds on each carbon grouped together Skeletal formula is a zigzag which shows only the bond between carbons (every point is a carbon): if something other than carbon or hydrogen is present, we write it down. The functional group of a molecule is an atom or group of atoms which give it distinct predictable properties. Nomenculture is a set of rules to name organic compounds. Prefixes indicate number of carbon attoms, suffixes show the functional group, locants show the position of atoms in a molecule. (Page 106) Alkanes Alkanes are saturated hydrocarbons and are generally unreactive and do not contain a functional group, they only undergo substitution reactions. Structural isomers (types) Chain isomerism: molecules with different carbon chains Ex: butane and methylpropane Take care to name molecules with their longest chain. Position isomerism: molecules with the same functional group in different positions on the same carbon chain Cycloalkanes: these come up often in multiple choice questions. 17 These all have the same number of carbons and hydrogrens Each carbon atom has 4 bonds, and the total number of hydrogen atoms is 10. Alkanes as fuels Fractional distillation: Crude oil is a mixture which contains mostly hydrocarbons. It is heated in a furnace which turns the crude oil into vapour (mostly) The column has a temperature gradient (hot at bottom cool at top) Fractions with longer chains & high boiling points are near bottom Fractions with shorter chains & lower boiling points are near top Cracking Demand for short-chain hydrocarbons is higher so long chain hydrocarbons are cracked Alkane -> alkane + alkene or alkane + hydrogen Reforming Some structures burn more efficiently/smoothly than others, (cyclic compounds > branched compounds > straight chain compounds) so this process is used to convert straight chain alkanes into branched or cyclic hydrocarbons by heating them with a catalyst (ie platinum). Complete combustion Alkane + Oxygen -> Carbon dioxide + Water Incomplete combustion 18 Alkane + Oxygen -> Carbon/Carbon monoxide + Water Some molecules in crude oil contain atoms of sulfur. During combustion, the atoms form sulfur dioxide which then reacts with oxygen in the atmosphere to form sulfur trioxide. These oxides are acidic. When they dissolve in water in the atmosphere they form sulfurous acid and sulfuric acid which contribute to acid rain. Some molecules contain nitrogen which can form nitrogen oxides. Nitrogen dioxide is acidic and can dissolve in water to form nitrous acid and nitric acid. Both contribute to acid rain. Note: although this can happen, when asked in exams to mention the main contributor to acid rain always mention sulfur and not nitrogen. Catalytic converters (same as IGCSE) Catalytic converters in car engines turn harmful gases into less harmful gases. They use catalysts such ass platinum, palladium and rhodium. The metals are spread thinly over a honeycomb mesh (increases surface area for reaction and in turn saves money) Alternative fuels Needed because crude oil and other fossil fuels are non-renewable (finite) Global warming and climate change caused by the pollutants produced when burning them. Carbon neutrality The idea is that if a fuel emits as much carbon as it uses up, it is carbon neutral – ie it has no effect on the net amount of carbon dioxide in the atmosphere 19 Comparison between fossil and biofuels: Types of reactions in organic chemistry 1. Addition reactions 2 reactants join to form 1 product 2. Substitution reactions 2 reactants combine to form 2 products 3. Oxidation In organic chemistry this is to do with the addition of oxygen or the loss of hydrogen. When writing out the reactions, write oxygen as [O] and not O2 because the oxidising agent is not oxygen from the air. 4. Reduction The addition of hydrogen or loss of oxygen. 20  5. Homolytic fission Breaking of a covalent molecule where the shared pair of electrons is divided equally between the atoms. 6. Heterolytic fission Breaking of a covalent molecule where the shared pair of electrons is divided unequally between the atoms. Heterolytic fission usually occurs when there is a difference in electronegativity between the atoms (the one with higher electronegativity keeps the electrons) Note: the curly arrows show the movement of electrons, they are not always necessary in showing these reactions (question will specify). Electrophiles are chemical species that like negative charges/are attracted to negative charges or a region of high electron density (δ-) Tip: All positive ions are electrophiles but not all electrophiles are positive ions. Free radical substitution reactions of alkanes You need to memorise all the steps for the mechanism of substitution and be able to draw them out. (Example) The chlorination of methane 1. Initiation UV light breaks the chlorine molecule into chlorine free radicals (a species with an unpaired electron) through homolytic fission. 2. Propagation A free radical is extremely reactive. Chlorine free radicals collide/attack the methane molecules and remove a hydrogen to form a methyl free radical and hydrogen chloride. The methyl free radical attacks a chlorine molecule, making chloromethane and reforming a chlorine free radical. 21  3. Termination The free radicals in step 2 keep forming and the reaction is a cycle until two reactive species turn to unreactive species. There are different possibilities depending on the free radicals formed. Further substitution o It is hard to prevent these reactions from occurring and because of that the desired product is not always formed. This reaction has a low atom economy (need to separate desired product) and a low percentage yield. Alkenes Alkenes are unsaturated hydrocarbons (double bond is present) so they are very reactive. The carbon-carbon double bond Alkenes contain one σ bond and one π bond. Note: electrons in π bonds are further from the carbon atoms than those in σ bonds so they are more likely to undergo reactions. 22  Geometric isomerism Geometric isomers differ because their groups are attached at different positions on opposite sides of the C=C bond. Note: The C=C bond means there is restricted rotation so groups can only be in one of 2 positions. Cis and trans naming system: Cis-but-2-ene The CH3 groups are both on the same end of the molecule (both on top) Trans-but-2-ene The CH3 groups are on the opposite ends of the molecule. The cis-trans naming system only works for some molecules and doesn’t work when there are different groups attached on the C=C bond, so we instead use the E-Z naming system for more complex molecules. E-Z naming system: 1 2 Steps to name the molecules: 1. Figure out the part of the name that is used for both (1-bromo-1-chloro-2-fluoroethene) 2. Use priority rules to decide which atoms have higher priority for each side. Higher priority means higher atomic number. F > H and Br > Cl 3. Where are the atoms with higher priority in relation to each other? In the first molecule, Fluorine and Bromine are both below so it is the Z isomer. In the second molecule they are across from each other so it is the E isomer. Tip: E for enemies 23 Z for friendZ ( ) Skeletal formulae Z isomer E isomer Addition reactions of alkenes: Notes: o Curly arrows used in this reaction must start from a bond and move to an atom or start from a lone pair and move to an atom. o Alkenes have pi bonds so they have an area of high electron density around those bonds. 1. Hydrogenation Adding hydrogen Forms an alkane Catalysed by nickel 2. Halogenation Adding a halogen Forms dihalogenoalkanes 3. Hydration Adding steam Forms alcohol – adds OH and H to the C=C bond Catalysed by H3PO4 4. Addition of hydrogen halide Forms a halogenoalkane 5. Oxidation to diols Involves both addition (of water) and oxidation 24  The oxidising agent is potassium manganate (VII) Just write [O] in the equation – no need to mention the agent Can be used as a test for alkenes because it turns from purple to colourless – alkanes do not react this way Electrophilic addition of hydrogen halides Hydrogen bromide is a polar molecule and is attracted to the pi bond in alkenes because it is electrophilic. An electrophile attacks centres of negative charge (pi bonds) (Example) 1. In the first step, the hydrogen bromide molecule is separated by heterolytic fission and the electrons go to bromine (more electronegative) The hydrogen is partially positive and the bromide is partially negative. A carbocation intermediate (because the positive hydrogen ion bonds to it) and a bromide negative ion form. 2. In the second step the bromide ion is attracted to the carbocation and forms a covalent bond with it. 25 Electrophilic addition of halogens Bromine by nature is not a polar molecule but the electrons in the pi bond in C=C repel the electrons in bromine, inducing polarity in it. Asymmetrical molecules When the atoms on either side of the C=C bond are different there are multiple possible products cations (and therefore products) that can be formed. A is a secondary carbocation and B is a primary carbocation. The carbocation A is more stable than B because the positive charge is spread to more alkyl groups, so it is formed more often and it is what forms the major product. Sometimes there is a tertiary carbocation in which case it is the most stable and forms the major products. Addition Polymerisation Same as IGCSE good luck ‫توتة توتة خلصت الحدوتة‬ 26 27