As Chemistry Unit 1 Notes (edexcel) b486z

Chapter 1.1 Formulae, equations and amounts of substances

THE FOUNDATIONS OF CHEMISTRY: What’s in an atom?  The nuclei of atoms has protons (+) and neutrons (0) [called nucleons]  Electrons (-) occupy empty space outside the nucleus  Element: substance that cannot be broken down chemically into simpler substances  Atomic number: number of protons of an atom of the element  Number of protons is equal to the number of electrons, so the atom is electrically neutral  Protons and neutrons have mass but electron’s mass is negligible  Mass number: number of protons + number of neutrons Isotopes:  Atoms with the same atomic number but a different numbers of neutrons are called isotopes  Some heavy isotopes are radioactive  They may have different physical properties but the same chemical properties as the number of electrons is the same Holding atoms together:  Bonding between metals and non-metals is ionic  An atom loses or gains an electron to have a complete outer shell, this makes the atom, an ion  Strong electrostatic forces hold the oppositely charged ions in a giant lattice  Covalent bonding: non-metals bond together by sharing pairs of electrons Ions in solutions:  Ions that are not involved in an equation are called spectator ions  Molecular equation shows complete formula of every substance Relative atomic mass:  The average mass of its isotopes compared with the mass of an atom of the carbon-12 isotope  It is an average of the differing isotopes Counting and Weighing atoms:  Mole: the amount of substance that contains as many particles (atoms, ions or molecules) as there are atoms in exactly 12g of carbon 12  Any mole of a substance contains 6.23x10^23 particles (Avogadro’s number)  Avogadro’s constant has the unit particles per mole (mol-1) and sometimes has the symbol L or NA

Relative formula mass and relative molecular mass:  The sum of the relative atomic masses of all the atoms within a chemical formula  Molar mass: the RMM or RFM in grams per mole; it tells you the number of grams of a substance that makes up one mole Molar Volume:  Molar volume (of the gas) [Vm]: one mole of any substance must occupy the same volume under the same conditions  STP: standard temperature and pressure [1 atmospheric pressure and 298K] is used to compare molar volumes  Under STP: 1 mole of any gas occupies 24dm3 The empirical formula:  Gives the ratio of the different atoms present; it is the simplest formula for a compound showing the whole number ratio of each atom of an element present  Does not say how many atoms are bound together in one molecule of a compound Moles in solution:  Solution: solute that has dissolved in a solvent  Molar solution: solution of concentration 1M  Concentration [molarity] is measured in mol dm-3 Other units of concentration Percentage by mass:  Percentage by mass = mass of solute/mass of solution x 100 Percentage by volume:  Percentage by volume = volume of one component/total volume x 100 Parts per million:  Concentration= mass of component/mass of solution x 1 000 000 ppm  Used for levels of pollution in air or water; used when percentages are not very useful The yield of a reaction:  Yields may not be 100% as: reactants may not continue until all reactants are used up (may be an equilibrium reaction); some product may be left on apparatus; volatile products may evaporate; human error Double salts:  Crystals that contain two different salts in a 1:1 ratio Atom economy:  Atom economy: mass of atoms in desired product/mass of atoms in reactants x 100%

Chapter Two: Energetics and Enthalpy Change Energy and Energetics:  Energetics: study of energy transfers between chemicals and their surroundings  Thermochemistry: study of these energy transfers Exothermic and Endothermic:  Endothermic reaction: reaction that needs energy from heating to take place  Exothermic reaction: reaction that releases energy  Bond breaking requires energy whilst bond making releases energy Energy changes in exothermic reaction:  Energy released by bond formation is greater than the energy needed to break the bonds in the reactant  Neutralisation reactions are exothermic Energy changes in endothermic reactions:  The energy required to break the bond is greater than the energy released when new bonds are formed in the product  Photosynthesis and thermal decomposition is endothermic Enthalpy changes:  An exothermic reaction releases energy to the surrounding in the form of heat (enthalpy change of reaction)  System: the reaction in which the changes are happening  Surrounding: everything outside the system  Boundary: separates system from the surroundings  Closed boundary: prevents particles leaving or entering the surroundings  Isolated boundary: prevents energy from leaving or entering the surroundings  Energy cannot be created or destroyed  Energy is transferred and spread through the surroundings (dissipated)  Principle of conservation of energy: total energy content of the universe is constant How much energy is transferred?  Enthalpy (H): energy content in a system held at a constant pressure  Enthalpy change (H): the amount of heat given off or absorbed in a reaction carried out at constant pressure  H = Hproducts – Hreactants Enthalpy level diagrams: In an exothermic reaction – the enthalpy change is negative as H[products]
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The energy that the system loses is given out as heat and the surroundings warm up (there is a net decrease in the potential energy of the system) The temperature of the surroundings is lower than the temperature of the products so energy is transferred to the surroundings as products cool

In an endothermic reaction, H is positive  Energy content of system has increased as it has gained energy from the surroundings  Energy needed to break the bonds in the reactants is more than the energy needed to form the bonds in the reactants  Energy is transferred from the surroundings and there is a net increase in potential energy of the system The temperature of the surroundings is higher than the temperature of the products, so energy is transferred to the products

Heat capacities and Calorimeters:  The heat capacity (C): of a body of liquid or air is the amount of energy required to raise its temperature by 1K  The specific heat capacity: amount of energy in joules required to raise the temperature of 1kg of a substance by 1K  Energy transferred (J) = mass (kg) x specific heat capacity (J kg-1 K-1) x temperature change (K) The bomb calorimeter:  A calorimeter used to determine the energy change during a reaction accurately  Useful for studying the enthalpy changes when fuel burns and to find calorific value  Heat is transferred to the surroundings: makes it slightly inaccurate and unreliable (however heat loss can be calculated and corrections can be calculated and compensated)  You can use an electrical heater to create the exact same temperature change in the calorimeter  So electrical needed to bring about same temperature change can be calculated  Tells you more accurately the energy change that occurred the reaction as measured energy change duplicates any heat losses form the calorimeter Important standard enthalpy changes:  Standard enthalpy changes refer to reactions done under standard conditions and everything in their standard states Standard enthalpy of combustion (Hөc)  When one mole of a substance is burnt completely in oxygen under standard conditions  Complete combustion of carbon only forms carbon dioxide and NOT carbon monoxide so the enthalpy of combustion is the first process  C(s) + ½O2(g)  CO(g)  but C(s) + O2(g)  CO2(g)   The standard enthalpy changes of combustion varies due to the number of bonds that need to be made and broken and the types of bonds involved  One mole of hexane will release more energy than one mole of methane as it has more as there is more bond breaking and more bonding forming hexane  Energy released from combustion is due to bond making with the oxygen so the more oxygen a fuel has in its molecule, the less energy it will give out  Bonds between different elements have different bond energies so chemical make-up of the fuel and the bonds will affect the enthalpy of combustion Standard enthalpy of formation (Hөf)

The enthalpy change when one mole of the compound is formed form its elements under standard conditions  The units for enthalpy of formation is kJ mol-1  Enthalpy of formation for en element is 0 Standard enthalpy change of atomisation (Hөa)  Enthalpy change when one mole of its atoms in the gaseous state is formed from the element under standard conditions  Endothermic as it involved increasing the separation between atoms (requires energy)  May be measured by spectroscopic means in cases of gases Standard enthalpy of neutralisation (Hөn):  Enthalpy change when one mole of acid is neutralised by an alkali in their standard states at 25 degrees and in solutions containing 1 mol dm-3  Enthalpy of dilute acids with strong bases and strong alkali will always be close ot -58kJ mol-1 as they are fully ionised Hess’s law:  Total enthalpy change for a reaction is independent of the route taken Bond enthalpies and mean bond enthalpies:  Bond dissociation enthalpy: the energy needed to break a particular covalent bond or the energy released when the bond is formed  Same type of bonds in different compounds can have slightly different bond enthalpies depending on the combinations of other atoms in the molecule and their effect on the bonds Mean bond enthalpy:  The mean value of the bond dissociation enthalpy of a particular type of bond over a wide range of different compounds  Apply to gaseous states so can lead to inaccuracy using a different state  They allow you to compare the strengths of bonds in different atoms  Allow you to estimate the enthalpy changes in reactions  Help in building up an understanding of the structure and bonding of compounds  Develop an understanding of the mechanisms of chemical reactions  They can be used to predict which bond will break first and how easy it is to break the bond  Bonds with high bond enthalpy will need more energy so it is likely it will not break first (and vice versa)  Bonds with very low bond enthalpies could even take place at room temperature Stability: 

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The lower down the energy diagram something is, the more energetically stable it is Bond breaking takes energy and there is a minimum amount of energy needed before a reaction can start: this is called the activation energy The mixture can be kinetically stable even though it may be energetically (thermodynamically) unstable

CHAPTER THREE - ATOMIC STRUCTURE AND THE PERIODIC TABLE   

RAM: of an element is the average mass of the isotopes compared with the mass of an atom of Carbon- 12 Atomic mass unit: mass of a carbon 12 divided by 12 Relative formula mass: sum of RAM of all the atoms in the chemical formula (referred to as RMM in molecular compounds)

ELECTRONS IN THE ATOM       

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MEASURING THE MASS OF AN ATOM  1. 2.

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To find the RAM, you need to measure its mass and compare it to the mass of a carbon 12 atom VAPOURISATION: Sample must be in a gaseous state to move through the machine. It is vaporised and then injected IONISATION: Bombarded with a beam of high energy electrons which collide with the atoms of the sample. Knocks of electrons to form positively charged ions. Allows them to be accelerated in an electric field ACCELERATION: Accelerated using an electric field DEFLECTION: es through velocity selector (all will be the same velocity). Any difference in effect of magnetic field would be due to charge/mass not speed. Enter a uniform magnetic field which deflects. Heavier and small positive charged ions are deflected less than small, more positively charged ions. Strength of magnetic field is steadily increased. At particular settings, only ions of one particular mass:charge ratio will through – other ions will be deflected too little or too much to go through DETECTOR: Measures how many ions through at each different magnetic field and velocity selector setting. Shows how many ions of each mass:charge ratio there are Results are obtained in mass spectrum. Relative heights of peaks show relative abundance of different ions present The peaks on the mass spectrum of an element show the different isotopes of the elements

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Only a limited number of energy changes or transitions can take place within the atom Electrons are arranged in shells (name of shell = principal quantum number [n]) tells us about the size of the shell. Larger the shell = further away from nucleus Ionisation: complete removal of an electron from an atom thermic as work has to be done to overcome attractive between it and the nucleus) Ionisation energy: amount of energy needed to remove an electron from an atom Can be measured by gradually increasing voltage applied to a gas until it conducts electricity and emits light (tells electron has been freed) Ground state: an atom that is at its lowest energy level First ionisation energy: the amount of energy required to move the first electron in the outermost shell. The less tightly it is bounded, the lower the IE will be (thus more reactive) Each shell may contain a number of subshells (s,p,d,f,g) Shell Subshell 1 1s 2 2s, 2p 3 3s, 3p, 3d 4 4s, 4p, 4d, 4f Shell 1 is the closest to the nucleus so will need the most energy to remove electron. Subshells have different energies (electrons with lowest energy will be closest to nucleus) : s (lowest energy) < p < d Subshells contain orbitals: the regions where the electrons are most likely to be found As the number of shells increase, the energy gap between successive shells gets smaller. Thus orbitals in neighbouring shells may overlap: e.g. 3d has an energy level above that of 4s but lower than 4p Atom is in ground state when electrons are in the orbitals with the lowest possible energy levels Electron spin: a moving charge creates a magnetic field (electron can spin either clockwise or anticlockwise). Two electrons in the same orbital cannot have the same spin (thus orbital can only have 2 electrons with opposite spin) Hund’s rule: when electrons are placed within a set of orbitals with equal energy, they spread out to maximise the number of unpaired electrons. Electrons will fill the lowest energy orbitals first, then the remaining orbitals in order of increasing energy Nodes: Areas where the electron wave has zero amplitude thus the electron density is zero (electrons will not be present) S orbitals are spherical; P orbitals are dumbbell shaped and lie at ninety degrees to each other. Size of orbitals increase with principal quantum number

ELECTRONIC CONFIGURATION AND CHEMICAL PROPERTIES  Hydrogen does not fit naturally into any group so is isolated at the top  All elements in a period have the same number of electron shells The s-block elements:  Outermost electron is in the s subshell  Lose electrons easily to form positive ions. Includes reactive metals which can form stable ionic compounds with non metals  Lower boiling and melting point and lower density than most other metals  Conduct electricity due to metallic bonding resulting in sea of delocalised electrons  Helium and Hydrogen are also s block. But the characteristics of them are so different to the other s-block elements that they are put in a separate group The d-block elements:  Contains elements with successive electrons being added to the d- subshell  Transition metals (but zinc does not fit this description/ and mercury)  Much less reactive as outer subshell s is full whilst d subshell is being filled  Can conduct heat and electricity; shiny, hard and malleable # The f-block elements:  Successive electrons are being added into the f subshell  Top row (lanthanides)= very similar metals; bottom row (actinides) = radioactive [actinides up to uranium are naturally occurring; others are artificial and are very unstable – short half-lives) The p-block elements:  Groups 3,4,5,6,7 and 8. Electrons are being added to p subshell in outer shell  Metals: (e.g. tine and lead) – form positive ions and ionic compounds with non-metals. Can conduct heat and electricity, but do not have strong metallic characteristics. Post transition metals are relatively unreactive  Metalloids: diagonal block; act like non-metals but can conduct electricity (albeit poor)  Non-metals: covalent with non-metals and ionic with metals. Majority do not conduct electricity but some do. Some form giant covalent structures and many exist as small molecules

PATTERNS IN THE PERIODIC TABLE  Periodicity: repeating trends of properties of elements in the periodic table Trends in the periodic table  Atomic radius increases going down a group due to extra electron shells being added, despite it having extra protons, the electrons are further away and are increasingly screened from the nucleus (shielding increases thus nuclear charge decreases)

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Across the period, atomic radius generally decreases: nuclear charge becomes increasingly positive as number of protons increase; the number of electrons also increase but they are all in the same shell so they are more strongly attracted to the nucleus  Positive ions have a smaller radius as loss of electron(s) means that remaining electrons will have a greater share of positive charge from nucleus thus are more tightly bounded. In addition, when a ion is formed, the whole shell is lost  Negatively charged ions: even though additional electrons are within the same shell, there is extra negative charge thus electrons will be less tightly bounded to the nucleus so the radius will be larger Periodic trends in ionisation energy:  The more tightly held the outer electrons; the higher the first ionisation energy  The further the outermost electron is away from the nucleus (atomic radius), the smaller the attraction it will have thus the ionisation energy decreases  A more positive charge nucleus will have a greater attraction to the outermost electron which results in a high ionisation energy  Inner shells of electrons repel the outermost electron shielding it away from the nucleus, the further (and more electron shell) the outermost electron is from the nucleus, the less attraction it will have to the electron so a lower ionisation energy  Ionisation increases as you go across a period due to increasing positive nuclear charge but no addition of electron shells to screen the outer shells  Atomic radius decreases as electrons are held together more tightly so ionisation energy will be high Ionisation energy trends of groups:  Group 8: have high ionisation energy due to stable electronic structure (has a full outer shell) which makes them very unreactive  Group 3: The increased distance results in a reduced attraction and so a reduced ionisation energy; The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy  Group 6: There are paired electrons in the subshell as opposed to a half full subshell; the repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be Patterns in physical properties:  The melting temperature: is the temperature where the pure solid is at equilibrium with the pure liquid at atmospheric pressure  High melting point of metals are due to their metallic structure, atoms are held together tightly by a sea of negative electrons

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Giant molecular structure (such as carbon and silicon) had high melting points as they are held in by covalent bonds in a crystal structure; it is very difficult to remove individual atoms  Non –metals found on the right often have simple molecular structure so are found as simple, individual molecules. The atoms are ed by strong covalent bonds but the molecules themselves are bonded by weak intermolecular forces: so they can be separated easily resulting in a low melting temperature Chapter Four: Bonding What is a chemical bond?  The forces holding an atom together Ionic bonding:  Bonding between metals and non-metals  Forms oppositely charged ions which are held together by strong electrostatic forces of attraction  Ions are held together in a giant ionic lattice which can then form ionic crystals  Can be thought of as the net electrostatic attraction between the ions  Lattice structure of ionic compound is the arrangement of ions in a way that maximises the attractive forces of the oppositely charged ions and minimises the repulsion between the similarly charged ions  The forces that are exerted by the ions act equally in all directions and so hold the ions tightly  Isoelectronic: when an ion gains or loses electron(s) to have the same electronic configuration of a noble gas (e.g. Na+ and Ne are isoelectronic)  Transfer of electrons forms the ionic (electrovalent) bond#  Octet rule: when elements react; they tend to do so in a way that results in an outer shell containing eight electrons Giant ionic lattices:  High melting point  Can dissolve in water  Can conduct electricity when molten or in aqueous solution Trends in ionic radii:  Is the radius of an ion in a crystal  Radius of positive ion is smaller than atom of the element as remaining electrons are more strongly attracted to the positive nucleus; negative ions are larger as the extra electrons mean they are less bounded to the nucleus Types of lattice structure:  Electro n density map: show the exact arrangement of ions in an ionic lattice which varies depending on the relative sizes of the different ions presen  The coordination number: number of nearest neighbouring ions

 Face centred cubic structure and body centred cubic structure Lattice energy:  The enthalpy of formation of one mole of an ionic compound from gaseous ions under standard conditions  This is a bond forming process so the enthalpy change is negative as it is exothermic What affects lattice energies?  The lattice energy becomes less negative as the size of the ions increase. In smaller ions, the attractive force of the positive nucleus holds the outer electrons more tightly because they are closer to the nucleus  Lattice energy becomes more negative with ions of greater charge as they are attracted to other ions more strongly  Increasing the charge on either ion increases the attractive force between two oppositely charged ions in the lattice  Decreasing the size of one or both ions decreases the distance between them and so increases the attractive force Polarisation in ionic bonds:  Attraction of the positive cation for the outer electons of the negative cations. Electron cloud of anion is attracted to the positive cation  If distortion is great, it will lead to a charge cloud which is similar to a covalent bond  Polarising power depends on charge density (so ionic charge and radius)  Cations with small ionic are much more polarising as they have a higher charge density (effect of positive nucleus is felt more strongly)  Cations with large positive charge are more polarising than cations with a smaller charge as they have a stronger attraction for the outer shell  The larger the anion, the more easily it is polarised  Theoretical models always assume that the ions are spherical and separate and electron charge is evenly distributed across the ion  Polarisation of bond, distortion in the shape of the ion and increasing covalent character in bonds reduce the lattice energy Covalent Bonds:  A shared pair of electrons; electron density exerts and attractive force on each nucleus  Between two non metals  Sharing two pairs of electrons results in a double bonds; three pairs is a triple bond Dative Covalent bonds:  Sometimes the pair of electrons that make up the covalent bonds come from the same atom  Compounds that have unshared electron pairs readily to form dative covalent bonds  Dative covalent bonds are the same length and strength as any other covalent bonds Metallic Bonding:

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Good conductors of heat and electricity Metal crystal: positive metal ions surrounded by a sea of delocalised electrons, there is a strong attraction between positive metal ions and electrons

Chapter Five: Introductory organic chemistry: Hazards and Risks: - Hazard presented by a substance or an activity is its potential to do harm (potential is absolute) - Some chemicals are flammable and some are toxic - Tendency to burn or to poison are texted and calculated and will always be the same Risk: - Risk associated with a particular hazard is the chance that it will actually to cause harm - Affected by nature of hazard involved and the level of exposure to it - Level of exposure is dependent on the expertise of the person working with the chemical, volume being used, the conditions its used in and the protective clothing and equipment available Ways of reducing risk: - Working with smaller amounts: when quantities are smaller, it is easier to contain the reaction in closed apparatus - If heat is given out, the smaller the quantity, the less heat given out - Smaller quantities are also easier to transfer from one equipment to another (e.g. less chance of spillage) - Taking specific precautions or using alternative techniques depending on the properties of the hazardous substances you are using - E.g. using a low concentration as possible (at low concentrations, the chemicals are more likely to be irritant than corrosive so hazard is lowered) - If hazard risk is high, the risk can be kept low by careful planning and risk assessment - Careful useful of safety measures: fume cupboards to remove toxic or flammable fumes; personal protection such as safety googles

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Changing the conditions in which a reaction takes place: e.g. lowering temperature slows down the rate of reaction, this would reduce risks of reaction mixture overheating and producing fumes; although the cooling may change equilibrium position and will affect proportion of reactants and products in the final mixture - Using alternative methods with less hazardous substances: possible to substitute chemicals and still study the same basic reaction, but they may not be as effective as original one but originals are far more hazardous. For example, cyclohexane can be used instead of tetrachloromethane, however plastic screw tops or bungs will be affected if it is stored for over a month Why does carbon form four covalent bonds?  Arrangement of the four bonds around the carbon atom is nearly tetrahedral  Electronic configuration indicates that it needs four electrons to reach a noble gas configuration  Once bonds are made, the outer electrons form a fully shared octet (inner shell complete, no lone pairs or empty orbitals so it does not need to with any other atom)  Can form strong covalent bonds with itself and other elements such as hydrogen Organic families:  Aliphatic molecules: straight or branched chain carbon skeleton  Alicyclic molecules: closed rings of carbon atoms which may contain C=C bonds  Arene: derived from benzene molecules and contain a benzene ring with six carbons in structure  Functional group: atom or group of atoms which is typical for a particular organic family and which determines the same chemical properties of the molecule  Functional group is affected by the environment (e.g. as hydrocarbon gets bigger, it has increasing effect on the chemistry of the molecule and influence of functional group decreases) The different functional groups:

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General formula: describes the number of carbons and its relationships to other atoms  Functional group, shape and size of carbon chain affects how it reacts The shapes of molecules:  Isomers: when two or more compounds have the same molecular formula but the atoms are connected together differently  Structural isomerism: remain part of the same homologous series although boiling points may vary  Brought about by the different shapes of molecules which affect the intermolecular forces between the molecules due to the way in which the molecules can pack  Stereoisomerism: when 3D arrangements of the bonds in a molecule allow different possible orientations in space The Alkanes:  Hydrocarbons: group of organic compounds containing several homologous series (aliphatic; alicyclic and arenes)  All insoluble in water and burn completely in oxygen to product CO2 and H20 General properties:  Are saturated: so have no double or triple bonds between carbon atoms so contain the max possible number of hydrogen atoms  Can occur as both and straight chained carbons Where are they obtained from?  Found in fossil fuels alongside with other product of decomposition (methane)  Plants and animals died and sank to the bottom, deep layers of decomposing material formed which became encapsulated in rock  Over time it was exposed to extreme heat and pressure to form crude oil/petroleum Using crude oil:  Crude oil is extracted from the ground.  Primary distillation is the basic process by which petroleum is turned to a valuable liquid into useful chemicals with whole variety of properties  Liquid collected over each range of temperature is called a fraction 1. 1-2% refinery gas: lightest fraction of crude oil with the lowest boiling point ; contains gaseous alkanes; used as a fuel 2. 15-30% gasoline: mixture of straight chained or branched liquid hydrocarbons; used for engine combustion

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10-15% kerosene: mainly C11 or C12 – used for aircraft fuel, can be cracked down to more useful hydrocarbons 4. 15-20% diesel oil; gas oil: used in industrial boilers, can be cracked down using catalytic cracking 5. 405-% residue: complex mixtures of hydrocarbons; very viscous and high boiling point; can be used as fuel for power stations or cracked further to produce lubricating oils Different fractions in demand:  There is a higher demand for short chain hydrocarbons than long chain hydrocarbons  Heating long chain alkanes can cause the molecule to split into a shorter chain molecule  Alkanes are very unreactive, so cracking them to make alkenes (reactive) is very efficient  Using high temperatures to break it down is very expensive so often a catalyst is used (catalytic cracking)  They are non corrosive with metals (good lubricating oils) – harmless to skin  At room temperature, they are not affected by concentrated mineral acids or alkali  They are no affected by oxidising agents or the most reactive metals  As C-C and C-H bonds evenly share electrons since electronegative of carbon and hydrogen are very close so the bonds are not polar to any extent so there are no charges to attract other polar or ionic species  Formation of alkanes are due to free radicals which have unpaired electrons; this means they have very high activation energy but once this is overcome, they proceed to react very rapidly in the gas phase Breaking bonds:  Breaking bond is called bond fission  When the bonds are broken, there are two ways in which the paired electrons shared between the two atoms can go Homolytic fission:  Equal sharing out of the electrons in the bonds, so each of the participants in the bond receives one electron when the bond splits  Despite neither atoms having overall charge, the free radicals with their unpaired are extremely unreactive as the unpaired electron has a strong tendency to pair up with another electron from another substance  Mainly occurs when there is little or no ionic character in the covalent bond Heterolytic fission:  Involves an unequal sharing of the electrons in the covalent bond so that both electrons go to one electron  Results in two charged particles, the atoms which gain the electrons become negatively charged and the atoms which loses the electrons become positively charged  This happens when there’s a degree of polarity in the covalent bond Heating Alkanes: Cracking:

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When alkanes are heated in high temperatures in the absence of oxygen, they split into smaller molecules (C-C bonds are broken and form C=C bonds – it is a thermal decomposition reaction)  Cracking of methane produces powdered carbon which is used in car tyres… etc Combustion:  Will combust to produce water and carbon dioxide, they are very energetically unstable  Will only burn in their gaseous state so liquids are less flammable (must be vapourised)  An exothermic reaction: energy needed to break the bonds in the chemical reaction of combustion is less than the energy returned when the new bonds are formed in the products of combustion  Methane, propane and butane are all commonly used fuels, many areas, natural gas is piped into homes  Propane canisters can be supplied; great advantage is that it can be readily liquefied under pressure and at low temperatures so that a large quantity of gas can be stored in a small space  It will not burn until it is a gas so it can be stored and transported relatively easily and safely