Chemistry Yr11 Syllabus Notes PDF

Chemistry Yr11 Syllabus Notes Scientific terminology: Module 1- Properties and Structure of Matter Properties of Matter explore homogeneous mixtures and heterogeneous mixtures through practical investigations: – using separation techniques based on physical properties (ACSCH026) – calculating p...

Chemistry Yr11 Syllabus Notes Scientific terminology: Module 1- Properties and Structure of Matter Properties of Matter explore homogeneous mixtures and heterogeneous mixtures through practical investigations: – using separation techniques based on physical properties (ACSCH026) – calculating percentage composition by weight of component elements and/or compounds (ACSCH007) Define/describe diagram homogenous Have a uniform composition through-out e.g., air, sugar waters heterogenous Have a non-uniform composition through- out – able to recognise different particles in mixture e.g., pasta sauce, sand and water, trail mix Filtration -Technique that separate solids from liquids -relies on difference in state and size of components -pass mixtures through a filter with holes or pores of a particular size – allows particles to pass through while others cannot -e.g., separating sandy water  Evaporation -used to separate fluids or soluble substances from a solution -relies on difference in boiling points e.g., wet clothes drying in the sun, drying hair with a hair dryer Distillation -boiling points of different liquids -one evaporates and then condenses in a cooler area to be collected decantation -pouring away a liquid while leaving the solid behind -relies on density Percentage Composition Formula: Mass of component in sample (g) *100 Total mass of sample (g) Nomenclature (IUPAC) Ionic Compounds 1. Cations named first, followed by the 2nd 2. Monatomic anions are followed by “-ide” 3. Polyatomic anions have unique names 4. Cations – indicate oxidation state in roman numerals (e.g. iron (II)) Covalent Compounds 1. More electronegative element is named lasts 2. Last element is followed by “-ide” 3. Prefix denoting the number of atoms (e.g. monoxide, dichloride) investigate the nomenclature of inorganic substances using International Union of Pure and Applied Chemistry (IUPAC) naming conventions -classify the elements based on their properties and position in the periodic table through their: – physical properties – chemical properties Periodic table -elements listed in ascending order of organisation atomic number -divided into rows (periods) and columns (groups)- elements that behave similarly tend to be grouped together -periods correspond to the number of electrons shells of element -groups corresponds to number of valence electrons (and reactivity) Metals -mainly on left of periodic table -form cations, metallic luster, good conductors, malleable and ductile, high melting point, mostly solid at room temp. Non-metals -mainly on right of periodic table (except Hydrogen) -forms anions, dull, colourless to colourful, poor conductors, brittle, low melting point, often liquids or gases at room temperature Atomic Structure and Atomic Mass Diagram Position Isotopein periodic description - An atom of an element with different numbers of neutrons tables -some have unstable nuclei and undergo radioactive decay for Distribution of -Atomicthe number >82 nucleus to are all unstable become – more stable electrons, protons, and too heavy neutrons -Atomic number as volume increases, moles increase Molar volume formula: n1 = V/Vm (the number of moles of gas in a sample) = (to the amount of space the gas takes up) / (the amount of space 1 mole of gas takes up) Standard conditions:  Standard temperature and Pressure (STP) o Temperature: 0°C (or 273.15K) o Pressure: 100kPa o Molar volume: 22.71 L  Standard Laboratory Conditions (SLC/RTP) o Temperature: 25°C (or 281.15K) o Pressure: 100kPa o Molar volume: 24.79 L Gas Stichometry:  Volume ratios are the same as Molar ratios for gases at the same condition Module 3- Reactive Chemistry Chemical Reactions investigate a variety of reactions to identify possible indicators of a chemical change physical change vs chemical change:  creation of a new substance ONLY in chemical change  physical  involves change in state/shape Indicators of Chemical Change: Things caused by chemical bonds being broken and formed  Change in temperature  Change in colour   A noticeable odour  Precipitate  Bubbles (formation of a gas) use modelling to demonstrate – the rearrangement of atoms to form new substances – the conservation of atoms in a chemical reaction (ACSCH042, ACSCH080) Modelling Chemical Reactions:  Reinforces that chemical reactions involve the rearrangement of atoms and adhere to the law of conservation conduct investigations to predict and identify the products of a range of reactions, for example: synthesis:  The combination of multiple reactant to form a single product  A + B + C  ABC  E.g. 2H2(g) + O2(g)  2H2O(l)  Release energy (exothermic) decomposition  Involves a single compound breaking down (decomposing) into multiple products  Breaking of bonds – requires energy (endothermic)  AB  A + B  Example: 2H2O(l)  2H2(g) + O2(g) combustion  Hydrocarbon + oxygen  carbon dioxide + water  Produce energy (exothermic)  Complete combustion: hydrocarbons all burn in sufficient air or oxygen to produce carbon dioxide and water  Incomplete combustion: occurs in low oxygen environments where carbon monoxide or soor (carbon) may result precipitation   Involve the formation of a precipitate (insoluble salt) from two aqueous solutions  Reactants are both soluble  Cloudiness of a solution indicates a precipitate has formed (solid suspended in solution)  Anions and cations switch partners  Double replacement reactions  Apply solubility rules  AB + CD  CB + AD acid/base reactions  Neutralisation reactions  Acid + base  salt + water  Acids are substances that produce H+ ions  Bases are substances that produce OH- ions (most of these are hydroxides) acid/carbonate reactions (ACSCH042, ACSCH080)  Acid + carbonate  salt + carbon dioxide + water  Carbonates are a group of ionic salts containing CO32 investigate the chemical processes that occur when Aboriginal and Torres Strait Islander Peoples detoxify poisonous food items  Detoxification of Cycad Fruit o Cycad fruit is highly toxic and carcinogenic o Leaching   Grinding or pounding of kernels  Increases surface area of fruit and increases toxin removal in later stages  Moved into a mesh-like bag and soaked in water (for at least several days) since many toxins in fruit were water-soluble and would leech out of fruit and into water  Further ground and then baked as damper  Relays on physical and chemical processes – disillusion of toxins, grounding, etc. o Fermentation   Stored for months in moist environment (often in a trench)  When fruit was frothy and mouldy, the toxins were no longer present  o Actively cultivated these plants with fire to increase production  Detoxification of Black beans o Black beans are highly toxic o Seeds were collected and heated in fire which breaks down toxins and allowed easy removal of the peel o Seeds were scraped – increased surface area and facilitates leaching o Placed in dilly bag and left to soak in water – removes soluble toxins o Mash further ground and then baked as dampers construct balanced equations to represent chemical reactions Predicting Reactions of Metals conduct practical investigations to compare the reactivity of a variety of metals in: – water – dilute acid (ACSCH032, ACSCH037) – oxygen – other metal ions in solution Reaction of Metals and Oxygen:  All metals except silver, platinum and gold react with oxygen to form oxides  Metal + Oxygen  Metal oxide Reaction of Metals and Water:  Some metals react with water or steam while others do not  Metal + water  Metal hydroxide + Hydrogen Gas Reaction of Metals and Acid:  Metals can react with acid to form a metal salt and hydrogen gas  Metal + dilute acid  metal salt + hydrogen gas Vigour of Reaction: A more reactive metal loses electrons more readily and will react more vigorously.  this is effected by other factors e.g., temperature, surface area  a more vigorous metal is more reactive Practical Investigation Aim: Compare the reactivity of a variety of metals and construct a metal activity series. Risk Assessment Risk: Broken glassware Assess: Moderate Control: Handle glassware with caution Risk: Bunsen burner Assess: Moderate Control: Handle Bunsen burner with caution Risk: Dilute Acid Assess: low control: wear protective clothing (eyewear, lab coat) Risk: High energy metal reactions Assess: Moderate control: allow teacher to conduct reactions behind safety screen Method a) Reaction of metals and water 1. A similar sized sample of each metal was placed in a test tube 2. Water was added to each test tube and observations were recorded 3. If no reaction was observed, steps 1-2 were repeated at a higher temperature in a hot water bath b) Reaction of metals and 0.1M Hydrochloric acid 1. A sample of each metal was placed in a sperate test tube 2. 2-3mL of dilute acid was added to each test tube and observations were recorded c) Reaction of metals and oxygen For Ca, Mg, Al, Zn, Fe, Sn, Pb, Cu, Ag – 1. samples of each metal were heated directly by a Bunsen burner 2. Observations were recorded For Na and K - 1. Samples of each metal were cut to expose the non-oxidised metal surface 2. Observations were recorded Results a) Reaction of metals and water  b) Reaction of metals and 0.1M Hydrochloric acid c) Reaction of metals and oxygen Metals that react vigorously with one reactant also tend to react vigorously with others. construct a metal activity series using the data obtained from practical investigations and compare this series with that obtained from standard secondary-sourced information (ACSCH103) Metal activity series: A list that orders metals from most active to least active. - used to predict displacement reactions Limitations: 1) relies on qualitative observations – these terms don’t allow us to make accurate comparison of different reactions 2) reaction rate was impacted some conditions e.g., surface area, temperature analyse patterns in metal activity on the periodic table and explain why they correlate with, for example: – ionisation energy (ACSCH045) – atomic radius (ACSCH007) – electronegativity (ACSCH057) Metal activity increases left and down the periodic table. Ionisation energy: the amount of energy it takes to remove 1 electron from a stable atom.  The reactivity of metals increases as their ionisation energy decreases – the lower the ionisation energy, the easier it is to remove an electron and there the more reactive a metal is  As you move across the periodic table, the number of valence electrons increase and the ionisation increases  As move down the table, the number of valence electrons decreases so the ionisation energy decreases because the attraction between the valence electron and the nucleus decreases Atomic radius: the measure of distance between the nucleus and the outer shell electrons  Increases down a group as an extra electron shell is added  Decreases across a period due to the increasing size of the attractive charge from the nucleus  For active metals in groups 1 and 2, atomic radius increases down the group as does reactivity Electronegativity: the strength of attraction between the nucleus and the valence electron  Increases form left to right across a period  Decreases from top to bottom  Opposite to reactivity apply the definitions of oxidation and reduction in terms of electron transfer and oxidation numbers to a range of reduction and oxidation (redox) reactions If metal atoms have lost electrons, then some other species must have gained electrons. Reactions where one or more electrons are transferred from one atom to another are called electron transfer reactions or REDOX REACTIONS.  When an atom loses one or more electrons, it has been oxidised  When an atom gains one or more electrons, it has been reduced Oxidation and reduction occurs simultaneously – there cannot be an overall loss or gain of electrons Redox  reduction-oxidations reactions An oxidising agent or oxidant is a substance that oxidisers another agent and therefore is reduced. A reducing agent or reductant is a substances that reduces another agent and therefore is oxidised. Redox reactions occur:  during combustion  during metal displacement reactions  during the corrosion of iron Oxidation States 1. Oxidation state changes when electrons are lost or gained  Oxidation state increases during oxidation  Oxidation states decreases during reduction 2. Use of oxidation states  For determining if a reaction is a redox reaction  To describe a redox reaction in terms of which species have been oxidised and reduced Oxidation State rules: 1. ELEMENTS  substances in their elemental form have an oxidation state of 0 2. IONS  The charge on the ion is the oxidation state of the ion 3. OXYGEN  oxidation of -2 in compounds, except in peroxides where the oxidation state is -1 4. HYDROGEN  oxidation state of +1 in compound, except in compounds with metals where the oxidation state is -1 a. Metals – try to lose electrons to form stable outer shell so Hydrogen accepts electrons from metals – takes a negative oxidation state b. Non-metals – try to gain electrons to form a stable outer shell so Hydrogen donates electrons – takes a positive oxidation state 5. METALS  group 1,2, and 13 metals take same oxidation state as their ions (e.g. group 1 takes a +1 oxidation state, while group 2 takes +2 and group 13 takes a +3) 6. MOLECULAR IONS AND MOLECULES  the sum of oxidation states equal the charge on the molecular ion, while the sum is equal to zero for molecules AND the most electronegative element takes the negative oxidation number  Example: conduct investigations to measure and compare the reduction potential of galvanic half-cells Galvanic half-cells: electrochemical cells – these convert chemical energy into electrical energy Electrochemical cells = portable or fixed devices that store chemical energy. Galvanic cells separate the reducing agent and oxidising agent of a spontaneous redox reaction into two different half cells to produce energy. o The reducing agent is oxidised and loses electrons o Electrons travel to the oxidising agent o The oxidising agent is reduced and gains electrons This means there is a flow of electrons from the negative to the positive electrodes. This creates electrical energy. Half cells:  contains an electrode (metallic conducting plate that carries electric current into out of galvanic half- cells)  contains an electron donor and its corresponding acceptor  metal is generally used as the electrode and its submerged in its own ion  a gas can also be used as a reactant instead of a metal, where the gas is bubbled over an inert electrode and the solution containing its ions Anode: the electrode where oxidation occurs, the negative electrode cathode: the electrode where reduction occurs, the positive electrode Electron flow:  a circuit allows electron flow between half cells  electrons will flow from the anode to cathode  flow of electrons generate an electrical current which can be used by a load on the wire Salt bridges:  often a tube of potassium nitrate gel or filter paper soaked in solution of potassium nitrate (but can be soluble ionic compound)  provide ions that are free to move and balance the charge formed in the half cells, and create a complete circuit  the anode ions are balanced by the movement of nitrate ions from the salt bridge into the anode half cell the loss of positive cathode ions is balanced by the movement of potassium ions from the salt bridge into the cathode half cell  without the constant balance of charge, there would be an accumulation of charge in each half cell and the reaction would stop Galvanic Cell Notation:  label reducing agent first and corresponding oxidising agent second  separate the first half cell from the other half cell by a double line – this represents a physical separation of the two half cells  then label corresponding reducing agent and the oxidising agent if gas cell:  label the inert electrode, followed by the reducing agent in square brackets  mark corresponding oxidising agent after a line (this is a single ion) construct relevant half-equations and balanced overall equations to represent a range of redox reactions Half equations:  allows us to represent redox reactions in terms of the component reduction and oxidation reactions Example: oxidation: Mg(s)  Mg2+ + 2e- reduction: O2(g) + 4e-  2O2-s Balanced overall equations: combination of two half equations  electrons must be balanced example: 2Mg(s) + O2(g) + 4e-  2Mg2+ + 2O2- + 4e- which then becomes: 2Mg(s) + O2(g)  2MgO(s) Conjugate Redox pair: an oxidising/reducing agent and its corresponding reducing/oxidising agent  typically two per redox reaction example: 2Mg2+ as oxidising agent/ Mg(s) as reducing agent In oxidation half-equations, electrons appear on the product side of the equation because they involve a loss of electrons. Transforming an element to ionic form means the element has lost electrons. predict the reaction of metals in solutions using the table of standard reduction potentials Table of Standard Reduction Potentials  Lists the reduction potentials of a range of reactions and galvanic cells  Reduction potentials: a measure of the tendency of a reduction reaction to occur  in the reverse reaction, the reduction potential becomes the oxidation potential  developed under standard conditions – 100kPa, 1 mol/L concentration, 25°C  ranks both metals and non-metals  species towards the bottom left of the table are stronger oxidising agents   species towards the top right are stronger reducing agents - we know that the strongest reducing agent will be oxidised in a reaction and lose electrons  the table is generated in reference to a standard hydrogen half cell (which is given an E 0 of 0V)  some half cells are given a negative value to indicate that oxidation, rather than reduction, has occurred  the reaction higher up is the oxidation reaction – the cell potential is positive and the reaction is spontaneous Predicting cell reactions:  Identify the two half equations from the table  Identify which species is the strongest reducing agent, and which is the strongest oxidising agent, and write the two half equations that will occur  Write the overall cell equation  Identify the anode and the cathode, and direction of the electron flow  some of the cell potential comes from the oxidation reaction, and some from the reduction reaction predict the spontaneity of redox reactions using the value of cell potentials (ACSCH079, ACSCH080) Rates of Reactions conduct a practical investigation, using appropriate tools (including digital technologies), to collect data, analyse and report on how the rate of a chemical reaction can be affected by a range of factors, including but not limited to: Rate of reaction: how quickly or slowly a chemical reaction happens  the speed at which a chemical reaction proceeds Rate of reaction = (amount of product formed or amount of reactant used up) / time   Measuring the rate of reaction is assessed through how fast a product is formed or a reactant is used up  if a product is a gas, measure the change in mass as the reaction progresses  measure results as time progresses and product forms  measure with a control first – just the normal reaction under normal conditions  graph using time on x axis and product formed (or reactant used up) on y axis  rate of product formation at any point in the reaction is given by the gradient at the point of the graphs  if gradient is = 0.4, then rate of reaction = 0.4 cm3 / sec (or whatever units used on graph)  this is a comparable value that shows difference in rate of reactions when variables are changed  gradient = (change in y) / (change in x) temperature  As temperature increases, rate of reaction increases surface area of reactant(s)  As surface area increases, rate of reaction increases concentration of reactant(s)  As concentration of 1 or more reactant increases, rate of reaction increases  Increased concentration of gas means increased pressure catalysts (ACSCH042)  Catalysts = substances that speed up the rate of reaction without being used up investigate the role of activation energy, collisions and molecular orientation in collision theory Collisions in reactions  Particles must collide to react  A successful collision results in reaction Collison theory: Particles will react successfully if they collide under specific conditions 1. Particles must collide with at least a minimum amount of kinetic energy 2. Particles must collide with the correct orientations Activation energy: The minimum amount of kinetic energy particle need to collide and react  This is because breaking bonds takes energy  Normally in units kJ/mol  how much energy is needed to break the bonds in one mole of substance to begin the reaction  Reactions with low activation energy will happen much more easily then reactions with high activation energy  Reactions with high activation energy can be given energy by heating them Molecular orientation:   Particles need to collide in the right direction to react  so atoms line up in a way that enables the breaking and remaking of bonds  Gas and liquid particles are in random constant motion – there is always a probability that the particles will collide with correct orientation explain a change in reaction rate using collision theory (ACSCH003, ACSCH046) Temperature:  Particles increased in kinetic energy and move around faster  Particles collide more frequently because more particles fulfill their activation energy for collisions  More collisions  increases number of successful collisions even if probability of success remains the same  Hence more successful collisions occur and the rate will increase Concentration:  More particles  particles collide more because probability of collision occurring in that given space increases  More collisions occur meaning more successful collisions occur, increasing rate of reaction Surface Area:  Increased surface area increased number of particles exposed to collisions – increases number of possible collisions and hence increases collisions and rate of reaction Catalyst:  Lowers activation energy of reaction by providing a different way for the bonds to be broken and remade Module 4- Drivers of Reactions Energy Changes in Chemical Reactions Standard Enthalpy Changes: - All of the different enthalpy change describe reactions preformed under standard conditions Types of Standard Enthalpy changes: Enthalpies of combustion are always negative – burning always releases heat. Neutralisation: conduct practical investigations to measure temperature changes in examples of endothermic and exothermic reactions, including: – combustion – dissociation of ionic substances in aqueous solution (ACSCH018, ACSCH037) investigate enthalpy changes in reactions using calorimetry and 𝑞 = 𝑚𝑐𝛥𝑇 (heat capacity formula) to calculate, analyse and compare experimental results with reliable secondary-sourced data, and to explain any differences Calorimetry  measures heat changes Solid/liquid reactants:  put reactants in any vessel  measure the temperature before and after the reaction  use temperature in equation to find enthalpy change Gas reactant:  burn the fuel  let the flame heat the water  measure temperature change of water Limitation:  ideally the heat absorbed or released would go where it is measured, but heat loss or gain is inevitable - can minimise heat variation by insulating materials  incomplete combustion may occur – forms soot which doesn’t release as much energy output of a reaction  the consequences of these limitations is that the measured temperature change is not as high in theory Enthalpy change equation: q = enthalpy change, measured in joules (j) m = mass of substance being heated or cooled, measured in grams (g) c = specific heat capacity of the substance, measured in joules/grams/Kelvin (j/g/k) ∆ T = a change in temperature between the end and beginning of a reaction – always subtract the beginning temperature from the final temperature, can be measured in °C construct energy profile diagrams to represent and analyse the enthalpy changes and activation energy associated with a chemical reaction (ACSCH072) Exothermic: reaction gives out heat – when ∆H is negative – heat lost to surroundings Endothermic: absorbs heat - ∆H is positive – heat is taken from surroundings Enthalpy Level Diagrams:  shows the change in heat energy  y-axis represents energy  horizontal lines represent energy levels of the reactants and the product  vertical difference between the two lines represents the enthalpy change Things with low enthalpy are more stable than things with high enthalpies – chemicals like to be in a lower energy state.  the lower the horizontal line on an enthalpy diagram, the more difficult it is to react Reaction profiles:  show the enthalpy change during the reaction  there is never an instantaneous switch from reactants to products – sometimes there are intermediate compounds formed in the middle  y-axis represents the energy change  x-axis represents the reaction progress Exothermic profile: x Endothermic profile: Bond Enthalpies:  breaking bonds requires energy – so is an endothermic process  forming bonds releases energy – so is an exothermic process  different chemical bonds have different strengths and energies so the energy required to break bonds is not equal to the energy released when forming bonds  this change in energy is referred to as an enthalpy change  bond enthalpies are often referred to as average bond enthalpies  the energy required to break a chemical bond is also affected by the atoms and other bonds that surround it  the difference across molecules in a chemical bonds enthalpy means that bond enthalpy is quoted as an average value - this values gives the average amount of energy needed to break this particular bond Calculating enthalpy changes:  (energy put in to break bonds) – (energy released to form bonds) = enthalpy change Energy to break bonds > energy released  overall enthalpy is +  endothermic Energy to break bond < energy released  overall enthalpy is -  exothermic model and analyse the role of catalysts in reactions (ACSCH073) catalyst = a substance that speeds up the rate of reaction without itself being used up by providing an alternative reaction pathway with a lower reaction energy  less activation energy required to break bonds in reaction  speeds up rate of reaction as more of the reactant particles will have enough energy for their collisions to result in a reaction  only needs a tiny amount of catalyst in order to catalyse a large amount of the stuff you want to react  most catalysts will only work with a specific reaction Catalysts and Reaction Profile  catalysts do not impact the start and end products, or enthalpy, of reactions  only reduces the threshold energy needed for collisions to be successful Enthalpy’s Law explain the enthalpy changes in a reaction in terms of breaking and reforming bonds, and relate this to: – the law of conservation of energy Chemical bonds have energy associated with them – energy exists in bonds to hold atoms together breaking these bond requires energy from the surroundings  breaking bonds releases energy to the surroundings  bonds represent a lower energy configuration A bond’s energy is the amount of energy given up compared to the unbonded state.  Each bond broken has a different strength (and hence a different energy) – the overall reaction then must have either energy inputted or energy released  So, reactions result in energy changes The overall change in energy in a reaction is called the ENTHALPY CHANGE. Enthalpy change is in the form of heat energy. Enthalpy change = the heat change of a reaction held at a constant pressure  symbol is delta H (∆H) – delta means ‘a change in’ and H means ‘enthalpy’  unit is joules or kilojoules  commonly found in kJ/mol If a reaction gives out energy, ∆H is negative – the reactant energy is greater than the product energy. If a reaction absorbs energy from the surroundings, ∆H is positive – the heat energy of the chemicals increase and the product’s energy is greater than the reactants Standard conditions for enthalpy changes:  temp 25°C (or 289K)  pressure of 100kPa  all chemicals in reaction must be in their usual physical state If a reaction absorbs energy, the pre-existing bonds must have been stronger than the new ones. This is because it has taken more energy to break the existing bonds than was released in forming the new ones. investigate Hess’s Law in quantifying the enthalpy change for a stepped reaction using standard enthalpy change data and bond energy data, for example: (ACSCH037) – carbon reacting with oxygen to form carbon dioxide via carbon monoxide Hess’s Law = the enthalpy change of a reaction is independent of the route taken  reactions do not usually go straight form reactant to product energies, but rather other chemicals are produced and then react to form the products (an intermediate set)  the same products will always end up at the same enthalpy, no matter how they got there  Each chemical species has an inherent enthalpy due to the bonds holding them together apply Hess’s Law to simple energy cycles and solve problems to quantify enthalpy changes within reactions, including but not limited to:  – heat of combustion  – enthalpy changes involved in photosynthesis  – enthalpy changes involved in respiration (ACSCH037) Enthalpy cycles  show the routes a reaction could take while ending up at the same products  cycles used for when ∆H of a reaction is unknown  the number of intermediate steps does not impact enthalpy Enthalpies of formation in cycles  the formation enthalpy of an element is zero when doing calculation with cycles  the starting point of formation enthalpy is the element  when all the intermediates are elements, reverse the first cycle arrow (between the reactants and intermediates) – arrows represent the formation of compounds from their elements  arrows represent enthalpy change of formation so are labelled Hf  swapping the arrow makes the reaction go from decomposition to formation (as decomposition is opposite to formation)  swapping the direction of an arrow means we add a negative (so ∆Hf2 = -∆Hf3)  Combustion enthalpies cycles  even if the products are the same, different amounts of energy is released along the two paths – this can be used to find enthalpy of formation reaction enthalpy change of the formation reaction = the enthalpy of combustion for the reactants – the enthalpy of combustion for products Entropy and Gibbs Free Energy analyse the differences between entropy and enthalpy Entropy: a measure of disorder of a system  a system is disordered if there are many ways of arranging the particles that make it up  ordered if molecules are fixed in place  relating to energy, this means how many different ways can energy in a system be distributed between the particles that make it up Entropy changes  a process will always result in an increase in entropy unless work is done  in spontaneous chemical reactions, the entropy of the system will always increases  Characteristics of a chemical reaction that impact entropy change:  Gaseous moles – gases have a much higher entropy that solids or liquids o The side of an equation with more moles of gas has the highest entropy  Temperature - increasing the temperature increases the energy of our particles o Particles vibrate more vigorously in solids, flow more easily in liquids, and have a higher speed and collide more in gases  causes more disorder as there are more arrangements of the particles  Particle number – more ways to arrange the particles Entropy Laws  universe moves from ordered to disordered states – any action will make the universe more disordered  decreasing the entropy in one place must increase it even more in another place  if a reaction results in a decrease in entropy, work has to be done on the system – this work has a cost elsewhere still that increases the entropy of the universe Every chemical or substance has a particular amount of disorder at a particular temperature and pressure. use modelling to illustrate entropy changes in reactions predict entropy changes from balanced chemical reactions to classify as increasing or decreasing entropy Entropy = S Units : J/K/mol – this is because entropy is related to the energy of particles and depends on the number of particles and temperature Standard entropy conditions: 298 K and 1 atm of pressure ∆S = change in entropy – useful in determining whether a reaction is likely to happen Calculating entropy change: ∆Ssystem = Sproducts – Sreactants The entropy change of an isolated system is always positive (increasing) - includes the universe as it is not in contact with any body that can do any work on it. Reaction system and surroundings are non-isolated systems because they can interact with and transfer energy to each other – they can positive or negative. Entropy of the surroundings change based on whether we give or take energy away from it. ∆Ssurroudnings = -∆H / T ∆Ssurroudnings: entropy change of surroundings ∆H: enthalpy change of the reaction T: temperature of the surroundings (K) – Room Temp: 298K explain reaction spontaneity using terminology, including: (ACSCH072) – Gibbs free energy – enthalpy – entropy solve problems using standard references and 𝛥𝐺𝑜 = 𝛥𝐻 𝑜 − 𝑇𝛥𝑆𝑜 (Gibbs free energy formula) to classify reactions as spontaneous or nonspontaneous Gibbs free energy determines if a reaction can feasibly happen by considering both entropy and enthalpy symbol: G units: J/mol Free energy change: ∆G ∆G = ∆H – T∆S ∆G = Gibbs free energy change ∆H = enthalpy change T = temperature ∆S = entropy change A reaction is feasible (spontaneous) if ∆G is negative or zero.  this means enthalpy needs to be negative (-∆H)  entropy change needs to be positive (+∆S)  does not guarantee reaction will happen A positive ∆G value is not feasible so must input external energy to begin reaction. predict the effect of temperature changes on spontaneity (ACSCH070) Entropy and enthalpy changes can be positive or negative so there are different combinations of the effect of temperature on feasibility. Critical value: the temperature at which ∆G goes between positive and negative (feasible and not feasible) where ∆G = 0  can be found when ∆G is substituted with 0 so the equation becomes T = ∆H /∆S