Chemistry Past Paper PDF

Chemistry Working scientifically Chapter 1 Observation: using all your senses as well as instruments to observe things that your senses alone cannot detect Hypothesis: a tentative explanation for an observation that is based on evidence and prior knowledge (testable and falsifiable) Inquiry questi...

Chemistry Working scientifically Chapter 1 Observation: using all your senses as well as instruments to observe things that your senses alone cannot detect Hypothesis: a tentative explanation for an observation that is based on evidence and prior knowledge (testable and falsifiable) Inquiry question: Method (procedure): a step-by-step description of how the hypothesis will be tested Validity: is the experiment testing the set hypothesis? (only one variable is changed at a time) Reliability: experiment can be repeated many times with consistent results Accuracy: obtaining correct results, reducing systematic errors Precision: to consistently obtain the same measurement by reducing random errors Risk assessment: helps to identify, assess and control hazards PPE: should be worn Errors Mistakes: avoidable errors Systematic: consistent and will occur again if the experiment is repeated in the same way and result in bias Random: unpredictable and follow no random pattern Uncertainty = ± maximum difference from the mean Qualitative Quantitative Correlation and causation (causing/influencing an outcome) Module 5: Equilibrium and acid reactions Static and dynamic equilibrium What happens when chemical reactions do not go through to completion? Gibbs free energy - ∆𝐺 < 0 (spontaneous) - ∆𝐺 > 0 (non-spontaneous) - ∆𝐺 = 0 (equilibrium) - ∆𝐻 < 0, ∆𝑆 > 0 (spontaneous) - ∆𝐻 > 0, ∆𝑆 < 0 (non-spontaneous) Open and closed systems Irreversible and reversible systems Irreversible Reversible - Combustion - Evaporation/condensation of water - Photosynthesis - Formation of saturated sugar solution - Neutralisation (acid/base) - Reaction between hydrated cobalt (II) - Burning magnesium chloride and dehydrated cobalt (II) - Burning steel wool chloride - Reaction between iron (III) nitrate and potassium thiocyanate EQUILIBRIUM Occur in closed systems Forward reaction: when the reactants form the products Reverse reaction: when the products re-form the reactants Experiment 1 Reaction of hydrated cobalt (II) chloride and dehydrated cobalt (II) chloride Solid dehydrated cobalt (II) chloride blue reacts with water colourless to produce hydrated cobalt (II) chloride pink - Spectator ions removed - Therefore REVERSIBLE Experiment 2 Reaction of iron (III) nitrate and potassium thiocyanate The two solutions are mixed to produce iron (III) thiocyanate red  In an energy profile diagram, the forward reaction is endothermic while the reverse reaction is exothermic Collision theory - Frequency of successful collisions (to overcome E ) A - Correct orientation  - Sufficient kinetic energy (KE), equal to or greater than E A Dynamic equilibrium - Rate of forward reaction is equal to rate of reverse reaction - Bonds are constantly being broken and new bonds are formed as products and reactants continue to convert from one form to the other - Concentration remains constant - Occur in a closed system Formation of ammonia gas from 1 mole nitrogen gas and 3 mole hydrogen gas Interpreting the rate of reaction / time graph - Forward: from collision theory, the concentrations of nitrogen and hydrogen decrease → the frequency of collisions decrease → rate of production of ammonia decreases  - Reverse: at the same time ammonia is being formed, some ammonia molecules collide and decompose to reform nitrogen and hydrogen - Equilibrium: the forward and reverse reactions proceed at the same rate (ammonia is being formed at the exact same rate that it is breaking down). The concentrations of all three remain constant. No macroscopic change can be observed Interpreting the concentration / time graph - Equilibrium is first established when there is no change in ANY of the concentrations Static equilibrium - The rates of the forward and reverse reactions are almost zero - No conversion of reactants to products or products to reactants - E.g. conversion of graphite into diamond Modelling Static Dynamic Non-equilibrium systems Combustion - Exothermic → negative enthalpy (ΔH < 0) - Disorder increases → increase entropy (ΔS > 0) - Negative gibbs (ΔG < 0) → spontaneous reaction Photosynthesis - Endothermic → positive enthalpy (ΔH > 0) - Disorder decreases → decrease entropy (ΔS < 0) - Positive gibbs (ΔG > 0) → non-spontaneous Factors that affect equilibrium What factors affect equilibrium and how? Le Chatelier’s principle - “If an equilibrium system is disturbed, the equilibrium will shift/adjust to minimise the disturbance” - Law of opposites Position of equilibrium depends on the reaction conditions - Concentration - Temperature - Pressure/volume Pressure and volume 𝑘 𝑃𝑟𝑒𝑠𝑠𝑢𝑟𝑒 = 𝑣𝑜𝑙𝑢𝑚𝑒 An equilibrium system will respond to an increase in pressure by SHIFTING to reduce the pressure, therefore moving the equilibrium position to the side with the LEAST gas particles (moles) Collision theory: as volume decreases, gas molecules are closer and successful collision become more frequent - The initial direction of the reaction (e.g. forward) will become the net forward reaction  The concentration of gases will increase simultaneously, until they gradually change to accommodate the disturbance, arriving at a new equilibrium Spontaneous increase (spike) For a reaction with equal numbers of reactant and product particles, a change in pressure will not shift the equilibrium position The volume decrease (pressure increase) causes the rates of both the forward and reverse reactions to increase equally The addition of an inert gas will keep the system at equilibrium (does not change concentration of reactants or products) but at a higher total pressure. Collision theory states that any collisions with an inert gas will not produce a reaction Dilution The addition of water instantaneously decreases the concentration of both products and reactants - Adding water decreases the number of particles per volume, shifting the position of equilibrium towards the side that produces the greater number of dissolved particles - Therefore, net reverse reaction Temperature The effect of a change in temperature depends on whether the reaction is exothermic or endothermic Taking the exothermic reaction (above), brown to colourless: - Increasing the temperature increases the KE of substances in the mixture = increase in frequency of successful collisions (overcoming EA) - The reaction will minimise the increase in KE by absorbing energy (endothermic) - Favouring the net reverse reaction Increase in temperature = increase in KE = gradual change in concentration = net reverse reaction Gradual change (curve) A temperature increase causes the rates of both the forward and reverse reactions to increase equally Catalyst A catalyst lowers the activation energy (EA) for both the forward and reverse reactions equally - Increasing rate of both forward and reverse reactions, given the lower EA increases the frequency of successful collisions, arriving at equilibrium faster - Catalyst will not change concentration at equilibrium OR the equilibrium constant (Keq) Concentration Concentration vs time graph Rate vs time Calculating the equilibrium constant (Keq) How can the position of equilibrium be described and what does the equilibrium constant represent? Equilibrium law Equilibrium law states: - Equilibrium constant (Keq) is the concentrations of products divided by reactants, at equilibrium - Aqueous and gases only - The index for each component concentration is the coefficients for the substances in the balanced equation - Equilibrium constant @ equilibrium while reaction quotient (Q) @ any stage during the reaction Q < Keq: right shift to form more products Q = Keq: system is at equilibrium Q > Keq: left shift to form more reactants  Heterogeneous reactions - Concentration of pure solid or liquid is 1 (i.e. neglected) Equilibrium constant - Determines the equilibrium yield (conc. of products divided by conc. of reactants at equilibrium) Only temperature affects equilibrium constant Calculating equilibrium constant Ratio 1 1 2 Initial 0.232 0 0 Change –x +x +2x Equilibrium 0.12 0.5 1 [𝑃𝑟𝑜𝑑𝑢𝑐𝑡𝑠] 𝐾𝑒𝑞 = [𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] Solution equilibria How does solubility relate to chemical equilibrium? Solubility rules NO3- soluble NH4+ soluble Group 1 soluble Acetate soluble SO42- mostly soluble CO32- mostly insoluble PO43- mostly insoluble OH- mostly insoluble Module 6: Acid/base reactions Module 7: Organic chemistry Module 8: Applying chemical ideas