Chemistry - Units 3&4 Exam

understand that an acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours; the acidic form is of a different colour to the basic form

* Acid - Base indicators (also known as pH indicators) are substances which change colour with pH. They are usually weak acids or bases, which when dissolved in water dissociate slightly and form ions * An acid-base indicator is either a weak acid or weak base that exhibits a colour change as the concentration of hydrogen (H+) or hydroxide (OH-) ions changes in an aqueous solution. Acid-base indicators are most often used in a titration to identify the endpoint of an acid-base reaction

recognise that indicators change colour when the pH = pKa and identify an appropriate indicator for a titration, given equivalence point of the titration and pH range of the indicator

* Acid-base indicators are usually weak acids or bases, which dissociate slightly in water and form ions * Indicators change colour in acidic and basic solutions, so they can provide information about the pH of a solution - Indicators change colour when the solution contains equal amounts of both forms: * This means that the colour of the indicator changes when pKa = pH - equivalence point: the point in a titration when the reactants have reacted in the molar ratio of the balanced chemical equation - end point: the point in a titration when the indicator changes colour

analyse experimental data to determine and compare the relative strengths of acids and bases

* The strengths of Brønsted-Lowry acids and bases in aqueous solutions can be determined by their acid or base ionisation constants * Stronger acids form weaker conjugate bases * Weaker acids form stronger conjugate bases

understand that galvanic cells, including fuel cells, generate an electrical potential difference from a spontaneous redox reaction which can be represented as cell diagrams including anode and cathode half-equations

* general rule (redox table): - if a species is on the top right, it will spontaneously react with a bottom left species

distinguish between the terms end point and equivalence point

- The equivalence point is when the amounts of two substances are just sufficient to cause complete consumptions of both reactants (equivalent quantities of acid and base - not always the same moles) - The endpoint is when the indicator changes colour (the pH unique to the indicator where the colour change occurs) * it is important to match the equivalence point of the reaction with the endpoint of the indicator, or the result will be under or over titrated

analyse experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration of reactants and product against time, to identify the position of equilibrium

- at the start of the reaction, the forward reaction has a higher rate - at the start of the reaction the concentration/amount of reactants are high and there are no products - as the reaction proceeds, the concentration of the products increases * As a reaction is approaching equilibrium: - Reactant concentration: As the reaction goes to the right, the reaction concentration decreases. - Products concentration: As the reaction goes from left to right, the concentration of the products increases. - Forward reaction rate: The reaction concentration decreases and therefore, there are fewer reactant collisions causing the forward rate to decreases. - Reverse reaction rate. The concentrations of the products increases, therefore there are more product collisions causing the reverse reaction rate to increase. - the forward and reverse rates are equal at equilibrium - the reactant and product concentrations and the macroscopic properties are constant at equilibrium * Although the concentrations are constant when equilibrium is reached, the forward and reverse reactions continue forever and the reaction is never 'complete'. measuring concentration (two methods): - pH (reaction involves an acid or base, measure change in pH - hydrogen ion concentration to determine the concentration from the initial reaction to when equilibrium is reached using stoichiometry) - observation (colour used to indicate that the reaction has moved forward)

use appropriate mathematical representation to solve problems for hydrogen ion concentration [H+ (aq)], pH, hydroxide ion concentrations [OH- (aq)] and pOH

- calculating pH of a strong acid: * pH = -log[H+ ] - calculating [H+ ] from pH: * [H+ ]= 10^-pH - the pOH scale (calculating pOH of strong base): * pOH = -log[OH-] - relationship between pH and pOH * pH + pOH = 14 the pH of mixtures: - when a strong acid and base are mixed, neutralisation (ph of about 7) takes place, but only if there's no excess reactant - if excess reactant is 'leftover' it will affect the pH eg. nHCl = 0.01M nNaOH = 0.005 - n = c x v (how those are found) therefore, NaOH is limiting reagent * this will be a limiting reagent problem * amount of HCL left over = 0.005 therefore c = n/v = 0.005 /0.11 = 0.045M * pH = -log(0.045) = 1.35

explain the effect of changes of concentration and pressure on chemical systems at equilibrium by applying collision theory to the forward and reverse reactions

- concentration of products or reactants: note: only substances in aqueous or gas states can change concentration (both solid and liquid will not change by adding more of it) - concentration of a substance (aq/g) can be increased by: * adding more solute to solution * increasing the pressure of the gas (decreasing volume) - concentration of a substance (aq/g) can be decreased by: * removing some of substance * adding more solvent to the solution * decreasing pressure of the gas (increase volume) example of a change in concentration: A(s) + B(aq) --> C(l) + D(g) - A or C changed = no change in equilibrium - more B added/B increased (counteract - forward reaction favoured = shift right and new equilibrium established) - D is removed/lowered = shift right - more pressure increased = shift left (D increased) - more water added = shift left (B lowered) - pressure * pressure is caused by the number of collisions - when volume is changed, the concentrations and the pressures of both reactants and products are changed (gas or aqueous) - the effect of change in pressure is only applicable to gaseous equilibrium * increase in pressure = equilibrium shift to the side of reaction with fewer moles of mass (oppose change by lowering pressure, as to counteract, the system aims to do the opposite) * decrease in pressure = equilibrium shift to the side with more moles - if there are the same number of moles on each side of the reaction, an increase or decrease in pressure will not cause a change - the addition of an inert (unreactive) gas eg. argon, can increase the pressure of the system but will have no effect on an equilibrium (as the concentrations of all substances remain the same)

deduce the extent of a reaction from the magnitude of the equilibrium constant

- if K is large (>10^3) then equilibrium shifts to the right to favour the products - if K is small (<10^-3) then equilibrium shifts to the left to favour the reactants * if K = 1 then equilibrium has similar concentrations of reactants and products present

understand that physical changes are usually reversible, whereas only some chemical reactions are reversible

- irreversible reactions: reactions in which the reactants covert to products and the products cannot convert back to reactants eg. combustion, gas to solid etc. - reversible reactions: reactions in which the reactants covert to products and the products can convert back to reactants two reactions: - forward: A + B --> C + D - back/reverse: C + D --> A + B the combined effect is the reaction: A + B ⇌ C + D - physical changes are usually reversible, whereas only some chemical reactions are reversible

recognise that the strength of acids is explained by the degree of ionisation at equilibrium in aqueous solution, which can be represented with chemical equations and equilibrium constants (Ka)

- strong acids ionise completely (Ka>1) - weak acids don't ionise completely (Ka<1) - the first ionisation/dissociation constant is greater than the second Ka1 > Ka2 because the first proton to dissociate is the most strongly acidic, followed by the next most strongly acidic proton eg. H2SO4 can donate 2 protons in a solution - first dissociation step will occur completely (strong acid), but the second step is only weakly dissociating

explain and predict the effect of temperature change on chemical systems at equilibrium by considering the enthalpy change for the forward and reverse reactions

- temperature * endothermic - energy required (is absorbed and stored in bonds) for the reaction to take place (ΔH = +) * exothermic - energy is released by the reaction (ΔH = -) - for a reversible reaction, one reaction is endo and the opposite is exo - the ΔH given always represents the forward reaction - energy can be written as a reactant (endo) or a product (exo) eg. an endothermic reaction - if temp is increased = shift right (endo reactions are favoured by the increased temperature) * the reaction shifts to the side with less energy (trying to counteract the increased temp - shifts away from the energy) - if temp is decreased = shift left (exo reactions are favoured by a decreased temp) * shifts to the side with more energy (trying to counteract the decreased temp) eg. exothermic reaction - temp is increased = backward favoured (shift left) - temp decreased = forward favoured (shift right)

appreciate that observable changes in chemical reactions and physical changes can be described and explained at an atomic and molecular level

- when equilibrium is reached, no further observable change takes place

recognise that the relationship between acids and bases in equilibrium systems can be explained using the Brønsted-Lowry model and represented using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs

Bronsted-Lowry model defines: - acids as any species that donates H+ protons - bases as any species that accepts H+ protons - neutralisation: exothermic reaction where the proton is transferred (change in pH (power of hydrogen ion) is measured by an indicator)

deduce the equilibrium law expression from the equation for a homogeneous reaction and use equilibrium constants (Kc), to predict qualitatively, the relative amounts of reactants and products (equilibrium position)

Equilibrium constant: aA + bB ⇌ cC + dD - net ionic equation - only the reacting species are shown in the equation * for this equation, the equilibrium constant is Kc = [C]c [D]d / [A]a [B]b - a reaction's equilibrium constant (K) measures the extent to which reactants are converted to products - if K is large (>10^3) then equilibrium shifts to the right to favour the products - if K is small (<10^-3) then equilibrium shifts to the left to favour the reactants * if K = 1 then equilibrium has similar concentrations of reactants and products present * the only factor that changes K is temperature in an exothermic reaction: - if the temperature decreases, the equilibrium will shift to favour the exothermic reaction (shift right to oppose change) and K is, therefore, increased - if the temperature increases, the equilibrium will shift left to oppose change and K is, therefore, decreased in an endothermic reaction: - if the temperature decreases, the equilibrium will shift to favour the endothermic reaction (shift left to oppose change) and K is, therefore, decreased - if the temperature increases, the equilibrium will shift right to oppose change and K is, therefore, increased Reaction quotient (Q): - when Kc is the equilibrium constant *Q is a quantity that changes as a reaction system approaches equilibrium (shows you where the reaction is in terms of equilibrium position) - As soon as equilibrium is reached, Q = K K = Q = products/reactants - If Q < K (more reactants present, shift right until the system reaches equilibrium) - If Q > K (more products present, shift left until the system reaches equilibrium)

deduce the oxidation state of an atom in an ion or compound and name transitional metal compounds from a given formula by applying oxidation numbers represented as roman numerals

Identifying the reducing agent and oxidising agent is given an equation: - oxidation numbers/states (oxidation rules): * often redox reactions are complex with more than two species, difficult to identify RA and OA, therefore, you assign oxidation state/number * oxidation number = number of electrons gained or lost by an atom - when atoms gain electrons during a reaction the oxidation state decreases, when they lose electrons the oxidation state increases * gain (electrons) = reduction - reduces oxidation number left to right (reactant to product) * loss (electrons) = oxidation - increases oxidation number Table of rules: - elements have an oxidation number of 0 - certain elements when present in compounds have common oxidation numbers (group 1 is +1, group 2 is +2, H is +1 and O is -2 except in peroxide = -1) - monatomic ions (charge on the ion) - polyatomic ions (sum of oxidation numbers = charge) - neutral compound (sum of oxidation numbers = 0) - most electronegative element is (-) and the other is (+)

determine the expression for the dissociation constant for weak acids (Ka) and weak bases (Kb) from balanced chemical equations

Ka = acid ionisation constant: Ka = [H3O+] [A-] / [HA] - products over reactants, not water as it is a liquid state Kb = base dissociation constant: Kb = [BH+] [OH-] / [B] - large Kb value indicates high level of dissociation of a strong base Calculating the pH of a weak acid using Ka: - weak acids don't ionise completely; small Ka values; reaction lies to the left - therefore, [H+] is not the same as the concentration of the acid (it is less) - hence, pH value needs to be calculated differently (using the Ka to determine [H+] - not the normal formula) - same process for the Kb value to find [OH-] for a weak base

use appropriate mathematical representation to solve problems, including calculating dissociation constants (Ka and Kb) and the concentration of reactants and products

Kw = Kb x Ka 1x10^14 = Kb x Ka You can calculate the acid dissociation constant (Ka) of a weak acid if you know the pH and concentration of a solution of the acid. - Ka = [H3O+] [A-] / [HA] = [H+]^2 / HA ([H+] becomes squared as the concentration of the reactants are is the same) * use c = n / v and pH = -log[H+], [H+] = 10^-pH You can calculate the pH of a solution of a weak base if you know the base dissociation constant (K ). Firstly, find the concentration of [OH-] from the expression for K and then you can calculate the pOH and pH. - Kb = [BH+] [OH-] / [B] = [OH-]^2 / [B] (OH-] becomes squared as the concentration of the reactants are is the same)

apply Le Châtelier's principle to predict the effect changes of temperature, concentration of chemicals, pressure and the addition of a catalyst have on the position of equilibrium and on the value of the equilibrium constant

Le Chatelier's Principle: - when chemical equilibrium is disturbed by changing the conditions, the system will react in a way to counteract the change (new equilibrium is then established) - if an equilibrium system is subjected to a change, the system will adjust itself to partially oppose the effect of the change these could be: - concentration of products/reactants - temperature (endo/exo - reaction favours one) - pressure - catalyst - if a catalyst is added: both reactions (forward and backwards) will be sped up * therefore, the equilibrium is reached sooner, but it is not disrupted (no reaction will be favoured/no shift in equilibrium)

understand that acids are substances that can act as proton (hydrogen ion) donors and can be classified as monoprotic or polyprotic depending on the number of protons donated by each molecule of the acid

Monoprotic acid: - acid that donates only one H+ ion Polyprotic acid: - donates more than one proton per molecule: * Diprotic: two separate steps to donate one H+ at a time eg. H2SO4 - sulfuric acid (Ka1 > Ka2) * Triprotic: three separate steps to donate one H+ at a time eg. H3PO4 - phosphoric acid The distinction between strength and concentration should be covered.

use appropriate mathematical representation to solve problems, including calculating equilibrium constants and the concentration of reactants and products

RICE calculations: - a way to work out the concentrations of other substances in equilibrium in order to work out Kc R (ratio) I (initial) C (change) - use ratios E (equilibrium) - use E to find Kc by subbing into the Kc equation to find whether the reaction has equal concentrations or lies more to the right or left

understand that water is a weak electrolyte and the self-ionisation of water is represented by Kw = [H+ ][OH- ]; Kw can be used to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution

Self ionisation of water H2O(l) + H20(l) ⇌ H3O+(aq) + OH-3(aq) therefore, *Kw = [H+ ][OH- ] = 1x10^-14 - the equilibrium lies far to the left (indicating that only a small fraction of ions (products) are present the pH of water = -log(1x10^-7) = 7 (neutral) - as Kw increases, pH decreased - pH = power of hydrogen * the negative of the logarithm (base 10 scale which increases tenfold) of the [H+] pH = -log[H+ ]

mandatory practical - acid-base titration to calculate the concentration of a solution with reference to a standard solution

Standard solutions: - a solution of accurately known concentration to calculate the concentration of an unknown solution - primary standard (a substance of sufficient purity and stability that a standard solution can be prepared by weighing out the desired mass and making up an appropriate volume in a volumetric flask) eg. sodium carbonate, sodium hydrogen carbonate (bases) and potassium hydrogen phthalate (acid) - not suitable if substances react with water Accuracy: - high level of precision (appropriate glassware, rinsing glassware, reading meniscus at the bottom of the curve at eye level, measuring to two decimal places and using correct indicators) Rinsing procedures: - conical flask (rinse using distilled/ionised water only - can be left wet as it doesn't change the moles of solute) - volumetric flask (same as above) - burette (rinse with water then with the solution of the same concentration twice) - pipette (same as above) - need x3 close measurements of 0.1 difference to be able to conclude an accurate method has been followed

sketch the general shapes of graphs of pH against volume (titration curves) involving strong and weak acids and bases. Identify and explain their important features, including the intercept with pH axis, equivalence point, buffer region and points where pKa = pH or pKb = pOH

Titration curves: - it shows how pH changes during the course of the titration [H+] - strong acid v strong alkali * neutral pH (alkali titrated into acid = down to up, acid titrated into alkali = up to down) - strong acid v weak alkali * acidic pH (alkali titrated into acid = down to up, acid titrated into alkali = up to down) - weak acid v strong alkali * basic pH (alkali titrated into acid = down to up, acid titrated into alkali = up to down) weak acid v weak alkali * neutral pH (alkali titrated into acid = down to up, acid titrated into alkali = up to down) Buffer region and pKa values on titration curves: - buffers = weak acids/bases that act like "shock absorbers" to resist change in pH * area on a titration curve that remains relatively unchanged is called the buffer region * half equivalence point is the point in the buffer region that is halfway to the equivalence point half equivalence point (pH = pKa)

recognise that acid-base titrations rely on the identification of an equivalence point by measuring the associated change in pH, using chemical indicators or pH meters, to reveal an observable end point

Titrations: - analytical technique - pipette used (accurately measure solutions going into a conical flask) - burette allows accurate measurement of solution - The equivalence point is when the amounts of two substances are just sufficient to cause complete consumptions of both reactants - An indicator is used in acid-base titration to detect equivalence point - The endpoint is when the indicator changes colour (pH change) Titrations/Volumetric Analysis: - a technique in which a solution of a known concentration is added to a solution of unknown concentration until they have reacted in equal amounts (neutralised) * standard solution (burette) - known solution, unknown volume (purpose of titration) * unknown concentration, known volume (volumetric flask) Steps: 1. titrate until the colour of the solution in flask changes (based on pH indicator) 2. read volume of titrant used 3. calculate moles for the standard solution 4. find the molar ratio from the equation 5. calculate unknown concentration - diluting: to minimise error, concentrated solutions should be diluted to minimise error before titrating * remember to take the dilution into consideration when calculating the concentration - aliquot = samples from the same solution being tested

use appropriate mathematical representations and analyse experimental data and titration curves to solve problems and make predictions, including using the mole concept to calculate moles, mass, volume and concentration from volumetric analysis data

Volumetric analysis: - determine the concentration of one solution when the exact concentration of another is known - accurate measurements (reliable results)

identify and deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid)

acid + base = neutralisation (water and a salt) acid + base ⇌ conjugate base + conjugate acid eg. NH4+ + OH- ⇌ NH3 + H2O - strong base forms weak conjugate acid acid + base ⇌ conjugate acid + conjugate base eg. HNO3 + H20 ⇌ H30+ + NO3- - strong acid forms a weak conjugate * refer back to Brønsted-Lowry definitions

distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity and distinguish between the terms strong and concentrated for acids and bases

acids: - chemicals that donate H+ ions in chemical reactions - they conduct electricity - for example, HCl ionises (forms ions) in water - hydrogen ions attach onto the co-ordinate covalent bond of water common acids: - weak: phosphoric acid - H3PO4, acetic acid - CH3COOH, carbonic acid - H2CO3 - strong: sulfuric acid - H2SO4, nitric acid - HNO3, hydrochloric acid - HCl strong vs weak acids: - strong acids completely ionise in a solution (all molecules become ions) - weak acids ionise partially in solution (only some become ions)

recognise that amphiprotic species can act as Brønsted-Lowry acids and bases

amphiprotic species: - substances that can act as either an acid or a base eg. the water molecule has hydrogen atoms and can, therefore, act as an acid. However, acids can donate a proton to water and it becomes a proton acceptor.

appreciate that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added; Le Châtelier's principle can be applied to predict how buffer solutions respond to the addition of hydrogen ions and hydroxide ions.

buffers: - "chemical shock absorbers" - buffers are solutions that resist a change in pH when a small amount of acid/base is added * usually they are weak acid/bases and their conjugates * important solutions for when a stable pH is required eg. swimming pools or blood general equation: HA + H2O ⇌ H3O+ + A- (weak acid) (conjugate base) or a weak base and conjugate acid * this makes the buffer solution resist change in pH when a small amount of acid/base is added - adding a small amount of acid = [H+] increases and equilibrium shifts left (therefore H+ lowers so pH doesn't change drastically) - adding a small amount of base (OH-) = OH- reacts with H+ to form water = [H+] decreases and equilibrium shifts right (therefore H+ is re-established so pH doesn't change drastically) * buffers control pH by removing excess hydronium/hydroxide ions eg. how does the blood buffer system work: CO2 + H2O ⇌ H2CO3 ⇌ HCO3- + H+ - too much acid (H+) in the blood, reaction shifts left and absorbs the excess H+ ions. H2CO3 is unstable and forms carbon dioxide and water, breathe out CO2 (pH does not change much) - if too much base (OH-) is in the blood, OH- reacts with H+, equilibrium shifts right and produces H+ (acid) (pH does not change)

symbolise equilibrium equations by using ⇌ in balanced chemical equations

chemical equilibrium: - when the rate of the forward reaction equals the rate of backward reaction, the system is in equilibrium - dynamic equilibrium is only possible within a closed system - the rate of two opposing reactions are equal - at equilibrium, the concentration of all reactants and products remain constant (but not equal) - when equilibrium is reached, no further observable change takes place

recognise that chemical systems may be open or closed

chemical systems may be: - open: allowing matter and energy to be exchanged with the surroundings (no equilibrium) - closed: allow energy, but not matter, to be exchanged with the surroundings (process eventually achieves equilibrium) - no reactant/product can escape

understand that, over time, physical changes and reversible chemical reactions reach a state of dynamic equilibrium in a closed system, with the relative concentrations of products and reactants defining the position of equilibrium

equilibrium: a reaction which proceeds both forward and backward - Chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time, so that there is no observable change in the properties of the system - Reaction is reversible (occurs forwards and backwards) characteristics of equilibrium: - Forward rate is equal to the reverse rate - The concentration of reactants and products are constant (not equal) - Macroscopic (observable) properties are constant (colour, mass, density, pressure, concentrations) - no VISIBLE changes conditions of equilibrium: - Must have a closed system (no reactants or products lost) - Must have a constant temperature - Ea (activation energy required to start the reaction) must be low enough to allow a reaction * at any given time, both reactants and products are present in the system (equilibrium), this reaches dynamic equilibrium when: 1. the reaction is incomplete (not only products remain) 2. bonds are constantly broken and reformed (particles always moving) 3. rates of forward and backward reaction are equal (concentration of substances are constant but not equal) - dynamic is the same as equilibrium dynamic equilibrium: the state a reaction (have to undergo physical change) reaches when the rates of the forward and backward reaction are equal DYNAMIC (moving) - changing on an ATOMIC level The word dynamic means that forward and reverse reactions continue to occur - dynamic equilibrium is only possible within a closed system - when equilibrium is reached, the concentrations of reactants and products remain constant but not equal

explain the reversibility of chemical reactions by considering the activation energies of the forward and reverse reactions

forward reaction rate decreases as equilibrium is approached: - As the reaction goes to the right, the reaction concentration decreases and therefore, there are fewer reactant collisions causing the forward rate to decrease. reverse reaction rate increases as equilibrium is approached: - The reverse reaction rate increases as equilibrium is approached because as the reaction goes from left to right, the concentrations of the products increases, therefore there are more product collisions causing the reverse reaction rate to increase.

understand that equilibrium law expressions can be written for homogeneous and heterogeneous systems and that the equilibrium constant (Kc), at any given temperature, indicates the relationship between product and reactant concentrations at equilibrium

homogeneous equilibrium = substances are in the same state heterogeneous equilibrium = substances are in different states

identify the species oxidised and reduced, and the oxidising agent and reducing agent, in redox reactions

oxidation: - loses electrons - gets oxidised - reducing agent (reductant) = reactant that causes another reactant to gain electrons and be reduced, and is itself oxidised reduction: - gains electrons - gets reduced - oxidising agent (oxidant) = reactant that causes another reactant to lose electrons and be oxidised, and is itself reduced * do practice questions

understand that the pH scale is a logarithmic scale and the pH of a solution can be calculated from the concentration of hydrogen ions using the relationship pH = -log10 [H+ ]

pH = -log[H+ ] [H+ ]= 10^-pH

explain the relationship between the pH range of an acid-base indicator and its pKa value

pKa: - Ka can be used to calculate pKa (pKa = -log(Ka) ) - pKa is an indication of the degree of ionisation of an acid (how strong the acid is) pKb: - Kb can be used to calculate pKb (pKb = -log(Kb) ) - pKb is an indication of the degree of dissociation of a base (how strong the base is) *as pKa decreases (degree of ionisation), Ka increases (acid dissociation constant) pKa + pKb = 14 To summarise: - strong acid: high Ka, low pKa - weak acid: low Ka, high pKa - strong base: high kb, low pKb - weak base: low Ka, hight pKb What is the difference between pH and pKa? - pH is an indication of how many H+ are present (more H+ = lower pH) - pKa is an indication of the degree of ionisation (more ionisation, lower pKa)

recognise that a range of reactions, including displacement reactions of metals, combustion, corrosion and electrochemical processes, can be modelled as redox reactions involving oxidation of one substance and reduction of another substance

redox = reduction and oxidation redox: chemical reaction involving the transfer of electrons from one reactant to another (occurs simultaneously - one can't happen without the other) - oxidation: loss of electrons (oxidise = to lose) - reduction: gain electrons (reduce = to gain) * forms new products after transfer OILRIG (oxidation is loss, reduction is gain) redox reaction modelling (reactions that involve the transfer of electrons): - displacement * more reactive metal ion eg. K, Na, Li (stronger reducing agent) replaces less reactive metal ion eg. Cu, Ag, Au (weaker reducing agent) in a compound - stronger reducing agent = oxidised - weaker reducing agent = reduced - combustion * reaction with oxygen to form a metal oxide, covalent compound or CO2 and H2O - oxygen = reduced - metal = oxidised - corrosion * degradation of a metal to form a more stable metal oxide when exposed to liquids/gases * metal is more stable

understand that oxidation can be modelled as the loss of electrons from a chemical species, and reduction can be modelled as the gain of electrons by a chemical species; these processes can be represented using balanced half-equations and redox equations (acidic conditions only)

redox reactions: eg. displacement Mg (s) + CuSO4 (aq) --> MgSO4 (aq) + Cu (s) Mg = reducing agent (oxidised) Cu = oxidising agent (reduced) MgSO4 = spectator ion half equations: oxidation = Mg (s) --> Mg2+ (aq) + 2e- reducation = 2e- + Cu2+ (aq) --> Cu (s) * general rule (redox table): - if a species is on the top right, it will spontaneously react with a bottom left species

understand that the ability of an atom to gain or lose electrons can be predicted from the atom's position in the periodic table, and explained with reference to valence electrons, consideration of energy and the overall stability of the atom

usually a metal = oxidised usually a non-metal = reduced ] redox: chemical reaction involving the transfer of electrons from one reactant (reducing agent) to another (oxidising agent) - oxidation: loss of electrons (has a few electrons in the outer shell it wants to lose) - reduction: gain electrons (has almost a full outer shell and it wants to gain a full shell to be more stable) * top right of the periodic table: - these hold onto electrons and accept electrons more easily/readily (reduce easily) = strong oxidising agents - high electronegativity (like to make stable negative anions) * bottom left of the periodic table: - provide their valence electrons more readily (oxidise easily) = strong reducing agent - low ionisation energy (like to make stable positive cations) * valence electrons structure affects the ability to oxidise/reduce