chemistry - unit 3 aos 1

random

• density: p = m/v. also can find density of molecule using periodic table (g/mol)

• when a molecules goes from gas to liquid state, atoms are brought closer together by forming more bonds → forming bonds releases energy, so this is an exothermic process

• check in the molar enthalpy of combustion table what state things are at SLC

2.1 - energy + enthalpy

• all chemicals contain chemical energy stored in the bonds between atoms.

• H - enthalpy/heat content - the chemical energy stored within the bonds of a system, measured in kJ.

  • H_r - enthalpy of reactants

  • H_p - enthalpy of products

• ∆H - enthalpy change - in a chemical reaction, reactants → products. the enthalpy of the products differs from the enthalpy of the reactants, meaning the enthalpy of the system changes. this enthalpy change results in heat energy being absorbed or released to the surroundings. to calculate enthalpy change:

  • ∆H = H_p - H_r

2.1

A - enthalpy of reactants

B - activation energy of forward reaction (Ea)

C - activation energy of reverse reaction

D - change in enthalpy

E - enthalpy of products

2.1 - energy profile diagrams + enthalpy

• activation energy (E_A) - the energy required to break the bonds in the reactants, and therefore commence the reaction. breaking bonds always absorbs energy. the bonds are then formed again to form products, which releases energy. both endothermic and exothermic reactions require activation energy.

2.1 - exothermic reactions

• exothermic reactions

  • reactant energy > products energy → negative ∆H value

  • release energy to the environment

  • releases heat as a byproduct

2.1 - endothermic reactions

• endothermic reactions

  • products enthalpy > reactants enthalpy → positive ∆H value

  • absorb energy from the environment

  • feel cold

2.1 - thermochemical equations

• thermochemical equation - a full, balanced equation, with ∆H value written at end

• if coefficients are multiplied by a factor, the ∆H value is also multiplied by this factor

• the reverse reaction has the same ∆H value but with the sign reversed

2.1 - limiting reactant

• limiting reactant - the reactant that is used up first + determines how much product can be made

• in order to find out how much product is made, you need to first identify the limiting reactant to see how many moles of each reactants react

converting between J, kJ, MJ

2.2 - types of fuels

• fuels - substances with high amount of energy stored in its bonds that can be easily released

• renewable fuel - can be replenished by natural processes within a relatively short period of time → sustainable source of energy

• non-renewable fuel - used at a rate faster than they can be replenished

• two main types of fuel - fossil fuels (non-renewable) and biofuels (renewable)

2.2 - greenhouse gases + climate change

• greenhouse gases

co2 emmissions - excess release of CO₂ into atmosphere → adds to net CO₂ emmissions → enhanced greenhouse effect → extreme weather events + climate change

2.2 - fossil fuels

• 3 main - coal, crude oil/petrol, natural gas

• formed from the decomposition of dead plant and animal matter over millions of years

advantages

• higher energy content than biofuel → reduces cost

• easier to transport + we already have infrastructure for obtaining, transport, processing → reduces cost

• lots of jobs involved → good for economy

disadvantages

• non-renewable → not sustainable

• production + combustion releases CO₂ into atmosphere → adds to net CO₂ emmissions → enhanced greenhouse effect → extreme weather events + climate change

• release more pollutants - ex. photochemical smog such as NO₂, SO₂

2.2 - coal

• coal - fossil fuel formed from the fossilisation of wood + plant material into solid coal. with increased time, pressure + temperature, carbon content increases + energy content increases. (also changes from brown to black)

• brown coal energy content - 17 kJ/g

• black coal energy content - 32 kJ/g

• once obtained, coal is often combusted (coal + O₂ → CO₂ + H₂O) to produce electricity

2.2 - crude oil

• crude oil - fossil fuel often formed from the fossilisation of aquatic life. a liquid mixture of hydrocarbons of varying length. crude oil has no use as a fuel, but it can be separated through fractional distillation, and can be refined further, as many of it’s hydrocarbons can be used as fuels.

• fractional distillation - separates crude oil into fractions depending on boiling point (which depends on chain length, as longer chains have stronger intermolecular forces → require more energy to break → higher boiling point). each fraction still isn’t pure, it is a mixture of hydrocarbons, but they have similar boiling point + therefore chain length. some fractions include petrol and petrodiesel, which are fuels.

  1. crude oil is heated in a furnace until most of it vaporises

  2. vapour then enters a fractioning column that is hot at the bottom and cool at the top

  3. vapour rises up the column, cooling down as it rises due to temperature of fractioning column

  4. condensation at different levels - each substance recondenses at it’s boiling point → high boiling point = high condensation point → condenses lower down.

• once obtained, petrol is often combusted to power vehicles

2.2 - natural gas

• natural gas - gaseous fossil fuel found in deposits in the earth’s crust, which can be accessed by drilling.

• main = methane. CH₄

• once obtained, natural gas is often combusted (natural gas + O₂ → CO₂ + H₂O) to produce electricity, heat homes + cooking

2.2 - biofuels

• 3 main - biogas, biodiesel, bioethanol

• derived from plant materials (biomass)

advantages

• renewable - uses natural processes + made quickly

• although CO₂ is released into the atmosphere during combustion of biofuels, the same amount of CO₂ was removed from the atmosphere during the growth of crops used in the biofuel’s production through photosynthesis → CO₂ removed from atmosphere during plant’s life compensates for release of CO₂ when combusted. means process can be carbon neutral (but realistically, energy and CO₂ emmissions are involved in farming + transport). but still far less greenhouse gases produced than fossil fuels.

• can be produced from material that would otherwise have been waste

disadvantages

• lower energy content than fossil fuels

• less investment in infrastructure → more expensive

• increased demand for biomass → can impact food stability. if crops grown specifically for use as biomass → deforestation, land degradation, loss of biodiversity

2.2 - biogas

• biogas - a mixture of gases (mostly methane) released during the anaerobic breakdown of organic waste by bacteria. bacteria decompose complex molecules in substances such as protein + carbohydrates into the simpler molecular compounds such as CO₂ and CH₄ (methane).

• a digester is a large tank containing bacteria, and is supplied with ex. manure + farm waste to produce biogas. (particularly suited to farms as waste from digester can make a rich fertiliser),

• unfortunately, the infrastructure required to produce biogas is expensive

• once obtained, biogas is often combusted (biogas + O₂ → CO₂ + H₂O) to produce electricity + heat homes

• energy content - 26 kJ/g (lower than natural gas, because methane content is lower)

2.2 - biodiesel

• biodiesel - a mixture of fatty acid methyl esters produced by a chemical reaction between vegetable oil (soybean, canola, palm, animal fats, recycled) and an alchohol such as methanol.

  • straight vegetable oil is too viscous to be used for fuel → has to undergo transesterification to become biodiesel

  • transesterification - process in which the glycerol is split from a triglyceride (one of above oils), and methyl groups (CH₃) are added to the fatty acids to make fatty acid methyl esters, biodiesel. using methanol (alchohol) + KOH (catalyst)

  • triglyceride + methanol → (KOH) → glycerol + fatty acid methyl esters (biodiesel)

• once obtained, biodiesel is often combusted (ester + O₂​→CO₂ ​+ H₂​O) and used to power diesel engine vehicles

• energy content - 41 kJ/g

2.4 - bioethanol

• bioethanol - ethanol (C₂H₅OH) produced from fermentation + distillation of glucose in biomass

  • fermentation - exothermic anaerobic process in which enzymes + microorganisms (yeast) catalyse the fermentation of glucose into an alcohol + carbon dioxide. glucose is partially oxidised and partially reduced. process is carried out at 35℃ (optimal temp for enzymes + microorganisms), and in the absence of oxygen.

  • C₆H₁₂O₆(aq) → 2C₂H₅OH(aq) + 2CO₂(g)

  • after fermentation, the ethanol kills any remaining yeast, however it produces a solution with a concentration of only 10%v/v of ethanol. to separate the ethanol from the water, distillation must be done.

  • distillation - utilises the two different boiling points of the two liquids to separate them. the mixture is heated, and ethanol evaporates first due to it’s lower boiling point than water. this ethanol vapour is then cooled in a condenser and condensed into liquid ethanol. process is often repeated for higher purity.

• once obtained, bioethanol is often combusted (C₂​H₅​OH + 3O₂ ​→ 2CO₂ ​+ 3H₂​O) + used for heat, electricity, or blended with petrol

• energy content - 30 kJ/g

2.3 - fuel sources for the body

• fuel sources that provide the body with energy are known as nutrients. humans acquire nutrients through eating food, macromolecules are then digested, and enter cellular respiration (at different stages), where they are oxidised and converted into energy

• energy value - the energy released when the food is oxidised + converted into energy in the body. (after digestion). it is measured in kJ/g, but can be kJ/mol if the substance is pure (ex. glucose).

• energy content - energy released if the food was combusted

• for most foods, the energy value is similar to the energy content

2.3 - energy from glucose in plants + animals

• glucose - glucose is a carbohydrate and is one of the most important biomolecules for organisms - glucose and it’s polymer starch are more rapidly digested than other foods. it contains chemical energy + is the main source of energy in plants + animals. it can be formed in plants through photosynthesis.

• in plants - plants make their own glucose through photosynthesis. they can use it to produce energy through cellular respiration, or can polymerise it into starch or cellulose to store in their roots (an endothermic process)

• in animals - animals primarily use glucose for short term energy production through transporting it in the blood to body cells for cellular respiration, however they can also polymerise it into glycogen, to store in their cells (an endothermic process).

2.3 - digestion

• animals must digest food before it can undergo cellular respiration to produce energy.

• digestion - an exothermic process that breaks down nutrients into smaller units (monomers) through hydrolosis. carbohydrates → monosaccharides (mainly glucose), fats + oils → glycerol + fatty acids, proteins → amino acids

• hydrolosis -?

• dietary fibre cannot be digested by humans

2.3 - cellular respiration

• glucose (the main form of energy in plants + animals), is converted to energy through cellular respiration (mainly aerobic respiration, where oxygen is required).

• cellular respiration - an exothermic reaction in which oxygen is reduced and glucose (in this case) is oxidised to form ATP, carbon dioxide and water. it is the reverse reaction of photosynthesis. occurs in both plant and animal cells.

• equation (aerobic)

  • glucose + oxygen → carbon dioxide + water + ATP

  • C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O (l) ∆H = -2803 kJ

2.3 - photosynthesis

• plants create their own glucose through photosynthesis

• photosynthesis - an endothermic reaction in which carbon dioxide is reduced and water is oxidised to form glucose and oxygen (biproduct). carried out in the chloroplasts of cells of green leaves. absorbs energy from the sunlight

• equation

  • carbon dioxide + water → sunlight → glucose + oxygen

  • 6CO₂(g) + 6H₂O (l) → C₆H₁₂O₆(aq) + 6O₂(g) ∆H = +2803 kJ

2.3 - energy from carbohydrates

• structure

  • carbohydrates refer to monosaccharides, disaccharides + polysaccharides

  • monomer - monosaccharide

  • polymer - disaccharide/polysaccharide

  • formation of monosaccharides → disaccharides/polysaccharides

    • monosaccharides (simple sugars) are carbon rings, containing C, H and O. includes glucose, fructose and galactose. monosaccharides link together through condensation reactions, forming glycosidic bonds between monosaccharides.

    • if only 2 join → disaccharide. maltose = 2 glucose, sucrose = 1 glucose + 1 fructose, lactose = 1 glucose + 1 galactose

    • if multiple join → polysaccharide. starch = many glucose, glycogen = many glucose, cellulose = many glucose (many carbohydrates are polymers of glucose)

• carbohydrates

  • carbohydrates eaten from food → during digestion, enzymes in saliva + small intestine break polysaccharides (often starch) down into glucose → if not used immediately, glucose is polymerised + stored as glycogen, and later broken back down to glucose (but this doesn’t happen often as the body isn’t great at it) → glucose is transported in the blood to body cells for cellular respiration to be oxidised and converted to energy

  • energy value - 16 kJ/g

  • the body uses carbohydrates for immediate energy needs (bc glucose can be rapidly digested, and the body isn’t great at storing glucose as glycogen).

  • lower energy value - however releases less energy than fats + oils, because carbohydrates are already partially oxidised, and mainly contain C-C bonds which store less chemical energy than C-H bonds in oils + fats.

2.3 - energy from fats + oils

• structure

  • fats and oils are a mixture of triglycerides that are nonpolar (no carbon-carbon double bonds → saturated).

  • named a fat if solid at 25℃, and named an oil if liquid at 25℃. triglycerides obtained from animals are usually fat, and triglycerides obtained from plants are usually oil, hence animal fat and vegetable oil

  • triglycerides - composed of a glycerol molecule connected to 3 fatty acids (carboxylic acids with long hydrocarbon chains). contains C, H, O.

  • formation of triglycerides - glycerol molecule and 3 fatty acids join through condensation reaction, where each fatty acid is connected by an ester bond.

• fats + oils

  • available energy content after digestion - 37 kJ/g

  • high energy value - due to lots of C-H bonds, which store more chemical energy than ex. C-C bonds found in carbohydrates. also, they are stored tightly because the hydrophobic nature of fats + oils prevents water from attaching + disrupting tightly packed formation.

  • the body stores fats + oils in lipid form → long term energy storage (used in endurance exercise, fasting), as well as insulation + padding for the body. this is because although fats + oils produce more energy, they digest slower than carbohydrates

2.3 - energy from proteins

• structure

  • monomer - amino acids

  • polymer - polypeptide chains

  • formation of amino acids → polypeptide chains - amino acids join together via condensation reactions, forming a peptide/amide bond. linked sequence of amino acids forms polypeptide chain. contains C, H, O, N

  • → folded (1 or multiple polypeptide chains) into functional protein

• protein

  • available energy content after digestion - 17 kJ/g

  • body rarely uses protein as a source of energy, due to protein having many other functions in the body, however, when intensive exercise is depleting stored glycogen and lipids → protein can be used.

2.5 - combustion reactions

• combustion reactions - exothermic reactions, often represented with a thermochemical equation, with a negative ∆H value

• burning fuels is a combustion reaction

• complete combustion - oxygen present, makes carbon dioxide and water. as the fuel is completely oxidised, more energy is released than incomplete

• incomplete combustion - oxygen is limited, and carbon monoxide (CO) and/or carbon is produced instead

• balancing complete combustion reactions (always assume complete combustion)

  1. x + O₂ → CO₂ + H₂O

  2. balance C

  3. balance H

  4. balance O

2.5

• energy content/enthalpy - the potential chemical energy of a substance stored in it’s bonds.

• enthalpy change - difference in enthalpy of the products compared to the reactants

• molar enthalpy - the enthalpy of a substance given per mole. H = q/n

• heat of combustion - a specific thermodynamic value. max amount of energy released when 1 mole of fuel undergoes combustion, (exothermic → always negative ∆H), measured in kJ/mol (q/n)

• q/energy released (in kJ) = n x heat of combustion → often need to convert to J

3.1 - Stoichiometry involving combustion of fuels

• stoichiometry

  1. calculate amount in moles of known substance

     • n = m/M

     • n = cV (for a solution)

     • n = V/Vm (for gases, or solutions? in data booklet)

  2. use mole ratio from balanced chemical equation to determine amount in moles of unknown substance

     • x

  3. answer the question

     • m = n x M (finding mass of solids)

     • density: p = m/v

3.2 - Determination of limiting reactants or reagents

• if there are multiple quantities provided in a question, must determine limiting reactant for stoichiometry calculations

• limiting reactant - less moles once divided by it’s coefficient, is completely consumed in reaction

• reactant in excess - more moles once divided by it’s coefficient

• reactant - a starting material that undergoes change during a chemical reaction

• reagant - a substance added to a system to causea chemical reaction.

1. calculate the number of moles of each reactant present

2. divide each amount of reactant by it’s coefficient in the balanced equation, and identify what is smaller (limiting reactant)

3. use the amount of limiting reactant to determine the amount of product formed

3.3 - Calculating heat energy released

• specific heat capacity (c) - a measure of the amount of energy (in joules) required to increase the temperature of a specific quantity (usually 1 gram) by 1℃.

• the specific heat capacity of water is 4.18 J/g/℃ (very high, due to intermolecular forces being strong hydrogen bonds)

• when an exothermic reaction that releases heat energy (such as combustion of a fuel) is carried out under a container of water (can also use bomb calorimeter, less heat loss due to insulation), some of the heat released is transferred to the water → if temp change of water is measured and mass of water is known, water’s known specific heat capacity can be used to determine the amount of energy that the water absorbed and therefore, (approximately) was released in the exothermic reaction = can determine heat of combustion (assuming no energy is lost)

• q = mc∆T

  • q - energy transferred to water, in J (because c is in J)

  • m - mass of water (density is 1g/mL → 1g = 1mL)

  • c - specific heat capacity of water (amount of energy it takes to heat 1 gram of water by 1 degree) = 4.18 L/g/℃

  • ∆T - difference in temperature of water

• the coefficients in a balanced equation show the mole ratio of the amount, in moles, of reactants and products involved in the reaction.

after finding energy released → find heat of combustion/energy content

• energy content = q/∆m

  • q - energy transferred to water

  • ∆m - change in mass (not necessary for all mass to be used up)

• remember energy value (amount of energy the body can turn into energy) is not the same as energy content → determining the heat of combustion of a food can determine it’s energy content, but not it’s energy value.

heat loss → inaccurate results

• substantial energy loss occurs, mainly through heat loss (through air, especially in the experiment done in the open, and also from surface if no lid in bomb calorimeter) + also not all energy released is heat (light?) → energy content is only an estimation

but not a problem in calorimetry…

q_reaction = -q_water

3.4 - Solution calorimetry

heat loss → inaccurate results

• substantial energy loss occurs, mainly through heat loss (through air, especially in the experiment done in the open, and also from surface if no lid in bomb calorimeter) + also not all energy released is heat (light?) → energy content is only an estimation

calorimetry

• calorimeter - an instrument that measures energy changes in a reaction. has a stirrer so water is heated equally. minimises heat loss by having insulated walls and lid. solution calorimeter is one that holds a known volume of water for a reaction to be carried out. bomb calorimeter ?

• in calorimetry, energy content can be more accurately measured, by finding the calibration factor of the calorimeter. calibration involves adding a known quantity of heat energy from an electrical source or from a chemical source to find the realtionship between the energy transferred to the water and the temperature change in the calorimetr (thus accounting for heat loss). CF = how much heat is needed to change the calorimetre temp by 1 ℃, in J/℃

• electrical calibration

  • CF = q/∆T

    • CF - calibration factor, in J/℃

    • q/E - (energy, in J) = V (voltage, in volts) x I (current, in amps) x t (time, in seconds)

• after calibrating and finding CF, a reaction can be carried out, and the energy it releases can be found accurately, accounting for heat loss

  • q = CF x ∆T

    • q - energy, in J, absorbed by calorimeter/released from reaction BUT if the calorimeter absorbs it, that means the reaction released it → q_reaction = -q_calorimeter. double check by seeing if temp decreased or increased

    • CF - calibration factor, in J/℃

    • ∆T - change in temperature of the calorimeter

  • →to find enthalpy of products in reaction

  • ∆H = q/n

3.5 - Energy from fuels and food

• % energy efficiency = (useful energy/energy input) x 100

• the energy value of fuels and foods is measured in kJ/g

4.1 - redox reactions

• redox reactions involve the transfer of electrons

  • OIL RIG

    • oxidation is loss of electrons / or reaction with oxygen (glucose in cellular respiration)

    • reduction is gain of electrons

• oxidation and reduction occur simultaneously as electrons are transferred between the reactants

• oxidant/oxidising agent - substance that causes oxidation + is reducing. reduces in oxidation number

• reductant/reducing agent - substance that causes reduction + is oxidising. increases in oxidation number

4.1 - redox equations

• to balance full equations, you need to pull out conjugate pairs and first balance each half equation using KOHES method

• if there is one product, split into ions for each half equation

1. K - key atoms - balance key atoms (not H or O)

2. O - oxygen - add H₂O atoms to balance oxygens (reaction in solution means H₂O available)

3. H - hydrogen - by adding H₂O → unbalances hydrogen

   • in acidic conditions (H⁺ ions available, assume) - add H⁺ atoms

   • in basic conditions (OH^- ions available) - indirectly add H⁺ atoms - for every H⁺ atom required, add a H₂O atom to this side, and an OH^- atom to the other → then cancel out waters that appear on both sides

4. E - electrons - calculate the overrall charge of the LHS and RHS (remember to account for coefficients) → add electrons to more positive side to balance charge

5. S - states

combining

• full equations can’t have electrons, because it must show that all the electrons have been transferred

• → half equations must have same amount of electrons so they can cancel out in the full equation. → entire half equation/s must be multiplied by a factor to get same amount of electrons → combine equations together + cancel out electrons

• after combined, cancel out anything that appears on both sides (often H₂O and ions), if there’s less on one side, remove that amount from both sides

4.2 - oxidation numbers

• oxidation numbers - assigning imaginary numbers to individual atoms in an equation

  • purpose - by comparing oxidation number of an atom before + after the reaction → determine if a redox reaction occurred + determine which is the oxidant and reductant (particularly helpful when it is difficult to see ex. SO₂ → SO₃)

  • oxidant/substance reducing - reduces in oxidation number

  • reductant/substance oxidising - increases in oxidation number

oxidation numbers rules

• all atoms are treated individually, even if in a polyatomic molecule or ionic compound

• disregard coefficients, but not subscripts

• while charges have the sign after (ex. 2-), oxidation numbers have the sign before (ex. -2)

• if it asks you what is reducing and oxidising but it’s an atom in a polyatomic ion, say ex. S in H₂SO₄. it is always the reactant not the product oxidising/reducing!!

• if it is asking you the reductant or oxidant, name the whole molecule

• O and H are not usually oxidised/reduced, so unnecessary to find their oxidation numbers unless no other atoms change oxidation number

• if an element appears in multiple molecules on one side, then find each individual oxidation number. if all increase → oxidation. if all decrease → reduction. if mixed, disproportionation (don’t need to know)

• free element: oxidation number = 0

  • Na = 0

  • Cl₂ = 0

• monatomic ions: oxidation number = charge of the ion

  • Na⁺ = +1

  • Cl^- = -1

  • Mg²⁺ = +2

• polyatomic ions: if there are brackets, first determine the charge of each ion

  • Fe(ClO₂)₃ = Fe has a charge of +3, and ClO₂ has a charge of -1

• polyatomic ions: sum of oxidation numbers = charge of whole

  • sum of NaCl = 0

  • sum of SO₄ ^2- = -2

• polyatomic ions: must assign oxidation numbers following an order

  1. most electronegative (F = -1, O= -2, Cl = -1, N = -3, S = -2)

  2. group 1 and 2 elements (group/column 1 ex. H all have 1 valence electron → +1)

  3. last element is what you make it get back to overall charge

  • O in SO₄^-2 = -2 → -2 × 4 = -8 → S must be +8 but bc overrall charge is -2 → S = +6

• polyatomic ions: general rules

  • H is always +1, except it’s -1 in hydrides (when bonded to metals ex NaH)

  • O is always -2, except it’s +2 in OF₂ and it’s -1 in peroxides (molecules with O-O single bond, such as H₂O₂)

  • F is always -1

5.1 - galvanic cells

• galvanic cell - generates electricity from spontaneous, exothermic redox reactions, converting stored chemical energy into electrical energy.

• chemical energy → electrical energy vs

• chemical energy → heat energy → mechanical energy → electrical energy (combustion)

• =galvanic cells are more energy efficient than other fuels that undergo combustion because there are no intermediate steps where energy is lost. also, most of the energy in a galvanic cell is captured as electricity. fuel cells are even more energy efficient.

• REDCAT ANOX - reduction at cathode, oxidation at anode

• consists of two half cells, each containing an electrode in contact with a solution. the species in each half cell form a conjugate redox pair. if one reactant in a conjugate pair is a solid metal, then that will be the electrode. if neither reactant is a solid metal, then an inert/non-reactive electrode (ex. platinum, graphite) will be used

• if there’s multiple reactants present, the two reactants that are furthest apart and happen spontaneously will react / the strongest reductant will react with the strongest oxidant

• to calculate maximum voltage of the cell, find the potential difference (E⁰_cell) of the reactants

• heat can be used as a catalyst to speed up reaction

• primary galvanic cells - non-rechargeable, because products migrate away from the electrodes after the reaction

• secondary galvanic cells - rechargeable

• galvanic cells are used in non-rechargeable (primary) alkaline batteries. they are small + portable, and power many things such as tv remotes + smoke detectors

how it works

• if reactants came into physical contact (direct spontaneous redox reaction)

  • electrons are directly transferred

  • charges are balanced automatically at the reaction site (unlike galvanic)

  • but lots of heat/thermal energy is produced as a byproduct

• galvanic cell (indirect spontaneous redox reaction)

  • half equations are separated by making them occur in separate half cells

  • electrons flow through an external circuit from the anode to the cathode, so electrical energy can be captured

  • but this causes charge inbalance:

    • at anode (oxidation) - metal atoms lose electrons → form positive ions → solution becomes more positively charged

    • at cathode (reduction) - positive ions gain electrons → positive ions removed from solution → solution becomes more negatively charged

  • = internal circuit prevents the accumulation of charge in either half cell, by providing cations to the cathode and providing anions to the anode.

    • salt bridge must contain free-moving ions from a soluble salt, that is relatively unreactive (doesn’t react with one of reactants, or form a precipitate)

  • both the external circuit and internal circuit carry electric current. external circuit = electron current. internal circuit = ionic current.

• progression of reaction in galvanic cell

  • at start, big difference in chemical potential energy. electrons want to move → current flows (+ salt bridge allows for continued flow)

  • over time, electrons are transferred + difference shrinks. electrons no longer want to move → no electron flow

  • = in galvanic cells, the reaction doesn’t last forever. → the cell must be disposed (usually contain toxic chemicals) or recharged, until they eventually die + need to be disposed.

observations

• when something is reducing, the mass of the electrode increases because solid forms on it (if solid in half equation), and the mass/volume of the solution decreases

• when something is oxidising, the mass of the electrode decreases and the mass/volume of the solution increases

5.2 - electrochemical series

• electrochemical series - a set of reduction half equations ordered by reactivity. each equation reversed is it’s oxidation reaction.

  • top = most readily reduced/strongest oxidant, least readily oxidised/weakest reductant. highest E⁰ value

  • bottom = least readily reduced/weakest oxidant, most readily oxidised/strongest reductant. lowest E⁰ value

• E⁰ value - standard reduction potential of a particular reduction half equation. measured in volts. relative to the H⁺/H₂ reduction reaction set at 0V

• E⁰_cell - the potential difference of a redox pair. measured in volts. use to find maximum cell voltage in a galvanic cell.

  • equation: top E⁰ value - bottom E⁰ value.

• a redox reaction will only occur spontaneously if the reducing half equation (forward) is above the oxidising half equation (backward), meaning it has a higher E⁰ value

• if there are multiple reactants in a cell, the strongest oxidant will react with the strongest reductant (if possible)

• limitations of the electrochemical series

  • gives no indication of the reaction rate → reaction rate could be so slow that the reaction can’t be observed, but the observers wouldn’t know if the reaction is just very slow or not happening at all. also, when comparing two redox reactions, can’t determine what is faster (one may have more potential but have a much slower rate of reaction/release less energy in a period of time)

  • the series is written assuming SLC (100 kPa, solids are pure, solutions at 1M concentration, 25℃) → changing concentrations of solutions and pressure pf gases changes E value → if conditions are changed and E value changes, the order of half-reactions in the electrochemical series may also be different → predictions of reactions based on the standard electrode potentials may not be reliable

most reactive = bottom

5.3 - fuel cells

• fuel cell - a type of galvanic cell (generates electricity from spontaneous redox reactions) directly converts chemical energy → electrical energy. the reactants (usually gases) are supplied continuously, allowing constant production of electrical energy.

• structure

  • two gas compartments

    • must be separate (thanks to electrodes + electrolyte solution) to ensure hydrgoen and oxygen gas don’t come into contact, because when they do they react explosively.

  • two electrodes - anode at the fuel compartment + cathode at the oxygen compartment. electrodes must be:

    • unreactive - musn’t interfere with reaction. ex. graphite, platinum.

    • (often) catalytic - as rate of reaction is slow, electrodes must be catalytic to increase rate of reaction. platinum is effective at this. (galvanic cells do not use catalysts!)

    • porous - so that reactants can come into contact with the electrolyte. also increases surface area and efficiency

    • conductive - so that electrons can flow through the external circuit

  • electrolyte solution (contains water + ions) - carries ions from one electrode to the other

    • the ions produced at one electrode will migrate to the other through the electrolyte (internal circuit), thus changing the pH of the electrode. once the rate at which ions are produced and consumed is equal, the pH of the electrodes will stay constant.

• fuels - always oxidised at the anode. hydrogen is commonly used as a fuel, as the only products are electricity, heat and water. but carbon-based fuels can also be used, however they produce CO₂.

• oxygen - always reduced at the cathode.

• fuel cells are mainly classified by the acidity of the environment

acidic fuel cell

• H⁺ available in electrolyte → oxygen reduction half equation is O₂(g) + 4H⁺(aq) + 4e^- → 2H₂O(l)

alkaline fuel cell

• OH^- available in electrolyte → oxygen reduction half equation is O₂(g) + 2H₂O(l) + 4e^- → 4OH^-(aq)

5.3 - fuel cell advantages

• can continuously produce electricity as long as gases (fuel + oxygen) are supplied

• high energy efficiency → less fuel used to generate same amount of energy

  • convert chemical energy → electrical energy - less intermediate steps than traditional engines in which energy is lost. (also in galvanic cells)

  • most energy is converted into electrical (also in galvanic cells), less heat loss than combustion

  • water produced can be reused or used as a coolant

  • some modern fuel cells also use reuse waste heat by turning it into steam

• environmentally friendly

  • main product of process is water (if fuel is hydrogen). no greenhouse gases released, while combustion of fossil fuels releases lots of greenhouse gases

  • hydrogen can be produced in an environmentally friendly manner (if electrolysis of water is powered by renewable energy), while oil extraction and refining is very damaging.

• no moving parts → quiet + reliable

• other fuels (such as methanol) can also be used

• electricity is generated on site → no connection to electricity grid is needed

• hydrogen cars recharge faster than electric cars + can store more energy bc hydrogen is so light)

hydrogen sustainability

5.3 - fuel cells disadvantages

• rate of reaction is generally slow

• catalytic electrodes (platinum) are expensive

• requires a constant supply of fuel

• storage + transport of hydrogen is difficult - compressing hydrogen to store or storing as liquid is expensive + difficult. also unsafe, as hydrogen is extremely flammable + explosive.

• much of hydrogen production is powered by fossil fuels → energy loss + greenhouse gas emmissions

• if carbon based fuel used, CO₂ is produced

• some types of cells use toxic electrolytes, and their electrodes contain rare, expensive or harmful materials.

• there is no existing infrastructure to support fuel cells → fuel cells are not used widespread

5.4 - most environmentally friendly fuels