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Chemistry - Coggle Diagram

- - - - Acid- donates a proton
        Base- accepts a proton
        
        Requires an acid base pair
        
        Conjugate acid-base pairs
        
        To summarise, in the forward reaction:
        The NH4+ ion is the conjugate acid of the base NH3
        The OH− ion is the conjugate base of the acid H2O
        
        In the reverse reaction:
        NH3 is the conjugate base of the acid NH4+
        H2O is the conjugate acid of the base OH−
    - - Substance that can act as an acid or a base
    - - metal + acid → hydrogen gas + salt
    - - Bases include metal oxides, metal hydroxides and ammonia
        Alkalis are limited to soluble in water
        
        Metal oxide + acid → salt and water
        
        Metal hydroxide + acid → salt and water
    - - Acid + base → salt + water + energy
        
        Antacids
        
        Antacid tablets can be bought at the supermarket or pharmacists to ease the symptoms of excess stomach acid that causes indigestion and heartburn. The active ingredients of most antacid tablets are metal carbonates or hydrogen carbonates, NaHCO3, CaCO3, MgCO3, for instance, or insoluble metal hydroxides such as Mg(OH)2 or Al(OH)3
    - - Metal carbonate + acid → salt + water + carbon dioxide
        Metal hydrogen carbonate + acid → salt + water + carbon dioxide
        
        This is because H2CO3 is the product of the double replacement but is unstable so it decomposes to water and carbon dioxide
  - - - -ite → -ous acid
      - -ide → hydro- -ic acid
      - -ate → -ic acid
  - - - this type of reaction can be carried out using a simple calorimeter. Heat is released when an acid and base react together, which means it is an exothermic reaction. The enthalpy change that occurs is known as the enthalpy change of neutralisation; it is the enthalpy change when one mole of water is formed in the reaction of an acid and a base.
  - - - pH = −log[H+(aq)]
        
        Conversely, the lower the [H+], the higher the pH value. In terms of hydroxide ions, a higher concentration of OH− ions, [OH−], corresponds to a higher pH value and a lower [OH−] corresponds to a lower pH value
        
        pOH = −log[OH−(aq)]
        
        an aqueous solution with a pH of 3 is 10 times more acidic than an aqueous solution with a pH of 4; and 100 times (or 10 × 10) more acidic than an aqueous solution with a pH of 5.
    - - Universal indicator
        
        Universal indicator is actually a
        mixture of indicators that
        produces different colours
        in solutions of different pH
        
        A pH probe and meter produces
        a more accurate method of
        measuring pH, provided
        they are correctly calibrated
      - Litmus papere
        
        Acid - red
        Alkali - Blue
      - Indicators are weak acids or bases with the dissociated and undissociated forms having different colours. Like all weak acids and weak bases, they exist in an equilibrium between the two forms, which are represented as HIn (the undissociated form) and In− (the dissociated form)
        
        HIn (aq) ⇌ H+ (aq) + In− (aq)
        Colour A Colour B
        
        the position of equilibrium can be changed by adding an acid (H+ ions) or a base (OH− ions)
        
        phenolphthalein is an appropriate indicator for a titration of a strong base with a weak acid, methyl orange is not. This is because the pH range over which it changes colour (its end-point) does not fall within the vertical region of the curve.
    - - If the acid is weak, the equilibrium concentration of the undissociated acid molecules can be taken as the initial concentration of the acid, [HA]eq = [HA]initial.
        
        The assumption here is that very few acid molecules have actually dissociated.
  - - - The position of equilibrium for the dissociation of H2O lies very far to the left, so the concentration of the water is effectively constant.
        
        At 298 K (25˚C), pure water has [H+] and [OH–] of 1.00 × 10–7 mol dm–3
        kw= 1*10^-7
        
        By applying Le Chatelier’s principle, we can deduce that increasing the temperature favours the forward reaction, as the forward reaction is endothermic
        
        Hence, at temperatures above 25oC, the pH of pure water will be less than 7 (a value that we normally think of as slightly acidic), even though the solution is still neutral as [H+] = [OH–].
        
        pKw, pH and pOH
        
        By taking the log10 of both sides of the Kw expression, we get the following relationship:
        pKw = pH + pOH
        
        At 298 K pH + pOH = 14.00
  - - - HA (g or l) → H+ (aq) + A– (aq)
        This reaction can also be written involving the formation of the hydronium ion (H3O+):
        HA (g or l) + H2O (l) → H3O+ (aq) + A– (aq)
        
        Strong acids:
        
        Chloric acid: HClO3.
        
        Hydrobromic acid: HBr.
        
        Hydrochloric acid: HCl.
        
        Hydroiodic acid: HI.
        
        Nitric acid: HNO3.
        
        Perchloric acid: HClO4.
        
        Sulfuric acid: H2SO4.
    - - In other words, the equilibrium position lies very much to the left and the value of the equilibrium constant, Kc, for the dissociation is small.
        HA (aq) ⇌ H+ (aq) + A– (aq)
        HA (aq) + H2O (l) ⇌ H3O+ (aq) + A– (aq)
        
        Distinguishing between weak and strong acids and bases
        
        Weak and strong acids and bases of the same concentration (equimolar solutions) can be distinguished due to their different properties such as pH, electrical conductivity and, in the case of acids, by their reactions with active metals.
        
        Weak acids and bases only partially ionise in solution, therefore, The weaker acid will have a smaller pH
        
        Weak acids and bases have a lower concentration of mobile ions in solution because they only partially ionise in solution. For this reason, they are poor conductors of electricity
        
        Finally, we have the reaction with an active metal such as zinc or magnesium. Active (or reactive) metals are those above hydrogen in the activity series. When the reaction of a weak acid with magnesium is compared to that of a strong acid with magnesium, the reaction with the strong acid will be a faster, more vigorous reaction
        
        dissociation constants
        
        Acid
        
        Base
        
        pKa and pKb
        pKa = –log10Ka and pKb = –log10Kb
        Ka = 10–pKa and Kb = 10–pKb
        
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        Conjugate pairs
        
        Strength of conjugate acids and bases
        
        The stronger the acid the weaker the conjugate base will be.
    - - The general equation for the dissociation of a strong base, MOH, is shown below.
        MOH (aq) → M+ (aq) + OH– (aq)
        
        All bases, except for the hydroxides of groups 1 and 2, are weak
      - Weak bases are made up of molecules that react with water to release hydroxide ions. The general equation for the dissociation of a weak base, B, is shown below
        B (aq) + H2O (l) ⇌ BH+ (aq) + OH– (aq)
  - - - Nitrogen monoxide
        
        Radical formation - naturally occurs by lighting storm, also byproduct of internal combustion engines- diesel engines contribute more due to higher temperatures
        
        N2 (g) + O2 (g) → 2NO• (g) (requires high temperatures)
        2NO• (g) + O2 (g) → 2NO2• (g)
        Acid rain formation
        2NO2• (g) + H2O (l) → HNO3 (aq) + HNO2 (aq)
        or
        4NO2• (g) + 2H2O (l) + O2 (g) → 4HNO3 (aq)
        
        Acid rain
        
        Rain naturally has pH of 5.6 (due to CO2 in the atmosphere)
        CO2 (g) + H2O (l) ⇌ H2CO3 (aq)
        Acid rain is when less than 5 pH caused by NOx and SO2,
        
        Acid deposition effects
        
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      - Sulfur dioxide
        
        Formation of sulfur dioxide- naturally occurs when released from volcanic eruptions. Also from coal and oil power stations
        
        S (s) + O2 (g) → SO2 (g)
        Photochemical oxidation (net reaction):
        2SO2 (g) + O2 (g) → 2SO3 (g)
        Acid rain formation
        SO2 (g) + H2O (l) → H2SO3 (aq)
        SO3 (g) + H2O (l) → H2SO4 (aq)
  - - - . In this complex ion, the ligands (water molecules) are acting as Lewis bases and the central metal ion (Cu2+) is acting as a Lewis acid
        
        Species called electrophiles act as Lewis acids and nucleophiles act as Lewis bases.
  - - - The curve starts at pH 3, which is the pH of the weak acid used to produce this pH curve. At first, the addition of the strong base results in the pH staying fairly constant; this is known as the buffer region. Then there is a sharp increase in pH from 7 to 11 and the curve flattens out at pH 13
        
        The half-equivalence point can be found on the x-axis at the point which is half the volume added at the equivalence point. On the pH curve in Figure 3, this corresponds to the 25 cm3 point. At the half-equivalence point, exactly half the acid has reacted with the base and been neutralised, with the other half remaining unreacted.
    - - more complex possibilities do exist, such as the titration of sodium carbonate solution with hydrochloric acid (two equivalence points), or the reaction between phosphoric acid and sodium hydroxide solution (three equivalence points
      - conductometric titration. As the acid and base react, the conductivity of the solution decreases as H+ and OH− ions react to form water molecules. This produces a V-shaped line
  - - - Some salts, however, dissolve in water to produce either acidic or alkaline solutions. This occurs because one of the constituent ions of the salt reacts with water to release an excess of either hydroxide or hydrogen ions
        
        Salts of strong acids and strong bases
        
        When a strong acid and strong base react, the resulting conjugate acids and bases are both weak, which means that neither is able to hydrolyse water. The solution will, therefore, have a pH of 7 at 298 K
        
        Salts of weak acids and strong bases
        
        An example of hydrolysis by the salt formed from a weak acid and strong base is that of sodium ethanoate (NaCH3COO). This salt fully dissociates in solution to form sodium ions and ethanoate ions:
        NaCH3COO (aq) → Na+ (aq) + CH3COO− (aq)
        
        The ethanoate ion (CH3COO−) is the conjugate base of a weak acid and is strong enough to hydrolyse a water molecule:
        CH3COO− (aq) + H2O (l) ⇌ CH3COOH (aq) + OH− (aq)
        
        The increasing concentration of OH− ions causes an increase in the pH of the solution. Therefore, a solution of a salt formed by the reaction of a weak acid and a strong base will have a pH greater than 7 at 298 K.
        
        Salts of strong acids and weak bases
        
        In the reaction of a strong acid and a weak base, the conjugate acid of the weak base is strong enough to hydrolyse water. For example, the reaction of HCl and NH3 produces the ammonium ion according to the following reaction:
        HCl (aq) + NH3 (aq) → NH4+ (aq) + Cl− (aq)
        
        The ammonium ion is strong enough to hydrolyse water:
        NH4+ (aq) + H2O (l) ⇌ NH4OH (aq) + H+ (aq)
        
        The pH of the solution decreases because of the increase in concentration of H+ ions. A solution of a salt formed by the reaction of a strong acid and a weak base will have a pH less than 7 at 298 K.
        
        If the cation is a metal ion with a high charge density (Fe3+, Al3+ or Cr3+) at the centre of a hydrated complex ion, the solution will be acidic as a result of hydrolysis of a ligand water molecule (Figure 1):
        [Fe(H2O)6]3+ (aq) + H2O (l) ⇌ [Fe(H2O)5(OH)]2+ + H3O+ (aq)
        
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        Salts of weak acids and weak bases
        
        The pH of a salt formed by the reaction of a weak acid and a weak base depends on the Ka and Kb values of the acids and bases involved. If the parent acid and base have approximately the same value, neither ion will hydrolyse and the salt produced will have a pH of around 7.0
  - - - A buffer solution resists changes in pH when small amounts of acid or base are added.
        
        Composition of an acidic buffer
        
        An acidic buffer can be made by mixing together a weak acid and the salt of a weak acid and a strong base
        An example is mixing equimolar solutions of ethanoic acid and sodium ethanoate.
        
        The equation for the dissociation of ethanoic acid is shown below, with the position of equilibrium lying to the left.
        CH3COOH (aq) ⇌ CH3COO– (aq) + H+ (aq)
        
        Sodium ethanoate fully dissociates according to the following equation:
        NaCH3COO (aq) → Na+ (aq) + CH3COO– (aq)
        
        The mixture therefore contains relatively high concentrations of both undissociated CH3COOH molecules and CH3COO− ions (from the salt), which react with the added H+ or OH− ions. When a small amount of acid is added, the H+ ions will combine with the CH3COO– ions to form CH3COOH, removing most of the added H+ ions and resisting a change in pH
        CH3COO– (aq) + H+ (aq) ⇌ CH3COOH (aq)
        
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        Composition of a basic buffer
        
        A basic buffer can be made by mixing together the same amounts of a weak base and the salt of a weak base and a strong acid, for example, ammonia (NH3) and ammonium chloride (NH4Cl)
        
        Ammonia reacts with water according to the following equation (note that like the equilibrium lies to the left as NH3 is a weak base):
        NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH– (aq)
        
        The salt, ammonium chloride, fully dissociates into ammonium ions and chloride ions:
        NH4Cl (aq) → NH4+ (aq) + Cl– (aq)
        
        When a small amount of acid is added, the H+ ions will combine with the NH3 to form NH4+, removing most of the added H+ ions and resisting a change in pH:
        H+ (aq) + NH3 (aq) → NH4+ (aq)
        
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        Alternate ways of making a buffer solution
        
        Generally, buffer solutions can be made by mixing a weak acid or weak base with an equimolar solution of a salt containing its conjugate acid or base
        
        An alternative approach, which results in an acidic buffer solution, is to start with a weak acid and add half as many moles of strong base. Similarly, a basic buffer can be prepared by starting with a weak base and adding half as many moles of strong acid
        
        For example, if 1.0 mol of ethanoic acid is reacted with 0.50 mol of sodium hydroxide, the resulting solution would contain 0.50 mol of ethanoic acid and 0.50 mol of sodium ethanoate:
        
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- - - - This depends on the context
        
        Link Title
        
        Sankey diagram
      - Other possible requirements
        
        Requires:
        Reasonable rate of transfer
        Minimum pollution
        Prefers:
        Cheap
        Easy
        Plentiful
      - Renewable energy source
      - Nonrenewable energy source
  - - - The energy density of a fuel is the energy stored per unit volume.
        
        Energy density = Energy released from fuel / Volume of fuel consumed
    - - The specific energy of a fuel is the energy produced per unit mass. The higher the specific energy, the more energy can be stored per unit mass.
        
        Specific energy = Energy released from fuel / Mass of fuel consumed
  - - - Pros
        
        100 year supply
        
        World wide
        
        Convertible
        
        Gas and liquid fuel
        
        feedstack
      - Cons
        
        CO2 emmission
        
        Fly ash
        
        Acid rain
        
        difficult to transport
        
        Mining process
        
        Dangerous
      - Coal classification
        
        the process of converting coal into syngas or synthetic natural gas (CH4)
        
        Syngas (synthesis gas) is a mixture of carbon monoxide (CO) and hydrogen gas (H2)
        
        C(s) + H2O(g) → CO(g) + H2(g)
        
        Syngas can also be converted to synthetic natural gas (CH4). Synthetic natural gas (CH4) is formed by reacting carbon monoxide with hydrogen gas
        
        CO(g) + 3H2(g) → CH4(g) + H2O(g)
        
        Coal can also be reacted directly with steam to form synthetic natural gas. The equation below shows the reaction of carbon (coal) with steam to form synthetic natural gas.
        
        2C(s) + 2H2O(g) → CH4(g) + CO2(g)
        
        Pros
        
        coal gasification are the lower SOx, NOx and particulate emissions from burning coal-derived gases
        
        Syngas is also a useful intermediate for the production of compounds such as methanol and ammonia. It can also be converted to liquid hydrocarbons via a multi-step mechanism known as the Fischer-Tropsch process
        
        Coal combustion reactions vary:
        
        C+O2→CO2
        
        2C+O2→2CO
        
        C + CO2 → 2 CO
        
        C+ H2O → CO + H2
        
        C + 2H2 → CH4
        
        CO is extremely toxic; to reduce it there are two primary methods:
        
        CO + H2O → CO2 + H2
        
        CO +3 H2 → CH4 + H2O
      - Coal liquefaction
        
        Covert coal from coal → liquid hydrocarbons
        
        Direct
        
        reacting coal with hydrogen in the presence of heat and a catalyst, producing liquid hydrocarbons
        
        nC(s) + (n+1)H2(g) → CnH2n+2(l)
        
        Indirect
        
        coal is first converted into syngas and then into a liquid hydrocarbon via the Fischer-Tropsch process
        
        nCO(g) + (2n+1)H2(g) → CnH2n+2(l) + nH2O
      - Carbon Footprint
        
        the total amount of greenhouse gases produced to directly and indirectly support human activities
        
        In reality, calculating a carbon footprint is difficult for the following reasons:
        
        It is difficult to decide which specific carbon-based greenhouse gases to include in the calculation and whether or not non-carbon based greenhouse gases should be included.
        
        Whether or not carbon monoxide (CO) from incomplete combustion should be included, as it can be oxidised to form carbon dioxide.
        
        Lastly, the difficulty in summing up the steps in a process that results in the production of carbon dioxide.
        
        Calculate the mass of CO2 produced by the combustion of 1.00 × 106 kg (1000 tonnes) of anthracite coal (assume that anthracite is 80% carbon by mass).
    - - Fractional distillation of crude oil
        
        https://app.kognity.com/study/app/2024-chemistry-hl/sid-48-cid-170435/book/fractional-distillation-of-crude-oil-id-11997/
        
        Octane number
        
        The octane number is a measure of a fuel’s ability to resist auto-ignition under pressure (and knocking) in a car engine. The higher the octane number, the greater the resistance to auto-ignition and knocking
        
        High octane number include feature:
        
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        System based on heptane (0) to trimethyl pentane (100)
      - Pros
        
        Easy to transport
        
        Good for cars and planes, etc
      - Cons
        
        CO2 emmissions
        
        Acid rain
        
        Damages envrionment
        
        Only decades left of it
        
        Limited location for extraction
      - Oil and natural gas were formed from ocean-dwelling plants and animals. When these organisms died, they sank to the bottom and became buried under layers of sediment. As the layers of sediment increased, the heat and pressure also increased, which, combined with anaerobic conditions, produced oil and natural gas.
        
        Coal, was formed from the remains of prehistoric plants that grew in vast swamps. When the plants died, they became buried under layers of mud and sand. Over millions of years, heat and pressure, combined with the anaerobic conditions, converted the decaying plant material into coal.
    - - Pro
        
        Clean fuel
        
        Easy transportation
        
        Pipes or tankers
        
        High specific energy
        
        Low CO2/KJ energy release
      - Cons
        
        Less than 100 years supply
        
        Limited location
        
        explosion risks
        
        Poor storage
  - - - an example of an artificial (induced) transmutation - where one element is transmuted into another through artificial means.
    - - he unified atomic mass unit (u) using a neutral carbon-12 atom as our standard of precisely 12.000000 u.
        
        1u=1.661*10^-27kg=931.5MeVC^-2
    - - What E = mc2 means is that mass and energy are interchangeable.
        
        The difference in mass between the reactants and the products is called the mass defect ∆m.
        
        2(1H) + 2n → 4He.
      - Binding energy
        
        The binding energy Eb of a nucleus is the amount of work or energy that must be expended to pull it apart.
        
        Disassembling 4He will require energy.
        
        Assembling 4He will release energy.
        
        Binding energy per nucleon
        
        This number tells us about how difficult it is to remove each nucleon from the nucleus.
        
        Half Life
        
        But random though the process is, if there is a large enough population of an unstable nuclide, the probability that a certain proportion will decay in a certain time is well defined.
        
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        Decay rates are measured using various devices, most commonly the Geiger-Mueller counter.
        ∙Decay rates are measured in Becquerels (Bq).
        
        1Bq= 1 beat per minute
        
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        The decay constant, symbol lambda (λ), is the probability per unit time that a given nucleus will decay
        
        t1/2 = ln2/λ
        
        Rearranging gives the equation:
        
        λ = ln2/t1/2
        
        ionising radiation
        
        This is dangerous radiation that can damage cells in the human body.
        
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        Link Title
    - - In the animation, 235U is hit by a neutron, and capturing it, becomes excited and unstable:
        ∙It quickly splits into two smaller daughter nuclei, and two neutrons.
        
        If each of these neutrons splits subsequent nuclei, we have what is called a chain reaction.
        
        Subcritical mass, if neutrons escape the rxn
        
        Critical mass, when the mass of radioactive isotope is high enough to self sustain
        
        isotope separation
        
        two types of isotope separation used in uranium enrichment are gaseous diffusion and centrifugation
        
        In centrifugation, gaseous uranium hexafluoride (UF6) is fed into a centrifuge and rotated at very high speeds of up to 90,000 rpm
        
        The heavier 238UF6 molecules are forced to the edge of the centrifuge by the centripetal force and are removed.
        
        both of these processes result in enrichment off uranium 235
        
        The lighter 235UF6 molecules remain at the centre of the centrifuge.
        
        Gaseous diffusion uses Graham's law of effusion to separate the isotopes according to their masses
        
        Effusion is the process by which a gas escapes through a small hole in a container. Graham's law of effusion states that the rate of effusion of a gas is inversely proportional to the square root of its molar mass
        
        Recall from topic 5 that the absolute temperature of a gas is directly proportional to the average kinetic energy of its particles. Therefore, two different gases at the same absolute temperature have the same average kinetic energy
        
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        Comparison of uranium dioxide and uranium hexafluoride
        
        Before enrichment, UO2 must be converted to gaseous uranium hexafluoride (UF6), which has a much lower boiling point
        
        UO2
        
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        UF6 is octahedral cova;ent structure that
      - Note that the splitting was triggered by a single neutron that had just the right energy to excite the nucleus.
    - - is the combining of two small nuclei into one larger nucleus.
        In order to fuse two nuclei you must overcome the repulsive Coulomb force and get them close enough together that the strong force takes over.
      - The reason fusion works on stars is because their intense gravitational fields can overcome the Coulomb repulsion force between protons.
        
        The energy released by fusion prevents the gravitational collapse from continuing.
        
        As the fusion progresses different elements evolve, and therefore different fusion reactions occur.
        ∙At the same time, the mass of the star decreases a bit according to E = mc2
        
        As the star evolves along the fusion side of the Eb / A curve, it slowly runs out of fuel to fuse and gravity again wins out.
        For a star with the mass of the Sun, the shrinking will eventually stop and the sun will become a white dwarf and slowly cool down.
        
        iron catastrophe.
        
        For a more massive star, there is enough gravity to fuse the elements all the way up to iron.
        Since the radiation pressure now ceases, gravity is no longer balanced and the star collapses into a neutron star.
    - - Function off of nuclear fission reactions
        
        235U has a very high energy density and specific energy, which make it a very useful fuel. In a nuclear power station, the heat produced by a nuclear fission reaction is used to boil water. Similar to a conventional power station, the steam produced is used to drive a turbine which drives a generator that generates electricity
        
        Contain control rods that absorb neutrons to prevent and uncontrolled chain reaction. They keep the average number of neutrons to the right level to maintain the power plant but not lose control
        
        Fast breeder reactors
        
        a nuclear reactor that generates more fissile material than it consumes
        
        reactions take place in a fast breeder reactor to convert non-fissile 238U into fissile 239Pu.
        
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        were developed because the supply of the fissile isotope 235U is limited.
      - Nuclear Power problems
        
        Nuclear waste
        
        Sources: powerstation, research facilities, military establishments, industry, and hospitals
        
        Low level: low activity or short half-life, can be stored on site or shipped for specialized disposal, can be burned and ashes disposed, can be buried underground
        
        High level: high activity or long half-life, initially stored in cooling pods then can be sent to dry storage casks made of thick concrete or reprocessed to make fuel. No long term solution for the stored waste
        
        Health issues
        
        Meltdown risks
        
        Nuclear weapon fuel generation
  - - - Requires: Sunlight, CO2, H2O, chlorophyll
        Sunlight + 6CO2 + 6H2O → C6H12O6 + 6O2
        
        Note the presence of many alternating single and double bonds, which is known as a conjugated system, or conjugation. This is the part of the molecule responsible for absorbing visible light.
        
        Pigment molecules absorb light in the visible region of the electromagnetic spectrum, between 400 and 700 nm. Chlorophyll appears green because it absorbs wavelengths of light that correspond to blue and red light and reflects the remaining wavelengths
        
        Pigment molecules have extensive conjugation which consists of many alternating single and double bonds. The electrons in the conjugated system are delocalised, which means they are free to move within the conjugated system
        
        Sunlight causes the electrons to become excited and transition to higher energy levels. The energy absorbed when the electrons transition to higher energy levels corresponds to the wavelengths of visible light.
        
        Longer conjugation = longer wavelength = lower energy
      - Biofuels
        
        Biofuels are fuels produced from plant material. The energy in biofuels is obtained from biological carbon fixation. Carbon fixation is a process that takes inorganic carbon (CO2) and converts it into organic compounds such as glucose (C6H12O6)
        
        The carbon dioxide that is absorbed during photosynthesis is released when the biofuels are used. It is for this reason that, in theory, biofuels are carbon neutral
        
        fossil fuels are used to fuel the farm machinery and to make the fertilisers that are necessary to grow the crops. Therefore, biofuels are not completely carbon neutral, but they do add less carbon dioxide to the atmosphere
        
        Fermentation of glucose
        
        General reaction for generating biofuels is:
        C6H12O6 → 2 C2H5OH + 2 CO2
        
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        Transesterification
        
        Biodiesel is a biofuel that can be produced from vegetable oil, fish oil, animal fats and waste cooking oil
        
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        Advantages
        
        biodiesel is renewable and, like bioethanol, it is theoretically carbon neutral as the CO2 that is absorbed during photosynthesis is released when the biodiesel is burned. Biodiesel is also non-toxic and biodegradable, and it reduces the emission of atmospheric pollutants such as particulates, carbon dioxide and sulfur dioxide
        
        Disadvantages
        
        disadvantage of biodiesel is that, as with the production of bioethanol, land that could be used for food production is used to grow crops to produce biodiesel. In addition, a loss of biodiversity could result from the planting of the crops used to produce biofuels in general.
    - - Photovoltaic (PV) cells convert light energy from the Sun into electrical energy. PV cells are made of semiconductor materials such as silicon
        
        The bonds between the atoms of silicon are strong covalent bonds, which means the electrons are held in fixed positions, unable to move within the structure. This is the reason for the poor electrical conductivity of silicon at low temperatures.
        
        Effect of increasing temperature on electrical conductivity
        
        The electrons in silicon are not free to move within the giant covalent structure; instead they are held in fixed position between the silicon atoms. As the temperature increases, semiconductors ionise forming free moving electrons that allow a small current to flow.
        
        Semiconductor doping
        
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        When the temperature of a metal (tungsten) increases, the increased thermal energy causes the positive metal ions to vibrate more. The increased vibrations of the positive metal ions impede the movement of the delocalised electrons within the structure, causing the electrical conductivity to decrease.
        
        Pros and cons
        
        ilicon advantage: highly efficiency = greater than 20%
        
        disadvantages: more complex and expensive
      - How photovoltaic cells produce an electric current
        
        When the light energy contacts the surface of a photovoltaic cell, the electrons within the semiconductor material are excited from the valence band to the conduction band.
        
        At 0 K, the valence band is completely full of electrons and the conduction band completely empty, therefore silicon does not conduct electricity at this temperature. As the temperature increases, electrons are excited to the conduction band, leaving behind holes in the valence band. The electrons in the conduction band are free to move and allow a semiconductor such as silicon to carry an electric current.
        
        A photovoltaic cell is composed of two layers: n-type silicon and p-type silicon. Between the two layers is a region called the pn junction, which is a boundary between the two types of semiconductor materials.
        
        The n-type silicon has a greater number of mobile electrons than the p-type silicon. These mobile electrons naturally diffuse across the pn junction from the n-type to the p-type silicon, creating positive charges in the n-type material.
        
        Electron holes can also diffuse across the pn junction, creating negative charges in the p-type material. The separated positive and negative charges in the n-type and p-type materials results in the formation of an electric field across the pn junction.
        
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        Pro and cons
        
        DSSC advantage: low cost, simple to make, absorb longer wavelengths
        (lower energy) so they work better on cloudy conditions
        
        disadvantage: electrolyte can freeze if it is cold or expand if it is hot
        Less efficient= 11%
  - - - the process that warms the surface of the Earth
        
        The enhanced greenhouse effect is caused by greenhouse gases such as carbon dioxide being added to the atmosphere by human activities, and is thought to be contributing to the warming of the Earth's surface
        
        Greenhouse gases in the atmosphere are transparent and allow the passage of incoming short wavelength solar radiation to reach the surface of the Earth. The Earth’s surface reradiates some of this radiation back into the atmosphere as longer wavelength (infrared) radiation. Greenhouse gases in the atmosphere absorb this longer wavelength radiation and reradiate it back to the surface of the Earth, causing the temperature at the surface to increase.
        
        Green house gases
        
        Greenhouse gases are gases that absorb infrared radiation
        
        When discussing the potency of a greenhouse gas, both of these factors (the greenhouse factor and the relative abundance) must be taken into account.
        
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        Link Title
  - - - Once a battery is no longer able to produce an electric current, it must be disposed of, which can cause harm to the environment. Primary cells can only be used once and cannot be recharged as the chemical reactions that take place within the cell are not reversible
        
        Oxidation of zinc atoms at the anode: Zn(s) → Zn2+(aq) + 2e-
        
        Reduction of copper ions at the cathode: Cu2+(aq) + 2e- → Cu(s)
        
        Net ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
      - The salt bridge completes the circuit and prevents a build-up of charge by allowing ions to migrate between the two half-cells.
        
        A battery is one or more voltaic cells connected in series. The voltage produced by a battery primarily depends on the nature of the materials, specifically the electrodes and the electrolyte used
        
        Figure 1, the electrons flow from the half-cell with the more negative standard electrode potential to the half-cell with the more positive standard electrode potential. The voltage produced by a voltaic cell can be calculated using the standard electrode potentials of the half-cells
        
        In a voltaic cell, the current is carried by ions that diffuse through the electrolyte. Ions can only diffuse at a finite speed, which results in the internal resistance of a cell
        
        The higher the voltage the higher the inter resistance as well, so its an balancing act, like most things are
        
        EƟcell = EƟred − EƟox
        
        don't change the value of the metal that undergoes oxidation as the equation accounts for this, so just plug in the value that you see for the reduction reaction in data booklet table 24
    - - A secondary cell is a voltaic cell that can be recharged. The chemical reactions that occur within a secondary cell are reversible
        
        Lead Acid Battery
        
        They are typically used in motor vehicles due to their ability to produce the large electric current required to 'start' an engine. The cell potential of a lead-acid battery is 2.02 V. A typical battery has six cells in series, producing a total of 12 V.
        
        Discharged
        
        At the anode: Pb (s) + HSO4- (aq) → PbSO4 (s) + H+ (aq) + 2e-
        
        At the cathode: PbO2 (s) + 3H+ (aq) + HSO4- (aq) + 2e- → PbSO4 (s) + 2H2O (l)
        
        Overall equation: Pb (s) + PbO2 (s) + 2H2SO4 (aq) → 2PbSO4 (s) + 2H2O (l)
        
        Recharged
        
        At the anode: PbSO4 (s) + 2H2O (l) → PbO2 (s) + 3H+ (aq) + HSO4- (aq) + 2e-
        
        At the cathode: PbSO4 (s) + H+ (aq) + 2e- → Pb (s) + HSO4- (aq)
        
        Overall equation: 2PbSO4 (s) + 2H2O (l) → Pb (s) + PbO2 (s) + 2H2SO4 (aq)
        
        Advantages
        
        Lead-acid batteries are simple to manufacture, cheap and have low self-discharge rates
        
        Disadvanteges
        
        Lead-acid batteries have low energy densities and the use of lead can cause environmental concerns.
        
        Overcharging can lead to the production of hydrogen gas, which is an explosion risk.
        
        Lead is a toxic metal that can affect the central nervous system.
        
        Nickel-cadmium battery
        
        Discharged
        
        At the anode: Cd (s) + 2OH− (aq) → Cd(OH)2 (s) + 2e−
        
        At the cathode: 2NiO(OH) (s) + 2H2O (l) + 2e− → 2Ni(OH)2 (s) + 2OH− (l)
        
        Overall equation: Cd (s) + 2NiO(OH) (s) + 2H2O (l) → 2Ni(OH)2 (s) + Cd(OH)2 (s)
        
        Recharged
        
        At the anode: 2Ni(OH)2 (s) + 2OH− (l) → 2NiO(OH) (s) + 2H2O (l) + 2e−
        
        At the cathode: Cd(OH)2 (s) + 2e− → Cd (s) + 2OH− (aq)
        
        Overall equation: 2Ni(OH)2 (s) + Cd(OH)2 (s) → Cd (s) + 2NiO(OH) (s) + 2H2O (l)
        
        Advanteges
        
        NiCd batteries charge quickly and
        have a high number of
        charge/discharge cycles.
        
        Disadvanteges
        
        The use of cadmium and nickel can
        cause environmental concerns
        during disposal.
        
        Lithium-ion battery
        
        Lithium-ion (Li-ion) batteries are widely used in portable electronic devices such as smartphones and portable music players.
        
        Discharged
        
        At the anode: Li (graphite) → Li+ (electrolyte) + e-
        At the cathode: Li+ (electrolyte) + e- + CoO2 (s) → LiCoO2 (s)
        
        Recharged
        
        At the anode: LiCoO2 (s) → Li+ (electrolyte) + e- + CoO2 (s)
        At the cathode: Li+ (electrolyte) + e- → Li (graphite)
        
        Advanteges
        
        Li-ion batteries have high energy densities and low discharge rates.
        
        Disadvantages
        
        Li-ion batteries are expensive when compared to other types of batteries.
        Li-ion batteries have been known to explode, causing
        fires that are very difficult to extinguish.
    - - Fuel cells convert chemical energy into electrical energy. They differ from primary cells in that they need a constant source of fuel to produce an electric current.
        
        Proton exchange membrane (PEM)
        
        At the anode: 2H2 (g) → 4H+ (aq) + 4e-
        
        At the cathode: O2 (g) + 4H+ (aq) + 4e- → 2H2O (l)
        
        Overall: 2H2 (g) + O2 (g) → 2H2O (l)
        
        Electrons flow in the external circuit from the anode to the cathode, producing an electric current. At the cathode, oxygen gas is reduced, producing pure water. Certain types of fuel cell contain a proton exchange membrane (also called a polymer electrolyte membrane) that consists of a semipermeable layer which allows the passage of hydrogen ions, or protons, but prevents mixing of the reactants.
        
        Alkaline fuel cells
        
        used by NASA in the Apollo mission series due to their high efficiencies and the fact that they produce pure water, which could be used by the astronauts while in space.
        
        At the anode: 2H2 (g) + 4OH- (aq) → 4H2O (l) + 4e-
        
        At the cathode: 2H2O (l) + O2 (g) + 4e- → 4OH- (aq)
        
        Overall: 2H2 (g) + O2 (g) → 2H2O (l)
        
        Direct methanol fuel cells
        
        Direct methanol fuel cells do not have the same problems associated with the storage of high-pressure gases as PEM fuel cells. Methanol, being a liquid fuel, is easier to store and transport than compressed hydrogen gas.
        
        At the anode: CH3OH (l) + H2O (l) → CO2 (g) + 6H+ (aq) + 6e-
        
        At the cathode: 6H+ (aq) + 3/2O2 (g) + 6e- → 3H2O (l)
        
        Overall equation: CH3OH (l) + 3/2O2 (g) → CO2 (g) + 2H2O (l)
        
        Microbial fuel cells
        
        A microbial fuel cell converts chemical energy to electrical energy using bacteria that live inside the anode (called the bioanode) of the fuel cell. At the anode, the bacteria break down organic matter to produce hydrogen ions, electrons and carbon dioxide.
        
        At the anode: CH3COO- (aq) + 2H2O (l) → 2CO2 (g) + 7H+ (aq) + 8e-
        
        At the cathode: 2O2 (g) + 8H+ (aq) + 8e- → 4H2O (l)
        
        Overall equation: CH3COO- (aq) + H+ (aq) + 2O2 (g) → 2CO2 (g) + 2H2O (l)
        
        Unlike PEM fuel cells, they do not require catalysts made of expensive metals such as platinum (Pt).
        
        Carbohydrates such as glucose, C6H12O6, can also be used as a fuel for microbial fuel cells.
        
        C6H12O6 (aq) + 6H2O (l) → 6CO2 (g) + 24H+ (aq) + 24e–
        
        At the cathode, oxygen is reduced to form water
        
        24H+ (aq) + 6O2 (g) + 24e– → 12H2O (l)
        
        Overall equation
        
        C6H12O6 (aq) + 6O2 (g) → 6CO2 (g) + 6H2O (l)
      - Comparison to other fuels
        
        In addition to being less harmful to the environment, fuel cells are more efficient than petrol and diesel engines.
        
        Fuel cell vehicles are considerably more expensive than the gasoline-powered equivalents. There is also a lack of hydrogen infrastructure, especially when compared to conventional fuels such as petrol or diesel.
      - Thermodynamics efficiency of fuel cells
        
        Thermodynamic efficiency = ΔG / ΔH × 100 %
    - - Report feedback on content
        A concentration cell has the same electrodes in each half-cell but the concentration of the electrolyte in each half-cell is different.
        
        A concentration cell produces an electric current by decreasing the concentration of the solution in the concentrated half-cell and increasing the concentration of the solution in the dilute half-cell. The electrons flow from the dilute half-cell to the concentrated half-cell until an equilibrium is reached, at which point the cell will cease to produce an electric current. At the point of equilibrium, the electrolyte concentrations in each half-cell are equal.
        
        At the dilute half-cell (anode): Cu(s) → Cu2+(aq) + 2e-
        
        At the concentrated half-cell (cathode): Cu2+(aq) + 2e- → Cu(s)
        
        Net ionic equation: Cu(s) + Cu2+(aq) ⇌ Cu2+(aq) + Cu(s)
      - The Nernst equation
        
        E = EƟ – (0.0257 / n) ln Q
        
        Example
        
        Calculate the voltage produced by the concentration cell in Figure 1 at a temperature of 298 K.
        
        The EƟ of a concentration cell is 0 V, as under standard conditions of 1.00 mol dm-3 no voltage is produced.
        
        Q, the reaction quotient, is calculated by dividing the concentration of the half-cell at the anode by the concentration of the half-cell at the cathode.
        In the reaction quotient, pure solid reactants and products are not included in the expression.
        If the coefficients of the aqueous ions are the same, then they are left out of the expression, if they are different, then they are included.
        
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        cadmium half-cell and a lead half-cell with electrolyte concentrations of 0.0500 mol dm-3 and 0.500 mol dm-3 respectively. Calculate the voltage produced by the voltaic cell at 298 K.
        
        The net ionic equation for the cell is:
        Cd(s) + Pb2+(aq) → Cd2+(aq) + Pb(s)
        EƟcell = Ered – Eox
        EƟcell = –0.13 – (–0.40)
        
        Q = [Cd2+] / [Pb2+]
        Q = 0.0500 / 0.500
        Q = 0.100
        
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- - - - Enthalpy is the term used for the heat released or absorbed by a system at constant pressure
        
        Exothermic reaction
        
        Energy being realesed
        
        Endothermic reaction
        
        Energy being absorbed
        
        To find the change of enthalpy we can use a calorimeter and use this equation
        
        q = mc∆T
        
        But to compare it to other compounds we divide this value by moles
        
        Graphing results
        
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        Reactions with solutions
        
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    - - Kelvine scale = celcius +273
  - - - Hess’s law states that the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same.
        
        Hess’s law by drawing enthalpy cycles
        
        ΔH1 = ΔH2 + ΔH3
        
        Enthalpy change of formation (ΔH⦵f)
        
        Enthalpy of formation values are useful in that they indicate the stability of compounds in relation to their elements - the more negative the value, the greater the stability of the compound
        
        Hess's Law is that the enthalpy change of any reaction can be calculated provided that we know the enthalpy of formation values of the reactants and products
  - - - The energy required to break a chemical bond is called the bond enthalpy, H or bond dissociation energy, E
        
        It is also important to note that these values refer to molecules in the gaseous state
        
        When new bonds are formed, the energy released is equal to the energy that was absorbed to break the bond
        
        Average bond enthalpy
        
        the enthalpy change when one mole of bonds are broken in the gaseous state averaged for the same bond in similar compounds
        
        the data booklet are calculated from a range of different compounds, in which the same type of bond has different bond enthalpy values.
        
        Calculating enthalpy changes using bond enthalpy values
        
        ΔH = ΣE(bonds broken) − ΣE(bonds formed)
  - - - Eg: NaCl (s) → Na+ (g) + Cl− (g) ΔH⦵ lat = +771 kJ mol−1
        This is an endothermic process
        
        Magnitude of the lattice enthalpy
        
        The magnitude of the lattice enthalpy of an ionic compound can be thought of as a measure of the strength of the ionic bonds between the ions
        
        The magnitude of the lattice enthalpy depends on two factors:
        the charge of the ions (ionic charge)
        the size of the ions (ionic radii).
        
        Eg: ΔH⦵lat MgO (s) = +3791 kJ mol−1
        ΔH⦵lat NaCl (s) = +790 kJ mol−1
    - - It is not possible to determine the lattice enthalpy value of an ionic compound directly by experiment. Instead, we can use an energy cycle called a Born–Haber cycle.
        
        The different steps in the cycle are as follows:
        
        Enthalpy of formation, ΔH⦵f: the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
        
        Enthalpy of atomisation, ΔH⦵at: the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.
        
        Bond dissociation energy, E: the energy required to break one mole of bonds in the gaseous state.
        
        Ionisation energy, ΔH⦵IE: the energy required to remove one mole of electrons from one mole of gaseous atoms.
        
        Electron affinity, ΔH⦵EA: the energy released when one mole of electrons are added to one mole of gaseous atoms.
    - - The first step in the dissolving process is the break-up of the lattice structure, by which the ionic solid is converted to gaseous ions. The enthalpy change for this process is the lattice enthalpy, ΔHlat
        
        gaseous ions are hydrated by the surrounding water molecules (Figure 1a). During hydration, ion-dipole forces are formed between the gaseous ions and the partial charges of the water molecules.
        
        The change in enthalpy that occurs when an ion becomes hydrated is the enthalpy of hydration, ΔH⦵hyd.
        
        The sum of these two enthalpy changes is the enthalpy change of solution (ΔH⦵sol)
        
        ΔH⦵sol = ΔH⦵lat + ΔH⦵hyd
        
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        Factors affecting the enthalpy of hydration
        
        The magnitude of the enthalpy of hydration for a particular ion depends on two factors: the ionic radii and the charge on the ion
        
        the charge density of an ion increases with greater ionic charge and a decrease in ionic radius. A higher charge density results in a stronger ion-dipole force between the ion and the water molecule, and a greater, more negative value for the enthalpy of hydration.
  - - - Changes in entropy (ΔS)
        
        When considering the sign of the entropy change (positive or negative), it is important to consider the changes in the state of the reactants and products. For example, evaporation results in an increase in entropy (+ΔS) as gases have higher entropy than liquids
        
        Standard entropy values
        
        A perfect crystal at absolute zero (0 K) has a standard entropy of zero; therefore, every other substance has a positive standard entropy value. Standard entropy values can be used to calculate the standard entropy change (ΔS⦵) for a reaction, using the following equation:
        ΔS⦵ = ΣS⦵ (products) − ΣS⦵ (reactants)
      - A spontaneous process occurs without the addition of energy, other than that required to overcome the initial energy barrier (also known as the activation energy
        
        For a spontaneous process, the total entropy of the system and surroundings (ΔStotal) must increase. This can be represented in equation form as:
        ΔStotal = ΔSsystem + ΔSsurroundings ≥ 0
        
        Certain chemical reactions have a decrease in the entropy of the system and yet are still spontaneous. This is because the entropy of the surroundings increases to a much greater extent, which gives an overall increase in entropy for the process
        
        Free energy and work
        
        Gibbs free energy is the energy associated with a process (such as a chemical reaction) that can be used to do work, specifically non-expansion work
        
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- - - - At the same temperature, all gases have the same average kinetic energy, therefore, the distribution is dependent on the molar mass of the gas
  - - - Higher activation energy means slower reaction
  - - - Graphs
        
        The gradient of the line represents the rate of the reaction
        
        Average rate
        
        Change in concentration in a time period/ interval
        
        instantaneous rate
        
        determined graphically from a change in reactant or product concentration against time
        
        Experimental techniques for measuring rates of reactions
        
        Production of gas
        
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        Change in ion concentration by conductivity
        
        Measuring time for precipitate to form
        
        Change in concentration by titration
        
        Or find the derivative at that point in time
    - - 2A + B → A2B
        Rate=k[A]2[B]
        
        k is the rate constant, dependent on experimental conditions and the reaction
        
        Order of reaction
        
        shows the relationship of the concentration of reactants to the change in rate. (most common
        
        First order
        
        reactants are directly proportional to the rate
        
        Second order
        
        reactants are exponentially proportional to the rate
        
        zero order
        
        Reactants concentration has no effect on the rate
        
        If reactant A is 1st order and reactant B is 2nd order the overall order is 3rd (1+2=3)
        
        Molecularity
        
        unimolecular, bimolecular, termolecular (A, A+B, A+B+C)
      - The Arrhenius equation
        
        k=Ae^(-Ea/RT)
    - - Mechanisms are the detailed steps of a reaction, shows the process of the molecules rearrangement to form the products.
        
        The slow step determines the rate of reaction.
        
        This is only a theory about how a reaction takes place; it cannot be proved. However, a reaction mechanism can be disproved if it does not agree with the experimental data.
        
        For it to be considered plausible it must:
        
        the elementary steps must add together to give the overall balanced equation for the reaction, and
        the reaction mechanism must be consistent with the experimentally determined rate expression (in other words, the rate expression deduced from the rate-determining step must be consistent with the experimentally determined rate expression).
        
        Eg: The balanced equation of the reaction from before is
        NO2 (g) + CO (g) → NO (g) + CO2 (g)
        
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        Eg:Step 1 NO2 (g) + NO2 (g) → NO3 (g) + NO (g) slow
        Step 2 NO3 (g) + CO (g) → NO2 (g) + CO2 (g) fast
        so rate = k [NO2]^2
      - Elementary reactions (single step reactions)
        
        Non elemne... (multiple
    - - Concentration
      - Temperature
      - Surface area
      - Catalyst
      - Agitations (stirring)
- - - - Oxidation half reaction: reaction where electrons are lost
      - Reduction half reaction: reaction where electrons are gained
      - Oxidizing agent: reactant of the reduction reaction
        Oxidizing agents readily accept electrons
        Common example are halogens, oxygen and hydrogen peroxide
      - Reducing agent: reactant of the oxidation reaction
        Reducing agents readily donate electrons
        Common examples are metals, C, H2, S, P, hydracids, low oxidation states.
      - Oxidation state: number and sign that indicates the gain or loss of electron control of an atom during reaction
        
        Rules for determining oxidation states
        
        Free elements such as O2, Cl2, N2, are assigned an oxidation state of zero.
        
        The sum of the oxidation states of all the atoms in a compound must be equal to the net charge on the compound.
        
        The alkali metals (Li, Na, K, Rb and Cs) in compounds are always assigned an oxidation state of +1.
        
        Fluorine in compounds is always assigned an oxidation state of −1.
        
        The alkaline earth metals (Be, Mg, Ca, Sr, Ba and Ra) and Zn in compounds are always assigned an oxidation state of +2.
        
        Hydrogen in compounds is assigned an oxidation state of +1, except in certain metal hydrides (e.g. NaH), where it is −1.
        
        Oxygen in compounds is assigned an oxidation state of −2, except in peroxides (e.g. H2O2), where it is −1, or when it is combined with fluorine, in which it is +2.
        
        Chlorine in compounds is assigned an oxidation state of −1 unless it is combined with oxygen or fluorine.
        
        The charge on a metal ion is the same as its oxidation state, e.g. Zn2+ has an oxidation state of +2.
        
        The sum of the oxidation states in a polyatomic ion (such as CO32−) must add up to the charge on the ion.
        
        Disproportionation reactions
        
        when the same species is oxidised and reduced simultaneously during a reaction to form two different products
        
        eg: 2H2O2 (aq) → 2H2O (l) + O2 (g)
  - - - Used to detect the biochemical oxygen demand (BOD) . Which is the amount of dissolved oxygen required to decompose the organic matter in a water sample over a set period of time.
        Polluted water will have a higher BOD
        
        Chemicals added:
        Manganese (II) sulfate (fixes oxygen- requires basic conditions)
        Acid (required for Iodine reaction)
        Potassium iodide (generates iodine)
        Sodium thiosulfate (titrate excess iodine)
        Starch (indicactor)
        
        Color Progression
        Pale brown precipitatie (MnO2)
        Light yellow (I-)
        Dark blue (NaS2O3 with KI)
        Clear
        
        To measure the BOD of a river, two 300.0 cm3 samples of river water were collected. The Winkler method was used to determine the initial concentration of dissolved oxygen in one sample at day zero. 30.00 cm3 of 0.0100 mol dm−3 of Na2S2O3 (aq) was required to reach the end-point of the titration. The second sample was then stored in a dark place to be tested after 5 days.
        (a) Calculate the initial concentration of dissolved oxygen in the sample of water in mg dm−3.
        
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        To Determine the biochemical oxygen demand, BOD, you have to find difference between the initial and the final concentration of O2 in the body of water
        Eg: (b) After 5 days, the second sample of river water contained 1.88 × 10−5 mol of O2. Determine the biochemical oxygen demand, BOD, of the river water
        
        Final amount of dissolved O2 = 1.88 × 10−5 mol
        Convert from amount (in mol) to mass (in mg):
        (1.88 × 10−5 mol × 32.00 g mol −1) × 1000 = 0.602 mg
        
        Calculate the concentration of dissolved O2 after 5 days:
        0.602 mg / (300.0 / 1000) = 2.01 mg dm−3 = 2.01 ppm.
        
        Finally, subtract the concentration of dissolved O2 after 5 days from the initial concentration of dissolved O2:
        8.00 mg dm−3 – 2.01 mg dm−3 = 5.99 mg dm−3 = 5.99 ppm
    - - Voltaic cell
        
        a redox reaction to produce an electric current
        It consists of two half-cells: a zinc electrode in a solution of zinc sulfate, and a copper electrode in a solution of copper(II) sulfate.
        
        The voltage produced by a voltaic cell depends on the difference in reactivity between the two metals in the half-cells. The circuit is completed by a salt bridge that allows ions to flow in order to complete the circuit and prevent build-up of electric charge. It;S A SPONTANEOUS PROCCESS
        
        When combined in a voltaic cell the more reactive metal 'drives' the direction of the equilibrium. The more reactive metal is oxidised, and the less reactive metal is reduced
        
        The difference in the tendency to either undergo oxidation or reduction creates a potential difference between the two half-cells, also known as the electromotive force (EMF) or cell potential (E⦵cell).
        
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        Cell diagram convention
        
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        A simple salt bridge can be made from a strip of filter paper soaked in saturated potassium nitrate (KNO3), although more complex salt bridges can be made using a glass tube filled with agar gel. Potassium nitrate is chosen because its constituent ions do not react with the other ions present in the two half-cells, or with the electrodes
        
        Standard electrode potential
        
        Eocell is the standard electrode potential that is the difference of the standard electrode potential of the two half cells.
        
        standard hydrogen electrode) which has been assigned 0.00V
        This can act as the cathode or anode depending on if the other half cell is higher or lower on the electrochemical series.
        Half reaction: 2H+ (aq) + 2e– → H2 (g) or H2 (g) → 2H+ (aq) + 2e–
        
        hydrogen gas, H2 (g), at a pressure of 100 kPa and temperature of 298 K
        acidic solution with [H+] of 1.0 mol dm–3
        inert platinum electrode.
        
        Metals in the half-cells at the top of the electrochemical series, with large negative electrode potential values, are stronger reducing agents – they have the greatest tendency to lose electrons (undergo oxidation) and form positive ions in aqueous solution.
      - Electrolytic cell
        
        the process by which a compound is broken down into its constituent elements using electricity
        
        Characteristics
        
        Reduction occurs at the negative cathode
        
        Involves an endothermic non-spontaneous redox reaction
        
        Cell converts electrical energy into chemical energy
        
        The cathode is negative and the anode is positive during electrolysis
        
        The electrolyte is a molten liquid (or an aqueous solution)
        
        Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte
        
        electrolysis with aqueous solutions
        
        The electrolysis of aqueous solutions of ionic compounds is more complicated than that of molten ionic compounds because the water itself can be electrolysed. Water can either be oxidised at the anode or reduced at the cathode according to the following equations:
        
        Oxidation: 2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
        Reduction: 2H2O (l) + 2e– → H2 (g) + 2OH– (aq)
        
        The process of selective discharge occurs when one species is preferentially discharged over another at either the anode or cathode of an electrolytic cell.
        
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        At the anode:
        At low solution concentrations (a dilute solution), H2O (l) is oxidised to form oxygen gas
        At high solution concentrations (a concentrated or saturated solution), chloride ions are oxidised to form chlorine gas
        
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        Electrolysis of water
        
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        Electrolysis of aqueous copper(II) sulfate
        
        the cathode reveals a pinky-brown coating of copper caused by the reduction of copper(II) ions to form solid copper according to the following equation:
        Cu2+ (aq) + 2e– → Cu (s)
        
        This reaction is used in the refinement of impure copper. The anode is made from impure copper and pure copper is deposited on the cathode. Effectively, there is a net transfer of copper from the impure sample to the cathode via the copper(II) sulfate solution
        
        At the anode, water is oxidised to form oxygen gas:
        2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
        If the graphite electrodes are replaced by copper electrodes, the reaction taking place at the cathode remains the same – copper(II) ions are reduced – but at the anode the copper electrode itself undergoes oxidation:
        Cu (s) → Cu2+ (aq) + 2e−
        
        Electroplating
        
        At the cathode, silver ions undergo reduction to form silver atoms which coat the spoon:
        Ag+ (aq) + e− → Ag (s)
        
        The electrolyte solution is a salt containing the ions of the metal to be plated onto the object.
        
        At the anode, the silver atoms that make up the piece of silver undergo oxidation:
        Ag (s) → Ag+ (aq) + e−
        
        Calculation
        
        ΔGo = – nFEo
        
        G= -, reaction is spontaneous
        Eo= standard cell potential
        Reminder if ΔG is at zero the system is at equilibrium
        
        I=Q/t
  - - - This reaction list is useful for displacement reaction
        
        the reactivity of the element is measured in ( -) Eored or in voltage, more ( -) Eored results in higher reactivity
        
        Rust prevention
        
        The process of rusting, which we referred to at the very start of this subtopic, requires the presence of oxygen and liquid water. Rusting is a redox reaction in which iron atoms are oxidised to form iron(II) ions:
        Fe (s) → Fe2+ (aq) + 2e−
        
        Molecular oxygen is reduced to form hydroxide ions:
        O2 (g) + 2H2O (l) + 4e− → 4OH− (aq)
        
        ionic equation:
        2Fe (s) + O2 (g) + 2H2O (l) → 2Fe2+ (aq) + 4OH− (aq)
        
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- - - - H2O (l) ⇌ H2O (g)
  - - - aA + bB ⇌ cC + dD
        where a, b, c and d represent the stoichiometric coefficients in the balanced equation, is written as:
        
        Solids and liquids are not included, because it doesn't make sense to measure solid in concentration
        Dissociation expressions and constants are a type of Equilibrium expressions
        
        KC TO KP
        
        Kc is concentration expression equilibrium: unit is M
        Kp is pressure expression equilibrium: unit is atm
        
        To convert Kc to Kp use the following formula:
        Kp=Kc(RT)^(△n)
        
        ICE CHARTS
        
        The method involves making what is referred to as an 'ICE box'. ICE stands for:
        the initial concentration, I
        the change in concentration, C
        the equilibrium concentration, E.
        
        5% rule
        
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        Gibbs free energy and equilibrium
        
        ΔG⦵ = −RT ln K
        
        Gibbs free energy and entropy at equilibrium
        
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        Not at equilibrium
        
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        KINETICS TO EQUILIBRIUM
        
        MANIPULATING K VALUES
        
        Mathematically whatever happens with the reaction “one up” for the K values. Meaning:
        -if you add the equations, multiply the K values
        -if you multiply the equations, raise K to that exponent
        Kc=K’c*K”c
    - - Q is the same expression as K, however K is AT equilibrium, Q is when the values are not at equilibrium or it is unknown
        
        Q>K: Q needs to decrease so reaction moves to the left
        Q=K: the reaction is at equilibrium
        -Q<K: Q needs to increase so reaction moves reverse to the right
  - - - Pressure/Volume (gases only)
        Increase pressure reaction moves to the side with less moles of gas
        Decrease the pressure reaction moves to the side with more moles of gas
        K value doesn't change
      - Concentrations
        Increase reactants reaction moves to the right
        Increase products reaction moves to the left
        K value doesn't change
      - Temperature
        If exothermic: Increase temperature reaction moves to the left (because heat is a product)
        If endothermic: Increase temperature reaction moves to the right (because heat is a reactant)
        K value does change
        
        For an exothermic reaction, increasing the temperature shifts the equilibrium position to the left and the value of Kc decreases. For an endothermic reaction, increasing the temperature shifts the equilibrium position to the right and the value of Kc increases.
        
        This is when the forward reaction is endothermic
      - Catalyst
        
        Interestingly, the presence of a catalyst has no effect on the equilibrium position or the value of the equilibrium constant Kc, because a catalyst increases the rate of both the forward and reverse reactions by the same amount
    - - The Haber process
        
        Nitrogen is highly important for plant growth. However, most plants cannot ‘fix’ nitrogen directly from the air.
        
        The chemical process that produces ammonia gas involves the following reaction:
        N2 (g) + 3H2 (g) ⇌ 2NH3 (g) ΔH = –92 kJ mol^–1
        
        However, industrially a temperature of 450°C is actually used. This is an optimum temperature (Figure 5) – a compromise between producing sufficient ammonia and producing it at a reasonable rate of reaction. If too low a temperature is used, the reaction is too slow
      - The Contact process
        
        Sulfuric acid (H2SO4) is the single most produced chemical worldwide and, although there were several alternative historic methods of production, the Contact process is now almost solely and universally used worldwide
        
        2SO2 (g) + O2 (g) ⇌ 2SO3 (g) ΔH = –197 kJ mol–1
        
        Applying Le Châtelier’s principle suggests the highest yield of sulfur trioxide would be obtained using low temperatures and high pressures.
        
        very low temperatures cannot be used, as the rate of reaching equilibrium would be too slow, and therefore uneconomic. A further factor to be considered here is that the vanadium(V) oxide catalyst becomes effective only at temperatures above 400°C.
        
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- - - - Relative atomic mass
        
        This is the mass of an element relative to a 12th of the mass of 12C.
    - - At standard temperature and pressure
        
        (STP, T = 273K, P = 1.01x105 Pa):
        
        Molar Volume of Gas = 2.27x10-2 m3 mol-1
        = 22.7 dm3 mol-1
      - At room temperature and pressure
        
        (RTP, T = 298K, P = 1.01x105 Pa):
        
        Molar Volume of Gas = 2.45x10-2 m3 mol-1
        = 24.5 dm3 mol-1
      - Avogadro’s Law
        
        equal volumes of all gases, at the same temperature and pressure, have the same number of molecules
      - representing gases
        
        Ideal gases
        
        Perfectly elastic
        
        No intermolecular forces
        
        They take up almost no volume
        
        In reality, the opposite is true, but this assumptions makes the maths much easier and all these things are very minute
        
        Identical perfect spheres.
        
        Equation: PV = nRT
        
        Dalton’s Law of Partial Pressures
        
        Ptotal= P1 + P2 + P3 ...
    - - Dilution
        
        When a solution is diluted, more solvent is added to it, the
        
        C₁V₁=C₂V₂ or M₁V₁=M₂V₂
      - Titration
        
        (C₁V₁)/n₁=(C₂V₂)/n2
        
        Back Titration
        
        Adding a known but excess (too much) amount of one standard reagent to a known mass of the substance being determined (the analyte).
        
        When direct titration endpoint is hard to find (eg. weak acid and weak base)
        When the reaction occurs super slowly
    - - Molecular
        
        The number of atoms of each element in a molecule
      - Empirical
        
        The ratio of the atoms of each element in a compound in its lowest terms
        
        % Composition by Mass
        
        percentage of an atom in a molecule by mass
        
        Hydrates
        
        Compound that contains water with a definite mass in the form of H2
        
        Hydrates are often in the form of a crystal that can be heated, and the water can be “burned off” by turning it into steam
        
        The substance that is left over after a hydrate is has lost its water is called anhydrate
    - - The Limiting Reagent
      - Percentage yield
        
        Theoretical yield
        ---------------------- *100
        Actual yield
  - - - Eg: Colour change:
        
        temperature change
        
        -gas or solid form
      - Synthesis
        
        A + B → AB
      - Single replacement
        
        A + BC → AC + B
      - Double displacement
        
        AB + CD → AD + CB
        
        Net Ionic Equations
        
        Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq)+ NO₃⁻(aq)
        
        Spectator ions
      - Decomposition
        
        AB → A + B
      - Combustion
        
        CxHy + O2 → CO2 + H2O
    - - Deposition(Gas to solid) and sublimation
      - Compounds
      - Mixtures
        
        Homogeneous
        
        Heterogenous
        
        you can see the different components of the mixture
        Eg: oil and vinegar
        
        Collied
        
        These are mixtures where the module isn't fully dissolved but is also not divided like it is in heterogenous mixtures. The spread of these molecules is spread out evenly like in homogenous. Eg: Milk
  - - - Mass Spectrometry
        
        Relative atomic mass= % abundance mass + % abundancemass + % abundance * mass…
    - - Energy levels
        
        number of electrons in an energy level =2n2
        
        As the energy level increases the difference between the levels decreases
        
        spectroscope.
        
        Absorbtion spectrum
        
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        Emission spectrum
        
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        electrons can move up or down
        between these energy levels
        when absorbing or realising photons
        
        This graph is for Hydrogen
        
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        E = hf
        
        h = 6.63×10^(-34) Js
      - Orbitals/ sub levels
        
        Hybridisation
        
        In order for bonding to occur a given orbital must have a single electron. Each bond requires its own orbital. So electrons may have to move to a higher state. But then they are unequal energies which does not work. So they hybridize their energies.
        
        The three main types of hybridisation are sp, sp2 and sp3, with carbon undergoing sp3 hybridisation in the formation of methane, for example
        
        sp3 hybridisation
        
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        sp2 hybridisation
        
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        sp hybridisation
        
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        geometry for other compounds
        
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        Pauli exclusion principle- no two electrons can occupy the same space with the same spin
        Aufbau rule- fill the lowest energy orbital first (sometimes the lowest energy orbital is not from the lowest energy level i.e. 4s fills before 3d)
        
        Hund’s rule- fill one electron into each equal orbitals before placing a second one there
        
        Noble Gas Notation
        
        Orbital diagrams
        
        Exceptions
        
        Ex. Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5
        Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10
        
        Ion configurations
        
        For ions: continue to add in the same pattern as before, however remove from the outermost sub-level. (4s is more outer than 3d)
      - Chemical bonding and Structure
        
        Ionic Bonds
        
        Compounds with a metal and a nonmetal, cations bonded to an anion due to electrostatic attraction
        
        The atoms are trying to fill the outer shell to create a full octet. This is in accordance with the “octet rule” that states atoms are more stable if their electron configuration matches that of a nobel gas.
        
        Polyatomic ions
        
        Structure
        
        Lattice structure
        
        Properties
        
        Soluble in water
        
        solid in room temp
        
        highmelting and boiling points
        
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        electronegativity difference is approximately 1.8 or greater.
        
        Covelent
        
        an electronegativity of less than 1.8.
        
        electrostatic attraction between positively charged nuclei and shared pairs of bonding electrons.
        
        Polarity
        
        pure covelent
        
        Electrons are perfectly shared equaly
        
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        Non polar (weak polar)
        
        electrons are shared equally, quantified as 0.1-0.4 electronegativity difference. (C-H bonds)
        
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        Polar covalent
        
        Electrons shared unequally
        
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        Ionic bonds
        
        electrons are so unequally shared it is considered donated and accepted, quantified as 1.8 or greater (metal and nonmetal)
        
        covalent networks
        
        Carbon Allatropoes
        
        Fullerenes
        
        Sigma and pi bonds
        
        Sigma - direct head-on axial overlap of atomic orbitals (initial bond)
        
        Pi- sideways overlap of atomic orbitals (double and triple bonds)
        
        Metalic Bonding
        
        Lewis structures
        
        Resonance structures
        
        Multiple possible arrangements for electrons in a compound
        
        VSPER Structures
        
        Bonding Domains
        
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        Formal charge
        
        Formal charge is used to determine which lewis structure is preferred.
        Formal charges will always sum to the charge of the compound/ion
        
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        Coordinated covalent bonding
        
        A covalent bond that has both bonding
        electrons coming from one atom.
        
        Bond strength, length and order
        
        Triple bond is stronger than double bond is stronger than single bond
        Stronger bond will be a shorter bond
        
        Order
        
        Bond order is in accordance with single/double/triple bond
        
        Single bond- 1
        Double bond- 2
        Triple bond- 3
        
        Resonance structure will have fraction bond order:
        Example O3 bond order of 1.5
        
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    - - the occurrence of similar properties of elements after a definite interval
      - Groups, families
        
        Alkali metals, alkaline earth metals,
        transition metals, metalloids, nonmetals,
        halogens, noble gases, lanthanides, actinides
      - Trend with the periodic table
        
        atomic radius
        
        The distance between the center of the nucleus and the outer electrons
        
        Across periods: decreases (because effective nuclear charge is increasing)
        Across groups: increase (because it adds a new shell/energy level)
        
        For Ionic radii:
        
        Add an electron: increases (because the e>p, and increase electron repulsion)
        
        Remove an electron: decreases (because p>e)
        
        ionization energy
        
        Energy required to remove an electron from an atom or ion
        
        Across periods: increasing
        Down groups: decreases
        
        electron configuration of ionised transition metal
        
        Co: [Ar] 4s2 3d7
        
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        when you break a full subshell there is a minor peak
        
        electron affinity
        
        The energy change that occurs when an electron is added.
        
        Across periods: increases (becomes more negative)
        Down groups: decreases (becomes less negative)
        
        Effective nuclear charge
        
        The force felt by the outside electrons from the nucleus of an atom
        
        Across periods: it increases
        Across groups: it stays the same
        
        Metallic Characteristics
        
        he higher the nuclear charge and smaller the atomic radius of an atom, the lower its metallic character
        
        Melting points
        
        Trend with in groups
        
        Alkali metals
        
        going down decrease the melting point, and reactivity increases.
        
        Halogens
        
        going down the molar mass increases which increases the intermolecular forces, increases the melting point and boiling point, reactivity decreases
        
        Transition Metals
        
        d-block metals are dependant on the electron configuration of the d-block
        
        The d-block elements are not necessarily transition elements
        
        Transition elements are elements with an incomplete d- orbital in one or more of its atoms (ex. Cu+2 is a transition element but Zn or Zn+2 is not)
        
        Radii
        
        There is little change in radius of d-block elements because the electrons are being added and removed from inner sub-level (3d) so the effective nuclear charge is remaining fairly unchanged
        
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        oxidation states
        
        Note that all transition elements (with the exception of scandium) can have a +2 oxidation state because they lose their two 4s electrons to form 2+ ions. It should be noted that chromium and copper lose one electron from the 4s sub-level and the second electron from the 3d sub-level when they form 2+ ions because they only have one electron in the 4s sub-level.
        
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        Ionization energy trends on ions of d-block elements
        
        There is little change in radius of d-block elements which results in little change in ionization energy because a similar size (effective nuclear charge) results in a similar attraction of nucleus on the outer electrons
        
        complex ions
        
        Complex ion + counter ion = balanced compound.
        The counter ion is called coordinate compound, balances the charge of the complex ion Ex. [Fe(H2O)6]Cl3
        
        Vocab
        
        Ligand- the compound or ion attached or associated with the central atom/ion (require lone pair electron to act as the lewis base)
        
        Central metal ion: metal ion surrounded by ligands
        
        Complex ion: When ligands attach to a central metal ion
        
        Coordinate covalent bond: when the bonding pair of electrons is donated by just one of the bonding atoms.
        Coordination number: number of coordinate covalent bonds formed with ligands
        
        Monodentate, didentate, polydentate….
        
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        Magnetism
        
        -diamagnetic- only paired electrons: exhibit weak repulsion an external magnetic field, we however think of it being non-magnetic
        
        -ferromagnetic- contain unpaired electrons: that can align with an external magnetic field and can keep this alignment when the external magnetic field is removed. Only certain metals (iron, cobalt, nickel and few rare earth elements) can be ferromagnetic
        
        -paramagnetic- contain half-filled atomic orbitals: exhibit attraction to an external field, however they do not retain their magnetic properties outside of the external field. More unpaired electrons result in stronger attraction
        
        Collision theory and Catalysts
        
        catalyst increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy (Ea)
        
        heterogeneous catalysts
        
        catalyst is in a different state to that of the reactants
        
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        homogeneous catalyst
        
        the same physical state as the reactants. Transition metals are also effective as homogeneous catalysts in redox reactions due to their ability to have variable oxidation states
        
        they appear in the elementary steps of a reaction mechanism and in the rate-determining step but not in the overall balanced equation for the reaction
        
        Collision theory
        
        Collision theory states that in order for a reaction to occur the reactants must have sufficient energy(activation energy) and correct orientation
        
        Colored Complexes
        
        Partially filled d-orbitals when bonded with ligands will result in non-degenerate d-orbitals. Electrons that then get excited between the non-degenerate d-orbitals result in an emission of a visible wavelength.
        
        The color visualized is the opposite color of what is absorbed by the complex.
        
        The change in ΔE will affect the color (higher split = lower wavelength)
        
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        degenerate, meaning that they are of equal energy. When ligands bond to the central metal ion, repulsion between the electrons in the ligands and those in the d orbitals of the transition metal ion causes the five d orbitals to split into two sets of different energy (non-degenerate orbitals)
        
        Beer's Law
        
        Each complex has a unique 𝝀max which is found by scanning the all the visible wavelengths with a sample.
        
        Once the 𝝀max is found multiple solutions of the complex is made with various concentrations. The absorbance of these solutions is plotted against the concentration.
        
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        Electron negativity
        
        The force from the nucleus to attract electrons on a neighboring atom
        
        Across periods: increases (Smaller radius results in higher electronegativity)
        Down groups: decreases
      - https://www.youtube.com/watch?v=0VM1ZdLCFxc&t=1s