Unit 2 Chem Revision.docx

Unit 2 – What makes water such a unique chemical? Physical properties: - Change of state - No new substances are formed Chemical properties: - Disappearance/appearance of a substance - Gas/odour is released - Colour or temperature change - A solid is precipitated - Light is emitted IMPORTANT: Always balance chemical equations Properties of Water Latent Heat – Energy needed to change the physical state of a substance at 0°C Latent heat of fusion – changing water from solid to liquid Latent heat of vaporisation – changing water from liquid to gas at 100°C Specific Heat capacity – energy (KJ) needed to raise temperature by 1°C Water does not heat as quickly as other liquids

Writing Ionic Equations 1) Write balanced chemical equation 2) Dissociate all soluble compounds into their free ions 3) Replace acids and bases that react with water with their product ions 4) Cancel free ions that remain constant in the equation 5) Write net ionic equation Solubility Saturated – when solute stops dissolving at max amount Unsaturated – when solute dissolves Supersaturated – more solute maximum amount of solute Increase in temperature = increase in kinetic energy = increase in solubility Increase in temperature = increase in kinetic energy = liquid becomes gaseous and leaves solution = gases in water decrease High air pressure above liquid creates difficulty in boiling – boiling point is increased Concentration Measure of the amount of a particular solute in a given amount

Q = mc(∆T) Q = Energy (KJ) m = mass of substance (g/kg as given) c = specific heat capacity ∆T = change in temperature Density – the amount of mass within a volume of space 𝑚 d= 𝑉 Solid water (ice) is a covalent lattice – less dense than liquid water Water’s strong hydrogen bonds: - Relatively high melting and boiling points - High latent heat values - Relatively high specific heat capacity of liquid water Electrical conductivity: - Self-ionises and becomes hydroxide - Covalently bonded with no charged particles Dissociation – Water (catalyst) separates ions from ionic substances Ionisation – Polar molecules react with water to produce ions Hydrolysis – Water breaks a covalent H-O bond

%m/m

grams in every 100g

%m/v %v/v Ppm Ppb

grams in every 100ml mL in every 100ml mg per litre micrograms per litre 𝑚𝑜𝑙𝑒𝑠

Concentration = 𝑣𝑜𝑙𝑢𝑚𝑒(𝐿)

𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔)

𝐶(𝑔𝐿−1 ) = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝐿)

𝐶(𝑔𝐿−1 ) = 𝐶(𝑚𝑜𝑙𝐿−1 ) × 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠

Dilutions C1 V1 = C2 V2 Types of reactions Precipitation - Ions in solution combine to form a solid compound in water - Water is unable to dissociate ionic lattices of non-polar precipitates

Unit 2 – What makes water such a unique chemical? Soluble Ions Group 1 metals Sodium (Na+) Nitrate (NO3-) Ammonium (NH4+) Potassium (K+) Ethanoate (CH3COO-) Carbonate (CO32-) Hydroxide (OH-) Oxide (O2-) Phosphate (PO43-) Sulphite (SO32-) Chloride (Cl-) Bromide (Br-) Iodide (I-) Insoluble Ions: - Silver - Lead (II) - Mercury (I) - Calcium - Barium

Soluble salts All All All All All All Group 1 Group 1 Group 1 Group 1 Group 1 Most Most Most Ag+ Pb2+ Hg+ Ca2+ Ba2+

Combustion Hydrocarbon + (plentiful) oxygen -> carbon dioxide (CO2) + water Hydrocarbon + (limited) oxygen -> carbon monoxide (CO) + water Acid-base (Neutralisation) Properties of Acid - Donates a proton (loses H+) (bonds with H are highly polar) - Produces H+ in H3O+ metal - Ionises in water to produce electrolyte - Tastes sour - pH less than 7 - Turns blue litmus to red Properties of Bases - Accepts a proton (gains H+) - Produces OH- ions in water - Dissolves in water to produce an electrolyte - Taste bitter and chalky - Feels soapy and slippery - pH greater than 7

acid + base -> salt + water acid + metal -> salt + hydrogen excludes Cu, Hg and Ag acid + metal carbonate -> salt + water + carbon dioxide acid + metal oxide -> salt + water acid + metal hydroxide -> salt + water acid + metal sulphide -> salt + hydrogen sulphide Conjugate pairs - Acid and bases which differ by the presence or absence of a proton (H+) Acids to conjugate base: REMOVE H+ and ADD electron Bases to conjugate acid: ADD H+ and REMOVE electron Water as a base: HSO4-(aq) + H2O(l) -> SO42-(aq) + H3O+(aq) Water as an acid: HSO4-(aq) + H2O(l) -> H2SO4(aq) + OH-(aq) Polyprotic – donates/accepts more than one proton Diprotic – donates/accepts two protons Amphiprotic – can donate or accept a proton Homogenous – mixture that is evenly spread and dissolved (miscible) Heterogeneous – mixture that has uneven, undissolved clumps (immiscible) Strength vs Concentration Strength – % of protons available to donate/ionisation in water Concentration – ratio of water to solute in solution Strong acids - hydronium H3O+ - hydrochloric acid HCl (monoprotic) Weak acids - sulfuric acid H2SO4 (diprotic) - ethanoic acid CH3COOH - nitric acid HNO3 Strong bases Weak bases - hydroxides OH- Ammonia NH3 2 - oxides O - NaCO3 pH measure of strength of acid or base Neutral – 7 pH = -log[H3O+] [H3O+] = 10-pH Acidic - <7 10−14

pOH = -log( 𝑂𝐻− )

[OH-] = 10-pOH

Basic - >7

pH + pOH = 14.00 IMPORTANT – Remember negative in log and multiply by no. of H

Unit 2 – What makes water such a unique chemical? To find pH/pOH 1) Find moles 𝑛 2) Find molarity (M= 𝑉) 3) Log molarity to find pH/pOH Water’s most neutral stage is 7pH at 25°C Redox Reactions The transfer of electrons in a reaction. Two reactions that occur simultaneously. Oxidation Reduction - Reductant - Oxidant - Gain oxygen - Loss oxygen - Loss hydrogen - Gain hydrogen - Loss electrons - Gain electrons To balance redox half-equations K – potassium ions are balanced O – oxygen ions are balanced (add H2O) H – hydrogen ions are balanced (add H+) E – Electrons charges are balanced S – Write the states of each substance IMPORTANT – Beware of +/- signs and molecule coefficient when adding charges Overall reaction: Multiply half-equations to cancel out added electrons then restate the reaction. Oxidation Numbers - A count of the electrons transferred in the formation or breaking of chemical bonds Rules Hydrogen +1 -1 when with electropositive elements Oxygen -2 +1 or +2 with Fluorine Fluorine negative Aluminium +3 Copper +2 Pure elements 0 REMEMBER - FONCl G1 elements +1 G2 elements +2 G17 elements -1 Compounds sum of oxidation numbers = 0 Ions ion charge Multiply with number of atoms in compound

Analysing water for salts Gravimetric Analysis The isolation of an ion in a solution by a precipitation reaction 1) Dissolve mixture and filter insoluble material 2) Form and filter the precipitate 3) Wash the precipitate to remove soluble impurities 4) Dry the precipitate until all water ions are removed 5) Weigh the precipitate until a constant mass is achieved The precipitate should: - Have low solubility - Be stable when heated - Should have a known formula Errors: - Filtered/washed/dried ineffectively – result mass is higher - Not all chemicals have precipitated – result mass in lower - Other chemicals have precipitated – result mass is higher - Precipitate is slightly too soluble – result mass is lower To find %m/m of analyte in precipitate 1) Convert given mass of precipitate into moles 2) Find mass of molecule to be analysed using molar mass 3) Find % of this of whole substance To find concentration of analyte in original substance 1) Convert given mass of precipitate into moles 2) Use mole to mole ratio from balanced equation to find moles of analyte 3) Find concentration of analyte using volume of original substance 𝐷𝑖𝑓𝑓𝑒𝑟𝑒𝑛𝑐𝑒 𝑖𝑛 𝑣𝑎𝑙𝑢𝑒𝑠 (𝑚𝑎𝑠𝑠) %difference from manufacturer’s value = 𝑚𝑎𝑛𝑢𝑓𝑎𝑐𝑡𝑢𝑟𝑒𝑟′ 𝑠 𝑣𝑎𝑙𝑢𝑒 × 100 Higher: - Not all water was removed from precipitate - Other ions may have precipitated - Not all impurities were removed from precipitate Lower: - Not all of substance was dissolved - Some ions were lost in filtration or transfers - Not all ions have precipitated Chromatography To separate and analyse small amounts of mixtures. Different compounds travel at different but constant speeds. Most effective in identical conditions. Adsorption – sticks and forms a bond with surface Desorption - breaks bonds

Unit 2 – What makes water such a unique chemical? Paper: - Soluble substances move upwards with solvent/mobile phase - Substances stick to stationary phase – least adsorbed moves more distance, most adsorbed moves less distance - Compounds with similar polarity to stationary phase are most strongly adsorbed - Compare unknown substances with known substances to determine chemicals within them Column: - Soluble substances move downwards with solvent and appear as bands - Substances move and adsorb similarly to paper - Fast and gives better separation

To improve separation of bands: - Increase length of column - Pack column more densely - Increase time allowed for separation - Increase temperature Retardation factor =

𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑡𝑟𝑎𝑣𝑒𝑙𝑙𝑒𝑑 𝑏𝑦 𝑐𝑜𝑚𝑝𝑜𝑛𝑒𝑛𝑡 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑡𝑟𝑎𝑣𝑒𝑙𝑙𝑒𝑑 𝑏𝑦 𝑠𝑜𝑙𝑣𝑒𝑛𝑡

Increases when compound is least strongly adsorbed (distance is further) Retention time – time taken to travel through the column Increases when compound is most strongly adsorbed (distance is shorter) To increase retention time: - Reduce length of column - Increase amount of sample used - Pack column less tightly - Reducing pressure of mobile phase Under similar conditions, compounds will have unique Rt and Rf values

Displaying the data Chromatogram - Each compound is produced as a peak - Retention times are used to identify compounds - Area under peaks show the amount of compound present Calibration curve - A line a best fit drawn to scatterplot - Used to interpolate data - Dilute mixture to make extrapolation more accurate

Colorimetry - To determine concentration of coloured compounds - Uses different parts of colour spectrum - Colour of light used is most effective when it is complementary to colour of test cell Colourimeter – measures absorbance of a specific wavelength of light that passes the test cell Conditions: - Substance being analysed should be coloured - Substance should be free of other coloured species that could also absorb the light UV-visible spectroscopy - Measures absorbance of light at a specific frequency - Less likely to suffer interferences from similar coloured compounds - Uses monochromators to give better control of wavelength - Has a wider range - Intensity of light decreases as concentration increases Atomic Absorption Spectroscopy (AAS) - Analyses specific wavelengths and adsorption of radiation

Unit 2 – What makes water such a unique chemical? -

Atoms and molecules absorb and emit electromagnetic radiation when moving from excited states to ground states Only metals ions of interest get excited so very low levels (ppb) can be detected Different parts of the spectrum affect different parts of atom

To calculate concentration 1) Find moles in volumetric flask using 𝑛(𝑣. 𝑓) = 𝐶 × 𝑉(𝑣. 𝑓). Remember C(sample) = C(diluted v.flask) 𝑛(𝑣.𝑓.) 2) Find concentration using 𝐶(𝑜𝑟𝑖𝑔𝑖𝑛𝑎𝑙) = 𝑉(𝑜𝑟𝑖𝑔𝑖𝑛𝑎𝑙) To calculate mass of substance 𝑚𝑎𝑎𝑠𝑠 𝑜𝑛 𝑐𝑢𝑟𝑣𝑒 𝑥 (𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑎𝑚𝑝𝑙𝑒) 1) Find mass in sample using 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑛 𝑐𝑢𝑟𝑣𝑒 = 𝑠𝑎𝑚𝑝𝑙𝑒 𝑣𝑜𝑙𝑢𝑚𝑒

Conical and volumetric flask: - Rinse with distilled water to avoid increasing solution - Any impurities or excess substances may react with chemical or take up volume and result in a lower titre To find concentration of the original 1) Find moles of titre 2) Use mole to mole ratio to find moles of the aliquot 3) Find undiluted moles in original by 𝑚(𝑎𝑙𝑖𝑞𝑢𝑜𝑡) × 4) Find concentration of analyte in original using 𝐶 Titration Curve Strong base/strong acid = 7 Weak base/strong acid = <7 Strong base/weak acid = >7 Weak base/weak acid = around 7

2) m(diluted aliquot) = m(aliquot) 𝑉 𝑖𝑛 𝑣.𝑓. 3) Find mass in undiluted volumetric flask by 𝑚(𝑎𝑙𝑖𝑞𝑢𝑜𝑡) × 𝑉 𝑖𝑛 𝑎𝑙𝑖𝑞𝑢𝑜𝑡 4) m(original) = m(v.flask) IMPORTANT – Remember to check the units of mass/volume

Substance starting in conical flask is first half of graph Substance added to conical flask is second half of graph

Titration Volumetric analysis to find the concentration of another solution Aliquot – the volume of sample delivered by a pipette Titrant – the reactant delivered by the burette Titre – volume of solution delivered by burette to reach the end point Analyte – substance being analysed Concordant titre – three titres that differ by a max of 0.10ml Equivalence point – point where reactants are mixed in stoichiometric ratio with the balanced equation End point – The point where indicator changes colour Pipette and Burette: - Rinse with same chemical - Water in pipette will dilute aliquot and result in a lower titre - Water in burette will dilute titrant and result in higher titre

𝑉 𝑖𝑛 𝑣.𝑓. 𝑉 𝑖𝑛 𝑎𝑙𝑖𝑞𝑢𝑜𝑡 𝑛 = 𝑉

Limiting Reagents - Mole ratio in chemical equation show exact number of moles needed of each chemical - Some moles left unused = excess reagent - All moles completely used up = limiting reagent

Galvanic Cell

Unit 2 – What makes water such a unique chemical? -

converting chemical energy to electrical energy One half cell contains a conjugate pair

2 electrodes – anode and cathode (2 metals) Electrolyte – aqueous conductors; neutral solutions of ionic substances Internal circuit – electrolyte solution does not interfere/react with reaction; allows movement of positive and negative ions External circuit – allows of movement of electrons from anode to cathode Salt bridge: Anion -> Anode KNO3 is the typical salt bridge

Cation -> Cathode

Reduction -> Oxidant/Cathode Oxidation -> Reductant/Anode Higher voltage – lower voltage = overall voltage Electrochemical series - Oxidants at top – Reductants at bottom - Lowest = backward reactions - Reductants must be lower than Oxidants for a strong reaction to occur Significant Figures - Count all non-zero numbers - Count all zeros between ^ - Count zeros after a decimal point if significant numbers are present on the left - Do not count any zeros left of a non-zero number - Do not count zeros right of a non-zero number without a decimal point Use scientific notations to make leading/trailing zeros significant Adding/subtracting - Add/subtract normally - Round to lowest number of decimal places given Multiplication/division - Multiply/divide normally - Round to lowest number of significant figures given 10𝑎 = 𝑏

decimal places of b = sig.fig. of a