ACIDIC ENVIRONENT Indicators • classify common substances as acidic, basic or neutral

- Acid = Substance which in solution produces H ions (H+) specifically hydronium ions (H3O+)o Sour taste o Sting/burn the skin o Solution which conducts electricity o Turns blue litmus red o E.g. vinegar (acetic acid), aspirin (acetyl salicylic acid), lemon (citric acid)

- Base = Substance which contains the oxide (O2-) or hydroxide (OH-) ion or which in solutions

produces hydroxide ion o Alkali is a compound which contains or produces OH- and is a soluble base o Soapy feelo Bitter taste o Good conductor of electricity in solution o Turns red litmus blue o E.g. washing soda (sodium carbonate), drain cleaners (sodium hydroxide), lime

(calcium carbonate)- Indicators used to classify substances as acids or bases

o Litmus paper was the first indicator used • identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour

- Indicator = a substance, usually a vegetable dye, that in solution changes colour depending on whether the solution is acidic or alkaline

- Litmus = dye extracted from lichens mixed with algae first used in 18th C- The range over which indicators change varies for each

o E.g. litmus has 1 colour change whereas universal indicator has 14 • identify and describe some everyday uses of indicators including the testing of soil acidity/basicity

- Practicality = Cheap & convenient way of determining acidity/alkalinity of substances - Everyday uses;

o Testing acidity/alkalinity of soils some plants need acidic soil (flowers) whereas others need and alkaline or neutral soil (vegetables)

o Testing swimming pools Pools need to be approx. neutral however by adding chemicals for sanitation changes this

o Monitoring wastes from labs anything discharged into the sewerage system must be nearly neutral however some solutions (e.g. photographic solutions) are often highly alkaline

o Testing fish tanks similar to testing soil some fish need a specific pH level in order to breed successfully and decrease stress

• perform a first-hand investigation to prepare and test a natural indicator Cabbage and rose petal prac.

- Shredded cabbage was placed in a beaker and covered with distilled water before being boiled and the purple liquid being poured into a separate beaker.

- Shredded pink rose petals were placed in a beaker and covered with distilled water before being boiled and the slightly pink liquid being poured into a separate beaker

- A chart was used with various substances being put in squares - Each substance was then tested for its acidity/alkalinity via the cabbage, rose petal and

universal indicator - Red cabbage juice is red/pink in acidic solutions and purple/green (as H+ decreases) and

yellow in OH- solution- Pink rose petal juice is pink in acidic solution and yellow/green in alkaline solutions

• identify data and choose resources to gather information about the colour changes of a range of indicators

• solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic

- Cleaning products are alkaline - Drinking water is neutral - Fizzy drinks are acidic

Oxides and pollution• identify oxides of non-metals which act as acids and describe the conditions under which they act as acids 

- Acidic oxide; Covalent o Reacts with water to form acido Or reacts with base to form salt o Or does both o E.g. CO2, SO2

- Basic oxide; Ionic o Reacts with acid to form salt o Or Does not react with alkali solutions

- Amphoteric oxide;o React with acids to form saltso And react with alkalis

- Neutral oxides don’t react with acids or bases • analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides 

- Elements to the left of the table are mostly acidic- Elements to the right are mostly basic

- There are 5 amphoteric elements • define Le Chatelier’s principle 

- Le Chatelier’s principle states that if a system at equilibrium is disturbed the system adjusts itself to minimize the disturbance.

- When a system is at equilibrium there are no observable changes in the macroscopic properties.

• identify factors which can affect the equilibrium in a reversible reaction - Change in partial gas pressure;

o Increase partial pressure of a reactant gas shifts equilibrium to favour the products o Increase partial pressure of a product gas shifts the equilibrium to favour the reactants

- Change in temperature; o Endothermic: increase in temperature causes equilibrium to shift to favour the

products o Exothermic: increase in temperature causes the equilibrium to shift to favour the

reactants - Change in concentration;

o Increase concentration of reactants or decrease concentration of products shifts the equilibrium to favour the products

o Decrease concentration of reactants or increase concentration of products shifts equilibrium to favour the reactants

• describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle 

- CO2(g) + H2O(l) H2CO3(aq) represents the equilibrium equation for carbon dioxide dissolved in water.

- Pressureo If pressure of CO2 is increased some will dissolve into the water as H2CO3 to

counteract it therefore the equation will favour the right ()o If pressure of CO2 is decreased then H2CO3 will decompose to CO2 and the

equilibrium will favour the left () - Temperature

o As the temperature increases the solubility of carbon dioxide decreases o Exothermic = heat is liberated o Add heat reaction moves left () to counteract this as it (H2CO3 side) is the one

which absorbs heat • identify natural and industrial sources of sulfur dioxide and oxides of nitrogen  

- Sulfur Dioxide;o N

atural:

Geothermal hot springs and volcanoes

o Industrial: Burning fossil fuels and extracting metals from sulfide ores- NO;

o Natural: Bacteria in soilo Industrial: Nitrogen fertilisers

- NO2;o Natural: Lightningo Industrial: Combustion

- N2O;o Natural: From NO & O2

o Industrial: NO from combustion

• describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen 

- Production of sulfur dioxide o Combustion of coal = S(s) + O2(g) SO2(g) o Extraction of metals from sulfide ores = 2ZnS(s) + 3O2(g) 2ZnO(s) + SO2(g)

- Production of NOo Biological process o Lighting = O2(g) + N2(g) → 2NO(g)

- Production of NO2

o Nitric oxide reacting with oxygen = 2NO(g) + O2(g) → 2NO2(g)

- Production of N2Oo 2NO(g) + O2(g) 2NO2(g)

• assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen 

- Increase in atmospheric concentrations of sulfur and nitrogen oxides due to industrialisation - Direct Evidence;

o No conclusive direct measurements of SO2 or nitrogen oxides over a long period of time due to

Overall levels being extremely low Analysis techniques that are able to detect them have only recently been

developed They’re washed out by the rain

- Indirect evidence;o Acid rain increased during high manufacturing activity which release oxides of sulfur

& nitrogen o Effects of acid rain have become more pronounced

- According to evidence there has not been a significant increase in the atmospheric concentrations of oxides of sulfur or nitrogen

- Don’t have any data beyond 30yrs ago insufficient to claim increase - Higher occurrence of acid rain means there has been an increase in atmospheric

concentrations o Plausible to state an increase in cities & mining areas

- Enhanced greenhouse effect can suggest increase in atmospheric concentrations • calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0˚C and 100kPa or 25˚C and 100kPa 

- V = n x Vmo V= volume of gas (L)o n = Number of moles of gas (mol)o Vm = molar volume of gas (L/mol)

-• explain the formation and effects of acid rain 

- Formation = Caused by nitrogen oxides & sulfur dioxides - Effects

o Erosion of metals, marbles o Damage to forests and vegetation o Acidified soil = inhibit growth of seedlings & basic minerals dissolved

Mineral nutrients required for plant growth removed Insoluble minerals dissolved releasing toxic levels of metal ions = interferes

with normal root uptake o Acidified bodies of water e.g. Scandinavian Lakes

Aquatic organisms become stressed High levels of toxic heavy metals Aquatic invertebrates can’t reproduce

• identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25˚C and 100kPa 

- Aim: To determine the amount of carbon dioxide gas in a can of soft drink.- Method: A sealed coke bottle, beaker, filter paper and stirring rod were all weighed before

starting the experiment. A weighed amount of salt was added to a beaker filled with coke and stirred before leaving to rest for 30 mins and then weighed again. Any measurements taken were recorded throughout the experiment.

- Mass of CO2 released: Material in bottle –final mass of coke = mass of CO2

o 300.71 – 298.92 = 1.79g of CO2 • analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment 

- Sulfur dioxide:o Industrial origin

Processing/burning fossil fuels Extracting metals from sulfide ores

o Reasons for concerns Irritant & Poisonous Causes breathing difficulties Respiratory problems

- Nitrogen oxides:o Industrial origin

Combustion in cars & power plants (NO & NO2) Use of nitrogenous fertilisers

o Reasons for concern Enhanced greenhouse effect Irritates respiratory tract & eyes Causes tissue damage Causes photochemical smog which forms ozone, pollutants (PANS) and haze

pH and Strength• define acids as proton donors and describe the ionisation of acids in water

- Acid = a proton (hydrogen ion) donor - Strong acid = total dissociation & ionisation- Weak acids = ionise to variable extents - Acids dissolve in water = ionisation become ions in and around the water molecules

• identify acids including acetic (ethanoic), citric (2- hydroxypropane-1,2,3- tricarboxylic), hydrochloric and sulfuric acid

- Hydrochloric acid = HCl Strong monoprotic acid used in the production of PVC- Sulfuric acid = H2SO4 Strong diprotic acid used in fertilisers, industrial catalyst- Acetic acid = CH3COOH Weak monoprotic acid used in Soft drinks, glue, vinegar and as a

food additive- Citric acid = C6H8O7 Weak triprotic acid used as a food additive and preservative

• describe the use of the pH scale in comparing acids and bases- Acids = below 7 - Bases = above 7 - Neutral = 7 - pH scale is a measure of concentration of H+ ions in solution - determines acidity/alkalinity of substance based on H+ concentration - Formula for pH = –log10[H+]

• describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute

- Strength of acid determined by ionising ability strength not based on concentration - Strong acid = ionises completely - Weak acid = partially ionises - Concentrated solution = total concentration of solute is high (e.g. more than 5mol/L)- Dilute solution = total concentration of solute is low (e.g. less than 2mol/L)

• identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+]- pH = -log10[H+]

o H+ = concentration of H ions in solution - pH of substance represents negative power of negative log function

o e.g. 10-11 = pH 11- If pH changes by 1 then [H+] changes by 10 (add extra 0)

• compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules

- Degree of ionisation determines an acid’s strength as weak/strong - Formula = [H+] ÷ concentration × 100

Acid Concentration (mol/L)

Strength [H+] Degree of ionisation (prac.) %

Degree of ionisation (theoretical) %

HCl 0.1 Strong 0.058 58 100Citric 0.1 Weak 0.005 5 8.00Acetic 0.1 Weak 0.0003 0.3 1.30

• describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions

- Strong acid ionises completely in solution therefore there will be no HCl intact molecules- Weak acid partially ionises in solution and is thus in equilibrium between its intact molecules

& ions • solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals

- Acid-base indicators used to determine acidity of an unknown substance based on colour change

o Can be affected by the colour of the substance - pH metres & probes are more accurate

o accuracy of 0.01 compared to the 1 of an acid-base indicator o not affected by colour

- pH probes/metres must be calibrated & rinsed before use due to sensitivity • plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids

- Aim: To investigate the pH of equal concentrations of strong and weak acids

- Method: First a 0.1mol/L of citric acid was made up then using this a 0.01 and 0.001 were made. This process was repeated for each acid. (A lot of shaking was involved). A pH strip was placed in each beaker and a pH probe was used in order to have 2 results for the pH of the substances.

- Result: The strong acids were closer to a pH of 1 and degree of ionisation of 100% whereas the weaker acids were not. As the concentration of each acid decreased the degree of ionisation increased

• gather and process information from secondary sources to write ionic equations to represent the ionisation of acids

- HCl = HCl(g) + H2O(l) H3O+(aq) + Cl-

(aq)

- Citric = C3H5O(COOH)(COO)22-

(aq) + H2O(l) H3O+(aq) + C3H5O(COO)3

3-(aq)

- Acetic = CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-

(aq)

• use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids

- Molymod kits were used - Computer generated photos were also looked at

• gather and process information from secondary sources to explain the use of acids as food additives- Acids are frequently added to food in order to:- Improve taste give food a sour/tart taste e.g. soft drinks- Preserve food increase acidity as most bacteria can’t survive in acidic conditions e.g.

canned fruit• identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition

- Citric acid (C6H8O7) = citrus fruits used to control acidity of food products & act as a neutralising/buffer agent. Used in baking powder etc.

- Propanoic acid (C3H6O2) = sugars used as a preservative in bread - Sodium erythobate (C6H7NaO6) = sugar can, beets and corn used to reduce growth of

mould, bacteria & yeast. Used in cookies etc. - Ammonia (NH3) = stale urine and decomposing matter used as a food additive and an

antimicrobial agent. Used in baked goods and meat products etc. • process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations

- pH = -log10[H+]Conjugates, Buffers and Titration• outline the historical development of ideas about acids including those of:

○ Lavoisier: ○ Davy○ Arrhenius

- Lavoisier; 1780o Acids = substances containing oxygeno Disproved many oxygen containing substances are basic e.g. NaOH

Also some acidic substances contained no oxygen e.g. HCl- Davy; 1815

o Acids = substances containing H that could be partly or totally replaced by metalso When acids reacted with metals they formed salts (TRUE)o Bases = substances that reacted with acids to form salt + water (TRUE)

- Arrhenius; 1884o Acids = substances that ionised in solution to produce H+ ionso Determined strength of acid based on ionisation complete ionisation = strong and

partial ionisation = weak (TRUE) o Base = substance that in solution produced OH- ions

Excludes metallic oxides that are basic o FAILED

No recognition to role of solvent ionisation is a reaction b/w acid & solvent

Strength of acid depends on nature of acid AND solvent Circumstances of acid-base reactions occur in solvents in which the ‘acid’

isn’t ionised e.g. reacts to form salt but doesn’t ionise • outline the Brönsted-Lowry theory of acids and bases

- 1923 they separately proposed definitions for acids & bases in terms of proton donors & acceptors

- Acid = proton donor o Substance more likely to give up protons in solvent

- Base = proton acceptor o Substance more likely to accept protons in solvent

- Theory relates acidity & basicity to structure of substance AND solvent o Production of H+ ions due to properties of acid relative to solvent

• describe the relationship between an acid and its conjugate base and a base and its conjugate acid- Acid gives up a proton to form its conjugate base

o acid + H2O H3O+ + Conjugate base- Base accepts a proton to form its conjugate acid

o base + H2O Conjugate acid + OH-

• identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature

- Strong acid + Strong base neutral salto E.g. HCl + NaOH NaCl + H2O

- Strong acid + weak base acidic salto E.g. NH4Cl + H2O NH3Cl + H3O+

- Weak acid + strong base basic salt o E.g. Na2CO3 + H2O Na2CO2 + 2OH-

• identify conjugate acid/base pairsAcid Conjugate base Base Conjugate acidHCl Cl- NH3 NH4

+

H2SO4 HSO4- HCO3

- H2CO3

H3O+ H2O OH- H2O• identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions

- Amphiprotic = acts as a proton donor and acceptoro E.g. HCO3

- can act as a base (form H2CO3) or acid (form CO32-)

- Most common amphiprotic substance is water o Can take a proton (H3O+) mixed with acid or lose a proton (OH-) mixed with base

• identify neutralisation as a proton transfer reaction which is exothermic- Neutralisation = acid + base water + heat

o H3O+(aq) + OH-

(aq) 2H2O(l) + heat - It’s a proton transfer which is exothermic

• describe the correct technique for conducting titrations and preparation of standard solutionsADD TO ALREADY MADE PALM CARD

- Properties of a primary standard:o High level of purityo Accurately known compositiono Free of moistureo Stable & unaffected by air during weighing o Readily soluble in distilled water o Solid with high molar weight to reduce percentage error in weighing

o Reacts instantaneously and completely • qualitatively describe the effect of buffers with reference to a specific example in a natural system

- Buffer = solution that resists changes in pH when an acid/base is added o Contains comparable amounts of weak acid & conjugate base

- General equation = HA + H2O H3O+ + A-

- Example in a natural system;o In blood there are 4 equilibriums o Carbonic acid – hydrogen carbonate buffer linked to haemoglobin – oxyhaemoglobin

equilibrium keeps pH b/w 7.35 & 7.45 In lungs this equilibrium shifts to form more oxyhaemoglobin due to high

amounts of oxygen o As it reaches tissues respiration removes oxygen

Equilibrium shifts so that more haemoglobin is formed as there is less oxygen left in the blood

o This equilibrium is used as a constant buffer to keep blood in the right pH range to ensure correct bodily functions occurring

o basic version of the equilibrium • gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions

- Arrhenius stated when an acid reacts with a base it is the H+ and OH- ions that react to form neutral water

o acid + base = salt + water - For metallic oxides & carbonates

o Do neutralise acids o BUT many carbonates are insoluble in water o ALSO Carbonates & metallic oxides don’t contain OH- which must dissociate into

water o THEREFORE Arrhenius failed

- Arrhenius couldn’t account for salts (with one part the same) had different acidities o E.g. NaCl is neutral whilst Na2S is basic

• choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions

- Aim: To select equipment & perform an experiment to determine the pH of various salt solutions

- Equipment: 7 Salts listed, 7 test tubes, Universal indicator, 7 pipettes & 2 test tube racks- Method: Small amount of each salt solution was placed in a separate test tube and then pH

strip & universal indicator was used to determine pH• perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases• perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies

- Standardised NaOH - Used this to determine quantity of acetic acid in vinegar - Concentration was 0.68mol/L

• analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills

- Spillages of strong acids and bases can occur & thus must be neutralised

o E.g. strong acid spilt, weak base used to neutralise it- Neutralisation reactions are exothermic and thus produce heat - Sodium bicarbonate is used to neutralise acid & base spills because it is;

o A solid thus easily cleaned o Easily & safely stored o Not as dangerous when used in excess compared to other acids/bases

Esterification • describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds

- Alkanols are compounds with an OH functional group- Alkanoic acids contain a carboxylic acid functional group (-COOH) - DIFFERENCE: Alkanes contain a double bonded oxygen atom which allows for 2 hydrogen

bonds whereas alkanols only have a single bonded oxygen and thus only one hydrogen bond o EFFECT: alkanes have a higher melting & boiling point due to extra H-bond

- • identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8

- Esters = compounds formed when alkanoic acids react with alkanols- Naming: remove “-anol” and replace with “-yl” on the alkanol

o Remove “-oic acid” and replace with “-oate” on the alkanoic acid - EXAMPLE : Hexanoic acid + propanol propyl hexanoate

• explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures

- Alkanes, alkenes and alkynes are not polar NO hydrogen bonding o ONLY weak dispersion forces = low boiling point

- Alkanols = 1 centre of polarity, dipole-dipole bonding and hydrogen bonding Higher boiling points

- Alkanoic acids = 2 centres of polarity, dipole-dipole bonding and hydrogen bonding Highest boiling points

- 2 centres of polarity make it harder to separate the molecule & thus higher MP/BP - As number of C atoms increase the MP/BP increases

• identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification

- Esterification = process of making an ester due to the reaction between an alkanol and alkanoic acid

o Alkanoic acid + alkanol alkyl alkanoate + water - EXAMPLE : 1-butanol + ethanoic acid butyl ethanoate

o CH3CH2CH2CH2OH(l) + CH3COOH(l) CH3COOCH2CH2CH2CH3(l) + H2O(l)

o

• describe the purpose of using acid in esterification for catalysis

- It’s a slow process and does not react to completion- Rate of reaction can be increased by

o Adding a catalyst H2SO4 used as it’s a dehydrating agent & when water is removed the equilibrium will counteract this by shifting to the right thus increasing the yield of the ester

NOTE catalysts don’t increase yield they decrease time to reach equilibrium

o Heating the mixture to increase kinetic energy and produce more of the ester • explain the need for refluxing during esterification

- Reactants & products are volatile & readily vaporise Avoid loss of material a condenser is used to circulate cold water cooling the vapours

- Reactants & products have low BP & thus vaporise at low temp. system is open so there’s no build-up of pressure

- Reaction vessel heated by electrical heating mantle clamped to round-bottom flask organic liquids & vapours are flammable

- Vaporisation of reactants occurs quickly small boiling chips prevent ‘bumping’ & provide a large SA for vaporisation to occur without superheating & explosion

• outline some examples of the occurrence, production and uses of esters- Flavouring: characteristic smells & tastes of fruits & flowers e.g. Octyl acetate = orange &

butyl butanoate = pineapple- Perfume/fragrance: Some smells can be extracted to be used in cosmetic perfumes & lotions

and air fresheners e.g. phenyl methyl acetate = jasmine & Methyl salicylate = wintergreen - Solvents & coatings: Many esters used in wood lacquers, thinners and other coatings &

surface finishes. E.g. ethyl ethanoate = nail polish remover, solvent and coating & 1-butyl propanoate = enamels, lacquers and coating

• identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux

- Aim: To prepare an ester using reflux equipment- Method: The alkanol & alkanoic acid were poured into a round bottomed flask with

concentrated sulfuric acid & a few boiling chips. The flask was then set up with the condenser for reflux and was heated (heating mantle) for 30mins.

- Result: Ester produced smelt like banana but the solution smelt like nail polish