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HSC Chemistry Module 9.3 SummaryIndicators were identified with the observation that the colour of some flowers depends on soil compositionClassify common substances as acidic, basic or neutralExamples of common acidic, basic, or neutral substancesAcidicNeutralBasicVinegarPure waterCaustic soda solutionCarbonated soft drinksSugar solutionDrain cleanerCitrus juicesPure alcohol solutionAmmonia cleanserLemon juiceSodium chloride solutionMilk of magnesiaStomach acidEthanolWashing sodaAspirinOilsSoapCar battery acidLactose solutionsLime waterIdentify 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 An indicator is a substance that changes colour depending on the degree of acidity of alkalinity of the solutionMost indicators change colour over a small pH range, and produce 2 colours: one for acidic solutions, and one for basic solutionsExamples of indicators are litmus, phenolphthalein, methyl orange, and bromothymol blueEach of the above substances changes colour over a limited range:Litmus: 4.5-8.5Phenolphthalein: 8.2-10.0Methyl orange: 3.1-4.4Bromothymol blue: 6.0-7.6Indicators can be used in combination to obtain a more exact pH valueIdentify data and choose resources to gather information about the colour changes of a range of indicatorsBelow are the pH colour change ranges for the above indicators. MEMORISE!!!An indicator changes colour according to the concentration of H3O+ in the solution, as it shifts the equilibrium of the indicator solution, were an=anionIn an indicator, Han and an- have different coloursThe H3O+ concentration shifts the equilibrium, and thereby producing more Han or an-, and thus the solution changes colour according to pHIdentify and describe some everyday uses of indicators including the testing of soil acidity/basicitySoil testingSoil acidity is an important factor in determining what plants and crops can be grown, so knowing the pH of soil is vitalA sample of soil is taken, and mixed in distilled water to make a slurryA neutral white insoluble powder (barium sulfate) is sprinkled over the sample, and a suitable indicator is added. The white powder allows the colour to be observed more clearly.If the soil is too acidic, chemicals such as lime (calcium oxide) are added to make the soil more basicPool testingThe acidity of swimming pools needs to be monitored and controlled to prevent the growth of microbes, whilst avoiding skin and eye irritationWater is sampled, and a few drops of indicator is addedPool chlorine or hydrochloric acid is added to achieve a suitable pHMonitoring pH of chemical wastesWastes from laboratories or photographic film centres can be highly acidic or alkalineDischarges into the sewerage system need to be nearly neutral, so the pH of wastes are monitored, and neutralised if necessarySolve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basicEXAMPLEA drain cleaner’s pH was tested using litmus paper and phenolphthalein. The litmus turned blue, and the phenolphthalein turned pink. What can be said about the pH of the drain cleaner?As the litmus paper turned blue, the drain cleaner is basicAs the phenolphthalein turned pink, it must be strongly basicSome dilute vinegar was also tested using methyl orange, which turned yellow, and phenolphthalein, which turned pink. What can be said about its pH?The test is inconclusive, as the ranges do not definitively determine whether the vinegar is acidic, basic, or neutral.Perform a first-hand investigation to prepare and test a natural indicatorMETHODOne large beetroot was peeled and chopped, then placed into a beaker of 100mL of distilled boiling water, heated with a Bunsen burner. 5mL of HCl, NaOH, distilled water, and NaCl each were added to four test tubes, and then 3mL of beetroot juice was added to each test tube. Any colour changes were observed.SAFETY: Acids are corrosive, bases are caustic => use low concentrations (1M), wear safety glassesRESULTSThe beetroot juice remained red in water, NaCl, and HClThe beetroot juice turned yellow in NaOH solutionThus beetroot juice can be used to distinguish basic substances by changing colour from red to yellow.A beetroot was used as its pigmentation can be easily extractedVALIDITY/ACCURACY/RELIABILITYResults were compared to other class results, which were similarHCl and NaOH were used to test the aim, by measuring a wide range of the pH scaleDistilled water minimised error in the preparation of the beetroot juiceA fresh beetroot was used instead of a canned one as canned beetroots contain preservatives (such as acids) that may alter the resultsThe experiment was limited because…The size of the beetroot was not controlledThe exact pH transition was not measured => could be measured by titration with data loggers and sensorsWhile we usually think of the air around us as neutral, the atmosphere contains acidic oxides of carbon, nitrogen and sulfur. The concentration of the acidic oxides have been increasing since the Industrial RevolutionIdentify oxides of non-metals which act as acids and describe the conditions under which they act as acidsACIDIC OXIDESCompounds that contain oxygen that act as acidsReact with water to produce an acidReacts with bases to produce saltsNon-metals oxides mainly behave as acidsEXAMPLES OF ACIDIC OXIDESSulfur dioxide (SO2) (sulfurous acid)Carbon dioxide (CO2) (carbonic acid)Nitrogen dioxide (NO2) (nitric and nitrous acid)Diphosphorus pentoxide (P2O5) (phosphoric acid)BASIC OXIDESReact with acids to form saltsDo NOT react with alkali solutions (an alkali is a water-soluble base)Metal oxides mainly behave as basesEXAMPLES OF BASIC OXIDESPotassium oxide (K2O) (potassium hydroxide)Magnesium oxide (MgO)(magnesium hydroxide)AMPHOTERIC OXIDESAmphoteric oxides can react with both acids and basesFor example, aluminium oxide (Al2O3) can react with hydrochloric acidBUT can also react with sodium hydroxideNOTE!!! => Amphoteric substances can REACT with both acids and bases, but do not necessarily ACT as both acids and bases. Such substances are called amphiprotic substancesNEUTRAL OXIDESSome elements form neutral oxides, such as hydrogen (H2O), carbon monoxide (CO), dinitrogen oxide (N2O), and nitric oxide (NO)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 oxidesGENERALISATION/TRENDSNon-metal oxides are acidicThe further to the right & top of the Periodic Table (except the noble gases), the more acidic its oxide isThis is because the more electronegative the element is, the greater ionisation occurs when in contact with water (~100% for S, N, Cl)Metal oxides are basicThe further to the left & bottom of the Periodic Table, the more basic its oxide isThis is because the less electronegative elements ionise easily in water to form oxide ions, which react with water to form hydroxide ionsAmphoteric oxides are between the acidic and basic oxidesNoble gases have no oxidesSome other elements (H, N, C) form neutral oxidesDefine Le Chatelier’s principleRecall that not all chemical reactions go to completion, i.e. the reaction goes in one direction (reactants to products)Not all reactions are one-directional => many reactions consist of a forward reaction (left to right) and a reverse reaction (right to left)These reactions are reversible reactionsOne example of a reversible reaction is a saturated solution of a salt in waterAnother example is the reaction of carbon dioxide in waterReversible reactions do not go to completion, but reach a chemical equilibriumAt chemical equilibrium, the concentration of reactants and products do not changeThe forward and reverse reactions still continue, but proceed at the same rate (the equilibrium is dynamic)Chemical equilibrium occur in closed systems, where the macroscopic (observable) properties are constantIf the conditions of the system change, the system will no longer be at equilibriumLe Chatelier’s principle states that:If a system at equilibrium is disturbed, then the system adjusts itself so as to minimise the disturbanceA system at equilibrium will readjust to oppose any changesIdentify factors which can affect the equilibrium in a reversible reactionSome of the factors that can affect the equilibrium in a reversible reactions are the concentration of particular chemical species, the total pressure of the system (only for reactions that involve gases), and the temperatureCONCENTRATIONBlueGreenConsider the following four scenarios guided by Le Chatelier’s principle:If the concentration of Cl- ions are increased, the equilibrium shifts to the right (i.e. the rate of the forward reaction increases), so the system becomes more green as more CuCl42+ ions are producedIf Cl- ions are removed, the equilibrium shifts to the left, and the system becomes more blueIf water is added, the equilibrium shifts to the left, and the system becomes more blueIf water is removed, the equilibrium shifts to the right, and the system becomes more greenGenerally…If the concentration of a particular species is INCREASED, the equilibrium point will shift to the OPPOSITE side of the equation to oppose the disturbance by reducing the concentration of that species.If the concentration of a particular species is DECREASED, the equilibrium poin till shift to the SAME side of the equation to oppose the disturbance by increasing the concentration of that species.PRESSUREBrownColourlessRecall that equal number moles of gases occupy equal volumesThe pressure of a closed system can be increased by reducing the volume of the system, and can be decreased by increasing the volume.If the pressure is increased, the equilibrium will shift to the right, as the right has less moles of gas (therefore occupies less volume), and the system becomes more colourlessIf the pressure is decreased, the equilibrium will shift to the left, as the left has more moles of gas (therefore occupies more volume), and the system becomes more brownGenerally…If the pressure of a system is increased, the equilibrium will shift to the side with fewer moles to decrease the pressure of the system.If the pressure of a system is decreased, the equilibrium will shift to the side with more moles to increase the pressure of the systemTEMPERATURERecall that an endothermic reaction absorbs heat, whilst an exothermic reaction releases heat.By using Le Chatelier’s Principle, the following predictions can be made about a closed system:An increase in temperature will increase the rate of the endothermic reactionA decrease in temperature will increase the rate of the exothermic reactionIf the ΔH of a reversible reaction is given, it is referring to the forward reactionRemember that ΔH is positive for endothermic reactions, and negative for exothermic reactionsEXAMPLEColourless WhiteIn the above reaction, the forward reaction is exothermic, and the reverse reaction is endothermicIf the temperature of the system is increased, the equilibrium shifts to the left (favouring the endothermic reverse reaction), and the system becomes more colourlessIf the temperature of the system is decreased, the equilibrium shifts to the right (favouring the exothermic forward reaction), and the system becomes more whiteDescribe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principleThe reaction between carbon dioxide and water is a reversible reaction that goes to equilibriumThe equilibrium between carbon dioxide and carbonic acid can be affected by concentration, pressure, and temperature.EFFECT OF CONCENTRATIONThe effect of concentration on the equilibrium can be explained by Le Chatelier’s principleIf the concentration of CO2 is INCREASED, the equilibrium will shift to the right to dissolve more CO2, thereby opposing the increase of CO2If the concentration of CO2 is DECREASED, the equilibrium will shift to the left to produce more CO2, thereby opposing the decrease of CO2The effect of acids (H3O+) and bases (OH-) can also be explained by considering Le Chatelier’s principleIf an acid is added, the concentration of H3O+ increases, so the equilibrium shifts to the left to decrease the concentration of H3O+.If a base is added, the OH- ions react with the H3O+ to produce water. Consequently, the equilibrium shifts to the right to produce more H3O+ ions to counteract the change.PRESSUREThe left side of the equation contains one mole of gas (CO2), whilst the right side of the equation contains zero moles of gas. Hence:If the pressure is INCREASED, the equilibrium shifts to the right, thereby reducing pressure by dissolving CO2 gas (decreasing the volume of gas decreases the pressure)If the pressure is DECREASED, the equilibrium shifts to the left, thereby increasing the pressure by producing CO2 gas.TEMPERATURENote that because ΔH is negative in the above equation, the forward reaction is exothermic, and the reverse reaction is endothermic. Hence:If the temperature is INCREASED, the equilibrium will shift to the left (i.e. the reverse reaction will be favoured), as it is an endothermic reaction, which will oppose the change by cooling the systemIf the temperature is DECREASED, the equilibrium will shift to the right (i.e. the forward reaction will be favoured), as it is an exothermic reaction, which will oppose the change by warming up the systemIdentify natural and industrial sources of sulfur dioxide and oxides of nitrogenSulfur dioxide (SO2)Most of the atmospheric sulfur dioxide comes from the oxidation of sulfur in compoundsNATURAL SOURCES include volcanic activity/gases, geothermal hot springs, bushfires [H2S], bacterial decomposition [H2S]INDUSTRIAL SOURCES include the smelting of metallic sulfide ores (e.g. Cu2S), and the combustion of fossil fuels (particularly coal) containing metallic oxides (e.g. FeS2) or carbon compounds containing sulfurOxides of nitrogen (NOx)NOx is the general for oxides of nitrogen. These oxides include nitric oxide (or nitrogen monoxide, NO), nitrogen dioxide (NO2), and nitrous oxide (or dinitrogen monoxide, N2O)Nitric oxide (NO)Nitric oxide is primarily produced through the reaction of nitrogen and oxygen gas in high-temperature combustion environments (particularly high voltage)NATURAL SOURCES include lightning, which is a localised, high-temperature, and high-voltage environmentINDUSTRIAL SOURCES include the combustion of fossil fuels, particularly in car engines and at power stationsNitrogen dioxide (NO2)Much of the atmospheric nitrogen dioxide is produced through the slow reaction of nitric oxide and oxygenNATURAL SOURCES include the conversion of nitric oxide produced by lightningINDUSTRIAL SOURCES include the direct emission of NO2 from motor vehicles and power stations, and the conversion of NO released from vehicles and power stations.Nitrous oxide (N2O)NATURAL SOURCES include the natural actions of nitrogen-fixing bacteria on nitrogenous materials in the soilINDUSTRIAL SOURCES include the increased use of nitrogenous fertilisers, which is more food for bacteria to convert to nitrous oxideDescribe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogenSulfur dioxideAs mentioned above, the general reaction that produces atmospheric SO2 is the oxidation of sulfur in sulfur-containing compoundsFor example, sulfur dioxide is produced when organic matter releases H2S during decomposition:The smelting of zinc/iron also releases sulfur dioxide into the atmosphereOxides of nitrogenAs mentioned above, nitrogen and oxygen gas react in high-temperature environments (such as lightning or combustion engines) to produce nitric oxideNitric oxide and oxygen can slowly react to produce nitrogen dioxideThe reverse equation can proceed in the presence of sunlightAssess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogenThere is extensive evidence to suggest that there has been an increase in the atmospheric concentration of oxides of sulfur and nitrogen.Qualitative evidence includes:The increased formation of photochemical smog in urban areas, which indicates an increased concentration of NO2e.g. in 1952, 4000 people were killed due to smog in London, leading to environmental controls in the U.K.Increased incidence of acid rain indicates an increase in atmospheric SO2 and NOxIncreased use of motor vehicles, combustion of fossil fuels for electrical production, and smelting of sulfide ores since the Industrial Revolution would logically suggest an increased concentration of SO2 and NOxTHIS EVIDENCE IS CIRCUMSTANTIAL HOWEVER => it does not prove with certainty that concentrations of oxides of sulfur and nitrogen have increased. Nonetheless, they do provide a strong indication of increased levels of oxides of sulfur and nitrogen.Quantitative evidence for increasing levels globally has been difficult to obtain for the following reasons:Oxides of sulfur and nitrogen occur in relatively low concentrations => the global average is 0.001ppm, whilst peak levels in polluted urban areas such as Los Angeles peak at 0.47ppm. For comparison, CO2 occurs at ~360ppm.Chemical instruments able to measure these very low concentrations have only been available since the 1970s, so not reliable data is available before thenSO2 and NOx are dissolved by atmospheric water, so measuring their levels globally is difficult.Quantitative measurements for oxides of sulfur and nitrogen locally is availableThe NSW EPA monitors atmospheric pollutant levels across Sydney, and has shown that the levels of SO2 and NOx have stabilised over the past two decades, and vary seasonally.This is due to emission controls that have been progressively implemented since the 1950s.The stabilised trend is similar in other Western cities, though levels in industrialising areas (notably China) have nearly tripled over the past twenty yearsExplain the formation and effects of acid rainFORMATION OF ACID RAINRecall that pure water has a pH of 7 (i.e. is neutral)Unpolluted rain water typically has a pH of 6.0-6.5 due to atmospheric CO2, which dissolves in water to form carbonic acid (as discussed above)Acid rain is rain that has a pH less than 5.0Acid rain in industrial areas typically occurs due to the presence of SO2 and NOx, which react with water to form acidsSulfur dioxide reacts with water to produce sulfurous acidSulfurous acid reacts with oxygen (catalysed by air particles) to produce sulfuric acidNitrogen dioxide reacts with water to produce nitric acid and nitrous acidNitrous acid reacts with oxygen (catalysed by air particles) to produce nitric acidBoth sulfuric and nitric acids are strong acids, thus both can significantly reduce the pH of rain water to unnatural levels (pH of 3-5)EFFECTS OF ACID RAINIncreasing acidity of lakes, which can kill marine life such as fish by killing eggs, and irritating skin and gillsDamage to forests through the leeching of toxic ions such as aluminium and mercury. Also, the decrease in the pH of soil makes it difficult for plants to absorb sufficient calcium, potassium or potassium, and can kill bacteria important micro-organisms in the soilErosion of marble and limestone decorations, as such materials contain carbonates (primarily calcium carbonate), which readily with acids.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 environmentSee above for the industrial origins of sulfur dioxide and oxides of nitrogen, and the relevant chemical equations.The reasons for their concern are significant, due to the resultant health problems, and environmental problems.Health problemsSulfur dioxide is a severe respiratory irritant, and causes breathing difficulties at concentrations as low as 1ppm. People suffering from asthma and emphysema are particularly susceptible.Nitrogen dioxide irritates the respiratory tract, and causes breathing discomfort at 3-5ppm. At higher concentrations it can destroy tissue, as it forms nitric acid, which is a strong acid.Nitrogen dioxide also leads to the formation of photochemical smog, which leads to the production of ozone, and other particulates (haze) and pollutants such as PAN (peroxyacylnitrates). Ozone has harmful effects on the body at concentrations as low as 0.1ppm (see Module 9.4 section 4)Environmental problemsBoth sulfur dioxide and oxides of nitrogen lead to the formation of acid rain, which leads to environmental destruction (see below)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***NOTE*** => ALWAYS note the number of significant figures given in a question. The final answer cannot have more significant figures than the data with the least significant figures.CALCULATING VOLUMES GIVEN THE MASSES OF SOME SUBSTANCESGiven the mass of a substance, the volume of a gas can be calculated by the following steps:Write the relevant chemical equationConvert mass to molesUse the stoichiometry of the equation to work out the number of moles of each chemical species as requiredConvert moles to volumeNOTE: The molar volume of gases at 0°C and 100kPa or 25°C and 100kPa is given on the chemistry data sheet in examsEXAMPLECalculate the volume of sulfur dioxide produced when 50g of hydrogen sulfide is oxidised at 25°C and 100kPa.First write the relevant chemical equationNext calculate the number of moles of H2SFrom the stoichiometry of the equation, we can see thatThusNow we can calculate the volume of SO2CALCULATING MASSES GIVEN GASEOUS VOLUMESGiven the volume of a gaseous substance, the mass of a particular species can be calculated by the following stepsWrite the relevant chemical equationConvert volume to molesUse the stoichiometry of the equation to work out the number of moles of each species as requiredConvert moles to massThe steps involved are the reverse of the example given above.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 100kPaMETHODA 300mL bottle of soda water was weighed on electronic scales. The can was then vigorously shaken, and the cap was opened. The bottle was shaken periodically for one hour, whilst ensuring no drink spilled, and the final mass of the drink was recorded.Make sure no drink spills, because it reduces the validity of the experiment, and poses a safety hazard.RESULTSInitial mass: 330.36gFinal mass: 328.50gMass loss: 1.86gVolume of CO2 released: 1.05L (at 25°C 100kPa)In the experiment, we assume that any mass lost is due to carbon dioxide lossTHEORYCarbon dioxide exists in soda water by the following equilibrium:By shaking the bottle, we introduce additional energy (kinetic) to the soda water.By Le Chatelier’s Principle, the equilibrium opposes the increase in energy by favouring the reverse reaction, as it is an endothermic reaction that absorbs heat, thus CO2 gas evolves.RELIABILITYThe experiment was not repeated, no average was taken, the range of data wasn’t identified, and no outliers were identified => unreliable methodTaking repeated results would improve reliability, but they must be conducted UNDER IDENTICIAL CONDITIONSVALIDITY/ACCURACYVariables that could affect the investigation (namely the conditions of the CO2/H2CO3 equilibrium) were controlled, ensuring a fair testThe use of electronic scales increased the accuracy of resultsThe validity of the experiment could have been improved by having a bottle of water open as a controlA more accurate result could have been obtained through a titrationFactors to consider in the validity…Record the original mass of the soft drink whilst it is in the bottleEnsure that no soft drink spills in the experimentMinimise the evaporation of water (do not heat, keep lid on top etc.), as this would alter the mass change and hence be an invalid testAttempt to evolve all the gas from the soft drinkUse soda water as opposed to cola or other soft drinks, as most soft drinks contain other acids (e.g. H3PO4) for flavouring, which may also exist in an equilibrium, and thus alter the resultsAcids occur in many foods, drinks and even within our stomachsDefine acids as proton donors and describe the ionisation of acids in waterAcids are called proton donors, as they donate protons (H+ ions) in reactionsEXAMPLEIn the above reaction, HCl has donated a proton (H+) to the OH- ion I NaOH to form H2OAcids are ionised in water, as the acid donates a proton to a water molecule to form a hydronium ion (H3O+)EXAMPLEHCl donates a proton H2O to form the hydronium ion (H3O+) and a chloride ion, which is solvated by the waterIdentify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acidACETIC ACID3171825114300Systematic name: Ethanoic acidMolecular formula: CH3COOHWeak monoprotic acid (4.2% ionisation at 0.010molL-1)Present in vinegarManufacture to make organic chemicals4000500318770CITRIC ACIDSystematic name: 2-hydroxypropane-1,2,3-tricarboxylic acidMolecular formula: C6H8O7Weak triprotic acid (27.5% ionisation at 0.010molL-1)Present in citrus fruitUsed as a food additiveHYDROCHLORIC ACIDMolecular formula: HCl3362325128270Very strong monoprotic acid (100% ionisation)Present as stomach acidUsed in swimming pools3018790133350SULFURIC ACIDMolecular formula: H2SO4Very diprotic strong acidUsed to make fertiliserIdentify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+]pH is defined by the following equation:OR As pH is a logarithmic scale base 10, a change in pH of 1 means a ten-fold change in [H+]e.g. A pH of 3 means [H+] = 1.0x10-3molL-1, whilst pH of 4 means [H+]=1.0x10-4molL-1The logarithmic scale was chosen for convenience of calculationsAnother scale used is pOH, whereThe relationship between pH and pOH areNOTE: In calculations, make sure that you use the concentration of H+ ions, NOT the number of molesProcess information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrationsCalculating the pH of strong acids:Calculate the concentration of the acid from the given data (if in doubt, convert any chemical data to moles, then divide by volume)Write the acid’s ionisation equation MonoproticDiproticTriproticUse the stoichiometry of the equation to calculate the concentration of hydrogen ions => multiply by the ratioUse the pH equation to calculate the pH of the substanceEXAMPLECalculate the pH of 0.50mol of sulfuric acid, if the volume of acid is 5.0LThe ionisation reaction isThusREMEMBER***VERY IMPORTANT NOTE*** SIGNIFICANT FIGURES!!! The rules for significant figures in pH is an EXCEPTION to other calculations => the number of significant figures in the hydrogen ion concentration should equal the number of decimal places in the pH value (i.e. the whole number doesn’t count in significant figures)For example, if [H+]=0.00100, then pH=3.000 not and I repeat NOT 3.00Use log10 (log base ten) on the calculator, NOT the natural logarithm (loge or ln)When using the pH formula, ensure that you are using the concentration of H+ ions, NOT the number of molesIf the acid is polyprotic, and the concentration of the acid is given, you must multiply the concentration according to the number of ionised hydrogen ions to find the hydrogen ion concentration => writing the ionisation equation significantly helpsIf the concentration of hydroxide ions is given, calculate the pOH of the solution, and find the pH by using the relationship pH+pOH=14If the concentration of a weak acid is given, and the degree of ionisation, multiply the concentration of the acid by the degree of ionisation (as a percentage) to find the concentration of ionised acid particles => continue from step 2Describe the use of the pH scale in comparing acids and basesThe pH scale is a logarithmic scale to determine the acidity or basicity of a substanceThe pH scale is a measure of the hydrogen ion [H+] or hydronium ion [H3O+] in a solutionA substance is classed as acidic or basic according to its pH in reference to the pH of pure waterPure water exists in an equilibrium with hydronium and hydroxide ionsCareful measurements have shown thatThus neutral substances have a pH of 7Acidic substances have a pH < 7 => the lower the pH, the more acidic a substance isBasic substances have a pH > 7 => the higher the pH, the more basic a substance isDescribe acids and their solutions with appropriate use of the terms strong, weak, concentrated and diluteSTRENGTHS OF ACIDS => DEGREE OF IONISATIONA strong acid is an acid that completely ionises in water solutionOnly acids that achieve 100% ionisation are classed as strongThe ionisation reaction goes to completionEXAMPLEThe following table provides a list of strong acids and bases. MOST OTHER ACIDS OR BASES ARE WEAK => MEMORISE THIS TABLEACIDSBASESNameFormulaNameFormulaHydrochloric acidHClPotassium hydroxideKOHNitric acidHNO3Sodium hydroxideNaOHSulfuric acidH2SO4Lithium hydroxideLiOHHydrobromic acidHBrRubidium hydroxideRbOHHydroiodic acidHICaesium hydroxideCsOHPerchloric acidHClO4In a weak acid, only some of the acid molecules ionise to form hydronium ions => they only partially ioniseA weak acid is releases less protons than a strong acid of the same concentrationWeak acids ionise to variable extents => they have a degree of ionisationA weak acid reaches an equilibrium between ionised and intact moleculesEXAMPLESTRENGTH IS NOT RELATED TO CONCENTRATIONCONCENTRATION OF ACIDS => CONCENTRATION OF PARTICLESA concentrated acid has a high concentration of acid particles (>~5molL-1)A dilute acid has a low concentration of acid particles (<~2molL-1)pH measures the concentration of an acid, hence a strong acid could still have a relatively high pH (i.e. a strong acid can be dilute)BE CAREFUL TO DISTINGUISH BETWEEN STRENGTH AND CONCENTRATION => a strong acid can be both concentrated and dilute, so can a weak acidCompare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their moleculesHCl => Citric => Acetic (decreasing strength)The following table shows the pH of each acid at 0.100molL-1AcidpH at 0.100mol-1Hydrochloric acid (HCl)1.0Citric acid (C6H8O7)2.1Acetic acid (CH3COOH)2.9By using the pH formula, we can calculate the concentration of hydrogen ions [H+] in each solutionHCl: [H+] = 0.1molL-1Citric: [H+] = 0.0079molL-1Acetic: [H+] = 0.00126molL-1From the above data we can calculate the degree of ionisation of each acid dividing [H+] for each acid by the total concentration of substanceHCl: 0.1/0.1 = 100% ionisationCitric: 0.0079/0.1 = 7.1% ionisationAcetic: 0.00126/0.1molL-1 = 1.26% ionisationThus acetic acid is the weakest acid, as only 1.26% of the acid particles ionise (i.e. it has the lowest degree of ionisation) => it is a very weak acidCitric acid is slightly stronger than acetic acid, as a greater percentage (7.1%) of its molecules ionise in solution => it is a weak acidHydrochloric acid by far the strongest acid, as it achieves 100% ionisation of its molecules => it is a strong acidNOTE => the degree of ionisation of a weak acid depends on the concentration of the acid => changing the concentration of a weak acid solution can change the acid’s strength.Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ionsThe strength of an acid depends on its degree (100%) ionisationA strong acid will completely ionise (100 in solution => the ionisation reaction goes to completionA weak acid will partially ionise in solution => the ionisation reaction goes to an equilibriumThe greater proportion of an acid’s ions to intact molecules, the stronger the acid isIn other words, the faster the rate of the forward reaction, the stronger the acid isIn the diagram below, HA is a stronger acid than HB, as there is a greater ratio of A- ions to HA molecules than B- to HB in their respective equilibriumsFor example, a solution of acetic acid reaches an equilibrium between intact molecules and ionsTypically acetic acid’s degree of ionisation is ~1.3%An equilibrium exists between intact acetic acid molecules and its ions, but the proportion of ions to intact molecules is very low ([CH3COOH]/[CH3COO-]), hence acetic acid is a weak acidWeak acid equilibriums adhere to Le Chatelier’s principle, thus the pH of an acid can change according to the factors that influence the equilibriumGather and process information from secondary sources to write ionic equations to represent the ionisation of acidsRecall that strong acids go to completion. Generally…EXAMPLEWeak acids on the other hand reach an equilibrium. Generally…EXAMPLENOTE: Polyprotic ionise in steps, and the degree of ionisation for each step significantly decreaseEXAMPLEUse available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acidsAbove is a model of the molecular nature of acids, and their subsequent ionisationThe beaker on the left contains a strong acid. As can be seen, there are no intact molecules, as the acid has completely ionised, releasing one H+ ion per moleculeThe beaker on the right contains a weak acid of equal concentration. The acid has only slightly ionised, with most of the molecules remaining intact. The concentration of H+ ions is significantly less than in the left beaker, and hence its pH would be much higher.The ionisation of strong and weak acids can also be modelled using molecular model kits, which can demonstrate the breaking and formation of chemical bonds during the ionisation processGather and process information from secondary sources to explain the use of acids as food additivesWeak/dilute acids are used as food additives for two primary reasons: as preservatives (as many microbes can’t survive in acidic conditions), and as flavouring (adds a certain tartness, or sharp/sour flavour)PRESERVATIVESAcetic/ethanoic acid (CH3COOH) => Used as vinegar (4% solution) to preserve food by ‘pickling’Propanoic acid => controls bacteria and mould growth, particularly in bread, potato crisps, and cake mixesCitric acid => natural preservative, often added to jams and conservesTartaric acid => preservative in jams, fruits, pickles, and soft drinksFLAVOURINGSPhosphoric acid (H3PO4) => added to soft drinks for a tartness of flavourCarbonic acid (H2CO3) => adds the ‘fizz’ to soft drinksAcetic/ethanoic acid => also used as acidic flavouring, such as salad dressingMalic acid => flavour enhancer particularly in fruit fillings; adds a savoury tasteIdentify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical compositionNaturally occurring acidsNaturally occurring basesNameChemical formulaNameChemical formulaStomach acid (hydrochloric acid)HClAmmoniaNH3Vinegar (acetic acid)CH3COOHMetallic oxidesCuO, Fe2O3Citric acidC6H8O7Limestone (calcium carbonate)CaCO3Lactic acidC3H6O3NicotineC8H14N2Acid rain (carbonic acid)H2CO3Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicalsMETHODA range of substances dissolved in 20mL solutions were placed in separate beakersThe substances tested were washing powder, an antacid tablet, table salt solution, milk, a vitamin C tablet, and aspirinEach substance was tested with a pH probe and data logger to determine its pH2 drops of universal indicator was added to each test tube, and the colour was checked against a pH chartSAFETY: General risks of using acids => use low concentration, and use a dropper when transferring solutionsRESULTSSubstancepHColourClassificationWashing powder8.0-9.0BlueBaseAntacid tablet8.0-9.0BlueBaseTable salt6.5-7.5GreenNeutralMilk6.5-7.5GreenNeutralVitamin C tablet5.0-6.0YellowAcidAspirin5.0-6.0YellowAcid RELIABILITY/ACCURACY/VALIDITYUniversal indicator was used as it could distinguish between acids, bases, and neutral substancesThe use of a pH probe provided more accurate quantitative resultsEqual quantities of each substance was dissolved in 20mL of water to make a fair test of the substance’s pHA wide variety of substances were tested, and results matched expected results and other results in the classPlan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acidsRefer to practical 9.2.3j)METHOD50mL of 0.1M solutions of sulfuric acid, hydrochloric acid, citric acid, and acetic acid were placed into four separate small beakers. The pH of each solution was tested with a pH probe.NOTE: Ensure that the pH probe is always wetSAFETY: Acids are corrosive substances => use dilute concentrations, use a dropper to transfer acid, wear safety gogglesRESULTSSulfuric acid (H2SO4): 2.20Hydrochloric acid (HCl): 2.37Citric acid (C6H8O7): 2.85Acetic acid (CH3COOH): 3.44Sulfuric acid had the lowest pH, because it is a strong diprotic acid, whilst hydrochloric acid had a higher pH because it is a strong monoprotic acid. Citric acid and acetic acid are both weak acids, but citric acid ionises to a greater extent than acetic acid, and so has a lower pH.VALIDITY/ACCURACY/RELIABILITYComputer technologies were used to maximise accuracy and reliability.Results were compared to other results in the class and were found to be similar, and matched expected resultsOther variables (e.g. temperature, pressure) were controlled, ensuring a fair testThe pH values did not match the expected theoretical results because...The theoretical values are taken at standard laboratory conditionsThe standard solutions prepared for the classroom are not accurate => taken to one significant figureBecause of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refinedOutline the historical development of ideas about acids including those of:LavoisierDavyArrheniusLAVOISIER => proposed that acids contained OXYGENStated that acids are corrosive, sour-tasting substances containing oxygenOxygen was only a recently-discovered elementSince most known acids contained oxygen (e.g. CH3COOH, H2SO4, H2CO3), he believed all acids must contain oxygenLavoisier only experimented with oxyacidsHe believed that oxygen was the source of acidityDAVY => proposed that acids contained HYDROGENDavy demonstrated that HCN and HCl did not contain oxygen, thus disproving Lavoisier’s hypothesisOther hydrohalic acids had recently been discovered (e.g. HBr, HF), leading Davy to propose that hydrogen gave acids their acidic propertiesThis led to his hypothesis that all acids contain hydrogenGerman scientist von Liebig later extended Davy’s theory on acids to state that acids contained replaceable hydrogen, thus explaining why methane (CH4) was not acidicARRHENIUS => proposed that acids IONISED in WATER to produce H+ IONSDavy’s hypothesis was limited as it did not explain why many acidic properties, such as the production of NO2 instead of hydrogen gas when an acid reacted with a metalAfter Arrhenius’s extensive work on hydrolysis, he discovered that hydrogen gas evolved at the cathode during the electrolysis of waterThis lead him to propose that acids dissociated in water to produce H+ ions, which then reacted during hydrolysis to produce hydrogen gasFor example, hydrogen chloride gas ionised in solution to produce H+ ions.He also proposed that bases dissociate in water to produce hydroxide ions (OH-)He recognised that some acids were weaker than others (e.g. acetic acid), and proposed that weaker acids only partially ionised in water => this led to the development of the pH scaleGather and process information from secondary sources to trace developments in understanding and describing acid/base reactionsThe definition of acids/base reactions have changed throughout history, due to advances in technology and understanding have altered the direction of scientific thinkingBoth the Lavoisier and Davy model of acids were limited, and are not used todayThe Arrhenius model of acids/bases is used as a simple model todayThe current models of acid/base reactions are the Br?nsted-Lowry model, and the Lewis model (not required for HSC)TheoryAcid definitionBase definitionDevelopmentLimitationsLavoisierContained oxygenNo definitionInvestigated non-metallic oxides that produced oxyacids in water (e.g. H2SO4). Oxygen had recently been discovered.Substances not containing oxygen (e.g. HCl, HCN) were shown to be acidicDavyContained hydrogenNo definitionInvestigated hydrohalic acids (e.g. HCl, HBr)Did not explain the production of other gases (e.g. NO2) in acid/metal reactionsArrheniusIonised in water to produce H+ ionsIonised in water to produce OH- ionsInvestigated electrolytes and electrolysis of acids. Explained how acids (H+) and bases (OH-) react to produce water in neutralisation reactions.Does not explain why metallic oxides and carbonates are basic. Does not explain acidic/basic salts. Does not explain anhydrous neutralisation reactions.Br?nsted-LowryProton (H+) donorProton (H+) acceptorInvestigated the problems with Arrhenius’s definition of acidsLimited to reactions containing hydrogen => Lewis acidsEach definition has had an impact on our understanding of the nature and properties of acids/bases and their reactionsLavoisier’s definition of acids was wrong, but did increase awareness of the need to define an acidDavy’s definition helped classify substances as acidic, and directed scientific thinking on acids towards the study of hydrogen’s roleArrhenius’s definition helped interpret acid properties in terms of the hydrogen ions produce, though his definition of bases was limited => this was a significant increase in our understanding of acids, and in the development of the concept of an acidBr?nsted-Lowry’s definition further increased our understanding by allowing for more accurate quantitative analysis of acids (e.g. pH, treatment of acid/base equilibriums). It also changed scientific thinking by demonstrating the importance of the solvent in acid/base reactions, and showed that the acidic and basic salts were due to simple acid or base reactions.Outline the Br?nsted-Lowry theory of acids and basesThe Arrhenius definition of acids/bases was limited for the following reasons:Did not explain why metallic oxides and carbonates were bases, despite their ability to neutralise acidsDid not explain acidic salts (e.g. zinc chloride) or basic salts (e.g. sodium sulfide)Did not explain why anhydrous neutralisation reactions would occur (e.g. hydrochloric acid dissolved in benzene and ammonia reacting to produce ammonium chloride)Did not explain why some substances can act as both an acid and a baseNEW UNDERSTANDINGS IN ACIDS AND BASES LED TO THE DEVELOPMENT OF THE BR?NSTED-LOWRY THEORYThe Br?nsted-Lowry theory of acids states:Acids are proton (H+) DONORSBases are proton (H+) ACCEPTORSA neutralisation reaction is a proton-transfer reactionThe Br?nsted-Lowry theory applies to non-aqueous environments (e.g. non-aqueous solvents, gas-phase reactions) => acids and bases can be solids, gases, or anhydrous solutionsEXAMPLEUnder the Arrhenius definition of acids, the above reaction would not be classed as an acid/base reaction, as there are no free H+ ions presentBut under Br?nsted-Lowry theory, the above reaction is classed as an acid/base reaction, as HCl had donated protons, whilst NH3 has accepted protonsThus HCl has acted as an acid in an anhydrous environment, and NH3 has acted as a base => IMPROVEMENT ON THE PREVIOUS DEFINITION ON ACIDS AND BASESThe Br?nsted-Lowry theory also has implications for the role of the solventWhen an acid is ionised in water, water acts as a Br?nsted-Lowry base, as it accepts protons. For example…When a base is ionised in water to produce an alkali solution, water acts as a Br?nsted-Lowry acid, as it donates protonsIn the self-ionisation of water, water acts as both an acid and a baseDescribe the relationship between an acid and its conjugate base and a base and its conjugate acidUnder Br?nsted-Lowry theory, an acid donates a proton in a reaction, resulting in the acid species becoming deprotonatedThe deprotonated acid can then accept protons, hence it can act as a baseSimilarly, a protonated base species can donate protons, hence act like an acidThis concept is known as conjugate acids and basesA CONJUGATE BASE is the original acid with a hydrogen ion removedA CONJUGATE ACID is the original acid with a hydrogen ion addedConjugate means ‘linked with’ => under Br?nsted-Lowry theory, each acid has a conjugate base, and each base has a conjugate acidBy convention, conjugate acids and conjugate bases are on the right-hand side of an acid-base equationEXAMPLESH2SO4 is the acid; its conjugate base is HSO4- (H2SO4 minus a proton)HCl is the acid; its conjugate base is Cl- (HCl minus a proton)H2O is the base; its conjugate acid is H3O+ (H2O plus a proton)NH3 is the base; its conjugate acid is NH4+ (NH3 plus a proton)H2O is the acid; its conjugate base is OH- (H2O minus a proton)RELATIVE STRENGTHSGenerally…The stronger the acid, the weaker its conjugate baseThe stronger the base, the weaker its conjugate baseThe relationship between an acid and its conjugate base can be considered in the equilibrium belowIf HA is a strong acid, then the equilibrium will strongly favour the forward reactionThus the reverse reaction will only very weakly proceed (i.e. A- will not accept protons well) => A- is thus a weak conjugate baseSimilarly for a base and its conjugate acid…If A- is a strong base, the equilibrium favours the forward reactionThus reverse reaction does not proceed well (i.e. HA does not deprotonate well) => HA is thus a weak conjugate base.The diagram below illustrates the relationship between the strengths of an acid or base and its conjugateIdentify conjugate acid/base pairsConjugate acid/base pairs can be easily identified from a chemical reactionRemember…A conjugate base is the original acid minus a protonA conjugate acid is the original base plus a protonConjugate acid/base pairs can be identified by the following method:Identify the proton donor and proton acceptor in the forward reactionClassifying accordingly the reactants as B-L acids and basesIdentify the conjugate acid and baseRecord the acid/base pairEXAMPLESAcid/base pairs are H2CO3/HCO3- and HS-/S2-Acid/base pairs are H2O/OH- and H3O+/H2OBelow is a table of common acid/base pairsIdentify a range of salts which form acidic, basic and neutral solutions and explain their acidic, neutral or basic natureDefinitions…A salt is an ionic compound produced by a neutralisation reaction, consisting of an anion and a cationThe hydrolysis of a salt is the reaction of a salt and water, producing a pH changeUnder Br?nsted-Lowry theory, any ion can act as an acid or a base, and react with waterConsequently, many salts can form acidic or basic solutions in waterThe pH of a salt in solution depends on the relative strengths of the anion and cation as either acids or basesThe acidity or alkalinity of a salt can be predicted by considering the reactants of the neutralisation reaction producing the saltACIDIC SALTSOccurs when the cation’s acidity is greater than the anion’s alkalinityThe cation reacts with water to a greater extent than the anion, thus producing more H3O+ than OH- => the pH of the solution decreases, thus the salt undergoes hydrolysisAcidic salts result when a strong acid reacts with a weak baseThe cation from the base acts as a weak acid by reacting with water to produce H3O+, whilst the anion reacts to a much lesser extent, if at allEXAMPLESNH4Cl (ammonium chloride) is an acidic salt, as the NH4+ ion acts as a weak acid, whilst the Cl- does not reactZnSO4 (zinc sulfate) results from the reaction of zinc hydroxide and sulfuric acid. The hydrated zinc ions act as a weak acid when dissolved in water, whilst the sulfate ions do not react.BASIC SALTSOccurs when the anion’s acidity is greater than the cation’s alkalinityThe anion reacts with water to a greater extent than the cation, thus producing more OH- than H3O+ => the pH of the solution increases, thus the salt undergoes hydrolysisBasic salts result when a weak acid reacts with a strong baseThe anion from the acid acts as a weak base by reacting with water to produce OH-, whilst the cation reacts to a much lesser extent, if at allEXAMPLESKF (Potassium fluoride) is produced from potassium hydroxide and hydrofluoric acid. The fluoride ion acts as a weak base in water, whilst the potassium ion does not reactNaCH3COO (sodium acetate) is produced from sodium hydroxide and acetic acid. The acetate ion acts a weak base in water, whilst the sodium ion does not reactNEUTRAL SALTSThe pH of a neutral salt in water is 7Neutral salts can be produced from the reaction of…A strong base and a strong acid (e.g. NaCl, Ba(NO3)2) => neither the cation or anion of the salt reaction with water sufficiently to alter its pH => the ions of such salts do not hydrolysisThe anions of strong acids are all the halide ions (except F-), and strong oxyanions such as NO3- and ClO4-The cations of strong bases are those from Group 1 and Ca2+, Sr2+, and Ba2+ from Group 2A weak base and a weak acid (e.g. NH4CH3COO) => the cation and anion react with water to approximately same extent, thus causing no net change in pH => the ions of such salts hydrolysis to the same extentNOTE: The pH of a weak acid/weak base salt MUST be tested to verify its neutrality, because the two ions must be hydrolysis to the same extent => any imbalance will cause the pH of the salt solution to changeIn summary…A strong acid and a strong base will produce a neutral saltA weak acid and a weak base will generally produce a neutral saltA strong acid and a weak base will produce an acidic saltA weak acid and a strong base will produce a basic saltIdentify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutionsAn amphiprotic substance is one that can act as BOTH a proton donor and a proton acceptor, i.e. can act as both a Br?nsted-Lowry acid and baseThey act as an acid or base depending on the conditions of the reaction, usually the acidity of other substances in the reactionEXAMPLESWater (H2O) is amphiprotic, as seen in its self-ionisationIf a substance (e.g. HCl) has a greater tendency to lose protons than water, then water acts like a base, producing an acidic solutionIf a substance (e.g. NH3) has a lesser tendency to lose protons than water, then water acts like an acid, producing a basic solutionThe hydrogen carbonate ion (HCO3-) is also amphiproticWhen placed in an alkaline solution, HCO3- acts as an acidWhen placed in an acidic solution, HCO3- acts as a baseIn the above reaction, the H2CO3 decomposes to H2O and CO2, thus carbon dioxide gas evolvesThe hydrogen sulfite ion (HSO3-) acts in a similar way to HCO3-, releasing chocking sulfur dioxide gas when it reacts with a hydroxide ionHydrogen phosphate (HPO4-) also acts similarly, except that it forms a slightly basic solution when added to waterIdentify neutralisation as a proton transfer reaction which is exothermicNeutralisation reactions are reactions between acids and basesUnder Br?nsted-Lowry theory, and acid is a proton-donor and a bases is a proton-acceptor, thus in a reaction of acids and bases, a proton is transferred from and acid to a baseEXAMPLEThe reaction of hydrochloric acid and sodium hydroxide is a neutralisation reaction:Whilst a proton transfer may not be immediately evident, by considering the net ionic equation…The above equation is the same for all Arrhenius neutralisation reactions (i.e. hydronium ion and hydroxide ions)We can see that a proton has been transferred from the hydronium ion to the hydroxide ion, thus it is a proton transfer reactionGenerally, all neutralisations are proton transfer reactions (see below for more examples)Also note that the above reaction is exothermic => the change in enthalpy for all neutralisations is around -57kJmol-1, depending on the strength of the acid or base (a weak acid/base won’t fully ionise in solution)OTHER EXAMPLESA proton is transferred from the hydronium ion to the ammonium moleculeNote that the enthalpy change is less than the reaction above => acetic acid is a weak acid, so does not ionise 100% in solutionAnalyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spillsIt is important to safely clean up any spills involving acids and bases, because they are corrosive, thus can damage equipment or the environment, and pose a serious safety riskIf an acid or base is spilt on skin, first aid must be administered immediately by washing the skin with copious volumes of waterIf an acid or base is spilt on the floor, the following procedure should be followed: The area first needs to be isolated to minimise further damage and the possibility of inhaling toxic fumes => this involves spreading sand around the spill site to minimise the spill from spreadingThe spill should then be cleaned up first by neutralising the acid/base, then once safe, the spill should be cleaned up with paper towels and disposedThe preferred chemicals used in neutralising spilt acids/bases are stable (safely handled), solid in a powdered form (easily transported and stored), cheap, and amphiprotic (so it is weak, and can clean up both acid and base spills)The substance most commonly used to spill up laboratory acid/base spills is sodium hydrogen carbonate (NaHCO3)The HCO3- ion is amphiprotic, cheap, and readily availableAdditionally, the neutralisation reaction with an acid produces CO2 gas, thus the fizzing can be used to monitor the progress of the neutralisation reactionThus NaHCO3- has been very effective as a safety measure or to minimise damage in accidents or spillsIt is very important that the chemical used to neutralise spills is a weak acid/base, as neutralisation reactions are exothermic, so the quantity of heat produce needs to be minimal.In addition, if excess chemical is used, the spill will be weak, thus does not pose further safety risks => if excess of a strong acid/base is added, the spill will become dangerous again.NOTE: As proton transfer reactions are exothermic, special care must be taken when diluting acids => a small volume of concentrated acid should be added to a large volume of water, so that the water can dissipate the generated heat. If water is added to acid, the acid’s temperature can quickly reach boiling point, which could be ejected from a beaker and pose a very serious safety issue.Qualitatively describe the effect of buffers with reference to a specific example in a natural systemBuffers are solutions that resist changes in pH when small quantities of acids or bases are added to themA buffer solution contains comparable quantities of a weak acid and its conjugate base, which exist in the following equilibriumBy considering Le Chatelier’s principle, we can see that the equilibrium can resist changes in pHIf an acid is added, the concentration of H3O+ is added, then the concentration of H3O+ ions increases. The equilibrium resists the change by favouring the reverse reaction to reduce the concentration of H3O+, thereby maintaining the original pH.If a base is added, it reacts with the H3O+ ions to produce water, thus increasing the concentration of water and decreasing the concentration of H3O+. The equilibrium resists the change by favouring the forward reaction, thus the pH of the solution is maintained.As the buffer contains comparable quantities of the weak acid and its conjugate base, the system can resist large changes in pH.The action of buffers can be observed by considering the titration curve of a weak acid and a strong base, and the corresponding titration curve of a weak base and a strong acid, as the buffer zone lies in the flat region of the curveEXAMPLE OF A NATURAL BUFFERSystem: The carbonic acid system in freshwater lakes and riversThe carbonic acid molecules are formed when carbon dioxide dissolves directly in the water, or from the dissolved carbon dioxide in rain waterThe hydrogen carbonate ions are leeched out of rocks and minerals in the lakeThe additional HCO3- pushes the above equilibrium to the left, thus raising the pH to between 6.5 and 7.5.The buffer resists the addition of acids (such as acid rain) or bases, thus a neutral pH is maintained.For example, the absence of the above buffer in Scandinavian lakes due to the absence of carbonate rocks (thus no additional HCO3-) led to these lakes being the first to detect a falling pH from acid rainThe natural carbonic acid buffer in lakes and rivers is important as marine life requires a neutral pH to live => for example, fish eggs are destroyed if water is acidic.Other possible buffers are shown below:Describe the correct technique for conducting titrations and preparation of standard solutionsBACKGROUND THEORYTitration, also known as volumetric analysis, is a chemical technique used to determine the concentration of an unknown solution through experimental first-hand dataIt is a technique of quantitative chemical analysis, which requires high degrees of accuracyThe concentration of an acid or a base can be determined by titration, using a neutralisation reactionA certain volume of and acid/base of unknown concentration is slowly reacted with another solution of known concentration and volume, until the endpoint is reachedBy measuring the volume of the acid/base that was required to reach the endpoint, the concentration of the unknown solution can be calculatedIt is important to distinguish between the equivalence point and endpoint of an acid-base titrationAt the equivalence point, equal volumes of the two reactants have been reacted to cause the complete consumption of both reactions (which depends the stoichiometry of the reaction)At the endpoint, there is a permanent colour change in the indicatorFor titration, the equivalence point and endpoint should coincideEXAMPLETitration of sulfuric acid against a standardised 0.100molL-1 sodium hydrogen carbonateFirst write the relevant neutralisation reaction:Next determine the molar ratio of reactantsFor every 1 mole of acid, there are 2 moles of basePerform the titration with a methyl orange indicator, and determine the volume of NaHCO3 required to react with H2SO4The endpoint was reached after 25.00x10-3L of NaHCO3 was titrated against 28.35x10-3L of H2SO4Ideally, repeat the titration as many times as possible, discount outliers, and average the reliable resultsCalculate the unknown concentrationThere are many rules to follow to achieve a highly accurate titration, including the preparation of a standard solution, indicator choice, and the titration itselfPREPARATION OF STANDARD SOLUTIONA standard solution (also known as titrant) is a solution of accurately known concentrationSelecting an appropriate standard solutionWe cannot accurately determine the concentration of many common acids or bases (e.g. HCl, H2SO4, NaOH), as they’re concentration changes over time for various reasonsWe can accurately determine the concentration of such acids or bases through a titration against a primary standard (a solution prepared from dissolution, which has an accurately known concentration)To prepare a primary standard, we first need to use a suitable substance that will not cause error in the titrationThe criteria for choosing a primary standard are…High level of purityAccurately known compositionHigh molecular (to reduce errors associated with weighing)Free from moistureStable and unaffected by air in weightingReadily soluble in pure (distilled) waterReacts instantaneously and completelyNon-hydroscopic (does not absorb water from surroundings) and non-efflorescent (does not release water to surroundings)The table below lists some suitable substance for use as a primary standardAcid standardBase standardPotassium hydrogen phthalate (KHC8H4O4)Sodium carbonate (Na2CO3)Benzoic acid (C7H5O2)Sodium hydrogen carbonate (NaHCO3)Oxalic acid (H2C2O4.2H2O)Borax (Na2B4O7.10H2O)Other common acids/bases are unsuitable, as they don’t fit the above criteriaHCl is efflorescentH2SO4 is severely hydroscopicNaOH reacts with gases in the airNa2CO3.nH2O contains moisturePreparing the standard solutionTo prepare a primary standard, the procedure must be followed with high accuracy and precisionEnsure that all glassware is clean, and rinsed with distilled unless indicated otherwiseIn the method below, analytical-grade anhydrous sodium carbonate (Na2CO3) is used (99.9% purity after roasting for 30 mins at 150°-180°C in a drying oven, then cooled to form crystals in desiccator)Calculate the mass of primary standard required to achieve a desired molarity. The result is a guide only; it is the final concentration that it is important.Weigh the anhydrous sodium carbonate (1.325g) in a dry 50mL beaker on an electronic scale. Remember to first place the beaker on the scale, and then zero it. Use a plastic wash bottle to add a little distilled water to the beaker. Stir the solution with a fine, short glass rod to dissolve the sodium carbonate, making sure that the glass rod does not touch foreign substances. Ensure all solid has dissolved.Transfer the sodium carbonate solution from the beaker into a clean 250mL volumetric flask by placing a small, clean glass funnel over the neck of the flask. Rinse the beaker, glass rod, and funnel into the volumetric flask with the wash bottle to ensure that all the sodium carbonate has been transferred.Fill the flask with distilled water until it reaches about 1cm below the graduation mark. Continue to fill the flask up to the graduation mark drop-by-drop using a dropper or Pasteur pipette until the bottom of the meniscus is aligned with the graduation mark. Avoid parallax error (error can be minimised by marking a piece of paper with a black line and placing it at the graduation mark to aid locating the graduation mark.Stopper the flask, and invert the flask 10 times to ensure a homogenous solution. Label the flask as 0.0500molL-1 Na2CO3.TITRATIONFirst we’ll define relevant terms and equipmentAliquot – A known volume of a liquidTitrant – A solution that is added from the burette (typically standard solution)Burette – A piece of cylindrical glassware, held vertically, with volumetric divisions on its full length and a precision tap (stopcock) on the bottom. It is used to dispense precise volumes of the liquid reagent, with the volume dispensed being the difference between volumetric divisions. Burettes are very precise, with accuracy up to ± 0.05mL.Pipette – A glass tube used to transfer precise volumes of liquid reagants. They are usually dsigned to trnasfer a set volume, such as 25mL. The reagant is drawn up the pipette using a pipette filler (e.g. a rubber bulb)Volumetric flask – A glass flask with a long neck with a graduation mark to measure specific volumes. Volumetric flasks are used to prepare and hold standard solutionsConical flask – Typically used to hold the reactants during titration. Its shape prevents the reactants from spilling as they are swirled together.The diagram below shows the above glassware. The flask with the yellow stopper is the volumetric flask, and then clockwise is the burette, pipette and bulb, and the conical flask.Next we must choose an appropriate indicatorRecall that not all salts are neutralA strong acid and a weak base produce an acidic saltA weak acid and a strong base produce a basic saltHence at the equivalence point, the pH of the solution may not be neutralTo ensure the equivalence point and endpoint align, we must choose an appropriate indicatorThe table below summarises the commonly used indicatorsIndicatorInitial ColourRangeFinal colourDiagramMethyl orangeRed3.1-4.4YellowBromothymol blueYellow6.0-7.6BluePhenolphthaleinColourless8.2-10.0PinkConsider the titration curve for a strong acid/ strong base titrationThere is a steep increase in pH from 3 to 11 within a range of about a drop. Thus any of the above indicators could be used, as they would all register the pH change. The equivalence point lies at around a pH of 7, so use bromothymol blue if available, but methyl orange or phenolphthalein can be used too.Strong acid/weak baseThe equivalence point lies at around 5.3 for HCl/NH3, so the most suitable indicator would be methyl orange. Bromothymol blue could also be use, but phenolphthalein cannot be used, as its pH transition range is outside the steep increase in pH shown aboveWeak acid/strong baseThe equivalence point lies at around 8.7 for NaOH/CH3COOH, so phenolphthalein is the most suitable indicator, but bromothymol blue could also be used too. Methyl orange cannot be used, as its pH transition range lies outside the steep increase in pH shown above.Weak acid/weak baseAs shown above, there is no sharp change in pH for a weak acid/weak base reaction, thus these reactions should be avoided in titrations.pH probes and data loggers can also be used in titrations to determine the equivalence point by analysing the titration curves similar to those above.Now we must rinse the glassware appropriatelyThe volumetric flask (including stopper) should be rinsed thoroughly with distilled water. The flask can be left wet, as water is going to be added to it later.The conical flask should be rinsed thoroughly with distilled water, and left wet for similar reasonsThe burette should first be rinsed with distilled water by filling the burette, then opening the tap to clean the tip. Fill the burette again with distilled water, swirl the water to rinse the sides, then poured out the top. Repeat this THREE times, and then do it once more with the solution it is going to contain.The pipette should be rinsed three times with distilled water, then once with the solution it is going to contain.Now we can set up the titrationUsing a funnel, pour the titrant into the burette until above a suitable marking (e.g. 30mL) to facilitate measurements. Open the tap slightly until the meniscus sits just above the mark. Holding a white card with a black line makes the meniscus clearer.Using the pipette, transfer a fixed volume (typically 25mL) of the other reagent into the conical flask. UNDER NO CIRCUMSTANCES SHOULD THE FINAL DROP OF THE PIPETTE BE EXPELLED => pipettes are designed with this feature in mind.Add three drops of the appropriate indicator to the solution in the conical flaskSet up the burette above the conical flask using a retort stand and clamp, and place the conical flask on a white tile to make the indicator colour clearNow we can conduct the titrationFirst conduct a trial run, where the volume required to reach the endpoint is roughly determined by running the burette at a steady speed until the endpoint is reached. Keep this flask so that the final colour can be compared to in subsequent titrationsWith one hand, open the tap so that the titrant runs steadily into the conical flask. With your other hand, swirl the flask to homogenise the solution.Once within 10mL of the expected volume, close the stopcock. Wash down the conical flask with a wash bottle as the solution may have splashed. Open the stopcock slightly, so that the titrant slowly drips into the conical flask.Stop as soon as the endpoint is reached, and record the final volume reading on the burette. If unsure if the indicator has changed colour, record the volume, then add drop-by-drop, recording each reading until the endpoint is definitely reached. The diagram below shows methyl orange just before, at, and just after the endpoint has been reached.Repeat the titration as many times as possible, keeping each flask for colour comparison. For increased accuracy, add partial drops from the burette by quickly turning the stopcock on and off so that a drop hangs from the tip, then resting the drop on the conical flask.We now calculate the unknown volumeCompare the readings from each titration => the volume of titrant used is calculated by subtracting the final reading of the burette from the initial reading.Discount any outliers, and ideally average the three closest results.Use the average value in subsequent calculations. The following formula can be used if the molar ratio is 1:1Choose equipment and perform a first-hand investigation to identify the pH of a range of saltsMETHODThe pH of various salts was determined by using a pH meter, and then confirmed by using universal indicatorThe salt solutions tested were sodium carbonate, sodium hydrogen carbonate, sodium chloride, ammonium chloride, and ammonium sulfate.All salt solutions were 0.1M, and 20mL of each salt was testedSAFETY: Ammonium chloride releases toxic fumes if heated => keep away from flame. Ammonium chloride, sodium acetate, and sodium carbonate are slightly toxic if ingested and are skin irritants => wear eye protection, use a dropper when transferring solutionsRESULTSSalt (0.1M)pHColourNa2CO311-12Dark blueNaHCO38-9BlueNaCl6.5-7.5GreenNH4Cl5-6Orange(NH4)2SO44-5Light pinkVALIDITY/RELIABILITY/ACCURACYThe results confirmed expected resultsThe use of technology (pH probes) provided more accurate quantitative analysis, and more reliable resultsThe results may have been inaccurate due to the inaccuracies of the concentration of solutions prepared in the school lab.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 basesMETHODSee above for extensive notes on preparing standard solutions and conducting titrationsA standard solution of sodium carbonate (Na2CO3) was prepared, and was titrated against an HCl solution to determine the concentration of HClSAFETY: The concentration of HCl is unknown, thus it could be of high concentration => clean up spills immediately, wear safety glasses, use a pipette to transfer HCl solutionsRESULTSThe following reaction occurred in the titrationm(Na2CO3)=14.56g[Na2CO3] = 0.549molL-1 (Na2CO3 was dissolved in a volumetric flask)22.3mL of Na2CO3 was required for 25mL of HCl solution to reach endpoint (methyl orange indicator was used)Thus the concentration of HCl solution was 0.98molL-1VALIDITY/RELIABILITY/ACCURACYSee titration notes above for a detailed account of conducting valid titrationsAll glassware was rinsed appropriately prior to useElectronic scales were used for accuracy, providing mass up to 2 decimal placesMultiple titration runs were conducted, the first one discounted as a rough guide, and any outlier results were discounted before taking an average => increases reliabilityPerform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologiesMETHODThe concentration of household vinegar was determined through titrations involving computer-based technologiesSee above for extensive notes on titrationA primary standard of potassium hydrogen phthalate was prepared accuratelyThe potassium hydrogen phthalate was titrated against a prepared solution of NaOH to produce a secondary standard solutionThe NaOH secondary standard was titrated against the vinegar solution to determine the concentration of vinegar.A pH probe and data logger were used in both titrations to determine the equivalence point, and the pH probes were calibrated using pH 4 and 7 buffersRESULTSThe data logger produced titration curves similar to those in the titration notes above (weak acid/strong base), from which the equivalence point could be determinedThe concentration of vinegar was calculated to be 0.79M (see above for notes on calculations)JUSTIFICAITON OF METHODA secondary standard of NaOH was produced to increase the reliability of results, as weak acid/ weak base titrations provide unreliable resultsAll glassware was rinsed appropriately prior to useTrial runs were conducted for both titrations, and average results were taken, discounting outliersEsterification is a naturally occurring process which can be performed in the laboratoryDescribe the difference between the alkanol and alkanoic acid functional group in carbon compoundsRecall that a functional group is a group of atoms responsible for particular characteristics of a compoundAlkanols contain the hydroxyl group (-OH) attached to a carbon atomFor example, below is the structural formula of ethanolAlkanoic acids contain the carboxyl group (-COOH) attached to a carbon atomOne of the oxygen molecules is double-bonded to the central carbon, whilst the OH group is single-bonded to the carbon (shown below)NOTE: The –OH group in the carboxyl group is NOT a hydroxyl groupThe hydrogen atom of the carboxyl group can ionise in solution, hence the group is acidicFor example, below is the condensed structural formula of ethanoic acidExplain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structuresThe melting/boiling point of a molecular substance reflects the strength of the intermolecular bonds within a given substanceThe stronger the intermolecular bonds, the more energy is required to weaken (melting) or break (boiling) the bonds, thus the melting/boiling point is higherALKANOLSAlkanols contain the hydroxyl (-OH) group an C-O bonds, both of which are polarThis allows for strong hydrogen bonds to be formed between alkanol molecules, in addition to dispersion forcesAlkanols thus have a higher melting/boiling point than many carbon compounds of similar mass, such as alkanesALKANOIC ACIDSAlkanoic acids have a higher molecular mass than their corresponding alkanol, since the carboxyl group has a higher mass than the hydroxyl groupThus alkanoic acids have stronger dispersion forces between molecules than the corresponding alkanolAlkanoic acids also have more extensive hydrogen bonding than alkanolsThe C=O bond in alkanoic acids is also polar in addition to the C-O and O-H bondsThis allows two hydrogen bonds to be formed per molecule, compared to 1 hydrogen bond per alkanol moleculeThus alkanoic acids have stronger intermolecular forces between molecules than alkanols of similar molecular massAs a result, alkanoic acids have higher melting/boiling points than alkanols of similar size, which in turn have higher melting/boiling points than similarly-sized alkanes and alkenesIn summary, alkanoic acid > alkanol > ester > alkaneIdentify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterificationGenerally, esterification is the reaction between an acid and an alcohol to produce an esterIn the HSC course however, we only consider the reactions between straight-chained alkanoic acids, and straight-chained primary alkanols => these produce carboxylate estersNote that esterification is NOT an acid-base reaction => the –OH group of the acid reacts with the hydrogen atom of the hydroxyl group of the alkanol.Carboxylate esters contain the -COOC- structural unit:Generally…The equilibrium at room temperature lies much to the left at room temperatureSee below for an example of esterification, and naming of alkyl alkanoatesIdentify the IUPAC nomenclature for describing esters produced by reactions of straight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8First let’s review the IUPAC nomenclature for describing alkanoic acids and alkanolsALKANOIC ACIDSCount the number of carbons, take the parent alkane name, drop the final –e, and add on the suffix ‘-oic acid’Since the –COOH contains only one free carbon valence, no number is needed to signify its locationEXAMPLEThere are three carbon atoms, thus the parent alkane is propane. By dropping the suffix –e and adding –oic acid, we have ropanoic acidALKANOLSCount the number of carbons, take the parent alkane name, drop the final –e, and add the suffix ‘-ol’Remember to number the carbon where the hydroxyl group is by counting from the end that gives the smallest numberEXAMPLEThere are three carbons, thus the parent alkane is propane. By dropping the ‘-e’, adding –ol, and number the carbon atom where –OH lies (2), we have 2-proponolESTERSEsters from alkanols and alkanoic acids are named as alkyl alkanoatesThe alkyl comes from the alkanol, and the alkanoate comes from the alkanoic acid => hence the alkanol is named first, then the alkanoic acidEsters can be named by considering the reaction that produces the reaction, or the given structural formula of an esterIf the alkanol and alkanoic acid are given, name the ester by the following procedure:Replace the –anol of the alkanol with –yl (i.e. the alkyl is named according to the corresponding alkanol)Replace the ‘-oic acid’ of the alkanoic acid with ‘-oate’Place the two words togetherEXAMPLEFrom the 1-proponol we get ‘propyl’, and from the pentanoic acid we get ‘pentanoate’, thus the name of the resulting ester is propyl pentanoateNOTE: The preferred IUPAC name for methanoic acid and ethanoic acid are formic acid and acetic acid respectively. Thus methanoate may sometimes be named formate, and ethanoate may sometimes be named acetate.If the structural formula is given, follow the following method:Split the ester across the C-O-C bond => the side with the C=O bond is the alkanoate, the side without is the alkyl groupCount the number of carbons in the alkyl and alkanoate groups respectively, and name similarly to aboveEXAMPLEWe can see that the alkyl group is to the right of the C-O-C bond, whilst the alkanoate is to the leftThere are three carbons in the alkyl group, hence it is named propylThere are five carbons in the alkanoate group, hence it is named pentanoateThus the name of the ester is propyl pentanoateNOTE: As the dot point only asks for the reactions involving straight-chained primary alkanols, numbering is not important for esters encountered in the HSC courseDescribe the purpose of using acid in esterification for catalysisConcentrated sulfuric acid is often added as a catalyst in esterification (1-3mL of concentrated acid)As a catalyst, concentrated sulfuric acid is not consumed in the reaction, but allows the equilibrium point to be reached faster (esterification is naturally a slow reaction)Concentrated sulfuric acid is added because…It speeds up the rate of reaction by lowering the activation energy of the forward reaction, thus the equilibrium point is reached fasterSulfuric acid is a dehydrating agents, thus the equilibrium shifts to the right, and this increases the yield of esterExplain the need for refluxing during esterificationRefluxing involves heating up a mixture with a cooling condenser placed above, so any volatile reactants and products are condensed and returned to the reaction mixture => this reduces loss of substanceA refluxing apparatus consists of a water condenser mounted on top of a reaction vessel (normally a round-bottomed flask). The circulating cold water cools any vapours, causing them to condense and run back to the reaction vessel.The refluxing apparatus is open to the atmosphere, so there is no pressure build-up inside the apparatusHeating is desirable in esterification because…The higher temperatures increase the rate of reaction, thus the equilibrium is reached fasterEsterification is endothermic in the forward reaction, so a higher yield of ester is obtained if the mixture is heatedThe reactants and products of esterification are volatile however, so heating the reaction mixture causes the mixture to evaporateRefluxing causes any vapours from the mixture to condense and run back to the reaction vessel, thus allowing the reaction to proceed at higher temperatures without fear of losing reactionsThe vapours produced are also flammable, so refluxing increases the safety of esterificationThe reaction vessel is generally heated with a water bath or hot plate for safety reasons by reducing the risk of fire (as opposed to a Bunsen burner)Boiling chips (crushed ceramic) are also added to the reaction vessel => these prevent the risks of sudden superheating and explosive ejection of vapours by providing a large surface area on which vaporisation can occurOutline some examples of the occurrence, production and uses of estersNATURAL OCCURRENCEEsters occur naturally as flavouring agents or as scents, particularly in flowers and fruitsUsually it is a mixture of esters that creates a characteristic smell or taste in nature, though it can sometimes be due to the presence of a single esterFor example, the characteristic scent of pineapple is due to ethyl butanoateFats and oils are also triesters of glycerol and various long-chained carboxylic acidsINDUSTRIAL PRODUCTIONEsters are commonly produced as synthetic food flavours and perfumes to mimic those found in natureFor example, methyl butanoate is often used as apple flavouringShort esters are also used as solvents, such as ethyl acetate (ethanoate) as nail polish removerFor instance, 1-pentyl acetate is used as an extraction solvent in pharmaceuticals and in cleaning fluidsHigh molecular mass esters (e.g. dialkyl phthalates) are used as plasticisers to soften hard plastics such as PVCAcetyl salicylic acid is a major active substance in aspirinProcess information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmeticsMany esters are synthetically produced to mimic natural flavours and scentsSynthetic esters are often cheaper than natural extracts, and represent little health hazard provided they contain only substances that occur in natural flavoursEsters can be added to processed foods to produce desired flavours (see below for common flavours and esters)For example, iso-pentyl acetate is used in banana-flavoured milkMany cosmetics, particularly perfumes, comprise almost exclusively a mixture of esters in a solvent (often ethanol) to produce desired scentsEsters are also added to soaps, hand-lotions, or other cosmetics to give a pleasant smellEsterFlavourMethyl butanoateAppleEthyl butanoatePineappleButyl ethanoateRaspberryPentyl ethanoatePear1-octyl acetateOrangeIdentify data, plan, select equipment and perform a first-hand investigation to prepare an ester using refluxMETHODThe following apparatus was set upEthanol and ethanoic acid were used as reactants to produce ethyl ethanoate40mL of ethanol, 30mL of ethanoic acid, and 10 drops of 18M sulfuric acid (teacher use only) were added to the round-bottomed flaskBoiling chips were also added to the flaskThe condenser attachment of the reflux apparatus was attached to the flask, and secured onto a retort stand. Cold water was passed into the condenser.The hot plate was used to bring the reaction mixture to a gentle boil (~100°C), and was left for 30 minsHeating was stopped, and the entire apparatus was left to cool, with water still flowingAfter the apparatus had been carefully disassembled, the refluxed contents were poured into a beaker, 70mL of water was added, and the mixture was stirred. Any distinctive features were observedTo obtain a relatively pure solution of ethyl ethanoate, pour the mixture into a separating funnel, wash with 50mL of distilled water, and drain off the lower aqueous layer (ester is immiscible with the reaction mixture as it is not water soluble, and is also less dense). Repeat several times, then add 15mL of saturated NaHCO3 solution and 35mL of distilled water, shake the mixture, and drain off the aqueous layerSAFETYConcentrated sulfuric acid is highly corrosive => use a drop bottle, clean up spills immediately, have the teacher transfer contentsThe reactants are flammable => use the condenser to condense flammable vapours, do not use an open flame (i.e. use a water bath or a hot plate)Gas pressure can build up => make sure that the top of the condenser is openRESULTSThe final solution had a distinctive sweet smell => this indicates that an ester has been producedJUSTIFICATION OF METHODBoiling chips were added for slow and gentle heatingA hot plate allowed for gentle heating, and was more safe to use than an open flameRefluxing was conducted using a condenser to minimise the loss of reactants, and allow for the reaction mixture to be heatedConcentrated sulfuric acid was added as a catalyst for the reactionThe produced mixture was separated to allow us to obtain relatively pure ester, which aided in making accurate observations. Multiple washings and adding sodium hydrogen carbonate increased the purity of the final solution. ................
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