HSC 9.3 The Acidic Environment

1.      Indicators were first identified with the observation that the colour of some flowers depends on soil composition

Outcome: Classify common substances as acidic, basic or neutral.

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

Outcome: 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 colour.

                 Identify data and choose resources to gather information about the colour changes of a range of indicators.

Outcome: Perform a first-hand investigation to prepare and test a natural indicator.

Take a sample of red cabbage, grind it with pestle and mortar with some distilled water.  Filter this solution which becomes a natural indicator.  Test it with HCl and NaOH to determine its colour changes in acid and basic solutions.

HCl – deep cherry red

NaOH – green, then brilliant yellow

Outcome: Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity

Testing the pH of soil to determine land’s ability to support certain crops, check the environmental effects of industrial actions.  Check the pH suitability of streams, rivers, pools for water safety, both consumption and swimming.

1.      While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur.  The concentrations of these acidic oxides have been increasing since the Industrial Revolution.

Outcome: Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids.

Oxides of non-metals tend to be acidic.  The more non-metallic the element, the more acidic its oxide is.  They must be dissolved in water to contain acidic qualities.

            SO_3(g) + H₂O_(l)              H₂SO_4(aq)

Outcome: Analyse the position of these non-metals in the Periodic Table and generalise about the relationship between position of elements in the Periodic Table and acidity/basicity of oxides.

Groups 3 – 8 form acidic oxides.  They dissolve in water to form acids, and also react with hydroxide ions in solution.

            Groups 1, 2, H and He form basic oxides.  When dissolved in water, they form hydroxide ions which cause the basicity of the solution.

            Elements in between, the transition metals, are amphoteric oxides, which neutralises both HCl and NaOH in solution.  They act as acids or bases depending on conditions.

For example, Period 3

The strengths of the acid increases as it moves towards the non-metals.  This corresponds to a steady fall in the electronegativity difference between O₂ and elements.

Outcome: Define Le Chatelier’s principle

If a system is at equilibrium and a change is made that upsets the equilibrium, then the system alters in such a way as to counteract the change and a new equilibrium is established.

Henry Loius le Chatelier proposed this general principal in 1885 to predict the effect of changes to systems at equilibrium.

Outcome: Identify factors which can affect the equilibrium in a reversible reaction.

If there is a high yield of products, the equilibrium lies to the right.

If reactants dominate the equilibrium mixture, the equilibrium lies to the left.

Increases in the concentration of materials on the left will increase the forward rate, and vice versa.

Increases in gas pressure (effective concentration of particles) increase the forward rate.  The removal of some products decreases the pressure and decreases the reverse rate.

Generally, increasing the temperature shifts an endothermic equilibrium to the right (products) and shifts an exothermic equilibrium to the left (reactants).

Increases in the surface area achieves a faster equilibrium but the position of the equilibrium remains unchanged.  A catalyst also allows for a faster rate of equilibrium but does not affect the position of the equilibrium.

Outcome: Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and relate this to Le Chatelier’s principle.

High concentrations of CO₂ drive the equilibrium to the right, increasing the acidity of the water as the system moves to reduce CO₂ concentration.  Soda water contains CO₂ gas under pressure and this equilibrium is maintained in a sealed/closed environment.  CO₂ is lost in an open system and the equilibrium shifts to the left to make more CO₂.  This is an exothermic reaction, so the heating of the solution shifts the equilibrium to the left, opposing the creation of more heat.

            Solubility of CO₂ decreases as H₂O temperature increases, equilibrium shifts left.  Solubility of CO₂ increases with pressure because there is a higher concentration of CO₂, equilibrium shifts right.

Outcome: Identify natural and industrial sources of sulfur dioxide and oxides of sulfur

Iron sulfide sediments have been gradually deposited in low oxygen environments, such as salt marshes.  This occurs as iron pyrite (FeS₂) and on exposure to oxygen, it oxidises to form acidic sulfates.  Sulfuric acid solutions can be found in mineral springs from this oxidation reaction of organic decay.

            2H₂S + 3O₂                 2SO₂ +2H₂O

In industry, the burning of fossil fuels releases their small amounts of sulfur compounds which combine with oxygen during combustion to form sulfur dioxide gas.  The smelting of sulfide ores also produces SO₂.  SO₂ can oxidise in the atmosphere to form SO₃.

Outcome: Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen.

Outcome: Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen.

Process and present information from secondary sources to describe the properties of sulfur oxide and oxides of nitrogen.

Our energy is supplied mainly through the burning of fossil fuels, and the pollutants are released by industrial action.  The burning of these fuels releases sulfur and nitrogen oxides.  Cars are another main source of nitrogen oxides.

Sulfur oxide and nitrogen oxides both dissolve in water to acidic solutions.  In areas of high urban development, sulfur and nitrogen oxides have combined with water vapour to produce airborne acidic droplets, dropping the pH levels of rain water to levels of 4.0 – 5.6.

Outcome: Explain the formation and effects of acid rain.

Choose resources, gather and analyse information from secondary sources to summarise the industrial origins of the above gases and evaluate reasons for concern about their release into the environment.

Acid rain is formed when oxides of sulfur and nitrogen combine with airborne water droplet, acidifying the rain water.

Acid rain dissolves calcium carbonate from marble and limestone which is used in buildings and monuments, corroding them and causing safety concerns.  Acid also attacks sandstones in nature as it attacks their cementing agent of marble CaCO₃.

It removes protective waxes from leaf surfaces, killing many plants.  Acids dissolve and remove important nutrients needed for plant growth in ion exchanges, dissolving minerals and releasing toxic levels of heavy metals into the soil.  Marine life cannot survive in low pH levels.  Bronchitis and asthma are worsened by acid rain.

Outcome: 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 101.3kPa or 25°C and 101.3kPa.

1 mole of gas at STP (0°C and 101.3kPa) = 22.4 L

1 mole of gas at SLC (25°C and 101.3kPa) = 24.5 L

Step 1: Write a balanced chemical equation.

Step 2: Convert the known quantity into moles.

Step 3: Find out how many moles of the unknown are used using ratios.

Step 4: Convert moles of the unknown into units asked for.

Examples:

1.      What volume of carbon dioxide at SLC is formed when 2.33g of copper carbonate is reacted with sulfuric acid?

1.      4.0L of nitrogen dioxide at STP is produced by the reaction of copper metal with nitrous acid.  Copper(II) nitrate and water are also formed.  How much copper is used originally in grams?

Outcome: 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 101.3kPa

When opened, the equilibrium shifts left as CO₂ escapes from the water and enters the atmosphere.  When the amount of CO₂ in the soft drink has reached a balance with the amount of CO₂ in the atmosphere, the soft drink has been decarbonised, or goes flat, losing its acidity.

            The volume of CO₂ calculated could have been slightly inaccurate because of water evaporation.  This was minimised by placing the coke can in an artificially created high humidity environment containing water in an almost sealed plastic bag.  The experiment was also stored in a dark, cool place away from direct sunlight.  To ensure that mass loss is only due to CO₂ and not water evaporation, the volume of coke after the experiment could be measured, the mass of water loss calculated and added to the mass of the coke can after.

1.      Acids occur in many foods, drinks and even within our stomachs

Outcome: Define acids as proton donors and describe the ionisation of acids in water

Bronstod-Lowry defined an acid as a proton donor or something which donated a H⁺ ion in solution.  Bases were thus defined as proton or H⁺ acceptors.

            In a neutralisation reaction, a proton is transferred from the acid to the base.  A new conjugate base and conjugate acid is formed as a result.

If the equilibrium lies to the right, A is a stronger proton donor than CA.  HCl is the acid because it donates H⁺ to the water, and H₂O is a base because it accepts the proton.

If the equilibrium lies to the left, H3O⁺ acts as an acid by donating a proton to the Cl^-, which is a base because it accepts the proton.

Relative strengths of acid-base pairs can be established by finding out where the equilibrium lies.  In this example, the equilibrium lies to the right.

Due to the amphiprotic nature of water, ionisation processes can occur.  In all cases, the product of the hydrogen and hydroxide concentrations is constant, and is called the ionisation constant for water, K_w.

Strong acids and bases completely dissociate or ionise in water, and are strong electrolytes.  Weak acids and bases are compounds which only partially ionise in water.

Outcome: Identify acids such as acetic acid (ethanoic acid), citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid), vitamin C and hydrochloric acid as naturally occuring acids, and acids such as sulfuric acid and hydrobromic acid as manufactured acids.

                 Identify data, gather and process information from secondary sources to identify examples of naturally occuring acids and bases, their chemical composition and describe their pH in their naturally occuring form

Acetic acid, CH₃COOH, is present in vinegar, usually about 4% in solution.  It is a fairly weak acid.  Citric acid (C₆H₈O₇) is found in citrus fruits and is an important acid produced in the metabolism of carbohydrate.  It is an extremely weak acid.  Vitamin C, or ascorbic acid, C₆H₈O₆, is found in fruit and vegetables, and is stronger than vinegar in its first stage.  Hydrochloric acid, HCl, is found diluted 0.1M in gastric juices, with a pH1.

Sulfuric acid, H₂SO₄, is the most widely used industrial acid and is usually manufactured by the contact process:

Hydrobromic acid, HBr is usually gaseous:

Outcome: Describe the use of the pH scale in comparing the concentrations of acids and alkalis.

Outcome: Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute

Strong acids – completely ionise in water

Weak acids – partially ionise in water

Concentrated – have a low pH from high molarity

Dilute – have a higher pH from low molarity

Outcome: Identify pH as –log₁₀[H⁺] and explain that a change in pH of 1 means a ten-fold change in [H⁺]

(From 4.)  Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations

Formulae: pH = –log₁₀[H⁺]

                  pH + pOH º 14

Because of the log₁₀, a change in pH of 1 corresponds to a ten-fold change in [H⁺].

Outcome: Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and relate this to the degree of ionisation of their molecules

K = [H⁺] [A^-]                           where, K is acid strength

          [HA]                                         [HA] is undissociated acid concentration

Outcome: Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecules and its ions

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

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

For a strong acid, such as HCl, the equilibrium would lie strongly to the right, as there are very few intact molecules.  The more ions there are compared to the number of intact molecules, the stronger the acid.

Outcome: Gather and process information from secondary sources to explain the use of acidic oxides such as sulfur dioxide as food additives

A low pH < 4, will prevent the development of some dangerous pathogens in food, allowing it to be preserved.  Acidic oxides, for example sulfur dioxide, is the most effective inhibitor of the deterioration of dried fruits and fruit juices.

Outcome: Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals

Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids

1.      Because 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 refined.

Outcome: Outline the historical development of ideas about acids including those of:

–        Arrhenius

–        Lavoisier

–        Davy

Assess the importance of each definition in terms of understanding

Antoine Lavoiser hypothesised in 1776 that the presence of oxygen in compounds formed from non-metals causes acidity.  As it was inaccurate, his contribution precipitated further experiments which challenged his views and continued discoveries.

            Humphry Davy in 1810 decomposed muriatic acid (HCl) and found that only hydrogen and chlorine were produced, and renamed the acid hydrochloric acid.  He proposed that all acids contain the element hydrogen.  It was a step towards finding out the reasons for why solutions were acidic, basic or neutral.

            Svate Arrhenius (1859 – 1927)proposed a conceptual definition for acids and bases: An acid is a substance that provides H⁺ in aqueous solution.

            A base is a substance that provides OH^- ions in aqueous solution.

Many substances were left out of this list, but it set some groundwork for future developments.

Outcome: Outline the Bronsted-Lowry theory of acids and bases

Many acidic or basic compounds did not fit Arrhenius’s defnition.  JN Bronsted and TM Lowry independently proposed a new theory of acids and bases, that is now known as the Bronsted-Lowry definition: An acid is a proton donor

                                                A base is a proton acceptor.

This is one of the most important contributions, and explains the relationships between acids and bases.

Outcome: Describe the relationship between an acid and its conjugate base and a base and its conjugate acid

Identify conjugate acid/base pairs

Once an acid has reacted, it becomes a conjugate base as it can now accept a proton.  A base accepts a proton, and becomes a conjugate acid because it can donate a proton.

Outcome: Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature

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

Hydrogen salts contain hydrogen ions that can be neutralised by a base.  The anions of these salts in water are amphiprotic as they can both donate and accept protons.

Sodium phosphate is the salt formed on complete neutralisation of the phosphoric acid in step 3.  In steps 2 and 3, the hydrogen salt is behaving as an acid as it neutralises a base.

Basic salts contain oxide or hydroxide ions that can be neutralised by an acid.  They are the chemical opposites of hydrogen salts.  For example, Copper(II) hydroxychloride (Cu(OH)Cl) is formed when copper(II) hydroxide is partially neutralised by hydrochloric acid.

Neutral salts do not have either H⁺ or OH^- ions, for example NaCl, and thus do not form either acidic or basic solutions.

Outcome: Identify amphiprotic substances and write equations to describe their behaviour in acidic and basic solutions

Amphiprotic substances act either as acids or bases depending on conditions.  Water, HCO₃^- and HSO₄^- are common examples.

Outcome: Outline the Lewis definition of an acid

Gather and process information from secondary sources to trace and describe developments in understanding and describing acid/base reactions and use available evidence to explain the expanded view of an acid as developed by Lewis

Gilbert Newton Lewis proposed a theory of acids and bases that is useful in aqueous and non-aqueous systems.  Substances that are proton acceptors have unshared pairs of electrons.  Protons can thus bind to these unshared pairs of electrons with coordinate covalent bonds.

            An acid is an electron pair acceptor.

            A base is an electron pair donor.

Outcome: Identify neutralisation as a proton transfer reaction which is exothermic

Analyse information from secondary sources to predict combining volumes and masses of acids and bases and products in neutralisation reactions

Perform a first-hand investigation to measure temperature change during neutralisation and calculate the molar heat of reaction

A neutralisation reaction is exothermic because of the energy released in transferring a proton.

            HCl + NaOH               NaCl + H₂O

The molar heat of reaction is the energy given off when 1 mole of the chemical completely reacts: DH = mCDT  J

            Where C_water = 4.18 J.ml^-1

                        m = mass in grams

                        DT = change in temperature

Outcome: 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

In chemical spills, neutralisation can easily prevent corrosive, harmful or dangerous chemicals from destroying the environment and posing a threat to human and natural habitants.  By chemically reacting with the harmful chemical, the chemical is converted into other, less harmful substances which can then be cleaned away.

Outcome: Describe the correct technique for conducting titrations and preparation of standard solution

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

Process information from secondary sources to:

–        visualise the rearrangement of particles, and change in electrical conductivity and pH, which occurs in the reaction vessel during titration

–        explain the purpose of accuracy during titration

–        identify and describe modern analytical method used to accurately measure concentrations

Perform a first-hand investigation to determine the concentration of a commercial acidic or alkaline substance such as vinegar, orange juice or window cleaner

Outcome: Qualitatively describe the effect of buffers with reference to a specific example in a natural system

Buffers are mixtures present in our bodily fluids.  A buffer solution is a solution that can resist changes in pH on the addition of small quantities of acid or base.  Our blood system is a well-studied buffer solution.

            Blood pH is maintained in the range of 7.35 – 7.45, by the presence of a weak carbonic acid and its conjugate base, hydrogen carbonate ion.  The carbonic acid is produced by carbon dioxide dissolved in body fluids.

When the concentration of CO₂ is high, the equilibrium is shifted right in the direction of producing more carbonic acid.  This produces more hydrogen ions in the blood, dropping the pH.  This condition is acidois, caused by problems with the respiratory or metabolic system.  This condition can be reversed in a healthy person by the production of hydrogen carbonate ions which shift the equilibrium back to the left, restoring the pH.

            The reverse is alkalosis, when there is a drop in carbon dioxide levels during hyperventilation or rapid breathing.  The pH rises as the equilibrium shifts left, and the pH is restored by producing more carbonic acid through slower breathing.  Increased carbonic acid shifts the equilibrium to form more hydrogen carbonate ions and hydrogen ions.

1.      Esterification is a naturally occuring process which can be modelled in the laboratory

Outcome: Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds

Alkanols are alcohols formed from alkanes, and have OH- as their functional group.  Alkanoic acids are alkyl groups with COOH- as their functional group.

Outcome: Explain the difference in melting point and boiling point caused by the alkanoic acid and alkanol functional groups

The first two members of these groups, Methanol, Ethanol, Methanoic acid and ethanoic acid, have unusually high melting and boiling points when compared to the rest of the series.  This is because of the hydrogen bonding that occurs between their functional groups.  The alkanoic acids have stronger hydrogen bonds than alkanols.

Outcome: Identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification

Alkanols react with alkanoic acids in an acid-catalysed condenssation reaction to form alkyl ethanoates or esters.  Water is formed as a by-product.

Outcome: Describe the purpose of using concentrated sulfuric acid in esterification for catalysis and absorption of water

Concentrated sulfuric acid acts as a dehydrating agent and removes the water that is formed.  It is also a catalyst because it removes a product, shifting the equilibrium to the right.  The minimisation of water present and an excess of one reactant will favour the formation of the ester.

Outcome: Explain the need for refluxing during esterification

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

Reflux is the process of heating a reaction mixture in a vessel attached to a cooling condenser which prevents any loss of vapour.  The reaction mixture continuously boils and vaporises.  The vapours recondense and return to the reaction flask.  Molecules that return to the reaction flash continue to react and so the system will ultimately achieve equilibrium without the loss of hot vapours.  This speeds the whole process of producing esters.  (See prac)

Outcome: Outline some examples of the occurrence, production and uses of esters

Aspirin (Acetylsalicylic acid) is one of the most famous ester.  It is manufactured from salicylic acid (2-hydroxybenzoic acid), its alcohol group condensing with the ethanoic acid to produce the ester.  Salicylic acid is extracted from willow leaves.  It is used as a painkiller.

Outcome: Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics

Esters have very distinctive smells and many are responsible for the flavour and odour of fruits.  They are now commercially manufactured as artificial flavouring and perfume.

Ethyl ethanoate                                     Solvent for paints and lacquers

Butyl ethanoate                                     Nail lacquers