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

Classify common substances as acidic, basic or neutral

Classification Properties Colour Change (litmus) SubstancesAcid Sour-tasting

Corrosive to active metalsDissolves limestoneUnreactive towards fats

Colour changes from blue to red

VinegarSoda waterSoft drinksLemon juiceBattery acid

Neutral Not much tasteUnreactive towards most metalsUnreactive towards limestoneUnreactive towards fats

No colour change. Stays blue/purple

WaterGlucose (sugar)Sodium chloride (table salt)

Base Bitter-tasting (soapy)Unreactive towards most metalsUnreactive towards limestoneReacts with fats

Colour changes from red to blue

Baking soda solutionLime waterAmmonia solutionsBleachDetergent

Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour;

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

Acid-base indicators are substances that change colour depending on the pH

The working/effective range of an indicator is the pH range in which a colour change is observed.

Litmus paper(working range: 5-8) can be used to provide a general classification for a substance to identify whether it is an acid or base. It is best used in solutions that are moderately/strongly basic or acidic.

Methyl orange (working range 3.1-4.4) Bromothymol blue (working range 6.0-7.6) Phenolphthalein (working range 8.3-10.0)

pH Methyl orange Litmus Bromothymol Blue

Phenolphthalein

14 base Yellow Blue Blue Crimson (red)13 Yellow Blue Blue Crimson (red)12 Yellow Blue Blue Crimson (red)11 Yellow Blue Blue Crimson (red)10 Yellow Blue Blue Deep pink9 Yellow Blue Blue Pink8 Yellow Blue-purple Green-blue Light pink7 neutral Yellow Blue-purple Green Colourless6 Yellow Blue-purple Yellow-green Colourless5 Orange Bluish red Yellow Colourless4 Orange-red Red Yellow Colourless3 Orange-red Red Yellow Colourless2 Red Red Yellow Colourless1 Red Red Yellow Colourless0 acid Red Red Yellow Colourless

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

Indicators can be used for various purposes in everyday life.

Testing the acidity of water- In swimming pools, the acidity levels must be monitored in order to effectively kill

microbes. Samples of pool water can be tested using a pool test kit which consists of phenol red indicator. If the pH is below 6.8, the indicator turns yellow; if the indicator is above 8.4, the indicator turns red purple. If the indicator is pink or orange, the acidity of the pool water is appropriate.

Testing the acidity of soil- Soils consisting of different pH are required for different plants/crops. Soil pH can be

measured using electronic instruments or universal indicator. The soil is first mixed with water in a tube and then indicator is added. Sometimes the colour of the soil can hide the indicator colour change; to counter this, a neutral white powder (e.g. barium sulphate) can be added to the top layer of the damp soil before adding indicator. This then allows the colour of the indicator to be clearly visible.

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

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

React with water to form acids. React with bases to form salts.

Non-metal oxides behave as acids. Some non-metal oxides that act as acids are: SO2, NO2 and P2O5 (phosphorus pentoxide).

CO2 (g) + H2O (l) H2CO3 (aq) SO2 (g) + H2O (l) H2SO3 (aq) 2NO2 (g) + H2O (l) HNO3 (aq) + HNO2 (aq) (nitrous acid)

HNO2 (aq) + O2 (g) HNO3 (aq)

P2O5 (g) + H2O (l) 2H3PO4 (aq) (phosphoric acid)

*Note: Exceptions include N2O, CO and NO; these are neutral oxides

Basic oxidesBasic oxides:

React with water to form bases React with acids to form salts

Metal oxides behave as BASES. Some metal oxides that act as bases are: K2O, Na2O and MgO. In solution, they tend to form basic hydroxides

K2O (s) + H2O (l) 2KOH (aq) (potassium hydroxide) Na2O (s) + H2O (l) 2NaOH (aq) (sodium hydroxide) MgO (s) + H2O (l) Mg(OH)2 (aq) (magnesium hydroxide)

Amphoteric oxidesAmphoteric oxides are oxides that can act as both acids and bases (e.g. Al2O3)Their behaviour depends on the reaction they are put in.The only elements that combine to form amphoteric oxides are beryllium, aluminium, zinc, tin and lead.

Analyse the position of these non-metals in the Period Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides

Non-metallic oxides are generally acidic. Acidity generally increases across a period and increases down a group.

Metallic oxides are generally basic (sources of O2+). The basicity of metal oxides generally decreases across a period and increases down a group

Transition metal oxides display variable acid-base behaviour and are often amphoteric. Oxides of metalloids are often amphoteric

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen

Sulfur dioxide:

Natural sources include: volcanoes, bushfires and bacterial decomposition

Industrial sources: burning of coal in electricity production, smelting of metals

Oxides of nitrogen:

Natural sources include: lightning storms

Industrial sources: internal combustion within cars and other heavy machinery

Describe, using equations, examples of chemical reactions which release sulphur dioxide and chemical reactions which release oxides of nitrogen

Sulfur Dioxide:

The burning of sulfur-rich coal and other fossil fuels directly combines sulfur with oxygen:S (s) + O2 (g) SO2 (g)

The extraction of metals from metal sulfides also releases sulfur dioxide. E.g. smelting of galena for lead: 2PbS (s) + 3O2 (g) 2PbO (s) + 2SO2

Oxides of Nitrogen: When nitrogen and oxygen react at high temperature (e.g. in engines of motor vehicles),

nitric oxide is formed:N2 (g) + O2 (g) 2NO (g)

Nitric acid is neutral; however it reacts with oxygen in the air to form nitrogen dioxide, which is acidic.

2NO (g) + O2 (g) NO2 (g)

SO2 (sulfur dioxide) NO (nitric oxide) NO2 (nitrogen dioxide)Properties Colourless

Pungent colourSoluble in water**

ColourlessNo smellInsoluble

Reddish-brownChoking odourSoluble in water **

Uses Food preservativeBleachingFumigant

Nitric acidFertilisersExplosives

Issues: Respiratory irritant (1ppm) – asthmatics are particularly susceptible

Synthesising NO2 Respiratory irritant (3-5ppm). At higher concentrations it can cause tissue damage

Assess the evidence which indicates increases in atmospheric concentration of oxides of sulphur and nitrogen

Since the Industrial Revolution of the 19th Century, vast quantities of nitrogen and sulphur oxides have been released in the air. The main sources of these oxides are from the smelting of metals, burning of coal/oil and internal combustion within engines.

Oxides of sulphur and nitrogen dioxide were difficult to measure due to its relatively low concentrations in the atmosphere. Chemical instruments able to measure such low concentrations were only commercially available since the 1970s. Most of the evidence for the increases in sulphur oxides was in heavily industrial areas, where the increased effects of acid rain (which is formed when the oxides react with rainwater) were observed.

Explain the formation and effects of acid rain

‘Acid rain’ generally describes rain with moderately low pH of around 4-5. It occurs when the atmosphere is polluted with acidic oxides (e.g. sulfur dioxide and nitrogen dioxide) which have high solubility in water. The main causes of acid precipitation are related to industrial emissions.

The following reactions account for the formation of acid rain:

Nitric oxides 2NO2 (g) + H2O (l) HNO3 (aq) + HNO2 (aq) (nitric, nitrous acid) HNO2 (aq) + O2 (g) HNO3 (aq)

Sulfur oxides

SO2 (g) + H2O (l) H2SO3 (aq) (sulfurous acid) 2H2SO3 (aq) + O2 (g) 2H2SO4 (aq)

(OR)

2SO2 (g) + O2 (g) 2SO3 (g)

SO3 (g) + H2O (l) H2SO4 (aq)

Effects of acid rain

Has a corrosive effect on limestone and marble. For example, the calcium carbonate in the marble is attacked by sulfuric acid in acid rain to form insoluble calcium sulfate:

Can attack metallic structures of iron or steel. The iron is oxidised by the acid and becomes chemically weathered:

Acidifies soils, which can inhibit the growth of certain plants. Aquatic organisms are also affected by acid rain as it changes the pH of many lakes. As the

lake water becomes more acidic, the presence of hydronium ions interferes with the carbon dioxide/carbonate equilibrium. Plants that require CO2 and crustaceans (with shells that require carbonate) consequently suffer.

Analyse information from secondary sources to summarise the industrial origins of sulphur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment

Oxides of sulphur and nitrogen tend to cause acid rain (with the exception of NO, which is a neutral oxide, but it reacts with oxygen to form NO2). Also, all oxides of sulphur and nitrogen are respiratory irritants. Both of these effects (acid rain/respiratory irritant) can damage ecosystems if the concentrations of the oxides are not moderated

The acid rain caused by dissolution of SO2 and SO3 in rainwater (H2SO4) lowers the pH of lakes and soils, interfering with the balance of nutrients (e.g. CO3

2+). The acidity can also affect the soil and inhibit growth of certain plants.

Sulfur and nitrogen oxides produced from car exhausts pose a health hazard on population areas due to irritation of the lungs, particularly asthmatics.

NO2 forms photochemical smog. Sunlight reacts with nitrogen dioxide, hydrocarbons and oxygen to form ozone. Ozone is harmful in the troposphere at concentrations as low as 0.1ppm- NO2 NO + O·- O· + O2 O3

Define Le Chatelier’s principle

The French chemist Henri Le Chatelier suggested that:

“When a system at equilibrium is disturbed, the equilibrium position will shift in the direction which tends to minimise, or counteract, the effect of disturbance”

Identify factors which can affect the equilibrium in a reversible reaction

Concentration – if the concentration of a solute reactant is increased, the equilibrium position shifts to use up the added reactants by producing more products. (Note concentration is actually referring to C=n/V)

Pressure – if the pressure of an equilibrium position is increased, then the equilibrium position shifts to reduce the pressure (the side with less moles of gas

Volume – corresponds to pressure. If volume is increased, pressure is reduced; if volume reduced then pressure is increased

Temperature – if the temperature of an endothermic equilibrium system is increased, the equilibrium position shifts to use up the heat by producing more products. If the temperature of an exothermic equilibrium system is increased, the equilibrium position shifts to use up heat by producing more reactants

* For more detail, refer to notes

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

Carbon dioxide in water:

1. CO2 (g) CO2 (aq) [Δv = -ve]

2. CO2 (aq) + H2O (l) H2CO3 (aq)

3. H2CO3 (aq) H+ (aq) + HCO3- (aq)

4. HCO3- (aq) H+ (aq) + CO3

2- (aq)

Like all dissolution reactions, this dissolution reaction is exothermic

Pressure – if we increase the gas pressure, equilibrium 1 shifts to the right, thus dissolving more carbon dioxide. This causes equilibrium 2 to shift to the right (due to increase in concentration of reactant) - causing an increase in the production of carbon acid. The increase in concentration of carbonic acid shifts equilibrium 3 and 4 to the right. The solution becomes more acidic (increased concentration of H+ ions). If we open a bottle of soft drink, the pressure decreases and the opposite effect occurs. Bubbles of CO2 gas can be seen coming out of the solution, and the soft drink becomes ‘flat’ as gaseous carbon dioxide escapes

Temperature – Since equilibrium 1 is exothermic, if we warm the soft drink, the equilibrium shifts towards the left (since heat is thought of as the ‘product’ of equilibrium 1). This leads to the release of CO2 from solution.

pH – If an additional source of acidity was added, the additional H+ would drive equilibrium 3 and 4 towards the left, and eventually lead to the release of CO2 from solution. Conversely, the addition of a base would decrease the concentration of H+ (due to the neutralisation reaction it creates), thus shifting the equilibrium towards the right – dissolving more CO2 and producing more carbonate ions.

Calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at standard temp/pressure

This is based on Avagadro’s deduction that equal volumes of gases, under the same conditions (standard) would

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

Strong acids: Hydrochloric, sulfuric, nitric

Strong bases: Hydroxides of group I, II metals (NaOH)

Weak acids: Acetic, citric, carbonic, hydrogen fluoride, sulfurous, nitrous, phosphoric

Weak bases: Ammonia

Define acids as proton donors and describe the ionisation of acids in waterAn acid can be defined as a substance, that in solution, produces hydrogen ions (H+), or protons. However, in aqueous solution, free H+ does not actually exist; as the hydrogen ion actually combines with a water molecule to form a hydronium ion (H3O+)

An ionisation reaction essentially involves an acid being ionised, for example:HNO3 (l) + H2O (l) H3O+

(aq) + NO3-(aq)

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

Strong acidsHydrochloric acid(HCl)

Hydrochloric acid (HCl) is a strong acid made by passing hydrogen chloride gas into water. The final concentrated solution is about 12 mol/L.

Sulfuric acid(H2SO4)

Sulfuric acid is another strong acid; in concentrated solution it is about 18 mol/L. However, unlike hydrochloric acid, it is diprotic, meaning it can donate two protons per molecule of acid. In contrast, a monoprotic acid is one that can only donate one proton per molecule. Sulfuric acid ionises in two stages:

*note degree of ionisation decreases at second stepNitric acid(HNO3)

Nitric acid is another strong acid, and when concentrated, is about 16 mol/L. Its ionisation equation:

Weak acidsAcetic acid(CH3COOH)

Acidic acid, also known as ethanoic acid, is a weak acid – therefore it is only weakly ionised. Even though it has 4 hydrogen atoms, it is monoprotic. Acetic acid is formed naturally when microbes ferment sugars in fruit juices. It is also used as the major component of vinegar.

Citric acidC3H5O(COOH)3

Citric acid is a weak, tripotic acid (stronger than acetic) found in citrus fruits. Its chemical name is 2-hydroxypropane-1,2,3-tricarboxylic acid.

It ionises in three steps (though the degree of ionisation decreases at each step)

Describe the use of the pH scale in comparing acids and basesThe pH scale is used to determine the acidity or basicity of a substance. It is numbered from 0-14.

A pH of 7 is attributed to neutral substances A pH <7 refers to acidic substances (strong acids closer to 0) A pH >7 refers to basic substances (strong bases closer to 14)

Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and diluteA concentrated solution is one in which the total concentration of solute species is high A dilute solution is one which the total concentration of solute species is lowA strong acid is one in which all the acid present in the solution has ionised to hydronium ionsA weak acid is one in which only some of the acid molecules in solution have ionised to form hydronium ions

Identify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+]The pH of a solution is defined of the negative of the logarithm (base 10 ) of the hydronium ion concentration:

A change in pH of one means a 10-fold change in [H+]. For example a solution with a pH of 3 has 10 times the hydronium ion concentration than a solution with pH 4.

Furthermore, there is a similar scale for [OH-]; the relationships between pH and pOH are:

[H+][OH-] = 10-14

i.e. the product of the ion concentrations must equal to 1.0 x 10-14 (known as the water constant- Kw)Note that water undergoes self-ionisation according to the equation: H2O H+ + OH-. In pure water the [H+] = [OH-] = 10-7.

pH = -log10[H3O+]

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

Degree of ionisation is calculated using: concentration of hydronium ions/ concentration of initial acid, as a percentage.

The strength of an acid is dependent on the ability of the acid to ionise in water. A stronger acid has a higher degree of ionisation than a weaker acid. Strong acids are completely ionised (100%), whilst weak acids have degree of ionisations less than 100%.

The following table represents the degree of ionisation when 0.10 mol/L solutions of hydrochloric, citric, and acetic acid were prepared:Acid (0.01M) Hydrochloric Citric AceticDegree of ionisation (%) 100 27.5 4.2

From the above table, we can conclude that HCl is the strongest acid, followed by citric acid and then acetic acid. Between citric and acetic acid, it is difficult to compare them as citric acid is triprotic whilst acetic acid is monoprotic. However, we know that the higher value for citric acid is mainly from the first step of ionisation.

Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ionsA strong acid is 100% ionised in solution (i.e. there is no unionised molecules of the acid present). This means that its ionisation reaction goes to completion. For example:

A weak acid is not completely ionised in solution, and some of the unionised acid molecules still exist. Thus, its ionisation reaction is written in equilibrium. For example:

Gather and process information from secondary sources to explain the use of acids as food additives:Acids are added to food for 2 reasons: as preservatives, and to add flavour.

Preservatives- Ethanoic acid (in the form of vinegar) is used as a preservative in ‘pickling’. - Propanoic acid is often used as a preservative in bread.- Sulfur dioxide is added to food as a preservative, as it forms sulfurous acid, which kills

bacteria in food.- Citric acid is a natural preservative, often added to jams and conserves.

Flavourings- Carbonic acid is added to soft drinks to add ‘fizz’.- Phosphoric acid is also added to soft drinks to add ‘tartness’ of flavour.- Ethanoic acid, as vinegar, is also used as flavouring.

Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition:Natural acids

Hydrochloric acid: Aqueous HCl; it is produced naturally by the lining of our stomachs. It aids in the digestion of food.

Citric acid: C6H8O7; occurs naturally in large quantities in citrus fruits. Ethanoic acid: CH3COOH; it is found naturally in vinegar, which is produced by the natural

oxidation of ethanol. Lactic acid: C3H6O3; it is formed in the body during strenuous exercise.

Natural bases Ammonia: NH3; it is present in the stale urine of animals. It is also formed through the

anaerobic decay of organic matter. Metallic oxides: E.g. iron(III) oxide, copper(II) oxide and titanium(IV) oxide. These insoluble

oxides are solid bases found in minerals. Calcium carbonate: CaCO3; it is found naturally as limestone.

Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations

An important thing to note is when calculating the pH of acids that are not monoprotic (i.e. are diprotic or triprotic). Write out a balanced equation when calculating the pH resulting from a reaction.

Strong, diprotic acid

Diprotic acids release 2 protons, e.g. H2SO4.H2SO4 2H+ + SO4

2-

For every mole of acid, it releases 2 protons.E.g. Calculate the pH of 0.1 M sulfuric acid:

It has 0.1 mol/L; therefore its [H+] = 2 × 0.1 = 0.2 mol/L So, pH = -log10(0.2) = 0.7

Weak acid

The question should give the degree of ionisation as a percentage.

E.g. Calculate the pH of 0.1 mol/L ethanoic acid if only 1.3% ionises:

[H+] = 1.3% of 0.1 = 0.0013 pH = -log10(0.0013) = 2.9

* For more practice, refer to Faulder’s sheets or past HSC questions

4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, acid and base have been refined

Outline the historical development of ideas about acids including those of: Lavoisier, Davy, and Arrhenius

Antoine Lavoisier was the first chemist to create theory of acids. He showed that many non-metal compounds containing oxygen produced acids as they dissolved in water

Lavoisier hypothesised that the presence of oxygen in the non-metal compounds gave the compounds their acidic properties

However, his theory of acids did not explain why oxides of metals were not acidic

Humphry Davy electrolysed samples of hydrochloric acid solution and showed that it produced hydrogen and chlorine gas; no oxygen was formed.

Thus, Davy proposed that the presence of hydrogen in acids gave them their acidic properties

However, his theory did not explain why many compounds of hydrogen were not acidic (e.g. methane)

Svante Arrhenius noted that all acidic solutions were formed when acids ionised/dissociated into ions (hydronium ion and an anion) as they dissolved in water

Arrhenius recognised that some weaker acids did not ionise as completely as the strong acids in water

He also proposed that a base was a substance that produced hydroxide ions when dissolved in water

However, his theory also only applied to aqueous solution It also did not explain why ammonia (NH3), metal oxides (e.g. MgO) and carbonates are

basic; but these compounds do not contain hydroxide ions.

Outline the Bronsted-Lowry theory of acids and bases

In 1923, Johannes Bronsted (a Danish chemist) and Thomas Lowry (an English chemist) independently proposed a new theory of acids and bases now known as the Bronsted-Lowry theory.

Bronsted and Lowry defined acids and bases as follows:

Acids are proton donors Bases are proton acceptors

Thus, a substance cannot act as an acid (proton donor) without another acting as a base (proton acceptor).

The Bronsted-Lowry theory also explains why acids dissolve in water to produce ions. A proton is donated from the acid to the water molecule to produce a hydronium ion. The presence of the hydronium ion gives the solution its acidic properties. Water is not only behaving as a solvent, but also as a BL-base.

Similarly, a base can be dissolved in water to produce ions. In this instance, water acts as the BL-acid , donating a proton.

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

In an equilibrium reaction between a BL-acid and a BL-base, the BL-acid donates a proton to the BL-base, thus becoming proton deficient after the reaction. The proton deficient species is the conjugate base, as it can accept protons to return back to its previous state. The original BL-base has accepted a proton, and the new species is known as its conjugate acid (as it can donate a proton to revert back).

Conjugates are the acids/bases on the right-hand side of the reaction. A conjugate acid/base pair differs only by one proton (e.g. HF and F-)

Strong acids (e.g. hydrochloric acid and nitric acid) have very weak conjugate bases. The conjugate acids of strong bases such as the hydroxide ion are also very weak. Weaker acids/bases have relatively strong conjugate acids/bases. (Only incompetent/very weak acids and bases have actual ‘strong’ conjugates)

Identify conjugate acid/base pairs

Acid Conjugate base Base Conjugate acidHCl Cl- O2- OH-

HNO3 NO3- OH- H2O

H3O+ H2O S2- HS-

HSO4- SO4

2- CO32- HCO3

-

HF F- NH3 NH4+

HNO2 NO2- HCO3

- H2CO3

CH3COOH CH3COO- HS- H2SH2CO3 HCO3

- CH3COO- CH3COOHNH4

+ NH3 F- HFHCO3 CO3

2- SO42- HSO4

-

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

Basic salts refer to those that form basic solutions; Acidic salts refer to those that form acidic solutions. Other salts are neutral.

The acidity/basicity of salts can be shown through hydrolysis reactions (see if the salt acts as a BL-acid or BL- base as it reacts with water – neutral salts do not react with water). Hydrolysis refers to the breaking of a bond with the addition of water

Weak Acid + Strong Base Basic Salt + Water

Weak Base + Strong Acid Acidic Salt + Water

Strong Acid + Strong Base Neutral Salt + Water

Neutral salts include NaCL, KBr, NaNO3. Acidic and basic salts can be identified through the strength of the acid/base reactants. An alternate method would be to analyse the origins of the cation and anion to determine whether they are neutral, acidic of basic.

Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions

An amphiprotic species is one that can both donate or accept a proton.

H2O is amphiprotic

H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)

HCO3:is amphiprotic ־

- HCO3 ־

(aq) + H3O+ (aq) H2CO3 (aq) + H2O (l)

- HCO3 ־

(aq) + OH ־ (aq) CO3

־2 (aq) + H2O (l)

Identify neutralisation as a proton transfer reaction which is exothermic

All neutralisation reactions are exothermic; they all liberate heat energy

The amount of heat liberated per mole in neutralisation reactions is almost the same no matter how strong the acid/base: ΔH ≈ -56 kJ/mol. Although note that strong acid/base reactions produce slightly more heat per mole

This similarity is understandable as the same reaction occurs in each neutralisation reaction:

- H+ (aq) + OH ־

(aq) H2O (l)

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

It is important to immediately neutralise any chemical spills involving strong acids and bases, as they are corrosive and can be extremely dangerous. Neutralisation reactions are widely used as safety measures in cleaning up after such incidents.

It is important to note that neutralisation reactions are exothermic and thus release large quantities of heat. Care must be taken, including the use of safety goggles and lab coats (in the laboratory).

Strong or concentrated acids and bases must never be used to neutralise spills; if an excess is used, the spill will become dangerous again.

When neutralising an acid or a base the following procedure is followed: The most preferred agents of neutralisation has the properties of being stable, easily

transported, solid (powdered), cheap and amphiprotic (so it can act as a WEAK acid or a WEAK base).

This is the safest material, as it can neutralise both acids and bases; even if an excess is used, it is very weak, and so does not pose any safety risks. The neutralised product (water, and a soluble salt) is then absorbed using paper towels and disposed.

The most common substance used to neutralise spills in laboratories is powdered sodium bicarbonate (NaHCO3); this is because the hydrogen carbonate ion (HCO3

-) is an amphiprotic species, and it is cheap and readily available substance.

In large chemical spills, inert sand or vermiculite is used to prevent the spillage area from further contamination. The acidic sand/vermiculite is then placed in a chemical waste container and removed for neutralisation off site (using sodium bicarbonate).

Describe the correct technique for conducting titrations and preparation of standard solutions

Volumetric analysis is a quantitative technique that involves the determination of unknown concentration of a solution through a chemical reaction with a standardised solution. In acid-base analysis, the reaction involved is a neutralisation reaction.

Apparatus

Volumetric flasks are flasks that hold an accurately known volume of solution. This volume is indicated by a line etched into the neck of the flask (calibration mark). Volumetric flasks are used to the preparation of solutions whose concentrations must be accurately known (primary or secondary standards)

A pipette is used to accurately deliver a specific volume of solution such as 5.00mL, 10.00mL or 20.00mL. The pipette is used to deliver the appropriate volume of solution (an aliquot) into a conical flask prior to titration.

Burettes are used to accurately deliver variable volumes of solution Burettes are usually graduated from 0.0mL to 50.0mL and the reading on the scale is estimated to +-0.05mL. The difference between the initial and final readings on the burette indicates the volume of solution delivered in a titration.

*for diagrams, refer to book

Other terminology:

A known concentration volume of titrant reacts with a solution of titrand to determine concentration. The volume of titrant reacted is called the titre (usually delivered by the burette).

The equivalence point is the point at which the neutralisation reaction is complete – i.e. when a stoichiometric ratio of reactants have been added.

The end point of the titration is the point at which the indicator changes colour. It is best to have the end point as close to the equivalence point as possible.

Preparing standard solutions

A primary standard is prepared using chemicals that are pure and satisfy the following criteria:

High level of purity Accurately known composition Free of moisture Stable and unaffected by air during weighing Readily soluble in pure (distilled) water

High molar weight solid to reduce percentage error in weighing Reacts instantaneously and completely

Many pure substances do not meet such requirements: concentrated HCl fumes and loses HCl gas, concentrated sulphuric acid absorbs water from the atmosphere, NaOH absorbs moisture from the air (hygroscopic) and reacts with CO2.

Hydrated sodium carbonate is also unsuitable as it loses water as it Is being weighed; however, anhydrous sodium carbonate is suitable. It is firstly dried in a drying oven then cooled in a desiccator. The final product is free of water and able to use for a primary standard.

A secondary standard is a solution whose concentration has been determined by reacting with a primary standard.

Method:

1. Clean all apparatus (rinse volumetric flask with distilled water several times)

2. Weigh the watch glass

3. Weight out 1.3g of anhydrous sodium carbonate and record its exact mass

4. Transfer sodium carbonate to the beaker.

5. Wash the sodium carbonate into beaker using a wash bottle

6. Stir until all the sodium carbonate is completely dissolved

7. Pour the solution into the volumetric flask using a filter funnel

8. Rinse the beaker, stirring rod and funnel several times with the wash bottle, ensuring all solution goes into the volumetric flask

9. Add distilled water to the volumetric flask until the level is almost to the graduated mark on the flask (e.g. 5mL away)

10. Using a dropper, add distilled water drop by drop until the bottom of the meniscus is level with the graduation mark

11. Calculate the exact concentration of the solution you have made and write it on the flask

Titrating the primary standard against HCl acid

1. Clean all apparatus (rinse several times with distilled water – for burette and pipette rinse it twice after with small quantities with the solution to be transferred in them; any water droplets will dilute the solution. Conical flask can be left wet as any water remaining will not change the number of moles of solution it holds)

2. Set up the burette in its stand. Check it is vertical and tap is off

3. Pour hydrochloric acid through a funnel into the burette and take a reading

4. Pipette 25mL of standard solution into a conical flask

5. Place a white tile under the conical flask to make it easier to observe colour changes of the indicator

6. Add 3 drops of indicator to the conical flask

7. Adjust burette until the tip is just inside the top of the conical flask

8. Open the tap and slowly add hydrochloric acid to the conical flask. Use left hand to turn the stopper and right hand to gently swirl the conical flask as acid is added. Look down at the colour in the flask against the white tile

9. Stop as soon as there is a permanent change in the colour of the indicator

10. Read the level in the burette

11. Refill burette and repeat steps 5-10 until you have 3 titres within 0.1mL with each other

12. Ignore the first rough titre. Using the three consistent titres, calculate the average value

13. Calculate the concentration of the hydrochloric acid

Risks PrecautionsAcid/base in eyes Safety glassesGlassware can break and cut you Take careSpillages can cause slipping Need to wipe up as soon as possible

*Note that the unknown and known can be placed in either conical flask or burette – it doesn’t matter.

Choice of indicator

The indicator is chosen such that the pH at the end point matches as closely as possible with the pH at the equivalence point of the titration.

The equivalence point is approximately in the middle of the steep slope (inflexion point). Note that weak acid to weak base reactions are generally avoided due to the fact that there is no rapid change in pH even at the equivalence point.

Qualitatively describe the effect of

buffers with reference to a specific

example in a natural system

A buffer is a solution that contains comparable amounts (roughly molar equivalent) of a weak acid and its conjugate base, and is therefore able to maintain a relatively constant pH when small quantities of acid or base are added. The equilibrium involved is:

HA + H2O H3O+ + A-

If hydronium ions are added, then by LCP, the equilibrium shifts to the left; the base A - will combine with much of the added hydronium ions to form HA, in an attempt to minimise the change in hydronium ion concentration. More of the un-ionised (weak) molecular acid is produced, thus minimising pH change.

If instead, hydroxide ions are added, it would react with the hydronium ions, causing pH to rise; however, by LCP the equilibrium will shift to the right and more A - will be produced. Since more of the hydronium ions is then produced, the pH change is minimised.

Example in natural system

An example of a natural buffer system is the carbon acid system. It occurs natural in freshwater lakes and rivers (that contain carbonate rocks from which HCO3

ions can be formed), maintaining the ־constant pH of 6.5-7.5 needed for life to exist. The equilibrium involved is:

H2CO3 (aq) + H2O (l) H3O+ (aq) + HCO3

־(aq)

Carbon dioxide from the air dissolves in the water, forming carbonic acid:

CO2 (aq) + H2O (l) H2CO3 (aq)

Meanwhile, its conjugate base, HCO3 is present as ions leeched out of rocks and minerals ,־

containing the lake/river. These processes produce comparable amounts of H2CO3 / HCO3 to ־

produce a buffer system.

There needs to be a source of HCO3 .apart from dissolved CO2 in order to produce this buffer system ־

Therefore, rainwater is not a buffer (as the only source of HCO3(is through the ionisation of H2CO3 ־

5. Esterification is a naturally occurring process which can be performed in the laboratory

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

A functional group is a classified group of atoms in a molecule which is responsible for the characteristic properties of that molecule.

The same functional group with undergo the same/similar chemical reactions regardless of the size of the molecule. Alkanols and alkanoic acids are part of the functional group: organic compounds.

Alkanols

Alkanols are a homologous series (a series of compounds with the same general formulae) that are:

Derived from alkanes Contain the hydroxyl group (-OH)

They have the general formula:

CnH2n+1OH

Neutral

Alkanoic acids

Alkanoic acids are considered to be weak organic acids that are:

Derived from alkanes Contain the carboxylic functional group (-COOH) Acidic (H+ ions disassociate from COO- in solution)

Alkanol Alkanoic acid- Possesses a bent geometry around the

oxygen atom- Two bond dipoles: C-O and O-H bonds- Polar molecules

- An oxygen is double bonded to a central carbon, and an –OH group is single-bonded to the same carbon. (Note the –OH here is not called a hydroxyl group)

- The carboxyl group is always at the end of the chain

- More polar than alkanol due to polar C-O, O-H C=O bonds

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

The strength of intermolecular forces determines the boiling and melting points of a molecule. The stronger the intermolecular forces, the more energy needed to break these forces (i.e. higher MP and BP)

For the same number of carbons in a straight carbon-chain, the highest BP’s and MP’s belong to: the alkanoic acids, then the alcohols and then the parent alkanes.

There is a general trend of increasing BP/MP with increasing molecular weight. This is because as MW increases along a homologous series, the SA increases and therefore so too does the number and strength of dispersion forces.

Alkanes Alkanols Alkanoic acidsThe only intermolecular forces are dispersion forces. They are very weak; hence the BP and MP’s of alkanes are low.

Alkanols experience dispersion forces, but also dipole-dipole interactions due to the polar C-O and O-H bonds creating a net dipole. Furthermore, alkanols can interact through hydrogen bonding (due to the OH bond).The intermolecular forces in alkanol molecules require more energy to overcome – giving rise to higher MP and BP’s.

Like alkanols, alkanoic acids have dispersion forces, dipole-dipole interactions and hydrogen bonds; however the additional double bonded C=O means it is even more polar than alkanols – resulting in stronger dipole-dipole forces.The intermolecular forces in alkanoic acids require the most energy to overcome – thus it has the highest MP and BP’s.

*the hydrogen bonding in alkanols and alkanoic acids also allows for their high solubility in water

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

Naming esters:

Alkanol part first Alkanoic part last Write as two separate words (alkyl alkanoate)

Drawing esters (includes structural and condensed structural formulae)

Alkanoic acid first Alkanol part second

Example

Above – Propyl acetate (CH3COOCH2CH2CH3)

*Exceptions when naming alkanoic acids are: methanoic and ethanoic acid, for which the preferred IUPAC names are formic and acetic acid respectively

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

Esterification is the process which forms esters. It involves an acid-catalysed condensation reaction between an alkanoic acid and an alkanol (or more generally between a carboxylic acid and an alcohol)

Esters are sweet-smelling, volatile organic compounds which contain the ester functional group

“-COOC-”

Note, important steps when writing out equations of esterification

Write equilibrium arrow Write catalyst on top of the equilibrium arrow Write water as a product

Example of esterification

Describe the purpose of using acid in esterification for catalysis

Explain the need for refluxing during esterification

Esterification reactions are usually quite slow and the reaction does not proceed to completion as it is a reversible reaction in equilibrium. In order to increase the rate of reaction, we:

Add a strong, concentrated acid as a catalyst Heat the reaction mixture

Purpose of the acid

The purpose of the acid is to decrease the time for the reaction to reach equilibrium and to increase yield.

The concentrated, sulphuric acid acts as a catalyst as well as a dehydrating agent. The acid speeds up the rate of reaction by lowering the activation energy. Furthermore, by removing water from the reaction, the equilibrium is shifted towards the products; thus increasing the yield of product – although this increase is quite small.

The need for refluxing

By heating the reaction mixture (increasing the kinetic energy of the reaction), increasing the amount of collisions between molecules and thus increase the rate of reaction. Refluxing, therefore, maximises the rate of reaction by allowing the mixture to be heated at B.P. without losing any of the reactants.

The reactants and products of the esterification reaction are flammable and volatile, and so we cannot use a naked flame such as that of a Bunsen

burner (flame is too hot, and could ignite the reactants/products). These issues can be avoided by heating the mixture using a reflux apparatus.

The mixture is heated by a hot-water bath supported by an electric hot plate. A water condenser is mounted above the reaction flask (round bottomed flask), and cold water circulates to cool the hot, rising vapours. The vapours condense back to liquid state and drip back into the reaction flask. This process therefore allows the heating of the reaction at BP, without losing any of the reactants or products.

The system is also open to the atmosphere to avoid build- up of pressure due to the production of vapours. Small, boiling chips are dropped in the reaction mixture to enable a large SA for vaporisation to occur without the risk of superheating and the explosive ejection of vapours.

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

Esters occurs widely in both nature and are produced synthetically in industry.

In nature, esters are found in

ATP (important in reactions in cells) Fats and oils Natural waxes

Esters are also manufactured for a wide range of industrial applications

Lubricants Solvents Plasticisers Flavouring/fragrances

Occurrence Production UseNatural fatty acids Produced from natural oils such

as soybean and canola using transesterification. An ester is formed by reacting vegetable oil with sodium hydroxide dissolved in methanol. After processing, it can be used as biofuel in diesel engines.

Alternate source of fuel (biofuel)

Phthalate esters Produced from the reaction between Phthalic acid and a dialkyl.

Plasticiser (added to plastics to soften and increase flexibility)

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

The characteristic tastes and smells of fresh fruits and flowers are due to a complex mixture of chemicals, including esters. As a result, esters are often used as flavour enhancers and in fragrances in cosmetics.

For example:

The ester octyl acetate, which is found in oranges and other citrus fruits, is often used in lollypops or soft drinks to give a fruity orange flavour. The esterification reaction involves 1-octanol and acetic acid.

Methyl salicylate (wintergreen oil) is an organic ester produced by wintergreens, and is synthetically produced for use in lotions and creams to soothe sore muscles. It is an ester of salicylic acid and methanol.