1.1.1 Classify common substances as acidic, basic or neutral Acids Bases An acid is a substance which in solution produces hydrogen ions, H+ or more strictly H3O+, sometimes called the hydronium ion. Common properties of all acids are:

• Acids have a sour taste • Acids sting or burn the skin • In solution, acids conduct

electricity • Acids turn blue litmus (a

vegetable dye) red Examples of Acidic Substances

• Orange – Citric Acid • Vinegar – Acetic Acid • Soft Drink – Carbonic Acid • Stomach Acid –

Hydrochloric Acid • Car Battery Acid – Sulfuric

Acid

- A base is a substance which either contains the oxide O2- or hydroxide ion OH- or which in solution produces the hydroxide ion. A soluble base is called an alkali.

Common properties of alkalis are:

• Alkalis have a soapy feel • Alkalis have a bitter taste • In solution, alkalis are good

conductors of electricity • Alkalies turn red litmus blue

Example of Basic Substances

• Baking Soda - Sodium hydrogen carbonate Drain Cleaner – Sodium Hydroxide

• Window Cleaner – Ammonia • Toothpaste, Soap

Common examples of Neutral Substances

• Pure Water • Sugar Solution • Sodium Chloride Solution • Pure Alcohol Solution

1.1.2 Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour Indicator Colour change pH range Methyl orange Red – Yellow pH 3.1 – 4.4 Bromophenol Blue Yellow – Blue pH 3.0 – 4.6 Bromocresol Blue Yellow – Blue pH 3.8 – 5.4 Methyl Red Pink – Yellow pH 4.4 – 6.0 Bromothymol Blue Yellow – Blue pH 6.2 – 7.6 Phenol Red Yellow – Red pH 6.8 – 8.4 Thymol Blue Yellow – Blue pH 8.0 – 9.6 Phenolphthalein Colourless – Red pH 8.3 – 10

1.1.3 Identify and describes some everyday uses of indicators including testing of soil Soil Testing Some plants need an acidic soil while others need an alkaline soil Acidity is usually difficult to measure with indicator due to the darkness Soil is moistened with water and universal indicator added White barium sulfate powder places on surface of soil White powder absorbs soil water so that color can be seen against white background Chemical Analysis Chemists routinely use indicators in analytical work Used in titrations to signal point when the acid neutralizes the base Visible signal as most of these solutions are colorless Monitoring Pool Acidity Acidity levels need to be carefully controlled at 7.4 to reduce skin and eye irritation Regularly tested with indicator with a color range similar to that of universal indicator Monitoring wastes from photographic processing (discharges to the sewerage system must be nearly neutral: photographic solutions are often highly alkaline). 1.2.1 Perform a first-hand investigation to prepare and test a natural indicator Aim: to prepare and test a natural indicator (red cabbage) Method: Obtain a small handful of red cabbage Heat the cabbage in a 250ml baker with about 80ml of hot water Allow the water to boil Decant the solution into a clear beaker Test the solution with the following substances, HCl, Ammonia solution, acetic acid, NaOH Tabulate observations Results: Substance Acid/Base Indicator Colour HCl Acid Red/pink Ammonia solution Base Dark green Acetic Acid Acid Dark Pink NaOH Base Green Conclusions: the red cabbage produced a purple solution which was tested and found to be read in an acid solution and green in a basic solution.

1.2.2 Identify data and choose resources to gather information about the colour changes of a range of indicators Aim: To find the colours of various indicators in acidic basic and neutral solutions Materials: Indicators – litmus, methyl orange, bromomethyl blue, phenolphthalein. Method: Test each indicator by adding 2-3 drops to 1ml of HCl, NaOH, Na2CO3 and H2O respectively Tabulate colours of indicators in each solution Results: Reaction with Substances Substance L M.O B.B P U.I HCl Orange red Green colourless Red NaOH, Blue Orange Blue Red Purple Na2CO3 Purple Orange blue Red Purple H2O L. purple orange green colourless Green Conclusion: HCl is acidic, NaOH, Na2CO3 basic, H2O neutral 1.2.3 Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic Aim: To measure the pH range of various common substances using a pH meter and universal indicator. Method:

1. Check pH meter is correctly calibrated and review instructions for use. Ensure pH meter is properly handled and rinsed during testing.

2. Place a small amount of solution to be tested in a beaker (about 3cm depth). Insert the pH meter electrode and record the pH.

3. Measure the pH by adding a couple of drops of universal indicator to the beaker and record the results.

4. Repeat the process with all samples being sure to clean the pH meter electrode and beaker between samples.

Results: Sample

Universal Indicator Colour and pH

pH Meter Reading

Acidic/Alkaline/ Neutral

Distilled water 0.1Mol solution

Green 7 Dark purple

6.9 12.0

Neutral/Acidic Alkaline

NaOH Ammonia solution NH4OH Dishwashing Liquid Sodium Hydrogen Carbonate NaCl solution Effervescent Aspirin Carbonated Lemonade Apple Juice Lemon Juice Vinegar Hydrochloric Acid HCl

11 Purple 9.5 Pale Green 7.5 Purple 10 Green 7 Yellow 6.5 Pink 4 Pink 4 These substances went pink and read as 4, but if the scale kept going readings would have been lower.

11.2 8.3 9.5 7.4 6.2 2.9 3.8 2.8 2.7 1.7

Alkaline Alkaline Alkaline Neutral/Alkaline Acidic Acidic Acidic Acidic Acidic Acidic

2.1.1 Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids Background Information Acid react with hydroxides to for salt and water HCl + NaOH Na2Cl + H2O Acids react with metal oxides to form salt and water 2HCl + MgO MgCl2 +H2O Acids react with NH3 to form a salt HCl + NH3 NH4Cl Acids react with reactive metals to form a salt and hydrogen gas 2HCl + Mg MgCl2 + H2

Acids react with carbonates to form a salt, carbon dioxide and water 2HCl + 2Na2CO3 2NaCl +CO2 +H2O Oxides of Non-Metals Acidic Oxides Basic Oxides Reacts with water to form an acid or CO2 + H2O H2CO3 SO2 + H2O H2SO3 SO3 + H2O H2SO4 2NO2 + H2O HNO3 + HNO2 Reacts with bases to form salts (or does both) are oxides of non-metals Common acidic oxides are CO2, SO2, SO3, and NO2

Reacts with acids to form salt Does not react with alkali solutions (such as NaOH or KOH). Are oxides of metal and their ionic compounds Common basic oxides are Na2O, MgO, CuO

2.1.2 Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides. The basic oxides are oxides of metal and their ionic compounds. Elements to form basic oxides appear towards the left of the periodic table. Acidic oxides are generally oxides of non-metals. Hey are all covalent compounds. Elements that form acidic oxides appear towards the right and top of the periodic table. As the elements became more and more non-metallic in character, their oxides became more and more acidic. Aluminium oxide is classed as amphoteric because t shows acidic properties as well as basic properties. 3rd Period

2.1.3 Define Le Chatelier’s Principle Equilibrium - When reactions are reversible reactions. The reaction can go in the forward direction or in the reverse direction Characteristics - of a chemical systematic equilibrium. 1-It is a closed system 2-the rate of the forward reaction is equal to the rate for the reverse reaction 3- the concentration of each reactant remains constant, but not necessarily equal. 4- the system has constant macroscopic properties(eg. Color, temperature, and pressure)

Le Chatelier’s Principle If a systematic equilibrium is disturbed, then the system adjusts to minimize the disturbance. This means that either the forward or reverse reaction proceeds at a faster rate until equilibrium is re-established. A system is in equilibrium when the forward and reverse reactions are occurring at the same rate and therefore does not go to completion. For equilibrium to be established the system must be closed because the relative amount of the products and reactants at equilibrium can be affected by shifting the equilibrium in order to favor one side. 2.1.4 Identify factors which can affect the equilibrium in a reversible reaction Factors that can affect equilibrium are

• Temperature change • Pressure increase • Addition of Reactant or Product

How an equilibrium system responds to changes.

• Temp increase - shifts in direction of ENDO • Temp decrease – shifts in direction of EXO • Reactant conclusion increases – shifts towards products RHS • Reactant conclusion decreases – shifts towards reactants LHS • Product conclusion increases – shifts towards reactants LHS • Product conclusion decreases – shifts towards products LHS • Pressure increases – shifts towards side with lower no. of moles of gas • Pressure decreases– shifts towards side with higher no. of moles of gas • Changes in amount of solid and liquid substance have no effect

2.1.5 Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle CO2 dissolves in water producing carbonic acid (exothermic reaction) as shown in the equation

CO2 (g) CO2 (aq) + heat

CO2 (g) + H2O (l) H2CO3 (aq) H+ + HCO3- + heat

Disturbance Effect on equilibrium (right means inc. solubility) Change in pressure Increased pressure – shifts to the right – solubility increased -

more dissolved CO2 Decreased pressure – shifts to the left – solubility decreased - more CO2 gas.

Change in concentration of CO2 in system

Increased concentration of CO2- shifts right – solubility increased - more dissolved CO2. Decrease in concentration of CO2 – shifts left – solubility decreased - more CO2 gas

Change in temperature Increased temp – shifts left – solubility decreased - more CO2 gas Decreased temp – shifts right – solubility increased – more dissolved CO2.

Changing pH Adding NaOH to equilibrium- neutralizes carbonic acid- moves to the right Adding H+ shifts to the left

2.1.6 Identify natural and industrial sources of sulphur dioxide and oxides of nitrogen

Sources Sulfur dioxide (SO2) Natural • Geothermal hot springs

• Volcanoes industrial • Processing or burning fossil fuels

S compounds + O2(g) SO2(g)

• Extracting metals from sulfide ores 2ZnS(g) + 3O2(g) 3ZnO(s) + 2SO2(g))

Sources Nitrogen Dioxide (NO2) Natural • Action of sunlight on NO and O2

2NO(g) + O2(g) 2NO2(g) industrial • Combustion of fuel in motor vehicles

• Power Stations

Sources Dinitrogen monoxide (N2O) Natural • Produced by bacteria in soil industrial • Fuel for racing cars

Sources Nitrogen Monoxide (NO) Natural • Produced by bacteria in soil

• Lighting N2(g) + O2(g) 2NO(g)

industrial • Combustion of fuel in motor vehicles • Burning of Biomass

2.1.7 Asses evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen. The evidence which indicates increases in atmospheric concentrations of these gases have been gathered from measurements on bubbles of ancient air trapped in ice sheets of Antarctica and Greenland. This is then compared to more modern levels. This evidence successfully shows the increase in the atmospheric oxides of sulfur and nitrogen. The burning of coal and petroleum products as well as the smelting of minerals has led to a significant increase in the levels of acidic oxides in the atmosphere. This is particularly true for heavily industrialized nations, such as the USA. The annual average concentration of SO2 and NO2 in most large cities around the world is about 0.01 for each gas. This is about

ten times the value for clean air. In Sydney and Melbourne there is some concern about the number of days per year that the safe levels for these gases are exceeded. Globally, because SO2 and NO2 are washed out of the atmosphere by rain, there appears not have been any significant build-up of their concentrations over the last century or so (unlike CO2 which there has been a 30% increase and N2O increased as well). 2.1.9 Explain the formation and effects of acid rain The presence of acidic oxides into the atmosphere leads to the production of acids rain. Acidic oxides dissolve in water droplets in the atmosphere and these droplets fall to earth as rain Weak sulfurous acids (H3SO4) is formed when SO2 dissolves in water SO2(g) + H2O(l) H2SO3(aq)

Strong sulfuric acids (H2SO4) can form if the sulfurous acid is catalytically oxidized in the air 2H2O(l) + H2SO3(aq) 2H2SO4(aq) A mixture of weak nitrous and strong nitric acids is formed when NO2 dissolves in water droplets 2NO2(g) + H2O(l) HNO2 (aq) + HNO3(aq)

Nitrous acid is unstable and can decompose to form nitric acid and nitric oxide or it can be catalytically oxidised to form nitric acid HNO2 HNO3(aq) + H2O(l) + 2NO2(g)

HNO2 + O2 HNO3 2.2.2 Analyze 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 released make acid rain. Acid rain causes Weathering and erosion to marble and sandstone buildings Destroys forest by taking the waxy cuticle on the surface of the leaves, as well as damaging roots Increased soil acids Acidification of lakes and waterway causes fish and organism to die or prevents them from reproducing. Sulfur dioxide itself is a toxic gas. It irritates eyes, respiratory tract and causes lung damage and asthma

Nitric oxides is also damaging Is harmful to respiratory system, causing decreased lung function, susceptibility to respiratory infections and increased sensitivity to asthma triggers. 2.1.8 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

Gay Lussacs Law The volumes of reacting gases and the resulting products are always in simple number ratios. H2SO4 + 2NaOH → Na2SO4 + 2H2O Mole Ratio: 1 : 2 : 1 : 2 Avogadro’s Law Equal Volumes of equal gases at the same temp and pressure contain the same number of moles. Moles of Solids: Moles = Mass _ Molar Mass Moles of Gases: Moles = Volume _ Molar Volume Or Volume = Moles X Molar Volume Molar Volume is the same for all gases S.T.P - 0 oC (273 oK) and 100 kPa - Molar Volume = 22.71 L S.L.C -25 oC (298 oK) and 100 kPa - Molar Volume = 24.79 L Concentration Formula: Moles = Concentration Volume Dilution Formula: C1V1 = C2V2

2.2.1 Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25˚C and 100kPa Aim: To calculate the volume of carbon dioxide released from a can of soft drink Materials: can of soda water, NaCl, electronic balance Method:

1. Find the mass of a can of soda water 2. Finds the mass of a 500ml beaker 3. Pour the drink into the beaker 4. Find the mass of the empty can 5. Add a known mass of sodium chloride (e.g. 25g) 6. Stir the salt and drink mixture until no more gas is released 7. Find the final mass of the beaker and the drink (by minus-ing the weight of the salt)

Results Closed Can : 265.73g Empty Can: 15.2g Salt: 25.02g Beaker: 199.79g Mass of soda water: 250.53 Beaker + water +salt: 439.44 Soda water after CO2 release: 234.63 CO2 loss: 15.9 N(CO2)= 15.9 = .36 44 Volume = .36 X 24.79 = 8.96 3.1.1 Define Acids as proton donors and describe the ionization of acids in water Ionisation is the process of the formation of ions When an acid dissolves in water it ionizes as follows HA + H2O H3O+ + A-

any acid hydronium ion any ion During ionization, the acid molecule donates a proton (H+) to the water forming and hydronium ion H3O+ Eg. H2O + H+ H3O+ HCl + H2O H3O+ + Cl –

3.1.2 Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acid Acid Formula acetic (ethanoic), CH3COOH(aq)

found in vinegar citric (2-hydroxypropane-1,2,3-tricarboxylic) (COOH)CH2CH(OH)(COOH)CH2COOH(aq)

found in citrus fruits

hydrochloric HCl found in stomach juices sulfuric H2SO4(aq) found in volcanic springs 3.1.4 Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute A strong acid is one which completely ionises in solution to form H+. Eg. HCl , H2SO4, HNO3. A weak acid is one which is only partially ionized in solution. Acetic acid, CH3COOH, nitrous acid H2NO3, carbonic acid H2CO3, sulfurous acid H2SO3 Concentrated acid is one which has a high number of moles per litre (molarity). Dilute acid is one which has a low number of moles per litre. 3.1.7 Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions Strong acids completely ionise and the hydrogen ion released joins with water molecules to create the hydronium ion (H3O+) there is no equilibrium in the ionisation reaction for strong acids and if there were it would lie completely to the right. Weak acids only partially ionise and thus an equilibrium is formed, this occurs with all weak acids like ethanoic acid. In the equilibrium of the weak acid, the equation lies mostly on the left (partial ionisation). The molecule of the weak acid is in solution with few of its ions. CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO–(aq)

3.1.6 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 The strength of an acid is related to the degree of ionization. The strengths of acids can be compared by measuring the pH of different acids at the same concentration. Strengths of acids can also be compared by measuring their electrical conductivity at the same concentration. Acids ionise in water and become proton donors, forming [H+] ions in water. The greater the concentration of [H+], the greater is the strength of the acid. Hydrochloric acid ionises more completely than citric acid which ionises more than acetic acid. Hydrochloric > citric acetic > acetic 3.1.3/5 Describe the use of the pH scale in comparing acids and bases and identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+] The pH is the Potential (p) Hydrogen (H) or H3O+ concentration of a solution. It is a logarithmic scale that includes numbers 1 to 14 (the logarithm indicates that a change of 1 unit on the pH scale is a change to a power of ten in actual H3O+ concentration). A reading of more than 7 on the pH scale indicates the solution is basic, and a number below 7 is acid (7 is neutral).

pH [H+] [OH-] [H+] × [OH-] 0 100 = 1 10-14 10-14 1 10-1 10-13 10-14 2 10-2 10-12 10-14 3 10-3 10-11 10-14 4 10-4 10-10 10-14 5 10-5 10-9 10-14 6 10-6 10-8 10-14 7 10-7 10-7 10-14 8 10-8 10-6 10-14 9 10-9 10-5 10-14 10 10-10 10-4 10-14 11 10-11 10-3 10-14 12 10-12 10-2 10-14 13 10-13 10-1 10-14 14 10-14 100 = 1 10-14

From the table, the 0.1mol L-1 has a pH of 1, and an H+ concentration [H+] of 10-1mol L-1. The 0.01 mol L-1 has a pH of 2, but an [H+] of 10-2 mol L-1.

The difference between the two in only one pH value, but the concentration has changed by a power of ten. This is the logarithmic relationship between pH value, and the corresponding [H+] value. Basically a change of one in the pH scale means a ten fold change in the concentration of hydrogen ions because pH is based on a logarithmic scale (base 10). To work out [H+] from a known pH or to work out pH from a known [H+]:

Self ionization of water H20 OH- + H+ H20 + H20 OH- + H3O+

In pure water at 25 oC -7 [H3O+] = [OH-] = 1 X 10 molL-1 [H3O+] X [OH-] = 1 X 10-14 [H3O+] > [OH-] = acidic [H3O+] < [OH-] = basic Note: the number of decimal places for pH should equal the number of significant figures for [H3O+]

3.2.1 Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals pH meters can be used to measure pH they give fairly accurate values but are relatively expensive. A cheaper way of estimating the pH of solutions I s the use of indicators A pH meter is a non destructive way of testing whether a solution is acidic, basic or neutral, the solutions tested are unaffected, and using indicator solution is a destructive way of testing. Indicator will contaminate a portion of the solution tested.

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

Aim: To determine the pH of identical concentrations of weak and strong acids. Method:

1. Use known weak (CH3COOH) and strong (HCl) acids. 2. Pour equal quantities of each into equal quantities of

distilled water in separate beakers, to form .1 mol/L 3. Use a pH probe to determine the pH in the beakers. 4. Change the concentrations by removing 25ml of acid

solutions with a pipette and adding to 250mL of distilled water forming .01mol/L

Results:

Substance pH CH3COOH 0.1 L-1 3 HCl 0.1mol L-1 1

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

HCl(l) + H2O(l) H3O+(aq) + Cl- H2SO4(l) + H2O(l) H3O+(aq) + HSO4-(aq)

HSO4-(aq) + H2O(l) H3O+(aq) + SO42-(aq) CH3COOH(l) + H2O(l) CH3COO-(aq) + H3O+(aq)

CO2(g) + H2O(l) CO32-(aq) + 2H+(aq) 2H+(aq) + H2O(l) 2H3O+(aq)

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

To model ionization, transfer the appropriate hydrogen atom form an acid molecular model to a water model (hydronium ion) Advantages & Disadvantages Models help to visualize and understand the concepts involved, eg model are 3D and clearly show bonding. Constructing a model reinforces understanding of structure and formulae. Models however have limitation- they are simplifications which are not to scale, do not show eg, electrons and give no indication of concentration.

3.2.5 Gather and process information from secondary sources to explain the use of acids as food additives Food preservatives can be divided into two categories: those which stop the growth of micro organisms, and those which prevent food being spoiled by exposure to oxygen.

• Citric & tartaric acids are often added to jam to give a sharp taste. There acidity also prevents the growth of micro organisms

• Acetic acid in vinegar heaps preserve chutneys • Acidic sulphur dioxide is used in dried fruit and white wine to stop attacks from

microbes. • Ascorbic acid is an antioxidant thus is added to food to prevent spoilage by

oxidation. 3.2.6 Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition • Hydrochloric acid (HCl) is produced by glands in the lining of our stomachs to form an

acidic environment for the efficient operation of the enzymes that break complex food molecules into easily transportable small molecules that are absorbed into the blood stream when they pass into the intestine

• Acetic acid (systematic name ethanoic acid), CH3-COOH. Acetic acid is present in vinegar which is commonly made from wine by oxidation of ethanol.

• Citric acid (systematic name 2-hydroxypropane-1, 2, 3-tricarboxylic acid with molecular formula C6H8O7). It occurs in citrus fruit.

• Vitamin C or ascorbic acid which has the molecular formula, C6H8O6. It occurs widely in fruits and vegetables and is an essential part of our diet.

• Uracil C4H2OON2H2 is a base 3.2.7 Process information from secondary sources to calculate pH of strong acids, given appropriate hydrogen ion concentrations. You can use the formulae You can use pH meters or indicators You can measure electrical conductivity-strong acids are ore conductive as they fully ionise and have more charge carriers

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

Lavoisier Davy Arrhenius

Lavoisier Was called father of chemistry. He showed that combustion involves oxygen. His combustion experiments led him to believe that acids were made of 2 substances, one of which was oxygen. He believed that oxygen was present in all acids and caused their acidity. Davy Demonstrated that muratic acid (HCl) was a compound of hydrogen and chlorine and did not contain oxygen. He observed that metals could displace hydrogen form acids and concluded that acids contain hydrogen Arrhenius Suggested that acids are neural substances that dissolve in water and dissociate to give positive hydrogen ions and negative ions. Positive ions are called cations and negative ions are called anions According to this theory -an acid is a substance that ionises in water to produce H ions, as the only positive ions -a base is a substance that ionises in water to produce OH ions, as the only negative ions His theory has limitations -it applies only to aqueous solutions -does not account for why salts act as acids or bases - does not include amphoteric substances (ZNO) which can act as acids and bases. 4.1.2 Outline the Brönsted-Lowry theory of acids and bases An acid base reaction involves proton transfer from acid to the base Acid is a proton donor. Base is a proton acceptor An acid looses its proton to form its conjugate base. The base accepts a proton to form its conjugate acid A substance which can act as a proton donor and acceptor is amphoteric 4.1.3 Describe the relationship between an acid and its conjugate base and a base and its conjugate acid An acid gives up a proton(H ion) to form its conjugate base. A base accepts a proton to form its conjugate acid. A conjugate base of a weak acid is a strong base

A conjugate base of a strong acid is a weak base HCl (aq) + H2O (l) → H3O+ (aq) + Cl–(aq) ACID BASE CONJUGATE ACID CONJUGATE BASE 4.1.4 Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature Find the Sheet you lost. Identify Conjugate acid/base pairs. Conjugate base Acid + water → H3O+ + conjugate base HCl + H2O → H3O+ + Cl- CH3COOH + H2O H3O+ + CH3COO- Conjugate acid Base + water → conjugate acid + OH- NH3 + H2O NH4+ + OH- CO32- + H2O HCO3- + OH-

4.1.6 Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions Amphoteric substance- general term which refers to any substance displaying both acid and base behaviour. Amphiprotic substance -A substance which can act both as a proton donor and a proton acceptor. H2O, and HCO3 are examples. HCO3 is an amphiprotic substance As a base (accepts proton) HCO3–(aq) + H3O+ (aq) → H2CO3 (aq) + H2O (l) As a acid (donates proton) HCO3–(aq) + OH–(aq) → CO32–(aq) + H2O (l) 4.1.7 Identify neutralisation as a proton transfer reaction which is exothermic Neutralization is a proton transfer reaction (protons are transferred from one species to another to form a salt). All neutralization reactions are exothermic.

HNO3 (aq) + NH3 (aq) NH4+ (aq) + NO3-(aq) The acid transfers a proton to ammonia (to form the NH4+ ion) HCl (aq) + NaOH(aq) Na Cl(aq) + H2O(l) HCl (aq) + OH- (aq) Cl- (aq) + H2O(l) HCl transfers a proton to the base Oh to form water Net ionic equation of neutralisation reactions are H+ OH- H2O(l) General equation is Acid + Base Water + Salt + Heat 4.1.8 Describe the correct technique for conducting titrations and preparation of standard solutions A standard solution must

• Be available in pure form • Accurately know chemical formulae • Mass must not change when exposed to air(rules out all chemical that lose or gain

water or react with chemicals in the air) • Its molar mass should be relatively high. This minimising weighing errors • Must be soluble

Common solutions are Na2CO3(base) anhydrous sodium carbonate, H2C2O4.2H2O(acid) hydrated oxalic acid.

PHOTOCOPY

4.1.9 Qualitatively describe the effect of buffers with reference to a specific example in a natural system A buffer solution resists changes in pH when small amounts of acid or alkali are added to it An acidic buffer solution can be made by mixing a weak acid by its conjugate base. An example is a solution of ethanoic acid and ethanoate. 1. CH3COOH CH3COO- + H+ 2. CH3COOH Na CH3COO- + Na+ When H+ are added the equib shifts left When OH- are added equib shifts right A basic buffer solution is made by mixing a weak base with its conjugate acid.Eg. NH3/NH4Cl Carbonic acid and hydrogen carbonate buffer- This buffer system occurs naturally in lakes and rivers, due to the dissolution of CO2 from the air and CO32- minerals (limestone) into the water. 1. CO2 + H2O H2CO3 2. CaCO3 + H20 CaOH + HCO3- 1+2. H2CO3 + H2O H3O+ HCO3- Carbonic acid ionisation counteracts the addition of acid by reacting it with H+ to form H2CO3, the equib shifts to the left, where there is already H2CO3 so the equib shifts back to normal maintaining the pH. The CO3 acid ionisation counteracts any pH changes when bases are added to the buffer. The equib shifts to the right to produce more H3O+ ions. It was the absence of this buffering action in lakes in Scandinavia that resulted in the first detection of falling pH from acid rain. A buffer solution is a solution which contains comparable amounts of a weak aid and is conjugate base which is therefore able to maintain an approximately constant pH even when significant amounts of strong acid or strong base are added to it. 4.1.3 Describe the relationship between an acid and its conjugate base and a base and its conjugate acid An acid gives up a proton(H ion) to form its conjugate base. A base accepts a proton to form its conjugate acid. A conjugate base of a weak acid is a strong base A conjugate base of a strong acid is a weak base HCl (aq) + H2O (l) → H3O+ (aq) + Cl–(aq) ACID BASE CONJUGATE ACID CONJUGATE BASE

4.1.4 Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature Find the Sheet you lost. Identify Conjugate acid/base pairs. Conjugate base Acid + water → H3O+ + conjugate base HCl + H2O → H3O+ + Cl- CH3COOH + H2O H3O+ + CH3COO- Conjugate acid Base + water → conjugate acid + OH- NH3 + H2O NH4+ + OH- CO32- + H2O HCO3- + OH-

4.1.6 Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions Amphoteric substance- general term which refers to any substance displaying both acid and base behaviour. Amphiprotic substance -A substance which can act both as a proton donor and a proton acceptor. H2O, and HCO3 are examples. HCO3 is an amphiprotic substance As a base (accepts proton) HCO3–(aq) + H3O+ (aq) → H2CO3 (aq) + H2O (l) As a acid (donates proton) HCO3–(aq) + OH–(aq) → CO32–(aq) + H2O (l) 4.1.7 Identify neutralisation as a proton transfer reaction which is exothermic Neutralization is a proton transfer reaction (protons are transferred from one species to another to form a salt). All neutralization reactions are exothermic.

HNO3 (aq) + NH3 (aq) NH4+ (aq) + NO3-(aq) The acid transfers a proton to ammonia (to form the NH4+ ion) HCl (aq) + NaOH(aq) Na Cl(aq) + H2O(l) HCl (aq) + OH- (aq) Cl- (aq) + H2O(l) HCl transfers a proton to the base Oh to form water Net ionic equation of neutralisation reactions are

H+ OH- H2O(l) General equation is Acid + Base Water + Salt + Heat 4.1.8 Describe the correct technique for conducting titrations and preparation of standard solutions A standard solution must

• Be available in pure form • Accurately know chemical formulae • Mass must not change when exposed to air(rules out all chemical that lose or gain

water or react with chemicals in the air) • Its molar mass should be relatively high. This minimising weighing errors • Must be soluble

Common solutions are Na2CO3(base) anhydrous sodium carbonate, H2C2O4.2H2O(acid) hydrated oxalic acid.

PHOTOCOPY

4.1.9 Qualitatively describe the effect of buffers with reference to a specific example in a natural system A buffer solution resists changes in pH when small amounts of acid or alkali are added to it An acidic buffer solution can be made by mixing a weak acid by its conjugate base. An example is a solution of ethanoic acid and ethanoate. 1. CH3COOH CH3COO- + H+ 2. CH3COOH Na CH3COO- + Na+ When H+ are added the equib shifts left When OH- are added equib shifts right

A basic buffer solution is made by mixing a weak base with its conjugate acid.Eg. NH3/NH4Cl Carbonic acid and hydrogen carbonate buffer- This buffer system occurs naturally in lakes and rivers, due to the dissolution of CO2 from the air and CO32- minerals (limestone) into the water. 1. CO2 + H2O H2CO3 2. CaCO3 + H20 CaOH + HCO3- 1+2. H2CO3 + H2O H3O+ HCO3- Carbonic acid ionisation counteracts the addition of acid by reacting it with H+ to form H2CO3, the equib shifts to the left, where there is already H2CO3 so the equib shifts back to normal maintaining the pH. The CO3 acid ionisation counteracts any pH changes when bases are added to the buffer. The equib shifts to the right to produce more H3O+ ions. It was the absence of this buffering action in lakes in Scandinavia that resulted in the first detection of falling pH from acid rain. A buffer solution is a solution which contains comparable amounts of a weak aid and is conjugate base which is therefore able to maintain an approximately constant pH even when significant amounts of strong acid or strong base are added to it. 4.2.1 Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions 4.2.2 Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions 4.2.3 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

4.2.4 Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies

4.2.5 Analyze information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills Acids and base used in industry can often be determined as they are corrosive. Neutralization reactions are therefore widely used in laboratories and factories were acids and bases need to be discarded, this is hazardous to the environment due to varying pH levels and to individuals working with these chemicals. Neutralization is employed to ensure that chemicals which are spilled are neither acidic or basic. Amphiprotic substances are useful for neutralizing chemical spills. These substances contain amphiprotic ions. Such as the hydrogen carbonate ion in NaHCO3. If the chemical spill contains an acid. If the chemical spill contains a base This amphiprotic substance is used due to its stable state as a solid easy to store and handle. It is also cheap and does not cause major problems when used excessively. It is also found in solid form and hence can easily soak up toxin without spreading it. It is also a weak base and so will not heat up excessively with its exothermic reaction. The use of neutralization reactions a safety measure is effective. It is successful in making the acid or base less harmful. This also protects individuals involved around the toxin. Through the use of neutralization management of accident can be safely controlled. 5.1.1 Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds. Alkanols The functional group in alkanols in the hydroxyl-group (OH) – alcohol group The general formula is. There are some isomers of alcohols

Alkanoic Acids. The functional group in Alkanoic acids is the carboxyl-group The general formula is, this is applicable only after methanoic acid.

Alkanoic acids have no isomers

Polarity Both C-O and O-H bonds are polar this means Alkanols and Alkanoic acids are polar molecules However the C=O bond is Alkanoic acids is also polar so with these 3 polar bonds Alkanoic acids are the even more polar. The O-H bonds make hydrogen bonds, consequently there are strong intermolecular forces, in Alkanols and Alkanoic acids. 5.1.2 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 The alcohol (ending in yl without the ol) alkanoic acid-(ending with noate without ic) # of C Alkanol # of C Alkanoic Acid Ester Formed 1 3 5 7 9

Methanol Propanol Pentanol Septanol Nonanol

2 4 6 8 10

Ethanoic Acid Butanoic Acid Hexanoic Acid Octanoic Acid Decanoic Acid

Methyl Ethanoate Propyl Butanoate Pentyl Hexanoate Septyl Octanoate Nonyl Decanoate

5.1.3 Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary Alkanols structures Alkanoic acids like Alkanols have hydrogen bonding between molecules. Alkanols contain the hydroxyl group (-OH) but alkanoic acids contain the carboxyl group (-COOH) thus in alkanoic acids in addition to the (-OH) group there is another polar group (-C=O) which contributes further to the strength of the intermolecular forces between the acid molecules. The hydrogen bonding is stronger in alkanoic acids to the presence of this functional polar group.

Thus Alkanoic acids have a higher melting and moiling points than the corresponding Alkanols, due to the more extensive hydrogen bonding and an increase in dipole-dipole interactions. Hydrogen bonding in Alkanols: Hydrogen bonding in Alkanoic Acid:

.

Photocopy Picture from Fosters notes

5.1.4 Identify esterification as the reaction between an acid and an Alkanol and describe, using equations, examples of esterification Esters are formed when alkanoic acids react with Alkanols.

ALKANOL + ALKANOIC ACID ESTER + WATER Eg. Methanol + Ethanoic Acid MethylEthanoate + Water

5.1.5 Describe the purpose of using acid in esterification for catalysis Esterification is reversible and des not go to completion. That is it reaches equilibrium. Concentrate H2SO4 is used as a catalyse to increase the rate of reaction by lowering the activation energy. It also absorbs the product water and so forces the equib. to the right. The minimization of water present and an excess of one reactant will favor the formation of the ester. 5.1.6 Explain the need for refluxing during esterification Reflux is the process of heating a reaction mixture in a vessel attached to a cooling condenser which prevents any loss of vapor. The reaction mixture continuously boils and vaporizes. The vapors 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 vapors. This speeds the whole process of producing esters. The alternative is to use a closed system but the pressure would be to high. So basically 1. To prevent loss of volatile liquids. 2. To allow the reaction to be carried out at higher temps.

5.1.7 Outline some example so the occurrence, production and uses of esters. Flavours Octyl Ethanoate (orange flavour) Butyl Butanoate (pineapple flavour) Perfumes Be Processed foods Cosmetics

Esters occur naturally in the form of flavourings and scents. The odours and flavours of fruits are caused by the presence of esters. Animal fats and plant oils are also esters. Some esters are used in industry as solvents or thinners as they are able to dissolve many polar and non polar substances. Manufactured esters are used for flavouring food and colouring and flavouring cosmetics. Ethyl acetate is nail polish remover. Esters have pleasant odours, occurring in perfumes and flavouring, naturally in fruits, plants. They are manufactured as synthetic flavours and perfumes, often to mimic the flavours of plants (by first identifying the main esters present in a substance). Produced, these are much cheaper than natural extraction. Eg pentyl ethanoate = banana flavouring, octyl ethanoate = orange flavouring Esters also are used as solvents (eg Ethyl acetate) for polar and non-polar compounds and soaps. Aspirin (acetylsalicylic acid) is an example of an ester used in medicine. Due to their dissolving tendencies towards polar and non-polar organic substances, esters are used in varnishes, lacquers and paints and some are used as plasticisers (eg. for PVC).

Flavour Main Esters present Apricot Ethyl butanoate Banana Pentyl ethanoate Rum Ethyl methanoate Orange Octyl ethanoate

Esters occur naturally. The presence of esters gives flavour to fruits and vegetables, and scents to anything between flowers and rancid butter. Solid animal fats and plant oils are also natural esters. Mixing common natural and synthetic esters can produce scents in cosmetics, and colour and flavouring for food. Some esters are also used in industry as solvents or thinners as they are able to dissolve many polar and non-polar substances. Some examples of the use of esters: