9.3 the acidic environment

7/30/2019 9.3 the Acidic Environment

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The Acidic Environment

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

Classify common substances as acidic, basic or neutral- Common substances have varying pH, being classed as either acidic, basic or neutral- Acids, bases and neutral substances are usually classed based on their propertiesProperties:

Type of Substance Properties

Acidic - Tastes sour- Changes indicator colours (e.g. blue litmus to red)- Conducts electricity- Reacts with metals to produce hydrogen gas- Reacts with carbonates to form carbon dioxide and water- Reacts with metal oxides and hydroxides to produce salt and water

Basic - Tastes bitter- Changes indicator colours (e.g. red litmus to blue)- Conducts electricity- Reacts with amphoteric metals such as aluminium to produce hydrogen gas- Dissolves amphoteric metal hydroxides such as Al(OH)3

Neutral (water) - Almost none of the properties of acids or bases- Does not react with most metals- Does not conduct electricity

Common Substances:

Acidic Basic Neutral

Lemon Juice Cloudy Ammonia Water

Vinegar Detergents Salt Water

Gastric Juice Antacid Tablet Ethanol

Soft Drink Caustic Soda Glucose Solution

Solve problems by applying information about the colour changes of indicators to classify somehousehold substances and acidic, neutral or basic

- The acidity, basicity or neutrality of a substance can be determined by observing the colour change that itimposes on an indicator

- Multiple indicators are used if a single one is not enough to classify the substances pH- The range of colour change of each indicator is noted, and if a substance is within this range of colour change,

the colour will be a mixture of the two extremities and it narrows down the pH greatly

Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue canbe used to determine the acidic or basic nature of a material over a range, and that the range is

identified by the change in indicator colour

- Indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to measure thepH of a substance

- They are dyes and they do this by changing colours depending on the pH solution- This change in colour can be used to determine the pH of a solutionIndicator Colour Change pH Range

Litmus Red Blue 4.8 8.1


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Phenolphthalein Colourless Red/Magenta 8.3 10

Bromothymol Blue Yellow Blue 6.0 7.6

Methyl Orange Red Yellow 3.1 4.4 Identify data and choose resources to gather information about the colour changes of a range of

indicators

Summary of Investigation:

- An investigation was performed to gather information about the colour change of a range of indicators and theresults were compared to second hand sources

- The investigation involved preparing many batches of pH solutions ranging from 0-14 using serial dilutions,adding a different indicator to each

- The results varied slightly from the actual ranges of these indicators due to factors such as errors made duringthe serial dilution

Results:

pH/Test Tube

Colour of the Indicator

Phenolphthalein Methyl Orange Litmus Solution Bromothymol

Blue

Universal

indicator

0 Colourless Red Red Yellow Orange Red



3 Colourless Peach Red Lime Green Red

4 Colourless Yellow Purple Teal Orange

5 Colourless Yellow Purple Blue Light Green

6 Colourless Yellow Purple Teal Light Green7 Colourless Yellow Purple Blue Light Green

8 Colourless Yellow Purple Teal Green

9 Colourless Yellow Purple Blue Green

10 Colourless Yellow Purple Blue Teal

11 Colourless Yellow Purple Blue Purple

12 Pale Magenta Yellow Purple Blue Purple

13 Magenta Yellow Blue Blue Purple

14 Magenta Yellow Blue Blue Purple

Actual Values:

Indicator Colour Change pH Range

Litmus Red Blue 4.8 8.1

Phenolphthalein Colourless Red/Magenta 8.3 10

Bromothymol Blue Yellow Blue 6.0 7.6

Methyl Orange Red Yellow 3.2 4.4 Perform a first-hand investigation to prepare and test a natural indicator

Aim:

- To prepare and test a natural indicator found in red cabbageMethod:

Part 1Making the Indicator:

1. A small handful of shredded red cabbage was placed into a pestle


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2. A large pinch of sand and 5-10mL of ethanol was also added into the pestle3. The mixture was carefully ground with a mortar until the ethanol had become intensely coloured4. The solution was decanted into a small beaker and covered with plastic wrap for use in part 2, noting the initial

colour of the indicator

Part 2Testing the Indicator:

1. 15 test tubes were set up on a test tube rack and numbered 0-142. 20mL of 1M HCl was added into the test tube number 0, and 20 mL of 1M NaOH added to the test tube

numbered 14, and 18mL of distilled water was added to the test tube numbered 7

3. The 1M solutions were diluted inwards by a serial dilution of a factor of 10 using a pipette until all the test tubesfrom 0-14 were filled with a solution

4. 5 drops of the indicator was added to each test tube and shaken and the colour change was recordedResults:

Part 1Making the Indicator:

- The ethanol turned form a clear, colourless liquid to a cloudy deep purple liquid, which was the initial colour ofthe indicator

Part 2Testing the Indicator:

Analysis:

- Red cabbage contains two pigments that are responsible for its indicator properties (anthocyanin and flavonol)- The wide range of colours displayed when this dye was tested was due to its two pigments interacting with each

other in different pH ranges

Pigment Acidic Colour Neutral Colour Basic Colour

Anthocyanin Red Blue/mauve Blue

Flavonol Colourless Colourless Yellow

Discussion:

Justification of Method:

- The sand was added into the mixture as it is an abrasive material that can break the cell walls and membranes ofthe cabbage in order to release the dye

- The ethanol was added as it acts as a solvent to dissolve the indicator once it was released from the cabbage-

The serial dilution was done to ensure that there was a solution for every pH range, providing a full range toshow the full extent of the indicator and the extent of its colour change

Conclusion:


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- The natural indicator in red cabbage was prepared and tested over a range of pH

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

- Soil is very dark in colour, and will most likely absorb any indicator added to it, making viewing the colour changeimpossible

- In order to test the pH of soil, an insoluble, white neutral salt such as barium sulfate which can absorb moistureis sprinkled over the soil

- This provides a white background to view the colour change of the indicator, whilst not altering the pH in anyway as the salt is neutral

Pool Water:

- Pool water must be maintained in a constant and narrow range of around 7.4 to not irritate skin or eyes- This is done by adding a few drops of phenol red into a sample of pool water, with a peachy colour indicating a

pH of around 7.4

Aquarium Water:

- Fish can only survive in a very limited pH range so the pH must be maintained in order for the fish to survive- Different fish can survive on different pHs, so different indicators would be used depending on the fish living in

the aquarium water


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2. While we usually think of the air around us a neutral, the atmosphere naturallycontains acidic oxides of carbon, nitrogen and sulfur. 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 asacids

- Oxides of non-metals which act as acids include carbon dioxide, sulfur and oxides of nitrogen- These oxides only act as acids when they are dissolved with water, and then hydrolysing (reacting to create a

change in pH) to form carbonic, sulfurous/sulfuric and nitrous/nitric acids respectively

Equations:

Oxides of Carbon:

- Oxides of Sulfur:

- -

Oxides of Nitrogen:

- 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

- Many acidic oxides are non-metal oxides located on the right hand side of the periodic table- The left hand side consists of metals, which usually either form basic oxides or neutral insoluble oxides- There are a few elements in between in the semi-metals that form amphoteric oxides, oxides which can act as

both an acid or a base


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- Neutral oxides are those that do not react with either acids or bases and include oxides of hydrogen, carbon andnitrogen.

- The noble gases are stable and hence do not react to form oxides

Define Le Chateliers principle- Le Chateliers principle states that If a system is at equilibrium and a change is made that disturbs the

equilibrium, then the system responds in such a way as to counteract the change and eventually a new

equilibrium is established

Identify factors which can affect the equilibrium in a reversible reaction- Factors which can affect the equilibrium in a reversible reaction include temperature, concentration and total

gas pressure

Temperature:

- In a reversible reaction, normally one of the reactions will be exothermic, making the reaction that occurs in theopposite direction endothermic

- If the temperature is increased, the reaction shifts to favour the endothermic reaction in order to counteract theincrease in temperature, absorbing the temperature through the endothermic reaction

Concentration:

- In a reversible reaction, when the concentration of one of the reactants or products is increased or decreased,the equilibrium is disturbed

- The system shifts in a fashion in order to counteract this concentration change, such as favouring the forwardreaction if the concentration of the reactants is increased

- This shift will counteract the increase in the concentration of the reactants, decreasing the concentration until anew equilibrium is formed

Total Gas Pressure:


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- In a reversible reaction where gases are involved, gas pressure plays a part in the equilibrium- Since the same number of molecules of gas will occupy the same volume, if the total gas pressure is increased, to

counteract this change the reaction will shift to the side which produces the least molecules of gas, as to

decrease the pressure

- Conversely, if the pressure is decreased, the reaction will shift to the side which produces the most molecules ofgas, as to maintain the pressure

Catalyst:

- When a catalyst is added, it does not shift the equilibrium of the system- Instead, it speeds up both the forward and reverse reactions, so the equilibrium is reached faster

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium processand explain in terms of Le Chateliers principle

Carbon Dioxide:

- Carbon dioxide is a covalent molecular molecule that is non-polar due to its symmetrical structure and nounbonded electron pairs

- Due to its non-polar nature, it is only slightly soluble in water, and thus can be shown as an equilibrium reactionwhen it is being dissolved in water,

Reactions:

- These reactions show the process of carbon dioxide dissolving in water and forming carbonic acid1. 2. 3. 4. Changing Solubility:

Temperature Change:

- It can be seen from the first equation that the dissolution of carbon dioxide into water is an exothermic reaction- This means that as the water is heated up, the reaction will shift in favour to the endothermic reaction in

accordance to Le Chateliers principle, so the solubility of carbon dioxide will decrease

Pressure Change:

- If the pressure of the system was increased, the reaction will shift to the side with the least moles of gas, whichis to the right

- Therefore the solubility of carbon dioxide will increase if the pressure is increasedAdding Acid:

- If an acid is added to the solution, the hydronium ion concentration will increase, disturbing the equilibrium


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- This will shift reactions 3 and 4 to the left in order to remove the introduced ions, which will subsequently shiftreaction 1 to the left, and therefore decrease the solubility of carbon dioxide

Adding Base:

- If a base is added into the solution, the hydroxide ions will react with the hydronium ions to form water in aneutralisation reaction

- This removes the hydronium ions, as well as increasing the number of water molecules, causing reactions 2, 3and 4 to all shift to the right, causing reaction 1 to also shift to the right, increasing the solubility of carbon

dioxide

Adding Carbonate:

- If a carbonate is added into the water, the increase in carbonate ions will shift reaction 4 to the left tocounteract the increase

- This will subsequently shift equations 3, 2 and 1 to the left, and thus causing the solubility of carbon dioxide todecrease 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 25oC and 100kPa

Aim:

- To find the mass of carbon dioxide in a can of soft drink and calculate the volume of carbon dioxide at 25oC and100kPa

Method:

1. The mass of an empty bottle was weighed using a triple beam balance and recorded2. The mass of a full can of lemonade was weighed using a triple beam balance and recorded3. The lemonade was poured from the can into the empty bottle4. The mass of the empty can was weighed and recorded5.

The bottle was shaken to release the carbon dioxide, uncapping the lid often to release the gas

6. Measured amounts of salt were added to the bottle and shaken7. When the salt became saturated and no more carbon dioxide was being released, the bottle containing the

decarbonated lemonade and salt solution was weighed and recorded

Results:

Object Mass (g)

Empty Bottle 51.2

Full Can 417.2

Filter Paper 0.95

Empty Can 16.2Salt 141.895

Bottle with Final Solution 590.5

Analysis:


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- The mass of the carbon dioxide that was liberated from the solution was first calculated, being 3.55g- This value was then used to calculate the number of moles that was released, being 0.080663485 mol- The number of moles was then used to calculate the volume it would occupy at 25oC and 100kPa, being 1.9996LDiscussion:

Justification of Method:

- Many of the steps conducted in the method ensure the validity of the method as it ensured only the variablethat was intended to be measures was measured

- Weighing the can before and after removing the soft drink allowed for an accurate measurement on exactly howmuch soft drink was transferred into the bottle

- Weighing the empty bottle allowed for the mass of the contents to be known at all times- The weighing of the salt added ensures that it doesnt alter the result in any way, and it also does not invalidate

the experiment as it does not alter how much carbon dioxide was originally in the solution

Why Salt was Added:

- The salt can help liberate the carbon dioxide quicker as carbon dioxide is a non-polar molecule so it is not verysoluble in water

- Salt however, is an ionic substance and is therefore highly soluble in water- The water molecules will interact with the sodium and chloride ions rather than the carbon dioxide molecules

which displaces the molecules from the solution, turning it into a gas

Conclusion:

- The mass of carbon dioxide found in a can of soft drink was 3.55g, and this was calculated to be 1.9996L of gas at25

o

C and 100kPa

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogenSulfur Dioxide:

Natural:

- About of total sulfur dioxide emissions can be accounted for from natural sources- These include volcanoes, hot springs, burning organic matter, oxidation of hydrogen sulfide and the decay of

organic matter

Industrial:

- Industrial sources make up the remainder of the sulfur dioxide emissions- These include combustion of fossil fuels as they contain sulfur impurities, smelting of sulfide ores, manufacture

of sulfuric acid, petroleum refining and making coke from coal

Oxides of Nitrogen:

- The three main oxides of nitrogen released are Dinitrogen Monoxide (N2O, Nitrous Oxide), Nitrogen Monoxide(NO, nitric oxide) and Nitrogen Dioxide (NO2)


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Natural:

- Natural sources of oxides of nitrogen include lightning strikes causing oxygen and nitrogen in the air to react,creating nitric oxide, and nitrous oxide formed by soil bacteria

Industrial Sources:

- Industrial sources of oxides of nitrogen include combustion reactions that reach high enough temperatures like1300

oC causing a similar reaction to lighting, and agricultural practices that add nitrogenous fertiliser increases

the rate at which the bacteria produce nitric oxide to release

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

Sulfur Dioxide:

- Hydrogen Sulfide is oxidised in the decaying of natural matter, - Combustion of fuels with sulfur impurities, - Smelting zinc sulfide ores, Oxides of Nitrogen:

- Reaction from lighting or combustion reactions, - Nitric oxide reacting with oxygen to form nitrogen dioxide,

Assess the evidence which indicates increases in atmospheric concentrations of oxides of sulfur andnitrogen

- There is much evidence to support the increase in atmospheric concentration of oxides of sulfur and nitrogenOxides of Nitrogen:

Antarctic Ice Cores:

- Air bubbles trapped in Antarctic ice contain samples of air from many years ago- Through the analysis of these air bubbles, it can be seen that oxides of nitrogen remained steady until the

industrial revolution, where they have begun to grow at a rate of about 0.2 to 0.3% per year

Photochemical Smog:

- Oxides of nitrogen also contribute to a photochemical smog that can cover industrialised areas- This smog is caused by a reaction between oxides of nitrogen and ozone, creating an aerosol layer- Therefore the amount of photochemical smog in the air can provide information about the amount of the oxides

of nitrogen in the atmosphere, and can also be used as evidence


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Oxides of Sulfur:

- Sulfur dioxide levels are harder to measure as it readily decomposes into other materials and is removed fromthe air rather quickly, but there is still some evidence of its increase in the atmosphere

Acidity of Rain:

- Evidence for increasing amounts of sulfur dioxide in the atmosphere is the increasing acidity of rain- This is especially evident in more industrialised areas, where many of the forest and aquatic systems have

suffered greatly from the acid rain

- This is caused by the oxides of sulfur dissolving and reacting with water to form sulfurous and sulfuric acid, andcan be used as evidence for the increasing atmospheric concentration

Increased Acidity of Aquatic Systems:

- As a consequence of the increase in the amount of acid rain, many lakes and rivers will also have an increasedacidity

- This can also be used as evidence for the increase of acidic oxides in the atmosphere, as the acid rain causing thisis formed by the increasing concentration of acidic oxides

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

Summary of Origins:

- Industrial origins of sulfur dioxide and the oxides of nitrogen are the burning of fossil fuels, smelting of ores,petroleum refining, industrial combustion reactions and agricultural practices that release nitrogen into soil

Concern about their Release:

Sulfur Dioxide:

- The release of sulfur dioxide creates concerns as it can irritate the respiratory system, causing breathingproblems in people

- Also, even at low concentrations in the atmosphere, it can form acid rain, with many environmental impacts

Oxides of Nitrogen:

- There are many concerns with the release of oxides of nitrogen into the atmosphere, with nitrous oxide being apotent greenhouse gas

- Nitric oxide and nitrogen dioxide can lead to photochemical smog and acid rain, as well as nitrogen dioxide beingable to irritate respiratory tracts, leading to breathing discomfort

Explain the formation and effects of acid rainFormation:

- Acid rain forms when gases in the atmosphere such as carbon dioxide, sulfur dioxide and nitrogen dioxide,amongst other acidic oxides, dissolve in the water through ionisation to form weak acids

- This causes the water to be slightly acidic when it precipitates- Although this removes these gases from the atmosphere, acid rain can cause damaging effects on the ground,

especially in areas where there is a high concentration of these acidic oxides

- Sulfur dioxide forms weak sulfurous acid (H2SO3) whereas sulfur trioxide forms strong sulfuric acid (H2SO4)


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- Sulfurous acid can also be oxidised with a catalyst in order to produce sulfuric acid- Nitrogen dioxide produces weak nitrous acid (HNO2) and also strong nitric acid (HNO3) when it dissolves in water- Nitrous acid can also be oxidised with a catalyst to produce the stronger nitric acidEffects:

- The acidity from the acid rain can change soil acidity, making it difficult for plant roots to take up essentialminerals such as calcium and potassium, inhibit their growth

- Many micro-organisms in the soil that are essential for recycling of nutrients (decomposition) are also killed asthey can only survive in a specific pH range

- The waxy cuticle of leaves can be removed, causing leaf damage and damage to the plant- Many buildings made of limestone, marble or concrete can be damaged as the acid reacts with the stone and

dissolves it

- Photochemical smog and acid rain can combine to become a health hazard to people- Acid rain will increase the acidity of aquatic ecosystems, reducing the pH and killing the organisms, with most

fish and shellfish dying at a pH of 4.5 to 5.0, and killing fish eggs at a pH of below 5.5

- Increased acidity can also release high levels of normally insoluble, toxic elements such as aluminium andmercury in runoff, contributing to heavy metal poisoning in aquatic life

Calculate volumes of gases given masses of some substances in reactions, and calculate masses ofsubstances given gaseous volumes in reactions involving gases at 0

oC and 100kPa or 25

oC and

100kPa

- In each case, the number of moles of gas must first be calculated using either

or

- When the number of moles is determined, this can then be used to calculate the value the question asks

3. Acids occur in many foods, drinks and even with our stomachs Define acids as proton donors and describe the ionisation of acids in water

Acids:

- Acids are substances which donate a hydrogen ion, which consists of only a proton, so therefore acids are protondonors

- Usually this donated proton does not exist alone and reacts with other substances- In the ionisation of acids in water, the donated hydrogen ion reacts with water molecules to form hydronium

ions

Equations:


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- -

Gather and process information from secondary sources to write ionic equations to representionisation of acids

Equations:

- General Equation: - Ionic Equation: - Net Ionic Equation:

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

- Some common acids include acetic or ethanoic acid, citric acid, hydrochloric acid and sulfuric acidAcetic Acid:

Summary:

- Acetic acid, also known as ethanoic acid, is the acid that is commonly found in vinegar- It is naturally occurring in the decomposition of organic material, but is most commonly made industrially- Ethanoic acid is a weak and monoprotic acid

Structural Formula:

Hydrochloric Acid:

- Hydrochloric acid is produced industrially on a large scale and is a strong, monoprotic acidSulfuric Acid:

- Sulfuric acid is also produced industrially on a large scale as well as being produced naturally in the atmospherethrough the formation of acid rain

- Sulfuric is a strong and diprotic acid

Citric Acid:

Summary:

- Citric acid, also known as 2 hydroxypropane 1, 2, 3 tricarboxylic acid, is a naturally occurring acid in allcitrus fruits and is used as a preservative in many foods

- Citric acid is a weak and triprotic acidStructural Formula:


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Describe the uses of the pH scale in comparing acids and bases- The pH (potential of hydrogen) scale measures the concentration of hydrogen ions in a solution- Water self ionises to produce H+ and OH- naturally, with pure water producing these ions in equal amounts of

110-7

mol/L

- It has also been found that when the concentrations of H+ and OH- in water are multiplied together at 25oC, aconstant value is given, known as the water constant of 110

-14mol

2/L

2

- Since the hydrogen ions and the hydroxide ions will always multiply to give this constant, KW, therefore the pHscale, which measure the concentration of hydrogen ions in a solution, can be used to compare the acidity and

basicity of solution

pH [H+] [OH

-] [H

+][OH

-] Example

0 10-0

10-14

10-14

1M Hydrochloric Acid

1 10-1

10-13

10-14

0.1M Hydrochloric Acid

2 10-2

10-12

10-14

0.01M Hydrochloric Acid

3 10-3

10-11

10-14

Soda water, wine

4 10-4

10-10

10-14

Tomato juice, beer

5 10-5

10-9

10-14

Acid rain

6 10-6 10-8 10-14 Urine

7 10-7

10-7

10-14

Pure water without any dissolved gas

8 10-8

10-6

10-14

Sea water

9 10-9

10-5

10-14

Detergent solution

10 10-10

10-4

10-14

Concentrated detergent

11 10-11

10-3

10-14

Household ammonia

12 10-12

10-2

10-14

0.01M sodium hydroxide

13 10-13

10-1

10-14

0.1M sodium hydroxide

14 10-14

10-0

10-14

1M sodium hydroxide

Describe acids and their solutions with the appropriate use of the terms strong, weak, concentratedand dilute

Strong and Weak Acids:

Strong Acids:

- Strong acids are acids that completely ionise when it is in a solution with water- These reaction of the ionisation of strong acids in water can hence be written using a unidirectional arrow, as the

reaction only occurs in one direction


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Weak Acids:

- Weak acids, however, do not completely ionise, with only some of the acid donating a proton leaving thereaming acid as an unchanged molecule in aqueous form

- The reaction of the ionisation of weak acids in water can hence be written using an equilibrium arrow, as sincethey do not ionise completely, the reaction can occur in both directions

Concentrated and Dilute Acids:

- Whether or not an acid is concentrated or dilute is independent of the strength of the acid- The degree of ionisation is irrelevant, as only the total number of molecules of the acid within the unit volume is

counted for concentration

Concentrated Acids:

- A concentrated acid is when there are a large number of molecules of acid per unit volume of the solutionDilute Acid:

- A dilute acid is when there are only a few molecules of acid per unit volume of the solution Use available evidence to model the molecular nature of acids and simulate the ionisation of strong

and weak acids

Identify pH as and explain that a change in pH of 1 means a ten-fold change in [H+]- The pH scale runs on alog to the base ten scale, which can also be written as 10-pH = [H+]- Since there is a base of ten, when the pH of a solution changes by a value of one, the hydrogen ion concentration

of the solution changes by a multiple of ten

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

- The pH of strong acids will always assume 100% ionisation of hydrogen ions into solution- The only difference between the concentration of the acid and the concentration of the hydrogen ion sis

whether or not the acid is polyprotic

- The pH of an acid can be calculated by the formula 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

Acid Concentration (mol/L) pH Strength Degree of Ionisation (%)

Citric 0.1 2.1 Weak 8

Acetic 0.1 2.9 Weak 1.3


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Hydrochloric 0.1 1.0 Strong 100

Summary:

- Hydrochloric is the strongest acid out of the three, ionising completely, followed by citric acid and finally aceticacid being the weakest, with the degrees of ionisation being 8% and 1.3% respectively

Explanation:

- Hydrochloric acid completely ionises so the concentration of hydrogen ions is the same as the concentration ofthe acid

- Citric acid only partially ionises, so the concentration of the hydrogen ions is less than the concentration of theacid, with the lower concentration of hydrogen ions creating a higher pH

- Acetic acid is the weakest acid, with only a very small amount of the acid ionisingComplete Ionisation of Citric Acid:

-

- -

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

Aim:

- To measure the pH of identical concentrations of hydrochloric, acetic and citric acidsMethod:

1. The pH probe was calibrated using buffer solutions2. The pH probe was placed into 0.1 mol/L hydrochloric acid and was left until the pH stabilised3. The pH was recorded4. Steps 2 and 3 were repeated using 0.1 mol/L citric acid and 0.1 mol/L acetic acid5. The results were collected from two other groupsResults:

Groups

pH

Hydrochloric Acid Acetic Acid Citric Acid

1 1 2.72 1.99

2 1.07 2.89 2.09

3 1.1 2.77 1.99

Average 1.06 2.79 2.02

Analysis:

- Although each of the acids has the same concentration, there is a different pH


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- This indicates that there are different concentrations of hydrogen ions in each solution, so there is a varying levelof ionisation between the acids

- It can be seen that hydrochloric acid ionised the most, being the lowest pH and therefore having the highesthydrogen ion concentration

- After that comes citric acid, and then acetic, with acetic being the weakest acid out of the threeDiscussion:

Calibration of the pH Probe:

- The pH meter was calibrated to ensure that it was measuring the correct value of pH,- This was done by first cleaning the probe with distilled water- The data logger was then set to calibrate, and placed into a buffer solution of known pH- When the voltage value on the screen stopped fluctuating, the pH was typed in and this was repeated with

different pH solutions

- This validates the results as it ensures for correct measurementsConclusion:

- The pH of the same concentrations of hydrochloric, acetic and citric acids were measured and compared Describe the difference between a strong and a weak acid in terms of an equilibrium between the

intact molecule and its ions

- When an acid ionises, an equilibrium reaction is formed between the intact molecule and the ionised formsStrong Acids:

- In a strong acid, the molecules in the acid ionise completely, and are unlikely to recombine to reform the intactmolecule

- The reaction is therefore unidirectional, as the equilibrium is almost completely shifted to the right, and can berepresented by the unidirectional arrow in the ionisation reaction,

Weak Acids:

- In a weak acid, only very small amounts of the intact molecules ionise to form ions- These ions can also recombine to reform the complete molecule fairly often, and thus forms a dynamic

equilibrium

- This reaction is therefore reversible, occurring in both directions and is therefore represented using anequilibrium arrow in the ionisation reaction,

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

Aim:

- To use a pH probe and indicator to determine the acidity, basicity or neutrality of common household chemicalsMethod:

1. Beakers of each solution were prepared (for shared use)2. The pH probe was placed into each beaker and the value of the pH was recorded once it stopped fluctuating3. Some of the solution from each beaker was extracted into separate test tubes and universal indicator added,

noting colour change and determining the pH using an indicator chart


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Results:

Sample pH Meter Reading Universal Indicator Colour/pH Acidic/Basic/Neutral

Ammonia Solution 9.70 Green/7.5 Basic

Lemon Juice 2.45 Red/


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Preservatives:

- Acids can be used as a preservative as they lower the pH of foods-

This is significant as many micro-organisms and bacteria can only survive in a specific pH range, and if thesolution is too acidic or basic they may die, which makes adding acid an effective preservative

Flavour Enhancer:

- Acids are generally described to have a sour taste, as can be observed form the acetic acid in vinegar and sitricacid in citrus fruits making them sour

- These acids can be added to food as additives to enhance flavour, such as acetic acid used to make sweet andsour sources to create an overall pleasing combination of flavours

Identify data, gather and process information from secondary sources to identify examples ofnaturally occurring acids and bases and their chemical composition

Acids:

- Citric Acid (C3H5O(COOH)3) occurs naturally in citrus fruits such as oranges, lemons and limes- Hydrochloric Acid (HCl) occurs naturally in the stomachs as stomach acid to break down food- Carbonic Acid (H2CO3) occurs naturally as carbon dioxide in the atmosphere dissolves into rain- Ethanoic Acid (CH3COOH) occurs naturally in the oxidation of ethanolBases:

- Calcium carbonate (CaCO3) occurs naturally in rock such as limestone and marble, and in shells of shellfish- Ammonia (NH3) occurs naturally in volcanoes and the decomposition of plant matter


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

been refined

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

Early Definitions:

Grouping by Properties:

- The earliest definition of an acid was based on their properties and the effect they had on an indicator- They were deemed to have a sour taste and to turn red litmus blue- This was an early attempt to group acids, however there was little explanation about composition or how the

acids worked

Antoine Lavoisier (1779):

Grouping by Composition (Acids Contain Oxygen):

- Lavoisier discovered that some non-metal oxides that dissolved in water could form acids- He also noted that acids such as nitric, sulfuric, carbonic and acetic acids also contained oxygen- From these observations he concluded that acids were substances which contained oxygen- Thus he attempted to group acids by composition rather than properties

Problems:

- However, there were many problems with his theory, as not all oxides produce an acidic solution when dissolved,with metal oxides producing basic solutions instead

- It also limited acids to only substances containing oxygen, not accounting for later discovered acids that did notcontain oxygen such as hydrochloric acid

Sir Humphrey Davy (1810):

Grouping by Composition (Acids Contain Hydrogen):

- Davy first noticed that some metals can react with acids to displace hydrogen- Davy demonstrated that hydrochloric acid did not have oxygen, and thus disproving Lavoisiers theory- He stated that instead of oxygen, acidic substances contained hydrogen- This lasted for a long time, as even more acids were discovered that didnt contain oxygen (HF, HBr, HCN and H2S)

Problems:

- However his theory still did not provide an explanation on the behaviour of acids, and it also did not explain whysome compounds that contained hydrogen were not acidic, such as ammonia

Svante Arrhenius (1887):


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Grouping by Atomic Behaviour (Acids Ionise to Create Hydrogen, Bases Ionise to Create Hydroxide):

- Both Davy and Lavoisiers theories were based on observable properties, but Arrhenius proposed a definitionbased on concepts about particles that were too small to be observed

- He proposed that acids were substances that ionise to produce hydrogen ions in water, - He also proposed that bases were substances that ionised in water to produce hydroxide ions,

- This definition was beginning to explain the behaviour of acids and basesProblems:

- However his definition only applied to aqueous solutions, and also did not address the relative strengths ofdifferent acids and bases

- It does not account for insoluble metal hydroxide compound which react with acids, as they are not in solution- It also provides no explanation on the varying strengths of acids and bases, and cannot explain the acidity or

basicity of some salts when they dissolve in water

Outline the Brnsted-Lowry theory of acids and bases- The Brnsted-Lowry theory of acids and bases states that acids are proton or hydrogen ion donors, and bases

are proton or hydrogen ion acceptors

- In each reaction between an acid and a base, the proton is directly transferred from the acid to the base- When an acid dissolves in water, the water molecule acts as a base and accepts the proton to form hydronium- When a base dissolves in water, the water molecule acts as an acid and donates a proton to form hydroxide- This theory, unlike others, is not based on composition and works regardless of the solvent, with the acids and

bases not having to be in an aqueous state

Gather and process information from secondary sources to trace developments in understandingand describing acid/base reactions

Classification by Properties:

- At first, when acids and bases were grouped only by properties, little was known about their behaviour and whathappened during acid base reactions

- Acids were just classified as a substance that reacted with a base, and a base was a substance that reacted withan acid and these reactions produced a salt and water, but how this reaction occurred was not understood

Arrhenius Definition:

- The first attempts to classify acids and bases by behaviour were by Arrhenius, who stated that an acid produceda hydrogen ion in solution and a base produced a hydroxide ion

- This began to lead to an understanding of how acid base reactions worked, where the hydrogen ion from theacid and the hydroxide ion in the base combined to form water, with the remaining ions forming the salt

- However, this definition for acids and bases was very limitedBrnsted-Lowry Definition:

- A greater understanding of acid base reactions only arose with the Brnsted-Lowry theory of acids and bases, inwhich an acid was a proton donor and the base a proton acceptor

- This describes that in an acid base reaction, the proton is donated directly from the acid onto the base, allowingfor the reaction to occur

- This has allowed for an understanding of acid and base reactions Describe the relationship between an acid and its conjugate base and a base and its conjugate acid

Acid and its Conjugate Base:


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- When an acid donates its proton, it forms the acids conjugate base, forming a conjugate pair

Base and its Conjugate Acid:

- Conversely, when a base accepts a proton, it forms its conjugate acid and so forms another conjugate pair

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

Salts:

- When an acid reacts with a base in a neutralisation reaction, it forms a salt and water- However, depending on the strengths of the acid and base, the salt is not necessarily neutral- Some salts will hydralise (react with water to create a change in pH) in water, to either lower the pH to make the

solution acidic, or raise the pH to make the solution basic

- These are known as acidic and basic salts respectively, with neutral salts not altering the pH in any wayAcidic Salts:

- Acidic salts are formed when a strong acid reacts with a weak base to form a salt- The ions that make up this salt will usually be comprised of a weak conjugate base, and a strong conjugate acid- The strong conjugate acid will then donate its proton to produce hydronium ion sin water, and thus creating an

acidic solution, thereby making the salt acidic

- Examples of acidic salts include ammonium chloride and ammonium nitrateReactions:

- - Basic Salts:

- Basic salts are formed when a weak acid reacts with a strong base to form a salt- The ions in this salt will usually be comprised of a strong conjugate base, and a weak conjugate acid-

The strong conjugate base will then accept a proton from water to produce hydroxide ions in water, and thuscreating a basic solution, thereby making the salt basic

- Examples of basic salts include sodium ethanoate, sodium carbonate and potassium nitrateReactions:

- -

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

Aim:

- To determine the pH of a range of salt solutionsEquipment:


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- pH probe- Data logger- 0.1M salt solutions

- Test tube rack- Test tubes- Universal indicator

Method:

1.

2-3cm of 0.1M solution was poured into a test tube2. 4 drops of indicator was added to the test tube and the colour change was recorded3. The pH probe was placed into the beaker of 0.1M salt solution4. The reading of the pH was allowed to settle and was then recorded5. This was repeated with all of the solutionsResults:

Salt Solution Universal Indicator (pH) pH Probe Reading

Ammonium Nitrate (NH4NO3) Light Green 7.0 7.04

Sodium Phosphate (Na3PO4) Dark Green 7.5 6.40

Potassium Iodide (KI) Turquoise 8.5 8.50

Sodium Acetate (CH3COONa) Turquoise 8.5 8.74

Barium Chloride (BaCl2) Turquoise 8.5 9.61

Sodium Nitrate (NaNO3) Dark Green 7.5 7.05

Magnesium Oxide (MgO) Purple 10.0 10.46

Sodium Sulfate (Na2SO2) Dark Green 7.5 6.40

Sodium Hydrogen Carbonate (NaHCO3) Turquoise 8.5 9.07

Ammonium Chloride (NH4Cl) Light Green 7.0 7.19

Analysis:

- The salt solution varied in pH which is seen in the results, which indicates that not all salts are neutral- A hydrolysis reaction between the water and the ions that make up the salt result in the change in pH that is

seen in the results

Discussion:

pH of Salts:

- From the results it can be seen that salts can be acidic, basic or neutral- An acidic salt is formed between a strong acid and a weak base, a basic salt can be formed from a weak acid and

a strong base and neutral salts are formed from strong acids and strong bases

- Salts formed from weak acids and bases vary in pH as the strengths of the acids and bases vary, which alters thepH of the salt

Conclusion:

- The pH of a range of salt solutions was determined to show that they can be acidic, basic or neutral Identify conjugate acid/base pairs

Acid Conjugate Base Base Conjugate Acid

HCl Cl- OH- H2O

HNO3 NO3-

NH3 NH4+

H2SO4 HSO4-

CN-

HCN

NH4+

NH3 CO32-

HCO3-

H2O OH-

H2O H3O+

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


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- Amphiprotic substances can act as both an acid and a base, in other words, can either donate or accept a protonAmphiprotic Substances Acting as a Base (Proton Acceptor) Acting as an Acid (Proton Donator)

Hydrogen Carbonate (HCO3-)

Water

Other Examples:

- Hydrogen sulfate (HSO4-)- Dihydrogen phosphate (H2PO4-)- Hydrogen phosphate (HPO42-)

Identify neutralisation as a proton transfer reaction which is exothermic- According to the Brnsted-Lowry definition of acids and bases, a neutralisation reaction is a reaction between an

acid and a base, where the acid donates a proton and the base accepts

- Therefore this can be classified as a proton transfer reaction, as the proton is transferred from the acid to thebase

- This is an exothermic reaction, in that it releases energy in the form of heat, as the amount of energy releasedwhen the bonds form is much greater than the amount of energy required to break the bonds

- The amount of heat released is approximately -56kJ/mol, with this value varying slightly depending on thestrengths and concentrations of the acids and bases

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

Assessment:

- Neutralisation reactions are an effective method as a safety measure or to minimise the damage caused bychemical spills, especially acids and bases

- These can turn spills of potentially deadly corrosive or caustic chemicals into harmless water, salt and in somecases, carbon dioxide

Reasons for Use:

- Neutralisation is used as a safety measure in both large and small scale acid and base spills, both indoor andoutdoor, as well as for dumping industrial acidic and basic chemical wastes

- This is done to reduce the damage to the surrounding environment, as a change in pH of soil or aquatic systemscan be devastating to wildlife

- The chemicals that need to be neutralised are also often highly dangerous to people as it can cause chemicalburns, and neutralisation is used as a safety measure

Substance Used:

- Substances that are used to neutralise chemical spills are usually a weak amphiprotic substance in a powderedform, such as sodium hydrogen carbonate

- This is to reduce the heat released from the neutralisation as it is an exothermic reaction, as well as to reducethe spread of the chemical caused by adding a liquid to neutralise it

- An amphiprotic substance is useful as only a single substance is required for all situations, and can be used toneutralise solutions were acidity or basicity is unknown


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- Sand can also be used to first soak up any liquid in the spill to prevent it spreading, and then the sand can beneutralised

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

- Titrations and the preparation of standard solutions are used in the quantitative analysis of a substance and anyerrors in the analysis can lead to an incorrect and invalid result

- Therefore the method used to conduct these analyses are vitally important to ensure that the results areaccurate and valid

- Titrations involve first cleaning the equipment, preparation of standard solutions and then the titration itself

Cleaning:

Initial Rinsing:

- To ensure that there are no impurities that will alter the results in any way, all the equipment used must first bethoroughly cleaned, including burettes, pipettes, conical flasks, volumetric flasks, funnels and beakers

- All of the components must be first thoroughly rinsed with distilled water- They should, but do not have to be dried as this will not alter the results of the titration

The ette Rule:

- Finally, any piece of equipment ending in ette must be given a final rinse with the solution that it will contain,as they are used to measure volume of a substances where concentration is important

- Any water that is leftover in the equipment will alter this concentration and invalidate the result, so therefore itmust first be rinsed with the solution it will hold, as any residue will not alter the concentration

Preparation of a Standard Solution:

Summary:

- A standard solution is a solution of known concentration that can be used to determine the concentration ofanother solution by titration

- The easiest way to prepare a standard solution is by using a primary standard which can be directly dissolvedinto a known volume with the concentration being determined by a simple calculation

Properties of a Primary Standard:

- A high level of purity (must not be reactive or have reacted with other substances)- Accurately known composition (must not vary in composition such as holding crystals of water)- Free of moisture- Stable and unaffected by air during weighing- Readily soluble in distilled water- High molecular weight solid to reduce the percentage error in weighing- React instantly and completely

Procedure:


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- The solution is prepared by first zeroing a set of electronic scales with a piece of filter paper placed on it, themore precise the better

- Small amounts of the primary standard are added onto the filter paper, ensuring it is all on the paper, until thedesired mass is measured

- This is then carefully transferred into a beaker with the paper, ensuring none of the solid is lost, and is thendissolved with minimum distilled water

- The solution is transferred into the volumetric flask of a desired volume via a funnel to ensure none is lost- A distilled water squirt bottle is used to rinse the sides of the beaker, and the washings were then transferred

into the volumetric flask

- This was repeated a few times to ensure that all of the primary standard was transferred into the flask- The funnel is then rinsed thoroughly, allowing the washings to run into the flask, and is then removed- The volumetric flask is then filled to the neck by pouring distilled water into the flask- At the neck, an eye dropper is used to add water into the flask to prevent accidental overfilling, running the

water down the side of the neck to prevent any loss from splashing

-

The flask is filled until the bottom of the meniscus reaches the etched mark, placing a rubber stopper on thevolumetric flask and inverting multiple times to ensure the concentration is uniform throughout the solution

- The concentration is then calculated from the measured mass and the volume of the solution, and written onthe flask for future reference and use

Using a Secondary Standard:

- If a primary standard is not used, the mass of the substance added does not need to be measured- These do not possess the properties required to be classified as a primary standard, whether it be impure or

hygroscopic

- In order to make these into a standard solution, they must first be dissolved and titrated with a differentstandard solution in order to determine its concentration

- Once the concentration is known, it can then be used as a standard solution and is termed a secondary standardChoosing and Indicator:

- An indicator must first be chosen appropriate to the titration, making sure the end point of the indicator whereit just begins to change colour is very close to the equivalence point of the reaction so there is very little volume

difference between the endpoint and the equivalence point

- If the titration is between a strong acid and a strong base, the equivalence point will be close to neutral, somethyl orange, bromothymol blue and phenolphthalein are all suitable

- If it is a strong acid weak base titration, the equivalence point will be slightly acidic and hence phenolphthalein isunsuitable for this titration, with only bromothymol blue and methyl orange being suitable

- If it is a weak acid and strong base titration, the equivalence point will be slightly basic, and only phenolphthaleinis suitable

- Weak acid and weak base titrations are generally not performed due to the difficulty in determining theequivalence point

- Phenolphthalein is generally the most preferred indicator as it has the most distinct colour changeTitration:

Summary:

-

After the standard solutions have been prepared, the titration is ready to be performed- The titration being performed is an acid base titration, where and acid will react with a base until they have both

completely reacted


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- This point is known as the equivalence point, where all the acid and all of the base have stoichiometricallyreacted

- Terms that are often used in titrations are aliquot, which is the known volume in the conical flask, and titre,which is the variable volume being added via the burette

- It is generally not important which of the substances is the aliquot or the titre, but sodium hydroxide is generallynot use as the titre as it is able to dissolve glass, which can change the diameter of the burette, making the

measurements inaccurate

Preparing for the Titration:

- Firstly, a known volume of one of the substances is added to the conical flask via a volumetric pipette, leaving adrop at the end of the pipette after dispensing the solution as the pipette is calibrated for this

- A few drops of the selected indicator is then added into the solution in the conical flask, and a burette is set upusing a retort stand and clamps

- The burette is filled with the second solution after it is rinsed with it, ensuring that the tap is closed-

The tap is then opened for a short time to allow the solution to completely fill the end of the burette and is thenclosed again

- The burette does not have to be filled to the zero mark, as long as the initial volume in the burette is recorded- Finally the conical flask is placed underneath the tip of the burette, and the initial value of the burette is read,

ensuring to read off the bottom of the meniscus

Performing the Titration:

- The tap is opened to allow the solution to flow into the conical flask, ensuring that the flask is constantly beingswirled

- Any solution that is splashed onto the side of the flask can be washed down with distilled water as this does notchange the number of moles of substance in the flask

- This process continues until a colour change occurs in the conical flask, to indicate that the reaction has finished- The tap is then closed and the final value in the burette is recorded, with this first recording being known as a

rough run

- The process is then repeated as many times as desired to establish reliability, slowing the speed of the burettewhen it is about 5mL away from the value obtained in the rough run

- The adding solution is then slowed until it is drop by drop, and if there is a drop hanging off the end of theburette it can be washed off using distilled water, as this drop has already been counted as dispensed

- Ensuring that the flask is being swirled constantly, the solution is added from the burette until the instant thesolution changes colour, and then the amount is recorded

- Finally, if the values of the results are consistent, they are averaged and used in calculations to determine theconcentration of the unknown substance, thus completing the titration

Rough Run:

- The first titration is known as a rough run, as only a close value to the equivalence point is achieved due to nothaving an idea of how much solution must be added to reach the equivalence point, making it very easy to

overshoot this point

- This value is generally not used in any calculations, and will usually be a larger volume to the amount required toreach the equivalence point

Alternate Method using Technology:


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- An alternate method instead of using indicator alone is by using a data logger and a pH probe- The data logger can be used to generate a pH graph, and there will be a steep portion of the graph where the pH

changes rapidly

- The midpoint of this steep section will be the equivalence point, and can then be used to see how much volumewas added at this equivalence point which can then be used for calculations

Perform a first-hand investigation and solve problems using titrations and including the preparationof standard solutions, and use available evidence to quantitatively and qualitatively describe the

reaction between selected acids and bases

Aim:

- To prepare a standard solution to use in a titration to determine the concentration of Diggers HCl ConcentrateMethod:

- The method above was used to prepare the standard solution and the titration, with the primary standard usedbeing sodium carbonate

- The sodium carbonate was the aliquot, and the HCl was diluted 1 in 25 using a volumetric flask, and was the titre- When diluting the acid, it was ensured that the acid was added to excess water- The indicator chosen was methyl orange as the result was an acidic salt, and the colour of the endpoint of the

titration was peach

Results:

- The concentration of the standard solution was 0.150202849mol/L- On average, 23.267mL of HCl was required to react with 25mL of 0.150202849mol/L sodium carbonateAnalysis:

- HCl reacts with sodium carbonate in a 2:1 ratio, so twice as many moles of hydrochloric acid would be consumedthan sodium carbonate

- First the number of moles of sodium carbonate in 25mL of the standard solution was calculated- This was then multiplied by 2 to give the moles of HCl- This was divided by the average volume of the titre to provide the concentration for the diluted HCl- Finally this value was multiplied by 25 to give the value for the concentration of the concentrated HClDiscussion:

Selection of Indicator:

- The salt formed in the titration is between a strong acid and a weak base, which makes an acidic salt- Thus the selection of methyl orange as an indicator was appropriate as the equivalence point would have been

acidic, being in the same range as the end point of methyl orange

Adding Acid to Water:

- When concentrated acid like the concentrated HCl dissolved with water, the reaction is highly exothermic- If small amounts of water were to be added to the acid, the resulting reaction would vaporise the water and the

acid, being highly dangerous

- Adding the acid to excess water ensures that there is enough water to absorb this heat, to ensure no harmfulgases are released


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Conclusion:

- A standard solution was successfully prepared to be used in a titration which determined the concentration ofDiggers HCl concentrate to be 294.2927369g/L

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

Aim:

- To determine the concentration of a domestic acidic substance using computer based technologiesMethod:

- The method above for the titration with the data logger was used- This time, NaOH was used as the standard solution, being the titre, with the unknown concentration of HCl being

aliquot

- Phenolphthalein was used as the indicatorResults:

- The volume added add and the pH of the solution was recorded many times using the data logger in theexperiment, and was graphed in a pH vs volume of NaOH added graph

- The midpoint of the steep portion was used to be the equivalence point, with 54.7mL of 0.2489mol/L NaOHneeded to completely react with the 50mL of diluted HCl

Analysis:

- The same calculation as used as before except the acid base ration was 1:1

Discussion:

Use of Data Logger:

- Using the data logger and pH probe allowed for much more accurate and precise results as it is purelyquantitative analysis, rather than a portion of it being qualitative like the traditional titration method, where bias

can occur in determining if the colour had changed or not

- This method eliminates this bias, which can further validate the resultsConclusion:

- The concentration of the Diggers HCl concentrate was determined to be 294.0812829g/L using computer basedtechnologies

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

- Buffers are substances which prevent a change in pH in a solution- They usually consist of a weak acid and its conjugate base, or a weak base and its conjugate acid- These form equilibrium reactions which due to Le Chateliers principle will counteract an addition of an acid or

base to maintain a stable pH level


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Example (Carbonic Acid and Hydrogen Carbonate Ions):

- An example of a buffer would be the buffer formed between carbonic acid, a weak acid, and its conjugate base,hydrogen carbonate ions, which are usually present in equal amounts

Equations:

1. 2.

Nature of the Buffer:

- In the first equation, the hydrogen carbonate ion is a stronger base than water, as well as the hydronium ionbeing a stronger acid than the carbonic acid

- Likewise in the second reaction, the carbonic acid is a stronger acid than water, and the hydroxide ion is astronger base than the hydrogen carbonate ion

- Whilst the system is at equilibrium, due to the nature of the conjugate acid and base pairs, both the reactionswill tend to the left

Adding an Acid or a Base:

- When an acid is introduced into the system, the increase in hydronium ions will shift the first reaction to the left,producing unionised carbonic acid and water, thus removing most of the acid from the system

- Likewise when a base is introduced into the system, the increase in hydroxide ions will shift the second reactionto the left, with the carbonic acid transferring a proton directly onto the hydroxide ion to produce a hydrogen

carbonate ion and water

- It is due to these reactions quickly removing any hydronium or hydroxide ions introduced into the solution thatthey make effective buffers and are able to maintain the pH at a relatively stable level

Significance to Living Organisms:

- In living organisms, all metabolic reactions responsible for life are catalysed by biological catalysts known asenzymes

- These enzymes are protein globules that can only work in a very specific pH range, with a pH that is too high ortoo low changing the shape of the enzyme and destroying or denaturing it

- Therefore, buffers must be present in natural systems in order to prevent changes in pH that will risk destroyingthese enzymes and causing harm to the organism

Significance of the Carbonic Acid/Hydrogen Carbonate Buffer to Living Organisms:

- Carbon dioxide is a substance that is present in all cells, produced as a waste product from cellular respiration- Too much carbon dioxide is toxic for an organism, so it must be removed from the organism- This can only be done by dissolving into the blood plasma in the bloodstream, creating carbonic acid and the

hydrogen carbonate ions, creating a buffer system

- This system keeps the pH of blood at a constant level of around 7.4, and is important as it allows the enzymes towork at their most efficient rate at their optimum pH, thereby allowing the organism to survive


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Other Examples:

- Other examples of buffers include ethanoic acid and the ethanoate ion which maintains the pH at around 3.5 to5.5 and the ammonium ion and ammonia which maintains the pH at around 8.3 to 10.3

Limitations of Buffers:

- Obviously, buffers do have limitations, and will cease to work if excess amounts of acid or base are added to thesolution

- The excess amounts will drive the reactions to completion, and with the reagents which balance the pHcompletely depleted, the pH will no longer be buffered if the acid or base were to continue to be added


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

Describe the differences between the alkanol and alkanoic acid functional groups in carboncompounds

Alkanols:

Structure:

- Alkanols are carbon chains that contain a hydroxyl OH attached onto it- The homologous formula for an alkanol is CnH2n+1OH


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Properties:

- Due to the hydroxyl group at the end, this makes the molecular polar- The hydrogen bonded to the oxygen also allows for hydrogen bonding, and this allows the alkanols to be soluble

in water

- However as the non-polar carbon chain increases, the solubility in water decreases as the dispersion forces thatthe alkanol molecules have with each other overpower the hydrogen bonds formed with the water molecules

- Due to the polar nature, alkanols have higher melting and boiling points than corresponding alkanes- Alkanols are neutral as all the hydrogens are strongly bonded to the molecule, and it cannot accept any

hydrogens either

Alkanoic Acids:

Structure:

- Alkanoic acids have a carboxyl group at the end of the alkyl chain- This consists of a hydroxyl group like the alkanol, as well as a double bonded oxygen- The homologous formula is Cn-1H2n-1COOH

Properties:

- The extra oxygen as a strong partial negative charge, which increases the intermolecular forces between themolecules, making the melting and boiling points higher than the alkanols

- Being also a polar molecule, it is soluble in water, but as with alkanol, as the non-polar chain increases, solubilitydecreases

- Alkanoic acids are weak acids as they can donate the hydrogen at the end of the hydroxyl group Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained

alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8- Esters produced from straight-chained alkanoic acids and straight chained primary alkanols are named based on

the alkanol and the alkanoic acid that fromed the ester

- The alkanol is placed first in the naming, followed by the alkanoic acid- The prefixes are taken of both (meth, eth, prop, but, pent, hex, hept, oct), the suffix yl is placed after the prefix

of the alkanol, and the suffix oate is placed after the prefix of the alkanoic acid

- Therefore, an ester formed from butanol and ethanoic acid would be called butyl ethanoate

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

- The difference in melting and boiling points of alkanoic acids and alkanols can be explained by their structures- In an alkanol, only hydrogen bonding and dispersion occurs between molecules- However, due to the extra oxygen in the carboxyl group in alkanoic acids, this allows for a third type of bonding

to occur, dipole-dipole bonds- This increases the intermolecular forces between the alkanoic acids as compared to the alkanols, therefore

creating a higher melting and boiling point


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- Equal length carbon chains of alkanoic acid also have larger molecular masses than alkanols, which increase thedispersion forces, further increasing the strength of the intermolecular forces in the alkanoic acid

Hydrogen Bonding in Alkanols: Hydrogen Bonding + Dipole-Dipole bonding in Alkanoic Acids:

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

The Reaction:

- The process of esterification is the reaction between an alkanoic acid and an alkanol, which is a condensationreaction that is catalysed by heat and concentrated acid (usually sulfuric)

- The general equation is by - As seen, this is an equilibrium reaction which is catalysed and shifted to the right by heat and concentrated

sulfuric acid

- The heat provides sufficient energy in order for the particles to collide and react (collision theory) and thesulfuric acid catalyses the reaction, as well as shifting the reaction to the right being a dehydrating agent and

removing water from the system

Example (Butyl Ethanoate):

Chemical Equation:

- Word Equation:

- Structural Equation:

Describe the purpose of using acid in esterification for catalysts- The purpose for using the concentrated sulfuric acid is that it catalyses the reaction because it reduces activation

energy of the reaction by donating a proton to the unshared electron pairs of the alkanoic acid or the alkanol, toincrease the rate of reaction

- Sulfuric acid is also a strong dehydrating agent which removes water from the system, and causing the reactionto shift to the right


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- This shifting to the right will then increase the yield of the ester, so the use of concentrated sulfuric acid not onlyincreases the reaction rate, but also the yield of the reaction

Explain the need for refluxing during esterification- The reactants and the ester produced are all volatile liquids, so they change into gaseous state easily and

spontaneously

- This rate of vaporisation is only increased by the heat that is applied to catalyse the reaction- Therefore refluxing is required in esterification to recondense the ester back into liquid form to reduce the

amount of ester lost due to its volatility

- Refluxing is the process of attaching a condenser tube above the reaction flask in order to condense any vapoursthat leave the flask to allow it to return

- This allows minimal loss of the ester as most of it is condensed and returned to the flask, and therefore isrequired in esterification or else all of the ester would escape in the form of vapours

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

Aim:

- To prepare an ester using refluxEquipment:

- Retort stand- Bossheads- Clamps- Condenser tube-

Tubing- Pear shaped flask- Butanol- Ethanoic acid

- Concentrated sulfuric acid- Bunsen burner- Tripod- Heat mat-

Gauze- Water- Beaker

Method:

1. 7mL of 1-Butanol, 10mL of ethanoic acid, 5mL of concentrated sulfuric acid and 4 boiling chips were added to thepear shaped flask

2. The equipment was set up ensure that the glassware was securely clamped so it cannot fall3. The Bunsen burner was lit and allowed to heat the mixture for at least 30 minutes4. After the heating was finished, the setup was allowed to cool5. The flask was carefully detached and the contents poured into a separatory funnel6. 25mL of water was added to the funnel and it was shaken, draining off the lower aqueous layer after it had

settled, which was repeated twice

7. Finally 15mL of saturated sodium hydrogen carbonate and 15mL of water was added to the funnel, shaken anddrained

8. The ester was extracted into a small beaker

Results:

- Before refluxing, both the butanol and the ethanoic acid had foul smells


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- After refluxing however, the ester had a sweet smell of banana- When sodium hydrogen carbonate was added to the funnel, fizzing was apparentAnalysis:

-

The mixture fizzed when the sodium hydrogen carbonate was added as a neutralisation reaction occurredbetween any remaining acid and the sodium hydrogen carbonate, which released carbon dioxide gas

Discussion:

- The sodium hydrogen carbonate purifies the ester, as it is in an equilibrium reaction and there will still be someacid leftover in the acid

- Also the sulfuric acid used to catalyse the reaction would also be present in the mixture, so the sodium hydrogencarbonate neutralises this to form an aqueous layer which can be drained off

Conclusion:

- The ester butyl ethanoate was produced by refluxing Outline some examples of the occurrence, production and use of esters

Ester Alkanol Alkanoic Acid Smell Use

Butyl Ethanoate Butanol Butanoic Acid Banana Industrial solvent, synthetic

fruit flavouring

Butyl Methanoate Butanol Butanoic Acid Raspberry Flavouring and fragrance

Ethyl Ethanoate Ethanol Ethanoic Acid Sweet, Pear Drops Glue, nail polish remover,

cigarettes

Ethyl Butanoate Ethanol Butanoic Acid Pineapple Flavouring, solvent

Benzyl Acetate Benzyl Alcohol Acetic Acid Jasmine Natural in flowers, flower oils,

baiting bees

Methyl Phenylacetate Methanol Phenylacetic Acid Honey Flavouring, perfume

- All of these esters are produced industrially except for benzyl acetate which occurs naturally in flowers Process information from secondary sources to identify and describe the uses of esters as flavours

and perfumes in processed foods and cosmetics

- Esters occur naturally in a range of perfumes and flavourings in flowers and fruit- Particular esters responsible for scent or taste can be identified and then reproduced on an industrial scale for

use as flavours and perfumes in processed foods and cosmetics

Examples:

- Examples of these can be seen in the table above- In processed foods such as candy and ice cream, the ester is used as flavouring and can also provide a fragrance

to the food, such as banana flavours lollies

- In cosmetics, flavoured lip gloss makes use of esters for flavouring and perfume, as well as many creams andointments with fragrances coming from esters

9.3 the acidic environment

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