Module 6: Acid & Base Reactions

Module 6: Acid & Base Reactions 1

Module 6: Acid & Base ReactionsProperties Of Acids & Bases

📖

investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases ✅

Naming Inorganic Acids

Naming Inorganic Bases

Bases, unlike acids, are named using rules for naming ionic compounds.

Common Acid & Their Properties

The most common types of acids and their uses include:

Hydrochloric Acid ( )

Present in stomach to break down proteins. Cleaning agents for bricks and concrete. Production of a range of other chemicals

Sulfuric Acid ( )

Used in car batteries. Use in manufactures of fertilisers, dyes and detergents.

Nitric Acid ( )

HCl

H SO2 4

HNO3


Used in manufactures for fertilisers, dyes explosives.

Ethanoic Acid / Acetic Acid ( )

Found in vinegar. Used in food preservative. Manufactures of glues and plastics

Carbonic Acid ( )

Used to carbonate soft drinks and beer

Phosphoric Acid ( )

Used as flavouring and food preservative. Used to produce fertilisers and pharmaceutical products.

Citric Acid ( )

Found in citrus fruits. Used as a flavouring and food preservative.

Ascorbic Acid / Vitamin C ( )

Found in citrus fruits and vegetables. Used as a health supplement. Used as an antioxidant in food preperation.

Common Bases & Their Properties

The most common types of bases and their uses are:

Sodium Hydroxide ( )

Drain and oven cleaners. Soap making. Production of other chemicals

Ammonia ( )

Household cleaners. Production of fertilisers, explosives and plastics

Calcium Hydroxide ( )

Cement and mortar. Garden lime. Food preperation

Magneisum Hydroxide ( )

Antacids such as milk of magnesia to treat indigestion

Sodium Carbonate ( )

Manufactures of washing powders, soaps, glass paper. Water softener.

Sodium Hydrogen Carbonate ( )

Neutralising agent for spills. Baking. Antacids.

📖

conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions ✅

Properties Of Acids

CH COOH3

H CO2 3

H PO3 4

C H O6 8 7

C H O6 8 6

NaOH

NH3

Ca(OH)2

Mg(OH)2

Na CO2 3

NaHCO3


Common properties of acids include:

Corrosive

Turns blue litmus red

Reacts with bases

Solutions have pH below 7

Tastes sour

Solutions are electrically condutive

Properties Of Bases

Common properties of acids include:

Caustic and slippery

Turns red litmus blue

Reacts with acids

Solution have pH above 7

Tastes bitter

Solutions are electrically conductive

Chemical Indicators

Indicators are substances which change colour according to the pH of the substance they are immersed in. Each indicator has its own range in pH where, within that range, it will change from one colour to another.

By knowing the pH colour change range of indicators, it is possible to determine that a substance is above or below a specific pH, and as a result, whether they are acidic basic and to what degree.

Common indicators and their ranges include:

Litmus (Red - Blue for 4.8 - 8.1)

Phenolphthalein (Colourless - Pink for 8.0 - 10.0)

Methyl Orange (Red - Yellow for 3.1 - 4.4)

Bromothymol Blue (Yellow - Blue for 6.0 - 7.6)


📖

predict the products of acid reactions and write balanced equations to represent: – acids and bases ✅ – acids and carbonates � – acids and metals �

Neutralisation Reaction

When an acidic solution is combined with a basic solution, a neutralisation reaction occurs where the solutions are said to be neutralised - concentrations of hydronium and hydroxide ions within the mixture are equal.

Acid + Metal Hydroxide Reaction

When an acid reacts with a metal hydroxide (common base), it produces salt and water. Bases that contain hydroxide ions will accept the protons from an acidic solution to form water.

Example: solutions of phosphoric acid and potassium hydroxide react to form a potassium phosphate salt and water

Acid + Ammonia Reaction

When an acid reacts with ammonia, it only produces an ammonium salt.

Acid + Metal Hydroxide → Salt + Water

H PO + 3 4 (aq) 3KOH → (aq) K PO + 3 4 (aq) 3H O 2 (l)


Example: solutions of hydrochloric acid and ammonia react to form ammonium phosphate

Acid + Metal Carbonate Reaction

When an acid reacts with a metal carbonate or a metal hydrogen carbonate (bicarbonate), it produces a salt, water and carbon dioxide.

Example: solution of sulfuric acid and solid sodium carbonate react to form sodium sulfate, water and carbon dioxide.

Acid + Metal Reaction

When an acid reacts with a reactive metal, ir produces a salt and hydrogen gas.

Example: solution of hydrochloric acid and solid iron react to form iron chloride and hydrogen gas.

📖

investigate applications of neutralisation reactions in everyday life and industrial processes ✅

Neutralisation Reactions In Everyday Life

Brushing Teeth

Remnants of food form acid resulting from bacterial action that dissolves the enamel protecting teeth. Tooth paste contains a mix of mild bases such as sodium hydrogen carbonate and sodium fluoride which neutralises the acids in our mouth, protecting our teeth from decay

Antacids

Strong, concentrated hydrochloric acid is present in all our stomachs aiding in digestion. Condition called acid reflux causes acid from stomach to rise into

Acid + Ammonia → Ammonium Salt

H PO + 3 4 (aq) 3NH → 3 (aq) (NH ) PO 4 3 4 (aq)

Acid + Metal Carbonate → Salt + Water + Carbon Dioxide

Acid + Metal Hydrogen Carbonate → Salt + Water + Carbon Dioxide

H SO + 2 4 (aq) Na CO → 2 3 (s) Na SO + 2 4 (aq) H O + 2 (l) CO 2 (q)

Acid + Metal → Salt + Hydrogen Gas

2HCl + (aq) Fe → (s) FeCl + 2 (aq) H 2 (g)


oesophagus causing burning sensation called heartburn. To neutralise the acid, people take antacids containing base magnesium hydroxide or sodium bicarbonate.

Baking Powder

Baking powder is a commonly used ingredient in baking which contains sodium bicarbonate and a weak acid such as sodium aluminium sulfate. When it dissolves in water and is heated, they react together in neutralisation reaction to produce carbon dioxide which causes dough to rise as it bakes.

Treating Wasp & Bee Sting

A wasp venom is basic whilst a bee sting is acidic. The wasp sting can be neutralised by a weak acid like acetic acid and the bee sting can be neutralised by a weak base like sodium bicarbonate.

Pool / Soil Treatment

Soils need to at a specific pH level for optimal crop growth. Lime fertiliser can be used to neutralise acidic soil and rotting compost cna be used to neutralise basic soil. The formation of acidic gases like carbon dioxide can also neutralise the surroudning soil.

Neutralisation Reactions In Industrial Processes

Neutralising Chemical Spills

Chemical spills can occur which need to be safely contained and neutralised. A weak base like sodium carbonate can be used to neutralise an acid spill while a weak acid can be used to neutralise a basic spill.

Fertiliser Production

Production of certain fertilisers involve a neutralisation reaction between feedstock chemicals like the production of ammonium nitrate fertiliser.

Treating Acidic Effluent

Industries that discharge effluent (waste) into waterways are required to ensure they are not harmful to the environment. This involves neutralising acidic discharge.

Latex Production

To prevent the coagulation of latex caused by acids produced by bacteria, ammonium hydroxide is often added to neutralise the acidic culprits.

📖

conduct a practical investigation to measure the enthalpy of neutralisation ✅

Entropy Of Neutralisation

The entropy of neutralisation is the thermal energy released in the reaction when an acid and a base react in stoichiomateric proportions to make a salt and 1 mol of water represented by the symbol, . ΔHneut


Neutralisation reactions are exothermic and has a net ionic reaction involving hydrogen and hydroxide ions reacting together to produce water.

To measure experimentally measure the enthalpy of neutralisation. calorimatery is required. The assumptions considered when measuring enthalpy of neutralisation include:

Salt solution formed has the same heat capacity as water

Salt solution has the same density as water

All the heat energy released by the neutralisation reaction is transferred into the salt solution.

📖

explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to: – Arrhenius’ theory ✅ – Brønsted–Lowry theory �

Antoine Lavoisier Theory

Antoine Lavoisier proposed that acids were substances containing oxygen showing many metal oxide oxide formed aicds when dissolved in water. However, this was disproved as metallic oxides were definitely basic in nature and halogens acids were acidic despite containing no oxygen

Humphrey Davy Theory

Humphrey Davy suggested that acids were substances containing a 'replaceable hydrogen' by showing the production of hydrogen gas from acid - metal reaction. His theory was supported through the discovery of acids that contained hydrogen.

The concept of a 'replaceable hydrogen' is visible when metals react with acid. The metal replaces the hydrogen ion in the acid and metal salt and hydrogen gas is formed. He also suggested acids react with base to form salt and water.

However, he failed to explain the properties of acids.

Svante Arrhenius Theory

Svante Arrhenious suggested that acids where substances that when ionise in water produce hydrogen ions and bases were substances that when ionise in water produce hydroxide ions. He also explained the properties of acids and related acid strength to the degree of ionisation.

The limitation of the Arrhenius theory of acids and bases include:

H + +(aq) OH → −

(aq) H O ΔH =2 (l) neut −55.9 kJ ⋅mol−1


bases such as ammonia and sodium carbonate do not contain a hydroxide group.

insoluble bases could not be classified as bases by this thoery

this theory only applied to aqueous medium and could not explain neutralisation reactions in the gas phase

theory did not explain the important of the role of the solvent

theory did not explain the behaviours of amphoteric substance.

Bronsted Lowry Theory

Bronsted Lowry theory states that acids are proton donors and bases are proton acceptors. Therefore, an acid base reaction involves an exchange of protons from acid to a base. This theory managed to explain behaviour of amphiprotic substances - substances that can donate and accept protons.

Bronsted-Lowry Theory allows the reaction between an acid and base to be generalised. The reaction is considered reversible:

For fowards reaction, HA is an acid since it donates proton to base, B. In reverse reaction, HB+ is an acid that donates proton to base, A-.

The Bronsted Lowry theory gives rise to conjugate acids and bases. When an acid donates its proton, it becomes its conjugate base. Similarly, when a base accepts a proton, it becomes its conjugate acid. Conjugate acids and bases differ by one proton (H+).

Conjugate Acid

In the reaction between an acid and a base, the base accepts a proton (hydrogen ion) and becomes protonated. As the base is now protonated, it has the potential to donate a proton and act as an acid. This is called a conjugate acid.

HA+B⇋ HB ++ A−


A strong base will have a extremly weak conjugate acid. A weak base will have a moderately strong conjugate acid.

Conjugate Base

In the reaction between an acid and a base, the acid donates a proton (hydrogen gas) and becomes deprotonated. AS the acid is now deprotonated, it has the potential to accept a proton and act as a base. This is known as a conjugate base.

A strong acid will have an extremly weak conjugate base. A weak acid will have a moderately strong conjugate base.

Using Brønsted–Lowry Theory

📖

calculate pH, pOH, hydrogen ion concentration ([H+]) and hydroxide ion concentration ([OH–]) for a range of solutions ✅

Ionic Product Of Water

Self-ionisation of pure water occurs when water molecules react with each other to a very small extent. Water displays amphiprotic properties by behaving as a very weak acid and base to produce one hydronium and hydroxide ion.

H O(l) + 2 H O(l) ⇋ 2 H O (aq) + 3+ OH (aq)−


Concentration of hydronium and hydroxide ions are very low. In pure water at 25C, concentrations of both are at mol/L and multiply to give . If either concentration increases then the concentration of the other must decrease proportionally. This result is knows as the ionic product of water ( ).

Acidic & Basic Solutions

In acidic solutions, H3O+ ions are formed by reaction of acid with water and from self-ionisation of water so concentration of H3O+ ions will be greater than mol/L. Since Kw remains constant, concentration of OH- ions in an acidic solution must be less than

mol/L.

Conversly, for an acidic solution, OH- ions are formed by the reaction of bases with water and from self ionisation of water so concentration of OH- ions will be greater than mol/L. Since Kw remains constant, concentration of H3O+ ions in an basic solution must be less than mol/L.

The higher the concentration of H3O+ ions in a solution, the more acidic it is. Similarily, the higher the concentation of OH- ions in a solution, the more basic it is.

The pH Scale

The pH is a measure of acidity. Since acids produce hydrogen ions during ionisation in water, pH is a function of hydrogen ion concentration.

Alternatively

10 7− 1 × 10−14

Kw

K =w [H O ][OH ] =3+ − 1.0 × 10 at 25°C−14

10 7−

10 7−

10 7−

10 7−

pH = −log [H O ]10 3+

[H O ] =3+ 10 -pH


The pH of a solution decreases as the concentration of hydrogen ions increases.

The pOH Scale

The pOH is a measure of basicity. Since basic produce hyrodoxide ions during ionisation in water, pOH is a function of hydroxide ion concentration.

Alternatively

The pOH of a solution decreases as the concentration of hydroxide ions increases. At 25C, pOH and pH of a solution are linked through:

This means that the pOH scale is the reverse relationship of the pH scale. As the pH of a solution increases, the pOH decreases.

pH & pOH Of Acidic & Basic Solutions

Neutral, basic and acidic solutions can be defined in terms of their pH at 25 C:

Neutral solutions have a pH and pOH of 7

Acidic solutions have pH < 7 and pOH > 7

Basic solutions have pH > 7 and pOH < 7

Indicators

One properties of acids and bases is their ability to change colour of certain substances, called indicators.

pOH = −log [HO ]10−

[HO ] =− 10 -pOH

°

pH + pOH = 14

°


Indicators are weak acids or weak bases. The conjugate form of the indicator is one colour and the conjugate base form is another colour.

Universal indicator is widely uses to estimate pH of a solution since it is a micture of several indicators and changes through a range of colours.

📖

write ionic equations to represent the dissociation of acids and bases in water, conjugate acid/base pairs in solution and amphiprotic nature of some salts, for example: – sodium hydrogen carbonate ✅ – potassium dihydrogen phosphate �

Amphiprotic Substances

Amphiprotic substances are able to donate or accept protons, depending on the substances they are reacting with meaning they can behave as an acid or a base.

Water is an excellent example of an amphiprotic substance. When an acid reacts with water, hydronium ions are produced. When a base reacts with water, hydroxide ions are produced. If the solute is a stronger acid than water, then water will react as a base. If the solute is a stronger base than water, then water will react as an acid.

Monoprotic Acids

Monoprotic acids are acids that can donate only one proton. E.g. hydrochloric acid, hydrofluoric acid, nitric acid and ethanoic acid. Only the hydrogen that is part of the highly polar O-H bond is donated. This hydrogen is called the acidic proton.

Polyprotic Acids

Polyprotic acids are acids that can donate more than one proton from each molecule. The number of hydrogen ions an acid can donate depends on its structure. Polyprotic acids donate hydrogen ions in steps when reaction with a base.

Examples of polyprotic acids include diprotic acids - only donate two hydrogen ions such as sulfuric acid and acarbonic acid or triprotic - only donate 3 hydrogen ions such as phosphoric acid and boric acid.


Sodium Hydrogen Carbonate

Potassium Dihydrogen Phosphate

📖

construct models and/or animations to communicate the differences between strong, weak, concentrated and dilute acids and bases ✅

Strength Of Acids & Bases

Some acids can donate protons more readily than others. Bronsted-Lowry theory defines the strength of an acid as it ability to donate protons while the strength of a base is a measure of its ability to accept protons.

The stronger the acid is, the weaker its conjuaget base is. Similarily, the stronger a base is, the weaker its conjugate acid is.

The difference between a strong and weak acid/base is the degree of ionisation. A strong acid/base has all of its molecules undergo ionisation to form hydronium/hydroxide ions while a weak acid/base only has a few of its molecules undergo ionisation.

Strong & Weak Acids

NaHCO + 3 (aq) H O ⇌ 2 (l) OH + −(aq) Na + +

(aq) H CO 2 3 (aq)

KH PO + 2 4 (aq) H O ⇌ 2 (l) OH + −(aq) K + +

(aq) H PO 3 4 (aq)


Acids that readily and easily donate protons are called strong acids. Solutions of strong acids contains ions with no unreacted acid molecules remaining.

Acids that do not readily and easily donate protons are called weak acids. Solutions of weak acids contain ions with unreacted acid molcules. Partial dissociation of a weak acid is shows in an equation using reversible arrows.

Strong & Weak Bases

Bases that readily and easily accept protons are called strong bases. Solutions of strong bases contains ions with no unreacted base molecules.

Bases that do not readily an easily accept protons are called weak bases. Solutions of weak bases contains ions with uncreated base molecules.

Strength Versus Concentration

The strength (strong / weak) of an acid or a base is independent of its concentration (concentrated / dilute). Strength of an acid or a base is its ability to donate or accept protons while the concentration describes the amount of acid or base dissolved in a given volume of solution.

A concentrated acid/base solution is one with a large number of molecules of acid/base in a given volume. A dilute acid/base solution is one with few molecules of acid/base in a given volume.

https://www.youtube.com/watch?v=DupXDD87oHc

HA + H O → 2 A + − H O3+

HA + H O ⇌ 2 A + − H O3+

A + − H O → 2 HA + HO−

A + − H O ⇌ 2 HA + HO−

https://www.youtube.com/watch?v=DupXDD87oHc


Using Quantitative Analysis

📖

conduct practical investigations to analyse the concentration of an unknown acid or base by titration investigate titration curves and conductivity graphs to analyse data to indicate characteristic reaction profiles, for example: – strong acid/strong base ✅ – strong acid/weak base � – weak acid/strong base �

Volumetric Analysis

Volumetric analysis can be used to determine the amount or concentration of a dissolved substance in a solution using a standard solution - a solution with an accurately known concentration.

Primary standards are substances that are so pure that the amount, in mols, can be calculated accurately from their masses. A primary standard should:

Be readily obtained in a pure form

Have a known chemical formula

Be stable and easy to store

Have a high molar mass to minimise effect of errors

Be inexpensive

Examples of some common primary standards include:

Acids: anhydrous sodium carbonate ( ) and sodium borate ().

Bases: hydrated oxalic acid ( ) and potassium hyrogen phthalate ().

The steps for preparing a standard solution include:

Na CO2 3 Na B O ⋅2 4 7

H 02

H C O ⋅2 2 4 H O2KH(C H O )8 4 4


Titration

Titration is a form of volumetric analysis that is used to determine the concentration of an unknown solution (analyte) by measuring the volume of the sample that reacts with a known volume of another substance of known concentration (standard). When a solution is standardised by titration with a primary standard, it is known as a secondary standard.

Volumetric Flask

A volumetric flask is used to prepare the standard solution. An accurate weighted sample is placed in the flask and dissolved in distilled water to form a specific volume of solution.

Pipette

A pipette is used to accurately measure a specific volume of the solution of unknown concentration. This volume is called the aliquot which is then poured into a conical flask foranalysis.


Burette

A burette contains the titrant which delivers accurately known volumes of solution to the conical flask containing the acid or base. The volume of liquid delivered by the burette is called the titre.

It is very important to clean the glass wares above with distilled water and not tap water since it contains ions and other impurities that can taint results and cause inaccuracies. Pipettes and burettes should be firts rinsed with distilled water, then given a final rinsing with the solution they are to contain to minimise dilution as a result of left-over droplets of water.

Volumetric flask and conical flask should be rinsed with distilled water only and not the solution they will contain so no unknown additional amounts of thatr solution would be introduced into the flask.

Performing Titration

To determine the concentration of the analyte, the titrant is slowly added from the burette to the analyte in the conical flask until equivalence point has been reached. Equivalence point is when the reactants have both reacted completely as indicated by the mole ratio in the balanced equation.

In acid-base titration, an indicator that changes colour in the pH range of the equivalence point is used to indicate when acid and base have completely reacted. The end point is the point when the indicator changes colour permanently. For accuracy, the end point should be close to the equivalence point.

The steps to standardize an analyte using a tritration include:


1. An aliquot is measured using a pipette and transferred into a conical flask.

2. A few drops of an appropriate acid-base indicator are added so that a colour change signals when equivalence point has been reached.

3. The burette is filled with a titrant, and initial volume is noted.

4. The titrant is dispensed slowly into the conical flask from burette until the indicator changes colour permanently. The final volume of the titrant is noted.

Titration Curves

A titration curves or pH curve graphically represents the change in pH when an acid is titrated against a base. The shape of a titration curve depends upon the combination of weak or strong acids and bases used in titrations.

Strong Acid & Strong Base Titration

The titration curve produced when a strong acid is titrated with a strong base is shown below.


The graph starts at a high pH since a strong base is present in the conical flask. As the strong acid is initially added to the strong base, it causes a small decrease in pH. However, as more strong acid is added a sudden pH drop occurs. The equivalence point is the midpoint of the vertical section of the graph or where gradient of the titration curve is steepest.

The volume of acid at equivalence point indicates amount of acid needed to completely react with the base. As further strong acid is added after equivalence point, solution becomes highly acidic and pH does not change significantly.

Strong Acid & Weak Base Titration

The titration curve produced when a strong acid is titrated with a weak base is shown below.

The graph starts at a lower pH since a weak base is present in the conial flask. As the strong acid is initially added to the weak base, it causes a significant decrease in pH. However, as more strong acid is added a sudden pH drop occurs as equivalence point is reached.

As further strong acid is added after equivalence point, solution becomes highly acidic and pH does not change significantly.

Weak Acid & Strong Base Titration

The titration curve produced when a weak acid is titrated with a strong base is shown below.


The graph starts at a high pH since a strong base is present in the conical flask. As the weak acid is initially added to the strong base, it causes a small decrease in pH. However, as more weak acid is added a sudden pH drop occurs as equivalence point is reached.

As further weak acid is added after the equivalence point, solution becomes slightly acidic and pH does not change significantly.

Weak Acid & Weak Base Titration

The titration curve produced when a weak acid is titrated with a weak base is shown below.

The graph starts at a lower pH since a weak base is present in the conial flask. As the weak acid is initially added to the strong base, it causes a decrease in pH until equivalence point is reached. There is no clear change in pH at the equivalence point.

As further weak acid is added after the equivalence point, solution becomes slightly acidic and pH does not change significantly.

Selecting An Indicator

During titration, pH of solution in conical flask changes as standard solution is delivered from burette. The equivalence point can be detected using an indicator since they have different colours in acidic and basic solutions.

It is important to select an indicator that changes colour during steep section of pH curve, so that pH at end point is near the pH at equivalence point. It is generally acceptable to use:


Some more examples of indicators include:

For more accurate titration, a pH meter attached to a data logger can be used instead of an indicator.

Conductivity Curves

Conductivity is the ability of a solution to conduct an electric current. Change in conductivity is mearsured with a conductivity meter and is represented on a conductivity curve. Shape of graph depends on the strength of base and acid used. Conductivity of a solution depends on:

identity of ions - as ions size increases conductivity decreases as ions are less mobile.

charge of ions - Strong acids and bases have higher conductivity then weak acids and bases.


concentration of ions - electrical conductivity is proportional to the concentration of ions in solution. concentrated solutions have higher conductivity than dilute solutions.

mobility of ions - more mobile ions will have a higher conductivity since they flow faster.

temperature of ions - as temperature increases conductivity increases as mobility of ions has increased

Equivalence point is indicated by a sudden change in conductivity. Changes in conductivity are due to the fact that ions are replaced by another of different conductivity.

Strong Acid & Strong Base Titration

The conductivity curve produced when a strong acid is titrated with a strong base is shown below.

There is a high initial conductivity due to high presence of H+ ions.

Conductivity decreases as H+ and OH- ion react and get replaced by less conductive cations.

Equivalence point occurs at point with lowest conductivity where cations and anions ions are present

Conductivity increases as OH- ion concentration increases.

Strong Acid & Weak Base Titration

The conductivity curve produced when a strong acid is titrated with a weak base is shown below.


There is a high initial conductivity due to high presence of H+ ions.

Conductivity decreases as H+ and OH- ion react and get replaced by less conductive cations ions.

Equivalence point occurs when lowest conductivity is first reached.

Conductivity remains constant since weak base only dissociates to small extent.

Weak Acid & Strong Base Titration

The conductivity curve produced when a weak acid is titrated with a strong base is shown below.

There is low initial conductivity due to low concentration of H+ ions due to partial dissociation of weak acid

When base is added, initial drop occurs as H+ and H3O+ ions react to form water molecule.

Then conductivity increases as strong base reacts with unionised weak acid to form basic salt that disassociates into ions.

Equivalence point is reached where conductivity suddenly changes

Conductivity increases at greater rate as more strong base is added and increasing OH- concentration.

Weak Acid & Weak Base Titration

The conductivity curve produced when a weak acid is titrated with a weak base is shown below.


There is low initial conductivity due to low concentration of H+ ions due to partial dissociation of weak acid

When base is added, initial drop occurs as H+ and H3O+ ions react to form water molecule.

Conductivity increases as weak base reacts with unionised weak acid to form basic salt that disassociates into ions.

Equivalence point occurs when highest conductivity is first reached.

Conductivity remains constant since weak base only dissociates to small extent.

Calculations In Titrations

To calculate the concentration of the analyte using data from titrations you must have:

Standard solution concentration (mol/L)

Titre volume (L)

Aliquot volume (L)

The concentration of the aliquot is determined by following number of steps.

1. Write a balanced chemical equation for the reaction

2. Calculate the amount, in mol, of standard solution required to reach endpoint.

3. Calculate the amount, in mol, of the analyte that would have reacted with the given amount, in mol, of the standard solution using mole ratio.

4. Determine the concentration of the analyte using the amount, in mol, that reacted and the sample volume.

📖

calculate and apply the dissociation constant (Ka) and pKa (pKa = -log10 (Ka)) to determine the difference between strong and weak acids ✅

Dissociation Of Acids

The equilibrium constant can also be used to describe the dissociation of acids. When an acid ionises in water, it gives an to water which produces a hydronium ion ( ), conjugate base.

H+ H O3+


where:

is the acid

is the conjugate base

The conjugate base is the acid molecules minus a . The degree to which this happens depends on the strength of the acid. The stronger the acid, the greater the degree of ionisation, and the more this equilibrium is shifted to the right.

The equilibrium constant for the ionisation of an acid is called the acid dissociation constant, .

The higher the value of , the more the equilibrium is shifted right, which means that more of the acid ionised, which describes a stronger acid. Thus, a higher means a stronger acid.

Conversly, the lower the value of , the more the equilibrium is shifted left, which means that more of the acid is unionised, describing a weak acid. Thus, a lower means a weaker acid.

Acid Strength

An alternative way of representing acid strength is which is defined as:

This means that as the acid strength increases, the vlaue decreases and vice versa.

Dissociation Of Bases

The equilibirum constant can also be used to describe the dissociation of bases. When a base ionises in water, it takes a from water which produces hydroxide ions ( ) and a conjugate acid.

HA + (aq) H O ⇌ 2 (l) H O + 3+(aq) A − (aq)

HA

A−

H+

Ka

K =a [HA][H O ][A ]3

+ −

Ka

Ka

Ka

Ka

pKa

pK =a −log (K )10 a

pKa

H+ OH−


where:

is the base

is the conjugate acid

The conjugate acid is the base molecules plus an extra . The degree to which this happens depends on the strength of the base. The stronger the base, the greater the degree of ionisation, and the more the equilibrium is shifted to the right.

The equilibrium constant for the ionisation of a base is called the base dissociation constant, .

The higher the value of , the more the equilibrium is shifted right, which means that more of the base ionised, which describes a stronger base. Thus, a higher means a stronger base.

Conversly, the lower the value of , the more the equilibrium is shifted left, which means that more of the base is unionised, describing a weak base. Thus, a lower means a weaker base.

Base Strength

An alternative way of representing basestrength is which is defined as:

This means that as the base strength increases, the vlaue decreases and vice versa.

Calculations Involving Acidity Constants

Weak acids only dissociate to a limited extent. In calculation involving the dissociation of weak monoprotic acids, following assumptions can be made:

The only source of hydronium ions is from the dissociation of the weak acid and self ionisation of water does not contribute to concentration of hydronium ions.

A + −(aq) H O ⇌ 2 (l) OH + −

(aq) HA (aq)

A−

HA

H+

Kb

K =a [A ]−[OH ][HA]−

Kb

Kb

Kb

Kb

pKb

pK =b −log (K )10 b

pKb

[H O ] =3+ [A ]−

[HA] =equilibrium [HA]initial


The concentration of the acid does not change during the dissociation reaction and is equal to the concentration of the acid at equilibrium.

The fraction of the acid that has dissociated is called the percentage dissociation:

📖

explore acid/base analysis techniques that are applied: – in industries – by Aboriginal and Torres Strait Islander Peoples – using digital probes and instruments

Ã

conduct a chemical analysis of a common household substance for its acidity or basicity, for example: – soft drink – wine – juice – medicine

Applications Of Volumetric Analysis

Volumetric analysis is a basic technique used by industries including the food, mining, pharmaceutical, petrochemical and wine making industries.

Acid Content Of Wine

Wine is a complex mixture of chemicals including acids that contribute to the taste of wine and fermentation process. Acid-Base titrations are used to determine the acid content of wine.

A common procedure is to titrate an alliquot of wine against a standardized solution of sodium hydroxide. The total acid content of wine is in terms of percentage of tartaric acid - the major acid found in wine.

% dissociation = × HA

[A ]−100

H C O + 6 4 6 (aq) 2NaOH → (aq) Na H C O + 2 4 4 6 (aq) 2H O 2 (l)


Acid Content Of Fruit Juice & Soft Drink

Acid & Bases Content In Medicines

📖

conduct a practical investigation to prepare a buffer and demonstrate its properties ✅

Buffers

Buffers are important in the environment and living systems. In medicines, buffers are used to maintain the stability and effectiveness of the medicine. In food they preserve flavour and colour, and they are used during the fermentation of wine and in stages of textile dyeing. Buffers have a role in the maintenance and proper functioning of delicate chemical processes that can only occur within a narrow pH range.

Buffers solutions are able to resist a change in pH when small amounts of acid or base are added which is important for processes that require stable and narrow pH ranges.

A buffer consists of a weak conjugate acid-base pair meaning a buffer could be:

A weak acid and its conjugate base

A weak base and its conjugate acid

A weak acid donates a proton to a base to a limited extent. The acid's conjugate base contains one less proton than the acid. A weak base accepts a proton from an acid to a limited extent. The base's conjugate acid contains one more proton then the base. The conjugate acid-base pair chosen for preparing a buffer helps determmine the pH of the buffer

A weak acid and its conjugate base gives an acidic buffer solution while a weak base and its conjugate acid gives a basic Buffer solution. An acidic buffer solution has a pH less than 7. Acidic buffers can be made from a weak acid and one of its salts. An basic buffer solution has a pH greater than 7. Basic buffer can be made from a weak base and one of its salts.

Adding Acid / Base To Buffer


If a small amount of strong acid is added to an acidic buffer, the pH will decrease to a small extent but not as much as if the buffer were not present. The extra hydronium ions generated by the strong acid disturb the existing equilibrium of the buffer solution.

Le Chatelier's Principle states that the system will respns to oppose the change and restore equilibrium. Therefore, some of the additional hydronium ions will react with the weak conjugate base ions present in the buffer producing more weak acid and decreasing concentration of hydronium ions. Therefore, when acid is added the conjugate base combines with any hydronium ions that is added to maintain a stable pH.

When a strong base is added to a buffer, the hydroxide ions from strong base react with the weak acid in buffer to produce conjugate base ions which decreases the concentration of the hydroxide ions. Therefore, when a strong base is added the weak acid combines with any extra hydroxide ions produced which minimises effect on hydronium ion concentration maintaing stable pH.

Buffer Capacity

Buffer solutions have a pH range and capacity that determines how much acid/base can be added before pH changes. The concentration of the buffer influences the effectiveness of the buffer to resist change in pH.

A buffer is most effective when more HA and A- molecules are present to neutralise acid/base and when concentrations of HA and A- are equal. Buffer capacity is a measure of the effectiveness of a buffer solution at resisting change in pH.

📖

describe the importance of buffers in natural systems

Biological Buffer Systems


Carbonic Acid Buffer System

Respiration produces CO2 which dissolves and increases blood acidity:

CO2 (aq) + H2O (l) ←→ H2CO3 (aq)

The presence of buffer maintains pH values in the body. The carbonic acid buffer system consist of carbonic acid and hydrogen carbonate ion.

H2CO3 (aq) + H2O (l) ←→ HCO3- (aq) + H3O+ (aq)

By Le Chatlier's Principle, the system will respond to oppose the change and restore equilibrium and prevent alkalosis (to high pH) and acidosis (to low pH).

Documents

Module 6: Acid & Base Reactions