HSCChemistry
MODULE-6 ACID/BASE REACTIONS
PROPERTIES OF ACIDS AND BASES
Inquiry question: What is an acid and what is a base?
Students:
characteristics and properties of acids and bases and their reversible reactions (ACSCH101)
– acids and bases
– acids and carbonates
– acids and metals (ACSCH067)
– Arrhenius’ theory
– Brønsted–Lowry theory (ACSCH064, ACSCH067)
USING BRØNSTED–LOWRY THEORY
Inquiry question: What is the role of water in solutions of acids and bases?
Students:
– sodium hydrogen carbonate
– potassium dihydrogen phosphate
QUANTITATIVE ANALYSIS
Inquiry question: How are solutions of acids and bases analysed?
Students:
– strong acid/strong base
– strong acid/weak base
– weak acid/strong base (ACSCH080, ACSCH102)
– in industries
– by Aboriginal and Torres Strait Islander Peoples
– using digital probes and instruments
– soft drink
– wine
– juice
– medicine
OUTCOMES
ACIDS AND BASES
WHAT IS AN ACID?
Acids are a group of commonly used substances that share the following properties:
- Conductivity ? [ions]
- Strong Acids = Strong electrolyte
- Weak Acids = Weak electrolyte
WHAT IS A BASE?
Bases are another group of commonly used substances. The properties of bases include:
OUTCOMES
INDICATORS
The equation for calculating pH is: pH = − log10[H+]
HOW DO INDICATORS WORK?
Most indicators are organic weak acids or weak bases. They are a special case of a weak acid.
LOW PH & HIGH PH
o Equilibrium will lie towards the LHS
o Hind will dominate and colour observed will be Colour I
o Equilibrium will lie towards the RHS
o Ind- will dominate and colour observed will be colour II
Some limitations of indicators include:
Some advantages include:
SWIMMING POOLS
Swimming pool pH needs to be maintained close to 7.4 to avoid irritation of the eyes and mucous membranes (which are at approximately pH 7.4).
OUTCOMES
– acids and bases
– acids and carbonates
– acids and metals (ACSCH067)
REACTIONS OF ACIDS
ACID + METAL HYDROXIDE/OXIDE (NEUTRALISATION)
ACID + METAL CARBONATE / HYDROGEN CARBONATE
ACID + REACTIVE METAL (REDOX REACTION)
EVERYDAY USES OF ACIDS/BASES
Bee stings and ant bites are acidic in nature. They can be neutralised using alkaline medicine such as baking powder.
Wasp stings are alkaline in nature. Vinegar can be used to cure wasp stings because vinegar can neutralise the sting.
When stung by a stingray, concentrated vinegar can be used to stop the nematocysts from firing off such that you won’t get injected with more venom.
OUTCOMES
– Arrhenius’ theory
– Brønsted–Lowry theory (ACSCH064, ACSCH067)
LAVOISIER – 1776 “FATHER OF CHEMISTRY” (OXIDES)
Antoine Lavoisier was a French chemist who established the quantitative science of chemistry. He investigated oxides of different elements, and found that non-metal oxides reacted with water, producing acidic solutions. He concluded that an acid must contain oxygen. Eg:
DAVY – 1810 (HYDROGEN)
English chemist Humphry Davy (famous for his redox and electrolytes works) demonstrated that hydrochloric acid did not contain oxygen, thus disproving Lavoisier’s theory. Davy suggested that hydrogen must be the unifying component of acids rather than oxygen.
LIEBIG – 1838 (ACID + METAL)
The German chemist Justus von Liebig expanded on Davy’s idea. In 1838, he proposed that acids contain hydrogen which could be displaced by a reaction with a metal. Eg:
ARRHENIUS – 1884 (H+ AND OH-)
- Svante Arrhenius proposed the first concept of acids and bases we still use.
water into ions.
- An Arrhenius acid is a substance that produces a H+(aq) in water.
Arrhenius also notated that the most reactive acids also had the highest electrical conductivities. This led to the concept that the strongest acids were the most dissociated in aqueous solution.
An Arrhenius base is a substance that produces OH- (aq) in water.
LIMITATIONS OF THE ARRHENIUS DEFINITION
Arrhenius’ definition does not explain the basic behaviour of substances like ammonia, which do not contain hydroxide ions in their formulae and hence should not be able to produce OH-.
It does not explain why neutralisation reactions between some acids and bases produced solutions that were not neutral.
The Arrhenius definition only covers acids and bases in aqueous solutions.
BRONSTED-LOWRY ACIDS AND BASES (CONCEPTUAL DEFINITION)
In 1923, Danish chemist Johannes Nicolaus Bronsted and English chemist Thomas Martin Lowry independently
proposed a new definition of acids and bases.
Acids and bases are defined by their role in a reaction:
- The proton donor is the Bronsted-Lowry acid
- The proton acceptor is the Bronsted-Lowry base
The Bronsted-Lowry definition allows many more species to be defined as acids or bases. It can explain the basic behaviour of ammonia. The NH3(aq) is accepting a proton from HCl in the aqueous solution. NH3(aq) is a Bronsted Lowry base and HCl is a Bronsted Lowry acid.
Bronsted Lowry theory also explains the basic behaviour of ionic compounds in solution.
Soluble carbonates and hydrogen carbonates contain Bronsted Lowry bases. They produce basic solutions.
1) First the compound dissolves in water to produce aqueous ions. This step proceeds completely, since all Group 1 ionic compounds are soluble:
2) The dissolved carbonate or hydrogen carbonate ion is a Bronsted Lowry base which reacts with water to produce hydroxide ions:
The Bronsted Lowry definition is broad enough that some species like water can be classified both an acid and a base.
When HCL dissolves in water, water accepts H+ from HCl and is acting as a Bronsted Lowry base.
A substance that is capable of both donating and accepting protons (depending on the other reagent) is known as an amphiprotic substance.
HYDRONIUM IONS
STRENGTHS OF ACIDS AND BASES
STRONG/WEAK ACIDS
A strong acid is a strong electrolyte that ionises completely in water to produce hydronium ions in aqueous solutions. Eg:
All other acids are weak acids, which ionise partially in water to form hydronium ions in solution. They are weak electrolytes. Eg:
ACID DISSOCIATION CONSTANT (KA)
The higher the Ka the stronger the acid strength as it favours the RHS.
Most acids are monoprotic: they can only produce one proton. However, some acids have more than one Ka value.
These acids can produce more than one H+ ion and are known as polyprotic acid.
The number of protons an acid produces is unrelated to its strength.
STRONG/WEAK BASES
Strong bases include Group 1 and 2 metal hydroxides. When they dissolve, they produce free hydroxide ions.
Carbonates are weak bases.
- NaOH (Caustic soda, drain cleaner, oven cleaner)
- KOH
- Ca(OH)2 (Lime water)
- Ba(OH)2
Although Group 2 hydroxides are strong bases, they are poorly soluble in water.
A weak base is one that reacts partially with water to indirectly form hydroxide ions in solution.
STRENGTH VS CONCENTRATION
CONJUGATE ACIDS AND BASES
The ionisation of an acid can be represented by the following equation:
When the equation above is read in reverse, it shows A- accepting a proton and acts as a Bronsted Lowry base.
HA and A- are known as a conjugate acid/base pair.
RELATIVE STRENGTH OF CONJUGATE PAIRS
The two species in conjugate pair have inverse strength.
A strong acid has a strong tendency to give up a H+, and completely ionise in water.
A weak acid has a relatively strong conjugate base.
Rule: The weaker the acid, the stronger the conjugate base. (Vice versa for weak bases)
EXAMPLE QUESTION 1
Explain why the presence of nitrate ions in an aqueous solution will not make it basic. (i.e. will not produce extra OH-ions).
This reaction goes to virtual completion. The reverse reaction does not occur to any significant extent.
∴ NO3−(aq) has no tendency to accept a H+ and is a neutral ion. It will not accept any H+ from water to produce OH-ions. NO3− is also the conjugate base of a strong acid.
EXAMPLE QUESTION 2
Explain why an aqueous solution containing fluoride ions will be basic.
F- has a relatively high tendency to accept H+ as it is a conjugate base of a weak acid HF. It will accept H+ from water to produce OH- ions resulting in a basic solution of pH > 7.
OUTCOMES
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
NON-NEUTRAL SALT SOLUTIONS
Many ions are not neutral when dissolved in water. They will have acidic or basic properties.
- This means that in neutralisation reactions, the salt product is not necessarily neutral (pH = 7).
- pH depends on the nature of the salt.
To show that an ionic compound will form an acidic, basic or neutral solution:
- Strong acid + Weak base à Acidic
- Strong acid + Strong base à Neutral
- Weak acid + Strong base à Basic
EXAMPLE QUESTION 1
Will a solution of potassium fluoride be acidic or basic?
AMPHIPROTIC SUBSTANCE
Water can be classified as both a Bronsted-Lowry acid and base. A molecule or ion that can accept and donate protons is called amphiprotic. Amphiprotic species will lose or gain a proton depending on the other reactant.
- If the other species is a stronger base, it will act as a Bronsted-Lowry acid
- If the other species is a stronger acid, it will act as a Bronsted-Lowry base.
Examples include: HCO3-, HSO4-, H2PO4-, HPO42- and NH3
Amphiprotic substances are a subset of amphoteric substances, which are substances that can react with both acids and bases.
In laboratories, NaHCO3 is commonly used for neutralising acid spills as it is a weak, amphoteric base that produces CO2 during neutralisation.
It is a weak base therefore neutralisation is not as exothermic.
Amphoteric/amphiprotic therefore will absorb OH- if too much is added.
Produces CO2 therefore it can be seen when neutralisation is complete
Solid, therefore will not contribute to the size of the spill
ACID/BASE BEHAVIOUR OF OXIDES
An acidic oxide is one which either reacts with water to form an acidic solution or reacts with bases to form acidic salts. Common acidic oxides are CO2 and P4O10 (diphosphorus pentoxide) and SO2.
Non-metal from the RHS of the periodic table tend to form acidic oxides. These elements have high electronegativity and share electrons when bonding with oxygen, so non-metal oxides are covalent.
A basic oxide is one that reacts with water to form an alkaline solution or reacts with acids to form basic salts. Metals from the LHS of the periodic table tend to form basic oxides. These elements have low electronegativity, so metal oxides are ionic.
OUTCOMES
calculate and apply the dissociation constant (Ka) and pKa (pKa = -log10 (Ka)) to determine the difference between strong and weak acids (ACSCH098)
PH AND H+ CONNECTION
- pH 7 is neutral at 25℃ ([H3O+] = [OH- ])
The pH (potential of hydrogen) scale is logarithmic (base 10), not linear. Significant figures for logs are those after the decimal point.
PH OF STRONG AND WEAK ACIDS
pH can deduce the relative strength and concentration of different acidic solutions.
- ↑ Acid strength à ↑ Degree of ionisation à ↑ [H+] à ↓ pH
- ↑ [Acid] à ↑ [H+] à ↓ pH
- Acid strength and concentration both affect the pH of a solution.
For a monoprotic strong acid, the concentration of H+ is the same as the concentration of the acid (HX → H+ + X−)
For weak acids, the concentration of H+will depend on its strength and concentration of the acid solution.
POH OF STRONG AND WEAK BASES
pOH is a measure of OH- concentration that is similar to pH.
Strong bases are Group 1 and 2 hydroxides. The concentration of hydroxide ions depends on the number of hydroxide ions in the formula.
For weak bases, the concentration of OH- will depend on its strength and the concentration of the base solution.
AUTOIONISATION OF WATER
The concentration of H+ and OH- in any aqueous solution are directly related. This is because water is a very weak acid which forms the following exothermic equilibrium reaction.
- The equilibrium constant for this equilibrium is called the ionisation constant for water.
PKA AND PKB
Since Ka values are usually very small, pKa values are often cited instead:
OUTCOMES
calculate pH, pOH, hydrogen ion concentration ([H+ ]) and hydroxide ion concentration ([OH– ]) for a range of solutions (ACSCH102)
EXAMPLE QUESTION 1
Benzoic acid, a food preservative, is a weak monoprotic acid. The pH of 0.50M benzoic acid solution is 2.25. What is the Ka of benzoic acid?
EXAMPLE QUESTION 2
DEGREE OF IONISATION
All acids ionise in water, the degree of which is different for individual acids.
To calculate the percentage of any component in a sample, the formula is:
FOR A WEAK ACID HA
PH MEASUREMENTS
PH PROBE
A pH probe or pH meter can be used to measure pH.
- Advantages: Precision, non-destructive
- Disadvantages: Initial costs, requires calibration before use
OUTCOMES
calculate the pH of the resultantsolution when solutions of acids and/or bases are diluted or mixed
DILUTION CALCULATIONS
The total moles of solute in a concentrated solution and the diluted solution are the same.
EXAMPLE QUESTION 1
PH OF SOLUTIONS AFTER MIXING
EXAMPLE QUESTION 2
Calculate the pH of the resultant solutions when 50.0 mL of 0.020 M NaOH solution and 100 mL of 0.012 M Ba(OH)2
PH AFTER NEUTRALISATION
When a strong acid and strong base react and undergo a neutralisation reaction, the pH of the final solution will depend on the concentration of the reactant in excess.
EXAMPLE QUESTION 1
Calculate the pH of the resultant solution when 35.0mL of 0.250M HCl and 45.0mL of 0.350M NaOH are mixed.
OUTCOMES
HOW DO BUFFERS WORK?
A buffer solution is a solution that can resist pH change when small amounts of an acid or base are added. In order for it to work, it must be able to compensate for the addition of either of an acid or base, otherwise the pH will change significantly.
Buffers are commonly made by mixing a weak acid and its conjugate base (or vice versa) typically in equimolar amounts of each.
For example, a buffer solution containing CH3COOH and CH3COO− (in the form of NaCH3COO)
The mixture exists in equilibrium, so all of the species in the equation are present.
When a small amount of acid is added, the concentration of H3O+ increases.
When a small amount of base is added, the situation is slightly more complicated.
There are two acids that could react with the base CH3COOH and H3O+
Because the concentration of H3O+ is far lower than the concentration of CH3COOH, the majority of OH−will react with CH3COOH, producing water and CH3COOH−
Role of conjugate base is to react with excess H3O+. It ‘absorbs’ the additional H3O
BUFFER CAPACITY
The effectiveness of a buffer is known as buffer capacity.
It is defined as the moles of H3O+ or OH− necessary to change the pH of the buffer solution by one unit
Buffer capacity depends on both the pH of the buffer and the total concentration of the weak acid and conjugate base (or vice versa)
A buffer is most effective when the amounts of weak acid and conjugate base present are similar (equimolar)
When pH = pKa, the concentrations of weak acid and conjugate base are equal.
Therefore, a buffer solution is most effective when the pH is within 1 unit of its pKa.
When the pH is too high, there is not enough acid to react with the added OH− When the pH is too low, there is not enough conjugate base to react with any H3O+.
OUTCOMES
conduct practical investigations to analyse the concentration of an unknown acid or base by titration
DIRECT TITRATION CALCULATIONS
The steps for a titration calculation are:
For neutralisation reactions, the strength of the acid is irrelevant as the base is stronger than water. All of the protons in a polyprotic acid will be irreversibly removed by the base.
EXAMPLE QUESTION 1
Alex wanted to analyse a solution of NaOH. She first prepared 250 mL of a standard solution using 4.72 g of solid sodium carbonate. She then reacted 25.0 mL of HCl of unknown concentration with the standard solution. The following volumes of standard solution were required:
25.0 mL of the HCl solution was then titrated against the NaOH solution. The average volume of NaOH required to reach the equivalence point was 30.85 mL.
DILUTION – TITRATION CALCULATIONS
When analysing a substance, the concentration may be too high for a direct titration experiment tom be efficiently carried out. Instead, the substance would be diluted by a known amount, and then the diluted solution would be titrated.
WORKED EXAMPLE 1
21.55 mL of a hair dye product containing ammonia was diluted to 250.0 mL. 20.0 mL aliquots of the diluted dye were titrated against 0.048 M sulfuric acid. The average titre was 18.35 mL. The density of the hair product was measured to be 1.1 g/mL.
Calculate the percentage by weight of ammonia in the hair dye.
BACK TITRATION CALCULATIONS
A back titration, or indirect titration, is a two-stage analysis:
Reactant A (of unknown concentration) is reacted with an excess of Reactant B (of known concentration andvolume).
A titration is the performed on the excess Reactant B by determining the moles of Reactant C required to neutralise the excess.
Summary:
Back titrations are generally used when:
One of the reactants is volatile (e.g. ammonia)
An acid or base is an insoluble salt (e.g. calcium carbonate)
Direct titration would involve weak acid/weak base titration (making it difficult to determine the equivalence point).
EXAMPLE QUESTION 1
A student wanted to determine the calcium carbonate present in a 0.250 g sample of chalk. The student placed the chalk sample in a 250 mL conical flask and pipetted 50.00 mL of 0.100 M sulfuric acid into the flask. The excess acid was titrated with 0.340 M sodium hydroxide. The results for the volume of sodium hydroxide required are tabulated below:
Calculate the mass percentage of calcium carbonate in the chalk sample.
EXAMPLE QUESTION 2
A student was asked to determine the concentration of ammonia in a commercial sample of “cloudy ammonia cleaning solution”. The student first pipetted 25.00 mL into a 250 mL volumetric flask and filled it to the mark with distilled water. She then transferred a 25.00 mL aliquot to a clean conical flask, and added exactly 40.00 mL of 0.100 M HCl. She then titrated the excess HCl with 0.050 M sodium carbonate. 21.30 mL of sodium carbonate solution was required.
OUTCOMES
conduct practical investigations to analyse the concentration of an unknown acid or base by titration
TYPES OF ANALYSIS
- Qualitative analysis: Involves observations only
- Quantitative analysis: Involves measurements (mass, volume etc.)
GENERAL TITRATION PROCEDURE
In a titration experiment, the number of moles of a target material (the analyte) is determined. This can then be used to calculate the concentration.
A measured volume of the solution of unknown concentration, the analyte, is usually placed in a conical flask (titrand) with a burette containing the titrant above it.
A burette is a piece of volumetric glassware. It is a long tube with a tap at one end so measured volumes of titrant can be accurately added to the titrand.
Before the experiment begins, an indicator will normally also be added to the conical flask to determine the approximate equivalence point.
To perform the experiment the titrant is slowly added and stepwise to the conical flask, with swirling, until the indicator undergoes a permanent colour change (the end point).
Acid-base titrations are the most common titrations:
The pH changes rapidly towards the end point of the titration and it is easy to add too much titrant. The titrant must be added very carefully, in small volumes close to the end to successfully determine the exact amount required for complete reaction.
DETAILED TITRATION PROCEDURE
For titrating HCl against NaOH. HCl is placed into the conical flask and NaOH in the burette.
PREPARING THE TITRAND (IN THE CONICAL FLASK)
Rinse the inside of the pipette with a small with a small amount of HCl solution 3 times.
Use a pipette filler to fill the pipette with HCl until the bottom of the meniscus rests on the calibration line.
Hold the pipette so that its tip is resting against the inside of a clean conical flask. Let the solution run out.
Once the liquid level has stabilised, leave the pipette tip touching the flask for a few seconds before removing. (The pipette is calibrated to deliver the correct volume when a small amount of liquid remains in the tip – do not shake it into the flask).
Use a wash bottle containing distilled water to wash any solution that might be on the inside wall to the bottom of the conical flask.
PREPARING THE TITRANT (IN THE BURETTE)
TITRATION
WASHING
Burette and pipette: Used to deliver the solutions used in the titration. Final rinsing with the solution to be delivered.
Conical flask: Used to hold the aliquot or titrand. Final rinsing with distilled water.
The washing procedure affects the accuracy of the calculated concentration.
TITRATION TECHNIQUE
RINSING
It is important to rinse each piece of glassware with the appropriate solution after cleaning with distilled water and immediately prior to use.
The solution that is to be transferred using a pipette is of accurately known concentration, or its concentration is to be accurately determined. If droplets of distilled water are present in the pipette, it will dilute the reagent being delivered.
- Rinse burette and pipette with solutions
- Rinse conical flask and volumetric flask with distilled water
VOLUMETRIC ERRORS
All glassware in titration is calibrated to be accurate when measurements are taken at the bottom of the meniscus.
Any air bubbles in the liquid must be removed for volumes to be accurate.
PIPETTE CALBIRATION
The pipettes used in titration are calibrated to deliver the specified volume of solution with no additional force.
They are marked TD (to deliver) or EX (to expel)
This means that there should be a small volume of liquid left in the tip of the pipette after the aliquot has been accurately transferred. This should not be shaken out into the conical flask.
DEVIATION
Accessing accuracy – How close you are to the accepted value
STANDARD SOLUTION
A standard solution is a solution containing a precisely known concentration of a substance. They can be categorised as primary or secondary.
A primary standard is produced when a substance of high purity dissolved in a known volume of solvent.
Accurately weigh out a mass of solid close to the required mass in a beaker. Record the actual mass weighed.
Add enough distilled water to dissolve the solid.
Carefully transfer all the weighed mass to a clean volumetric flask of the approximate size, using a wash bottle and funnel. All the equipment that came into contact with weighed mass should be rinsed into the flask.
Add distilled water until the bottom of the meniscus is resting on the line on the neck of the flask. Add the last few drops with a dropper.
Stopper the flask. Firmly holding the stopper in place, invert several times to ensure the solution is homogeneous.
Label the flask with the exact concentration, solution, date and name.
A substance suitable for preparing a primary standard solution should have the following features:
- High purity
- Unaffected by exposure to air
- Non-hygroscopic (does not absorb water from air)
- Have a large molecular mass to reduce percentage errors
- Be a solid for easier weighing
- Cheap and readily available
- Have a high water solubility
A secondary standard is produced when its concentration is determined via stoichiometry.
- The process of producing a secondary standard is called standardisation
EQUIVALENCE POINT
The equivalence point of a titration is the point at which the amount moles of acid and bases added match the stoichiometric ratio.
PH OF THE EQUIVALENCE POINT
As the pH of water is neutral, the pH of the equivalence point will depend entirely on the salt produced: whether it is acidic, basic or neutral.
If an acidic or basic salt is produced by the neutralisation reaction in a titration experiment, the equivalence point will not be neutral.
- Neutral salts are formed when strong acids react with strong bases.
- Acidic salts are formed when strong acids react with weak bases.
- Basic salts are formed when weak acids react with strong bases.
Note: When strong acids react with metal carbonates (weak bases), the neutral salt is formed, but the resulting solution is still acidic. This is because the carbon dioxide dissolves in water to produce an acidic solution.
INDICATOR SELECTION
Indicators can be used to find the approximate equivalence point of a titration.
The equivalence point is when exactly enough moles of titrant have been added to react with all the titrand.
The end point is when the indicator first undergoes a permanent colour change.
An indicator should be selected so that the end point is as close as possible to the equivalence point. (Systematic error. Will impact validity and accuracy)
BROMOTHYMOL BLUE
PHENOLPHTHALEIN
METHYL ORANGE
EXAMPLE QUESTION 1
Explain why 0.20 M acetic acid and 0.20 M hydrochloric acid require the same volume of sodium hydroxide solution to reach equivalence point, but the pH values at their equivalence points are different.
Both acetic acid and hydrochloric acid are monoprotic acids that reacts to completion when reacted with a strong base. As the concentration are the same, the same amount is needed to neutralise the strong base, NaOH. CH3COO− is the conjugate base of a weak acid and will react with water to produce a basic solution. Therefore, resulting in a pH > 7. Cl− is the conjugate base of a strong acid and will not react hence the solution remains pH = 7.
STRONG ACID-STRONG BASE TITRATION
The equivalence point is located at the most vertical point (point of inflection). All three common indicators for titration are suitable for determining the equivalence point for a strong acid-strong base. This is because there is a large rapid change in pH near the equivalence point so all of the indicator would change colour when the same volume of based is added, therefore it is not critical which indicator is used.
WEAK ACID-STRONG BASE TITRATION
STRONG ACID-WEAK BASE TITRATION
WEAK ACID-WEAK BASE TITRATION
OUTCOME
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 (ACSCH080, ACSCH102)
CONDUCTIVITY GRAPHS
During a titration, the conductivity of the solution changes. The equivalence point may be located by plotting the conductance as a function of the volume of titrant added.
The electrical conductivity of a solution depends on:
- The concentration ions present
- The mobility of the ions present
o More mobile ions, the more conductive it is
o H+ and OH- are highly mobile
- H+ ions are more conductive than OH-ions
Conductometric titrations are useful for titrations of coloured solutions, analysis of dilute solutions, and when reversible reactions are used (e.g. weak acid-weak base titration).
General rule:
- Strong ➡ Linear
- Weak ➡ Curved
STRONG ACID + STRONG BASE
WEAK ACID + STRONG BASE
Initially, the conductance is low due to the low ionisation of the weak acid. On the addition of the strong base, there is a decrease in conductance due to the replacement of the H+ by Na+ but also supresses the dissociation of the acetic acid due to the common ion acetate.
The conductance increases on adding NaOH as it neutralises the undissociated CH3COOH to NaCH3COO which is a strong electrolyte. Conductivity increases due to the highly conductive OH- ions.
STRONG ACID + WEAK BASE
Before the equivalence point, conductivity decreases like in the strong acid-strong base graph.
After the equivalence point, the graph is almost horizontal as the excess weak base is not significantly ionised due to the presence of its conjugate acid.
Initially, the conductance is high due to the strong acid. The conductance decreases due to the replacement of H+.
After the equivalence point has been reached in the graph becomes almost horizontal, since the excess weak base (aqueous ammonia) is not easily ionised in the presence of the salt.
[Not shown on the graph]
1) CH3COOH is a weak acid and therefore only partially ionise.
2) [H+] gets used up. CH3COO- gets produced which supresses the ionisation of CH3COOH.
3) Production of more ions
4) Equivalence point
5) Excess NH3 is suppressed due to the common ion effect
OUTCOMES
– soft drink
– wine
– juice
– medicine
ANALYSIS OF COMMON HOUSEHOLD SUBSTANCES
The exact procedures will vary depending on the nature of the substance to be analysed:
- Indicators may not be suitable for intensely coloured substances
- Insoluble solids may need to be dissolved before titration
- Many substances will require dilution before titration.
OUTCOMES
– in industries
– by Aboriginal and Torres Strait Islander Peoples
USE IN THE WIDER WORLD
