Hsc chemistry

Module 6

Acid base reactions

Student notes

Introduction

· Acids and bases, and their reactions, are used extensively in everyday life and in the human body.

· The chemistry of acids and bases contributes to industrial contexts and the environment.

· Therefore, it is essential that the degree of acidity in these situations is continually monitored.

· By investigating the qualitative and quantitative properties of acids and bases, students learn to appreciate the importance of factors such as pH and indicators.

Properties of Acids and Bases

Inquiry question: What is an acid and what is a base?

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

· Naming Acids and Bases

Key Points

· Acids are named based on their anion — the ion attached to the hydrogen. In simple binary acids, one ion is attached to hydrogen. Names for such acids consist of the prefix “hydro-“, the first syllable of the anion, and the suffix “-ic”.

· Complex acid compounds have oxygen in them. For an acid with a polyatomic ion, the suffix “-ate” from the ion is replaced with “-ic.”

· Polyatomic ions with one extra oxygen (as compared to the typical polyatomic ion) have the prefix “per-” and the suffix “-ic.”

· Polyatomic ions with one fewer oxygen have the suffix “-ous”; ions with two fewer have the prefix “hypo-” and the suffix “-ous.”

· Strong bases with “-OH” (hydroxide) groups are named like ionic compounds. Weak bases are named like molecular compounds or organic compounds.

· Naming Acids

· Acids are named by the anion they form when dissolved in water. Depending on what anion the hydrogen is attached to, acids will have different names.

· Compounds are named starting with the prefix “hydro-,” then adding the first syllable of the anion, then the suffix “-ic.” For example, HCl, which is hydrogen and chlorine, is called hydrochloric acid.

Nomenclature of common acids. This chart provides the nomenclature of some common anions and acids

· More complex acids have oxygen in the compound. There is a simple set of rules for these acids.

1. Any polyatomic ion with the suffix “-ate” uses the suffix “-ic” as an acid. So, HNO3 will be nitric acid.

2. Any polyatomic ion with the suffix “-ite” uses the suffix “-ous” as an acid. So, HNO2 will be nitrous acid.

· Naming Bases

· Most strong bases contain hydroxide, a polyatomic ion.

· Therefore, strong bases are named following the rules for naming ionic compounds. For example, NaOH is sodium hydroxide, KOH is potassium hydroxide, and Ca(OH)2 is calcium hydroxide.

· Weak bases made of ionic compounds are also named using the ionic naming system. For example, NH4OH is ammonium hydroxide.

· Weak bases are also sometimes molecular compounds or organic compounds because they have covalent bonds. Therefore, they are named following the rules for molecular or organic compounds. For example, methyl amine (CH3NH2) is a weak base. Some weak bases have “common” names. For example, NH3 is called ammonia; its name isn’t derived from any naming system.

· Note the following equations when acids dissociate in water:

· Hydrochloric Acid

HCl(g) + H2O H+(aq) + Cl-(aq)

· Sulfuric Acid

H2SO4 + H2O 2H+(aq) + SO42-

· Nitric Acid

HNO3 + H2O H+(aq) + NO3-(aq)

· Phosphoric Acid

H3PO4 + H2O 3H+(aq) + PO43-(aq)

· Properties of Acids:

· Acids have a sour taste

· Conduct electricity in aqueous solution

· Turn blue litmus red

· Are corrosive

· Properties of Bases

· Bases feel soapy in aqueous solution

· Bases have a bitter taste

· Conduct electricity in aqueous solution

· Turn red litmus blue

· 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

Investigation 5.1

· predict the products of acid reactions and write balanced equations to represent:

· acids and bases

· acids and carbonates

· acids and metals

Revision from Year 11

· General Reaction Equations

1. Neutralisation

acid + alkali metal salt + water

Example

HCl + NaOH NaCl + H2O

2. Metal + acid

metal + acid metal salt + hydrogen

Example

Mg + H2SO4 MgSO4 + H2

3. Metal carbonate + acid

metal carbonate + acid metal salt + carbon dioxide + water

Example

MgCO3 + H2SO4 MgSO4 + CO2 + H2O

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

Research Assignment

· conduct a practical investigation to measure the enthalpy of neutralisation

Investigation 5.2

· 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

· In the 1780s, the French chemist, Antoine Lavoisier, undertook experiments on combustion and found that non-metal oxides reacted with water forming acidic solutions. He concluded that an acid must contain oxygen.

· In 1810, the English chemist, Humphry Davy, Davy demonstrated that muriatic acid (hydrochloric acid) was a compound of hydrogen and chlorine and did not contain oxygen.

In 1815, Davy observed that all known acids contained hydrogen that could be replaced by reaction with a metal. He also noted that compounds of metal with oxygen were bases.

· Lavoisier and Davy's definitions were based on observable properties.

· In 1884, the Swedish chemist, Svante Arrhenius, put forward definitions based on concepts about particles too small to be directly observed. Arrhenius proposed that:

An acid produced hydrogen ions H+ when dissolved in water.  A base produced hydroxide ions OH- when dissolved in water.

· Brönsted-Lowry theory of acids and bases.

· A theory, based on proton transfer, was independently outlined in 1923 by the Danish chemist, Johannes Bronsted, and the British chemist, Thomas Lowry. An acid is a proton donor and a base is a proton acceptor.

· An acid-base reaction involves proton transfer from acid to base.

· Lewis Definition

· An acid is an electron pair receptor

· A base is an electron pair donor

For example when BF3 reacts with NH3 to form BF3NH3, the boron is a Lewis acid as it accepts a non-bonding pair of electrons from nitrogen. Nitrogen is a base as it donates the electron pair.

Check Your Understanding p139 Q1, 3, 4, 5, 6 & 7

Chapter Revision Questions p141: Q2, 3, 4, 5, 6a, 7, 8

Using Brønsted–Lowry Theory

Inquiry question: What is the role of water in solutions of acids and bases?

Students:

· conduct a practical investigation to measure the pH of a range of acids and bases

Investigation 6.1

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

· The use of the pH scale in comparing acids and bases.

· The pH scale is used to compare the concentration of hydrogen ions in solutions of acids and bases.

· To a very small extent water undergoes auto- or self-ionisation:

2H2O(l) H3O+(aq) + OH-(aq)

· In the ionisation of water, equilibrium strongly favours the reactants. Hence, only small concentrations of hydrogen ions and hydroxide ions are formed. The equilibrium constant , Kw , for the ionisation of water is:

Kw = [H+] [OH-] = 1.0 x 10-14

pH + pOH = 14

pH = -log10[H+]

[H+] = 10-pH

pOH = -log10[OH+]

· Note: [ ] (square brackets) indicates concentration

· The following table relates pH to the hydrogen ion concentration, [H+], and provides examples of common aqueous solutions for each pH value.

pH

[H+]

[OH-]

[H+] x [OH-]

Aqueous solution example

0

100 = 1

10-14

10-14

1 M hydrochloric acid

1

10-1

10-13

10-14

0.1 M hydrochloric acid

2

10-2

10-12

10-14

0.01 M 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

0.01 M sodium hydroxide

13

10-13

10-1

10-14

0.1 M sodium hydroxide

14

10-14

100 =1

10-14

1 M sodium hydroxide

Worked Example 6.1and Try These Yourself p151

Check Your Understanding 6.2 Q2, 5, 6, 8 & 9

· conduct an investigation to demonstrate the use of pH to indicate the differences between the strength of acids and bases

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

· A concentrated solution contains a large amount of solute in a given amount of solution. A 10 mol L-1 solution would be called concentrated.

· A dilute solution contains a small amount of solute in a given amount of solution. A 0.01 mol L-1 solution would be called dilute.

· Strong acids essentially completely ionise in aqueous solution.

HCl(g) H+(aq) + Cl-(aq)

H2SO4(l) H+(aq) + HSO4-(aq)

HSO4-(aq) H+(aq) + SO4-(aq)

· Weak acids only partially ionise in aqueous solution.

For example: acetic acid

CH3COOH(aq) H+(aq) + CH3COO-(aq)

Or

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

· Note: Acetic acid only ionises a few percent in water and that the arrows indicate an equilibrium reaction.

· If you prepare 0.1 M solutions of each you should find the hydrochloric acid has a pH of 1 corresponding to a [H+] = 10-1 = 0.1 M. That is, practically every HCl molecule has ionised, producing an H+.

· By contrast a 0.1 M solution of acetic acid will have a pH close to 3 indicating a [H+] close to 10-3 = 0.001 M. Only about 0.001 / 0.1 = 1% of the acetic acid molecules have ionised producing an H+.

Investigation 6.2

· 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

· Note the following equations when acids dissociate in water:

Hydrochloric Acid (Monoprotic = 1H+)

HCl(g) + H2O H+(aq) + Cl-(aq) OR

HCl(g) + H2O H3O+(aq) + Cl-(aq)

Hydronium Ion (Coordinate covalent bond)

Sulfuric Acid (Diprotic = 2 H+)

H2SO4 + H2O 2H+(aq) + SO42-

Nitric Acid

HNO3 + H2O H+(aq) + NO3-(aq)

Phosphoric Acid (Triprotic = 3H+)

H3PO4 + H2O 3H+(aq) + PO43-(aq)

· Conjugate acid/base pairs

· When an acid donates a proton, it forms its conjugate base.

     HCl     +     H2O                Cl-           +     H3O+      acid                           conjugate base

· When a base accepts a proton, it forms its conjugate acid.      HCl     +     H2O         Cl-     +     H3O+                      base                   conjugate acid

· The stronger a particular acid, the weaker will be its conjugate base.

· Whenever an acid and a base react, they form their conjugates:      HCl     +     H2O             Cl-          +         H3O+     acid1          base2       conjugate base1    conjugate acid2

· Hydrochloric acid and chloride ion are a conjugate acid-base pair.

· Water and hydronium ion are another conjugate acid-base pair.

· Amphiprotic Substances

· An amphiprotic molecule (or ion) can either donate or accept a proton, thus acting either as an acid or a base. Water, amino acids, hydrogen carbonate ions and hydrogen sulfate ions are common examples of amphiprotic species. Since they can donate a proton, all amphiprotic substances contain a hydrogen atom.

· Water as an amphiprotic substance

· Water is normally considered a neutral substance. For instance, if water reacts with a base like ammonia, then it will act as an acid by donating a proton, a positively-charged particle, in the form of a hydrogen (H+) ion to ammonia (NH sub 3).

· Water can act as an acid or base:

H2O + H2O H3O+ + OH-

Water Acts as an Acid

· However, if water comes in contact with an acid, like hydrochloric acid (HCl), it will act as a base by receiving a proton in the form of a hydrogen (H+) ion from the hydrochloric acid (HCl).

Water acting as a base.

· Sodium hydrogen carbonate

· When sodium hydrogen carbonate (sodium bicarbonate) dissociates in water, it becomes a sodium ion (Na+) and a hydrogen carbonate (bicarbonate) ion (HCO3-).

· The bicarbonate ion is the amphiprotic species.

· This is because it can lose a hydrogen ion and become the carbonate ion (CO32-) or it can gain a hydrogen ion and become carbonic acid (H2CO3).

Potassium dihydrogen phosphate

· calculate the pH of the resultant solution when solutions of acids and/or bases are diluted or mixed

To calculate concentration of solutions after dilution or mixing the following formula is used:

C1V1 = C2V2

Where:

C = concentration in molL-1

V = volume (Note: volume can be in mL or L as long as it is the same on both sides of the equation.

Worked Example 6.2 Try these Yourself and Check Your Understanding 6.4 p160 – 161

Worked Example 6.3 and Try These Yourself and Check Your Understanding 6.4 p160 – 161

Worked Example 6.5 and Try These Yourself p164

Check Your Understanding 6.5 p165 Q2, 3, 4, 5, 6, 8

Quantitative Analysis

Inquiry question: How are solutions of acids and bases analysed?

Students:

· conduct practical investigations to analyse the concentration of an unknown acid or base by titration

· The correct technique for conducting titrations and preparation of standard solutions.

· The procedure of adding one solution to another until the reaction between them is complete is called titration.

· In titration, a definite volume of one of the solutions is placed in a conical flask using a pipette.

· A few drops of indicator are added to the flask so that a colour change will occur when the neutralisation reaction between the acid and the base is complete.

· The other solution is added from a burette.

· A solution of accurately known concentration is called a standard solution.

· A volumetric flask is used to make up the standard solution.

· A standard solution can be reacted with a solution of unknown concentration using titration technique. One reactant in solution is slowly added to another reactant in solution until an end point is reached.

· The end point of the titration is usually indicated by a change in colour of a small amount of indicator solution added to the mixture of reactants. For an acid-base titration an indicator is selected that changes colour at the pH of the salt solution formed at the point of neutralisation. This is known as the equivalence point.

Titration Equipment

· At senior high school level equipment such as burettes, pipettes and volumetric flasks give readings to three significant figures. Calculations are carried out to three significant figures.

· A pipette is used to accurately deliver the required volume to a conical flask. If the pipette is to be rinsed, it is rinsed with the solution it is to measure. If the conical flask is to be rinsed, it is rinsed with distilled water.

· The volume measured out by a pipette is called an aliquot.

· The solution in the burette is called the titrant.

· The solution to be analysed, the one with the unknown concentration, is called the analyte.

· Burettes are used to deliver variable volumes of solution. The difference between the initial and final volume in the burette is the volume of solution delivered in the titration. If the burette is to be rinsed it is rinsed with the solution it is to contain.

· To perform a titration:

· The glassware is appropriately rinsed.

· An aliquot is delivered to the conical flask.

· The burette is filled via a small funnel. The burette is filled above the top mark so that some can be drained to fill the tube below the stop cock. The meniscus is then lowered to the 0 mls mark.

· A few drops of the appropriate indicator is added to the conical flask.

· While swirling the conical flask the solution is slowly added to the conical flask with gentle swirling.

· This process is continued until the equilavence point is reached and the indicator changes colour. This is the end point for the reaction and indicates that the titration is complete.

· The equivalence point is of a chemical reaction is the point at which the amounts of the two reactants are just sufficient to cause complete consumption of both reactants without either being left over.

Investigation 7.1

Worked Example 7.2 and Try These Yourself p199

Investigation 7.2 and 7.3

Check Your Understanding 7.3 p202

Investigation 7.4

From Module 2

· Standard Solutions

· A solution of accurately known concentration is called a standard solution.

· A volumetric flask is used to make up the standard solution.

· For a chemical to be suitable to prepare as a standard solution, it must:

· be a water soluble solid

·  have high purity - usually Analytical Reagent (A.R.) grade

·  have an accurately known formula  

· be stable in air, i.e. it does not lose or gain water or react with oxygen or carbon dioxide in air. E.g. deliquescent

· The solution is prepared by:

· accurately weighing a calculated amount of solid

· dissolving it in de-ionised water

· transferring all of the dissolved solid to a volumetric flask

· adding water to the flask to prepare a fixed volume of solution.

· The concentration is calculated in mol L-1

Dilutions

The equation used for dilution calculations is:

C1V1 = C2V2

Where

C1 = Initial concentration

V1 = Initial volume

C2 = Final concentration

V2 = Final Volume

Investigation 7.2

Worded example 6.9 p179, 7.2 and Try These Yourself p179

Worded example 7.1 p198, 7.2 and Try These Yourself p199

Chapter Review Questions p187: Q3, 4,5, 6, 8, 9 & 10

· The ‘end point’ or ‘equivalence point’ of the titrations is when the stoichiometric quantity of each reactant is present. The end point is identified by a permanent colour change of the indicator used.

Investigation 7.3

· 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

· When you titrate a strong base with a strong acid, for example sodium hydroxide with hydrochloric acid, the equivalence point will be pH 7. Bromothymol blue changes colour around pH 7 so it is a good choice.

· When titrating a strong base with a weak acid, for example sodium hydroxide with acetic acid, the equivalence point has a pH>7 so phenolphthalein, which changes colour around pH 9 is a good choice.

· For titrations involving a weak base with a strong acid, for example ammonium hydroxide and hydrochloric acid, the equivalence point will have a pH<7 and methyl orange will be a suitable indicator.

· Titrating with an indicator is not used for weak acids and weak base as there is no sudden change in pH to indicate the equivalence point.

· Using Conductivity Graphs

· The equilavence point of a reaction can be identified by measuring the change in conductivity of the analyte using

· A titration that uses changes in conductivity to determin equivalence point is called a conductometric titration.

· The advantages of using a conductometric titration are that it can be used:

1. With very dilute solutions

2. When species are at trace levels

3. With coloured or turbid solutions

4. For acid-base, redox, precipitation and non-aqueous titrations

· The electrical conductivity of a solution is proportional the concentration of ions in solution.

· The equivalence point is determined graphically by plotting change in conductance against volume of titrant added.

Students summarise “Back titrations” p208 and work through Worked example 7.3

· model neutralisation of strong and weak acids and bases using a variety of media

· Students use diagrams, computer simulations and molecular model kits to model neutralisation of strong and weak acids and bases using a variety of media.

Students use molecular model kits to model neutralisation.

Homework

https://www.youtube.com/watch?v=ANi709MYnWg

Start at 1’09”

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

· The general equations for the ionisation of a week acid is:

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

· The equilibrium constant for the ionisation of a weak acid is called the acid dissociation constant and is represented by Ka.

· For example hydrofluoric acid ionises according to the following reaction:

HF(aq) + H2O(l) H3O+(aq) + F-(aq)

And the equilibrium expression is:

Worked Example 6.6 and Try These Yourself p167 - 168

Chapter Review Questions p141: 6, 9 & 10

· explore acid/base analysis techniques that are applied:

· in industries

Student research

· by Aboriginal and Torres Strait Islander Peoples

· using digital probes and instruments

Investigation 7.4 using Vernier equipment

Instrumental analysis of vinegar

· conduct a chemical analysis of a common household substance for its acidity or basicity (ACSCH080) , for example:

· soft drink

· wine

· juice

· medicine

Practical – Using litmus paper and pH probes.

· conduct a practical investigation to prepare a buffer and demonstrate its properties (ACSCH080)

· A buffer is a solution that can resist pH change upon the addition of an acidic or basic components. It is able to neutralize small amounts of added acid or base, thus maintaining the pH of the solution relatively stable.

1. The buffer capacity is the amount of acid or base that can be added before the pH begins to change significantly. It can be also defined as the quantity of strong acid or base that must be added to change the pH of one litre of solution by one pH unit.

· Types of Buffer Solutions

· Buffers are broadly divided into two types – acidic and alkaline buffer solutions. Acidic buffers are solutions that have a pH below 7 and contain a weak acid and one of its salts. For example, a mixture of acetic acid and sodium acetate acts as a buffer solution with a pH of about 4.75.

· Alkaline buffers, on the other hand, have a pH above 7 and contain a weak base and one of its salts. For example, a mixture of ammonium chloride and ammonium hydroxide acts as a buffer solution with a pH of about 9.25. 

· Restating the above information buffers are:

· Weak acid and its conjugate base or

· Weak base and its conjugate acid

Investigation 7.9

· describe the importance of buffers in natural systems (ACSCH098, ACSCH102)

Students summarise: Buffers in the environment, pages 224 to 226.

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