M6 ABR1

MODULE 6 ACID/BASE REACTIONS

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

Nomenclature of inorganic acids and bases

The system of nomenclature used for the naming of inorganic acids and bases follows the IUPAC system in the Red Book (Connelly, Damhus, Hartshorn, & Hutton, 2005). There are a number of common substances which may be found in a home or school laboratory which have the properties of an acid or a base. These include:

Common Acids

Hydrochloric acid (HCl)

Hydrofluoric acid (HF)

Sulfurous acid (H₂SO₃)

Sulfuric acid (H₂SO₄)

Nitrous acid (HNO₂)

Nitric acid (HNO₃)

Carbonic acid (H₂CO₃)

Phosphoric acid (H₃PO₄)

Common bases

Sodium hydroxide (NaOH)

Calcium hydroxide (Ca(OH)₂)

Barium hydroxide (Ba(OH)₂)

Sodium carbonate (Na₂CO₃)

Sodium bicarbonate (NaHCO₃)

Ammonia (NH₃)

Type of inorganic acid: Binary acids

These are acids that that contain hydrogen bonded with a non-metal element.

NOTE: You won’t find oxygen in binary acids. This is because that’s a different type of acids (oxyacids) which we explore later.

A common example of a binary acid that you have came across is hydrochloric acid, HCl.

To name binary acid, it is simple. Here are the steps.

Step 1: Name the hydrogen first. You do this simplifying the word to just ‘hydro’.

Step 2: Find the root name of the non-metal.

• You vocabulary list of root names of non-metals will grow as you gain exposure to them.
• We have attached a list below of the root names for common non-metals that you need to know when naming inorganic acids for the HSC Chemistry course.

Step 3: Combine ‘hydro’ and the root name of the non-metal. Start with hydro first, followed by the root name.

Step 4: Combine the hydro + rootname with ‘ic’

Step 5: Following the word you generated at the end of step 4, write ‘acid‘ after leaving a space between the two words.

List of root names for common non-metals (& transition metal)

Nitrogen (N) – Nitr

Phosphorus (P) – Phosph

Oxygen (O) – Ox

Iodine (I) – Iod

Fluorine (F) – Fluor

Chlorine (Cl) – Chlor

Bromine (Br) – Brom

Sulfur (S) – Sulf

Hydrogen (H) – Hydr

Carbon (C) – carb

Transition metal – Chromium (Cr) – Chrom

Transition metal – Manganese (Mn) – Mangan

Practical example of naming a Binary Acid

Let’s do an example. How do you name the HCl compound?

Step 1: ‘Hydro’

Step 2: The non-metal here is chloride. The root name for chloride is ‘chloro’

Step 3: Combine hydro and the root name. This gives ‘hydrochloro’

Step 4: Combine the root name with ‘ic’. This gives ‘hydrochloric’

Step 5: Add the word ‘acid’ following the name of the inorganic acid. This gives ‘hydrochloric acid’.

Hence, the answer to naming the compound, HCl, is hydrochloric acid.

Type of inorganic acid: Oxyacids

Another type of inorganic acid is oxyacids. Sometimes, you may see this type of acid being referred to as oxoacids.

Oxyacids contain oxygen atom(s).

In terms of common features shared between oxoacids and binary acids, both types have a hydrogen and non-metal component to them.

Step 1: Ignore the hydrogen. This means that we don’t need to name it, ie don’t write hydro.

Step 2: Find the root name of the non-metal group. Sometimes, instead of a non-metal, there may be a transitional metal instead. Such as chromium or manganese.

Step 3: Combine ‘ous’ or ‘ic’ with the rootname. Use ‘ous’ if the non-metal is in lower oxidation state. Use ‘ic’ if the non-metal is in the higher oxidation state. Generally, acids with ‘ic’ are more common than those with ‘ous’.

Step 4: Write acid following the acid’s name.

Practical example of naming an oxyacid

HNO3.

Step 1: We ignore the hydrogen.

Step 2: The root name of the non-metal, nitrogen, is ‘nitr’

Step 3:

The oxidation state of nitrogen here, in HNO3, is +5.

The alternative is HNO2 where nitrogen has an oxidation state of 3+.

This gives us ‘Nitric’, the one where the central nitrogen atom has the higher oxidation state.

• NOTE: We learnt about oxidation states in Preliminary HSC Chemistry Course.

Step 4: We add ‘acid’ to ‘nitric’. This gives us nitric acid.

So, our answer in naming the oxyacid, HNO3, is nitric acid.

Exceptions for the nomenclature of oxyacids

For oxyacids that have the phosphorus or sulfur non-metals, we need to modify their root names after Step 4.

For oxyacids containing phosphate central atom, we need to add ‘or‘ after its root name.

For oxyacids that have sulfur as the central atom, we need to add ‘ur‘ after its root name.

This means that INSTEAD OF:

H3PO4 = Phosphic acid, we call it phosphoric acid instead.

• That is, we added ‘or’ after phosphorus root name, phosph.

H2SO4 = Sulfic acid, we call it sulfuric acid instead.

• That is, we added ‘ur’ after sulfur’s root name, sulf.

Prefixes for oxyacids

Sometimes, you may encounter the need to use prefixes when naming oxyacids. However, in most cases, these oxyacids with prefixes are incorporated in a question without requiring you to name them.

So, either way, it is good to understand how the acids looks like if they carry a prefix in front of their name.

You will see two possible prefixes that is incorporated in front the name of oxyacids.

These prefixes are ‘hypo‘ and ‘per‘.

• By far, the most common oxyacid used in HSC Chemistry questions that have both of those prefixes are the oxyacid containing the chlorine atom.

Previously, we touched on the fact that we can have ‘ous‘ or ‘ic‘ after the root name of the non-metal based on its oxidation state. The one with the higher oxidation state have ‘ic’ after its root name whereas the other have ‘ous’.

• HClO2 = chlorous acid (oxidation state of chlorine atom is 3+)
• HClO3 = chloric acid (oxidation state of chlorine atom 5+)

If we remove an oxygen from chlorous acid, so that it becomes HClO, we need to add the prefix ‘hypo‘ in front of chlorous acid. Oxidation state of chlorine decreased from 3+ to 1+.

• So, HClO is called hypochlorous acid.

If we add an oxygen atom to chloric acid, so that it becomes HClO4, we add the prefix ‘per‘ in front of chloric acid. Oxidation state of chlorine increased from 5+ to 7+.

• So, HClO4 is called perchloric acid.

Naming Inorganic Bases

Typically, you will see bases either written with a metal or ammonium ion to begin with. This is then followed by a hydroxide ion.

In some cases, an oxide may be present rather than hydroxide ion. We will touch on oxides later.

Step #1: Name the metal or ammonium ion first, depending on the compound whether it has an ammonium ion or not.

Step #2: Add the word ‘hydroxide’ following the name of the metal or ammonium ion.

Example of naming inorganic base

Mg(OH)2.

Step #1: The name of the metal in the compound is magnesium, Mg.

Step #2: We add the word ‘hydroxide’ after magnesium.

So, the name of the base, Mg(OH)2, is magnesium hydroxide.

Other forms of acids and bases: Metallic and Non-Metallic Oxides

Note the following:

• Metallic oxides are basic because they can dissolve in water to form basic solutions.
• Example of metallic oxide: calcium oxide, CaO
• Non-Metallic oxides are acidic because they dissolve in water to form acidic solutions.
• Example: of non-metallic oxide: carbon dioxide, CO2

So how do you name these acids and bases? This was taught in preliminary HSC chemistry course. But we will quickly go through it again because we can.

So, we first name the metallic or non-metallic by stating the name of the element that is located more towards the left of the periodic table first. Take the metallic oxide, CaO as an example.

• In our example of CaO, calcium is written in the chemical formula first as well as named first because it is situated more towards the left of the periodic table than oxygen.
• Next, we write the root name of the second element in the compound. For oxygen, the root name is ox.
• Following this, we then add ‘ide‘ to the root name of the second element. So in our example, it will become ‘oxide’.
• After that, we add prefixes to the name of the first element and second element depending on the number of atoms it has.

For example, N2O4 is called dinitrogen tetroxide. ‘Di‘ prefix is used because there are two nitrogen atoms in a N2O4 compound and the ‘tetra‘* prefix is used because each N2O4 compound has four oxygen atoms.

• * Recall from Week 1 HSC Preliminary Notes: The prefix ‘tetra’ is shorten to ‘tetr’ if the prefix end with ‘a’ or ‘o’ AND the root name starts with ‘a’ or ‘o’.
• Here, tetra ends with ‘a’ and the root name ‘ox’ (for oxygen) starts with ‘o’.
• Hence, we cut off the last letter in the prefix ‘tetra’. Recall that this applies to other prefixes like penta, hexa, etc.

VIEW Videos:

• ABR#1 https://www.youtube.com/watch?v=HoMfv1eE2ck&feature=youtu.be [6.44 MINS] [ 6.44 mins]
• Introduction to Acids and Bases https://www.youtube.com/watch?v=vKn-3MyTU0U&list=PLIYT0Z0pGeLwJ6WjLN6kUdn1_FWySK-0E&index=2 [5.25 mins]
• How to name acids https://www.youtube.com/watch?v=VhgpkmAaiAk&index=27&list=PL0o_zxa4K1BWziAvOKdqsMFSB_MyyLAqS [10.51 mins]
• Acetic, citric and hydrochloric acids https://www.youtube.com/watch?v=edIFtQ_JvEs&index=31&list=PLB755F0C836855EAB [10.10 mins]

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

a) acids and bases

There are reactions involving acids and bases with which you MUST be familiar.

a) Acid + Base (neutralisation) reaction.

Acid and base reactions are called neutralisation reactions if they involve hydrogen (or hydronium ions) reacting with hydroxide ions to produce water. This can be expressed as:

H⁺_(aq) + OH^-_(aq) ⟶ H₂O_(l)

This is why water is produced as a product in the equation involving the reaction of hydrochloric acid with sodium hydroxide.

General word equation:

Acid plus Base gives Salt plus Water

Example:

Hydrochloric acid + sodium hydroxide ⟶ sodium chloride + water

Balanced species neutral chemical equation:

HCl_(aq) + NaOH_(aq) ⟶ NaCl_(aq) + H₂O_(l)

Net ionic equation:

H⁺_(aq) + OH^-_(aq) ⟶ H₂O_(l)

(Apply the NAAG SG CHPS solubility rules to determine whether the salt formed is soluble or insoluble.)

The salt formed from neutralisation can be neutral, acidic or basic depending on the strength of the acid and base used in the reaction.

To determine the acidity or basicity of a salt, dissolve it in water.

• If the salt give rise to a solution with a pH of less than 7, it is an acidic salt.
• If the salt give rise to a solution that has a pH of more than 7, it is a basic salt.
• If the resulting solution has a pH equal to 7, it is a neutral salt.

Suppose you mix the following acid and base solutions.

Equation 1: HCl_(aq) + NaOH_(aq) ⟶ NaCl_(aq) + H₂O_(l)

Sodium chloride will be the salt in the above scenario. It is a neutral salt as it yields a pH of 7 when you dissolve the ionic compound in water.

NOTE: It does not really matter if you use hydrogen ions or hydronium ions. This is because most of the acid and solutions that you will be handling are diluted with water.

For example, HCl_(aq) -> H⁺_(aq) + Cl^-_(aq)

The hydrogen ions can then react with water to form hydronium ions. This can be expressed as:

HCl + H₂O_(l) -> H₃O⁺_(aq) + Cl^-_(aq)

This is why pH can be expressed in two forms:

• pH = -log₁₀[H⁺] or
• pH = -log₁₀[H₃O⁺]

In general:

• A strong acid reacting with a strong base will give you a neutral salt (and water)
• A strong acid reacting with a weak base will give you an acidic salt (and water)
• A weak acid reacting with a strong base will give you a basic salt (and water)

VIEW Videos:

• ABR#3 https://www.youtube.com/watch?v=R8QpOu8noZ0 [6.57 mins]

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

b) acids and carbonates

General word equation:

Acid plus Carbonate gives Carbon dioxide plus Water plus a Salt

Example:

Hydrochloric acid + sodium carbonate ⟶ sodium chloride + water + carbon dioxide

Sodium chloride is formed as a salt. You will also see carbon dioxide bubbles as it evolves in the solution. Water will also be a product.

Balanced chemical equation:

2HCl_(aq) + Na₂CO_3(aq) ⟶ 2NaCl_(aq) + H₂O_(l) + CO_2(g)

Net Ionic Equation:

2H⁺_(aq) + CO₃^2-_(aq) ⟶ H₂O_(l) + CO_2(g)

Carbonates are substances that contain the carbonate polyatomic ion CO₃^2-.

Some examples of carbonates include:

• Calcium carbonate, CaCO₃
• Sodium carbonate, Na₂CO₃
• Magnesium carbonate, MgCO₃
• Ammonium carbonate, (NH₄)₂CO₃

When an acid reacts with a carbonate, you will get a salt, water and carbon dioxide gas as products.

Usually in the HSC course, you will see an acid reacting with a metal carbonate or ammonium carbonate.

If you have a metal hydrogen carbonate such as sodium hydrogen carbonate rather than sodium carbonate, the products produced are the same. That is, salt, carbon dioxide and water.

HCl_(aq) + NaHCO_3(s) -> NaCl_(s) + CO_2(g) + H₂O_(l)

VIEW Videos:

• ABR#4 https://www.youtube.com/watch?v=V3h8J9XDdfY [7.06 mins]

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

c) acids and metals

Reactions involving metals and acids are dependent on the activity of the metal and the concentration of the acid.

General word equation:

Dilute Acid plus Active Metal gives Salt plus Hydrogen gas

Example:

Hydrochloric acid + magnesium ⟶ magnesium chloride + hydrogen gas

Balanced chemical equation:

2HCl_(aq) + Mg_(s) ⟶ MgCl_2(aq) + H_2(g)

Net Ionic Equation:

2H⁺_(aq) + Mg_(s) ⟶ H_2(g) + Mg²⁺_(aq)

Acids will react with metals IF the metal is active enough to displace the hydrogen ions from the solution. This would mean that the solid metal will become metal ions by donating electrons, and the hydrogen ions will accept such electrons to become hydrogen gas. If the metal is not active enough to displace the hydrogen ions from solution, no reaction will occur.

The reaction between hydrochloric acid and magnesium metal can be expressed as follows:

2HCl _(aq) + Mg_(s) -> MgCl_{2 (aq)} + H_{2 (g)}

Magnesium chloride will be the salt in this case. The magnesium metal will displace the hydrogen ions from solution to form hydrogen gas, where the hydrogen ions are derived from the dissociation of hydrochloric acid.

HCl_(aq) -> H⁺_(aq) + Cl^-_(aq)

Each hydrogen ion will accept an electron donated from magnesium metal where two hydrogen atoms can bond to form hydrogen gas.

This means that, after displacement, the overall reaction will be:

Mg_(s) + 2H⁺_(aq) -> Mg²⁺_(aq)+ H_{2 (g)}

In preliminary HSC Chemistry, you have learnt about the metal reactivity series.

Reactive metals are those that readily and quickly react with acids. These reactive metals are positioned above hydrogen in the activity series table.

By losing electrons to form metal cations in solution, these metals achieve a more stable state (octet).

As you learnt in the Y11 course, as you move down a group on the periodic table, the number of electrons of elements increases significantly. These electrons occupy energy shells that are further away from the nucleus. As the number of energy shells increases, the effect of ‘shielding’ increases.

This means that the element can achieve a more stable state by losing those electrons that are loosely attracted to the nucleus in the high energy shells.

This means the metal is capable of donating electrons to hydrogen ions. This results in metal ions in solution and the displacing hydrogen ions into hydrogen gas.

Unreactive metals normally do not react with acids. The donation of electrons from these unreactive metals is less common. If the acid is a strong enough reducing agent, the unreactive metal may react with the acid.

VIEW Videos:

• ABR#5 https://www.youtube.com/watch?v=EQgBPnEvUQk [6.22 mins]

Notes

• Lemons are sour due to the presence of citric acid (C₆H₈O₇).
• Baking soda (NaHCO₃) is a weak base and tastes bitter.
• Some acids eg. HCl are strong electrolytes (conduct electricity well) whilst others eg. acetic acid - CH₃COOH are weak (don't conduct electricity well)
• Litmus is a natural dye (an indicator) used to identify acids and bases on the basis of its colour change. There are a number of different dyes which change colour in the presence of acids and/or bases, including phenolphthalein, methyl orange and bromothymol blue.
• The term salt refers to any ionic compound formed in a reaction between an acid and a base.

pH is a unitless value that is assigned to solutions in order to compare the relative acidity between solutions.

pH values can range from 0 to 14.

• Acidic solutions have a pH less than 7.
• Neutral solutions have a pH equal to 7.
• Basic solutions have a pH more than 7.

NOTE: pH values can exceed 14 or be less than 0 (negative), if the concentration of the strong acid or base is very high.

Indicators are used because their colour changes are useful in assisting in determining the approximate pH of a substance such as a solution. They are also useful to compare the relative acidity between different solutions.

Indicators are usually weak acids or bases, meaning they form an equilibrium in solution. In most cases, the indicator molecule reacts with water.

H₂In(aq) + H₂O(l) <-> 2H₃O+(aq) + In2- (aq)

‘ln’ is short for indicator. ‘H₂ln’ just means that the indicator has is bonded to two hydrogen atoms.

• In some cases, the indicator can have a valency of 1–.
• H₂In will have a unique colour compared to In2-.

If an acid is present or added into solution, the equilibrium position for the above reaction will shift to the left to minimise the increase in [H+]. The solution will develop the colour that H₂In gives in solution.

If a base is present or added in solution, the equilibrium position will shift to the right as the concentration of H3O+ ions decrease as they react with the hydroxide ion from the base to produce water.

The shift of the equilibrium position to the right will give the unique colour In2- which is observed as a colour change.

This is the reason for a colour change when doing titration experiments.

Some indicators change colour over a small pH range, other indicators change colour over a large pH range.

In the diagram below, Indicator A changes colour over a small pH range whereas Indicator B changes colour over a small pH range.

Which indicator is more accurate?

Indicator A changes colour (from red to blue and vice versa) over a smaller pH range than indicator B. It is a more accurate and more effective* indicator as, from the indicator’s colour change, the substance testing the pH for will have a pH of 5.5 – 6.5.

Since indicator B changes colour over a larger pH range (from 5.5 to 10) than indicator A, it is a less effective* and less accurate indicator than indicator A. It is harder to pinpoint the exact pH by observing the indicator colour.

• That is, the pH of the substance would be less certain due to large pH range over which the indicator changes colour.

* NOTE: The term ‘effective‘ for indicators depends on the purpose of using the indicator. To see if the substance is acidic or basic and nothing more specific, use an indicator that changes colour over the pH range between 7 to less than 7 to see if the substance is acidic. Alternatively, use an indicator that changes colour over the pH range from 7 to greater than 7 to see if the substance is basic. Another method of determining pH is using a litmus paper to easily distinguish if the substance is acidic or basic.

To use the indicator (A and B) to pinpoint the exact pH of the substance, then indicator A is more effective in doing the job than indicator B. However, indicator A still fails to pinpoint the exact pH as the substance can have a pH of anywhere from 5.5 to 6.5.

To pinpoint the exact pH, a pH probe can be used. It is more accurate than indicators.

Investigate the Gizmo below.

Suppose you have the following indicator dissolved in a beaker of water. This indicator can exhibit two colours which are red and yellow.

Equation 1: H₂In (aq) + H₂O (l) <-> 2H₃O⁺ (aq) + In^2- (aq)

• H₂ln colour = Red
• ln^2- colour = Yellow
• Note that ‘ln’ is short for indicator.
• ‘H₂ln’ just means that the indicator has is bonded to two hydrogen atoms.

Suppose that equation one proceeds to equilibrium when you dissolve the indicator in the solution containing the test solution or you wish to add the test solution into the indicator.

Either way, IF the test solution contains hydronium ions (e.g. adding acid to the beaker of water or the test solution already contains hydrogen ions when indicator was added), it will shift the equilibrium position of equation one to the left according to Le Chatelier’s Principle to minimise the increase in hydronium ion concentration. This would cause the solution change colour from yellow to red.

If the solution contains hydroxide ions, the hydronium ions formed from the dissociation of the indicator molecule will react with the hydroxide ion to form water.

This decreases the concentration of hydronium ions in solution and thus will cause the equilibrium position of equation one to shift to the right as per Le Chatelier’s Principle to minimise the decrease in [H₃O⁺]. This would cause the solution to change colour from red to yellow.

The indicator would undergo a colour change according to the equilibrium equation above.

Sometimes you may see this equation instead:

HIn _(aq) + H₂O _(l) <-> H₃O⁺_(aq) + ln^- _(aq)

• This just means that the indicator has a valency of -1. However, most indicators have a valency of -2 as shown above.

VIEW Videos:

• ABR#2 https://www.youtube.com/watch?v=7aNXb_sCqmk [5.23 mins]
• pH Indicators https://www.youtube.com/watch?v=G2JFXQZrOHc&list=PLIYT0Z0pGeLwJ6WjLN6kUdn1_FWySK-0E&index=3 [7.59 mins]
• pH of common substances: https://www.youtube.com/watch?v=b77Z64gYBHg&list=PLB755F0C836855EAB&index=3 [13.03 mins]
• Indicators and the pH scale: https://www.youtube.com/watch?v=zbsCY-IDkH4&list=PLB755F0C836855EAB&index=4 [14.24 mins]
• Indicators and their uses: https://www.youtube.com/watch?v=-V8kmrG6biw&list=PLB755F0C836855EAB&index=5 [10.09 mins]
• Testing a range of indicators: https://www.youtube.com/watch?v=njONTvCyVdU&list=PLB755F0C836855EAB&index=7 [12.49 mins]
• Solving problems: https://www.youtube.com/watch?v=MPxdjCxXXcA&list=PLB755F0C836855EAB&index=8 [8.44 mins]

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

Neutralisation reactions occur when an acidic substance reacts with a basic substance to form a salt and water. The hydrogen ion (proton) from the acid reacts with hydroxide ions from the base to form water.

This reaction has a number of important applications.

A. Applications in Everyday life

1. Antacids:

Neutralising stomach acids with an antacid is both an effective way to relieve heartburn as well as an example of an everyday application of a neutralisation reaction.

Our stomachs secrete an acid for the purpose of protein digestion. This acid is HCl. The protease enzymes in the stomach work most effectively at low pH, however if the stomach secretes HCl at the wrong time or in high quantities, stomach acid levels can rise, causing indigestion or heartburn. A quick way to ease the discomfort is to take an antacid. These may contain magnesium hydroxide Mg(OH)₂ or aluminium hydroxide Al(OH)₃ but may also contain a carbonate, eg sodium carbonate Na₂CO₃ in the alka-seltzer. Of course if a carbonate is used, then we know CO₂ will also be produced, which is why the drinking of an antacid solution may also bring up a burp.

Mg(OH)2(s) + H2O(l) -> MgCl2(aq) + 2H2O(l)

It is not only our stomachs that are sensitive to particular pH values. Many of the enzymes (biological catalysts) in our bodies have an optimum pH at which they function. If the pH is either too high or too low, the body must seek to neutralise the excess acid or base and restore the system to the correct pH range. This is also true for blood and urine.

2. Baking soda

Baking soda is used in the process of baking cakes and contains sodium hydrogen carbonate. When it makes contact with the buttermilk in the cake, a neutralisation reaction takes place. This is because NaHCO3 in baking soda can act as base and buttermilk is an acid.

3. Ant stings

Similarly, sodium hydrogen carbonate can be used to neutralise ant stings which contain formic acid. Formic acid is corrosive and can burn the skin. By neutralising the formic acid, the pain on the skin can be reduced.

NaHCO3(aq) + CH2O2(aq) -> NaCHO2(aq) + H2O(l) + CO2(g)

3. Gardening

Gardeners will often use their knowledge of neutralisation reactions to change the acidity of the soil. Some plants grow better in acidic or neutral soils whilst others prefer alkaline soils. This is particularly true of hydrangeas which act as living indicators of the soil pH. In alkaline soils, hydrangeas are pink, which the flowers are blue in acidic soils.

Slaked lime (calcium hydroxide) can be added to acidic soils, whilst gypsum or iron salts are added to alkaline soils.

The most common types of fertilisers are also produced in a neutralisation reaction; sulfate of ammonia (NH₄)₂SO₄ and ammonium nitrate.

B. Applications in Industry

Factories may dispose liquid waste that is often acidic. Calcium oxide (a base) is added to neutralise these acidic factory effluents before they are allowed to be disposed of outside the factory, as per government environmental regulations.

Some factories produces sulfur dioxide, which is treated by neutralising it with calcium oxide (lime) before it is allowed to be released from the factory and into the atmosphere. Sulfur dioxide is acidic.

In this case, the acidic oxide (SO2) reacts with an base (CaO) to form salt and water. This is an acid-base reaction.

SO2(g) + CaO -> CaSO3(s)

SIDE NOTE: This reaction does not produce water. According to Arrhenius’s theory, water must be present as a product for neutralisation reactions. However, according to Bronsted Lowry’s definition of acid-base reactions between acid and bases, water is not a compulsory product. The Bronsted-Lowry’s definition of neutralisation requires a conjugate acid and base to be produced.

2. Production of fertilisers that contains ammonium sulfate, (NH4)2SO4.

Ammonium sulfate is produced by reacting ammonia gas with sulfuric acid. This is a multi-step process as you need to combine hydrogen and nitrogen gas to form ammonia. You also need to process the air to extract hydrogen.

2NH3(g) + H2SO4(aq) -> (NH4)2SO4 (aq)

Other industrially important applications of neutralisation reactions include:

• Preparation of fabrics
• Preparation of esters
• Treatment of wastewater
• Cleaning of spills

VIEW Videos:

• ABR#6 https://www.youtube.com/watch?v=nZ9MMm9XuVM&index=9&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS [6.33 mins]
• Cleaning up spills: https://www.youtube.com/watch?v=GiMlKmrjGhs&index=33&list=PLeFSFSJ9WqSCKs3ELNcsclaXnKU0qvg7t [5.23 mins]
• Acid spills: https://www.youtube.com/watch?v=Pps9JIZXg2g&index=58&list=PLB755F0C836855EAB [11.07 mins]

Insert is from https://theacidicenviroment.weebly.com/everyday-uses-of-indicators.html

1.5 conduct a practical investigation to measure the enthalpy of neutralisation

All neutralisations are exothermic.

When the heat of neutralisation is measured for a range of strong acids and strong bases, the amount of heat released is always approximately 57 kJ per mole of water formed.

When a basic solution is added to an acidic solution, a reaction takes place that usually forms an ionic salt and water. The solutions are said to have been neutralised when the concentrations of the hydronium and hydroxide ions within the mixture become equal.

The net ionic equation for the reaction between a strong acid and a strong base, including the enthalpy change is:

H⁺ + OH^– → H₂O + 57kJ

A study of the heat energy released for various (strong) acid-base reactions reveal values which are very close to one another, suggesting that the same type of reaction is occurring each time, ie. the production of a salt and water, or more specifically, the transfer of a proton species from a proton donor to a proton acceptor.

A change in the enthalpy of a solution can be determined using the following equation: ΔH = -mcΔT. We will assume the final solution will have the same c value as pure water (a potential source of error).

We will be determining the standard enthalpy of neutralisation, that is, the experiment should be performed under standard conditions: surrounding conditions of 25 degrees Celsius and 1 atmospheric pressure or 100 kiloPascals (kPa).

Under these conditions, we measure change in enthalpy of neutralisation where exactly 1 mole of water is formed (the salt formed is just composed of spectator ions).

For example, suppose there is a neutralisation between sodium hydroxide and hydrochloric acid:

NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l)

Ionic equation:

• Na+ (aq) + OH– (aq) + H+(aq) + Cl– (aq) -> Na+(aq) + Cl–(aq) + H2O(l)
• All these reactants and products are mixed together in the same solution or system.

Removing the spectator ions:

H+(aq) + OH– (aq) -> H2O (l)

The formula to determine the change in enthalpy in a neutralisation reaction is:

ΔHneutralisation = - m x c x ΔT

• m is the total mass of acid and base combined, measured in grams
• c is the specific heat capacity of water, measured in Joules/gram/Kelvin (J/g/K)
• ΔT is the change in temperature (Final temperature after reaction – Initial temperature before reaction), measured in Kelvin
• ΔHneutralisation is the change in enthalpy, measured in kiloJoules (since specific heat capacity of water is in joules, divide answer by 1000 to convert into kiloJoules)
• The minus sign shows energy being released (T of surroundings will increase) = exothermic, the absence of - sign shows energy being absorbed ( - T change as energy from the surroundings is absorbed by the reaction) = endothermic
• NOTE: You may see the formula written as q or Δq. This is the energy of the system (water) before the molar calculation that converts it to ΔH

ΔH = Δq / n

VIEW Videos:

• ABR#7 https://www.youtube.com/watch?v=4VFEDDRwcYI&index=8&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS [4.42 mins]
• Neutralisation reactions: https://www.youtube.com/watch?v=psa0z2LW7t4&list=PLB755F0C836855EAB&index=58 [9.10 mins]

VIEW Videos:

• EAR#8 https://www.youtube.com/watch?v=demPq1ODYmA&index=7&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS [12.36 mins]
• The Arrhenius Theory https://www.youtube.com/watch?v=H7k6AXcE7K4&index=1&list=PLIYT0Z0pGeLwJ6WjLN6kUdn1_FWySK-0E [13.06 mins]
• Historical development: https://www.youtube.com/watch?v=3AvIr9bOJXE&list=PLB755F0C836855EAB&index=49 [5.59 mins]

VIEW Videos:

• EAR#9 https://www.youtube.com/watch?v=BUCvSBJleRU&index=6&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS [8.03 mins]

VIEW the Tour

Click on the box below for an overview of acid-base history

VIEW Videos:

• ABR#9 https://www.youtube.com/watch?v=BUCvSBJleRU&index=6&list=PLeFSFSJ9WqSB27YUnCfMsSpmNluCyhZbS [8.03 mins]
• Bronsted-Lowry theory: https://www.youtube.com/watch?v=4skFSXnZ-9g&list=PLB755F0C836855EAB&index=50 [6.06 mins]