Acids, alkalis, and salts

Example questions are taken from past papers.


4.1 describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Indicator	Colour in Acid	Colour in Neutral Solution	Colour in Alkali
Universal indicator	pH 0-3: Red (Strong) pH 4-6: Yellow (Weak)	Green (pH 7)	pH 8-10: Blue (Weak) pH 11-14: Purple(Strong)
Red litmus	Red	Red	Blue
Blue litmus	Red	Blue	Blue
Methyl orange	Red (Sometimes pink is accepted)	Yellow	Yellow
Phenolphthalein	Colourless	Colourless	Pink

Example question: State the colour of methyl orange in water contaminated with a small amount of nitric acid. (1m)

Mark scheme: Red (allow pink)

4.2 understand how the pH scale, from 0-14, can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline

pH 0-3: strongly acidic

pH 4-6: weakly acidic

pH 7: neutral

pH 8-10: weakly alkaline

pH 11-14: strongly alkaline

Example question: Suggest a possible pH value for water contaminated with a small amount of nitric acid. (1m)

Mark scheme: Any value in range 1-6.9. 

4.3 describe the use of universal indicator to measure the approximate pH value of a solution

pH 0-3: Red  (strongly acidic)

pH 4-6: Yellow  (weakly acidic)

pH 7: Green  (neutral)

pH 8-10: Blue  (weakly alkaline)

pH 11-14: Purple  (strongly alkaline)

Example question: Suggest why universal indicator is more suitable than methyl orange for comparing the acidities of samples of water. (1m)

Mark scheme: more than one colour in acid / indicates pH/ shows strongly or weakly acidic / shows how acidic the water is

Do not accept just more than one colour

My answer: Universal indicator can measure across the full pH scale and has colour change depending on acidity of the sample; but methyl orange is red in all acidic solutions, regardless of how acidic it is. 

4.4 define acids as source of hydrogen ions, H+, and alkalis as source of hydroxide ions, OH-

Example question: Hydrogen bromide and hydrogen chloride have similar chemical properties. 

(i) A sample of hydrogen bromide is dissolved in water.
A piece of blue litmus paper is placed in the solution. State, with a reason, the final colour of the litmus paper. (2m) 

Markscheme: 

Colour: Red/pink

Reason:  Hydrobromic acid formed/ H+ ions present

My answer: Red. They hydrogen atoms in hydrogen bromide dissociate in water to from H+ ions, which give the solution its acidity, thus forming hydrobromic acid. 

NB: Remember, it is the  H+ ions that are responsible for acidic properties. Hydrogen bromide by itself does not have  H+ ions, so it is not an acid.   

(ii) A sample of hydrogen bromide is dissolved in methylbenzene. 
A piece of blue litmus paper is placed in the solution. State, with a reason, the final colour of  

the litmus paper. (2m)

Markscheme: 

Colour: Blue

Reason:  No acid formed / no reaction / no H+ ions

My answer: Blue. Hydrogen bromide does not dissociate in organic solvents to from  H+ ions, so the solution is not acidic, and no acid is formed. 

4.5 predict the products of reactions between dilute hydrochloric, nitric and sulphuric acids; and metals, metal oxides and metal carbonates (excluding the reactions between nitric acid and metals)

Acid + alkali à salt + water (NEUTRALISATION REACTION)

• Hydrochloric acid + sodium hydroxide à sodium chloride + water

Acid + base (metal oxide) à salt + water (NEUTRALISATION REACTION)

• Sulphuric acid + copper oxide à copper sulphate + water

Acid + metal carbonate à salt + water + carbon dioxide

• Nitric acid + sodium carbonate à sodium nitrate + water + carbon dioxide

Acid + metal à salt + hydrogen

• Hydrochloric acid + magnesium à Magnesium chloride + hydrogen

4.6 recall the general rules for predicting the solubility of salts in water:

i. all common sodium, potassium and ammonium salts are soluble 

ii. all nitrates are soluble

iii. common chlorides are soluble, except silver chloride

iv. common sulphates are soluble, except those of barium and calcium

v. common carbonates are insoluble, except those of sodium, potassium and ammonium 

NB: any group 1 metal compound is soluble--they are alkali metals--dissolve in water to form alkalis!

4.7 describe how to prepare soluble salts from acids

acid + metal/metal oxide/metal carbonate

• Add acid to the excess metal/oxide/carbonate (NB: By using an excess of the metal or metal compound, you ensure that there is no acid left in the solution after reaction, so the filtrate will be a pure solution of the salt.)
• Stir
• Filter to remove excess solid
• Gently heat to evaporate the water to saturation point
• Dry the crystals with a paper towel/oven or in a dessicator 

4.8 describe how to prepare insoluble salts using precipitation reactions

Preparation of an insoluble salt:

1. Add the sodium salt solution of the anion to the nitrate salt solution of the cation until no more precipitate forms.
2. Filter to collect the residue
3. Wash the residue with cold water
4. Leave residue to dry on filter paper/dry in a warm oven

Example: Describe how to prepare a dry solid sample of silver chloride, AgCl, a salt which is insoluble in water. 

Sodium chloride + silver nitrate à silver chloride + sodium nitrate

NaCl (aq) + AgNO₃(aq) à AgCl (s) + NaNO₃ (aq)

1. Add dilute sodium chloride (chloride is the anion) solution to dilute silver nitrate (silver is the cation) solution until no more precipitate forms. 
2. Filter.
3. Wash residue with cold water.
4. Dry in a warm oven. 

4.9 describe how to carry out acid-alkali titrations

Titration:

1. fill the acid up to the mark in the burette
2. pipette 25.0cm³ sodium hydroxide into a conical flask
3. add a few drops of methyl orange indicator
4. add acid from the burette drop wise with swirling of flask
5. stop when colour change is permanent (turns pink/red)
6. note burette readings
7. repeat until concordant results are obtained (results are within 0.1 of each other)
8. take average of results

If they ask you how to prepare a soluble salt using an acid and an alkali, titration must be used. You first carry out a normal titration, and find out the exact amount of acid needed to neutralise the alkali. Then you repeat it without an indicator so that the salt is not contaminated with its colour. You remove the salt from the neutralised solution by evaporation, then you dry it. 

Rates of Reaction

Different Speeds of Reaction

Different chemical reactions take place at different speeds.

Very fast:

• explosion of a petrol-air mixture
• precipitation reactions
• fireworks going off
• coal burning

Moderately fast:

• reaction of metals or carbonates with dilute acids

Slow: 

• rusting of iron in air
• reaction of magnesium with cold water
• oil forming
• silver tarnishing

How do we calculate the speed of reaction? 

During a chemical reaction, the reactants get used up as products are formed. We can measure the speed of reaction by measuring the amount of a reactant used up per unit time, i.e.

Speed of reaction= amount of reactant used up/time taken

The speed of a reaction can also be measured in terms of the amount of a product formed/obtained per unit time, i.e. 

Speed of reaction= amount of product formed/time taken

For a chemical reaction that produces a gas, the speed of reaction can be found by measuring the volume of gas produced per unit time, i.e.

Speed of reaction= volume of gas produced/time taken

Measuring Speed of Reaction from Changes in Volume

The speed of a reaction can be found by measuring the following quantities at regular time intervals:

• volume of gas produced by the reaction
• the mass of the reactant that remains

Measuring Speed of Reaction from Changes in Mass

The speed of a reaction can also be found by measuring the changes in mass of a reaction mixture. This method works best for reactions which produce gases such as carbon dioxide. 

Factors Affecting Speed of Reaction

Many factors affect the speed of a chemical reaction. These include:

• the concentration of the reactants
• the pressure of the reactants (if the reactants are gaseous)
• the particle size or surface area of the reactants
• the temperature at which the reaction is occurring

Concentration:
Increasing the concentration means there's more of the reactants in the same volume. Thus with more particles there will be more collisions, and there will be higher chance of effective collisions, thus increasing the rate of reaction.

Pressure:
This is just like concentration, but it's more to do with gaseous reactants. With higher pressure--more particles--more collisions--higher chance of effective collisions--increase rate of reaction.

Surface area:
The bigger the surface area, the more chance of collisions, ditto above. Same thing. It's all about having more collisions, thus more effective collisions which increase rate of reaction. E.g. if you had a piece of metal, and you cut it into many smaller pieces, it would react faster with acid because the acid will have more area of the metal to collide with.

Temperature:
The higher the temperature, the more thermal energy is transferred to kinetic energy for the particles, thus they move faster and collide more often. If they have more energy, it is likelier that they will have minimum activation energy therefore there will be more effective collisions. And for a change, the rate of reaction increases... :P


For a reaction to occur between 2 particles: 

1. the reacting particles must collide with each other
2. they must collide with a certain minimum amount of energy known as the activation energy

In this way, collisions between reacting particles result in the formation of product particles. These collisions are known as effective collisions. 

Thus in a reaction between hydrogen and chlorine, only fast-moving molecules with energies equal to or greater than the activation energy will react on collision to form hydrogen chloride. 

• In general, when any factor increases the rate of effective collisions between reacting particles, it will also increase the speed of reaction. 

Energetics

4.13 understand the use of ΔH to represent molar enthalpy change for exothermic and

endothermic reactions

4.14 represent exothermic and endothermic reactions on a simple energy level diagram

4.15 recall that the breaking of bonds is endothermic and that the making of bonds is

exothermic

Just remember this! To break smth, obviously you need energy, so you need to take it in, hence breaking of bonds is endothermic. When substances form bonds, it's usually to become more stable, and thus they release energy. E.g. reactive elements like sodium react with chlorine to form sodium chloride, your common table salt, and this is a very stable ionic compound, and the reaction is exothermic. (less energy, less violently reactive...makes sense right?)

http://chemistry.about.com/cs/generalchemistry/a/aa051903a.htm

Heat Changes in a Reaction

heat change/enthalpy change: the amount of energy involved in a reaction, measured in kilojoules (kJ) which is 1000 joules, and is represented by the symbol DH. (D is the Greek letter for 'delta', which means change. H means energy content. 

For an exothermic reaction DH is negative. This is because the chemicals have lost energy to the surroundings. 

For an endothermic reaction, DH is positive. This is because the chemicals have gained energy from the surroundings. 

Examples:

1. When one mole of carbon is burnt in excess oxygen, 349 kJ of heat is produced. This is an exothermic reaction. So DH=-349kJ. 

 carbon + oxygen à carbon dioxide

C(s) + O₂ à CO₂ (g)   DH=-349kJ

2. When one mole of hydrogen reacts with one mole of iodine, 52 kJ of heat is absorbed from the surroundings. This is an example of an endothermic reaction. DH=+52 kJ

 hydrogen + iodine à hydrogen iodide

H₂ (g) + I₂ (g) à 2HI (g)   

Energy Level Diagrams for Exothermic and Endothermic Reactions

Energy level diagrams--convenient ways to express energy changes in a reaction

Exothermic Reaction 

Consider an exothermic reaction, heat energy is lost to surroundings. This means that total energy of the products is less than that of the reactants. 

The energy level diagram of an exothermic reaction

The difference between the energy levels of the products and the reactants is equal to the amount of energy given out by the reaction. 

i.e.  DH=H_products –  H_reactants  (as the energy of the reactants is larger, DH becomes negative. --> taking away a larger value from a smaller one gives a negative result. e.g. 2-4=-2) 

NOTE: USE THE SAME EQUATION FOR ENDOTHERMIC REACTIONS. It is always: 

 DH= H_products –  H_reactants

Endothermic Reaction

Since an endothermic reaction absorbs heat from the surroundings, the products will have more energy than the reactants. The difference between the energy levels of the products and reactants is the energy absorbed during the reaction.

The energy level diagram of an endothermic reaction

Activation Energy "E_a" - The minimum energy required to initiate (start) a chemical reaction. Both endothermic and exothermic reactions require activation energy.


So when particles collide and react, this is called an effective collision. This only happens when the particles have the minimum activation energy, if not, they may collide but it wouldn't result in a reaction.

Exothermic Reaction	Endothermic Reaction
Gives out heat to the surroundings	Takes in heat from the surroundings
Causes an increase in temperature	Causes a decrease in temperature
Has a negative DH	Has a positive DH
Has products lower in energy than the reactants	Has products that have higher energy than the reactants

Exothermic Reaction

Endothermic Reaction

Gives out heat to the surroundings

Takes in heat from the surroundings

Causes an increase in temperature

Causes a decrease in temperature

Has a negative DH

Has a positive DH

Has products lower in energy than the reactants

Has products that have higher energy than the reactants

This is a great, concise video to summarise everything, watch! 

Exothermic and Endothermic Reactions

Energy cannot be created or destroyed. However, it can be changed from one form to another. Energy changes occur in chemical reactions and even in some physical processes, such as when a solid dissolves in water.

4.10 recall that chemical reactions in which heat energy is given out are described as exothermic and those in which heat energy is taken in are endothermic

Exothermic Reactions

Reactions that give out heat energy to the surroundings are called exothermic reactions.

What are the characteristics of exothermic reactions?

When an exothermic reactions occurs,

• heat is liberated and is transferred from the chemicals to the surroundings, and 
• the temperature of the reaction mixture rises. The container feels hot.  

In an exothermic reaction, the temperature of the reaction mixture initially rises until the highest temperature is reached. When the reaction is completed, the temperature of the reaction mixture falls until it reaches room temperature.

Examples of exothermic reactions include:

• the combustion of fuels
• the rusting of iron
• the corrosion of metals
• the reaction between acid and alkali (neutralisation)
• respiration

Endothermic Reactions

Not all reactions give out heat. Reactions that absorb heat from the surroundings are called endothermic reactions. 

What are the characteristics of endothermic reactions?

When an endothermic reactions occurs, 

• heat energy is absorbed and is transferred from the surroundings to the reactants, and, 
• the temperature of the reaction mixture falls. The container feels cold. 

Initially, the temperature of the reaction mixture falls until the lowest temperature is reached. When the reaction is completed, the temperature of the reaction mixture rises until it reaches room temperature. 

Examples of endothermic reactions include:

• photosynthesis
• the action of light on silver bromide in photographic film
• thermal decomposition

Equilibria

4.22 recall that some reactions are reversible and are indicated by the symbol ⇌ in equations


4.24 understand the concept of dynamic equilibrium


4.25 predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions.


A reversible reaction is one that can occur in both directions. It means that the reaction can go both forwards and backwards-products can change back into reactants.

It is given by this symbol.

For example ammonia is made by a reversible reaction.

                        N₂ + 3H₂ ⇌ 2NH₃

Reversible reactions occur in both directions. Nitrogen and Hydrogen can make ammonia but ammonia can break down into nitrogen and hydrogen again. Each reversible reaction has a state of dynamic equilibrium, where there is an equal amount of products and reactants being made so it would look like there is no net reaction (no overall reaction)-forwards and backwards reactions is happening at the same rate. 

Le Chatelier's Principal states that:

If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or partial pressure, then the equilibrium shifts to counteract the imposed change and a new equilibrium is established.

Basically: the equilibrium will shift left or right to reduce this change. When they say the equilibrium shifts to the right, it means it is in favour of the forward reaction. Moving to the right=more products produced. And vice versa.

So with ammonia: there are 4 moles of gas in the reactants, but 2 moles of gas (ammonia) in the products. So the reactants have more pressure as there are more molecules. If you increase the pressure, you move the equilibrium to the right hand side, producing more ammonia, because you are trying to decrease the pressure. And the products have less pressure because there are only 2 moles there.

Temperature: If we increase the temperature the reaction will favour the endothermic reaction (the one that takes in heat). In this case the reverse reaction (making nitrogen and hydrogen) is endothermic and would increase. The equilibrium is shifting to the left hand side.

So therefore, to increase the yield of ammonia we can increase the pressure and decrease the temperature.

Increasing the temperature favours the endothermic reaction.

Decreasing the temperature favours the exothermic reaction. 

Titration

Specification 4.9 describe how to carry out acid-alkali titrations

Titration:

1. fill the acid up to the mark in the burette
2. pipette 25.0cm³ sodium hydroxide into a conical flask
3. add a few drops of methyl orange indicator
4. add acid from the burette drop wise with swirling of flask
5. stop when colour change is permanent (turns pink)
6. note burette readings
7. repeat until concordant results are obtained (results are within 0.1 of each other)
8. take average of results