Electrolysis: Potassium Iodide (Molten)

Aim:

To find the oxidising and reducing agents of molten potassium iodide

Equipment:

1x U-tube made out of copper
A power supply
2x wires connecting power pack to carbon electrodes
2x carbon inert electrodes
Molten potassium iodide

Safety:

·         Wear protective gear such as thick boots, high durable safety glasses, lab coat and extremely durable gloves

·         Avoid contact with any molten potassium iodide because 680^oC > is a hot temperature

Procedure:

1.       Collect all equipment needed and make sure all safety precautions are done or set.

2.       Setup the apparatus the same way as the previous experiment (Electrolysis: Potassium Iodide (solution))

3.       Add molten potassium iodide to the copper U-tube and turn the power supply on.

4.       Observe and record all results.

Predicted Results:

Anode	Cathode
Collecting the gas and cooling it down allows us to see a purple/black coloured substance at room temperature	Silvery solid starts forming over the carbon electrode until there is no molten potassium iodide is left

Redox:

2I⁻ _(l) → I₂ _(g) + 2e⁻   Oxidising

2K⁺ _(l) + 2e⁻   → 2K_(s) Reducing

2I⁻ _(l) +2K⁺ _(l) → I₂ _(g) +2K_(s) Redox Equation

Discussion:

From the procedures of this experiment we can see right at the start that it is very dangerous because of many parameters. First of all, the temperature of the molten potassium iodide is extremely high. This is because of its ionic bond creating a very strong bond within the ions, and breaking these bonds requires high amount of energy such as heat. Therefore, because of these dangerous temperatures, safety equipment is compulsory.

Our group could not experiment this theory as the extreme temperatures, lack of equipment and lack of space all stopped us because of the high risks.

To conclude, we can say that this experiment is not practically possible to do in a school environment; however, this experiment is done on a massive scale in the mining industry allowing the separation of a compound to its elemental form such as potassium iodide.

Electrolysis: Potassium Iodide (Solution)

Aim:

To find the oxidising and reducing agent in an aqueous solution of potassium iodide

Equipment:

1x U-tube
Phenolphthalein indicator
0.5 mol L⁻ Potassium Iodide Solution
About 15ml of distilled water
A power supply
2x carbon inhert electrodes
2x wires connecting power pack to carbon electrodes

Procedure:

1.       Collect all equipment needed and make create an observation and reaction table

2.       The diagram below shows the basis of the apparatus’ setup.

3.       Once the setup looks like the diagram below add a mix of 15 ml of potassium iodide and 15 ml of the distilled water

4.       Turn the power (6V supply should be enough)

5.       Once you see some reaction take place, add a few drops of the phenolphthalein indicator to check for any type of bases that may be in the solution

6.      

Record all results and observations.

Results:

Anode	Cathode
Solution around anode becomes yellowish/brown in colour.	Solution around cathode bubbles and is clear until Phenolphthalein indicator is added then solution becomes purple.

Redox:

2I⁻ _(aq) → I₂ _(aq) + 2e⁻   Oxidised
2H₂O _(l) + 2e⁻   → H₂ _(g) + 2OH⁻ _(aq) Reduced
2I⁻ _(aq) + 2H₂O (l) → I₂ _(aq) + H₂ _(g) + 2OH⁻ _(aq) Redox Equation

Discussion:

This experiment has shown and taught us a few lessons. First off we can see that water is a stronger oxidising agent than potassium. For this reason water reacts at the cathode instead of potassium, therefore forming hydrogen gas and hydroxide ions. To prove this we added a few drops of phenolphthalein indicator which resulted the solution to turn to purple colour at the cathode, in other words meaning hydroxide ions as the phenolphthalein turns purple when it comes in contact with a basic solution and colourless with an acidic solution.

The anode has a very different story as the water is not a stronger reductant than iodine, therefore, we can say that iodine reacts at the ₂ is formed at the anode however our group could not confirm this to be Iodine gas as the solution turned brownish/yellow, and the WACE WA data sheet stated that I₂ in an aqueous solution turns to a brownish/yellow colour. Therefore, this would mean that the I₂ is actually an ion.

anode instead of the water. We can see that I

Our experiment resulted in near perfect results it was conducted more than once and all steps being performed as perfectly possible.

To conclude, we can see that water is a stronger oxidising agent compared to potassium, however it is a weaker reducing agent compared to iodine 

Oxidation and Reduction

Aim:

To see which metals react with dilute hydrochloric acid and also other metal’s nitrate solutions.

Equipment:

Refer to Exploring Chemistry page 164

Procedure:

Refer to Exploring Chemistry page 165

Results:

Reactants	Observation	Products
Magnesium + Hydrochloric Acid	Magnesium dissolved a little while making a fizzing sound. On pop test it popped	Magnesium Chloride + Hydrogen
Zinc + Hydrochloric Acid	Bubbled a little however on the pop test, the flame only became brighter but did not pop	Zinc Chloride + Hydrogen
Lead + Hydrochloric Acid	Created 2-3 bubbles during experiment and could not get pop on pop test	Lead Chloride + Hydrogen
Copper + Hydrochloric Acid	Nothing happened	Copper + Hydrochloric Acid

Reactants	Magnesium Nitrate	Zinc Nitrate	Lead Nitrate	Copper Nitrate
Magnesium	NVR*	Magnesium strip becomes black	Magnesium strip becomes darker while making a dark flack on itself	Magnesium strip becomes black
Zinc	NVR*	NVR*	Zinc Strip becomes black	Zinc Strip becomes black
Lead	NVR*	NVR*	NVR*	N/A
Copper	NVR*	NVR*	NVR*	NVR*

*NVR:  No Visible Reaction

Discussion:

Redox Reactions:

1.      Mg + HCl  -> MgCl + H₂
Mg + 2H⁺ -> Mg²⁺ + H­­_2­
Mg -> Mg²⁺ + 2e^-
2H⁺ + 2e^- -> H₂
Mg + 2H⁺ -> Mg²⁺ + H_2
­Oxidising Agent: H⁺                         Reducing Agent: Mg

2.      Zn + HCl -> ZnCl + H₂
Zn + 2H⁺ -> Zn²⁺ + H­­_2­
Zn -> Zn²⁺ + 2e^-
2H⁺ + 2e^- -> H₂
Zn + 2H⁺ -> Zn²⁺ + H_2
­Oxidising Agent: H⁺                         Reducing Agent: Zn

3.      Pb + HCl -> PbCl + H₂
Pb + 2H⁺ -> Pb²⁺ + H­­_2­
Pb -> Pb²⁺ + 2e^-
2H⁺ + 2e^- -> H₂
Pb + 2H⁺ -> Pb²⁺ + H_2
­Oxidising Agent: H⁺                         Reducing Agent: Pb

4.      Mg + Zn(NO₃)₂ -> Mg(NO₃)₂ + Zn
Mg + Zn²⁺ -> Mg²⁺ + Zn
Mg -> Mg²⁺ + 2e^-
Zn²⁺ + 2e^- -> Zn
Mg + Zn²⁺ -> Mg²⁺ + Zn
_­Oxidising Agent: Zn                         Reducing Agent: Mg

5.      Mg + Pb(NO₃)₂ -> Mg(NO₃)₂ + Pb
Mg + Pb²⁺ -> Mg²⁺ + Pb
Mg -> Mg²⁺ + 2e^-
Pb²⁺ + 2e^- -> Pb
Mg + Pb²⁺ -> Mg²⁺ + Pb
_­Oxidising Agent: Pb                         Reducing Agent: Mg

6.      Mg + Cu(NO₃)₂ -> Mg(NO₃)₂ + Cu
Mg + Cu²⁺ -> Mg²⁺ + Cu
Mg -> Mg²⁺ + 2e^-
Cu²⁺ + 2e^- -> Cu
Mg + Cu²⁺ -> Mg²⁺ + Cu
_­Oxidising Agent: Cu                         Reducing Agent: Mg

7.      Zn + Pb(NO₃)₂ -> Zn(NO₃)₂ + Pb
Zn + Pb²⁺ -> Zn²⁺ + Pb
Zn -> Zn²⁺ + 2e^-
Pb²⁺ + 2e^- -> Pb
Zn + Pb²⁺ -> Zn²⁺ + Pb
_­Oxidising Agent: Pb                         Reducing Agent: Zn

8.      Zn + Cu(NO₃)₂ -> Zn(NO₃)₂ + Cu
Zn + Cu²⁺ -> Zn²⁺ + Cu
Zn -> Zn²⁺ + 2e^-
Cu²⁺ + 2e^- -> Cu
Zn + Cu²⁺ -> Zn²⁺ + Cu
_­Oxidising Agent: Cu                         Reducing Agent: Zn

9.      Pb + Cu(NO₃)₂ -> Pb(NO₃)₂ + Cu
Pb + Cu²⁺ -> Pb²⁺ + Cu
Pb -> Pb²⁺ + 2e^-
Cu²⁺ + 2e^- -> Cu
Pb + Cu²⁺ -> Pb²⁺ + Cu
_­Oxidising Agent: Cu                         Reducing Agent: Pb

Metal Reducing Strength (descending):

1.       Magnesium

2.       Zinc

3.       Lead

4.       Copper

From our experiment we can see multiple things from the results table. First off, we can see the reducing strength of each of the metals experimented on and their reactivity with dilute acids. We can say that Copper is the least reactive out of the metals as it does not react with the dilute hydrochloric acid. It’s also has the weakest reducing strength. We know this by the second results table. It shows us that it does not react with any other metal nitrates. Magnesium on the other hand can definitely be considered as the most reactive and strongest reducing strength. We know this because it extremely reacted with the dilute hydrochloric acid. It also reacted with all three other metal nitrates. From this we can say that copper is the least reactive and has the weakest reducing strength while magnesium has the strongest reducing strength and is the most reactive with dilute acids. From our workbook Chemistry for WA and the results from above, we can say that zinc has the second strongest reducing strength and reactivity with dilute acids. Lead comes in third out of fourth.

Acids and Bases Assignment

Task 1:

Acids and base solutions do conduct electricity based on their concentration and their strength on the pH level. For example solution of hydrochloric acid is a good conductor of electricity as all the compound completely ionises. This means that hydrochloric acid becomes hydrogen and chlorine ions. Since this acid is very strong, it completely ionises meaning that there will be no hydrochloric acid left in the solution after the reaction is complete.  As the ions are charged it allows the electricity current to flow through the ions which means the solution will conduct. This is also the same with strong bases such as sodium hydroxide. As the ionic compound completely ionises again allows the electricity current to flow through the sodium and hydroxide ions. However, weak bases and acids do not ionize completely. Therefore, this means that there are molecules in the solution for example acetate ions will be in the solution. This means that because there are acetate ions that have not completely ionised, does not allow the electricity to flow through very well. For this reason we can say that both acid and base solutions both conduct electricity. However, because some of the acids or bases such as acidic acid or ammonia are weak; they will conduct less electricity compared to acids and bases such as hydrochloric acid and calcium hydroxide because a weak acid or base does not ionise completely like a strong acid or base.

The relation between electrical conductivity and concentration is based on the difference between the solution. For example, in pure water, the concentration between the hydrogen ions and hydroxide ions are equal. Therefore, pure water cannot conduct electricity. However, in a solution of hydrochloric acid and water, the concentration between the hydrogens and the hydroxides are different. In this case, there are more hydrogens than hydroxides meaning that the solution is acidic. This is the same if a basic solution is used. There will be a difference in concentration between the hydroxides and hydrogens. Therefore, because there are more hydroxides than hydrogens, means that the solution is basic. Both of these solutions were between an acid or a base and water. However, if there is a solution between an acid and a basic (hydrochloric acid and calcium hydroxide), the solution will become neutral. This is because the hydroxides in the calcium hydroxide will balance out the hydrogens in the hydrochloric acid. This however, means that the concentration in both acid and base have to be the same and its ratio between the acid and the base have to be equal. If you have 1L of hydrochloric acid with a 1 mole per litre concentration, you will need to have 1L of calcium hydroxide with a 1 mole per litre concentration to neutralise the solution.

Ammonia will have a much less electrical conductivity compared to sodium hydroxide. This is because the sodium hydroxide completely ionises while the ammonia partially ionises. This means that in a solution of ammonia and water, only some of the ammonia molecules ionise into hydroxide and ammonium in water while most of it stays in a ammonia molecule. However, with a sodium hydroxide solution in water, the sodium hydroxide completely ionises to hydroxide ions and sodium ions. Because all of the sodium hydroxide solution ionises, means that if there is 1 mole of sodium hydroxide, all of it will ionise. This means that the solution has a strong electrical conductivity. However, with the ammonia solution, if 1 mole of it is added, only approximately 0.004 moles of ionises into hydroxide and ammonium ions.
-----------------------------------------------------------------------------------------Task 2:

Acids and bases have many different attributes such as being strong or weak and being concentrated or dilute. A strong acid or base are the ones usually on the ends of the pH table. A strong acid is usually between a pH reading of 1-3 and a strong base is between a pH reading of 13-15. A strong acid or base completely ionises which in theory is a one way reaction. It ionises so well that it is impossible to reverse the reaction. However, with weak acids and bases, its located between 8-11 if its a base and 4-6 if its an acid. These acids and bases are known to be weak meaning that they will not ionise as well as the strong acids or bases. This also means that the solution created after the partial ionisation is reversible as there is still more weak acid or base left which have not been ionised. For example hydrochloric acid which is a strong acid becomes chloride ions and hydrogen ions. However, in acidic acid which is a weak acid, only some of it becomes acetate ions and hydrogen ions. Most of it will stay as acetic acid molecule. Concentration is the measurement of how much of that substance is in the solution. For example, there could be hydrochloric acid in water which has a concentration level of 1 moles per litre. This means that there is 1 mole of hydrochloric acid in one litre of the solution. Dilution however, does not really mean anything. It is used when a concentration level is to be decreased. For example, if we add more water to 1 litre of 1 mole per litre hydrochloric acid, then the 1 mole per litre decreases. This is also the same with the substance becoming more concentrated. If we take water out of the solution this time, then the concentration of the hydrochloric acid increases from the 1 mole per litre.

Strong electrolytes are the solutions that ionise very well, weak electrolytes are the solutions that do not ionise very well, while non-electrolytes are solutions that do not ionise at all. Strong electrolytes such as hydrochloric acid and sodium hydroxide completely ionise in a solution with water. This means that they break down into ions of hydrogen ions and chloride ions (hydrochloric acid). For this reason, strong electrolytes can conduct electricity much better than weak electrolytes because the electricity can go through the ions. Strong electrolytes are usually strong acids and bases. Weak electrolytes can be described as the opposite of the strong electrolytes. Its usually the weak acids and bases that do not completely ionise. Once the solution does not completely ionise means that there are some ions in the solution while the rest is in molecule form. Ammonia and acetic acid are perfect examples of a weak electrolyte. This is because they do not completely ionise meaning that there are some acetate ions and hydrogen ions in the solution while most of it stays as a molecule (Acetic acid). Because of the limited number of ions in the solution, it will not conduct as well because there are much less ions to carry the electricity. Non-electrolytes can be described as solutions that do not ionise as all. For this reason they will not conduct electricity at all. Some example of this are glucose (sugar) and ethanol. They will not ionise meaning that they will stay in a molecular form.
------------------------------------------------------------------------------------------Task 3:

The term soil pH means how acetic or basic the soil is. A farmer uses different substances to balance the acidity of the soil to their needs. These substances include agricultural lime and ammonium nitrate. For all scenarios here we will want the farmer to have a soil pH of between 6 to 8. For example, if soil pH reading on a farmers ground says 2, it will mean that the soil is very acetic. In these cases the farmer will use substances such as agricultural lime to balance the acidity of the soil. Agricultural lime is mainly made up of calcium carbonate. Because calcium carbonate is a basic solution, it will balance out the acidity of the soil. However, the farmer will use substances such as ammonium nitrate if the soil tends to be more basic than neutral. As ammonium nitrate is an acid, it will balance out the basicity of the soil.

Soil acidification is the build up of hydrogen ions in the soil resulting in it to decrease its pH level. Some of the causes for this acidity are the fact that the trees are more basic than neutral. This means that once taken out, the soil becomes more acetic. Also with trees, as the leaves drop on to the ground creates organic acids such as acetic acid. For this reason, the soil pH level decreases the more leaves there are on the ground. Another cause for this is existence of sedimentary rocks in the soil such as coal. As they are rich in sulfides, once hydrated, it will react to become sulfuric acid. As sulfuric acid is considered as a strong acid, increases the soil’s acidity levels.

As the soil’s acidity levels increase, a farmer may use a basic solution to balance the soil’s pH level to neutral again. uses of chemicals such as calcium carbonate. Sodium hydroxide is not recommended even though it is a stronger base. This is because sodium hydroxide is toxic which would not be the best decision for use in soil where all the plants will grow.

Acid or Base? Investigation

Acid or Base?
Aim:

To check different solutions electrolyte strength and seeing if the solution is an acid, base or neutral.

Equipment:

• test tubes
• beakers

0.100 mol L-1 of the following solutions:

• ammonium chloride
• ammonium acetate
• potassium nitrate
• sodium carbonate
• sodium acetate
• sodium bicarbonate
• sodium sulfate
• hydrochloric acid
• sodium hydroxide
• acetic acid
• limewater
• DC power source
• 2 electrical conducting metal rods
• ammeter
• pH indicator

Method:

1. place 10 ml of each solution in separate test tubes
2. place 3 drops of the pH indicator in each of the test tubes
3. mix every test tube gently
4. leave the test tubes to rest for a couple of minutes
5. observe the colour of the indicator
6. compare the colour to the pH table
7. record the data
8. place 50 ml of every solution in separate beakers
9. connect the DC power source so that positive and negative create a complete circuit with the ammeter and the solution
10. place the rods inside on beaker on the edges so that the rods don’t touch each other
11. turn on power source and read the value on the ammeter
12. record the data from the ammeter
13. clean and dry rod before placing it in the next beaker

Safety Tips:

• have safety glasses on at all times
• have the DC voltage at 6V
• do not touch the solutions
• do not drink the solutions
• do not directly inhale the solutions
• wash hands if solution spills on you
• tell teacher ASAP
• do not touch both the rods together as they will spark
• if the ammeter indicates a reading too far to the right, break the power source

Observations and Results:

Solution	pH Level	Acid or Base	Electrolyte Strength in mA
Ammonium Chloride	6	Weak Acid	30
Ammonium Acetate	7	Neutral	25
Potassium Nitrate	7	Neutral	30
Sodium Carbonate	10	Base	40
Sodium Acetate	8 - 9	Weak Base	15
Sodium Bicarbonate	9 - 10	Base	20
Sodium Sulfate	8	Weak Base	35
Hydrochloric Acid	2	Strong Acid	150
Sodium Hydroxide	11	Strong Base	70
Acetic Acid	3 - 4	Weak Acid	5
Limewater	11	Strong Base	20

Discussion:

From the results table above we can see that stronger bases and acids are better electrical conductors. We know this as the strong bases and acids completely ionise. Therefore, creating a passage for the electricity to move through. Also from the table we can see that the weaker bases and acids are not as good as the strong bases and acids in conducting electricity. This is because the weaker bases and acids only partially ionises. This means that the solution has some ions but most of it still stays in molecular form. As there are much less ions to carry the electrical energy, means that it is not as good as a strong base or acid. For example in a solution of hydrochloric acid, it completely ionises to hydrogen ions and chloride ions.

HCl → H+ + Cl-

However in a acetic acid solution only some of the molecules ionise.

CH3COOH ⇆ H+ + CH3COO-

In a 1 mole per litre situation, the hydrochloric acid completely ionises meaning that the whole 1 mole of hydrochloric acid become hydrogen and chloride ions. However, for acetic acid in the same situation, only 0.004 mole of it ionises into hydrogen and acetate ions. The difference between the 1 mole and 0.004 mole means that the 1 mole which is the hydrochloric acid is a better electrical conductor.

Also from the results we can see that some neutral solutions conduct electricity. Because a solution is neutral does not mean that it will not conduct. As both the neutral solutions ionise, it allows the electricity to flow through. This means that for the solution to ionise it should be either a strong base, strong acid or an ionic compound.

A problem within the results is the electrolyte strength for limewater. From the indicator and formula it has to be a strong base. However its conductivity showed to be weaker than most weak bases. This is probably because of error in the experiment while test were being conducted. It should have the same electrolyte strength as sodium hydroxide.

Another problem within this experiment was the supply of the pH indicator table. It did not cover the full range of pH levels. For example, the once the pH indicator was added to the hydrochloric acid, it was more red compared to the lowest pH level of the table which was 3. For this reason this created a small confusion because the acetic acid had a reading of 3 which seemed to be a strong acid as it was right at the bottom of the pH table.

Conclusion:

To conclude, were able to get the electrolyte strength and pH level of each solution with the aid of a DC power source and a pH indicator. The results above are quite reliable with a few exceptions as they were all tested twice. Assumptions such as contamination will have to be taken into consideration as it is practically impossible to not contaminate a solution as some of the solution may not wash of the rods.