Precipitation reactions

Introduction
When both anions and cations in a solution combine, a precipitation reaction will occur to form a precipitate. Precipitates are ionic and insoluble solids that are the product of a precipitation reaction, an example of this is when carbonate ions in solution react with a solution containing magnesium ions, forming a precipitate, this can also be seen chemical equation form below.

Na2 CO3 (aq) + Mg SO4 (aq)  Na SO4 (aq) + Mg CO3 (s)

To determine whether a reaction produces a precipitate depends on the solubility rules. The solubility rules are guidelines on which ions form solids and which do not. If an ion is soluble it will remain as an aqueous form. If an ion is insoluble it will form a solid with a different ion from the other reactant. If all the ions are soluble in the reaction, then therefore no precipitate will be produced. The general solubility rules, are as follow:

Solutions which contain hydrogencarbonate ions react with hydrochloric acid, to form carbon dioxide gas, as seen below.

NaHCO3 + HCl  NaCl + H2O + CO2

Solutions containing carbonate ions also react with hydrochloric acid, to form carbon dioxide gas, as shown below.

Na2CO3 + 2HCl  2NaCl + H2O + CO2

To identify the amount of carbon dioxide produced from a reaction, a limewater test is used. The limewater will turn to a milky colour in the presence of carbon dioxide, which is a reactant of carbonate and acid or hydrogencarbonate and acid. The reason behind this is that when carbon dioxide reacts with limewater, it produces a white precipitate called calcium carbonate. To distinguish whether solutions contain hydrogencarbonate or carbonate, a solution containing magnesium ions can be added, as carbonate ions can form a precipitate in the presence of magnesium and hydrogencarbonate cannot.
In spite of this, some solutions can form a precipitate from being heated. An example of this can be seen as, no precipitate is formed from a mixture of a solution of hydrogencarbonate ions and a solution magnesium sulfate, However, when the solution of hydrogencarbonate ions and magnesium sulfate are heated, they form a precipitate. This is due to the fact that hydrogencarbonate changes to carbonate upon heating. This can be seen in chemical equation form below.

2NaHCO3  Na2CO3 + H2O + CO2
Na2CO3 + MgSO4  Na2SO4 + MgCO3

Because of this, if the solution is heated it can influence the reaction, as it increases the energy and frequency of collisions and therefore increases the rate of reaction.
Moreover, another factor that could influence the reaction is the concentration and amount of the hydrogencarbonate or carbonate in the solution. Because of collision theory, this means that both the frequency and the energy collisions would increase or decrease accordingly and therefore increase or decrease the rate of reaction.

Aim
To observe the precipitate formation in different carbonate reactions

Hypothesis
If carbonate and acid are combined they should form CO2, which can be distinguished with a limewater test.

If a reaction between a carbonate and metal sulphate is heated it will form a precipitate.

Materials
3 scoops of Sodium carbonate
3 scoops of Sodium bicarbonate
Distilled water
5ml of 2.0M Hydrochloric acid
Limewater
3.5 of 1.0M Magnesium sulfate
Test tubes
Test tube rack
Test tube holder
Stoppers for test tubes with plastic pipes
Marker
Bunsen Burner
Pipettes
Beakers

Procedure
1. Add 3 scoops of sodium carbonate to a test tube labelled ‘A’
2. Add 3 scoops of sodium hydrogencarbonate to test tube labelled ‘B’
3. Using a pipette add 5mL of 2M hydrochloric acid to each test tube.
4. Record observations
5. Then, swiftly assemble the apparatus for the limewater test.
6. Repeat steps 1, 2 and 3 but using the arrangement for the limewater test.
7. Record observations
8. Add 3 scoops of sodium carbonate to a test tube labelled ‘C’
9. Add 3 scoops of sodium hydrogencarbonate to test tube labelled ‘D’
10. Using a pipette add 3.5mL of 1.0M magnesium sulfate to each test tube.
11. Record observations
12. Heat the sodium carbonate and magnesium sulfate solution in test tube ‘C’
13. Heat the sodium hydrogencarbonate and magnesium sulfate solution in test tube ‘D’
14. Record observations

Results

Table 1: Observations of various Carbonate and Hydrocarbonate reactions

Carbonate Hydrocarbonate (Bicarbonate)
HCl added A moderate level of bubbles An array of violent bubbles
Limewater observations Appeared cloudy with a moderate level of bubbles Appeared cloudy and white
MgSO4 added Appeared cloudy and little particles could be observed Began to look cloudy but died down eventually
Heating observations Bubbled slowly from the bottom and no colour change Almost in an instant it became cloudy again and bubbled intensely

From this it shows that, when hydrochloric acid was added to a carbonate the reaction effervesced suggesting a gas was formed, this was also the case for when the hydrocarbonate. The limewater test with hydrogencarbonate was more cloudy and white than the carbonate. The results suggest that when the magnesium sulfate was combined with the carbonate it formed a precipitate, in contrast, the combination of hydrocarbonate and magnesium sulfate did not form a precipitate. When heated, carbonate and magnesium sulfate solution appeared to have remained the same, however, the hydrocarbonate and magnesium sulfate solution had formed a precipitate.

Discussion

When the hydrochloric acid was added to a carbonate the reaction effervesced suggesting a gas was formed, this is because solutions which contain carbonate ions react with hydrochloric acid, to form carbon dioxide gas, as is the same for solutions which contain hydrogencarbonate ions, hence, the solution hydrochloric acid and hydrogencarbonate also effervesced and produced carbon dioxide gas. This can be seen below.

NaHCO3 + HCl  NaCl + H2O + CO2

Na2CO3 + 2HCl  2NaCl + H2O + CO2

The limewater tests suggest that the hydrocarbonate produced the most carbon dioxide as the limewater test appeared more cloudy and white than the carbonate. This is because limewater turns to a milky colour in the presence of carbon dioxide, which is a reactant of carbonate and acid or hydrogencarbonate and acid.

The results suggest that when the magnesium sulfate was combined with the carbonate it formed a precipitate, in contrast, the combination of hydrocarbonate and magnesium sulfate did not form a sulfate, this is because carbonate reacts with magnesium sulfate to form a precipitate and hydrocarbonate does not react with magnesium sulfate. This can be seen below.

Na2 CO3 + Mg SO4  Na SO4 + Mg CO3

When heated, carbonate and magnesium sulfate solution appeared to have remained the same, however, the hydrocarbonate and magnesium sulfate solution had formed a precipitate, because the hydrocarbonate had changed to carbonate upon heating, as seen below.

2NaHCO3  Na2CO3 + H2O + CO2
Na2CO3 + MgSO4  Na2SO4 + MgCO3

The results reflected my hypothesis correctly, ‘If carbonate and acid are combined they should form CO2, which can be distinguished with a limewater test.’ and ‘If a reaction between a carbonate and metal sulphate is heated it will form a precipitate’. The hypothesis was supported because the carbonate and hydrochloric acid effervesced and formed carbon dioxide gas which was discovered using a limewater test. The other hypothesis was also supported as the combination carbonate and magnesium sulfate formed a precipitate.

Potential sources of error could consist of the concentration of carbonate/hydrogencarbonate in the solution, as the amounts of sodium carbonate/hydrogencarbonate are not the same due to the measurement units (scoops), therefore making it almost impossible to get 3 exact scoops into each test tube. This would result in different energy and frequency of collisions, and therefore increasing or decreasing rate of reaction depending on the concentration. This can be solved by measuring the amount of carbonate/hydrogencarbonate in a beaker or measuring and have a fixed measuring unit such as cm3 so that it can be more accurate, or even weigh the amount of carbonate/hydrogencarbonate so that it can be most accurate as possible. Another potential source of error could exist in spillage.

Source: Essay UK - https://www.essay.uk.com/essays/science/precipitation-reactions/


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