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The experiment aimed to discover how different types of substances interact together: water, hexane, ethanol, potassium carbonate, copper sulphate, sodium, iodine, and potassium.
Water could be mixed with ethanol and form hydrogen bonds. It is stated by Lister and Renshaw (2000; P296) that ethanol has an -OH function group attaching to a hydrocarbon chain. The -OH group has polar property and makes ethanol easily form bonds with water. Heat is usually released during bond making.
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When water meets hexane, the attractions within hexane between molecules are weak, temporary van der Waal forces (Lane, 2009). However, a large amount of energy is needed to break hydrogen bonds in water (Lister and Renshaw, 2000). Therefore, water and hexane could not mix spontaneously.
Lister and Renshaw (2000) describe that it is more likely for liquids with similar intermolecular forces to be mixed together. However, when ethanol (polar) meets hexane (non-polar), dissolution may still happen to a small degree because ethanol (CH3CH2OH) might have van der Waal forces hydrocarbon interactions with hexane.
Mixing liquids and solids
When potassium carbonate meet the solution of water and ethanol, water would easily break hydrogen bonds with ethanol and form bonds with potassium ions. Hence, ethanol separates from water and two layers would form in the solution with no colour change. One liquid is ethanol and the other one is water with potassium ions.
When copper sulphate (white powder) meets water, the ionic bonds between copper ions and sulphate ions are broken easily since water could hydrate copper ions and sulphate ions. This reaction produce CuSO4·5H2O, which could be a blue crystal and a solution. Therefore, the solution’s colour would change into blue and solid crystal would appear as well.
Copper sulphate does not dissolve in either ethanol or hexane because these two substances do not have the ability to hydrate copper and sulphate ions. Also, it is defined by Trader China (2009) that copper sulphate is insoluble in ethanol that has no impurities.
ChemiCool (n.d.) states that iodine hardly dissolves in water and the solution colour would usually change from clear to yellow brown. The solution colour changes because iodine is not totally non-polar. It could have dipole and temporary van der Waal forces with water and make the solution become light yellow or yellow brown, depends on its concentration.
When ethanol meets iodine, physical dissolution happens and iodine would dissolve in ethanol without any reaction. The colour of the solution would turn to black brown
Metals reacting with water
Sodium is an active substance and usually reacts intensely with water. It is because it loses electrons easily. Heat is usually released as energy exchange in this reaction. Lenntech (2009) suggests that sodium is ignited by the heat, and hydrogen gas is produced. The small ball of sodium might be pushed forward by hydrogen gas releasing and reacting with oxygen in the air (Huang and Lin, 2009). It would generally burn with an orange flame. This reaction produces basic sodium hydroxide which would make phenolphthalein indicator turn red.
Potassium reacts violently with water as well. It is more active than sodium because its outer electrons are further from the nucleus and therefore less attraction with it, which leads to higher energy and eventually less stability (Lane, 2009). Potassium is similar to sodium in this reaction except that it reacts more intensely and burns with a purple flame.
Therefore, liquids mixing with liquids would mostly require similar intermolecular forces while bonds have to be broken and formed during a mixture of liquids and solids; active metals would react with water intensely.
Test 1 was done beginning with the temperature of the water being measured in the beaker using a thermometer. Then 1 cm3 of water was added into a large test tube using a pipette. Using a different pipette, 1 cm3 of ethanol was added to the tube. Finally, the tube was shaken gently and observed, the temperature of the water was measured again.
1cm3 of water was poured into a test tube as a start for test 2. 1cm3 of hexane was added into the tube. The tube was shaken gently and observed. The temperature was measured.
In test 3, 1cm of hexane was added into a test tube using a pipette. 1cm3 of ethanol was added into the tube using a different pipette. Then, the temperature was measured after the tube was shaken gently and observed. Another 1 cm3 of ethanol was added into the tube after the temperature was measured.
Mixing liquids and solids
In test 4, potassium carbonate was added into the tube from step 1, using a spatula. The tube was shaken and observed.
Test 5 was started with 1cm water being added into a small test tube. 1cm ethanol and hexane was poured into 2 other test tubes. Then copper sulphate was placed into each tube. The tube was shaken and observed.
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Test 6 was started with 1 cm3 water being poured into a small test tube. 1 cm3 ethanol and hexane was poured into 2 other test tubes. Afterwards, iodine was added into each tube using a spatula. The tubes were shaken gently and observed again.
Metals reacting with water
75cm3 water was poured into the conical flask as the beginning of test 7. Then sodium was put into water using the tweezers and observed. Phenolphthalein indicator was added in the solution
Test 8 was repeated, as in test 7, using potassium.
Tests in section were successfully done and matched the theory. In test 1, heat was liberated as a temperature rise measurement which indicates bond making.
During section B, all tests were well done except for the reaction between ethanol and copper sulphate in test 5. As mentioned before, pure ethanol would not react with copper sulphate. A milk white solution might be the result of using impure ethanol which would probably react with copper sulphate.
Tests in section C appeared to be almost the same as theory predicted. However, sodium burned with a tiny yellow flame instead of an obvious orange flame. This might be a result of the poor purity of sodium and slow release of hydrogen.
Errors that occurred during this practical might be the imprecise mass of substances and their impurity. Even though an accurate amount of liquid was measured, when it was poured into the tube, the residue in the pipette would affect the amount of liquid being used. Similarly, the mass of solids were not measured and it might affect the dissolution rate. Also, since equipment was likely to be not perfectly cleaned, the result solution might not be completely pure as well.
To summarize, this experiment was mostly done well and its aim has been reached. To reduce errors, attention to the purity of substances and the residue of liquids remaining during measuring process would be needed as well as accurate measurement of solid mass.
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