We should be measuring pH for all cases because Ka (pKa) is related to pH, according to our definition of pH. ()
Planning (b):
Experiment for investigating temperature:
Experiment for investigating carbon chain:
Data Collection:
Acid used: HCOOH
Concentration used: 0.05 mol dm-3
Volume used: 150ml
- Mass of HCOOH used: 2.3g (Molar mass of HCOOH: 46)
-
Added into 1000 cm3 of distilled water
Data Processing and Presentation:
From our starting formula of pH:
So in order to get pKa we can simply substitute in the value of pH and [HA] into the formula where HA is the organic acid used in the experiment.
Percentage error of concentration of HCOOH:
= Percentage error of moles + Percentage error of volume of water
= 0.005 / (0.05*1*46) * 100% + 0.2 / 1000 * 100%
= 0.22% + 0.02%
= 0.24%
Highest percentage error of pH:
= 0.01 / 1.87 * 100%
= 0.53%
Highest percentage error of pKa:
= -log [ 10^-(percentage error of pH) + percentage error of concentration ]
= -log [ 0.03% + 0.24% ]
= 0.57%
Highest percentage error of temperature:
= 0.5 / 6 * 100%
= 8.33%
It is easier to see the relationship of pKa against temperature by plotting a graph.
Conclusion and Evaluation:
Conclusoin:
pKa decreases as temperature increases. pH decreases as temperature increases. The acidity of acid increases as temperature increases.
Evaluation:
This proves the hypothesis above. According to the Le Chatelier principle, the position of equilibrium will shift to counteract the external changes made onto the closed system. The dissociation of acid is always at equilibrium, with the forward reaction endothermic. This is because the forward reaction involves bond breaking (which takes in energy). So if we increase temperature, the equilibrium will shift to the left and thus more H+ is made. The acidity increases.
(From )
The literature value of pKa is 3.75. Our values of pKa differs from 2.44 ~ 4.36. It has been proven that pKa must be different in different temperatures (constants in equilibrium is only constant in a given temperature). Therefore the literature of pKa must be a value at a particular temperature only (at 25° C because Kw 1*10-14 is at this temperature). Since my values cover this literature value within a considerable range of temperatures, I can say my values are all valid, yet it is impossible to actually calculate the overall percentage error.
There are several systematic errors within the apparatus which can cause an error in our experiment. The pH meter is a new apparatus we use and is very depentent in this experiment. The pH meter is so sensitive that the pH reading becomes very unstable once there is a change in acidity. When we take out the probe and put it into the acid to measure the pH, the reading is keep on fluctuating and it takes qiute a while to stabilize. By the time it is stable, temperature will have dropped already. This means that there is a possible time delay between pH and temperature. Also, when we put the probe into the acid, the remainings on the probe is added into the acid which contaminates the acid. The pH and/or the concentration might differ and we will no longer be certain that the concentration is constant or not throughout the experiment. The more we put the probe into the acid, the higher this error.
As we heat up the acid above room temperature by adding hot water into the beaker, the rate of evaporation increases. The higher the temperature the faster the evaporation of water. As water evaporates the concentration increases. This means that as we heat up the acid the concentration changes. We are then dealing with two variables and thus out experiment can no longer be a fair test.
The thermometer we used is only an alcohol one, which is less accurate then mercury ones. The level of precision of temperature recordings decreases.
Improvements:
The pH meter is very percise in which we cannot improve on the pH reading by using a different one. However we can still increase the accuray by taking more readings then the average of them.
The concentration is very accutare as well (with a percentage error of less than 0.5%). It is because we used a digital balance and a volumetric flask in which both are very precise.
The change in oncentration of the acid is not significant. However we can still improve on this. We can reflux the water vapour so that the water vapour will return to the beaker. The concentration will then stay relativiely stable. We can also use a higher concentration to make the concentration constant. However the pH formula will only work in a low concentration, we cannot use a high concentration. This means in this experiment we will have to accept a higher percentage error of concentration and a higher percentage of impurities.
The reading of temperature is the one with the highest uncertainty. The percentage error of temperature, when approaching zero, will increase (as high as 8.3%). Then, the accuracy of the thermometer is much lower than that of the pH meter. So if we want a more accurate graph, we can use a more accurate apparatus to measure temperature e.g. thermocouple or a mercury-thermometer.
As mentioned above, the probe might contain some remaing buffer solution when we put it into the acid. So if we want to improve on this, we should either leave the probe in the acid without taking it out or we clean it every time when we are using it. Since the probe can be easily broken, it is better to leave it in the acid until the whole experiment is finished.
Although not likely, the buffer solution might not be exactly pH 4.00. However we assumed that it is 4.00 and calibrated the pH meter using this. If the buffer solution is not exactly 4.00, then it will affect all our readings of pH. Although this will only shift our graph up a little bit, it is still worth noting.