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Titrating Sodium hydroxide with an unknown molarity, against hydrochloric acid to find its' molarity.

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Introduction

Titrating Sodium hydroxide with an unknown molarity, against hydrochloric acid to find its' molarity. P2aP2aAim: - To find out the concentration of sulphuric (VI) acid by performing a titration with sodium carbonate solution. Introduction Sodium carbonate is an alkali, meaning that it has basic properties and has a pH that is greater than pH7, which is neutral. All pHs that are lower than pH7 are considered to be acidic. The pH scale runs from pH1, to pH14. A base is a solution/substance, which receives hydrogen ions (H+, also referred to as a proton) willingly, whereas acids are proton donators, this means that they give hydrogen ions off when in solution. The more basic something is, the more hydrogen ions it can receive and the more acidic something is, the more hydrogen ions it gives out, pH is a measure of this. In this particular titration, we will be neutralising the acid (H2SO4) with an alkali (Na2CO3), which is a carbonate. When an acid and a carbonate are reacted, the resulting products are a salt, carbon dioxide and water. These equations show what reactants and products are involved in this experiment: excellent shows you are thinking coherently about the problem evidence for P11a H2SO4 (?q) + Na2CO3 (?q) --> Na2SO4 (?q) + H2O (?) + CO3 (g) 1 mole + 1 mole --> 1 mole + 1 mole + 1 mole Hydrochloric acid (?q) + Sodium hydroxide (?q) --> Sodium chloride (?q) + Water (l) The acid (H2SO4, in this case) will release protons when the base (Na2CO3) is added to it, which will receive them. Once the solution has been neutralised, it will no longer be able to receive any more protons. We can test the pH of solutions by using various different types of indicators. Good chemical knowledge P11a Methyl orange is the indicator that will be used within this experiment. When methyl orange is present in acidic solutions, it is red; in basic solutions it is yellow. ...read more.

Middle

Mean titre = 23.30 + 23.25 + 23.25 + 23.30 4 = 93.1 4 = 23.275 cm3 H2SO4 The concentration of sodium carbonate solution that was used in the titration can be worked out by using two equations: 1. Number of moles = mass (g) � Relative Molecular Mass (gmol-1) 2. Number of moles = volume of solution (dm3) x concentration of solution (moldm-3) I know that I used 2.65g of anhydrous sodium carbonate in order to make up the standard solution of sodium carbonate. From this knowledge I can work out the number of moles of sodium carbonate that is in 2.65g of anhydrous sodium carbonate. To do this, I can use the first equation. Number of moles = 2.65g � 106gmol-1 = 0.025 moles of Na2CO3 However, the mass that was used was 2.65g; this was because I made up a solution that would be enough for ten titrations that were using 25cm3 of sodium carbonate solution. Therefore I need to divide the figure I obtained above by ten in order to find the correct number of moles of sodium carbonate that was used. Number of moles = 0.025 � 10 = 0.0025 moles of Na2CO3 To calculate the concentration of the sodium carbonate solution that was used I need to use the second equation. However, it needs to be rearranged before I can use it: Concentration of solution (moldm-3) = number of moles � volume of solution (dm3) Concentration of solution = 0.0025 � 0.025 = 0.10 moldm-3 So, the concentration of the sodium carbonate solution that I made up to use in the titration was 0.10moldm-3. H2SO4 (?q) + Na2CO3 (?q) --> Na2SO4 (?q) + H2O (?) + CO3 (g) 1 mole + 1 mole --> 1 mole + 1 mole + 1 mole This equation shows what happened within the titration and from it I can see that 1 mole of sodium carbonate will react with 1 mole of sulphuric acid, therefore within the titration the number of moles ...read more.

Conclusion

1 The nature of Science, published by the Royal society for Chemistry, section 1. Table to show the concentrations of sulphuric acid that the class calculated. Number Concentration (moldm-3) 1 0.10 2 0.11 3 0.11 4 0.11 5 0.11 6 0.11 Total 0.65 Average Concentration (moldm-3) 0.65 � 6 = 0.108 = 0.11 The table shows the concentrations that were calculated by the rest of my chemistry class as well as the average concentration. As you can see the average concentration is 0.11moldm-3, which is the molarity that I calculated the sulphuric acid to be. Also, the chemistry technician made up the sulphuric acid to be 0.10moldm-3, as the concentration that I calculated it to be is only 0.01moldm-3 out, I think that it is safe to say that the concentration of sulphuric acid I calculated is reliable. Improvements The experiment could be improved by changing some procedures and equipment, if this equipment were available, then the experiment would become more accurate and precise. As well as this the results obtained would too be more accurate and reliable. * Using a colorimeter to assess the end point of each titration. A colorimeter works by calculating the percentage of light either absorbed or allowed (transmitted) through a solution. As the colour of the indicator changes, so will the percentage of light transmitted or absorbed. In each of the titrations we could take a certain percentage as our end point, this way the readings for the end points should be more reliable. * Using grade A burettes, volumetric flasks and pipettes. These have a lower percentage uncertainty to them as they are more precise and so more accurate. These would lower the overall percentage uncertainties within the experiment, making it more precise. The experiment could be further extended by titrating the sulphuric acid against other alkali solutions with a molarity of 0.1moldm-3 and by using different indicators, which may give different end points to the titrations. ?? ?? ?? ?? Hanna Cheung ...read more.

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