gerald_mcdonald

but would the recorded concentration of HCl be lower, higher or unchanged...and how would you know? Still not getting it... i mean i know that the the HCl in the burette would be more diluted and hence be of lower in concentration but would this be the recorded concentration for HCl?

Washing soda is added to hard water to allow soap to lather. A brand of washing soda contains partially hydrated sodium carbonate solid. A 0.300 g sample completely reacts with 20.0 mL of 0.250 M hydrochloric acid. a) What mass of sodium carbonate was present? b) Calculate the percentage by mass of sodium carbonate in the washing soda. Please. Any help would be appreciated! How do I do this question? Would anyone be able to give me some hints/advice on how to tackle this question?

It would dilute the concentration of the solution. However what do you mean by "w/e"? What's the logic behind the question? I want to know particularly for future reference. It sounds like a typical exam type question. For question c, if there is water left over in the burette from a previous washing with water it would dilute each titre and as a result require a slightly greater amount of recorded HCl to be used in the titration. For it to react and reach the stochiometric mole ratio, more recorded HCl, in litres, will be used. But still not seeing the logic.

The concentration of sodium hydroxide in waste water from an alumina refinery was found by titrating 20.00 mL aliquots of the water against 0.150 M hydrochloric acid, using phenolphthalein as indicator. The average titre of several titrations was 11.40 mL. a) Why is an indicator used? What can I say to make a well response answer? I'm guessing I should use the words "end-point" and "equivalence point" and how they have to be relatively equal for the experimenter to obtain accurate results. But how can I explain in words why this is the case? Here is the answer I have written: "An indicator is used in the titration so that when the titration reaches a point where all the reactants have reacted with each other in their correct mole ratios (i.e. the equivalence point), a colour change is produced straight after. This is termed the endpoint and if the indicator had been chosen wisely for experiment then it should have a sharp endpoint. This allows..." However how can I make it better? How would you answer this question? b) Write an equation for the reaction that occurred. [latex]NaOH(aq) + HCl(aq)[/latex] ---> [latex] NaCl(aq) + H2O(l)[/latex] c) What was the molarity of sodium hydroxide in the waste water? [latex]c_{1}V_{1}=c_{2}V_{2}[/latex] [latex]c_{1}*0.02000=0.150*0.01140[/latex] [latex]c_{1}*0.02000=0.00171[/latex] [latex]c_{1}=0.855 M[/latex] d) What mass of sodium hydroxide would be present in 100 L of the waste water? [latex]c(NaOH)=0.0855 M[/latex] [latex]V(NaOH)=100 L[/latex] [latex]n=cV[/latex] [latex]n(NaOH)=0.0855*100=8.55 mol[/latex] [latex]m=n*M[/latex] [latex]m(NaOH)=8.55*(22.9898+15.9994+1.00797) = 341.9758035 = 342 g[/latex]

Could you do it? I don't understand what you're saying. And aren't hydroxides strong bases not strong acids?

What mass of sodium sulfate is produced when 25.0 mL of 0.100 M sulfuric acid is added to 20.0 mL of 0.15 M sodium hydroxide solution? [latex]n=cV[/latex] [latex]n(H_{2}SO_{4})=0.100*0.0250=0.00250=2.50*10^{-3} mol[/latex] [latex]n=cV[/latex] [latex]n(NaOH)=0.15*0.0200=0.0030=3.0*10^{-3} mol[/latex] [latex]\frac{n(H_{2}SO_{4})}{n(NaOH)} =\frac{1}{2}[/latex] [latex]n(H_{2}SO_{4})[/latex] required [latex]= \frac{1}{2} * (3.0*10^{-3}) = 0.0015 = 1.5*10^{-3} mol[/latex] [latex]n(H_{2}SO_{4})[/latex] is in excess by [latex](2.50*10^{-3}-1.5*10^{-3}) 1.0*10^{-3} mol[/latex]. NaOH is the limiting reagent. [latex]\frac{n(Na_{2}SO_{4})}{n(NaOH)} = \frac{1}{2}[/latex] [latex]n(Na_{2}SO_{4})=\frac{1}{2}*(3.0*10^{-3}) =0.0015=1.5*10^{-3} mol[/latex] [latex]m=n*M[/latex] [latex]m(Na_{2}SO_{4})=\frac{1}{2}*(1.5*10^{-3}) * (2*22.9898+32.064+15.9994*4)= (1.5*10^{-3})*142.0412[/latex] [latex]=0.2130618[/latex] [latex]=0.21 g[/latex] Did I make any mistakes?

What volume of 0.100 M sulfuric acid would be required to neutralise a solution containing 0.500 g of sodium hydroxide and 0.800 g of potassium hydroxide? Any help showing the step-by-step working out to the above problem would be appreciated!

A student titrated an aliquot of standard sodium carbonate solution with hydrochloric acid in a burette. State whether the concentration determined for the hydrochloric acid would be likely to be higher, lower or unchanged compared with the actual value if the student had previously washed with water, but not dried, the following apparatus: a) the pipette used to deliver the aliquot of sodium carbonate solution b) the flask containing the aliquot c) the burette. Any help to the above questions would be immensely appreciated!