Titration help handout:
The acid/base titration with NaOH and acetic acid is the oldest experiment in the book. I did it when I was in college (a long time ago) and both schools for which I work do the lab. I know they repeat it in higher level classes with a few more steps to make it more challenging. So it is to your advantage to try to master it at this level so you can handle more detail when they throw it at you.
Part I: Standardization of a base (NaOH)
Why would you standardize a solution? Generally it is because you need to precisely know the concentration of it.
What variables do you know in the first part? You know how many moles of acid that you started with. (KHP). You can calculate moles of acid in your starting material. You also know that the balanced chemical equation for the neutralization of KHP and NaOH reacts in a 1:1 mole ratio between acid and base.
KHP(aq) + NaOH(aq) → KNaP(aq) + H2O(l)
With your buret, you measure the volume of NaOH required to neutralize your KHP acid. You know the experiment is complete when you see your indicator turn a faint pink color. This indicates the pH of the titration is greater than 7 (slightly past the end point the pH can jump up sharply). The true equivalence between NaOH and KHP comes a few drops before you see pink in the solution but we approximate equivalence at the point the solution turns pink.
How do you calculate the concentration of the standard solution? You take moles of acid and set it equal to moles of NaOH. (You know this is true based on the balanced equation).
You can read the volume of NaOH required to complete the titration on your buret. This is the total mL (or Liters) you used to complete the titration.
Simply divide moles of base by liters of solution to get the concentration of the standard solution.
In part II of this experiment you are working backwords. What are your known and unknown variables? Let’s think about this carefully:
Generally you are told to pipette a certain amount of acetic acid in a flask. Let’s say you have a 5.00 mL calibrated pipette. You have 5.00 mL or 0.00500 L of acetic acid in your flask. (Notice I used three sig figs by adding two zeroes to the end of my liter value)
In this titration you now have a known concentration of NaOH from part A. (No, it is not 1:1 as I often see in my lab reports for this lab. It is the value you determined in part A). At this end of the titration, your acetic acid turns pink when moles of acetic acid is essentially equal to moles of base. How do you use this data to derive the concentration of your acid?
Let’s say you have a 0.2540 M concentration of NaOH from part A and you used 26.80 mL of your NaOH to titrate your acetic acid. You can determine moles of NaOH used for the titration by multiplying your concentration of NaOH by your volume of NaOH.
0.2540 moles/Liter NaOH X 0.02680 L NaOH= moles of NaOH (moles/Liter X Liters is always moles)
We know this is equal to moles of acetic acid by the balanced chemical equation.
What volume is this divided by to get the concentration? The original volume of acid you pipette into your flask- 5.00 mL or 0.00500 L.
The acid/base titration with NaOH and acetic acid is the oldest experiment in the book. I did it when I was in college (a long time ago) and both schools for which I work do the lab. I know they repeat it in higher level classes with a few more steps to make it more challenging. So it is to your advantage to try to master it at this level so you can handle more detail when they throw it at you.
Part I: Standardization of a base (NaOH)
Why would you standardize a solution? Generally it is because you need to precisely know the concentration of it.
What variables do you know in the first part? You know how many moles of acid that you started with. (KHP). You can calculate moles of acid in your starting material. You also know that the balanced chemical equation for the neutralization of KHP and NaOH reacts in a 1:1 mole ratio between acid and base.
KHP(aq) + NaOH(aq) → KNaP(aq) + H2O(l)
With your buret, you measure the volume of NaOH required to neutralize your KHP acid. You know the experiment is complete when you see your indicator turn a faint pink color. This indicates the pH of the titration is greater than 7 (slightly past the end point the pH can jump up sharply). The true equivalence between NaOH and KHP comes a few drops before you see pink in the solution but we approximate equivalence at the point the solution turns pink.
How do you calculate the concentration of the standard solution? You take moles of acid and set it equal to moles of NaOH. (You know this is true based on the balanced equation).
You can read the volume of NaOH required to complete the titration on your buret. This is the total mL (or Liters) you used to complete the titration.
Simply divide moles of base by liters of solution to get the concentration of the standard solution.
In part II of this experiment you are working backwords. What are your known and unknown variables? Let’s think about this carefully:
Generally you are told to pipette a certain amount of acetic acid in a flask. Let’s say you have a 5.00 mL calibrated pipette. You have 5.00 mL or 0.00500 L of acetic acid in your flask. (Notice I used three sig figs by adding two zeroes to the end of my liter value)
In this titration you now have a known concentration of NaOH from part A. (No, it is not 1:1 as I often see in my lab reports for this lab. It is the value you determined in part A). At this end of the titration, your acetic acid turns pink when moles of acetic acid is essentially equal to moles of base. How do you use this data to derive the concentration of your acid?
Let’s say you have a 0.2540 M concentration of NaOH from part A and you used 26.80 mL of your NaOH to titrate your acetic acid. You can determine moles of NaOH used for the titration by multiplying your concentration of NaOH by your volume of NaOH.
0.2540 moles/Liter NaOH X 0.02680 L NaOH= moles of NaOH (moles/Liter X Liters is always moles)
We know this is equal to moles of acetic acid by the balanced chemical equation.
What volume is this divided by to get the concentration? The original volume of acid you pipette into your flask- 5.00 mL or 0.00500 L.
I find this post very interesting, since it illustrates the very different ways we need to teach. I am not able to define it exactly though.
ReplyDeletecan you elaborate?
ReplyDeleteIt is very complicated. Note that I said "NEED to teach". I teach students who study for a BSc. In the past few, specially in the state I live in, most students try and get into a medical college. Those who study for a BSc are often those who did'nt make it into medical college. I cannot expect them to do things for themselves. I cannot expect them to go and find out about some new term or concept.I have to TELL them everything. I try to change this, but not very successfully.
ReplyDeleteOf course, the most important question is, "Is it pink yet?" I think sometimes it would be beneficial to students if I brought my friends 5-yr-old daughter into lab to field these questions - because if there is one thing she knows, it's how to correctly identify "pink"! :)
ReplyDeleteI just love all this participation in my blog!! Keep it coming ! Encourage others to stop by!
ReplyDeleteThanks for coming!
@cbeck, with phenolphthalein, while titrating with lower conc., the question "is it colourless yet?" is quite valid. I have a strong dislike of phenolphthalein since I use lower conc. and it really does not give a sharp end point. However, I have to stick to the prescribed indicator otherwise my students lose marks in their final exam conducted by the University.
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