Chemistry Lab/Titration Race
Titration Race was a topic of Chemistry Lab in 2009.
Titration Race (2009)
Despite the name of this portion of the 2009 Chem Lab event, this event can barely be considered a race. In fact, a recent rule clarification states that time will not be considered a tie-breaker at the national competition. However if time is considered, here are a few helpful hints to increase both your speed and accuracy in performing a titration.
-Begin with a microtitration. For example titrate with 2 mL of the base to get a ballpark figure of the concentration. Use this to decide how much of the base you will use in the future trials. Remember that the more you use to titrate with, the more accurate your results will be.
-Three trials are recommended. If the ask to show work for only 2 data points, still do three and show the two that are closest together UNLESS your first two seem so close that a third is unnecessary.
-Remember what you're working with. Understand how sulfuric acid might change the calculations to find its concentration.
-Do not overtitrate. This is what will seperate the teams the most. A healthy red color is not what you're looking for at the end of a titration. You are looking for your final solution to be barely tinged pink. If given paper, place it below your flask so that you can easily see ANY tinge of pink. If you do overtitrate, go back to the acid and add a drop or two to get to a good final point.
Acids and Bases
For more info on Acids and Bases, see Chem Lab/Acids and Bases
Titrations
Titrations are where acids and bases are mixed together to figure out some unknown quantity from multiple known quantities.
Titration Curves
Titration curves are a plot with the pH on the y-axis and the amount of acid (or base) added on the x-axis. They normally start out fairly flat. There is then a steep slope, and then more flatness.
Equivalence Point
The equivalence point is the point in a titration at which the amount of acid is equal to the amount of base. The steep slope in a titration curve is around the equivalence point.
End Point
The end point is the point in a titration at which the pH is 7.
Indicators
Indicators are an essential part of a titration. Indicators, when put in a solution, will change color depending on what the pH is. Indicators can be used to figure out when the equivalence point is reached. Indicator paper is paper with indicators mixed in. Universal indicator paper has a certain mix of indicators, making it so that from it's color you can tell the pH of a solution.
| Low pH color | Transition pH Range | High pH color | |
|---|---|---|---|
| Gentian violet (Methyl violet 10B) | yellow | 0.0-2.0 | blue-violet |
| Malachite green (first transition) | yellow | 0.0-2.0 | green |
| Malachite green (second transition) | green | 11.6-14 | colorless |
| Thymol blue (first transition) | red | 1.2-2.8 | yellow |
| Thymol blue (second transition) | yellow | 8.0-9.6 | blue |
| Methyl yellow | red | 2.9-4.0 | yellow |
| Bromophenol blue | yellow | 3.0-4.6 | purple |
| Congo red | blue-violet | 3.0-5.0 | red |
| Methyl orange | red | 3.1-4.4 | yellow |
| Screened methyl orange (first transition) | red | 0.0-3.2 | grey |
| Screen methyl orange (second transition) | grey | 3.2-4.2 | green |
| Bromocresol green | yellow | 3.8-5.4 | blue |
| Methyl red | red | 4.4-6.2 | yellow |
| Azolitmin | red | 4.5-8.3 | blue |
| Bromocresol purple | yellow | 5.2-6.8 | purple |
| Bromothymol blue | yellow | 6.0-7.6 | blue |
| Phenol red | yellow | 6.4-8.0 | red |
| Neutral red | red | 6.8-8.0 | yellow |
| Naphtholphthalein | colorless to reddish | 7.3-8.7 | greenish to blue |
| Cresol Red | yellow | 7.2-8.8 | reddish purple |
| Cresolphthalein | colorless | 8.2-9.8 | purple |
| Phenolphthalein | colorless | 8.3-10.0 | fuchsia(pink) |
| Thymolphthalein | colorless | 9.3-10.5 | blue |
| Alizarine Yellow R | yellow | 10.2-12.0 | red |
ICE Tables
ICE Tables are useful in titrations. They go something like this:
| [math]\displaystyle{ HA \to }[/math] | [math]\displaystyle{ H^+ }[/math] + | [math]\displaystyle{ A^- }[/math] | |
|---|---|---|---|
| Initial | concentration | concentration | concentration |
| Change | -x | +x | +x |
| End | concentration | concentration | concentration |
Henderson-Hasselbalch equation
[math]\displaystyle{ pH = pK_a + log_{10}\frac{[A^-]}{[HA]} }[/math]
Strong Acid - Strong Base Titrations
Strong Acid-Strong Base titrations are relatively simple to work with. Normally they are used when you are trying the figure out the concentration of either a strong acid or base.
Take a sample of the strong acid (or base). Measure and record the volume of it. Slowly start adding a strong base (or acid) for which you know the concentration. Find the equivalence point (the point at which the indicator should change color). Measure and record the volume of the strong base (or acid) needed to reach the equivalence point.
At the equivalence point the amount of acid is equal to the amount of base. So, the following equation should be true.
[math]\displaystyle{ M_aV_a = M_bV_b }[/math]
You know three of the variables and can solve for the fourth.
Weak Acid - Strong Base Titrations
Weak Acid-Strong Base titrations can be used to find the Ka of a weak acid. This can be used to identify the acid.
Take a sample of the weak acid. Measure and record the volume of it. Add 10 ml of the strong base. Measure and record the pH with pH paper. Plot it on a titration curve graph. Add another 10 mL and do the same thing. Continue doing so. After you have done a fair amount, connect the dots with a line and estimate the pH of the equivalence point and the amount of strong base needed to get there. (the equivalence point is around the middle of the steep part of the titration curve)
[math]\displaystyle{ M_aV_a = M_bV_b }[/math]
Where [math]\displaystyle{ V_b }[/math] is the concentration of strong base at the equivalence point.
Calculate the concentration of the weak acid in the original solution. The concentration with the strong base added will be [math]\displaystyle{ \frac{M_aV_a}{V_a + V_b} }[/math]. Lets call this [math]\displaystyle{ M_e }[/math]. We know that [math]\displaystyle{ M_b }[/math] represents [math]\displaystyle{ [OH^-] }[/math] since the base is a strong base. Thus, [math]\displaystyle{ [OH^-] = \frac{M_bV_b}{V_a + V_b} }[/math]. Let's call this [math]\displaystyle{ M_{ih} }[/math].
| [math]\displaystyle{ HA }[/math] + | [math]\displaystyle{ OH^- \to }[/math] | [math]\displaystyle{ A^- + }[/math] | [math]\displaystyle{ H_2O }[/math] | |
|---|---|---|---|---|
| Initial | [math]\displaystyle{ M_e }[/math] M | [math]\displaystyle{ M_{ih} }[/math] M | 0 M | ~ |
| Change | -x | -x | +x | ~ |
| End | [math]\displaystyle{ M_e }[/math] - x M | [math]\displaystyle{ M_{ih} }[/math] - x M | x M | ~ |
pH + pOH = 14 pOH = -log[[math]\displaystyle{ OH^- }[/math]] From the pH, calculate [math]\displaystyle{ [OH^-] }[/math]. Let's call this [math]\displaystyle{ M_{eh} }[/math].
[math]\displaystyle{ M_{eh} = M_{ih} - x }[/math]
Solve for x.
[math]\displaystyle{ \frac{[H^+][A^-]}{[HA]} * \frac{1}{[H^+][OH^-]} = \frac{[A^-]}{[HA][OH^-]} }[/math]
[math]\displaystyle{ \frac{K_a}{K_w} = \frac{[A^-]}{[HA][OH^-]} }[/math]
[math]\displaystyle{ {K_a} = \frac{x*K_w}{(M_e - x)(M_{eh})} }[/math]
Solve for [math]\displaystyle{ K_a }[/math]
Weak Base - Strong Acid Titration
A similar method to the Weak Acid-Strong Base titration is used for Weak Base-Strong Acid Titrations.