Grade 11 Chemistry Problem Solving Quiz
Date: Thursday March 27, 2014
This competition is geared towards providing a fun and educational environment to
promote Chemistry and in which students can be recognized for their problem-solving
and practical skills while engaging in friendly competition. The event also offers
students an opportunity to test their knowledge and mastery of the Mole Concept and
other key topics in Chemistry.
Each school will be represented by a team of four (4) students who will work
collaboratively to solve written chemistry problems and carry out a practical
experiment within a 2 hour period.
Sample Questions
Question 1
Minerals containing zinc are important to the human body, since the Zn2+ ion is
important in many enzymes. Zinc sulphate can be used as a dietary supplement in
cases of zinc deficiency. The compound crystallises as a hydrated salt, ZnSO4.7H2O,
and is soluble in water.
The recommended daily allowance (RDA) of zinc is 15·2 mg (0·0152 g).
Using the information given above, answer the following questions
a)
(i)
Calculate the relative molecular mass of hydrated zinc sulphate
crystals, ZnSO4.7H2O.
[1]
(ii)
Calculate the mass of hydrated zinc sulphate crystals which would
need to be taken to obtain the RDA for zinc.
(Assume that 1 mole of hydrated zinc sulphate contains
1 mole of zinc.)
[2]
(iii)
Normally this mass is taken via a 5 cm3 dose of an aqueous
solution of hydrated zinc sulphate, ZnSO4.7H2O.
Calculate the concentration of this solution in g dm-3.
[1]
b) Describe how you would obtain dry crystals of hydrated zinc sulphate from its
aqueous solution.
[2]
c) Minerals containing iron are also important to the human body.
4·07 g of a chloride of iron was formed by reacting 1·40 g of iron with
an excess of chlorine.
(i) Determine the empirical formula of the chloride.
[2]
(ii) Describe a test, by stating reagent(s) and observation(s), to show that an
aqueous solution of the chloride of iron contains Cl– ions.
[2]
Reagent(s) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
. . . . . . . . . . . . . ……………………………………………………………
Observation(s) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
...............................................................
...
[Total for Question 1 = 12 marks]
Question 2
The formula for iron (II) sulphate crystals is FeSO4.xH2O where x is a whole number.
A student determined the value of x using 0.0200 mol/dm3 potassium manganate (VII).
This was solution G.
a) A sample of iron (II) sulphate crystals was added to a previously weighed
container, which was then reweighed.
Calculate the mass of iron (II) sulphate crystals used in the experiment.
Mass of container + crystals
Mass of empty container
Mass of iron (II) sulphate crystals
= 12.38 g
= 5.42 g
= ……………… g
[1]
b) The sample was dissolved in 100 cm3 of dilute sulphuric acid and the solution was
made up to 250 cm3 with distilled water. This was solution H.
A 25.0 cm3 sample of H was measured into a titration flask. Solution G was run
from a burette into the flask containing H until an end-point was reached.
Potassium manganate(VII) is purple.
i) What was the colour change at the end-point?
From ............................................... to ..................................................[1]
ii) Three titrations were done. The diagrams below show parts of the burette
before and after each titration.
Use these diagrams to complete the table of results.
Tick (✔) the best titration results.
Using these results, the average volume of G was .................. cm3.
[4]
iii) G is 0.0200 mol/dm3 potassium manganate (VII).
Calculate how many moles of KMnO4 were present in the titrated volume of G
calculated in (c)
. ..........................................moles
[1]
iv) Five moles of FeSO4 react with one mole of KMnO4.
Calculate how many moles of FeSO4were present in 25.0 cm3 of H.
..........................................moles
[1]
v) Calculate how many moles of FeSO4 were present in the 250.0 cm3 of H.
..........................................moles
[1]
vi) Using your answers to (f), calculate the mass of FeSO4 in the
original sample of FeSO4.xH2O.
[Mr: FeSO4, 152.]
Mass = ......................g
[1]
vii) Using your answer to (a) and (g) calculate the mass of water in the
sample of FeSO4.xH2O. ..........................................g
[1]
viii) Using our answer to (h) calculate the number of moles of water in
the sample of FeSO4.xH2O.
[Ar:H,1; O, 16]
Moles = ..........................................
[1]
ix) Using your answers to (f) and (i) calculate the value of x in FeSO4.xH2O.
................................................................................................ .
[1]
[Total for Question 2 = 12 marks]
Practical Experiment
VITAMIN C IN TRU JUICE ORANGE & TRU JUICE CHERRY
( An Oxidation/Reduction Titration - Adapted from Microanalysis of Vitamin C in
Citrus Juices, Juniata College "Science in Motion" Project)
Introduction
Vitamin C (ascorbic acid) is a water soluble vitamin that is required daily since it is
not synthesized or stored by the human body. Vitamin C is required for cellular
respiration, for enzyme function, as a component in collagen formation, and has been
reported to increase iron absorption within the body.
The RDA of vitamin C in adults is 60 mg per day. This vitamin is nontoxic even at
high dosage levels (less than several thousand milligrams per day).
In this task, vitamin C is being oxidized by iodine.
Starch and vitamin C are mixed in a styrofoam cup. Iodine is added to the solution
drop by drop. Iodine will react with both starch and vitamin C, but it will react with
vitamin C first.
However, there is no visible evidence of this reaction. Once all of the vitamin C has
reacted, the next drop of iodine will react with the starch to produce a blue-black
color. Thus, the starch is used as an indicator to visually determine the end of the
reaction of iodine with the vitamin C.
This analysis is will allow you to determine the precise endpoint of a reaction and
therefore the precise quantity of reactant (vitamin C) in the titration flask (styrofoam
cup).
Goals
In this task, your team will use titration to determine the amount of vitamin C in two
flavours of Tru Juice – Cherry and Orange. The type of titration you will be doing is
an example of a redox titration.
Safety
1. Be very careful to keep chemicals away from your eyes
2. The iodine solution can stain your clothes and skin. Disposable gloves are
provided to protect your hands. Hand soap is also provided. Handle iodine
solution carefully
3. DO NOT drink the fruit juices in the lab.
Procedure
1. Use a dropper to place exactly 10 drops of the vitamin C standard solution in a
styrofoam cup
2. Use a second dropper to add exactly two drops of starch indicator solution to the
vitamin C standard solution in the Styrofoam cup
3. Titrate the Vitamin C standard solution by adding iodine solution (use a third
dropper) drop by drop. Mix after each drop by very gently shaking the Styrofoam
cup. COUNT THE NUMBER OF DROPS. When the solution changes from
colourless to blue and remains blue for 15 seconds, stop adding drops.
Record the number of drops of iodine solution needed to titrate vitamin C
standard _____.
Vitamin C Standard = 0.1% (Mass percent) = 1mg/mL
4. Use a new dropper to place 10 drops of the Tru Juice Orange Juice in a Styrofoam
cup
5. Add exactly two drops of starch indicator solution (use the dropper already used
with starch)
6. Titrate with the iodine solution as before. Remember to count your drops. Stop
adding drops when the solution changes from colourless to blue and remain for 15
seconds
7. Record the name of the fruit juice and number of drops in your data table.
8. Repeat the entire titration using Tru Juice Cherry juice
9. When finished, dispose of your solutions as instructed and wash your hands.
Data Table
Fruit Juice # Drops to
Titrate
Mass
Percent
Vitamin C
mg Vitamin
C per 4oz
serving
%RDA per 4
oz serving
(from
calculations)
%RDA per 4
oz serving
(from label)
QUESTIONS
Record your answers in the data table.
1. Calculate the mass percent vitamin C in the fruit juice. Remember that fruit juices that
have MORE VITAMIN C will require MORE DROPS of iodine to turn the solution
blue. The two are DIRECTLY proportional. You will use a proportion to figure out
the concentration of vitamin C in the fruit juice.
#drops for fruit juice titration
#drops for Vitamin C Standard titration =
mass % Vitamin C in fruit juice
mass % Vitamin C in standard
2. Which juice had more vitamin C?
3. A 0.1% solution means there is 1mg in every 1mL of water. Using this information,
calculate how many milligrams of vitamin C there are in a 1 ounce serving of each
juice. (1 ounce = 30 mL).
4. What is the recommended daily allowance (RDA) for vitamin C?
5. Calculate the percentage of the RDA (%RDA) in a 4oz serving of each fruit juice.
Record the values in your data table.
6. Record the %RDA vitamin C in each juice listed on the juice bottle (make sure to
calculate it in terms of a 4oz serving).
7. How does the information from your experiment compare to the information on the
label?
Label information (Vitamin C):
True Juice Orange: 180% RDA
True Juice Cherry: 770% RDA