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Chemistry Stage 1 Work Samples Concepts are introduced qualitatively, looking at the relationships of chemical interactions rather than going into mathematical quantitative analysis. Correct terminology, appropriate research skills and safe practical experimental methods are expected to be used Stage One program
Weeks Content Laboratory work Assessment
1–3
Macroscopic properties of matter � explain the differences between elements,
compounds and mixtures � use examples to describe homogeneous and
heterogeneous mixtures � explain the difference between a physical and a
chemical property � explain and describe the differences between
physical and chemical changes � describe the following methods for separating
mixtures: decantation filtration crystallisation evaporation distillation chromatography.
Mixtures (STAWA Experiment 1) Physical and chemical change Making pop corn Investigation How do we get clean water? (Introduction to how to conduct an investigation) Lolly chromatography
Task 1 Write up of the planning and evaluation of Chromatography with M&M’s Task 2 Practical Assessment Investigation to find out if seaweed contains chlorophyll
4–6
States of matter and Kinetic Theory � use the Kinetic Theory of matter to explain:
properties of gases, liquids and solids phase changes heating and cooling curves temperature and kinetic energy.
Gas laws � use qualitative data to explain the behaviour of
gases in response to changes in temperature, pressure and volume.
Solutions and solubility � describe the different solute/solvent combinations
that form different types of solutions � explain and apply the concept of solubility to
describe: unsaturated and saturated solutions concentration scales (g L-‐1)
� explain factors that affect solubility and crystallisation.
� analyse food labelling to determine concentration scale used to describe the nutrition information provided shelf life (length of use by dates).
Investigation What makes corn pop? (http://www.arborsci.com/CoolStuff/cool8.ht Practical Bubbles and Cartesian divers (Group conducting, individual processing and evaluation) http://www.arborsci.com/CoolStuff/cool8.htm Research Solvent extraction of plant dyes Practical Making a borax snowflake (http://chemistry.about.com/cs/howtos/ht/boraxsnowflake.htm)
Task 3 Write up of planning exercise for ‘What makes corn pop’? Task 4 Write up of processing data, evaluation of Bubbles and Cartesian divers Task 5 In-‐class assessment of Research topic Solvent extraction of plant dyes Task 6 Quiz on controlling variables in Making a borax snowflake practical Task 7 Topic test 1 Macroscopic properties of matter
Stage One sample assessment
SECTION 2 20 Marks Answer ALL questions.
1. A student is given a sample of a white solid that looks like sand. Observations with the aid of a hand lens show that the sand consists of an irregular mixture of small grains of two different crystals. (5 marks)
(a) Is the mixture heterogeneous or homogeneous? (b) What is the evidence for your conclusion in (a)? (c) If the mixture was sand and sugar, name a separation technique that could be used to recover the (i) sand (ii) sugar
(d) Are these separation techniques based on physical or chemical properties?
2. The graph below shows the change in the solubilities of potassium nitrate and lead nitrate as the
temperature of the water changes. Use the information given on the graph and what you have learnt about solutions to answer the following questions. (5 marks)
(a) At what temperature are both solutes equally soluble? (1 mark) (b) How much potassium nitrate would you expect to dissolve in 100 grams of water at
40°C? (2 marks) (c) Describe the solutions formed when the following masses of solute are added to water
and stirred: (i) 30 grams of potassium nitrate is added to 100 mL of water at 40 °C (1 mark) (ii) 70 grams of lead nitrate is added to 100 mL of water at 25 °C (1 mark)
Temperature (°C)
potassium nitrate
lead nitrate
3. A student was investigating the speed at which the temperature of cooking oil changes as it is being
heated in an electric deep fryer. The data that the student recorded is given in the table below. (10 marks)
Time (min) Temp (°C)
0 25 2 60 6 120 10 175 12 180 14 180
Plot the data on the graph paper below and then answer the questions below. (5 marks)
(a) From the graph, estimate the time that the oil began to boil. (1 mark) (b) Is the change that is occurring at t = 12 minutes a chemical or physical change? Give an explanation to support for your choice of answer. (2 marks) (c) Use the kinetic theory of matter to explain what is happening to the energy and the speed
of the oil molecules between time t = 8 minutes to t = 9 minutes.
energy (1 mark) speed (1 mark)
Chemistry Stage 2 Work Samples Key Concepts covered in Stage 2 Chemistry: • Macroscopic Properties of Matter – Kinetic Theory • Basic Calculations in Chemistry • Solutions • Atomic Structure and Bonding • Chemical Reactions and Rates of Reactions • Acids and Bases • Redox Reactions • Organic Chemistry
A Sample Program for a unit in Stage 2 Chemistry:
Week Content Text
Reference Exploring Chemistry
Assessment
8-‐10
Unit 3: MACROSCOPIC PROPERTIES: Solutions
identify, explain and give examples of saturated, unsaturated and supersaturated solutions
apply solubility rules to predict if a precipitate will form when two dilute ionic solutions are mixed
use the colour of ions to identify reactants and the products in chemical processes
explain colligative properties including the effect of concentration on vapour pressure, melting point and boiling point of a solution
describe characteristics of strong electrolytes & give examples including ionic compounds & strong acids
describe the characteristics of weak electrolytes and give examples including weak acids and bases
describe the characteristics of non-‐electrolytes and give examples including water and hydrocarbons
explain the differences between concentrated and dilute solutions of strong and weak electrolytes.
Pg. 61-‐62 Pg. 63-‐65
E4 pg 23-‐24 I2 pg 25 Set 2 pg. 35 Set 15 pg 110
Task 1: Investigation 2 pg. 25 E.C. Task 6: Assignment Unit 3
10
Applied Chemistry Solutions � describe and give examples of chemicals and their uses in and
around the home including vinegar, bleach, ammonia solution and caustic soda
� explain concentration units used in household mixtures (g 100g-‐1,
mL L-‐1, g L-‐1, percentage composition by mass)
Pg. 65-‐66
E5 pg 26-‐27 If time, I3
Task 11: Test Unit 3
SAMPLE ASSESSMENT OUTLINE FOR CHEMISTRY 2A
1. Investigating in chemistry 2. Structure, properties and uses of materials 3. Interaction and change 4. Problem solving and quantities in chemistry 5. Chemistry in action
Outcomes coverage Assessment
type & Weighting
Tasks Content 1 2 3 4 5
Weight %
Task 1: Investigation Investigation 2 from Exploring Chemistry Stage 2 – identifying unknown cations and anions
3
Task 2: Investigation Student investigate if an unknown substance is NaHCO3 and write a police report
4
Practical Assessments
and Investigations
(15-‐25%)
15%
Task 3: Practical Examination Examination to Validate write up of experiments and investigations conducted by students during 2ACHE
8
Task 4: Assignment Unit 1 In class assignment 1 3
Task 5: Assignment units 2 In class assignment 2 3
Task 6: Assignment Unit 3 In class assignment 3 3
Task 7: Assignment Unit 4 In class assignment 4 3
Assignments and
class work
(15-‐25%)
22%
Task 8: Mini Tests Accumulation of regular mini tests conducted throughout semester
10
Task 9: Test 1 Test on Unit 1 Macroscopic Properties
5
Task 10: Test 2 Test on Unit 2 Atomic Structure and Bonding
6
Task 11: Test 3 Test on Unit 3 Solutions 6
Task 12: Test 4 Test on Unit 4 Chemical Reactions and Stoichiometry.
6
Tests and exams
(50-‐70%)
63%
Task 13: 2ACHE Examination 2A Chemistry Examination on Semester’s work
40
TOTAL 100
In General:
Chemistry, the study of matter and its interactions, is an indispensable human activity that has contributed essential knowledge and understanding of the world around us. Chemical knowledge has enabled us to understand matter and devise processes for activities such as: cooking and preserving food; purifying air and water; recycling plastics; anaesthetising patients; creating and building computers; and communicating with others around the world about chemistry.
Each unit in Chemistry has a workbook which you work through with information, examples and problems to complete.
Example of a page from:
Volume-‐Volume Relationships in Chemical Reactions
Gay-‐Lussac’s Law
• Gay-‐Lussac’s Law of combining volumes states:
The volumes (measured at the same temperature and pressure) of different gases in a reaction are in the same proportion as their coefficients in the equation.
This law is a direct outcome of Avogadro’s hypothesis: equal volumes of gases contain equal number of molecules, thus equal number of moles.
• Gay-‐Lussac’s law only applies to gases in a reaction.
Example: In the production of ammonia gas, hydrogen gas and nitrogen gas are combined. If I have 85.12 L hydrogen and excess nitrogen, how many litres of ammonia gas can you produce. 1. Write balanced equation. N2 + 3H2 2NH3 2. Write in mole relationship 1 mol 3 mol 2 mol 3. Write in volume relationship (1 x 22.71 L) (3 x 22.71 L) (2 x 22.71 L) 22.71 L 68.13 L 45.42 L 4. Write in given data 85.12 L x
5. Calculate volume produced = 56.74666
x = 56.8 L ammonia gas produced.
Example for you to try: In the production of water, oxygen gas is combined with hydrogen gas
a. Write the balanced formula b. If I have 58.24 L of oxygen gas, how many litres of hydrogen gas do I need?
Experiments and Investigations: You also conduct investigation. An example is shown below. Investigation 2 -‐ Use this as a guide to write up your investigation (35 marks) Title: Identification please Purpose: Write a suitable purpose for this investigation. (1 mark) Equipment: Set 1: Nitrate solutions of Fe2+, Pb2+, Ag+, Cu2+, Mg2+ randomly labelled A, B, C, D and E.
Test solutions for above: OH-‐, Cl-‐, I-‐
Set 2: Sodium solutions of CO3
2-‐ , OH-‐, Cl-‐ , SO42-‐ ,
Test solutions for above: Pb2+, Ag+, Fe3+, Cu2+, H2SO4 Test tubes, test tube rack, droppers Procedure: Write a brief procedure explaining exactly how you conducted your investigation. Your procedure must be in point form and written such that someone who had not seen your investigation could repeat it exactly the way you preformed it. (3 marks) Results: Copy the tables below into your laboratory workbook (make the rows wide enough to record observations). Carry out the investigation and record your observation (ppt formed and colour). Where no observable reaction, write n.v.r. (5 marks) Set 1
A B C D E OH-‐ Cl-‐ I-‐
Set 2 F G H I
Pb2+ Ag+ Fe3+ Cu2+ H2SO4
Processing Data a. Correct identification of solutions A to I. (9 marks) b. Write equations for any reactions that occurred with each unknown solutions A to I. Do this by
putting the heading e.g. A, then write the equations for any reactions that occurred with solution A. Do the same for B to I. (10 marks)
Evaluating the Investigation:
1. Explain how effective your procedure was? (2 marks) 2. List improvements you would make to improve the investigation. (2 marks) 3. Discuss your confidence in your findings and reasons why you are or are not
confident with your findings. (2 marks)
Tests and Examinations
Tests and examinations form a reasonable proportion of physics assessments. Below are some examples of the types of questions you will need to answer as well as their answers. Multiple Choice Type Question:
Easy recall
Which of the following 0.100 mol L-‐1 water solutions would have the LOWEST electrical conductivity? 1. silver nitrate 2. ammonium chloride 3. sodium chloride 4. ethanoic acid Harder Consider the reaction between phosphoric acid and calcium hydroxide:
2H3PO4 + 3Ca(OH)2 Ca3(PO4)2 + 6H2O
How many moles of calcium hydroxide would react with 4 moles of phosphoric acid? a. 1.5 b. 2 c. 3 d. 6 Written Type Question:
Easy recall
Bottled gas, such gas that used for BBQs, is compressed propane gas (C3H8). when propane burns in air (oxygen) it not only produces heat but also carbon dioxide gas and water vapour. Is the burning of propane a physical or chemical change? Explain. (2 marks)
This is a chemical change (1 mark)
as the propane is turned into new substances and the reaction is not reversible. (1 mark)
Harder (using your understanding of chemistry to explain a situation)
While ethanoic acid and hydrochloric acid are both acids, hydrochloric acid will conduct electricity much better than ethanoic acid. Explain why. Appropriate equations will help your explanation. (5 marks)
HCl is a strong electrolyte so almost totally ionises into ions 1 mark
HCl(aq) � H+(aq) + Cl(aq) 1 mark
CH3COOH is a weak electrolyte and only partly ionises into ions 1 mark
CH3COOH(aq) � CH3COO-‐(aq) + H
+(aq) 1 mark
More ions in solution, the more the solution conducts electricity so HCl conducts more than CH3COOH. 1 mark
Answer: 2H3PO4 + 3Ca(OH)2 Ca3(PO4)2 + 6H2O 2 mol 3 mol 1 mol 6 mol 4 mol (4 x 3 / 2) = 6 mol
Mathematical Type Questions:
Simple
Calculate the mass of 0.810 moles of aluminium hydroxide. (2 marks) M(Al(OH)3) m = n x M = 26.98 + 48.00 + 3.024 = 0.81 x 78.004 = 78.004 g mol-‐1 m = 63.2 g Harder A solution containing 2.00 g of sodium carbonate was added to a solution containing 2.00 g of magnesium chloride. A white precipitate forms. The molecular equation is shown below.
Na2CO3 (aq) + MgCl2 (aq) � MgCO3 (s) + 2NaCl (aq)
What is the limiting reagent? Show working to justify your answer. (3 marks)
n(Na2CO3) = = 0.0189 mol n(MgCl2) = = 0.0210 mol
Na2CO3 (aq) + MgCl2 (aq) � MgCO3 (s) + 2NaCl (aq)
1 mol 1 mol 1 mol 2 mol have 0.0189 mol 0.0210 mol need 0.0189 mol 0.0189 mol 0.0189 mol (2 x 0.0189) excess NIL 2.1 x 10-‐3 mol as you can see from above, Na2CO3 is the limiting reagent Drawings:
Draw electron dot diagrams for the following. (4 marks)
Chemistry Stage 3 Work Samples
Key Concepts covered in Stage 3 Chemistry: These are the same as stage 2 Chemistry as stage 3 Chemistry builds on the concepts covered in stage 2.
A Sample Program for a unit in Stage 3 Chemistry (Term 1):
Week Content Exploring Chemistry
Assessment
1-‐2
Unit 1: Macroscopic properties of matter � interpret observations, such as the colour changes, of physical and chemical
systems at equilibrium � use observable properties, such as the colour of ions, to help predict and
explain the formation of products in chemical processes (see data sheet) � use the Kinetic Theory to explain the concept of absolute zero.
Invest: 1 pg. 41 Expt: 6 pg. 43
3
Solutions � apply the solubility rules to predict if a precipitate will form when two dilute
ionic solutions are mixed (see data sheet) � perform concentration calculations (mol L-‐1, g L-‐1, ppm, percentage
composition by mass) � calculate concentration of ions in solution for strong electrolytes � perform the calculation of concentration and volume involved in the dilution of
solutions and the addition of solutions.
Task 5: Assignment Unit 1
4
Unit 2: Atomic structure and bonding � explain structure of atom in terms of protons, neutrons & electrons � write the electron configuration using the shell model for the first twenty
elements e.g. Na. 2, 8, 1 � explain trends in first ionisation energy, atomic radius and electronegativity
across periods and down groups (for main group elements) in Periodic Table � explain the trend in successive ionisation energies � describe and explain the relationship between the number of valence electrons
and an element’s � bonding capacity and position on Periodic Table � physical and chemical properties.
Set 8: pg. 76 Set 9: pg. 78 Set 11: Pg. 82
Task 10: Test unit 1
5
� describe and apply the relationships between the physical properties and the structure of ionic, metallic, covalent network and covalent molecular substances
� use the Valence Shell Electron Pair Repulsion (VSEPR) theory and Lewis structure diagrams to explain and predict and draw the shape of molecules and polyatomic ions (octet only)
Expt 7: pg. 69 (modified) Set 10: pg. 80
6-‐7
� explain polar and non-‐polar covalent bonds in terms of the electronegativity of the atoms involved in the bond formation
� use relationship between molecule shape and bond polarity to predict and explain the polarity of a molecule
� explain differences between intermolecular and intramolecular forces � describe &explain the origin and relative strength of the following
intermolecular interactions for molecules of a similar size: 1. dispersion forces 2. dipole-‐dipole attractions 3. hydrogen bonds 4. ion-‐dipole interactions such as solvation of ions in aqueous solution
5. explain the relationships between physical properties such as melting and boiling point, and the types of intermolecular forces present in substances of similar size
Invest: 5 pg. 74 (see right) Expt: 8 pg. 72
Task 6: Assignment Unit 2
A Sample Assessment Outline for Stage 3 Chemistry: Weight % Assessment
type & Weighting
Tasks Content 3A 3B Total
Task 1: Independent Investigation
For example: Investigation 5 EC done at home but written up in class.
3 2 5
Task 2: Practical Examination 3A Examination to test understanding of experiments and investigations
5 5
Practical Assessm
ents and
Investigations
(15-‐25
%)
15%
Task 3: Practical Examination 3B Examination to test understanding of experiments and investigations
5 5
Task 4: Mini Tests and Revision Assignments
Regular mini tests and validation tests on revision assignments on any work covered.
3 4 7
Task 5: Assignment on Unit 1 Either validation questions on unit 1 assignment or in class assignment
1 1
Task 6: Assignment on Unit 2 Either validation questions on unit 2 assignment or in class assignment
1 1
Task 7: Assignment on Units 3 and 4
Either validation questions on unit 2 and 3 assignment or in class assignment
2 2
Task 8: Assignment on Units 5 and 6
Either validation questions on unit 5 and 6 assignment or in class assignment 2 2
Assignm
ents and
class work
(15-‐25
%)
15%
Task 9: Assignment on Units 7 and 8
Either validation questions on unit 7 and 8 assignment or in class assignment 2 2
Task 10: Test Unit 1 Test on Solutions 2 2
Task 11: Test Unit 2 Test on Atomic Structure and Bonding 3 3
Task 12: Test Units 3 and 4 Test on Chemical Reactions and Equilibrium.
5 5
Task 13: Test Units 5 and 6 Test on Chemical Acids and Bases and Titrations
5 5
Task 14: Test Units 7 and 8 Test on Oxidation & Reduction and Organic Chemistry
5 5
Task 15: 3ACHE Examination Semester One Chemistry Examination on 3A Chemistry
15 15
Tests and exam
s (50-‐70
%)
70%
70 %
Task 16: Semester 2 Examination
End of year Semester 2 Examination on whole years work
10 25 35
50 50 100
Example of a page from the Motion and Forces Workbook
APPLYING THE COLLISION THEORY
Nature of Reactants • slower reactions will have higher activation energies than faster reactions. • different activation energy for different reactions relating to the ease with which bonds are
broken or formed. • high activation energy -‐ reactions in which strong bonds are broken. Concentration • Increasing the concentration of a reactant, increase the collision rate. Greater the number of collision, faster rate of reaction. Sub-‐division • If the surface area of a reactant is increased, more molecules are exposed to collision. Temperature • increase in temperature increases the
average kinetic energy of the reacting molecules.
• increased velocities of the molecules -‐ greater rate of collision and rate of reaction.
• greater proportion of reactant molecules have sufficient kinetic energy to supply the activation energy needed for reaction.
Catalysts • provide an alternative pathway that has a lower
activation energy. • less collision energy is required for the
reaction -‐ greater proportion of collisions will be successful and the reaction rate will be greater.
• greater proportion of reactant molecules have sufficient energy to overcome the activation energy barrier.
• reactions can be carried out at much lower temperatures.
• form intermediate compounds in which the bond-‐breaking and forming process requires less energy than if the reactants alone were involved.
Experiments and Investigations: You also conduct investigation. An example is shown below.
Stage 3 Chemistry: Experiment 25 Reactions of Alcohols.
A: Reaction of Alcohols with Oxidising Agents:
Dichromate Ion:
A. Ethanol: Type of alcohol? _________________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
B. 1-‐butanol: Type of alcohol? _______________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
C. 2-‐butanol: Type of alcohol? _______________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
D. 2-‐methyl-‐2-‐propanol: Type of alcohol?________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
Permanganate ion:
A. Ethanol: Type of alcohol? _________________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
B. 1-‐butanol: Type of alcohol? _______________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
C. 2-‐butanol: Type of alcohol? _______________________________________________________
Observations: __________________________________________________________________
______________________________________________________________________________
Tests and Examinations Tests and examinations form a reasonable proportion of physics assessments. Below are some examples of the types of questions you will need to answer as well as their answers. Multiple Choice Type Questions Easy recall
The atomic radius of an atom: (a) increases down the group and increases across the row (b) increases down the group and decreases across the row (c) decreases down the group and increases across the row (d) decreases down the group and decreases across the row
Harder
0.50 mol of a certain metal reacts with excess hydrochloric acid to produce 0.50 mol of hydrogen gas. What is the charge on the metal ion produced in this reaction?
(a) 1+ (b) 2+ (c) 3+ (d) not enough information to determine charge. Written Type Question: 1. Give fully balanced equations for the reactions which occur (if at all) in the following experiments.
Use ionic equations where appropriate. Describe observations. (6 marks)
(a) Chlorine gas is bubbled through potassium bromide solution in a test tube. Cl2(g) + 2Br
-‐(aq) → 2Cl-‐(aq) + Br2 (aq)
Pale green gas dissolves and colour disappears Red substance forms in colourless solution 2. Write the two half equations then balance the following redox reaction. (3 marks) Cr2O7
2-‐ (aq) + I
-‐ (aq) � Cr3+ (aq) + I2 (aq)
Cr2O72-‐ + 14H+ + 6e � 2Cr3+ + 7H2O
6 2I-‐ � 3I2 + 6 2e x3 Cr2O7
2-‐(aq) + 6I
-‐ (aq) + 14H+ (aq) � 2Cr3+ (aq) + 3I2 (aq) + 7H2O (l)
3. Water is sometimes referred to as a ‘universal solvent’ because although it is polar itself, it is often able to dissolve both polar and non-‐polar solutes. Compare the way in which water dissolves an ionic solid such as sodium nitrate with the way it dissolves a covalent molecule such as sugar and state why sodium chloride can conduct in a solution but sugar can not.
Sodium nitrate: ionic bonds between ions in lattice are broken and water molecules surround sodium and nitrate ions forming ion-‐dipole interactions. [1 mark]
Formation of ions means solution can conduct electricity. [1 mark]
Sugar:
Forces between sugar molecules in crystal break and hydrogen bonds are formed between the H and O atoms in both the sugar and water. [1 mark]
No ions formed so solution doesn’t conduct electricity. [1 mark]
Answer: X+2 + 2HCl � XCl2 + H2 H2 same moles as metal 0.5 mol 0.50 mol so must be +2
Mathematical Type Questions:
Easier A solution of sodium carbonate was prepared by dissolving 3.34 g of Na2CO3 solid in a small amount of water then making the solution up to the mark in a 250.00 mL volumetric flask. The final mass of the solution is 242.4 g. Determine the concentration of the solution in the following units: a. ppm
ppm =
ppm = 1.38 x 104 ppm [2 mark] b. g L-‐1
c = [1 mark]
c = 13.4 g L-‐1 [1 mark] Harder 1. Iron is often found in the earth's crust as a hydrated iron oxide (Fe2O3.xH2O). A 0.668 gram sample of this ore was dissolved in excess sulfuric acid and all the iron was converted into an iron(II)sulfate solution. This solution was titrated against a 0.050 mol L-‐1 acidified potassium permanganate solution and exactly 25.00 mL of the purple solution was needed to complete the titration. Find the value of x in the hydrated formula above. (10 marks)
5Fe2+ + MnO4-‐ + 8H+ → 5Fe3+ + Mn2+ + 4H2O (2 mark)
n(MnO4-‐) = n(KMnO4) = c x V
= 0.05 x 0.025 = 0.00125 mol MnO4
-‐ (1 mark)
from the equation, n(Fe2+) = 5 x n(MnO4-‐) = 5 x 0.00125
= 0.00625 mol Fe2+ (1 mark)
n(Fe) = n(Fe2+) = 0.00625 mol Fe
n(Fe2O3) = ½ n(Fe) = 0.003125 mol Fe2O3 (1 mark)
m(Fe2O3) = n x M = 0.003125 x 159.7 = 0.4990625 g Fe2O3 (1 mark)
m(H2O) = m(sample) -‐ m(Fe2O3) = 0.6680 -‐ 0.4990625 = 0.1689 g H2O (1 mark)
n(H2O) = m/M = 0.1689 / 18.016 = 0.009377 mol H2O (1 mark)
(1 mark)
Hence, x = 3 (1 mark)
Drawings and/or Graphs:
Question 42 [15 marks] Answer A student investigates the effect of the concentration of hydrochloric acid on the rate of oxidation of zinc in the laboratory. She adds 40.0 mL of 1.00 mol L-‐1 hydrochloric acid to 20.0 g of zinc in a conical flask and measures the rate at which hydrogen is given off.
Time (min) 0 0.5 1.0 1.5 2.0 3.0 5.0 7.0 8.0 10.0
Loss in mass (g)
0 0.19 0.35 0.47 0.63 0.72 0.82 0.86 0.88 0.88
The flask and contents were immediately weighed and a stopwatch started. The mass of the flask and contents were noted as the reaction proceeded. The table indicates the loss in mass at various times. a. List two variables you would expect to control in this experiment. [2 marks]
Temperature mass of zinc
volume of acid that is added need TWO
b. List one variable you have to measure and one other variable that you could measure to determine the rate of reaction. [2 marks]
Variable you have to measure time
Variable that you could measure Loss of mass, volume of hydrogen evolved c. Write a suitable hypothesis for this experiment. [2 marks]
Any suitable hypothesis such as That the relationship between time and loss of mass will be linear until the zinc is consumed
d. Explain why the loss is constant from 8 minutes. [1mark]
One of the substances (most likely zinc) was a limiting reagent
It was and was totally consumed so the reaction stopped
NOTE: Must be based on chemistry understanding therefore reaction complete is not an acceptable answer.
Plot a graph of 'loss in mass' against time. [5 marks]
Marks: � axis correct 1 mark � labels on axis correct 1 mark � points plotted correctly 1 mark � line of best fit 1 mark � title 1 mark e. List two potential sources of uncertainty in experimental measurements in this investigation and how you could
minimise the error. [3 marks]
uncertainty in experimental measurement How could you minimise error � volume measurement � mass measurement � measuring time in minutes
� use pipette or burette � use electric balance that measures to 0.001 g � measure in seconds
