Download The Chemistry Handbook

V1.6
THE CHEMISTRY
AS/A-LEVEL
HAND BOOK
NAME:________________________ TEACHER:__________________
WHERE WILL IT TAKE YOU?
______________________________
CONTENTS
PAGE
Introduction
The Chemistry A-Level Commitment
3
4
Assessment Information & Record of Achievement
Internal Assessment
Record of achievement - End of unit tests & Internal mock exams
Record of Achievement - Homework
External Assessment information
6
7
9
11
General Information
Strategies for reaching and exceeding your target grade
Homework
15
16
Contact with your teacher
16
Tools
Reading Lists
Phone Apps
Internet
Definitions
Table of common ions and molecules
Properties of common acids/alkalis/bases
Equations
Balanced equations
Practical Techniques
Practical Observations
Practical Accuracy (experimental error)
Yields
Hazards /precautions
18
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19
20
21
22
22
24
25
26
28
29
AQA Provided information
Tabulating data
Significant figures
Uncertainties
Graphing
Glossary of terms
Colours of Metal Ions in Aqueous Solution
GCE Chemistry Data Sheet
The Periodic Table of the Elements
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33
34
39
48
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Introduction
Welcome to the Helsby High School Chemistry AS/A-level course. This handbook is aimed at providing
you with a reference source for some common queries on some common subjects. Keep it handy for the
duration of your course and refer to it often.
The specification that you will be following is the AQA GCE (AS & A Level) Chemistry specification
2015. The specification tells you what the exam boards expects you to know and do. The specification
is available on the Chemistry part of the shared file space.
You will study the A-level course over two years. At the end of the first year you will take external
exams for which you can achieve an AS-level, however, for the A-level qualification it will only be the
exams at the end of the two years study which will count towards the A-level. You will also be
completing a practical element of the course which is explained in the practical based Pupil Assessment
Booklet.
Come to the edge.
We might fall.
Come to the edge.
It’s too high!
Come to the edge!
And they came,
And he pushed,
And they flew….
Christopher Logue, ‘Come to the Edge’, 1969
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The Chemistry A-level commitment
Chemistry A-level is a highly sought after commodity by individuals, universities and employers. In
addition to teaching you a particular branch of science it helps you develop your transferable skills
such as problem solving, analysis, evaluation and communication skills. With Chemistry you can enter
medicine, law, banking, office based work, engineering, architecture and of course work in the Chemical
industry in a variety of careers.
But as with anything worth having in life you need to work at it. Chemistry A-level requires a
commitment that was not asked of you at GSCE level. If you commit you will succeed. The commitment
you make when choosing the AS & A-level Chemistry course is as follows;

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To commit to 9 hours per fortnight of independent study most of which will be directed by
your teacher in the form of homework.
To commit to using all the resources at your disposal while completing all your homework to the
best of your ability.
To commit to handing your homework in before or on the deadline given by your teacher.
To come to lessons fully equipped including pens, pencils, rulers, erasers, calculator, text book,
file, paper, planner, lab coat and a periodic table.
To keep an organised and complete file of notes, work completed & homework that is brought
to every lesson.
To attend every lesson and arrive on time and ready to work.
To revise fully and effectively for every internal unit test & mock exam as well as for all
external exams.
I have read and understood the commitment necessary to succeed at AS & A-level Chemistry.
Signed: __________________________
date:___________________
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ASSESSMENT
INFORMATION &
RECORD OF
ACHIEVEMENT
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Internal Assessment
In both Year 1 (AS- year) & Year 2 (A-level) between September & May you will be assessed through a
series of standard homework tasks, end of unit tests and mock exams. The result of these
assessments will contribute to your attainment level which will be reported to you at regular intervals
but particularly at consultation evenings and in written reports. Your reported attainment level is
aimed at giving you an indication of what level you are working at and whether you are on or above
target. It will also be used to inform the grades reported on your UCAS application forms.
The assessment tasks within each academic year are;
o
o
o
2 x mock exams
9 x end of unit tests
9 x standard homework tasks
You will also be given normal homework tasks which are aimed at helping you learn.
Each assessment task will be given a different weighting when applied to the calculation of your overall
attainment level for each module and your overall attainment.
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Record of achievement - End of unit tests & Internal mock exams
My Target Level is: _________
It is very important that you take all tests and exams seriously and revise well. These tests will
closely match what you will need to do in the external AS/A-level exams. From the results you will be
able to identify areas you need to work on – this could be knowledge based or exam technique.
Year 1 – AS-Level
Record the details of all your internal assessments in the table below.
Date
Subject of end of unit test or
internal exam
Mark &
grade
End of unit tests – Year 1
MOCKS
Mock 1
Mock 2
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Possible improvements
Year 2 – A-Level
Date
Subject of end of unit test or
internal exam
Mark &
grade
End of unit tests – Year 2
MOCKS
Mock 1
Mock 2
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Possible improvements
Record of Achievement - Homework
My Target Level is:________
It is very important that you put the maximum amount of effort and thought into your homework. It is
part of the learning process. Use the table below to record details of your standard homework tasks
and attainment.
Year 1 – AS-Level
Date
Task
Mark &
grade
Standard Homework tasks – Year 1
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Possible improvements
Year 2 – A-Level
Date
Task
Mark &
grade
Standard Homework tasks – Year 2
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Possible improvements
External Assessment information
AS-Level
Paper 1 and Paper 2 paper will be taken at the end of Year 1 and will be the assessment for the ASlevel only (these exams will not contribute to the A-level qualification taken at the end of the second
year).
The practical skills tested will relate to the skills, data analysis skills and practical techniques
undertaken in the following 6 required practicals.
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A-Level
Paper 1, Paper 2 & Paper 3 will be taken at the end of Year 2 and will be the assessment for the Alevel.
The exams will test content from both Year 1 & Year 2.
The practical techniques & data analysis tested will relate to the skills and practical techniques
undertaken in the 12 required practicals needed for the practical endorsement.
Practical Endorsement
The practical endorsement will be a Pass/Fail mark written on the qualification certificate. The details
are provided in the practical Pupil Assessment Booklet. The 12 required practicals are given in the
following table.
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GENERAL INFORMATION
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Strategies for reaching and exceeding your target grade
The following are a list of strategies proven to help students reach and exceed their target grade:
1.
Review work done in class by reading around the subject. Look up web sites on the internet and
read for understanding. Read (for understanding) your notes and text book. If it doesn’t make
sense keep researching and working through the problem until it does.
2. Complete all homework set to a high standard.
a. To do this you must first understand the area of chemistry it relates to (by looking at
notes and reading around the subject),
b. Secondly try the homework questions using the text book and notes to help you,
c. Thirdly check your answers,
d. Finally review your answers prior to handing it in.
In this whole process make sure you
 focus and concentrate on the work completely (no music, TV or chatting),
 correct any mistakes.
 Check : units; calculations; rearranging of equations; inputs into calculators; use of
words e.g. is it molecules or ions?; definitions; state symbols; chemical equations;
charges; is it logical?; do you contradict yourself?; does it actually answer the
question?; does it cover all the mark points?; have you written down all relevant working
out?; Have you made all numerical substitution into equations clear?; Have you written
down all relevant information or have you summarised it too much?; Have you used full
sentences?
3. If you have finished work set then devise your own study e.g. practise exam questions which
can be found in the text books, Moodle or on the AQA website.
4. Memorise important information such as definitions, equations, sequences, examples.
5. Organise your file.
6. Organise your life by organising your work time and play time so your play time does not affect
your work time.
7. Organise your study time so you make the most of it.
8. Learn definitions off by heart.
9. Don’t study in the break out space – you are wasting your time! Study in the library or the
other quiet study areas for Year 12 & 13.
10. Be strict with your work ethic at school. You should be working between 9am and 3.05pm. The
only time you should be taking a break from work in school is between 11am to 11.15am and
1.20pm to 2pm. This means you need to use your study periods productively.
11. Draw up mind maps & spider diagrams for each topic. Make brief notes using your text book
and internet searches.
12. Keep reviewing and revising past work – you will need it.
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13. Read for understanding again! Make connections. If anything doesn’t make sense don’t leave it
until tomorrow, try to understand it today by concentrating on trying to understand it. Read
for understanding again!
14. You must spend at least 9 hours a fortnight outside lessons on Chemistry homework & review.
15. Review marked homework. Analyse to understand why you got things right and why you missed
marks. Don’t make the same mistake twice!
Homework
You are expected to complete homework set by the deadline given. Follow the advice given in step 2 in
the section ‘Strategies for reaching and exceeding your target grade’ as well as the advice below.
Do
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Start the homework the day you are given it
Using all the resources available, including text books and notes, complete and review at
least twice more before handing it in
Leave time to be able to work on any problems you encounter
Arrange to see your teacher at least 2 days before the Homework is due if you are
having problems
Attempt all Homework questions
Do Not
 Go to your teacher the day the homework is due and tell them you couldn’t do it
 Hand in incomplete HW
 Fail to hand in the HW on time
Contact with your teacher
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
You have 9 lessons of Chemistry a fortnight in which you can ask anything. If you are struggling
outside of class, although it is always better to try and solve a problem yourself (even if it
takes you hours – you will learn from the experience) then ask because if you are having
problems then you can bet other people are having problems and your teacher will either go
over the relevant area again in class or arrange a separate tutorial at lunch time.
BUT: before asking your teacher they will expect you to have had a go at the problem and be
able to tell them what you have done to try and solve the problem yourself.
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TOOLS
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Reading List
Core text books (provided by the school on loan)
1. AQA Chemistry for A-level Year 1 and AS, Hodder Education (for Year 12)
2. AQA Chemistry for A-level Year 2, Hodder Education (for Year 13)
Support material
3. Chemistry Handbook
4. Pupil Assessment Booklet (Practical)
5. Practical Booklet Year 1
6. Practical Booklet Year 2
7. Practical Techniques Booklet
8. There are many revision books available, choose one that is appropriate to your learning style.
Ensure it is for the AQA 2015 syllabus being studied.
Other – for reading around the subject to widen your understanding & application
9. Bill Bryson: A short history of everything
10. Periodic Tales, Hugh Aldersey-Williams
11. The Elements – A very Short Introduction, Philip Ball
12. Molecules – A very Short Introduction, Philip Ball
13. Molecules at an Exhibition, John Emsley
14. The Periodic Kingdom, Peter Atkins
APPs for phones
Look out for any mobile phone apps. There are many periodic table and Chemistry quiz question apps
available for free.
Internet
1. Moodle (may be being replaced by RM Portico so discuss with your teacher)
You can access Moodle from any computer within school or outside of school. If you go to the school
website there is a link to the Moodle log in page. Your login details are the same as your normal login
details for accessing the school system.
On Moodle there are various resources that will be helpful to you over the course of the year such as
the Standard homework you will be set & past papers.
2. http://www.aqa.org.uk/subjects/science/as-and-a-level/chemistry-7404-7405
The AQA board website is where you will find the course specifications, examiners reports, past paper
questions & answers.
3. The following are a list of websites that contain information to help you with homework, review work
or revision. I have not vetted all of these web sites so when using them read for understanding – if
they don’t make sense then check the chemistry.
 Youtube – Crash Course Chemistry
 www.s-cool.co.uk/a-level/chemistry
 www.chemguide.co.uk
 alevelchem.com
 www.docbrown.info
 www.chembook.co.uk
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Definitions
It is essential that you learn all the definitions you come across (and there really is NO substitute to
just memorizing them). The text books you are provided with have a glossary of definitions at the back
of the book. Use the table below to record any that you want to have quick reference to. It is very
important in Chemistry to be precise in your definitions.
Unit
Word
Definition
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Table of common ions and molecules
You will need to know these common ions so you must memorise them. You will use them throughout
your Chemistry A-level. There are gaps in the table so you can add to the ions as you learn more.
Positive ions
H+
Li+
K+
Mg2+
Ca2+
Zn2+
Cu2+
Fe2+
Co2+
Al3+
Fe3+
Cr3+
NH4+
Pb2+
Hydrogen ion
Lithium ion
Potassium ion
Magnesium ion
Calcium ion
Zinc ion
Copper (II) ion
Iron (II) ion
Cobalt ion
Aluminium ion
Iron (III) ion
Chromium ion
Ammonium ion
Lead ion
Negative ions
FClBrIO2S2OHNO3NO2HCO3CO32SO42CrO42Cr2O72MnO4C2O42-
Fluoride ion
Chloride ion
Bromide ion
Iodide ion
Oxide ion
Sulfide ion
Hydroxide ion
Nitrate (V) ion
Nitrate (III) ion
Hydrogencarbonate ion
Carbonate ion
Sulphate ion
Chromate (VI) ion
Dichromate (VI) ion
Manganate (VII) ion
Ethandioate ion
You will need to know these common molecules so you must memorise them. You will use them
throughout your Chemistry A-level.
Formula
H2
F2
Cl2
Br2
I2
O2
N2
CH4
NH3
SO2
NO
NO2
CO
CO2
Name
Hydrogen
Fluorine
Chlorine
Bromine
Iodine
Oxygen
Nitrogen
Methane
Ammonia
Sulphur dioxide
Nitrogen monoxide
Nitrogen dioxide
Carbon monoxide
Carbon dioxide
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Properties of common acids/alkalis/bases
Acids
Bases
Slippery to touch
pH > 7
turns red litmus blue
turns phenolphthalein pink
Turn methyl orange to yellow
When soluble in water contain OH- ions
Sour
pH < 7
turns blue litmus paper red
corrosive
Turn methyl orange to red
Contain H+ ions when in solution
Alkali: Has two possible definitions
1) A basic (as in it is a base) salt of an alkali metal (Group 1) or alkaline earth metal (Group 2).
2) A base that dissolves in water
General chemical reactions of acids
 Acid + Metal → Salt + Hydrogen
e.g. 2HCl + Mg → MgCl2 + H2

Acid + Metal oxide → Salt + Water
e.g. 2HNO3 + MgO → Mg(NO3)2 + H2O

Acid + Metal hydroxide → Salt + Water
e.g. H2SO4 + Mg(OH)2 → MgSO4 + 2H2O

Acid + Carbonates → Salt + Water + Carbon Dioxide
e.g. 2HCl + MgCO3 → MgCl2 + H2O + CO2

Acid + Bicarbonates → Salt + Water + Carbon Dioxide
e.g. HNO3 + Mg(HCO3)2 → Mg(NO3)2 + H2O + CO2
Some common acids – with their common states
1) Hydrochloric acid – HCl gas dissolved in water to make a solution
2) Sulphuric acid – solution
3) Nitric acid – solution
4) Citric acid – solid
5) Ethanoic acid – liquid (in solution is vinegar)
6) Tartaric acid – solid
7) Phosphoric acid - solution
Some common bases - For the following then sodium can be substituted for other group 1 or group 2
metals
8) Sodium hydroxide – solid, dissolves to form a solution
9) Sodium Carbonate – solid
10) Sodium hydrogencarbonate – solid
11) Sodium Oxide - solid
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Equations
Equations represent what happens when chemical reactions take place.
There are different types of equation and you will need to be able to use the correct equation in the
correct context.
Word equations
Describes how a reaction progresses.
e.g.
Magnesium + oxygen → Magnesium oxide
Symbol equations - Formula equations
These are equations using the formulae of substances that show how many atoms of each element are
involved and where they go in the products.
e.g.
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
Symbol equations - Ionic equations
These equations only consider the ions taking part in the reaction
e.g.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
if the equation is split into just the reacting ions
H+ + Cl- + Na+ + OH- →
Na+ + Cl- +
H 2O
we can cancel the ions that don’t change and appear on each side of the equation so overall the ionic
equation is
H+ +
OH- →
H 2O
Remember: H2O is not ionic so cannot be split into ions (it is a molecule).
Symbol equations - Ion-electron equations / Half equations
These show the ionisation of the species.
Li+(g)
e.g. Li(g) →
+
e-
When asked for an equation in a question a balanced symbol equation, not a word equation, is
always expected.
Balanced equations
It is essential that you are able to balance full equations, ionic equations and half equations. Remember
these rules when balancing equations. The rules apply to most equation. During the course you will be
shown other ways of balancing complicated looking equations e.g. half equations for REDOX reactions.
1.
Write down the formulae for the reactants and products. The formula is fixed and cannot be
changed.
C3H8 + O2 →
CO2 + H2O
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2. The only way you can balance an equation is to place numbers in front of the formulae (you
cannot change the formula by changing the subscript numbers). You must place numbers in
front of the formulae until there are the same number of atoms of an element on the reactant
side of the equation as there are on the product side of the equation.
NO!
C3H8 + O3 →
C3O2
+
H 8O
YES!
C3H8 + 5O2 → 3CO2 + 4H2O
There are now 3 carbons, 8 hydrogens & 10 oxygens on each side.
3. You must also ensure that the charges balance.
H+ +
OH- →
H 2O
The positive and negative charges make an overall zero charge on the reactant side while the products
have no charge so the charges are balanced.
2Na+ +
SO42- →
Na2SO4
By placing a 2 in front of the sodium ion not only have we balanced the sodium but also the charges –
there are two positive and two negative charges on the reactant side resulting in an overall zero
charge and zero charge on the products side.
4. If required you then put state symbols in the balanced equation to show the state of the
reactants and products. Remember to think carefully; is the water produced a gas or a liquid in
your reaction?; Precipitates are solids; Is a gas given off?: Acids tend to be dissolved in water
therefore they are aqueous.
(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)
MgCO3(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) +
CO2(g)
5. NOW CHECK – do all the elements balance, do all the charges balance, are the state symbols
correct.
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Practical techniques
Throughout the course you will be shown and have the opportunity to practise practical techniques. You
are required to know how to carry out these techniques and why you use them. You may be asked about
them in the ISA exam. Use the following Youtube clips to help you better understand some of the
processes you will learn – especially before an ISA exam. The information presented below is not
exhaustive and may be added to.
1. Go to www.youtube.com
2. In the search box put in the following key words/titles to bring up the appropriate clip
a) RSCteacherfellows then
 Recrystallisation
 Hot filtration
 Melting point determination
 Running an Infrared Spectrum
 Distillation
 Vacuum Filtration
 Thin Layer Chromatography
 Heating under reflux
 Weighing compounds using a balance
b) Preparing a standard solution
c) Pipetting Technique
d) Titration technique using a burette
e) Organic Chem: How to Flute filter paper
f) Recrystallization of an impure compound (note: they refer to a conical flask as an ergimyer flask)
g) Reflux reactions
h) Distillation
i) Extraction
j) Calorimetry, heat of neutralization.avi
You may be asked as part of your exam (particularly the ISA exam) the reasons for carrying out a
particular practical technique. The following information will help but you will have to add to this
knowledge throughout the course.
Technique
Why it may be used
Filtration
To remove an insoluble impurity
or to collect a recrystallised product
or to collect a precipitate
or to separate a solid product from a solution.
Hot filtration
To remove an insoluble impurity from a sparingly soluble solute
or from a solute (in a solution) requiring recrystallization. It is
necessary to keep the solution hot to stop the product crystallising
Fluted filter paper
To increase surface area to allow a faster filtration
Recrystallisation
To purify the product / To remove a soluble impurity
Distillation
To remove / purify / separate a liquid
Fractional distillation To separate liquids with different boiling points
Reflux
Gives all reactants chance to react. Speeds up reactions without the
reactants and products escaping
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Practical Observations
It is important when making practical observations, or answering exam questions which ask for
observations that you are descriptive and precise. The following are key descriptive words and phrases
that you should use where appropriate.
a) Where there is the word ‘colour’ then insert the appropriate colour e.g. Blue solution
b) Where there is the word ‘specific’ then insert the appropriate word e.g. carbon dioxide gas
produced
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White precipitate produced
’colour’ precipitate produced
Clear, colourless solution
Clear, colourless liquid
‘Colour’ solution
Gas produced
Bubbles of gas produced
Colourless gas produced
‘Specific’ gas produced
Liquid condenses
No visible change
White solid produced
‘Colour’ solid produced
Increase / decrease in temperature observed
Fizzing / popping sound heard
Sharp / pungent smell detected
Vapour given off
Just saying Lime water goes cloudy is no longer enough detail, instead you would say a white
precipitate was produced.
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Practical Accuracy
In an experiment there is an experimental uncertainty (often called 'experimental error').
Experimental uncertainty arises because of:
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Limits in the precision of the measuring apparatus.
Imperfections in the experimental procedure itself.
Judgements made by the operator.
Operator error
To improve accuracy by you, the operator:


always record results to the precision allowed by the apparatus. For example, if the balance
reads to 0.01g, write 6.78g as 6.78g - make sure you don’t round it up to 6.8g or 7g.
always record results the moment you take them e.g. write down the mass used when next to
the balance, not after you have returned to your desk.
When doing calculations do not round up numbers too early – if you do you may loose accuracy (see also
section on significant figures).
Apparatus Errors
You must be able to calculate the size of errors in practical work.
Percentage error
=
Margin of error
Quantity measured
x
100
Mass
1. Consider weighing 1g of solid. If you use a two decimal place balance, the mass recorded will be to
the nearest 0.01g. In this example, the % error will be:
0.01 (margin of error) x 100 = 1%
1 (quantity measured)
2. Consider weighing the same 1g of solid on a three decimal place balance. The mass recorded will be
to the nearest 0.001g, and so the % error will be:
0.001 x 100 = 0.1%
1
There is much less error involved in this procedure.
3. Consider weighing 10g of solid on the two decimal place balance. In this case the % error will be:
0.01 x 100 = 0.1%
10
This error is less than weighing 1g on this balance
Choose the right balance for the amount of material to be weighed.
Volume
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4. Consider measuring 25cm3 in a 25cm3 measuring cylinder. The measurement will normally be to the
nearest 0.5cm3 (the accuracy will be quoted on the apparatus itself), so the % error will be:
0.5 x 100 = 2.0%
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5. Sometimes on equipment the accuracy is quoted as a ± a value e.g. on a burette the accuracy is given
as ±0.05cm3 so if 25cm3 of solution was dispensed then the % error will be:
0.05 x 100 = 0.2%
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However if 2cm3 of solution was dispensed then the % error will be:
0.05 x 100 = 2.5%
2
If choosing equipment to measure a volume choose the correct equipment based on the amount
to be measured and the possible error.
Other equipment
The errors for other equipment can be calculated in a similar way.
Total apparatus error
To work out the total error attributed to the apparatus for a particular experiment, then add all the
separate equipment errors together.
For example, imagine a pupil doing an experiment where she measured out 1.245 g of a base, made it up
to 250 cm3 of solution in a volumetric flask, pipetted 25 cm3 of that solution into a conical flask, and
then found that it reacted with 23.30 cm3 of acid in a titration using a burette.
Balance
( 0.001 g)
100 x (0.001/1.245) = 0.08%
Pipette
( 0.1 cm3)
100 x (0.1/25)
= 0.40%
100 x (0.1/250)
= 0.04%
3
Volumetric flask ( 0.1 cm )
Burette
3
( 0.15 cm )
100 x (0.15/23.30) = 0.64%
Total apparatus error
= 1.16%
This means that the result of the experiment should be within 1.16% of the correct value.
When you design experiments, you should aim to ensure that the total apparatus error is minimised by
working on a suitable scale and with suitable apparatus. A very small titre for example (e.g. 5 cm 3)
leads to a very large apparatus error for the burette (3%).
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Yields
1) Atom economy tells us in theory how many atoms must be wasted in a reaction or how many atoms
enter the desired product:
% atom economy = mass of desired product x 100%
Total mass of reactants
The mass is calculated from the balanced equation.
2) The yield tells us about the practical efficiency of the process:
Yield of a chemical reaction = The number of actual moles of a specified product x 100%
Theoretical maximum number of moles of the product
Or
=
actual yield in grams x 100%
Theoretical yield in grams
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Hazards / precautions
Hazards
New international hazard symbols are coming into use. You must know these symbols.
SYMBOL
HAZARD
Caution – used for less serious health
hazards like skin irritation.
Flammable
Dangerous to the environment
Explosive
Oxidising
Longer term health hazards such as
carcinogenicity
Corrosive
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Gas under pressure
Toxic
Precautions
You may be asked about how to minimise risk to do with hazards. When answering these sorts of
questions you must be specific to the question. E.g if the question asks ‘How can contact with the skin
be minimised?’ then it is no use answering ‘by washing hands’ because this is not a precaution – this is
something you do after contact has been made. The following table gives some precautions to be taken
against hazards but it is by no means exhaustive and you must continue to add to it as you gain
experience in the laboratory.
Hazard
Chemical substance
Possible harm
Contact with the person
causes irritation, burns,
poisoning, death (dependent
on hazard symbols)
Breaking glass e.g. breakage,
putting bungs in test tubes,
putting pipette holders on
pipettes
Cuts
Bunsen burner, heating
mantle, hot equipment
Burns / fires
Bags / stools
Trip hazard – someone falls
over
Possible contact with
substance causing irritation,
burns, poisoning (dependent
on hazard symbols)
Chemical spills
Precaution to reduce risk
Use gloves
Place lids on containers
Mop up spillages
Wear safety glasses
Wear a mask
Wear safety glasses
Use appropriate apparatus to hold
equipment e.g. test tube rack
Hold glassware close to where the
bung/holder is being inserted
Tie hair back
Leave hot equipment to cool
Do not have flammable substances
nearby
Bags and stools under the table out of
the way
Keep tops on bottles when not in use
Mop up spills straight away
When writing a risk assessment it is useful to set out a table such as the one above to record the
hazard, the possible harm and the precaution to reduce the risk.
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AQA PROVIDED
INFORMATION
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A. Tabulating data
It is important to keep a record of data whilst carrying out practical work. Tables should have clear
headings with units indicated using a forward slash before the unit.
Time
/ min
0
1
2
Temperature
/ oC
14.8
14.7
14.6
Although using a forward slash is the standard format, other formats are generally acceptable. For
example:
Volume in
cm3
15
25
35
Time
taken in s
23
45
56
Concentration
(mol dm-3)
1.0
1.5
2.0
Time (s)
152
93
54
It is good practice to draw a table before an experiment commences and then enter data straight into
the table. This can sometimes lead to data points being in the wrong order. For example, when
studying the pH change in an acid-base titration, a student may do a number of pH measurements at
10, 20, 25, 30 and 35 cm3 of reagent added, and then investigate the area between 20 and 30 further
by adding readings at 22, 24, 24.5, 25, 25,5, 26, 28. Whilst this is perfectly acceptable, it is
generally a good idea to make a fair copy of the table in ascending order of temperature to enable
patterns to be spotted more easily. Reordered tables should follow the original data if using a lab
book, data should not be noted down in rough before it is written up.
It is also expected that the independent variable is the left hand column in a table, with the
following columns showing the dependent variables. These should be headed in similar ways to
measured variables. The body of the table should not contain units.
Tabulating logarithmic values
When the logarithm is taken of a physical quantity, the resulting value has no unit. However, it is
important to be clear about which unit the quantity had to start with. The logarithm of a time in
seconds will be very different from the logarithm of the same time in minutes.
These should be included in tables in the following way:
Reading
number
1
2
3
time / s
log (time/s)
2.3
3.5
5.6
0.36
0.54
0.75
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B. Significant figures
Data should be written in tables to the same number of significant figures. This number should be
determined by the resolution of the device being used to measure the data or the uncertainty in
measurement. For example, a sample labelled as “1 mol dm-3 acid” should not be recorded in a table
of results as 1.0 mol dm-3.
There is sometimes confusion over the number of significant figures when readings cross multiples
of 10. Changing the number of decimal places across a power of ten retains the number of
significant figures but changes the accuracy. The same number of decimal places should therefore
generally be used, as illustrated below.
0.97
0.98
0.99
1.00
1.10
99.7
99.8
99.9
100.0
101.0
It is good practice to write down all digits showing on a digital meter.
Calculated quantities should be shown to the number of significant figures of the data with the least
number of significant figures.
Example:
Calculate the concentration, in mol dm–3, of a solution of sodium hydroxide that contains 0.28 mol
of NaOH in 465 cm3 of water.
Concentration = 0.28 x 1000 = 0.59
475
Note that the concentration can only be quoted to two significant figures as the number of moles is
only quoted to two significant figures.
Note: Counting the significant figures in a number
Read the number from left to right and count all the digits starting with the first digit that is not
zero. The examples below all have four significant figures:





0.09047
1.794
2.560 (because this number has a zero that is to the right of a decimal place and is shown
it is counted as a significant figure)
0.007503
2.085 x 105 (using scientific notation like this shows without ambiguity the number of
significant figures – if it was shown as 2085000 we would not know if this was 4,5,6 or 7 s.f.
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C. Uncertainties
Students should know that every measurement has some inherent uncertainty.
The uncertainty in a measurement using a particular instrument is no smaller than plus or
minus half of the smallest division or greater. For example, a temperature measured with a
thermometer is likely to have an uncertainty of ±0.5 °C if the graduations are 1 °C apart.
Students should be aware that measurements are often written with the uncertainty. An example of
this would be to write a voltage was (2.40 ± 0.005) V.
Measuring length
When measuring length, two uncertainties must be included: the uncertainty of the placement of the
zero of the ruler and the uncertainty of the point the measurement is taken from.
As both ends of the ruler have a ±0.5 scale division uncertainty, the measurement will have an
uncertainty of ±1 division.
area of uncertainty
object
ruler
For most rulers, this will mean that the uncertainty in a measurement of length will be ±1 mm.
Other factors
There are some occasions where the resolution of the instrument is not the limiting factor in the
uncertainty in a measurement.
Best practice is to write down the full reading and then to write to a fewer significant figures when
the uncertainty has been estimated.
Examples:
A stop watch has a resolution of hundredths of a second, but the uncertainty in the
measurement is more likely to be due to the reaction time of the experimenter. Here, the
student should write the full reading on the stop watch (eg 12.20 s) and reduce this to 12 s
later.
If a student measures the length of a piece of wire, it is very difficult to hold the wire completely
straight against the ruler. The uncertainty in the measurement is likely to be higher than the ±1 mm
uncertainty of the ruler. Depending on the number of “kinks” in the wire, the uncertainty could be
reasonably judged to be nearer ± 2 or 3 mm.
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Repeated measurements
If measurements are repeated, the uncertainty can be calculated by finding half the range of the
measured values.
For example:
1
12.3
Repeat
Time (s)
2
13.2
3
12.7
4
12.2
13.2 – 12.2 = 01.0 so
Mean time: (12.6 ± 00.5) s
Percentage uncertainties
The percentage uncertainty in a measurement can be calculated using:
The percentage uncertainty in a repeated measurement can be calculated using:
Titration
Titration is a special case where a number of factors are involved in the uncertainties in the
measurement.
Students should carry out a rough titration to determine the amount of titrant needed. This is to
speed up the process of carrying out multiple samples. The value of this titre should be ignored in
subsequent calculations.
In titrations one single titre is never sufficient. The experiment is usually done until there are at least
two titres that are concordant ie within a certain allowable range, often 0.10 cm3. These values are
then averaged.
For example:
Titration
Final reading
Initial reading
Titre / cm3
Rough
1
2
3
24.20 47.40 24.10 47.35
0.35
24.20 0.65 24.10
23.85 23.20 23.45 23.25
Here, titres 1 and 3 are within the allowable range of 0.10 cm3 so are averaged to 23.23 cm3.
Unlike in some Biology experiments (where anomalous results are always included unless there is
good reason not to), in Chemistry it is assumed that repeats in a titration should be concordant. If
they are not then there is likely to have been some experimental error. For example the wrong
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volume of solution added from the burette, the wrong amount of solution measuring the pipette or
the end point might have been misjudged.
The total error in a titre is caused by three factors:
Error
Reading the burette at the start of the titration
Reading the burette at the end of the titration
Judging the end point to within one drop
Total
Uncertainty
Half a division = ±0.05 cm3
Half a division = ±0.05 cm3
Volume of a drop = ± 0.05 cm3
± 0.15 cm3
This will, of course, depend on the glassware used, as some burettes are calibrated to a higher
accuracy than others.
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Uncertainties from gradients
To find the uncertainty in a gradient, two lines should be drawn on the graph. One should be the
“best” line of best fit. The second line should be the steepest or shallowest gradient line of best fit
possible from the data. The gradient of each line should then be found.
The uncertainty in the gradient is found by:
×
Note the modulus bars meaning that this percentage will always be positive.
Best gradient
Worst gradient could be either:
Steepest gradient possible
or
Shallowest gradient possible
In the same way, the percentage uncertainty in the y-intercept can be found:
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Combining uncertainties
Percentage uncertainties should be combined using the following rules:
Combination
Adding or
subtracting values
Operation
Add the absolute
uncertainties
Δa = Δb + Δc
Example
Initial volume in burette = 3.40 ± 0.05 cm3
Final volume in burette = 28.50 ± 0.05 cm3
Titre = 25.10 ± 0.10 cm3
Multiplying values
Add the percentage
uncertainties
Mass = 50.0 ± 0.1 g
εa = εb + εc
Percentage uncertainty in mass = 0.20%
Temperature rise (T) = 10.9 ± 0.1 oC
Percentage uncertainty in T = 0.92 %
Heat change = 2278 J
Percentage uncertainty in heat change = 1.12 %
Absolute uncertainty in heat change = ± 26 J
(Note – the uncertainty in specific heat is taken to
be zero)
Dividing values
Add the percentage
uncertainties
Mass of salt in solution= 100 ± 0.1 g
εa = εb + εc
Percentage uncertainty in mass = 0.1 %
Volume of solution = 250 ± 0.5 cm3
Percentage uncertainty in volume = 0.2 %
Concentration of solution = 0.400 g cm–3
Percentage uncertainty of concentration = 0.3 %
Absolute uncertainty of concentration = ± 0.0012 g
cm–3
Power rules
Multiply the
percentage
uncertainty by the
power
Concentration of H+ ions = 0.150 ± 0.001 mol dm–3
rate of reaction = k[H+]2 = 0.207 mol dm–3 s–1
(Note – the uncertainty in k is taken as zero and its
value in this reaction is 0.920 dm6 mol–2 s–1)
εa = c × εb
Percentage uncertainty in concentration = 0.67 %
Percentage uncertainty in rate = 1.33 %
Absolute uncertainty in rate = ± 0.003 mol dm–3 s–1
Note: Absolute uncertainties (denoted by Δ) have the same units as the quantity.
Percentage uncertainties (denoted by ε) have no units.
Uncertainties in trigonometric and logarithmic functions will not be tested in A-level exams.
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D. Graphing
Graphing skills can be assessed both in written papers for the A-level grade and by the teacher
during the assessment of the endorsement. Students should recognise that the type of graph that they
draw should be based on an understanding of the data they are using and the intended analysis of
the data. The rules below are guidelines which will vary according to the specific circumstances.
Labelling axes
Axes should always be labelled with the quantity being measured and the units. These should be
separated with a forward slash mark:
time / seconds
length / mm
Axes should not be labelled with the units on each scale marking.
Data points
Data points should be marked with a cross. Both  and  marks are acceptable, but care should be
taken that data points can be seen against the grid.
Error bars can take the place of data points where appropriate.
Scales and origins
Students should attempt to spread the data points on a graph as far as possible without resorting to
scales that are difficult to deal with. Students should consider:
 the maximum and minimum values of each variable
 the size of the graph paper
 whether 0.0 should be included as a data point
 how to draw the axes without using difficult scale markings (eg multiples of 3, 7, 11 etc)
 the plots should cover at least half of the grid supplied for the graph.
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This graph has well-spaced marking
points and the data fills the paper.
Each point is marked with a cross
(so points can be seen even when a
line of best fit is drawn).
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This graph is on the limit of
acceptability. The points do not
quite fill the page, but to spread
them further would result in the use
of awkward scales.
At first glance, this graph is well
drawn and has spread the data
out sensibly.
However, if the graph were to
later be used to extrapolate the
line, the lack of appropriate
space could cause problems.
Increasing the axes to ensure
sufficient room is available is a
skill that requires practice and
may take a couple of attempts.
Note: No zero on this graph.
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Lines of best fit
Lines of best fit should be drawn when appropriate. Students should consider the following when
deciding where to draw a line of best fit:
 Are the data likely to have an underlying equation that it is following (for example, a relationship
governed by a physical law)? This will help decide if the line should be straight or curved.
 Are there any anomalous results?
There is no definitive way of determining where a line of best fit should be drawn. A good rule of
thumb is to make sure that there are as many points on one side of the line as the other. Often the
line should pass through, or very close to, the majority of plotted points. Graphing programs can
sometimes help, but tend to use algorithms that make assumptions about the data that may not be
appropriate.
Lines of best fit should be continuous and drawn with a thin pencil that does not obscure the points
below and does not add uncertainty to the measurement of gradient of the line.
Not all lines of best fit go through the origin. Students should ask themselves whether a 0 in
the independent variable is likely to produce a 0 in the dependent variable. This can provide
an extra and more certain point through which a line must pass. A line of best fit that is
expected to pass through (0,0) but does not would some systematic error in the experiment.
This would be a good source of discussion in an evaluation.
Dealing with anomalous results
At GCSE, students are often taught to automatically ignore anomalous results. At A-level students
should think carefully about what could have caused the unexpected result - for example, if a
different experimenter carried out the experiment, similarly, if a different solution was used or a
different measuring device. Alternatively, the student should ask if the conditions the experiment
took place under had changed (for example at a different temperature). Finally, they can evaluate
about whether the anomalous result was the result of an accident or experimental error. In the case
where the reason for an anomalous result occurring can be identified, the result should be ignored.
In presenting results graphically, anomalous points should be plotted but ignored when the line of
best fit is being decided.
Anomalous results should also be ignored where results are expected to be the same (for example in
a titration in chemistry).
Where there is no obvious error and no expectation that results should be the same, anomalous
results should be included. This will reduce the possibility that a key point is being overlooked.
Please note: when recording results it is important that all data are included. Anomalous results
should only be ignored at the data analysis stage.
It is best practice whenever an anomalous result is identified for the experiment to be repeated. This
highlights the need to tabulate and even graph results as an experiment is carried out.
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Measuring gradients
When finding the gradient of a line of best fit, students should show their working by drawing a
triangle on the line. The hypotenuse of the triangle should be at least half as big as the line of best
fit.
The line of best fit here has
an equal number of points on
both sides. It is not too wide
so points can be seen under
it.
The gradient triangle has
been drawn so the
hypotenuse includes more
than half of the line.
In addition, it starts and ends
on points where the line of
best fit crosses grid lines so
the points can be read easily
(this is not always possible).
Δx
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The equation of a straight line
Students should be able to translate graphical data into the equation of a straight line.
Where y is the dependent variable, m is the gradient, x is the independent variable and c is the yintercept.
35
30
25
20
Δy
Δy
15
10
Δx
5
Δx
y-intercept
0
0
20
40
60
80
Δy = 28 – 9 = 19
Δx = 90 – 10 = 80
gradient = 19 / 80 = 0.24 (2 sf)
y-intercept = 7.0
equation of line:
y = 0.24 x + 7.0
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100
Testing relationships
Sometimes it is not clear what the relationship between two variables is. A quick way to find a
possible relationship is to manipulate the data to form a straight line graph from the data by
changing the variable plotted on each axis.
For example:

Raw data and graph
x
0
10
20
30
40
50
60
70
80
90
100
y
0.00
3.16
4.47
5.48
6.32
7.07
7.75
8.37
8.94
9.49
10.00
11
10
9
8
7
6
y
5
4
3
2
1
0
0
20
40
60
80
100
x
This is clearly not a straight line graph. The relationship between x and y is not clear.

Manipulated data and graphs
A series of different graphs can be drawn from these data. The one that is closest to a straight line is
a good candidate for the relationship between x and y.
x
y
0
10
20
30
40
50
60
70
80
90
100
0.00
3.16
4.47
5.48
6.32
7.07
7.75
8.37
8.94
9.49
10.00
√y
0.00
1.78
2.11
2.34
2.51
2.66
2.78
2.89
2.99
3.08
3.16
y2
0.00
10.00
20.00
30.00
40.00
50.00
60.00
70.00
80.00
90.00
100.00
y3
0.00
32
89
160
250
350
470
590
720
850
1000
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This is an idealised set of
data to illustrate the point.
The straightest graph is y
against x2, suggesting that
the relationship between x
and y is
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More complex relationships
Graphs can be used to analyse more complex relationships by rearranging the equation into a form
similar to y=mx+c.
Example one: testing power laws
A relationship is known to be of the form y=Axn, but n is unknown.
Measurements of y and x are taken.
A graph is plotted with log(y) plotted against log(n).
The gradient of this graph will be n, with the y intercept log(A).
Example two
The equation that relates the rate constant of a reaction to temperature is
This can be rearranged into
So a graph of
of
against
should be a straight line, with a gradient of
.
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and a y-intercept
E. Glossary of terms
The following subject specific vocabulary provides definitions of key terms used in AQA's AS and
A-level Biology, Chemistry and Physics specifications.
Accuracy
A measurement result is considered accurate if it is judged to be close to the true value.
Calibration
Marking a scale on a measuring instrument.
This involves establishing the relationship between indications of a measuring instrument and
standard or reference quantity values, which must be applied.
For example, placing a thermometer in melting ice to see whether it reads 0⁰C, in order to check if
it has been calibrated correctly.
Data
Information, either qualitative or quantitative, that have been collected.
Errors
See also uncertainties.
measurement error
The difference between a measured value and the true value.
anomalies
These are values in a set of results which are judged not to be part of the variation caused
by random uncertainty.
random error
These cause readings to be spread about the true value, due to results varying in an
unpredictable way from one measurement to the next.
Random errors are present when any measurement is made, and cannot be corrected. The
effect of random errors can be reduced by making more measurements and calculating a
new mean.
systematic error
These cause readings to differ from the true value by a consistent amount each time a
measurement is made.
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Sources of systematic error can include the environment, methods of observation or
instruments used.
Systematic errors cannot be dealt with by simple repeats. If a systematic error is suspected,
the data collection should be repeated using a different technique or a different set of
equipment, and the results compared.
zero error
Any indication that a measuring system gives a false reading when the true value of a
measured quantity is zero, eg the needle on an ammeter failing to return to zero when no
current flows.
A zero error may result in a systematic uncertainty.
Evidence
Data that have been shown to be valid.
Fair test
A fair test is one in which only the independent variable has been allowed to affect the dependent
variable.
Hypothesis
A proposal intended to explain certain facts or observations.
Interval
The quantity between readings eg a set of 11 readings equally spaced over a distance of 1 metre
would give an interval of 10 centimetres.
Precision
Precise measurements are ones in which there is very little spread about the mean value.
Precision depends only on the extent of random errors – it gives no indication of how close results
are to the true value.
Prediction
A prediction is a statement suggesting what will happen in the future, based on observation,
experience or a hypothesis.
Range
The maximum and minimum values of the independent or dependent variables;
For example a range of distances may be quoted as either:
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'From 10cm to 50 cm' or
'From 50 cm to 10 cm'
Repeatable
A measurement is repeatable if the original experimenter repeats the investigation using same
method and equipment and obtains the same results.
Reproducible
A measurement is reproducible if the investigation is repeated by another person, or by using
different equipment or techniques, and the same results are obtained.
Resolution
This is the smallest change in the quantity being measured (input) of a measuring instrument that
gives a perceptible change in the reading.
Sketch graph
A line graph, not necessarily on a grid, that shows the general shape of the relationship between
two variables. It will not have any points plotted and although the axes should be labelled they may
not be scaled.
True value
This is the value that would be obtained in an ideal measurement.
Uncertainty
The interval within which the true value can be expected to lie, with a given level of confidence or
probability eg “the temperature is 20 °C ± 2 °C, at a level of confidence of 95 %”.
Validity
Suitability of the investigative procedure to answer the question being asked. For example, an
investigation to find out if the rate of a chemical reaction depended upon the concentration of one
of the reactants would not be a valid procedure if the temperature of the reactants was not
controlled.
Valid conclusion
A conclusion supported by valid data, obtained from an appropriate experimental design and
based on sound reasoning.
Variables
These are physical, chemical or biological quantities or characteristics.
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categoric variables
Categoric variables have values that are labels eg names of plants or types of material or
reading at week 1, reading at week 2 etc.
continuous variables
Continuous variables can have values (called a quantity) that can be given a magnitude
either by counting (as in the case of the number of shrimp) or by measurement (eg light
intensity, flow rate etc).
control variables
A control variable is one which may, in addition to the independent variable, affect the
outcome of the investigation and therefore has to be kept constant or at least monitored.
dependent variables
The dependent variable is the variable of which the value is measured for each and every
change in the independent variable.
independent variables
The independent variable is the variable for which values are changed or selected by the
investigator.
nominal variables
A nominal variable is a type of categoric variable where there is no ordering of categories
(eg red flowers, pink flowers, blue flowers)
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V1.6
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