Unit 1: Chemical Changes and Structure
Controlling the rate of reaction
Lesson 1: Introduction(s)
How was everyone’s summer? See any good science?
What does everyone want to get out of this - why have they chosen Chemistry Higher?
Overview of what we’re covering this year - how to make chemistry make sense! Structure of the course, in three major units with increasing numbers of sub-topics in each one. What is required, practically speaking.
How do they see the relations between chemistry and everything else?
What did they find easy in Chemistry GCSE? What did they find difficult?
Equipment: They’ll need the textbook, paper, a ringbinder, a calculator. That’s this week’s homework.
Expectations: I expect them to do the work required; to think about chemistry and how it relates to other things; to arrive on time and not mess around too much.
Revisiting the things they should know already about rates of reaction
Controlling the rate of reaction
Lots of familiar territory here, but actually covering quite a bit of new ground
- Activation energy
- Activated complex
- Collision geometry
- Collision theory
- Enthalpy change
- Rate of reaction
Lesson 2: Applying this knowledge
Take turns to name one thing that influences the rate of a reaction
Questions from book and discussion of tricky points?
Experiment from http://www.rsc.org/learn-chemistry/resource/res00000413/rate-of-reaction-the-effects-of-concentration-and-temperature (CCE 29)
- Place 50 cm3 of solution A in a 250 cm3 beaker.
- Place the same volume of solution B in a second beaker.
- Mix the two solutions by pouring from one beaker into the other several times.
- Note the time required for a reaction to occur (formation of blue colour).
- Repeat, but use solution A that has been diluted to one half the concentration. Note
the time for the reaction to occur.
- Repeat using solution A warmed to 35 °C. Note the time for a reaction to occur.
Visible activated complex demonstration (CCD 1)
Hydrogen peroxide oxidises potassium sodium tartrate (Rochelle salt) to carbon
dioxide. The reaction is catalysed by cobalt(II) chloride. When solutions of hydrogen
peroxide and Rochelle salt are mixed, carbon dioxide is slowly evolved. The addition
of cobalt(II) chloride causes the reaction to froth, indicating a large increase in the
reaction rate. At the same time the colour of the cobalt(II) chloride turns from pink to
green (an activated complex), returning to pink again within a few seconds as the
reaction dies down. This indicates that catalysts actually take part in the reaction and
are returned unchanged when the reaction is complete.
▼ Bunsen burner, tripod, gauze and heat-proof mat.
▼ One 250 cm3
▼ One 0 –100 °C thermometer.
▼ One 25 cm3
▼ One dropping pipette.
▼ Access to overhead projector (optional).
The quantities given are for one demonstration.
▼ 5 g of potassium sodium tartrate-4-water (Rochelle salt, potassium sodium 2,3-
▼ 0.2 g of cobalt(II) chloride-6-water (CoCl2
▼ 20 cm3
of 20 volume (ie approximately 6 %) hydrogen peroxide solution
▼ 65 cm3
of deionised water.
▼ About 200 cm3
of crushed ice (optional).
Before the demonstration
Make a solution of 0.2 g of cobalt chloride-6-water in 5 cm3
of deionised water.
Make a solution of 5 g of Rochelle salt in 60 cm3
of deionised water in a 250 cm3
2 Classic Chemistry Demonstrations
Add 20 cm3 of 20 volume hydrogen peroxide to the solution of Rochelle salt and heat
the mixture to about 75 °C over a Bunsen burner. There will be a slow evolution of
gas showing that the reaction is proceeding. Stirring the solution makes the evolution
of gas more obvious. Now add the cobalt chloride solution to the mixture. Almost
immediately the pink solution will turn green and after a few seconds vigorous
evolution of gas starts and the froth will rise almost to the top of the beaker. Within
about 30 seconds, the frothing subsides and the pink colour returns.
Stand the beaker on a overhead projector to make the evolution of gas before the
addition of the catalyst more easily visible.
The green activated complex can be trapped if a sample of the green solution is
withdrawn with a dropping pipette and then transferred to a test-tube that is cooled in
crushed ice. The solution remains green for some time.
If the reaction is considered to be going too fast for easy observation, carry it out
at a lower temperature (although this will make it less easy to see the evolution of
before adding the catalyst).
The Periodic Table: bonding and structure
Lesson 3: The Periodic Table - bonding and structure (1)
The first 20 elements in the Periodic Table are categorised according to bonding and structure. Periodic trends and underlying patterns and principles.
Which parts of the periodic table tend to be solid at the highest temperatures?
- Quick recap of bonding types
- Intermolecular forces - focus on London dispersion forces for now
With help of Molymod:
- Types of covalent molecules
- Halogens, with single bonds
- Oxygen, with its double bond
- Sulfur also forms two or more bonds, but has much stronger London dispersion forces owing to its greater size, hence it is much less gassy than oxygen
- Nitrogen and phosphorus form three bonds, but where nitrogen usually forms a single, extremely strong triple bond, phosphorus makes tetrahedra, and in sharp contrast to nitrogen, is extremely reactive
- Carbon’s various exciting allotropes, and silicon’s related networks
- Boron’s unexpected but logical habit of forming icosahedra
Categorise the types of covalent structures
Lesson 4: The Periodic Table - bonding and structure (2)
Allotropes of sulfur
Demonstration of allotropes of sulfur following directions in Classic Chemical Demonstrations
(answering some questions from the textbook while the oil heats up)
Properties of sulphur, allotropy, the relationship between properties and structure.
About 20 min.
Sulphur is heated gently and the various forms are:
▼ an amber, mobile liquid consisting of S8 rings;
▼ a viscous liquid consisting of long tangled chains of sulphur; and
▼ a dark, mobile liquid consisting of shorter sulphur chains.
The first form is allowed to cool and needle-shaped crystals of monoclinic sulphur
The third form is plunged into cold water and plastic sulphur is formed.
Sulphur is dissolved in toluene or xylene and the solvent allowed to evaporate.
Crystals of rhombic sulphur form.
▼ Two boiling tubes.
▼ Holder to hold the tubes while heating them.
▼ One 250 cm3
conical flask with cork to fit.
▼ One 250 cm3
▼ One 1 dm3
▼ Petri dish or watch glass.
▼ Microscope – ideally a projection or video microscope.
▼ Bunsen burner, tripod and gauze.
▼ Heat-proof mat.
▼ Access to a fume cupboard.
The quantities given are for one demonstration.
▼ About 60 g of powdered roll sulphur. Note that ‘flowers of sulphur’ is not
suitable because it contains a lot of insoluble amorphous sulphur.
▼ About 700 cm3
of cooking oil or other high boiling point oil.
▼ About 100 cm3 of toluene (methylbenzene) or xylene (dimethylbenzene).
▼ Filter paper – about 18 cm in diameter.
Before the demonstration
Two thirds fill a 1 dm3
beaker with cooking oil and heat to about 130 °C. Half fill a
beaker with cold water.
Two-thirds fill two boiling tubes with broken/powdered roll sulphur (about 20 g in
each tube) and place in the oil bath. The sulphur will melt to a clear, amber,
transparent, mobile liquid. This will take about 15 min. Some teachers may wish to
pre-prepare at least one tube to save time. Remove one boiling tube and pour the
sulphur into a filter paper cone held together by a paper clip (Fig. 1) and supported in
a beaker. Allow the sulphur to cool and solidify. Break the crust with a spatula and
pour off any remaining liquid. Needle-shaped crystals of monoclinic sulphur will be
seen. When cool, the cone can be passed around the class.
Take the second tube and, using a test-tube holder, heat it gently over a small
Bunsen flame, keeping the contents of the tube moving to prevent local overheating.
The liquid gets darker and, fairly suddenly, becomes a viscous gel-like substance.
This occurs at about 200 °C. The tube can be inverted and the sulphur will remain in
it. Show that the mobile liquid re-forms on cooling.
Further heating beyond the gel-like stage liquefies the sulphur again to a very dark
red-brown liquid (the colour of bromine). Note that during heating it is probable that
the sulphur will catch fire and sulphur dioxide will be produced. Have a heat-proof
mat to hand to place over the mouth of the tube to extinguish the blue flames.
When the sulphur begins to boil, pour the liquid sulphur in a slow stream into a
beaker of cold water. A mass of brown plastic sulphur will form. Allow this to cool
thoroughly, taking care because the inside of the plastic sulphur may remain molten
after the outside has solidified. Remove the plastic sulphur from the water and show
that it is rubbery – it can be stretched and will return to its original shape. After about half an hour it will be noticeable that the shiny surface of the plastic sulphur is
beginning to dull and some of the elasticity is lost as it begins to turn back to the
more stable rhombic sulphur. Leave until the following lesson to monitor the progress
of this change. This will be very noticeable after a week or so but complete change
will take a long time.
Working in a fume cupboard, put about 10 g of powdered roll sulphur into a
conical flask and add about 100 cm3
of toluene or xylene. Leave the sulphur to
dissolve. This will take some several minutes; warming to about 50 °C will speed up
dissolution. Some teachers may wish to prepare the solution before the
demonstration to save time. Pour a little of the solution into a petri dish, watch glass
or microscope slide and leave it in the fume cupboard for the solvent to evaporate.
This will take about 10 min. Small crystals of rhombic sulphur will form. These can
be viewed under a microscope. The class will need to file past and view them in turn.
It is worth the teacher selecting a well-formed crystal for viewing.
A projection microscope or video microscope can be used to show the shape of the
rhombic crystals to the whole class.
Some stages of this demonstration are time-consuming, eg melting the sulphur in the
oil bath, dissolving the sulphur in the toluene or xylene and evaporating the solvent.
Some teachers may prefer to melt some sulphur before the lesson and to prepare
rhombic crystals before the lesson to save time. In the latter case, slower evaporation
(which can be brought about by covering the petri dish with filter paper with a few
holes in) will produce larger crystals. Particularly large and/or well-formed crystals
could be retained as examples for future use.
Discussion of the demonstration, what it shows etc.
Using ptable.com, webelements or similar, investigate which properties show clear trends as you go across or down the table. Try to explain these trends.
Lesson 5: Trends in the Periodic Table
Covalent radius, ionisation energy, electronegativity and trends in groups and periods, related to atomic structure.
Investigating the sulfur further - viewing with a microscope… Why was the sulfur that I poured in the water yesterday rubbery? Why and how has it changed now? Why and how is it different from the (monoclinic) sulfur that went in the filter paper, and the stuff that’s now in the watch glass?
- Discussion of patterns in Periodic Table - what did they find for their homework? Can they explain it?
- Chemical Misconceptions - Interactions worksheet, perhaps?
- Forming tin(IV) iodide? Or possibly some other covalent compound with metal?
Trends in the Periodic Table
Lesson 6: Investigating the sodium tin
Revisiting the alkali metals, introducing galvanic corrosion
- Here is a very old container of sodium, from the back of the store-room. Surprisingly, it’s made of metal rather than glass; I guess that was acceptable at some point. As you can see, it has a worrying hole in the side, with some kind of white crystals around the outside. Any suggestions how it could have got there?
One obvious suggestion would be plain oxidation. But aluminium is very resistant to this, and surely they wouldn’t use any metal for this that wasn’t? How can we figure out what kind of metal it is, assuming these crystals are the oxide the metal?
Another possibility is that the hole has been caused by acid or alkali reacting with the metal - aluminium turns out to be amphoteric, meaning it could be either. Perhaps some of the sodium reacted with water, producing sodium hydroxide, and reacted with the can?
Yet another possibility is that the hole was caused (at least partly) by galvanic corrosion. How could we work out whether this is a possibility?
Let’s figure out what those crystals are, and whether we can rule out any of our hypotheses for what happened to the tin.
First, write definitions of any new vocabulary introduced in the first part of the lesson.
Write down the aim of the experiment, any hypotheses, and your proposed method.
Record results and any conclusion.
Flame test should tell us what metal (or metals) are in the crystals. We may also need to do a flame test on some of the corrosion from the elsewhere around the can, in case it’s different.
We may be able to rule out galvanic corrosion by showing that the sodium, rather than the aluminium, would be the anode if that was what we were looking at.
How could we definitively show that the aluminium was attacked by sodium hydroxide?
Safety precautions: Tin itself behind a safety screen, goggles and lab coats all round
Discuss conclusions and any ideas for further tests that might tell us more.
Bonding in Compounds
Lesson 7: Bonding in Compounds (1)
Polar covalent bonding, and the bonding continuum
Explain the trends in electronegativity across and down the periodic table. Write down your answer.
Introducing and defining these concepts:
- Polar covalent bonding
- Bonding continuum
- Permanent dipole-permanent dipole interactions
- Hydrogen bonding
- Covalent bonding between atoms off different electronegativity, where electrons are unevenly shared between them
- The range of different bonds between pure covalent and pure ionic
- Intermolecular attraction caused by the polarity of the molecules involved
- A particularly strong form of the above, owing to hydrogen having only one electron to start with
Testing some different liquids for polarity using a charged rod. Noticing their viscosity, too.
Can we compare ethanol with a hydrocarbon that’s liquid at room temperature, like heptane?
Discuss some differences between polar and similar but non-polar substances. Generally, intermolecular forces are much greater in a polar substance, thanks to the interactions between the permanent dipoles - what effects would we expect to see with stronger intermolecular forces?
- Higher melting point
- Higher boiling point
- Higher viscosity
Explain your observations as completely as possible.
Lesson 8 (am): Bonding in Compounds (2)
Properties relating to intermolecular forces, particularly surface tension
Here’s an ice cube floating on oil. Will it stay floating once it melts? If not why not? Why does it float in the first place?
Maybe also beads of wax in molten wax?
Be ready to show hexagonal magnet structure with big gaps in it.
Investigating surface tension - floating a pin on the surface of some water; sinking it with soap; similarly, a layer of powder on the surface of some water immediately goes to the sides on addition of detergent; also, surface tension is responsible for the formation of a meniscus. Watch how many coins you can put into an alread-full beaker before it overflows. c.f. http://www.nuffieldfoundation.org/practical-chemistry/detergents-soaps-and-surface-tension
Discussion - why is there surface tension?
Video of surface tension in space
Lesson 9 (pm):
Solubility of different things in polar or non-polar liquids
What do you think it takes for something to dissolve? In particular, non-polar molecules?
Polar dissolves polar, non-polar dissolves non-polar. Polar solutions are all about charges; non-polar solvation happens through London forces. Polar solvents don’t usually dissolve non-polar things because they’re too attracted to themselves.
Experiments on solubility of different substances - salt, sugar, congealed paraffin, that sort of thing.
Lesson 10 (double):
Review of the unit
Discussion of yesterday’s experiments. What did they find? Does it back up what they learned at all?
Go through the unit summary noting down anything that you’re not really sure of and any questions you want to ask.
What are people unsure of? Discuss, fill in, make notes on unclear points.
Try out specimen exam questions on this unit. Attempt them in the designated time, but with open books and feeling free to ask each other questions.
Discuss how everyone thinks they did.
End of Unit Test
Unit 2: Nature’s Chemistry
Alcohols, Carboxylic Acids and Esters
Introduction and Recap
- Skeleton formulae
How many chemical suffixes can you think of, focusing on organic chemistry? How many can you define?
Discuss homework, then collect it in.
Explain skeletal formulae. Show some examples.
Recap of carboxylic acids, alcohols and esters.
Why would an ester be more volatile/have a higher melting point than its components?
Read about making an ester in the textbook.
Note the naming convention generally used for esters
Answer Question 1 on page 66.
Lesson 12 (am)
Uses of esters
Also, all edible oils are technically esters. Pleasingly, they’re used as lubricants in some industrial processes where you don’t want mineral oil getting in to food etc.
Making (esterification, a condensation reaction) and breaking (hydrolysis)
Textbook questions if necessary
Lesson 13 (pm)
Preparation of esters
A few drops of concentrated sulfuric acid are added to a mix of an alcohol (e.g. octanol) and a carboxylic acid (e.g. ethanoic acid).
They are heated in a water bath for several minutes.
Then they are poured into a beaker of sodium carbonate.
Full details at http://www.nuffieldfoundation.org/node/3104
Safety goggles, lab coats.
Chemicals dispensed in small quantities and handled by pipette-dropper
Lesson 14 (double)
Extraction of caffeine
Fats, Oils and Soaps
Lesson 15 (Mon)
Fats and oils
Check over homework
Revisit structure of triglycerides
Copy structure of glycerol.
Talk about saturated and unsaturated fats
Make models of saturated and unsaturated hydrocarbons
Question 1 on page 75 - challenging
Lesson 16 (Tues)
Hydrolysis and saponification
Going through homework corrections
Hand out photocopy about systematic name, with numbers
Discuss hydrolysis, relate to other -lysis processes they’ll have met. Notes should be taken if they’re not 100% sure about this, along with condensation reactions.
Learn about saponification. Polar molecules again! Amphiphilic ones! Emulsions, surface tension! Mono- and di-glycerides of fatty acids!
How to make a paperclip ‘float’ on water. How to sink it using detergent.
Lesson 17 (Wed a.m.)
Pick a term from the board and define it. Make sure everyone’s got all the definitions down.
Quite a bit of book work to get through the summary questions.
Lesson 18 (Wed p.m.)
Saturation of triglycerides
Testing degree of saturation of various oils and fats using bromine water
Lesson 19 (Thurs double)
Interlude: Unit 1 revision
Lesson 20 (Mon)
Go through some past exam questions
Read through sheet about proteins.
Go through exam questions.
Lesson 21 (Wed a.m.)
Going through exam questions, plus a bit on proteins
Lesson 22 (Wed p.m.)
Longer exam questions
Lesson 23 (Thu double)
Any Unit 1 questions/corrections of practice paper for 25-30 minutes, then test in exam conditions.
Homework - finish protein worksheet
Lesson 24 (Monday)
Lesson 25 (Tue)
Ionisation energy and more polarity
Ionisation energy is:
The energy required to remove one mole of electrons from one mole of atoms (or molecules) in the gaseous state.
It follows these trends:
- Higher towards the right of the table (even up to the noble gases)
- Lower towards the bottom
A polar bond between two atoms occurs if their electronegativity is different enough for one atom to lose part of its charge to the other.
The partial charge on an atom is written δ+ (delta plus) and δ− (delta minus).
A chemical is polar if it includes polar bonds in an asymmetric configuration.
Refer to diagrams of solvation by water molecules.
Lesson 26 (Wed a.m.)
Ensure the following terms are defined:
- Amino acid
A compound with an amine group, a carboxyl group and a side-chain which is specific to the amino acid.
22 ‘standard’ amino acids
8 ‘essential’ amino acids (for humans)
500 or so others!
Can be an energy source
A polymer chain consisting of amino acids.
Between about 2 and 50 amino acids, with peptide bonds (aka amide links) between them.
Formed by a condensation reaction (like the formation of esters).
A large biological molecule consisting of one or more long peptide chains
A bit more discussion of proteins - their role in cooking as well as life.
Things to get in there:
- Proteins as glue for food
- Fluorescence microscopy
- Royal jelly
Looking at 3D models of some proteins (and/or dipeptides/oligopeptides?). Spotting the amide links, identifying some of the constituent amino acids.
Check out ‘In Search of More Solutions’ for activity on aspartame and this page for amino acid chromatography
‘Reality of Nutrition’ activity is probably way too in-depth?
A couple of answers from the textbook
Lesson 27 (Wed p.m.)
Questions on proteins
Finish that sheet, and answer Summary Questions from the textbook.
Revisit reaction rates, briefly - check that everyone worked through the Checklists for Revision when they were revising for the test.
Lesson 28 (Thu double)
More on proteins??
Homework: all study questions from Bonding in Compounds chapter (?) and pick a flavour/aroma compound and research it. Write at least 200 words on:
- What type of compound it is
- Whether it’s polar
- How volatile, and why (molecular weight, polarity)
- What foods it’s found in
Also, read up on fluorescence microscopy now that the Nobel Prize has been awarded for it!
The Chemistry of Cooking and Oxidation of Food
Lesson 29 (Wed a.m.)
Cut apple. Why might lemon juice keep it from oxidising?
Go through homework on flavour compounds. Everyone shares their findings with the class.
Demonstrate some aromas - eugenol, thymol, vanillin, capsaicin?
Can I get some kind of a perfume in for this?
Look at and identify compounds used in various processed foods.
Lesson 30 (Wed p.m.)
Lesson 31 (Thu double)
Denaturing practical (coffee & soy milk)
- Coffee sometimes curdles soy milk. What does that mean, chemically?
- What factors might make it more likely that the soy milk would curdle?
- What could we do to minimise the chance?
- Add milk first or second?
- Allow coffee to cool?
- Change pH of coffee?
- Heat it to increase acidity?
- Make it with an Aeropress to decrease it?
- Dilute/concentrate it?
- Warm the milk?
- Stir thoroughly?
- Form hypotheses about these things. Test them.
Research the use of vitamins C and E in processed food. Why are they often used together?
Lesson 32 (Mon a.m.)
Flash card exercises - revision of unit to date
Lesson 33 (Tue a.m.)
Revising naming conventions
Lesson 34 (Wed a.m.)
Recap - mainly study questions
- Oxidising agents used to detect reducing substances
Lesson 35 (Wed p.m.)
Lesson 36 (Thurs double)
Between Units/Research Project
10 minutes of extra time on test
10 minutes’ discussion of antioxidant projects
Unit 3: Chemistry in Society
Getting The Most From Reactants
- Raw Material
- Environmental Considerations
- Designing Processes
Calculations From Equations
- Moles and Molar Volume
- Balanced Equations and Mass
- Molar Concentration
Lots to practise here - relations between different units and concepts.
Students may need quite a few examples before they feel confident.
Percentage Yield and Atom Economy
- Percentage Yield
- Atom Economy
- Thing 3
- Dynamic Equilibrium
- Closed System
- Le Chatelier’s Principle
- Haber Process
- Hess’s Law
The enthalpy change of a chemical reaction depends only on the chemical nature and physical states of the reactants and products and is independent of any intermediate steps
- Enthalpy of Combustion
- Enthalpy of Neutralisation
- Enthalpy of Solution
- Molar Bond Enthalpy
- Displacement Reactions
- Electrochemical Series
- Redox Reactions
- Oxidising and Reducing Agents
- Spectator ions
- Gas Liquid Chromatography
- Mobile Phase
- Stationary Phase
- Retention Time
- Difference in Separation based on Size and Polarity
- Quantities of Unknown Reagents
- Redox Equation
- End Point
- Standard Solution
- Indicator (‘a chemical dye added to a titration to detect the end point’)