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This really is the core of chemistry - and developing
understandings of chemical ideas depends upon the experience
of large numbers of examples in everyday life, not just in
science laboratories. Many elementary science courses
include, at an early stage, differences between chemical and
physical changes - this may be useful but we only
gather information and evidence ourselves about
changes to materials over a long period of time and it is
often not possible to decide definitely whether a particular
change can be categorised as chemical or physical (Many
everyday changes that take place during cooking, living,
dying or making tea are very complicated and often consist
of lots of different changes, some of which are chemical and
some physical!)
Even the criteria that are often taught to distinguish
between the two sorts of changes are problematic. The
following table is taken from my own science notes of almost
50 years ago and I certainly tried to use the ideas when I
began teaching chemistry in secondary school:
| |
Chemical |
Physical |
| 1 |
New substance(s) formed |
No new substance(s) formed |
| 2 |
Not easily reversed |
Usually easily reversed |
| 3 |
Often accompanied by a large energy
change |
Energy change usually small |
I could find examples that seemed to fit this
categorisation - a candle burning, an egg being cooked (in
boiling water) or a piece of iron going rusty seemed fairly
good examples of chemical changes. All seem to form new
substances - except the candle seems to disappear and
it is some time before convincing evidence can be provided
that water and carbon dioxide are the combustion products.
None of the processes is easily reversed although the energy
change for the egg and the iron are not obvious. Changes of
physical state are usually cited as examples of physical
change - e.g. water freezing or evaporating; candle wax or
chocolate melting - it is however, not obvious that
no new substance is formed. Children usually have to take on
trust that ice, water and water vapour are the same
substance (Actually I am no longer sure - since if one
considers that ‘hydrogen bonds’ are actual chemical
bonds then it really is probably more correct to suggest
that the three are chemically distinct since the bonding
linking the molecules is different in each case??!) Breaking
a drinking glass and powdering coal are indeed physical
changes (no new substance is formed - although some chemical
bonds must be broken during the process.) but in these cases
the processes are not easily reversed.
The only real criterion to distinguish between
chemical and physical changes is whether new substances are
formed but unfortunately it is not always clear whether this
is or is not the case. A few examples are discussed in Download
1 - these are useful for fairly experienced scientists
but should not be discussed in this way with beginners.
Until we are sure whether new substances are formed it is
probably better to be agnostic - and even when we think we
know a change is physical there may be evidence to change
our minds later! This makes a good discussion topic for your
trainee teachers.
Particle theory and chemical change:
- Only a limited number of types of atoms exist
(One type for each of the elements - just over 100) and
most of these can be combined in various ways to form the
huge range of chemical substances that can exist. (Most of
these contain only a small number of different kinds of
atom; at least 1 and rarely more than ten kinds.) The
forces that hold atoms together are called ‘chemical
bonds’.
- These atoms (Dalton’s ‘invention’ in
1810) are indestructible and cannot be created, destroyed
or inter-converted except in ‘nuclear’ reactions such
as those that take place inside stars, or during ‘supernovae’
explosions, in nuclear reactors or radioactive decay
processes.
- When a chemical change takes place the atoms
in the reacting substance(s) are rearranged to form the
product substances. And if, as is usually the case, the
bonds in the products are stronger or weaker than those in
the reactants, energy is also given out or absorbed during
the reaction. (The stronger, more stable a bond is the
lower the potential energy it has so that more energy has
to be provided to break a strong bond than a weak one.
Conversely more energy is given out when a stronger
bond is formed, as the atoms ‘rush’ together
under the strong electrostatic forces. Many pupils have
the misconception that energy is released as bonds break)
These are not trivial although they sometimes appear so
to more experienced chemists/scientists. Some of the
problems are explored below (and the work of Alex Johnstone
and Keith Taber [see references in section 5] are useful
sources of practical insight.)
In particular the following are pertinent - although some
may be controversial. (Please consider these critically. The
‘story’ of chemical change, chemical reactions and
materials is a developing one, not only from the perspective
of pupils, but also for the science teacher and scientists
themselves.):
- Children - and others - often think of atoms as ‘bits’
of ordinary ‘stuff’ that can melt or burn and have
colour (after all we all know that carbon atoms are
black and oxygen atoms are red!). These ideas are
clearly meaningless for separate atoms and apply only
to bulk matter. It is thus very important to help
children distinguish between bulk matter (of everyday
experience) and the world of these indestructible
atoms. See Download 2a the particle concept.
- Atoms are only rearranged during chemical reactions
- they can never be created nor destroyed. Atoms are
very small, and the very smallest part of an element
that can exist. (The electrons, protons and neutrons
that make up all atoms are of course much smaller but
they cannot be distinguished. All electrons are the
same whatever atom they are in.) Atoms of all elements
remain essentially intact during all chemical
reactions - it is only the outermost electrons that
are affected during chemical reactions. When atoms
approach each other it is these outer electrons that
determine whether the atoms will repel each other or
bond.
- Pupils get confused by the numbers we put in
chemical formulae. This has led to the use of word
equations at KS3, which only makes things worse -
for there is little logic in word equations. What is
needed is for the atoms to be drawn out or modelled.
Only when pupils can see that the number of atoms is
unchanged during a reaction, by actually counting
them, should we offer them formulae as a short cut to
drawing ‘blobs’ (of different colour) to represent
the atoms in the substances that react. See Download 2b Molecular Models and equations
- The chemical symbols we use to represent the
relative numbers of atoms involved in making up a pure
substance and that are rearranged during chemical
reactions need to become ‘commonplace’ and
meaningful to pupils if they are to be useful. The use
of symbols and formulae is implicit in the KS3
National Strategy in the Framework for teaching
science - ‘particles’ (DfES 2002), but it seems
unusual for sufficient emphasis to be placed on
learning and using the basic formulae for many pupils
to become fluent in the new language. However, rather
than giving pupils basic rules about the ways in which
atoms join together we could begin to show them that
the combining power of an atom (valency) is linked to
its position in the periodic table and can be
explained in terms of electrons and atomic structure.
This is part of the KS4 curriculum.
- It is remarkably difficult to relate, say, the white
crystals of substance (common salt = sodium chloride)
that we can see, with the conceptual picture of a huge
three dimensional array of sub-microscopic charged
atoms (ions) of sodium and chlorine in each crystal.
These we can infer (from evidence provided over a long
period of time) but cannot see (even through a
microscope) and at the same time we have to connect
these with the symbolic representation ‘NaCl’ (or, more accurately
'Na+Cl-'). Such interrelations make almost
impossible cognitive demands on most of us until such
time as these simple formulae and equations become
unproblematic and begin to ‘mean’ real substances
made up of the appropriate relative numbers of atoms.
(Johnstone, 1999)
- Quantitative aspects of chemistry become
accessible once we have chemical equations. Because
this involves numbers, relative masses and (albeit
simple) proportionality this is almost always
perceived as very difficult by pupils and is avoided
for all but the most able. There are a number of key
ideas that need to be grasped and mastered, but the
mathematics is not beyond an average student at KS2!
- The balanced equation shows exactly the
numbers of atoms of each element involved in the
reaction.
- The relative masses of all substances involved
can simply be shown by calculating the relative formula
mass of each substance and multiplying each by
the appropriate number from the equation
(that’s the big one in front of the formula
that tells how many moles there are of that
substance).
- If any of the substances is a gas under the
conditions of the experiment then it is useful
to remember that at room temperature and
pressure the volume of 1 mole of ANY GAS is
about 24 litres.
- The representation of elements and the formulae of
some compounds can cause concern initially (and
complications later). For example, the gaseous or
volatile elements exist in the form of stable
molecules that consist of pairs of atoms. It is these
molecules that take part in chemical reactions and
that we represent in chemical equations. These include
Hydrogen (H2); Oxygen (O2); Nitrogen
(N2); Fluorine (F2); Chlorine
(Cl2); Bromine (Br2) and Iodine
(I2).
When the structure of elements becomes more
complicated than this we usually ignore the bonding in
the element and just use the symbol as its formula,
thus implying to our students that the atoms are NOT
bonded together. E.g. Fe; Zn; C. (Download 2c
elements in chemical equations explores this a little
further.)
- Energy is usually given out during chemical changes
that we commonly experience in the laboratory. This
must be associated with the rearrangement of atoms. It
always requires energy to break the bonds between
atoms - and a rearrangement is not possible unless
some bonds are first broken. If the new bonds that are
formed are stronger than those in the reactants then
these bonds have ‘lower energy’ [see Note *] and some energy is
given out - usually as heat so the ‘particles’
move about faster and the temperature rises. Such
reactions are called ‘exothermic’. ‘Endothermic’
chemical reactions are also possible - but less easily
demonstrated - a very important example is
photosynthesis, the
reaction between carbon dioxide and water that takes
place in green plants where the energy is provided by sunlight.
This energy input is needed because the bond in
molecular oxygen (O2) is weak, so less
energy is released as the oxygen is formed than was
used to break apart the water and carbon dioxide.
Glucose, the other product of photosynthesis, is a
relative stable molecule, though, by convention, text
books and food packets suggest it ‘contains’
energy. It is the dangerous and reactive gas oxygen
that actually contains the weak bonds, and which ‘carries’
the energy potential that is available through
respiration. This conception is reviewed in download
2.4 in the Energy unit of Subject Knowledge.
[Note * Here lies the source of the misconception that bonds
'contain' energy. A 'lower energy bond' is a strong bond, one that
requires a large energy input to break it. The lower its energy
level, the stronger it is and the harder it is to break.
Conversely a 'high energy bond' such as the third phosphate in ATP
is actually a weak bond, somewhat 'unstable', and reactive,
because it doesn't require much energy to break, but as the new
bonds form, more energy is released than was used to break the
original bond, causing the original unstable ATP to be labelled,
somewhat confusingly, as a 'high energy' molecule]
At this point it is convenient to link KEY IDEAS of science
necessary for sustainability and environmental education. These
have been argued to be the following:
- Matter cannot disappear - atoms are indestructible
- Although energy is indestructible (First Law of
Thermodynamics) it spreads and becomes ‘degraded’ (Second
Law of Thermodynamics)
- Matter also scatters unless some energy source is available
to ‘tidy it up’ and keep it going in cycles (Second Law of
Thermodynamics again)
- Thus, on this planet of ours, energy is constantly being
degraded to enable matter to be re-cycled. This happens
naturally in the great natural cycles: rocks, climate, and
life*, but humans have recently failed to manage their
resources in a similar way
- There is Value in Structure - organised matter*, such as the
human brain, represents high quality
* Plants (strictly speaking we should say 'producers': algae,
chemosynthetic bacteria etc. as well as green plants) begin this process of producing Structure and Order by
using energy from the Sun, (allowing carbon dioxide and water to
form glucose and oxygen). Subsequently, through energy provided by
respiration (where oxygen re-joins with the glucose), Life on
Earth keeps this high state of order, represented by the global
ecosystem.
All of these are touched upon in this unit. Useful material of
this perspective can be gained from Download 3 Sustainability
Links which relates to ‘forum-for-the-future’ and the
international charitable organisation ‘The Natural Step’ (TNS).
It may thus be useful to keep in mind that
- the indestructibility of matter (and the limited supply of
atoms on Earth),
- the dissipation of energy (Finite fossil fuel resources and
problems of global warming),
- structure and photosynthesis
are basic requirements for discussion of the balance and
continuation of (human?) life on Earth, and should therefore form
an important part of a curriculum for scientific literacy in
school.
All these local and global environmental issues: pollution,
recycling, sustainable energy, food production, water supply,
transport etc. cannot be solved by science alone nor can science
answer questions as to what should be done - but hopefully
an understanding of the science behind these issues will
increasingly inform the human view of the problems and
possibilities and will enable us to maintain a hope for the
future.
N.B. As From Sept 2005 Sustainability Education is a required
constituent of ITT.
There are a limited number of reaction types and all involve
only the rearrangement of atoms and the associated energy changes. The various
classifications of reactions are covered in almost all school
science texts relating to chemistry - an interesting text is that
by Ryan (1996). See also Ross et al (2004) - especially Chapter
12. “Difficult ideas in chemistry” a summary of which is in
download 4c.
Some classifications that will be met:
- Combination and decomposition
- Double decomposition (Changing partners) - this includes
precipitation reactions, acid base (often treated separately
as ‘neutralisation.) and (probably) acid carbonate.
- Displacement
- Oxidation / Reduction
- Exothermic endothermic
- Catalysed reactions - especially important are reactions in
life (biochemistry) where reactions are controlled by enzymes.
Burning
A common example of a chemical reaction is that of a fuel
burning. Indeed, fire is a phenomenon that has interested human
beings for many thousands of years. It is, however, only since the
1770s, based upon the work of Priestly, Scheele and Lavoisier that
we have come to believe (know) that fire is the visible region
where energy is being given out (as light and heat) when ‘fuels’
react with oxygen from the air.
There is a most important parallel to be drawn between burning fuels
and aerobic respiration in living things on Earth. Animals, plants
and almost all living things use enzymes to break large molecules of fuel (such
as starch) into simple sugars. The small sugar molecules are now able to react
with oxygen within the living cells. This reaction with oxygen, called aerobic
respiration, releases energy, which is needed to drive endothermic
processes such as growth (e.g. protein synthesis) and movement.
The expression 'burning up calories' is a reference to this joining of oxygen
and the fuel in our bodies. Of course it is not the calories that are burnt, but
the fuel. The result is we breathe out carbon dioxide and water, and the energy
transferred is available for our bodies to 'use'. See the biology units for
further discussion on living things. Further perspectives on Fire and fire safety are given in Download
4a Fire Safety (Teachers notes) and some supporting
Power-point slides are in Download 4b Fire!
Other reactions
Download 4c Difficult ideas in chemistry summarises some
of the problems pupils might encounter as they try to come to
grips with some of the harder parts of chemistry to do with
reaction kinetics and energetics. It is taken from chapter 12 of
Ross et al (2004). It is discussed more fully in Materials unit
3 Patterns of Behaviour
Through their time in school children will
hopefully come to an understanding of their global environment. It
is an understanding based on the idea that matter (at least at an
atomic level) cannot be created or destroyed, therefore materials
are constantly being recycled in all processes on Earth. We see
this in the simplicity of the water cycle through to the most
complex cycling of materials by the processes of life itself.
At primary level these ideas are best
approached by getting children to think about where materials come
from, for example in
- a cup of milk (grass - cow - dairy - shops/milkman)
- a wooden table (carbon dioxide and water - growing tree - saw mill -
factory/carpenter - shops)
- a china plate (granite - erosion - river sediment - clay - pottery
works - shops)
- a metal spoon (metal ore - mining/quarrying - metal extraction (eg
iron and steel works) - factory/smith - shops)
This process of tracing things back
shows that materials don't come from nowhere - in the same way
they don't disappear when they are eaten or thrown away. For a
full discussion of the progression from reception to KS4 see
download 5:
- DfES (2002) ‘Key Stage 3 National Strategy - framework for
teaching science’
- Johnstone A (1999) The paper is available on the internet
at: <http://www.rsc.org/Membership/Networking/InterestGroups/ChemicalEducationResearch/Lectures.asp>
Accessed March 2006. This is Download 1a of the Materials 1:
classification Unit. (NOT this unit)
- McDuell B (2000) “Teaching Secondary Chemistry.” London,
John Murray.
- Ross K, Lakin L & Callaghan P (2004)“Teaching
Secondary Science - Constructing Meaning and Developing
Understanding.” Second edition London, David Fulton
Publishers.
- Ross, K.A., Lakin, L., Littledyke, M. and Burch, G (2005)
"The Science of Environmental Issues" CD-rom.
Cheltenham: University of Gloucestershire (available from: www.glos.ac.uk/science-issues)
- Ryan L (1996) ‘Chemistry for you.’ Stanley
Thornes Ltd., Cheltenham.
- Taber K (2002) ‘Chemical misconceptions - prevention,
diagnosis and cure. Volume 1: Theoretical Background. Volume
2: Classroom Resources. http://www.uoi.gr/cerp/2001_February/07.html
(full text available as PDF)
Downloads in this Unit:
Section Developed by:
Alan Goodwin, MMU
and Keith Ross, University of Gloucestershire
May 2006
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