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This is the third unit on Materials, and corresponds to
the strand in the Science National Curriculum for England
(1999) appearing at KS 3 and 4 entitled: Sc3 Materials and
their properties: Patterns of behaviour.
Finding patterns (or rules, laws and generalisations) and
then trying to understand and explain them (by means of
models, theories and inventive ideas) is perhaps the key
activity of science. As such understandings about our world
and the universe within which it exists develop, so does
mankind’s ability to predict the outcomes of
actions and exercise a measure of control over the physical
and chemical (including biological) systems within which we
live and we are.
In developing such understandings we have to learn to ask
questions that can be answered by making observations or by
doing experiments. These processes also test our theories,
models and understandings and thus, to varying degrees
current understandings of scientists have to be tentative.
We believe that there are a number of basic ideas that are
fundamentally correct and it is these that form the basis of
science education. These include:
- The particulate theory of matter (The kinetic theory)
- The atomic theory (Atoms with a small heavy nucleus of
protons and neutrons surrounded by a cloud of
electrons.)
- Conservation of matter/energy.
- Laws of motion (and relativity.)
- Principles of gravitational attraction and of magnetic
and electrical attraction/repulsion.
- Laws of electricity.
- Theory (Law) of evolution.
- Role of DNA in inheritance and cell function
Of course none of these is fully developed or fully
understood - however they have in general stood the test of
huge amounts of experimentation and critical thought
starting more than 300 years or so ago. From a school teaching
perspective these are still not common sense ideas for
pupils (and teachers) and must be explored in the contexts
of meaningful and significant issues and applications at
personal, local and global levels. (It is not a sufficient
reason for learning that these things are on the National
Curriculum and that they may be examined in SATs, GCSE or
other examinations.) All learners must of necessity make
their own sense of science - the system of education,
examinations, teachers etc should help them do this and
maximise students’ opportunity to take their learning as
far as is useful and significant for them. (Unfortunately
for many pupils - especially in secondary schools - the
experience of science in schools becomes less meaningful and
personally significant as the examinations (SATs at y9 and
GCSE at y11) begin to take their toll.)
Often science does not provide clear answers to questions
that involve complex interactions in the real world (the
environment; conservation, genetics, weather, global
warming, AIDS) and rarely does it provide a moral
imperative. It may, however, predict what could be done,
what (probably) will happen and even allow for costing of
actions, but science doesn’t and cannot determine what
should or should not be done.
The purpose of this section is not to provide material
that is available in chemistry text-books but rather to
fuel debate and to try to begin to explicate some of
the key ideas that contribute to a meaningful
understanding of the ‘patterns’ of chemistry. For this
reason the level of the discussion may be found to be a
little higher than in other sections. It is unlikely that
the material will have direct application in the classroom
before KS4 and then only selectively - and it is unlikely
that many of the patterns would be apprehended by students
at KS 1 & 2. However, this approach can inform decisions
at to what experiences and explanations are likely to be
more useful ‘in the future’. None-the-less, the
qualitative approach taken here is considerably
oversimplified and will need to be considered critically
before being adopted or adapted.
- Temperature is a measure of the average kinetic
energy of the particles within a system. Measured on the ‘Absolute’
scale (in Kelvins) the temperature is proportional to the
average kinetic energy of the particles moving randomly within
the system. This means that:
- Particles (molecules/atoms/ions) of equal mass will have
the same average kinetic energy at the same temperature.
- If particles have different masses then, on average at a
particular temperature, the heavier particles will move less
quickly. (Kinetic Energy = mv2/2.)
- The actual energy of individual particles in a system
cannot be predicted - the distribution of energies is
described by the ‘Maxwell-Boltzmann Equations’ (q.v.)
but some particles will have energies much higher
than this average and more will have energies lower than
average. (in this distribution the mode is always lower
than the mean.) (This is addressed further in Download
2.2.1)
- In reactions that are slow, it is only these very fast
moving particles that are able to initiate chemical
change, since only they have sufficient energy to overcome
the ‘activation energy’ for the reaction (enabling
bonds to be broken)
- Systems tend to change so that the stored potential
energy within them becomes a minimum. This change can only
take place if there is an accessible mechanism and pathway
by which the change can take place. (Water and balls etc. tend
to run down-hill - although this can be prevented or
controlled by a suitable barrier.) This is a useful general
principle that applies at all levels:
- To electrons in energy levels around isolated atoms of
the various elements. (This is used to ‘explain’ the
electronic structure of atoms and the pattern of the
Periodic Table of the Elements).
- To electron arrangements - chemical bonding - when atoms
link together either as elements or as chemical compounds.
- To changes in arrangements of atoms when chemical
reactions take place. (It is important to remember that,
at the atomic/molecular level, the process of change is
often one of random collisions between the particles
that are present in the system. The energy of collisions
cannot be predicted, although except at very low
temperatures, some collisions sufficient to cause
change are always possible. In living systems there
are sophisticated levels of organisation, involving, for
example, enzymes, that allow for very precise control of
pathways for reactions that cannot be emulated in flasks
and test-tubes at room temperature.)
- As the temperature increases each of these potential
energy ‘wells’ will begin to take up their quantum
share of the random ‘thermal’ energy of the system.
(The lower the well the higher the temperature needs to be
to get a full share for the particles in the well.) Thus
at low temperatures bonds will not fully vibrate, though
free molecules may already be rotating. At the high
temperatures in plasma, outer electrons are not longer
able to remain in their energy level in atoms.
- There is an important distinction between energetics and
kinetics. The first - more usually termed thermodynamics (energetics
is part of this) deals with systems that are at equilibrium
(Section 2.2). Things affecting systems at equilibrium are
fairly well understood and changes can often be accurately
predicted and calculated. Kinetics tell us about the path
towards that equilibrium
- For an outline of issues relating the electronic structure
of atoms of the elements with the pattern of the Periodic
Table - See Download 2.1.1.
- For an outline of the issues relating electronic structure
of the elements with chemical reactivity - See Download
2.1.2.
- There is an extensive section in the CD Science Issues (Ross
et al 2005) dealing with the electron pair bond principle
behind the formation of bonds between any two atoms, leading
to an understanding of the structure and properties of metals,
giant covalent materials, volatile (molecular) materials and
ionic materials..
The principle of moving towards a state of minimum potential
energy leads to the idea of ‘chemical equilibrium’ once this
minimum potential energy has been reached. This is a dynamic
situation with rates of reaction between the energy states
available to the system being the same for forward and back
reactions - and thus there being no apparent change with
time. It is important to realise that - in principle - all
chemical reactions are reversible. For example, when hydrogen and
oxygen react together to form water vapour - there is always a
possibility that molecules of water vapour will collide together
with sufficient energy to form hydrogen and oxygen again. Although
this is rare at ordinary temperatures so amounts of the elements
remaining are hardly detectable, as the temperature rises
increasing amounts of hydrogen and oxygen are present at
equilibrium. The position of a chemical equilibrium is usually
affected by changing the temperature, and often by changing the
pressure or other things. (Catalysts only provide alternative
pathways for change and do not change the equilibrium.)
Equilibrium Confusion
A comparison and contrast between ‘equilibrium’ as it is
understood, differently, in Physics and Chemistry is given in the
table below. Alex Johnstone contends that it is because students
usually become familiar with the static physics’ concept first
that they (we?) have so much more trouble with the contrasting
dynamic chemical equilibrium. Full details at: (http://www.rsc.org/Membership/Networking/InterestGroups/ChemicalEducationResearch/Lectures.asp).
Most of the alternative conceptions regarding chemical
equilibrium can be traced to the physics concept of equilibrium,
which is generally learned much earlier. The following table
highlights major points of comparison.
| Physicists’ Equilibrium |
Chemists’ Equilibrium |
| a. Masses/Moments equal on both sides |
a. Numbers of moles need not be the same on
both sides |
| b. Addition to one side makes balance tip to
that side |
b. Adding more reactants tilts the balance
towards products |
| c. Balance has sides |
c. No sides |
| d. Can do something to one side only |
d. T and P changes affect both sides |
| e. Static |
e. Dynamic |
Download 2.2.1: provides more information on energy
profiles for a reaction and the affect of catalysts.
Enthalpy, Free energy, Entropy and equilibrium
Download 2.2.2: Energy (Enthalpy), Free energy and Entropy,
discusses the conflict between the principle of minimisation of
potential energy and the constant random movement of particles
that represents their temperature. A passionate paper on the
understanding of, and teaching of, entropy at the undergraduate
level in USA is given in Download 2.2.3
Download 2.2.4: provides brief access to a number of
dynamical systems in which the chemist’s concept of equilibrium
is applicable (including some that would be classified as examples
of physical, rather than chemical, change). Where possible the
patterns of behaviour between different chemicals are emphasised.
(All can be explored further using conventional chemistry texts.)
Examples include:
- Evaporation, saturated vapour pressure (SVP) and boiling.
- Melting points (= Freezing points).
- Solubility in water.
- Oxidation and reduction.
- Acids, bases and neutralisation.
- Stability of hydrates, carbonates and nitrates.
- Some industrial processes.
Reaction kinetics
“All around us are the slow reactions of life, waiting to
be examined and explored. Our aim is to share with pupils our
ideas about what makes a chemical reaction sometimes go fast,
and at other times go slow. We hope to develop a collision
theory model to explain why reactions go faster when we
increase the surface area (of a solid reactant), the
concentration (of a reactant in solution) or the temperature.”
(from Ross et al chapter 12)
For a discussion on the way to introduce Kinetics at GCSE
see the final section of download 3 - difficult ideas.
Kinetics for A level
Students, at KS4 and below will not normally be expected to
be familiar with rate equations for chemical reactions. The
basic ideas that affect rates of reaction should however be
relatively clear, as we have stated above. That for a
chemical reaction to occur it is necessary for the particles
of the reactant to come into contact and to collide with
sufficient energy for reaction to occur. This is shown for
both the forward and back reactions in Download 2.2.2.
Download 2.2.2 is a very important link between
Thermodynamics and Kinetics there is however a very seductive
link between quantitative aspects of Equilibrium Constants
and Rate Equations. Seductive because it is easy and
gives the right answer BUT wrong because rate equations cannot
be written down from the stoichiometric equation for the
reactions. This is explored in Download 2.3.
At KS1 & 2 there is little evidence of pattern relevant
to materials although some useful tentative
generalisations will probably emerge. These have been covered
in Sections for Materials 1 & 2 and some examples are:
- Metals are strong, dense (‘heavy’) conduct
electricity, are malleable and thin pieces can be bent,
and stay bent, without breaking.
- If substances exist as gases then liquids will form if
they are cooled sufficiently and if cooling continues they
will eventually form solids.
- Heat is given out when things burn. (Burning requires
fuel and oxygen)
At KS3 a more formal introduction to chemicals, particles,
elements, compounds, formulae and chemical reactions is
possible. In general, if links are made between chemical
change and energy, it will probably be expected that chemical
reactions will be accompanied by evolution of heat. Some
discomfort may be arising with the awareness of ‘reversible
reactions’ and perhaps some spontaneous changes that are
endothermic.
KS4 will begin the links between atomic structure, the
formation of chemical bonds and the exploration of ideas
regarding reactivity, bonding and bulk properties and uses of
chemicals. The link between energy, electronic structure of
elements and chemical bonds is often lost when explanations of
bonding focus upon the various ways of ‘getting eight
electrons in the outermost shell’. (Incidentally, metallic
bonding is often omitted too since the focus is on the
extremes of sharing electrons - covalent bonding and
transferring electrons - ionic bonding between the linked
atoms. (For this reason the electron pair bond principle,
which applies to all bonding, may be a more helpful system -
see para 2.1 above)
Most of the ideas explored in this unit are directly
applicable only at KS5 (although fairly limited even here) and
at undergraduate level. However, if they are not explored the
teacher may feel vulnerable even when working with the ‘lower’
levels.
Download 3 Difficult ideas in chemistry contains a
useful review of some of the difficult ideas about patterns of
behaviour in chemistry. (It repeats Matter K3.2 Change -
Download 4.3)
- Atkins P (1995) ‘The Periodic Kingdom: a journey
into the land of the chemical elements.’ London,
Phoenix (Orion Books)
- Atkins P & Jones L (1999) ‘Chemical Principles:
the quest for insight.’ New York, W H Freeman
- Greenwood N N and Earnshaw A (1984) Chemistry of the
Elements) Oxford, Pergamon Press.
- Ross, Lakin and Callaghan (2004) ‘Teaching
Secondary Science’ Second Edition London: David
Fulton
Websites:
Downloads in this Unit:
Section Developed by:
Alan Goodwin, MMU
(with additional material from Keith Ross, University of Gloucestershire)
Oct 2006
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