K3.3 Patterns of Behaviour
This unit extends the ideas covered in the first two ‘materials’ units and is specifically aimed for teachers engaged with teaching chemistry at KS5 (A-level) although there is relevance at KS4 too (as well as for teachers’ personal knowledge development.) The introduction touches upon the natures of science and then focuses upon the patterns of ideas involved in modeling atomic structure and the Periodic Table of the elements and chemical bonding. This is followed by detailed consideration of chemical reactivity from both thermodynamic and kinetic perspectives. An outline of the progression of pupils’ ideas is given as well as some useful references.
This is one of 17 articles whose main aim is to support the processes of teaching/learning between the science education tutor and the trainee science teachers with a focus on “teachers’ knowledge and understanding”. During a primary or secondary BEd, PGCE or GTP we hope that those learning to become science teachers will be able to challenge their own understanding of science and scientific concepts. Unit K0 specifically explores general issues relating to all the knowledge units - to the learning of science.
Standards: This unit specifically addresses Q14 but, appropriately used can contribute to and provide evidence of competence for many others of the standards especially Q4,6,7,8,18, 22 and 25.
Keywords: Secondary, Teacher knowledge, Chemistry, A-level, Thermodynamics, Kinetics.
2.0 Key ideas to understanding patterns of behaviour in chemistry
2.1 Atomic structure and the Periodic Table
3.0 Progression in children's ideas about patterns of behaviour in chemistry
4.0 Useful references
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:
- All particles (molecules/atoms/ions) whether as aprt of a solid, liquid or gas structure, 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 and this is often less well understood.
- For an outline of issues relating the electronic structure of atoms of the elements with the pattern of the Periodic Table - See Download k3.3_2.1a.
- For an outline of the issues relating electronic structure of the elements with chemical reactivity - See Download k3.3_2.1b.
- 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.)
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 k3.3_2.1a: provides more information on energy profiles for a reaction and the affect of catalysts.
Enthalpy, Free energy, Entropy and equilibrium
Download k3.3_2.1b: 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 k3.3_2.2d: 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.
Download K3.3_2.2a 'Energy profiles for a reaction - the affect of catalysts'
Download K3.3_2.2b 'Energy (Enthalpy), Free energy and Entropy'
Download K3.3_2.2c 'Entropy by Frank Lambert'
Download K3.3_2.2d 'Dynamic Equilibria'
“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 k3.3_3.0a - 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 k3.3_2.2b.
Download k3.3_2.2b 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 k3.3_2.3a.
Download K3.3_2.3a 'Link between Chemical Equilibria and Rate equations'
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 k3.3_3.0a 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 k3.2_4.0c)
Download K3.3_3.0a 'Some difficult ideas in chemistry'
- 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 (1997) Chemistry of the Elements (2nd Ed) Oxford, Pergamon Press.
- Ross, Lakin and Callaghan (2004) ‘Teaching Secondary Science’ Second Edition London: David Fulton
- A very useful Chemistry resource developed recently by the Royal Society of Chemistry covers all stages of compulsory schooling and provides a wide variety of materials for use by teachers is: http://www.chemistryteachers.org/ (Accessed 23/11/09)
Section Developed by:
Alan Goodwin, MMU(with additional material from Keith Ross, University of Gloucestershire)Oct 2006
Published: 10 May 2007, Last Updated: 13 Sep 2008