Science Additional Specialism Programme

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Comments Linked to Criteria

Understanding of key ideas:

You show good knowledge and understanding of the key ideas, and have been able to set you work in a wider theoretical context. In order to improve further you would need to set your work in a much wider theoretical context.

Use of material:

Links between theory and practice have been clearly and effectively made. Your answer demonstrates a good understanding of the most relevant literature.In order to improve further you would needed to have very well developed critical arguments.

Range and comprehension of sources:

You have demonstrated evidence of an intelligent choice of reading.In order to improve further you would need to demonstrate initiative in the choice of primary and secondary sources.

Communication:

The structure of the answer is coherent, with a well constructed and clear central argument.In order to develop further the structure of your argument would need to be organised systematically around a central argument.

Level of citation:

The references and citations are in full and correct form.

General Comments:

An interesting piece of work. More developed critical arguments would have strengthened this assignment further.

First Marker: Second Marker: Date: 21 September 2010Mark: 65Classification: II(i)

Word count: 8396

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Science Additional Specialism Programme (SASP)

Chemistry

2009-2010Pg 2-9 Section A - Reflection and evaluation of the planning and delivery of a sequence of

Post 16 chemistry lessons.Pg 10-16 Section B – Critical reflection of how the learning process undertaken during the

course has impacted on teaching practice.Appendix:

Pg 17-18 Lesson planPg 19 Questionnaire analysisPg 20-25 Lesson activities – starter activityPg 26-27 Lesson activities – Electronic structure of the first 36 elementsPg 28 Lesson activities – Electronic configurationPg 29 Lesson activities – Ionisation energiesPg 30-31 Lesson activities – Ionisation energies – name the elementPg 32-37 Lesson activities – Practice examination questions for electronic structurePg 38-39 Lesson activities – Electronic structure homeworkPg 40-41 Lesson activities – Electronic structure testPg 42-63 Lesson activities – Electronic structure notesPg 64-74 Lesson activities – Ionisation energies notes

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Reflection and evaluation of the planning and delivery of a sequence of Post 16 chemistry

lessons.Year group 12Number of students 7Ability of students Mixed (Grade B - E)Exam board AQAChemistry topic 3.1.1 ‘Atomic structure’Background

The students I arranged to teach had previously being taught about the basic structure of the atom, about mass number and isotopes and about the mass spectrometer; therefore it was up to me to cover the remainder of the examination syllabus on atomic structure.Choice of topic

The reason why I chose ‘atomic structure’ with particular reference to ‘electronic configuration’ was because of the fact that I did not feel confident in my knowledge and hence teaching of this topic. I remembered ‘learning’ about the atom while studying for my A-levels 15 years ago, but not really understanding this abstract concept. During one of the sessions on the SASP (Science Additional Specialism Programme) course I began to fully appreciate the abstract concept of atomic structure and felt inspired to go away and read more on the subject matter. Now ‘clued up’, I felt that I needed to ‘try out’ my new found knowledge and understanding on a group of students. There has also been much criticism and many unanswered questions regarding how best to teach atomic structure. (Bent, 1984; Berry, 1986; Gillesphie, 1991; Hawkes, 1992; Shiland, 1995 and Tsaparlis, 1997a, 2002). Tsaparlis (1997a, p922) sums up the importance of this topic. According to him,

“It is no surprise that the atomic theory has fascinated chemistry teachers and constitutes the cornerstone of modern

chemistry curriculums even at the primary school level. As a matter of fact, the study of atomic and molecular structure – from

the elementary models to the old quantum theory and later quantum mechanical concepts – is considered a sine qua non in

chemical education.”

Lesson sequence

Having observed a series of A-level chemistry lessons in the high school, I decided to follow their successful structure in that I would begin with a starter activity which would recap some of the GCSE syllabus. The main part of the lesson would include an introductory discussion on the topic which would be led by me followed by various tasks which the students could attempt in an order of their choice and complete during their study periods. The plenary would involve a run though of the objectives to check that the success criteria had been met. To confirm that my teaching had been successful I decided that a brief test on the subject would be the starter for the subsequent lesson.Starter:

As a starter activity I decided to test the student’s prior knowledge of atomic structure from the GCSE syllabus. For this activity, I collated together some questions from the previous year’s examinations. The logic behind this came from the recent reading of some literature. Taber (2003) mentions that the models of the atom used in teaching have consequences for students. Scientists have not agreed on one specific model to use in their work; schools and colleges do not teach with one specific model; and in the minds of learners there is no one model of the atom. The GCSE recommended model of the atom is based upon the Bohr model – it has a nucleus with electron shells. It was because of this background reading that I became aware of the possible mind field that lay ahead of me. I needed to question the students somehow on what they had been taught and what conceptions or misconceptions they had before ploughing ahead with the A-level requirements as learning about the structure of the atom is difficult for many students. Cervellati and Perugin (1981), Taber (2000a,b) and Harrison and Treagust (2000) write about the fact that difficulty is often experienced by students when more sophisticated models of the atom are introduced post-16.Use of models

Hodgson (1995) mentions that hardly a lesson goes by without the use of one or more analogical models to explain the content. He talks about models providing the means for exploring, describing and explaining scientific and mathematical ideas and declares that they help make science relevant and interesting. According to Harrison and Treagust (2000) teachers should teach modelling skills and should encourage students to use multiple rather than isolated analogical models, and take the time to discuss and critique the models used in class. They state that highly abstract non-observable phenomena are explained using multiple models and that analogical models stimulates students’ curiosity and imagination and enhances their creative thinking. Grosslight et al., (1991) highlights an important fact that I needed to consider before ploughing ahead and using multiple models to explain electronic configuration. They point out that while the contradictory nature of some multiple models may be understood by the teacher and textbook writer, inexperienced modellers are unlikely to share this view. They mention

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that teachers should remain aware of the difference between expert modellers (themselves) and naïve modellers (their students). Harrison and Treagust (1996) make the point that atomic theory depends more than any other topic in chemistry on a variety of models to explain particulate behaviour. Justi and Gilbert (2000) have described a sequence of historical models of the atom starting with the ancient Greeks then leading onto Dalton, Thomson, Rutherford, Bohr and finally onto quantum-mechanical. Justi and Gilbert talk about these models having currency at some time in the development of atomic theory and mention that the modern chemical models (i.e. from Dalton on) played a role in research without necessarily being viewed as the ultimate model. It has been recognised that when teaching abstract and unfamiliar ideas, it is important for the teacher to find some way to ‘anchor’ the new knowledge among existing ideas using metaphor or analogies. Here lay a major stumbling block for myself; while attending the SASP course, on numerous occasions, reference to the ‘Royal Society of Chemistry‘ for resources to be used within school was mentioned. It was here that I originally looked for ideas on how teach electronic configuration and it was here that I came across a resource that asks students to compare the model of the atom with the solar system. Although at first I found it difficult to spot similarities and difference between the atom and the solar system, when reading through the answers, I could see the benefit of such an activity if the students had enough background knowledge. Since the students I would be teaching were not my regular students, I was not fully aware as to what they had previously been taught, therefore this could be an issue to which I would have to consider, Upon further background reading, I discovered that this activity may not have been as beneficial as I first thought. Fischler and Lichfeldt (1992) argue that initially teaching a model of the atom based on concentric electron orbits may be counter-productive. They found that the planetary orbit preconception seems to be especially resistant to change, Taber (2004) mentions how prior teaching about the structure of atoms seems to interfere with new learning. He found that students who are taught a model of the atom where electrons exhibit planetary orbits, in concentric shells, around nuclei, who then go on to study science at college or university level, then, the atomic model learnt at school acts as a ‘pedagogic learning impediment’. He found that when students are told about orbitals, they may interpret the new concept in terms of the familiar shells model. (More about orbital’s later).Gilbert (1998) suggests that a valid teaching model must be designed to lead towards the current consensus model(s) used in the field. The teaching model should encourage learners thinking towards a more sophisticated level of scientific understanding. Harrison and Treagust (1996) talk about their being significant problems in teaching and learning with analogical models in school science. They mention that teachers cannot predict how students will interpret the analogical models used by themselves, in textbooks, on videos, or in computer simulations. They point out that while the motivational benefit of models guarantees engagement, it exposes students to private or group interpretations that often lead to unexpected alternative conceptions.Misconceptions

It has been over 10 years since I have taught A-level, this, and the fact that chemistry is not my specialist subject led me into researching misconceptions that students may have on atomic structure. To overcome misconceptions is a difficult task. Even when students come close to realising their errors, they revert very easily to their previous ideas with which they are more comfortable. Difficulties and misconceptions in student understanding of atomic structure have been an important issue in the field of science education. Numerous studies have highlighted learning difficulties and misconceptions relating to atomic structure. Taber (2003) mentions that students often believe that the nucleus of an atom must have equal numbers of protons and neutrons as the neutrons have the role of neutralising the protons. ‘Neutr’ of neutron refers to a kind of neutralisation process rather than a ‘neutral’ charge. The starter activity I had planned would therefore raise this issue and would allow me to correct any such misconception. Taber (2003) also refers to some students believing that orbiting electrons push on the protons in the nucleus to hold the nucleus together. Harrison and Treagust (1996) found that language that is common to both biology and chemistry (e.g. nucleus and shells) was a major source of confusion for some students. For example, some students confused the nucleus of the atom with the nucleus of a cell and believed the nucleus to be the control centre of the atom. The biological influence was also evident when discussing shells, as some students saw shells as acting as a form of protection and, when asked for examples, listed items like sea shells, snail shells, clam shells etc. Harrison and Treagust (2000) mention that students prefer to think about abstract processes and concepts in concrete terms. Examples of this problem in school chemistry are the depiction of electrons as solid and static instead of diffuse and dynamic. Taber (2002c) mentions that post-16 students can harbour the alternative conceptions of: the nucleus is not attracted by the electrons; the nucleus attracts an electron more than the electron attracts the nucleus; the protons in the nucleus attract one electron each; and the electrons repel the nucleus. Orbitals?

Post-16 level students are often expected to move beyond ideas about electron shells and to learn something about electronic orbitals. Tsaparlis (1997b) states that serious consideration must be given to alternative ways of teaching chemistry that avoid orbitals. Gillesphie (1996) agrees with Tsaparlis and feels that orbital ideas are unhelpful prior to university level and that ‘electron pair domains’ are

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simpler and sufficient for school level study. However, examination stipulations may require learners to tackle orbital concepts. As I would be teaching students who would be sitting an AQA examination and the syllabus clearly mentions orbitals, I knew I would have to undertake this concept.Taber (2002c) affirms that abstract ideas such as orbitals are demanding even for able students, and time, reinforcement and practice are needed if learners are to show (and maintain) a good understanding of the orbital topic. Time, reinforcement and practice; 3 very important words I seriously needed to consider when planning the teaching fro this concept and the learning opportunities for the students. Taber (2004) refers to research that suggests that students adopt the term ‘orbital’ readily, but often tend to use it to re-label their existing conceptions of electron shells. Students will refer to orbitals as “round” or the “path the electrons takes” as it “circles the nucleus” and diagrams which did represent electrons in shells were said by students to show the orbitals. Taber points out that prior teaching about the structure of atoms seems to interfere with new learning particularly if students had previously been taught a model of the atom based upon the ‘solar system’ with electrons in concentric shells around the nuclei. These students, when told about orbitals, may interpret the new concept in terms of the familiar shells model. Summary of background reading:

One of skills of a good teacher is in finding ways to make complex ideas seem accessible, but done in a way that is scientifically valid, and provides a suitable platform for further learning. So, I needed to simplify for the students sufficiently, but not oversimplify to undermine their future needs. Taber (2004) declares that over simplifications have the potential to act as significant impediments to further learning. Taber (2003) points out that each teacher will interpret schemes of work and syllabuses in the light of their own scientific subject knowledge, professional pedagogic knowledge, and their personal evaluation of the readiness of students to take on new ideas. When I first looked at the A-level syllabus it seemed very daunting. It had been many years since I had taught at such an advanced level, and the fact that I was embarking on this in a subject that was not my natural specialism was even more overwhelming, From the SASP course I had gained enough background knowledge and understanding to have the courage to commit myself to further reading on the subject. Once I had the subject knowledge ‘cracked’, I then began researching how to deliver effective lessons to the students. My reading on the subject on the use of models, lead me down the path of using multiple models. When teaching using these models, I needed to ensure that I gave the students plenty of time to discuss and critique the models. Since the students I would be teaching were not my regular students, I needed to consider how they had previously been taught about the structure of the atom. During the starter activity I planned on circulating and discussing with the students what conceptions and misconceptions they currently had on the topic. If the students mentioned the structure of the atom being similar to that of the solar system, then instead of telling them that this analogy is not as beneficial as it first seems, I planned on spending some time looking at the similarities and differences of comparing the atom with the solar system so that the students themselves could then conclude that this is not a good analogy. (I did hope though, that this analogy was not raised by the students). Misconceptions was another major hurdle to begin with, but with the background reading, I felt confident knowing what the common misconceptions students may have on the subject. I do believe that once you know the likely misconceptions that students bring with them to the lesson, you (the teacher) will be able to, via discussion; diagrams etc. show the students the correct scientific concept Orbitals was my final hurdle, although I agree with Tsaparlis (1997b) and Gillesphie (1996) in that they should not be mentioned until university level, because the syllabus that the students would be examined on mentions orbitals, I knew that I would need to cover this idea. Taber (2004) mentions the various misconceptions and difficulties that students have with the term of ‘orbital’, but in an earlier article (2002c) he gives guidance on how to deal with this; time, reinforcement and practice. So, I needed to plan more than one activity for the students so they could reinforce and practice this concept. Feedback and analysis of the lessons

The delivery of the lessons on ‘electronic configuration’ was planned for a double afternoon lesson followed by a single lesson in the morning. I knew that the students had a ‘study’ period after the single lesson so I planned to stay in the laboratory so that the students could use me for support and guidance if they required. My strategy was to look at any misconceptions in the double lesson using the starter activity followed by the delivery of the content using the power point presentation and balloons. The remainder of the time would be used by the students to choose some of the activities which could then be completed in the subsequent single lesson and during their study time. I had also asked for the regular teacher to observe me teaching and give constructive feedback on the delivery of the lessons and the activities which I had planned. Finally, after the delivery of the lessons and the subsequent test, I asked the students to complete a questionnaire on the lessons and give some honest feedback. Unfortunately, the teacher that was due to observe my teaching was unable to stay for the majority of the double lesson in which I was to deliver the content, an emergency with a student meant that he had to leave the laboratory to deal with the problem. What he was able

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to do though, was to look at the power point presentation and activities I had set the students and gave verbal feedback on those as well as my interaction with the students during the single lesson.Teacher feedback:

The regular teacher of the students mentioned that the lesson plan was very detailed and covered the content extremely well. The starter activity he thought to be a very useful tool in diagnosing any misconceptions that the students may have, particularly with the power of discussion based upon the answers. The comment on the power point presentation was positive, he thought it to be well constructed and with sufficient detail. He particularly liked the fact that I planned to use balloons as 3-dimensional representations of shapes of orbitals along side the 2-dimenstional diagrams on the presentation, he thought this to be extremely beneficial to the kinaesthetic students. The activity ‘Electronic configuration of the first 36 elements’ in which the students would add arrows to the boxes to represent electrons he thought would be very valuable in that he ‘guaranteed’ that the majority of the students would make the classic errors when it came to the elements of Chromium and Copper. This is the reason why I had planned such an activity, thinking that the students would make the mistake, this would then allow me to discuss with them why such elements do not follow the ‘rules’ and as the classic saying goes ‘we all learn by mistakes’. He raised the issue of choice of activity and praised the amount of choice I had given to the students. He did mention though that perhaps I should have made a couple of the activities ‘compulsory’ and that these would be purposefully marked by me and that the students could self and peer assess the other activities. This is of course something I will do in the future The regular teacher also commented on the fact that I had set relevant homework, but pointed out that not all students would complete the task within the given time and that this would not be a reflection on me, it is a problem in general. On a more positive note, he did praise me highly on my willingness to stay after the single lesson for student support and guidance and was very pleased with the ‘uptake’ of this offer by the students. He congratulated me on my rapport with the students in that I was extremely friendly and welcoming towards the students, but done in a highly professional manner that the students still respected my position.On a final note, the teacher congratulated me on the use of a ‘test’ and commended me on the use of questions that would assess if the students had used their ‘study’ periods to further read up on the subject. The ‘test’ results of the students were extremely pleasing according to the regular teacher. Every student met their target and 2 students (28.6%) even exceeded their target. Student feedback:

The questionnaire that the students kindly completed for me and the ‘test’ on their learning of the subject has given me a means by which I can analyse the lessons in specific detail and has allowed me to feedback to their regular teacher any weaknesses/misunderstandings that he would need to revisit with them. 100% of the students enjoyed learning about ‘Electronic configuration’, although there were only 7 students, this is still a very pleasing result. From the starter activity, 42.9% of students either strongly agreed or agreed with the fact that it highlighted misconception(s) that they had, and of those, all of them strongly agreed that I (the teacher) explained the misconception(s) and helped turn them into the correct conception. 100% of the students found the power point presentation to be well delivered and useful and the various notes which I had produced on behalf of the students were equally well received. All students found the activity ‘Electron configuration of the first 36 elements’ to be a useful activity. As predicted by the student’s regular teacher and myself, the majority of the students made the classic error with Chromium and Copper. One student even commented: “At first I thought the activity to be useless, it seemed obvious what to do in terms of adding arrows to the boxes, it seemed a complete waste of time. When the teacher went through the answers, particularly with Chromium and Copper, I could then see the logic behind the activity. I will certainly remember the activity and will not make the same mistake again”. This was exactly the outcome that I had wanted, success!The feedback from the activities that I had planned for the students were well received. The only negative comment was that 14.3% of students (one student) found that the ‘Electronic configuration’ and ‘Ionisation energies’ activities were not useful. When questioned further the reason why they did not find them useful was “It involved too much writing, I like it when I have to answer straight forward simple questions”. When I brought this to the attention of the regular teacher, he commented on the fact that this particular student will probably get a grade E in their examination and will not continue with their study of chemistry after year 12 due to the fact that they do not like to write out long answers! It was this same student that disagreed with the homework and test activity being useful, they also disagreed with the activities being set at the correct level. Ever since the regular teacher explained their background I no longer feel upset with this criticism on my teaching of the subject. This one student agreed with all the other students (100%) that they had plenty of choice of activities, and all students (100%) felt I was knowledgeable about the subject, I was able to offer assistance when needed and that the help and guidance I gave was useful.

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One criticism made by 42.9% of the students was that felt they did not have enough time to complete the tasks. Although this is a high percentage, I do not feel disheartened as I pointed out to the students, the activities were not designed to be completed solely within the lessons, and they were designed to be completed during ‘study’ periods. One of the written comments made by a student was “I did not grasp the concept of orbitals when the teacher was explaining it using the power point. When she got out the balloons, it all made sense”. I am very pleased with myself for taking into account the various learning styles of students and feel that if I had not done so this student may have never understood the difficult concept of orbitals.Another comment that I was proud of was “I really liked the fact that the teacher gave us notes which we could use, it saved me a lot of time making my own notes. All I needed to do was highlight bits and add some extra notes. Thank you for this.” When students thank you for something, it really makes you feel worth while and I certainly felt valuable to these students. Self-criticism:

I believe that the students and their regular teacher have given me enough feedback to classify the lessons that I planned and delivered as ‘good’. I was pleased with my background reading and feel that it more than prepared me for the student’s misconceptions and difficulties in understanding such an abstract concept. What I am most pleased with though is that the students had not previously been taught the analogy of an atom being like the solar system. As a newly appointed ‘Head of Department’ I will be ensuring that my fellow teachers will never use this analogy in the future and have even gone so far as to mention it at Subject Leader Development Meetings to other colleagues. My rapport with the students was good, they respected me and felt confident to ask questions and were happy with the way in which their questions were answered. I was extremely proud of the fact that 71.4% of students (6 out of 7) volunteered to remain in the classroom after the single lesson to use me for help and guidance.I am extremely proud with the outcome of the ‘test’, the fact that every student had met their predicted grade was pleasing and the fact that 2 students had exceeded their predicted grade was most satisfying. Summary

Overall I am extremely proud of myself for what I have achieved with these students. From this challenge I have learnt that I need to teach modelling skills and should encourage students to use multiple models and take time to discuss and critique the models used in class. If I had been asked to teach even to GCSE level never mind A-level before attending the SASP course I would have ran a mile. The course has given me the subject and pedagogical knowledge to be able to teach students GCSE and even A-level in a subject that is not my specialism. Word count = 4298

References

Bent, H.A. (1984) Should orbitals be X-rated in beginning chemistry courses? Journal of Chemical Education, 61 (5), 421-423.Berry, K.O. (1986) What should we teach them in high school? Journal of Chemical Education, 63(8), 697-698.Cervellati, R. and Perugini, D. (1981) The Understanding of the Atomic Orbital Concept by Italian High School Students. Journal of

Chemical Education, 58 (7): 568-569.Fischler, H. and Lichtfeldt, M. (1992) Modern physics and students’ conceptions. International Journal of Science Education, 14: 181-190.Gilbert, J.K. (1998) Explaining with models. In: Ratcliffe, M (ed) ASE Guide to Secondary Science Education. Cheltenham: Stanley Thornes. 159-166.Gillesphie, R.J. (1991) What is wrong with the general chemistry course? Journal of Chemical Education, 68 (3): 192-194.Gillesphie, R.J. (1996) Bonding without orbitals. Education in Chemsitry, 33 (4): 103-106.Grosslight, L., Unger, C., Jay, E., and Smith, C. (1991) Understanding models and their use in science: Conceptions of middle and high school students and experts. Journal of Research in Science Teaching. 28 (9): 799-822.Harrison, A.G. and Treagust, D.F. (1996) Secondary Students’ Mental Models of Atoms and Molecules: Implications for Teaching Chemistry. Science Education. 80 (5): 509-534.Harrison, A.G. and Treagust, D.F. (2000) Learning about Atoms, Molecules , and Chemical Bonds: A Case Study of Multiple-Model Use in Grade 11 Chemistry. Science Education. 84 (3): 352-381.Hawkes, S.J. (1992) Why should they know that? Journal of Chemical Education, 69 (3): 178-181.Hodgson, T. (1995) Secondary mathematics modelling: Issues and challenges. School Science and Mathematics, 95 (7): 351-358.Justi, R. and Gilbert, J. (2000) History and Philosophy of Science Through Models: Some Challenges in the Case of ‘the Atom’. International

Journal of Science Education, 22 (9): 993-1009.Park, E.J. and Light, G. (2009) Identifying Atomic Structure as a Threshold Concept: Student mental models and troublesomeness. International Journal of Science Education, 31 (2): 233-258.

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Shiland, T.W. (1995) What’s the use of all this theory? – The role of quantum mechanics in high school chemistry textbooks. Journal of

Chemical Education, 72 (3): 215-219.Taber, K.S. (2002a) Conceptualizing Quanta – Illuminating the Ground State of Student Understanding of Atomic Orbitals. Chemistry Education: Research and Practice in Europe, 3 (2): 145-158.Taber, K.S. (2002b) Compounding Quanta – Probing the Frontiers of Student Understanding of Molecular Orbitals. Chemistry Education:

Research and Practice in Europe, 3 (2): 159-173.Taber, K.S (2002c) Chemical misconceptions – prevention, diagnosis and cure. Volume I: theoretical background. London: Royal Society of Chemistry.Taber, K.S (2003) The atom in chemistry education. Foundations of chemistry, 5 (1): 43-84 Taber, K.S (2004) Learning Quanta: Barriers to Stimulating Transitions in Student Understanding of Orbital Ideas. Science Education, 89 (1): 94-116Tsaparlis, G. (1997a) Atomic and Molecular Structure in Chemical Education: A Critical Analysis from Various Perspectives of Science Education. Journal of Chemical Education, 74 (8): 922-925.Tsaparlis, G. (1997b) Atomic orbitals, Molecular Orbitals and Related Concepts: Conceptual Difficulties among Chemistry Students. Research in Science Education, 27 (2): 271-287.Tsaparlis, G. (2002) Quantum-chemical concepts: Are they suitable for secondary students? Chemistry Education: Research and Practice

in Europe, 3 (2): 129-144

Critical reflection of how the learning process undertaken during the course has impacted on teaching practice.

Subject knowledge

I would say that before the course my subject knowledge of chemistry was not ‘weak’ but it was far from ‘strong’. Over the course of the year, I have improved my self-confidence enormously and developed a real interest in chemistry, I am now confident that I know chemistry. Last year, for certain topics, particularly atomic structure, I always felt ‘one step ahead’ of the students. Now, I am not spending my evenings ‘learning’ what I need to teach the next day, I now know the subject knowledge and I am confident with it. Eick (2002) suggests that self-confidence about one’s own teaching has an effect on student learning and I would most certainly agree. Thinking back to how I was a year ago, the students knew by the type of learning they experienced when I was not confident. Use of textbooks, watching of DVD clips and very little practical work meant I was ‘one step ahead’. Now, practical work, actual demonstrations (rather than watching them on a whiteboard) and varied activities that are thought about and differentiated for the various levels of students and various types of learners show the students that I am now equally confident with chemistry as I am with biology (my specialism). What I need to do in the future, is to try to become more confident with physics!! Before the course, I was always confused with the terminology used in chemistry, particularly with the terms ‘atom’, ‘element’, ‘molecule’ and ‘particle’, I would frequently get myself in a muddle and would use the terms interchangeably. I am now confident with the terminology, and because I am confident, I am able to correct the students if they use the incorrect term whereas before I would just ignore the use of the incorrect term. To try to alleviate the confusion for the students, this year I introduced a word wall. In my laboratory

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I now have a wall that is divided into sections for each year group that I teach. Each section is labelled, so the students know which part of the wall is relevant to them. For each topic, I then change the keywords so that the students know which words are important for the topic. This tool is a reference for the students to use to aid with the spellings of keywords. To integrate these words into everyday use, what I do for a starter from time to time, is to project the definitions of some of the topic words onto the whiteboard and ask the students to write down the word that corresponds with the definition, or for the lower ability pupils or the kinaesthetic learners I would ask them to physically go and choose the word from the board and place it next to the definition. Reflecting further on the success of this, wishing that I had done this earlier in my career, therefore helping more students, I then produced a ‘keywords’ page for each topic in the form of a table, the keyword being in the left hand column and the definition of the word being in the right hand column. I now give each student a copy of this ‘keywords’ page for their books at the beginning of the topic for them to use as a reference also. Since the use of the word wall and keywords page, what I have noticed is that more students are using keywords correctly and have the correct concept in terms of the definition of the word. Ausubel (2000) postulates that meaningful learning occurs when the learner’s appropriate existing knowledge interacts with new learning. Ausubel considered prior knowledge to be the most significant factor in planning teaching. An accurate estimation of the actual ‘prior learning’ that a learner will bring to bear on new material presented is therefore essential for effective teaching. Considering students prior learning is something that I did briefly when planning lessons, but upon spending some time discussing this with colleagues on the course I came to realise just how important this is. Early on in the course, a mention was made of a set of books that we might like to read ‘Key stage three – teaching the right science at the right time’ (Chemical and material behaviour (Gadd and Warren, 2009); Earth, environment and the universe (Petheram, Warren and Locke, 2009); Energy, electricity and forces (Petheram, 2009); Organisms, behaviour and health (Locke, 2009) and Interpreting how science works. (Gadd and Ellis, 2009)) I glanced through the books during one of the sessions and liked what I saw. I then purchased the full set for myself. Since then, I have read the set of books and use them as a reference source when planning lessons, particularly at key stage 3. What I particularly find useful about these books is the way in which they are presented. It begins with ‘the journey so far’ which summarises the prior knowledge, skills and understanding that you would expect pupils to have from key stage 2, the most significant factor in planning teaching according to Ausubel. It includes teaching issues which covers vocabulary and language, misconceptions and questions, other sections include application and implications, history and culture, hot topics, fact and figures and ‘the journey ahead’ which indicates the related science that students will encounter at key stage 4 and beyond. These set of books have now become part of my reference library which I use on a regular basis. During the course, what I became more aware of and have mentioned on numerous occasions in my reflective journal, is the progression of chemistry through the various key stages. I am now fully aware as to how it all fits together. I am now able to see the overall picture and regularly reinforce with my students how the various topics in chemistry link together. The students are now able to explain why, rather than just regurgitate the facts. This higher order thinking skill that some of the students are able to do is all down to the confidence I gained as a direct result of the reflection that I carried out at the end of each session. At the beginning of the course, although my subject knowledge was reasonable, I knew that when it came to looking at some topics at GCSE and virtually all of the A-level content, that my lack of knowledge would be of a hindrance. On the evaluation questionnaires that participants were asked to complete each half term, I mentioned that it would be helpful if we were given notice as to what would be the following week’s topic so that background reading could be explored if you wished. This advance notice of the topic in which we would subsequently examine was extremely beneficial to me. When we began looking at A-level, I always used to read the relevant chapters in the ‘Chemistry in Context’ textbook written by Hill and Holman (2000). This prior reading would give me a foundation on which my knowledge and confidence would grow the subsequent week when the tutor would discuss with us the content knowledge and give us ideas for practical’s to use with the students. Student ideas and misconceptions

As a chemistry teacher I have a challenging task - to help students make sense of a wide range of models and theoretical entities that we use in chemistry. To overcome this challenge, I have taken advice from K. Taber (2009) ‘to learn as much as possible about the way your

students understand chemistry and to teach accordingly’. Learning is an active process, and what students do with facts and ideas depends on what they already think and believe. Being able to recognise and work with these student-held ideas and conceptions is thus a key component of an effective scheme of work. As a teacher I need to understand what my students currently think, if I am to plan effective teaching that will not either ‘go over their heads’ or become significantly distorted. I therefore, should be able to avoid or reduce misconceptions by more careful planning and sequencing of my teaching.During the duration of the SASP course, what I have been doing on a regular basis is noting the misconceptions discussed during the ‘face to face’ days into my planning. At the beginning of my plans for each topic I have included a ‘common misconceptions’ page which lists the misconceptions for that topic. This is not a finished article, this is very much a work in progress in which I am able to add too if

9

necessary. As a recently promoted ‘Head of department’, I placed this misconceptions folder into the staff shared drive on the school network, each topic has a misconceptions page for staff to access. Various members of staff have given very positive feedback on this and have added to the files the misconceptions that they have come across also. This ‘common misconceptions’ file I have found particularly beneficial to jog my memory of what was discussed on the contact day. What I have also done, is add misconceptions that have been discussed in the various educational journals which I have read as I became more aware and more interested in the subject. ‘Beyond appearances: students’ misconceptions about basic chemical ideas’ written by Kind (2004) and ‘Chemical misconceptions – prevention, diagnosis and cure Volumes I and II’ written by Taber (2002) and ‘Student alternative conceptions in chemistry’ (Horton) are like the ‘bible’ of misconceptions, it is here where I always begin when reading further about misconceptions that students may hold and it is here that I look for ideas on how to deal with misconceptions. The literature on the subject of ‘misconceptions’ is certainly not complete. It is extensive and researchers continue to discover new misconceptions by asking new questions. Johnstone (2000) discusses that chemistry is regarded as a difficult subject for students due to the fact that the subject operates on and connects three levels of thought: the macro and tangible (what can be seen, touched and smelt); the sub micro (atoms, molecules, ions and structures) and the representational (symbols, formulae, equations, molarity, mathematical manipulation and graphs). Many misconceptions originate from the fact that students are introduced to all three levels of thought simultaneously, hence overloading. This journal also mentions that research literature has been dominated by work on misconceptions, but little has as yet appeared about how to reverse these or to avoid them altogether. The article also points out that most countries are struggling to find well qualified and competent teachers. I would have agreed with the statement in my non-specialism of chemistry. Before completion of the SASP course, I did not feel qualified and competent enough to teach A-level, now though, I would disagree with the chemistry aspect, I do feel qualified and competent. Looking back on the enjoyment and fulfilment I have achieved on completion of this course, I do believe that all science teachers should have the opportunity to complete such a course. I have become interested in educational research because of the course, this has influenced my thinking and has most certainly lead to better teaching and learning with myself, the teacher, the ‘lead learner’ What I have become more aware of during the course is that misconceptions can occur when students hold meanings for everyday words that differ from the scientific meaning. Two common examples being ‘heat’ and ‘temperature’. What I have discovered since the use of the word wall and the use of the keywords page (as mentioned in the ‘subject knowledge’ section) is that the number of students who give the incorrect definition of a word and hence have a misconception has greatly decreased. This aid to learning has not only worked in chemistry, it has also shown to significantly reduce misconceptions in biology and physics also. In biology, students frequently misunderstood the word ‘diet’, the majority thought it to be a word to describe ‘eating less food’. Since the use of the word wall and keywords page, the majority of students now correctly use the word and understand the word ‘diet’ to mean ‘what you eat’. In physics the misunderstanding over the words ‘mass’ and ‘weight’ has been significantly reduced since using word wall and keywords pages. Schmidt (1997) discusses how misconceptions form a meaningful and coherent alternative framework in student’s minds, which is very robust and difficult to change. What I have discovered during the duration of the course, is that with constant reinforcement and practice students can change their understanding of everyday words and can learn and understand the scientific meaning of such a word, hence reduce their misconceptions. Practical work

Practical work is the key component to science education. I am now able to increase the quantity and quality of practical work in my lessons. The main reasons I choose a practical activity in a lesson are that good quality practical work can engage students, help them to develop important skills, help them to understand the process of scientific investigation, and develop their understanding of concepts. Before the completion of the SASP course, my students would have very few opportunities to have a go at practical work themselves in chemistry topics and would witness very few demonstrations from myself, instead I would frequently show them ‘you tube’ videos of other teachers demonstrating. What my students experience now, is completely different, where equipment is available, the students themselves have the opportunity to investigate and if it is not realistic for the students to have a go themselves, then I am now more than willing to demonstrate to them, I cannot recall the last time the students experienced a ‘you tube’ video since attending the course. During the last week of term, I discussed with the students their experiences of science, all of the students questioned, mentioned that they have participated in more investigative and practical work this year and have enjoyed the demonstrations first hand rather than watching them on a screen. It is because of the learning I experienced myself on the course that I now understand the theory and relevance of practical work, it was on the course that I learned and was able to practice the correct practical techniques such as titrating accurately and using the apparatus correctly which I am now confident to pass on to my students. It was on the course that I had the opportunity to try new things, an opportunity that I would never have in the school environment, these opportunities were one of the most beneficial aspects of the course for me. Another positive aspect to practical work that I mentioned in my journal was the fact that I found out other ways to do the same experiment, this experience will allow me to adapt practical activities for various students, simplifying for the less able and extending for the gifted. During one of the sessions on ‘Elements and compounds’ we were given the opportunity to try out some

10

practical activities. When demonstrating the reactivity of Group 1 metals with water, I had always just placed a small piece of the metal into a water trough. During this session we were shown how to demonstrate that Sodium burns with an orange flame (place a piece of sodium on a piece of filter paper then place on water) and how to demonstrate that hydrogen gas is given off by the metals (place a combustion tube over the piece of metal, leave for a few seconds, then light the top of the tube). These demonstrations I had never seen before myself, hence I had a go during the session and then used my new found knowledge and skills to demonstrate to the students back in school, they were amazed!!During another session on ‘Acids and alkali’s’ we were shown how to demonstrate an effervescent universal indicator ‘rainbow’. Again, I had a go myself and then when teaching acids and alkalis I used this demonstration as an aid to discussion for the students. The higher ability students were able to explain what they observed which I was then able to mark and level accordingly and for the less able students they were able to draw a burette, colour in the appropriate parts and then label the parts in terms of the pH number and state if it was an acid, alkali or neutral and were appropriate, give the strength. From my journal I could go further and list numerous ways in which the practical experiences I came across during the course have improved my teaching skills and the learning experienced by my students, but unfortunately I do not have the time to list them all. Osborne et al (2003) conclude that engagement by students is raised by opportunities for students to take control of their learning and greater student autonomy; students desire more opportunities in science for practical work, extended investigations and opportunities for discussion. During this past academic year, the students have certainly been given more opportunities for practical work and extended investigations, all because of the new found knowledge and skills I have acquired. Hereby now lays a problem that I will have to give consideration. I could just simply incorporate as many practical activities for the students as possible, but this could lead to ‘overkill’. Abrahams and Miller (2008) have analysed ‘Does practical work really work?’ and have suggested that practical work in science could be significantly improved if teachers recognise that explanatory ideas do not ‘emerge’ from observations. Their two-domain model is a useful tool as a teacher in thinking about practical work. Level 1 task’s are where students do what the teacher intended and level 2 tasks promote the learning the teacher intended. In the context of practical work, there is a substantial difference in learning demand between tasks in which the primary aim is that students should see an event or phenomenon or become able to manipulate a piece of equipment (level 1) and tasks where the aim is that students develop an understanding of certain theoretical ideas or models that might account for what is observed (level 2). When incorporating more practical opportunities for the students, I am going to have to be very selective and thoughtful of the purpose of such an activity. I am going to have to ensure that practical opportunities are not solely level 1 tasks, I will have to ensure that level 2 tasks are integrated also. This analysis of practical activities in lessons will be completed in biology and physics as well as chemistry therefore improving science lessons as a whole. Discussion in the classroom

Newton, Driver and Osborne (1999) found that less than five per cent of science lesson time is given over to group discussions. In teaching science it is essential to build in opportunities for pupils to talk through ideas and listen to the ideas of others. Dialogue is central to learning science as it involves discussion of evidence and ideas. Through discussion pupils learn in the classroom that science is social, is not impersonal and is interesting. It makes science real and relevant. Osborne et al (2003) point out, as previously mentioned, that student engagement is raised when there is opportunities for discussion. Previously I would have planned very little opportunities for discussion in the chemistry topics, now, due to my increased confidence in my knowledge and a greater understanding of progression, I try to plan as many opportunities as possible for the students to discuss the topic / their understandings / their findings, this I have tried over the academic year in various ways, discussion in pairs, discussion in groups, discussion as a class and I have found all to be as equally engaging for the students. The way forward

DeWitt and Osborne (2008) indicate the importance of providing students with experiences that may engage them with science while they are at school. They mention the longstanding ASE (Association of Science Education) programme ‘Science Across the World’ (SAW) whereby students participate in science projects and share reports about their activities with classes in other countries. What emerged from their findings are some insights into the kinds of activities likely to engage secondary students with science topics and motivate their further study of science. Their findings suggest that students are engaged by activities that:

provide them with a much greater degree of autonomy in their learning which is both challenging and different; are hands-on and practical in nature; involve collaboration both with classmates and with students in other countries (using modern technology); focus on topical issues in science treating them in an in-depth manner (extend beyond a single lesson)

This article has led me to sign up the school to the ‘Science Across the World’ programme. We are planning on integrating this programme into our current scheme of work as of September 2010. In the first term, our Year 7 students will be studying ‘Acids and

11

alkali’s’ and we will be using the ‘Acid Rain’ topic of the programme to link with students from Nebraska, Lincoln, USA. This I hope will be the first of many opportunities for my students to collaborate with students around the world.Pasl, A.J. et al (2009) showed how a hands-on experimental approach to teaching improved student’s attitudes towards learning science, increased the students understanding and achievement and increased their self-confidence of being inventors. Their approach was to leave the students to conduct the experiment / task themselves without any explanation from a teacher; the students were also left to draw their own conclusions. Surveys and interviews were conducted after 2 years of experimentation. The results of the study showed that 73% of the students in the experimental group preferred the new approach as compared to 20% of the control group. Little difference was shown between the groups on knowledge-comprehension level based exams; however, the experimental group scored 63% higher on challenging exams that required higher thinking levels. In addition, 73% of the students in the experimental group expressed increased levels of self-confidence towards invention as compared to only 36% of the control group. These figures are certainly impressive and I cannot ignore them. Just before the school broke up for the summer break, I tried out this approach with positive results. As part of the ‘Solids, liquids and gases’ topic when looking at the process of ‘expansion’ and ‘contraction’ what I would have done in the past would be to demonstrate to the students how all three states expand and contract. What I proceeded to do after reading this literature was to allow the students to discover the processes themselves in a ‘circus’ of experiments. The students, working in groups of 3 or 4, visited each station where they were met with a set of instructions on what to do. Lower ability students had a set of questions to act as a prompt for discussion. This activity was very successful, the students enjoyed the experience, they were independent in their learning, were able to work in groups successfully and were able to discuss with each other what they thought was happening. This opportunity allowed the students to investigate and explore the everyday world and develop a deep understanding from their inquiries. In previous years most students could remember the demonstrations, but only the higher ability students could give a scientific explanation in terms of particles. This time, using the student hands-on experimental approach, more students were able to give a scientific explanation of what they observed in terms of particles. What I will do in the future, starting in September, is to look at increasing the number of opportunities that students have throughout the year to conduct an experiment / task themselves without any input from myself as well as increasing the opportunities for discussion as previously mentioned. Summary

Ratcliffe and Miller’s (2009) study on ‘Teaching for Understanding of Science in Context’ suggests that teachers’ practices can be changed, but this takes time and requires considerable support from teaching materials and other forms of professional development that encourage reflection on practice. This I would completely agree with, over the course of just a year, with support and inspiration from the SASP tutors I have discovered a new found interest in teaching. I have been motivated to research and read educational journals to study the latest findings and to try out new pedagogical techniques. As newly appointed ‘head of department’, I hope that I can inspire my colleagues just as much as I have been inspired by the tutors. Word count = 4098REFERENCESAbrahams, I. and Miller, R. (2008) Does practical work really work? A study of the effectiveness of practical work as a teaching and learning method in school science. International Journal of Science Education. 30 (14): 1945-1969Ausubel, D.P. (2000). The acquisition and retention of knowledge: A cognitive view. Dordrecht: Kluwer academic publishers.DeWitt, J. and Osborne, J. (2008) Engaging students with science: in their own words. School science review. 90 (331): 109-116Eick, C.J. (2002) Studying Career Science Teacher’s Personal Histories: A methodology for understanding intrinsic reasons for career choice and retention. Research in Science Education. 32 (3): 353-372.Gadd, K and Warren, D. (2009) Key stage three: teaching the right science at the right time ‘Chemical and material behaviour’. Bath: JMD Print ManagementGadd, T and Ellis, R. (2009) Key stage three: teaching the right science at the right time ‘Interpreting how science works’. Bath: JMD Print ManagementHill, G and Holman, J. (2000) Chemistry in Context. (5th Edition) Cheltenham: Nelson Thornes. Horton, C. Student alterntative conceptions in chemistry. www.daisley.net/hellevator/misconceptions/misconceptions.pdf. (accessed July 2010).Johnstone, A.J. (2000) Teaching of chemistry – logical or psychological? Chemistry education: Research and practice in Europe. 1 (1): 9-15. Kind, V. (2004) Beyond appearances: students’ misconceptions about basic chemical ideas. www.rsc.org/education/teachers/learnet/pdf/LearnNet/rsc/miscon.pdf (accessed July 2010).Locke, J. (2009) Key stage three: teaching the right science at the right time ‘Organisms, behaviour and health’. Bath: JMD Print Management

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Newton, P., Driver, R. and Osborne, J. (1999) The place of argumentation in the pedagogy of school science. International Journal of

Science Education. 21 (5): 553-576.Osborne, J., Simon, S. and Collins, S. (2003) Attitudes towards science: a review of the literature and its implications. International Journal

of Science Education. 25 (9): 1049-1079Pasl, A.J. et al. (2009) Autonomy in Science Education: A Practical Approach in Attitude Shifting Towards Science Learning. Journal of

Science Education and Technology. 18 (6): 476-486Petheram, L. (2009) Key stage three: teaching the right science at the right time ‘Energy, electricity and forces’. Bath: JMD Print ManagementPetheram, L., Warren, D. and Locke, J. (2009) Key stage three: teaching the right science at the right time ‘The environment, earth and universe’. Bath: JMD Print ManagementRatcliffe, M. and Miller, R. (2009) Teaching for Understanding of Science in Context: Evidence from the Pilot Trials of the Twenty First Century Science Courses. Journal of Research in Science Teaching. 46 (8): 945-959Schmidt, Hans-Jurgen. (1997) Students’ Misconceptions – Looking for a Pattern. Science Education 81 (2): 123-135.Taber, K. (2002) Chemical misconceptions – prevention, diagnosis and cure Volumes I and II. London: Royal Society of Chemistry. Taber, K. (2009) Misconceiving chemistry: the mismatch between chemical concepts and student thinking. School science review. 91 (335): 87-96.

Lesson plan

Atomic structure “Electron arrangement”

Aim know the electron configurations of atoms

and ions up to Z = 36 in terms of levels and sub-levels (orbitals) s, p and d

know the meaning of the term ionisation

Objectives To be able to draw the electron configurations

of atoms and ions up to Krypton (Z=36) using the arrow and box method

To be able to draw the electron configuration

13

energy. understand how ionisation energies in Period

3 (Na – Ar) and in Group 2 (Be – Ba) give evidence for electron arrangement in sub-levels and in levels

of atoms and ions up to Krypton (Z=36) using the orbital method

To be able to draw the electron configurations of atoms and ions up to Krypton (Z=36) using the shorthand method

To be able to give the meaning of the term ionisation energy

To be able to write an equation to illustrate the process of ionisation of an element

To be able to explain how the study of ionisation energies and atomic spectra give evidence for electron arrangement in levels and sub-levels.

To be able to identify an element using line emission spectra

Health and safety risk assessmentCare must be taken when using the interactive whiteboard. Avoid straining eyes and low light.Overall pictureStudents should know that: When atoms react, atoms are redistributed. Evidence for the electronic structure of atoms

was obtained by studying ionisation energies. Ionisation energies were obtained from

emission spectra. Each element has a characteristic line

emission spectrum which can be used to identify it.

Ionisation energies were used as evidence for shells and sub-shells.

Orbitals are the regions in which there is the greatest probability of finding particular electrons.

Each orbital can hold a maximum of two electrons and if they do, they will be spinning in opposite directions.

s sub-shells contain one orbital; p sub-shells contain three orbitals and d sub-shells contain five orbitals.

Principal quantum numbers (PQN) are used to denote the shells of electrons in atoms.

Links to AQA specification3.1.1 Atomic structure – Electron arrangementResources to be used

GCSE exam questions Power point presentation – Electronic

configuration Balloons (long and round) Worksheet; Electronic structure of the

first 36 elements Worksheet; Electronic configuration Worksheet; Ionisation energies Worksheet; Ionisation energies – name

the element Handout notes; Electronic structure,

Ionisation energy and Electronic structure of ions

Practice exam questions for electronic structure

Laptops Homework – Electronic structure Test – Electronic structure

14

Starter activityStudents to complete the GCSE exam questions on atomic structure (this will from a basic foundation to the lesson) followed by a brief discussion on the answers.Main activity

Using the power point presentation on Electronic configuration, talk through the Bohr model of the atom and how that was acceptable up to GCSE level. Discuss shells, sub-shells and energy levels followed by the rules of how electrons are arranged in atoms. Next discuss the shapes of orbitals (students might find it useful if balloons are used to show the shapes of orbitals 3-dimensionally) and the order in which orbitals are filled.

Students to then work through the worksheet ‘Electronic configuration of the first 36 elements’ where they add arrows to the boxes provided.

Using the power point for reference, discuss the answers to the worksheet. Look closely at Chromium and Copper (they do not follow the rules!)

Show pupils how to use the shorthand way of writing the electronic configuration of elements. Stress that if this method is used, then a noble gas is always used. e.g. Ca = 1s22s22p63s23p64s2 will become [Ar] 4s2

Using the Ionisation energy notes, discuss with the students how ionisation energies were used as evidence for sub-shells of electrons.

Using the power point presentation on Electronic configuration, talk through the electronic configuration of ions.

Give out notes on Electronic structure and Electronic structure of ions. Students to then work through activities of their choice

- Worksheet on Electronic configuration- Worksheet on Ionisation energies- Worksheet on Ionisation energies – name the element- Practice examination questions on electronic structure- Best choice website – AQA AS, 1.1 Atomic structure – ‘Shells and sub-shells’ and

‘Electronic configuration’Plenary

Have the students met the objective criteria? Go through each objective and check that students are able to successfully answer the objectives.

Give out homework on electronic structure. Explain that at the start of the next session there will be a quick test on electronic structure.

Questionnaire analysis

Electron configuration

(Please tick)strongly

agree agree disagree Strongly

disagree

1 I enjoyed learning about ‘Electron configuration’ 228.6%

571.4%

2 I understand why we learnt about ‘Electron configuration’

342.9%

457.1%

3 I am able to draw the electron configuration of atoms and ions up to Krypton using the arrow and box method

685.7%

114.3%

4 I am able to draw the electron configuration of atoms and ions up to Krypton using the orbital method

571.4%

228.6%

5 I am able to draw the electron configuration of atoms and ions up to Krypton using the shorthand method

571.4%

228.6%

15

6 I am able to explain the term ionisation energy 342.9%

457.1%

7 I am able to write an equation to illustrate the process of ionisation of an element

457.1%

342.9%

8 I am able to explain how the study of ionisation energies and atomic spectra give evidence for electron arrangement in levels and sub-levels

228.6%

342.9%

228.6%

9 I am able to identify an element using line emission spectra

228.6%

457.1%

114.3%

10 I am able to give a definition of the term orbital 671.4%

114.3%

11 I know how many electrons can be held in an orbital

671.4%

114.3%

12 I know how many orbitals are in each of the sub-shells (s, p and d)

671.4%

114.3%

13 The starter activity helped to recall my knowledge on atomic structure at GCSE level

457.1%

342.9%

14 The starter activity highlighted misconception(s) of my knowledge on atomic structure

114.3%

228.6%

457.1%

15

16

17

18

19 20

The teacher explained the misconception(s) that I had, and helped to turn them into the correct conception

342.9%

The power point presentation was well delivered, it helped to increase my knowledge

571.4%

228.6%

The activity ‘Electron configuration of the first 36 elements’ was a useful exercise

671.4%

114.3%

I know how Chromium and Copper do not follow the rules

342.9%

342.9%

114.3%

The teacher explained how ionisation energies were used as evidence for sub-shells of electrons 2

28.6%4

57.1%1

14.3%The notes on ‘Ionisation energy’ were helpful 6

71.4%1

14.3%21 The notes on ‘Electronic structure’ were helpful 6

71.4%1

14.3%22 The notes on ‘Electronic structure of ions’ were

helpful 671.4%

114.3%

23 The worksheet on ‘Electronic configuration’ (15 questions) was a useful activity 2

28.6%4

57.1%1

14.3%24 The worksheet on ‘Ionisation energies’ (11

questions) was a useful activity 228.6%

457.1%

114.3%

25 The worksheet on ‘Ionisation energies – name the element’ was a useful activity 4

57.1%3

42.9%26 The practice examination questions on electronic

structure was a useful activity 571.4%

228.6%

27 The ‘Best Choice’ website was a useful activity 671.4%

114.3%

28 The homework on ‘Electronic structure’ was a useful activity 2

28.6%4

57.1%1

14.3%29 The test on ‘Electronic structure’ was a useful

activity 228.6%

457.1%

114.3%

30 Overall, the lessons on ‘Electronic structure’ were set at the correct level for me 5

71.4%1

14.3%1

14.3%31 The activities set had plenty of choice 6

71.4%1

14.3%

16

32 The activities were set at the right level for me (I felt challenged) 2

28.6%4

57.1%1

14.3%33 I felt we had enough time to complete the tasks 4

57.1%3

42.9%34 The teacher was knowledgeable about the

subject 342.9%

457.1%

35 The teacher was able to offer assistance when I needed it 4

57.1%3

42.9%36 The teachers ‘help’ and ‘guidance’ was useful to

me 457.1%

342.9%

Any other comments about the lessons on ‘Electron configuration’:

________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

Lesson activities

GCSE Exam Questions.

January 2008 Lower

17

June 2008 Lower

June 2008 Higher

18

January 2009 Lower

19

June 2009 Lower

20

21

June 2009 Higher

22

Electronic structure of the first 36 elements.Atomic

Number (Z) 1s 2s 2p 3s 3p 3d 4s 4pH 1

He 2

Li 3

Be 4

B 5

C 6

N 7

O 8

F 9

Ne 10

Na 11

Mg 12

Al 13

Si 14

P 15

S 16

Cl 17

Ar 18

23

K 19

Ca 20

Sc 21

Ti 22

V 23

Cr 24

Mn 25

Fe 26

Co 27

Ni 28

Cu 29

Z 30

Ga 31

Ge 32

As 33

Se 34

Br 35

Kr 36

24

Electronic configuration

Write the electronic configuration of the following using the arrow and box method:

1. C

2. Cu

3. Mg+

Write the electronic configuration of the following using the orbital method:

4. N3-

5. Ar

6. Sc3+

7. Mn2+

8. Fe3+

9. V3+

Write the electronic configuration of the following using the shorthand arrow and box method:

10.Cl-

11.Fe

12.Br

Write the electronic configuration of the following using the shorthand orbital method:

13.Cr

14.Ga3+

15. Pb2+

Ionisation energies

25

1. Why does the first ionisation energy of atoms generally increase across a period?

2. Why is the first ionisation energy of boron less than that of beryllium?

3. Why is the first ionisation energy of oxygen less than that of nitrogen?

4. Why do first ionisation energies decrease down a group?

5. Why does helium have the highest first ionisation energy of all the elements?

6. Why is the second ionisation energy of an atom always greater than the first?

7. Why is the second ionisation energy of sodium much greater than the first?

8. Why does atomic size decrease across a period?

9. Why does atomic size increase down a group?

10.Why are cations always smaller than the corresponding atoms?

11.Why are anions always larger than the corresponding atoms?

Ionisation energies – name the element.Name the element, group and write down the electron configuration using the graphs below:

26

PRACTICE EXAMINATION QUESTIONS FOR ELECTRONIC STRUCTURE

1. (i) Complete the electron arrangement for a copper atom.

1s2 …………………………….........................................................................

(ii) Identify the block in the Periodic Table to which copper belongs.

........................................................................................…...............................

(iii) Deduce the number of neutrons in one atom of 65Cu

........................................................................................…...............................(3)

2. (i) Give the electron arrangement of an Fe2+ ion.

...........................................................................................................................

(ii) State why iron is placed in the d block of the Periodic Table.

27

...........................................................................................................................

........................................................................................................................... (2)

3. Complete the electronic configurations for the sulphur atom, S, and the sulphide ion, S2–.

S 1s2 ..............................................................................................................

S2– 1s2 ..............................................................................................................(2)

State the block in the Periodic Table in which sulphur is placed and explain your answer.

Block ....................................................................................................................

Explanation ..........................................................................................................(2)

4. In terms of s and p sub-levels, give the electronic configuration of an aluminium atom.

...............................................................................................................................(1)

How many neutrons are there in one 27Al atom?

...............................................................................................................................(1)

5. Explain why the first ionisation energy of rubidium is less than the first ionisation energy of krypton.

...................................................................................................................................

................................................................................................................................... (2)

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6. Aluminium, magnesium and vanadium are metals.

(a) Complete the electronic configurations for aluminium and vanadium.

Electronic configuration of aluminium 1s2 .................................................

Electronic configuration of vanadium 1s2 .................................................(2)

(b) State the block in the Periodic Table to which magnesium belongs.

...........................................................................................................................(1)

7. (a) Complete the following to show the electronic configuration of silicon.

1s2 2s2 ........................................................................................................................(1)

8. (a) Complete the electron arrangement for the Mg2+ ion.

1s2 ................................................................................................................................(1)

(b) Identify the block in the Periodic Table to which magnesium belongs.

.....................................................................................................................................(1)

(c) The Ne atom and the Mg2+ ion have the same number of electrons. Give two reasons why the first ionisation energy of neon is lower than the third ionisation energy of magnesium.

Reason 1 ......................................................................................................................

Reason 2 .......................................................................................................................(2)

(d) There is a general trend in the first ionisation energies of the Period 3 elements, Na – Ar

(i) State and explain this general trend.

Trend .................................................................................................................

Explanation .......................................................................................................

...........................................................................................................................

...........................................................................................................................

(ii) Explain why the first ionisation energy of sulphur is lower than would be predicted from the general trend.

...........................................................................................................................

...........................................................................................................................

........................................................................................................................... (5)

29

9. The diagram below shows the electronic structure of boron.

E n erg y

2 p

2 s

1 s(a) The electrons are represented by arrows. What property of the electrons do these ‘up’ and

‘down’ arrows represent?

.....................................................................................................................................(1)

(b) Suggest why electrons which occupy the 2p sub-levels have a higher energy than electrons in the 2s sub-level.

.....................................................................................................................................(1)

(c) Explain the meaning of the term first ionisation energy.

.....................................................................................................................................

.....................................................................................................................................(2)

(d) Explain why boron has a lower first ionisation energy than beryllium.

.....................................................................................................................................

.....................................................................................................................................(3)

(e) Explain why the first ionisation energy of helium is very large.

.....................................................................................................................................

.....................................................................................................................................(1)

10. (a) (i) State the general trend in the first ionisation energy of the Period 3 elements from Na to Ar.

...........................................................................................................................

(ii) State how, and explain why, the first ionisation energy of aluminium does not follow this general trend.

............................................................................................................................

............................................................................................................................

............................................................................................................................ (4)

(b) Give the equation, including state symbols, for the process which represents the second ionisation energy of aluminium.

30

...................................................................................................................................... (1)

11. There is a general trend in the values of the first ionisation energies of the elements Na to Ar. The first ionisation energies of the elements Al and S deviate from this trend.

(a) Write an equation, including state symbols, to represent the process for which the energy change is the first ionisation energy of Na.

...............................................................................................................................(2)

(b) State and explain the general trend in the values of the first ionisation energies of the elements Na to Ar.

Trend .....................................................................................................................

Explanation ...........................................................................................................

...............................................................................................................................

...............................................................................................................................(3)

(c) State how, and explain why, the values of the first ionisation energies of the elements Al and S deviate from the general trend.

How the values deviate from the trend .................................................................

Explanation for Al ................................................................................................

...............................................................................................................................

Explanation for S ..................................................................................................

................................................................................................................................(5)

12. (a) What is meant by the term first ionisation energy?

...............................................................................................................................

...............................................................................................................................

(b) The diagram below shows the variation in first ionisation energy across Period 3.

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F irs tio n isa tio ne n ergy

N a M g A l S i P S C l A rP er io d 3 e le m e n t

(i) What is the maximum number of electrons that can be accommodated in ans sub-level?

......................................................................................................................

......................................................................................................................

(ii) What evidence from the diagram supports your answer to part (b(i)?

......................................................................................................................

......................................................................................................................

(iii) What evidence from the diagram supports the fact that the 3p sub-level is higher in energy than the 3s?

......................................................................................................................

......................................................................................................................

(iv) What evidence from the diagram supports the fact that no more than three unpaired electrons can be accommodated in the 3p sub-level?

......................................................................................................................

......................................................................................................................

......................................................................................................................(5)

13. Explain why atomic radius decreases across Period 3 from sodium to chlorine.

.....................................................................................................................................

.....................................................................................................................................(2)

14. The diagram below shows the trend in the first ionisation energies of the elements from neon to aluminium.

32

N e N a M g A l S i P S

F irs t io n isa tio ne n erg y

E le m e n t

(a) Draw crosses on the graph to show the first ionisation energies of silicon, phosphorus and sulphur.

(3)

(b) Write an equation to illustrate the process which occurs during the first ionisation of neon.

.....................................................................................................................................(1)

(c) Explain why the first ionisation energy of neon and that of magnesium are both higher than that of sodium.

Explanation for neon ..................................................................................................

.....................................................................................................................................

.....................................................................................................................................

Explanation for magnesium .......................................................................................

.....................................................................................................................................

.....................................................................................................................................(4)

(d) Explain why the first ionisation energy of aluminium is lower than that of magnesium.

.....................................................................................................................................

.....................................................................................................................................(2)

Electronic structure - homework

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1. (a) In terms of sub-levels, give the complete electronic configuration of the nitrogen atom,

N, and of the nitride ion, N3–.

N...............................................................................................................................

N3– ..........................................................................................................................(2)

(b) Complete the electronic configurations for the metals sodium and iron.

Electronic configuration of sodium 1s2 ..................................................................

Electronic configuration of iron 1 s2 ..................................................................... (2)

(c) Complete the following electronic configurations. [Ar] represents the electronic configuration of the argon atom.

A copper atom [Ar] ................................................................................

A copper(II) ion [Ar] ................................................................................(2)

(d) Give the electronic configuration of the F– ion in terms of levels and sub-levels.

....................................................................................................................................

(1)

(e) Complete the following to show the electronic configuration of silicon.

1s2 2s2 ....................................................................................................................(1)

(Total 8 marks)

2. The values of the first ionisation energies of neon, sodium and magnesium are 2080,

34

494 and 736 kJ mol–1, respectively.

(a) Explain the meaning of the term first ionisation of an atom.

.....................………………………………………………………………………….

.....................………………………………………………………………………….

.....................………………………………………………………………………….(2)

(b) Write an equation to illustrate the process occurring when the second ionisation energy of magnesium is measured.

.....................………………………………………………………………………….

.....................………………………………………………………………………….(2)

(c) Explain why the value of the first ionisation energy of magnesium is higher than that of sodium.

.....................………………………………………………………………………….

.....................………………………………………………………………………….

.....................………………………………………………………………………….(2)

(d) Explain why the value of the first ionisation energy of neon is higher than that of sodium.

.....................………………………………………………………………………….

.....................………………………………………………………………………….

.....................………………………………………………………………………….(2)

(Total 8 marks)

ELECTRONIC STRUCTURE TEST

1. (a) Using arrows and to represent electrons, complete the energy-level diagram below to show the electronic arrangement in an atom of carbon.

2 s2 p

1 s(2)

35

(b) In terms of sub-levels, give the electronic configuration of the carbon ion C2+

...................................................................................................(1)

(c) Give the complete electronic configuration of krypton in terms of s, p and d sub-levels.

...................................................................................................(1)

(d) Complete the following electron configurations.

(i) Cu 1s22s22p63s23p6 ..................................................(1)

(ii) Cu2+ 1s22s22p63s23p6 ..................................................(1)

(Total 6 marks)

2. (a) Write equations to show the chemical processes which occur when the first and the second ionisation energies of lithium are measured.

First ionisation energy equation ..............................................

Second ionisation energy equation ..........................................(3)

(b) (i) Explain why helium has a much higher first ionisation energy than lithium.

...........................................................................................

36

...........................................................................................

...........................................................................................

(ii) Explain why beryllium has a higher first ionisation energy than boron.

...........................................................................................

...........................................................................................

...........................................................................................

(iii) Explain why the first ionisation energy of krypton is greater than the first ionisation energy of bromine.

...........................................................................................

...........................................................................................

(iv) Explain why the second ionisation energy of beryllium is greater than the first ionisation energy.

...........................................................................................

...........................................................................................

...........................................................................................(8)

(Total 11 marks)ELECTRONIC STRUCTURE

i) Energy levels

Electrons do not orbit the nucleus randomly; they occupy certain fixed energy levels. Each atom has its own unique set of energy levels, which are difficult to calculate but which depend on the number of protons and electrons in the atom.

Energy levels in an atom can be numbered 1,2,3,…. To infinity. 1 is the lowest energy level (closest to the nucleus) and energy level infinity corresponds to the energy of an electron which is not attracted to the nucleus at all. The energy levels thus converge as they approach infinity:

37

ii) Orbitals and sub-levels

Electrons do not in fact orbit the nucleus in an orderly way. In fact they occupy areas of space known as orbitals. The exact position of an electron within an orbital is impossible to imagine; an orbital is simply an area of space in which there is a high probability of finding an electron.

Orbitals can have a number of different shapes, the most common of which are as follows:

s-orbitals: these are spherical.

Every energy level contains one s-orbital.An s-orbital in the first energy level is a 1s orbital.An s-orbital in the second energy level is a 2s orbital, etc

p-orbitals: these are shaped like a 3D figure of eight. They exist in groups of three:

Every energy level except the first level contains three p-orbitals. Each p-orbital in the same energy level has the same energy but different orientations: x, y and z.A p-orbital in the second energy level is a 2p orbital (2px, 2py, 2pz)A p-orbital in the third energy level is a 3p orbital (3px, 3py, 3pz), etc

38

In addition, the third and subsequent energy levels each contain five d-orbitals, the fourth and subsequent energy levels contain seven f-orbitals and so on. Each type of orbital has its own characteristic shape.

S, p and d orbitals do not all have the same energy. In any given energy level, s-orbitals have the lowest energy and the energy of the other orbitals increases in the order p < d < f etc. Thus each energy level must be divided into a number of different sub-levels, each of which has a slightly different energy.

The number and type of orbitals in each energy level can thus be summarised as follows: Energy level Number and type of orbital

1st sub-level

2nd sub-level

3rd sub-level

4th sub-level

5th sub-level

1 1 x 1s2 1 x 2s 3 x 2p3 1 x 3s 3 x 3p 5 x 3d4 1 x 4s 3 x 4p 5 x 4d 7 x 4f5 1 x 5s 3 x 5p 5 x 5d 7 x 5f 9 x 5g

iii) Shells

Since the different sub-levels have different energies, and the energies of the different levels get closer together with increasing energy level number, the high energy sub-levels of some energy levels soon overlap with the low energy sub-levels of higher energy levels, resulting in a more complex energy level diagram:

Starting with the lowest energy, the orbitals can thus be arranged as follows:

39

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

Many of these sub-levels have similar energy, and can be grouped together.A collection of sub-levels of similar energy is called a shell.

1s│2s 2p│3s 3p │ 4s 3d 4p │5s 4d 5p│6s 4f 5d 6p

The arrangement of shells and the maximum number of electrons in each can be summarised as follows:

Shell number Orbitals in shell

1 1 x1s2 1 x 2s, 3 x 2p3 1 x 3s, 3 x 3p4 1 x 4s, 5 x 3d, 3 x 4p5 1 x 5s, 5 x 4d, 3 x 5p6 1 x 6s, 7 x 4f, 5 x 5d, 3 x 6p

iv) Electrons

Electrons repel each other. In a small space such as an orbital, it is impossible to put more than two electrons.

Since electrons are charged particles, and moving charges create a magnetic field, it is possible to create a small magnetic attraction between two electrons if they are spinning in opposite directions in the same orbital. This is the reason two electrons, and not one, are permitted in the same orbital.

It is thus possible to calculate the maximum possible number of electrons in each sub-level, and thus in each energy level:

Shell Number of electrons in each sub-level Max. no of electrons1 2 x 1s 22 2 x 2s, 6 x 2p 83 2 x 3s, 6 x 3p 84 2 x 4s, 10 x 3d, 6 x 4p 185 2 x 5s, 10 x 4d, 6 x 5p 186 2 x 6s, 14 x 4f, 10 x 5d, 6 x 6p 32

v) Electron arrangement in orbitals

There are three rules which determine the way in which electrons fill the orbitals

1. Aufbau/building principle: electrons always fill the lowest energy orbitals first.

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2. Hund's rule: electrons never pair up in the same orbital until all orbitals of the same energy are singly occupied, and all unpaired electrons have parallel spin.

3. Pauli exclusion principle: only two electrons may occupy the same orbital, and they must do so with opposite spin.

The arrangement of electrons in an atom is known as its electronic configuration. It can be represented in two ways:

The arrow and box method represents each orbital as a box and each electron as an arrow. The direction of spin is shown by the orientation of the arrow.

41

The electronic configuration of the first 18 elements using the arrow in box method is as follows:

1s 2s 2p 3s 3pH ↑

He ↑↓

Li ↑↓ ↑

Be ↑↓ ↑↓

B ↑↓ ↑↓ ↑

C ↑↓ ↑↓ ↑ ↑

N ↑↓ ↑↓ ↑ ↑ ↑

O ↑↓ ↑↓ ↑↓ ↑ ↑

F ↑↓ ↑↓ ↑↓ ↑↓ ↑

Ne ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

Na ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

Mg ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

Al ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

Si ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑

P ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑

S ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑

Cl ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

Ar ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

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The orbital method indicates the number of electrons in each orbital with a superscript written immediately after the orbital.

The electronic configurations of the first eighteen elements can be shown with the orbital method as follows:

H: 1s1

He: 1s2

Li: 1s22s1

Be: 1s22s2

B: 1s22s22p1

C: 1s22s22p2 or 1s22s22p63s23px13py

1

N: 1s22s22p3 or 1s22s22p63s23px13py

13pz1

O: 1s22s22p4 or 1s22s22p63s23p23px23py

13pz1

F: 1s22s22p5

Ne: 1s22s22p6

Na: 1s22s22p63s1

Mg: 1s22s22p63s2

Al: 1s22s22p63s23p1

Si: 1s22s22p63s23p2 or 1s22s22p63s23px13py

1

P: 1s22s22p63s23p3 or 1s22s22p63s23px13py

13pz1

S: 1s22s22p63s23p4 or 1s22s22p63s23px23py

13pz1

Cl: 1s22s22p63s23p5

Ar: 1s22s22p63s23p6

A shorthand form is often used for both the above methods. Full shells are not written in full but represented by the symbol of the element to which they correspond, written in square brackets.

Eg. 1s22s22p6 is represented as [Ne] and 1s22s22p63s23p6 is represented as [Ar].

The shorthand electronic configuration of the elements with atomic numbers 18 to 36 can be written as follows:

43

4s 3d 4pK [Ar] ↑

Ca [Ar] ↑↓

Sc [Ar] ↑↓ ↑

Ti [Ar] ↑↓ ↑ ↑

V [Ar] ↑↓ ↑ ↑ ↑

Cr [Ar] ↑ ↑ ↑ ↑ ↑ ↑

Mn [Ar] ↑↓ ↑ ↑ ↑ ↑ ↑

Fe [Ar] ↑↓ ↑↓ ↑ ↑ ↑ ↑

Co [Ar] ↑↓ ↑↓ ↑↓ ↑ ↑ ↑

Ni [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑

Cu [Ar] ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

Zn [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

Ga [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

Ge [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑

As [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑

Se [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑

Br [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

Kr [Ar] ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

44

Note the unusual structures of chromium and copper. The difference in energy between the 3d and 4s electrons is very small, and in chromium the energy required to promote and electron from 4s to 3d is recovered in the reduced repulsion which results from the fact that they are no longer paired. Thus the 4s13d5 structure in Cr is preferred.In copper, the 3d orbitals are actually lower in energy than the 4s orbital, so the 4s13d10 structure in Cu is preferred.

v) Electron arrangement in ions

The electronic configuration of ions can be deduced by simply adding or removing the appropriate number of electrons. The order in which electrons are to be removed can be deduced from the following rules:

- remove outer shell electrons first- remove p-electrons first, then s-electrons and then d-electrons- remove paired electrons before unpaired electrons in the same sub-level

vi) Effect of electronic configuration on chemical properties

The chemical properties of an atom depend on the strength of the attraction between the outer electrons and the nucleus. These in turn depend on the number of protons and on the electronic configuration, and so it follows that these two factors are instrumental in determining the chemical properties of an atom.

This is in contrast with the neutron number however, which has no effect on the chemical properties of an atom. Neutrons have no charge and hence exert no attractive force on the nucleus.

Isotopes, therefore, tend to have very similar chemical properties since they have the same atomic number and the same electronic configuration. They differ only in number of neutrons, which do not directly influence the chemical properties of an element.

45

IONISATION ENERGIES

i) First ionisation energy

The first ionisation energy of an element is the energy required to remove one electron from each of a mole of free gaseous atoms of that element.

It can also be described as the energy change per mole for the process:M(g) M+(g) + e

The amount of energy required to remove an electron from an atom depends on the number of protons in the nucleus of the atom and on the electronic configuration of that atom.

The first ionisation energies of the first 20 elements in the periodic table is shown below:

Variation of first ionisation energy with atomic number for the first twenty elements

0

500

1000

1500

2000

2500

0 5 10 15 20

atomic number

first

ioni

satio

n en

ergy

(kJ

per m

ole)

There are various trends in this graph which can be explained by reference to the proton number and electronic configuration of the various elements. A number of factors must be considered:

- Energy is required to remove electrons from atoms in order to overcome their attraction to the nucleus. The greater the number of protons, the greater the attraction of the electrons to the nucleus and the harder it is to remove the electrons. The number of protons in the nucleus is known as the nuclear charge.

- The effect of this nuclear charge, however, is cancelled out to some extent by the other electrons in the atom. Each inner shell and inner sub-shell electron effectively cancels out one unit of charge from the nucleus. This is known as shielding.- The outermost electrons in the atom thus only feel the residual positive charge after all inner shell and inner sub-shell electrons have cancelled out much of the nuclear charge. This residual positive charge is known as the effective nuclear charge.

- Electrons repel each other, particularly when they are in the same orbital. The degree of repulsion between the outermost electrons affects the ease with which electrons can be moved.

46

When considering trends in ionisation energies, it is thus necessary to consider 4 factors:

- nuclear charge

- shielding

- effective nuclear charge

- electron repulsion

The trends in first ionisation energies amongst elements in the periodic table can be explained on the basis of variations in one of the four above factors.

Trend across period 1

Compare the first ionisation energies of H and He. Neither have inner shells, so there is no shielding. He has two protons in the nucleus; H only has one. Therefore the helium electrons are more strongly attracted to the nucleus and hence more difficult to remove.

The first ionisation energy of He is thus higher than that of H.

Since H and He are the only atoms whose outer electrons are not shielded from the nucleus, it follows that He has the highest first ionisation energy of all the elements. All elements (except H) have outer electrons which are shielded to some extent from the nucleus and thus are easier to remove.So Helium has the highest first ionisation energy of all the elements.

Trends across period 2

Compare now the first ionisation energies of He (1s2) and Li (1s22s1). Li has an extra proton in the nucleus (3) but two inner-shell electrons. These inner-shell electrons cancel out the charge of two of the protons, reducing the effective nuclear charge on the 2s electron to +1. This is lower than the effective nuclear charge on the He 1s electrons, +2, and so the electrons are less strongly held and easier to remove.

The first ionisation energy of Li is thus lower than that of He.

Compare the first ionisation energies of Li (1s22s1) and Be (1s22s2). Be has one more proton in the nucleus than Li, and no extra inner-shell electrons, so the effective nuclear charge on Be is higher and the Be electrons are more strongly attracted to the nucleus.

The first ionisation energy of Be is thus higher than that of Li.In general, the first ionisation energy increases across a period because the nuclear charge increases but the shielding remains the same.

Compare the first ionisation energies of Be (1s22s2) and B (1s22s22p1).B has one more proton in the nucleus than Be but there are also 2 extra inner sub-shell electrons. These cancel out the charge of two more of the protons, leaving an effective nuclear charge of only +1. This is less than Be (+2) so the electrons are less strongly attracted to the nucleus and thus less difficult to remove.

47

The first ionisation energy of B is thus lower than that of Be.Ionisation energies decrease from group II to group III because in group III the electrons are removed from a p-orbital, so it is shielded by the s-electrons in the outer shell. Thus the effective nuclear charge decreases.

From B (1s22s22p1) to N (1s22s22p3) the proton number increases, but the number of electrons shielding the nuclear charge remains the same at 4. Thus the effective nuclear charge increases from B to N and the electrons become progressively harder to remove.

The first ionisation energy thus increases from B to N.

So far the concepts of effective nuclear charge and shielding have been used to explain the trend in first ionisation energies for the first 7 elements. They cannot, however, explain the fall between N and O. The electronic configurations of N and O must be considered more carefully:

1s 2s 2pN ↑↓ ↑↓ ↑ ↑ ↑

O ↑↓ ↑↓ ↑↓ ↑ ↑

Note that in N the electron is removed from an unpaired orbital, but in O it is removed from a paired orbital. In a paired orbital, the two electrons share a confined space and so repel each other. They are therefore less stable and easier to remove. This repulsion effect outweighs the higher effective nuclear charge in O.

The first ionisation energy of O is thus lower than that of N.First ionisation energies decrease from group V to group VI, since the electron removed from the group VI atom is paired, so there is more repulsion between the electrons and the electron is easier to remove.

The first ionisation energies increase as expected from O to Ne, due to the increase in effective nuclear charge.

------------------------

The trend in first ionisation energies across period 2 can thus be summarised as follows:

1. There is a general increase across the period as the nuclear charge increases and the shielding remains the same.

2. There is a drop from Be to B because in B a 2p electron is being removed and the extra shielding from the 2s subshell actually causes a fall in the effective nuclear charge.

3. There is also a drop from N to O because the electron in O is being removed from a paired orbital. The repulsion of the electrons in this orbital makes them less stable and easier to remove.

The same trend can also be found in Period 3 (Na - Ar). There is a general increase, but a drop between Mg and Al and also between P and S.

------------------

48

Trend down a group

The above graph also shows a clear decrease in first ionisation energy on descending a group. This can be explained in the following way:

On descending a group, the effective nuclear charge stays the same but the number of inner shells increases. The repulsion between these inner shells and the outer electrons makes them less stable, pushes them further from the nucleus and makes them easier to remove.

ii) Successive ionisation energies

The second ionisation energy of an atom is the energy required to remove one electron from each of a mole of free gaseous unipositive ions.

M+(g) M2+(g) + e

Other ionisation energies can be defined in the same way:

The third ionisation energy of an atom is the energy required to remove one electron from each of a mole of bipositive ions.

M2+(g) M3+(g) + e

The nth ionisation energy can be defined as the energy required for the process

M(n-1)+(g) Mn+(g) + e

It always becomes progressively more difficult to remove successive electrons from an atom; the second ionisation energy is always greater than the first, the third always greater than the second and so on. There are two satisfactory explanations for this:

As more electrons are removed from an atom, the number of electrons remaining in the atom decreases. The repulsion between these electrons therefore decreases, while the number of protons remains the same. The remaining electrons are thus more stable and increasingly difficult to remove.

The difference in successive ionisation energies, however, varies widely and depends on the electronic configuration of the atom in question. The difference in successive ionisation energies of an atom can be predicted qualitatively by consideration of the effective nuclear charge on the electron to be removed and the shielding of that electron by the inner shell and inner sub-shell electrons.

Consider the successive ionisation energies of aluminium, 1s22s22p63s23p1:

The 1st ionisation energy is fairly low because the 3p electron is shielded by all the other electrons, and the effective nuclear charge is only +1.

The 2nd and 3rd ionisation energies are significantly higher than the 1st because 3s electrons are being removed and the effective nuclear charge on these electrons is +3.

49

1st: 578 kJmol-1, 2nd: 1817 kJmol-1, 3rd: 2745 kJmol-1

There is a huge jump to the 4th ionisation energy, since a 2p electron is now being removed. The shielding has fallen and the effective nuclear charge has risen to +9.The 5th and 6th ionisation energies are also high.

4th: 11578 kJmol-1, 5th: 14831 kJmol-1, 6th: 18378 kJmol-1

There is another significant jump to the 7th ionisation energy, since an unpaired 2p electron is now being removed.

7th: 23296 kJmol-1, 8th: 27460 kJmol-1, 9th: 31862 kJmol-1

The next significant jump is between the 9th and 10th ionisation energies, since the 10th requires the removal of a 2s electron.

10th: 38458kJmol-1, 11th: 42655 kJmol-1

There is a huge jump to the12th ionisation energy, since a 1s electron is now being removed.

12th: 201276kJmol-1, 13th: 222313kJmol-1.

These ionisation energies could be plotted on a graph as follows:

Variation of ionisation energy with number of ionisations for aluminium

100

1000

10000

100000

1000000

1 2 3 4 5 6 7 8 9 10 11 12 13

number of ionisations

ioni

satio

n en

ergy

Note that the largest jumps by far occur between the 3rd and 4th ionisation energies, and between the 11th and 12th ionisation energies. In practice only large jumps such as this are visible on such a graph.The relative values of successive ionisation energies are therefore a direct indicator of the electronic configuration of the atom in question.

50

The trends can be summarised as follows:

1. The successive ionisation energies of an atom always increase. The more electrons that are removed, the fewer the number electrons that remain. There is therefore less repulsion between the electrons in the resulting ion. The electrons are therefore more stable and harder to remove.

2. By far the largest jumps between successive ionisation energies come when the electron is removed from an inner shell. This causes a large drop in shielding, a large increase in effective nuclear charge and a large increase in ionisation energy

By applying the above principles in reverse, it is also possible to predict the electronic structure of a species by analysis of the successive ionisation energy data:

Eg Si:

Variation of ionisation energy with number of ionisations in silicon

100

1000

10000

100000

1000000

1 3 5 7 9 11 13

number of ionisations

ioni

satio

n en

ergy

Large jumps occur between 4th and 5th and between 12th and 13th.

Therefore there are three shells: The first contains 2 electrons, the second 8 and the third 4.

ELECTRONIC STRUCTURESThis page explores how you write electronic structures for atoms using s, p, and d notation. It assumes that you know about simple atomic orbitals - at least as far as the way they are named, and their relative energies.

The electronic structures of atoms

Relating orbital filling to the Periodic Table

51

The first period

Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2.

The second period

Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s22s1. Beryllium adds a second electron to this same level - 1s22s2.

Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first.

B 1s22s22px1

C 1s22s22px12py

1

N 1s22s22px12py

12pz1

Note:  The orbitals where something new is happening are shown in bold type. You wouldn't normally write them any differently from the other orbitals.

The next electrons to go in will have to pair up with those already there.

O 1s22s22px22py

12pz1

F 1s22s22px22py

22pz1

Ne 1s22s22px22py

22pz2

You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases. There are two ways around this, and you must be familiar with both.

Shortcut 1: All the various p electrons can be lumped together. For example, fluorine could be written as 1s22s22p5, and neon as 1s22s22p6.

This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level (those on the outside of the atom), they are sometimes written in shorthand, sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for the electronic structure of something in an exam, write it out in full showing all the px, py and pz orbitals in the outer level separately.

52

For example, although we haven't yet met the electronic structure of chlorine, you could write it as 1s22s22p63s23px

23py23pz

1.

Notice that the 2p electrons are all lumped together whereas the 3p ones are shown in full. The logic is that the 3p electrons will be involved in bonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really of any interest.

Shortcut 2: You can lump all the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means the electronic structure of neon - in other words: 1s22s22px

22py22pz

2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s2.

On this basis the structure of chlorine would be written [Ne]3s23px23py

23pz1.

The third period

At neon, all the second level orbitals are full, and so after this we have to start the third period with sodium. The pattern of filling is now exactly the same as in the previous period, except that everything is now happening at the 3-level.

For example:

short versionMg 1s22s22p63s2 [Ne]3s2

S 1s22s22p63s23px23py

13pz1 [Ne]3s23px

23py13pz

1

Ar 1s22s22p63s23px23py

23pz2 [Ne]3s23px

23py23pz

2

The beginning of the fourth period

At this point the 3-level orbitals aren't all full - the 3d levels haven't been used yet. But if you refer back to the energies of the orbitals, you will see that the next lowest energy orbital is the 4s - so that fills next.

K 1s22s22p63s23p64s1

Ca 1s22s22p63s23p64s2

There is strong evidence for this in the similarities in the chemistry of elements like sodium (1s22s22p63s1) and potassium (1s22s22p63s23p64s1)

The outer electron governs their properties and that electron is in the same sort of orbital in both of the elements. That wouldn't be true if the outer electron in potassium was 3d1.

s- and p-block elements

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The elements in group 1 of the Periodic Table all have an outer electronic structure of ns1 (where n is a number between 2 and 7). All group 2 elements have an outer electronic structure of ns2. Elements in groups 1 and 2 are described as s-block elements.

Elements from group 3 across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements.

d-block elements

Remember that the 4s orbital has a lower energy than the 3d orbitals and so fills first. Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect.

d-block elements are elements in which the last electron to be added to the atom is in a d orbital. The first series of these contains the elements from scandium to zinc, which at GCSE you probably called transition elements or transition metals. The terms "transition element" and "d-block element" don't quite have the same meaning, but it doesn't matter in the present context.

If you are interested:  A transition element is defined as one which has partially filled d orbitals either in the element or any of its compounds. Zinc (at the right-hand end of the d-block) always has a completely full 3d level (3d10) and so doesn't count as a transition element.

d electrons are almost always described as, for example, d5 or d8 - and not written as separate orbitals. Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up.

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d5 means

d8 means

Notice in what follows that all the 3-level orbitals are written together, even though the 3d electrons are added to the atom after the 4s.

Sc 1s22s22p63s23p63d14s2

Ti 1s22s22p63s23p63d24s2

V 1s22s22p63s23p63d34s2

Cr 1s22s22p63s23p63d54s1

Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one electron in each orbital. It would be convenient if the sequence was tidy - but it's not!

Mn 1s22s22p63s23p63d54s2 (back to being tidy again)Fe 1s22s22p63s23p63d64s2

Co 1s22s22p63s23p63d74s2

Ni 1s22s22p63s23p63d84s2

Cu 1s22s22p63s23p63d104s1 (another awkward one!)Zn 1s22s22p63s23p63d104s2

And at zinc the process of filling the d orbitals is complete.

Filling the rest of period 4The next orbitals to be used are the 4p, and these fill in exactly the same way as the 2p or 3p. We are back now with the p-block elements from gallium to krypton. Bromine, for example, is 1s22s22p63s23p63d104s24px

24py24pz

1.

SummaryWriting the electronic structure of an element from hydrogen to krypton

Use the Periodic Table to find the atomic number, and hence number of electrons. Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p - until you run out of

electrons. The 3d is the awkward one - remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up.

Remember that chromium and copper have electronic structures which break the pattern in the first row of the d-block.

Writing the electronic structure of big s- or p-block elements

First work out the number of outer electrons. This is quite likely all you will be asked to do anyway.

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The number of outer electrons is the same as the group number. (The noble gases are a bit of a problem here, because they are normally called group 0 rather then group 8. Helium has 2 outer electrons; the rest have 8.) All elements in group 3, for example, have 3 electrons in their outer level. Fit these electrons into s and p orbitals as necessary. Which level orbitals? Count the periods in the Periodic Table (not forgetting the one with H and He in it).

Iodine is in group 7 and so has 7 outer electrons. It is in the fifth period and so its electrons will be in 5s and 5p orbitals. Iodine has the outer structure 5s25px

25py25pz

1.

What about the inner electrons if you need to work them out as well? The 1, 2 and 3 levels will all be full, and so will the 4s, 4p and 4d. The 4f levels don't fill until after anything you will be asked about at A'level. Just forget about them! That gives the full structure: 1s22s22p63s23p63d104s24p64d105s25px

25py25pz

1.

When you've finished, count all the electrons to make sure that they come to the same as the atomic number. Don't forget to make this check - it's easy to miss an orbital out when it gets this complicated.

Barium is in group 2 and so has 2 outer electrons. It is in the sixth period. Barium has the outer structure 6s2.

Including all the inner levels: 1s22s22p63s23p63d104s24p64d105s25p66s2.

It would be easy to include 5d10 as well by mistake, but the d level always fills after the next s level - so 5d fills after 6s just as 3d fills after 4s. As long as you counted the number of electrons you could easily spot this mistake because you would have 10 too many.

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The 4 rules:

Electrons enter and fill orbitals according to four rules:

Pauli Exclusion Principal Orbitals can contain a maximum of two electrons which must be of opposite spin.

Aufbau or Build-up Principle

Electrons enter and fill lower energy orbitals before higher energy orbitals.

Hund's Rule When there there are degenerate (equal energy) orbitals available, electrons will enter the orbitals one-at-a-time to maximise degeneracy, and only when all the orbitals are half filled will pairing-up occur. This is the rule of maximum multiplicity.

Madelung's Rule Orbitals fill with electrons as n + l, where n is the principal quantum number and l is the subsidiary quantum number. This rule 'explains' why the 4s orbital has a lower energy than the 3d orbital, and it gives the periodic table its characteristic appearance.

Certain 'magic' numbers of electrons of electrons exhibit energetic stability: 2, 10, 18, 36, 54, and 86 are associated with the Group 18 inert or noble gases: He, Ne, Ar, Kr, Xe, and Rn.

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ELECTRONIC STRUCTURES OF IONSWorking out the electronic structures of ions

Ions are atoms (or groups of atoms) which carry an electric charge because they have either gained or lost one or more electrons. If an atom gains electrons it acquires a negative charge. If it loses electrons, it becomes positively charged.

The electronic structure of s- and p-block ions

Write the electronic structure for the neutral atom, and then add (for a negative ion) or subtract electrons (for a positive ion).

To write the electronic structure for Cl -:

Cl 1s22s22p63s23px23py

23pz1 but Cl- has one more electron

Cl- 1s22s22p63s23px23py

23pz2

To write the electronic structure for O2-:

O 1s22s22px22py

12pz1 but O2- has two more electrons

O2- 1s22s22px22py

22pz2

To write the electronic structure for Na+:

Na 1s22s22p63s1 but Na+ has one less electronNa+ 1s22s22p6

To write the electronic structure for Ca2+:

Ca 1s22s22p63s23p64s2 but Ca2+ has two less electronsCa2+ 1s22s22p63s23p6

The electronic structure of d-block ions

Here you are faced with one of the most irritating facts in A'level chemistry! You will recall that the first transition series (from scandium to zinc) is the result of the 3d orbitals being filled after the 4s orbital.

However, once the electrons are established in their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital. The reversed order of the 3d and 4s orbitals only applies to building the atom up in the first place. In all other respects, the 4s electrons are always the electrons you need to think about first.

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You must remember this:

When d-block elements form ions, the 4s electrons are lost first.

Provided you remember that, working out the structure of a d-block ion is no different from working out the structure of, say, a sodium ion.

To write the electronic structure for Cr3+:

Cr 1s22s22p63s23p63d54s1

Cr3+ 1s22s22p63s23p63d3

The 4s electron is lost first followed by two of the 3d electrons.

To write the electronic structure for Zn2+:

Zn 1s22s22p63s23p63d104s2

Zn2+ 1s22s22p63s23p63d10

This time there is no need to use any of the 3d electrons.

To write the electronic structure for Fe3+:

Fe 1s22s22p63s23p63d64s2

Fe3+ 1s22s22p63s23p63d5

The 4s electrons are lost first followed by one of the 3d electrons.

The rule is quite simple. Take the 4s electrons off first, and then as many 3d electrons as necessary to produce the correct positive charge.

IONISATION ENERGYDefining first ionisation energy

Definition

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The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

This is more easily seen in symbol terms.

It is the energy needed to carry out this change per mole of X.

Things to notice about the equation

The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.

Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).

All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.

Patterns of first ionisation energies in the Periodic Table

The first 20 elements

First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.

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Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.

The size of that attraction will be governed by:

The charge on the nucleus.

The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.

The distance of the electron from the nucleus.

Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

The number of electrons between the outer electrons and the nucleus.

Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!)

If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

Whether the electron is on its own in an orbital or paired with another electron.

Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

Explaining the pattern in the first few elements

Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).

Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.

Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the

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nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.

You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).

If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.

The patterns in periods 2 and 3

Talking through the next 17 atoms one at a time would take ages. We can do it much more neatly by explaining the main trends in these periods, and then accounting for the exceptions to these trends.

The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies in period 3 are all lower than those in period 2.

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Explaining the general trend across periods 2 and 3

The general trend is for ionisation energies to increase across a period.

In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p. These are all the same sort of distances from the nucleus, and are screened by the same 1s2 electrons.

The major difference is the increasing number of protons in the nucleus as you go from lithium to neon. That causes greater attraction between the nucleus and the electrons and so increases the ionisation energies. In fact the increasing nuclear charge also drags the outer electrons in closer to the nucleus. That increases ionisation energies still more as you go across the period.

In period 3, the trend is exactly the same. This time, all the electrons being removed are in the third level and are screened by the 1s22s22p6 electrons. They all have the same sort of environment, but there is an increasing nuclear charge.

Why the drop between groups 2 and 3 (Be-B and Mg-Al)?

The explanation lies with the structures of boron and aluminium. The outer electron is removed more easily from these atoms than the general trend in their period would suggest.

Be 1s22s2 1st I.E. = 900 kJ mol-1

B 1s22s22px1 1st I.E. = 799 kJ mol-1

You might expect the boron value to be more than the beryllium value because of the extra proton. Offsetting that is the fact that boron's outer electron is in a 2p orbital rather than a 2s. 2p orbitals have a slightly higher energy than the 2s orbital, and the electron is, on average, to be found further from the nucleus. This has two effects.

The increased distance results in a reduced attraction and so a reduced ionisation energy.

The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well. That also reduces the pull from the nucleus and so lowers the ionisation energy.

The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the 3-level rather than the 2-level.

Mg 1s22s22p63s2 1st I.E. = 736 kJ mol-1

Al 1s22s22p63s23px1 1st I.E. = 577 kJ mol-1

The 3p electron in aluminium is slightly more distant from the nucleus than the 3s, and partially screened by the 3s2 electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

Why the drop between groups 5 and 6 (N-O and P-S)?

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Once again, you might expect the ionisation energy of the group 6 element to be higher than that of group 5 because of the extra proton. What is offsetting it this time?

N 1s22s22px12py

12pz1 1st I.E. = 1400 kJ mol-1

O 1s22s22px22py

12pz1 1st I.E. = 1310 kJ mol-1

The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from an identical orbital.

The difference is that in the oxygen case the electron being removed is one of the 2px2

pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.

The drop in ionisation energy at sulphur is accounted for in the same way.

Trends in ionisation energy down a group

As you go down a group in the Periodic Table ionisation energies generally fall. You have already seen evidence of this in the fact that the ionisation energies in period 3 are all less than those in period 2.

Taking Group 1 as a typical example:

Why is the sodium value less than that of lithium?

There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening.

Li 1s22s1 1st I.E. = 519 kJ mol-1

Na 1s22s22p63s1 1st I.E. = 494 kJ mol-1

Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it. The 2s1 electron feels the pull of 3 protons screened by 2 electrons - a net pull from the centre of 1+.

The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner electrons. The 3s1 electron also feels a net pull of 1+ from the centre of the atom. In other words, the effect of the extra protons is compensated for by the effect of the extra screening electrons. The only factor left is the extra distance

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between the outer electron and the nucleus in sodium's case. That lowers the ionisation energy.

Similar explanations hold as you go down the rest of this group - or, indeed, any other group.

Trends in ionisation energy in a transition series

Apart from zinc at the end, the other ionisation energies are all much the same.

All of these elements have an electronic structure [Ar]3dn4s2 (or 4s1 in the cases of chromium and copper). The electron being lost always comes from the 4s orbital.

As you go from one atom to the next in the series, the number of protons in the nucleus increases, but so also does the number of 3d electrons. The 3d electrons have some screening effect, and the extra proton and the extra 3d electron more or less cancel each other out as far as attraction from the centre of the atom is concerned.

The rise at zinc is easy to explain.

Cu [Ar]3d104s1 1st I.E. = 745 kJ mol-1

Zn [Ar]3d104s2 1st I.E. = 908 kJ mol-1

In each case, the electron is coming from the same orbital, with identical screening, but the zinc has one extra proton in the nucleus and so the attraction is greater. There will be a degree of repulsion between the paired up electrons in the 4s orbital, but in this case it obviously isn't enough to outweigh the effect of the extra proton.

Ionisation energies and reactivity

The lower the ionisation energy, the more easily this change happens:

You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form.

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The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process.

For example, you wouldn't be starting with gaseous atoms; nor would you end up with gaseous positive ions - you would end up with ions in a solid or in solution. The energy changes in these processes also vary from element to element. Ideally you need to consider the whole picture and not just one small part of it.

However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy of the reactions. Remember that activation energy is the minimum energy needed before a reaction will take place. The lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction are.

The fall in ionisation energy as you go down a group will lead to lower activation energies and therefore faster reactions.

SUCCESSIVE IONISATION ENERGIESDefining second ionisation energy

Second ionisation energy is defined by the equation:

It is the energy needed to remove a second electron from each ion in 1 mole of gaseous 1+ ions to give gaseous 2+ ions.

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More ionisation energies

You can then have as many successive ionisation energies as there are electrons in the original atom.

The first four ionisation energies of aluminium, for example, are given by

1st I.E. = 577 kJ mol-1

2nd I.E. = 1820 kJ mol-1

3rd I.E. = 2740 kJ mol-1

4th I.E. = 11600 kJ mol-1

In order to form an Al3+(g) ion from Al(g) you would have to supply:

577 + 1820 + 2740 = 5137 kJ mol-1

That's a lot of energy. Why, then, does aluminium form Al3+ ions?

It can only form them if it can get that energy back from somewhere, and whether that's feasible depends on what it is reacting with.

For example, if aluminium reacts with fluorine or oxygen, it can recover that energy in various changes involving the fluorine or oxygen - and so aluminium fluoride or aluminium oxide contain Al3+ ions.

If it reacts with chlorine, it can't recover sufficient energy, and so solid anhydrous aluminium chloride isn't actually ionic - instead, it forms covalent bonds.

Why doesn't aluminium form an Al4+ ion? The fourth ionisation energy is huge compared with the first three, and there is nothing that aluminium can react with which would enable it to recover that amount of extra energy.

Why do successive ionisation energies get larger?

Once you have removed the first electron you are left with a positive ion. Trying to remove a negative electron from a positive ion is going to be more difficult than removing it from an atom. Removing an electron from a 2+ or 3+ (etc) ion is going to be progressively more difficult.

Why is the fourth ionisation energy of aluminium so large?

The electronic structure of aluminium is 1s22s22p63s23px1. The first three electrons to be

removed are the three electrons in the 3p and 3s orbitals. Once they've gone, the fourth electron is removed from the 2p level - much closer to the nucleus, and only screened by the 1s2 (and to some extent the 2s2) electrons.

Using ionisation energies to work out which group an element is in

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This big jump between two successive ionisation energies is typical of suddenly breaking in to an inner level. You can use this to work out which group of the Periodic Table an element is in from its successive ionisation energies.

Magnesium (1s22s22p63s2) is in group 2 of the Periodic Table and has successive ionisation energies:

Here the big jump occurs after the second ionisation energy. It means that there are 2 electrons which are relatively easy to remove (the 3s2 electrons), while the third one is much more difficult (because it comes from an inner level - closer to the nucleus and with less screening).

Silicon (1s22s22p63s23px13py

1) is in group 4 of the Periodic Table and has successive ionisation energies:

Here the big jump comes after the fourth electron has been removed. The first 4 electrons are coming from the 3-level orbitals; the fifth from the 2-level.

The lesson from all this:

Count the easy electrons - those up to (but not including) the big jump. That is the same as the group number.

Another example:

Decide which group an atom is in if it has successive ionisation energies:

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The ionisation energies are going up one or two thousand at a time for the first five. Then there is a huge jump of about 15000. There are 5 relatively easy electrons - so the element is in group 5.

Exploring the patterns in more detail

If you plot graphs of successive ionisation energies for a particular element, you can see the fluctuations in it caused by the different electrons being removed.

Not only can you see the big jumps in ionisation energy when an electron comes from an inner level, but you can also see the minor fluctuations within a level depending on whether the electron is coming from an s or a p orbital, and even whether it is paired or unpaired in that orbital.

Chlorine has the electronic structure 1s22s22p63s23px23py

23pz1.

This graph plots the first eight ionisation energies of chlorine. The green labels show which electron is being removed for each of the ionisation energies.

If you put a ruler on the first and second points to establish the trend, you'll find that the third, fourth and fifth points lie above the value you would expect. That is because the first two electrons are coming from pairs in the 3p levels and are therefore rather easier to remove than if they were unpaired.

Again, if you put a ruler on the 3rd, 4th and 5th points to establish their trend, you'll find that the 6th and 7th points lie well above the values you would expect from a continuation of the trend. That is because the 6th and 7th electrons are coming from the 3s level - slightly closer to the nucleus and slightly less well screened.

The massive jump as you break into the inner level at the 8th electron is fairly obvious!

To plot any more ionisation energies for chlorine needs a change of vertical scale. The seventeenth ionisation energy of chlorine is nearly 400,000 kJ mol-1, and the vertical scale has to be squashed to accommodate this.

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This is now a "log graph" - plotted by finding the logarithm of each ionisation energy (press the "log" button on your calculator). This doesn't simply squash the vertical scale. It distorts it as well, to such an extent that the only useful thing the graph now shows is the major jumps where the next electron to be removed comes from an inner level. The distortion is so great in the first 8 ionisation energies, for example, that the patterns shown by the previous graph are completely (and misleadingly) destroyed.

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