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Lesson objectives • Define first ionisation energy and successive ionisation energy.

1.2.1 Evidence for shells 28-Oct-14. Lesson objectives • Define first ionisation energy and successive ionisation energy. • Explain the factors that influence ionisation energies.

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Lesson objectives • Define first ionisation energy and successive ionisation energy.

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  1. 1.2.1 Evidence for shells28-Oct-14 • Lesson objectives • Define first ionisation energy and successive ionisation energy. • Explain the factors that influence ionisation energies. • Predict the number of electrons in each shell as well as the element’s group, using successive ionisation energies.

  2. Evidence for shells Using your knowledge from GCSE draw the electronic structures for the following atoms: Calcium Chlorine Aluminium Sodium

  3. Ions • Draw electronic configurations for the ions Ca2+ Al3+ Cl- O2-

  4. Forming Ions The first ionisation energy (1st I.E.) of an element is the amount of energy required to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous 1+ ions. X (g) X+(g) + e- Example Na (g) Na+(g) + e- 1st I.E. = + 496 kJ mol-1 Don't forget state symbols

  5. This is what happens inside a plasma TV screen • The screen consists of 100’s of 1000’s of tiny cells which each contain a mixture or neon and xenon between two plates of glass. • The gas is electrically ionised into a mixture of +ve ions and –ve electrons. • The formation of these ions required energy . The mixture is called a plasma which causes the screen to emit light Xe (g) Xe+(g) + e-

  6. Forming Ions Consider the atomic structure and discuss in pairs what factors you think will affect ionisation energies.

  7. Ionisation energy is affected by:1 - Atomic Radius Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

  8. 2 - Nuclear Charge Helium Hydrogen 1st I.E = 1310 KJ mol-1 1st I.E = 2370 KJ mol-1 Why is the 1st IE greater for helium than for hydrogen? The greater the nuclear charge, the greater the attraction of the nucleus for the outer shell electron, therefore it requires more energy to remove it.

  9. 3 - Electron Shielding Sodium Potassium 1st I.E = 494 KJ mol-1 1st I.E = 418 KJ mol-1 There is an increase in nuclear charge from Na to K. However, for potassium there is an additional shell of electrons which shield the outer electron from the attraction of the nucleus. So the attraction of the nucleus for the outer shell electron is less, so is held less strongly. This is called electron shielding.

  10. Key definition • Electron shielding is the repulsion between electrons in different inner shells. Shielding reduces the net attractive force from the positive nucleus on the outer-shell electrons.

  11. True or false? • I played rugby at University and in my final year was awarded half colours (Palatinate – Durham) • One of my hobbies is breeding pythons. Currently I have 4 rock pythons, 10 eggs incubating which are due to hatch at the end of September

  12. Successive Ionisation Energies This is a measure of the energy required to remove each electron in turn. Example, Li (g) Li + (g) + e- Li + (g) Li 2+ (g) + e- Li 2+ (g) Li 3+ (g) + e- 1st I.E. = + 520 kJ mol-1 2nd I.E = +7298 kJ mol-1 3rd I.E. = +11815 kJ mol-1 Why does it take more each energy for each successive ionisation?

  13. Three successive ionisation energies of lithium

  14. Successive ionisation energies of nitrogen

  15. Write an equation to represent the 4th ionisation energy of chlorine • The successive ionisation energies for carbon are shown below. Draw its electronic configuration and use this information to explain any trends shown. 1st = 109 2nd = 235 3rd = 462 4th = 622 5th = 3780 6th = 4730 (to the nearest 10 kJ mol-1) 3. The graph below shows the successive ionisation energies of sodium, explain the trends shown. 4. How does the graph confirms the suggested simple electronic configuration for sodium of (2,8,1)

  16. Lesson Objectives • State the number of electrons that can fill the first four shells of an atom. • Define an orbital. • Describe the shapes of s- and p-orbitals.

  17. At GCSE : electrons in shells Lowest shell holds ____ electrons Higher levels hold ____ electrons Fill from the centre outwards At A level . . . • Electrons travel far from the nucleus, but there are main areas where they are commonly found. These are the principal shells and have distinct energy levels (represented by rings at GCSE). The innermost ring is level 1 and contains 2 electrons. The next level contains 8 electrons and so on. The levels are given numbers 1, 2, 3 etc which are also known as quantum numbers.

  18. Energy levels or shells • Electrons are constantly moving, and it is impossible to know the exact position of an electron at any given time. However, measurements of the density of electrons as they move around the nucleus show us there are areas where it is highly probable to find an electron. These regions of high probability are called ORBITALS. Each orbital can hold 2 electrons. • There are 4 different types of orbitals - s, p, d and f and they all have a different shape!!

  19. Key definition • ORBITAL – an atomic orbital is a region within an atom that can hold 2 electrons, with opposite spins

  20. s-orbitals • An s-orbital is spherical in shape • Each shell has an s orbital • This gives a total of 1 x 2 = 2s electrons in each shell

  21. p-orbitals • From n=2 upwards (ie. the second shell), each shell contains 3 p-orbitals, px, py, and pz • This gives a total of 3 x 2 = 6p electrons

  22. d-orbitals and f-orbitals • These are more complex . . . • From n=3 upwards each shell has five d-orbitals. This gives 5 x 2 = 10 d electrons • From n=4 upwards each shell has seven f-orbitals. This gives 7 x 2 = 14 f electrons

  23. Let’s revisit our lesson objectives • State the number of electrons that can fill the first four shells of an atom. • Define an orbital. • Describe the shapes of s- and p-orbitals.

  24. With orbitals having different shapes, chemists often represent an orbital as a box. We call this - Electrons in boxes

  25. Electrons within an orbital will repel each other. An electron has a property called spin and each electron in an orbital will always have opposite spins.

  26. Lesson objectives • State the number of orbitals making up s, p and d sub-shells. • State the number of electrons that occupy s, p and d sub-shells. • Describe the relative energies of s-, p- and d-orbitals for shells 1, 2 and 3. • Deduce the electron configurations of atoms for the first two periods.

  27. Shells, sub-shells and energy levels

  28. Filling the 2p-orbital

  29. • Describe the relative energies of s-, p- and d-orbitals for the shells 1, 2, 3 and of the 4s- and 4p-orbitals. • Deduce the electron configuration of atoms and ions up to Z = 36. • Classify the elements into s-, p- and d-blocks.

  30. Overlap of 4s- and 3d sub-shells

  31. Filling the orbitals in a potassium atom

  32. The Periodic Table, sub-shells and blocks

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