_______ ___ ___ _____ __ ___ ____ ___ _______ ____ ___ _____ ___ ___ __ _________ __ __ __ _________ ____ __________ _________ __ ______ ________ _______ ___ ___ ________ ___ ________ __ __________ Describe how the model of the atom has changed over the years, and how it continues to do so. Understand that scientific knowledge is always evolving. Describe how new theories are accepted by scientists. Week 1 © Pearson Education Ltd 2008 This document may have been altered from the original

Atomic Structure The history: 5BC Democritus: Atoms solid particles held together by hooks. 1800s John Dalton: Atoms cannot be divided. All atoms of a given element are the same. Atoms of 1 element are different from those of every other element.

J.J. Thomson: Atoms made up of smaller particles.‘Plum pudding’ model of the atom. Atoms contain equal numbers of negative and positive particles Rutherford: Nucleus contains positive charge and mass. Negative electrons in orbit.Atom mostly empty space Neils Bohr: Planetary model

1913 Rutherford discovered the proton De Broglie: Wave particle duality 1926 Schrodinger: Atomic orbitals Chadwick: Discovered the neutron 1964 and onwards Gell Mann and Zweig: Quarks – make up the other 3 fundamental particles.

The Atom for Practical Use Today we use the model of the nuclear atom with orbiting electrons in fixed paths as proposed by Niels Bohr in school since it allows us to explain the chemistry we need in relatively simple terms and leave the more complicated theories to universities and particle physics. Each of the previous models has limitations- what are they?

Electrons Electrons in atoms are the key to chemistry. Protons and neutrons give an atom its MASS but do not affect chemical reactions. Only electrons on the outside of atoms are involved in chemical change. If orbiting electrons are attracted to the nucleus of an atom by electrostatic forces the energy required to remove electrons should give information about their arrangement.

_____ _____ __________ ______ ___ __________ __________ ______ ______ ___ _______ ____ _________ __________ ________ ______ ___ ______ __ _________ __ ____ _____ __ ____ __ ___ _______ _ _____ _____ __________ __________ ________ 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. Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Ionisation Energy This is a measure of the energy required to remove electrons from an atom. What will affect the amount of energy needed to remove these electrons? Distance of electron from nucleus (Atomic radius). Nuclear charge. Electron shielding. Shielding reduces the nuclear charge felt by outer electrons and so the term Effective Nuclear Charge can be used to cover both.

Using Ionisation Energy Since each electron in an atom can be removed if enough energy is supplied to overcome electrostatic forces attracting it to the nucleus there are as may values for ionisation energy as there are electrons in an atom. 1 st Ionisation energy: The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. E.g. Na (g) → Na + (g) + e -

Successive Ionisation Energies These are a measure of the energy required to remove EACH electron from an atom IN TURN. Looking at data for any atom shows that the values increase as more electrons are removed. Why? As each electron is removed there is less repulsion between the electrons and those remaining will be pulled closer to the nucleus. As the distance from the nucleus decreases nuclear attraction increases. As electrons are removed the atom becomes an increasingly positive ion so also increasing the effective nuclear charge felt by the remaining electrons.

____ __________ __________ ________ __ _______Three successive ionisation energies of lithium Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Evidence for Shells Successive ionisation energies provide evidence that electrons are arranged in shells. Large jumps (factor of 10) in ionisation energy seem to occur at regular intervals. E.g. the last 2 electrons are always much harder to remove than the previous 8 (if 8 were there) etc. These big jumps are taken to suggest that electrons are being removed from shells closer to the nucleus with less shielding. The data allows us to predict the number of electrons in each shell of an atom and so the group in the Periodic Table in which the element occurs.

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Successive ionisation energies of nitrogen

Ionisation energy/KJ per mole Atomic number

Ionisation Energy Across the Periodic Table Plotting a graph of first ionisation energy for a large number of elements sequentially suggests that there is a periodic pattern in ionisation energy which indicates that electron shells may be more complicated than the GCSE model which successive ionisation energies for a particular atom support. General trend This is UP from left to right across a period. BUT The trend is uneven suggesting some electrons are easier to remove than might be expected. This is taken to mean that there are SUB SHELLS within each main electron shell.

____ ___ ______ __ _________ ____ ___ ____ ___ _____ ____ ______ __ __ ____ _____ __ _______ _______ ___ ______ __ _ ___ _ ________ 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. Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Shells and Orbitals Energy levels or shells Each period in the periodic table represents an electron shell. These are numbered 1-7. These numbers are the PRINCIPAL QUANTUM NUMBERS, n which are used to label the main electron shells. The larger the value of n, the further the shell is from the nucleus and the higher the energy level. Each of the principal quantum shells holds 2n 2 electrons.

Shells and Orbitals nShellElectrons 11 st shell2 22 nd shell8 33 rd shell18 44 th shell32

Sub Shells and Orbitals The energy levels of the principal quantum shells do not have precisely the same energy values for all electrons in that shell. Each contains a set of SUB SHELLS which contain ATOMIC ORBITALS with different energy values. An ORBITAL is a region around the nucleus where there is a 95% probability that an electron may be found. Each orbital has its own approximate 3D shape.

Orbitals and ‘Spin’ An atomic orbital is able to hold up to 2 electrons. These 2 negatively charged electrons repel one another but are still in the same region of space. This is possible because they are each spinning in opposite directions. Each of these spinning electrons behaves as a small magnet and opposite spins has the effect of making magnets with opposite poles.

Orbitals and ‘Spin’ Thus the repulsions are reduced (since N and S magnetic fields actually attract) and ‘spin pairing’ of electrons allows the 2 electrons to occupy the same small region of space. Looking back at the plot of ionisation energy across the periodic table we see that: Each n shell has an orbital which holds 2 electrons, described as an ‘s’ orbital. From n=2 all n shells also have 3 ‘p’ orbitals, each of which holds 2 electrons. The ‘p’ sub shell thus can hold up to 6 electrons

_ _ _______ __ _________ __ _____An s-orbital is spherical in shape Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Three p-orbitals

d and f Orbitals From n=3 upwards each shell has 5 ‘d’ orbitals, each of which can hold 2 electrons The d sub shell thus can contain up to 10 electrons. From n=4 upwards each shell has 7 ‘f’ obitals, each of which can hold 2 electrons. The ‘f’ sub shell thus can contain up to 14 electrons.

Atomic Orbitals

Shell Filling In each successive element of the Periodic Table the order of filling of shells and orbitals is in order of their relative energy. The ELECTRON CONFIGURATION of an atom is the one that gives as low an energy state as possible to the atom as a whole. As the atom increases in size the energies of the orbitals get closer together and so the fill pattern needs to be learned.

Shells, sub-shells and energy levels Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Shell Filling and Electron Configuration As you can see from the previous slide the 4s sub shell fills before the 3d sub shell. The ELECTRON CONFIGURATION for an atom can be deduced using the Aufbau principle. Electrons are added 1 at a time to build up the atom The lowest available energy level must be filled first Each energy level must be filled before the next, higher energy level starts to fill. Sub shells fill singly (parallel spin) before spin pairing begins.

Representing Electron Configurations E.g. hydrogen has 1 electron in an s orbital in a shell with a principle quantum number n=1. We show this as: 1s 1 Helium has 2 electrons, both in the 1s orbital so is shown as 1s 2. Using this method F has 9 electrons arranged: 1s 2 2s 2 2p 5 This method of representing electron configuration is described as sub shell or ‘s,p,d,f’ notation.

Box Notation or Electrons in Boxes This is a convenient way of not only showing where electrons are in an atom, but also how the orbitals are filling. Each box represents an orbital. Each orbital can hold up to 2 electrons with opposite spins. The spin of an electron is shown by an arrow for each electron present and each is given a direction (‘up’ or ‘down’) to show the spin. Boxes are filled in order of energy and p and d orbitals fill in parallel spin then spin pairing order.

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Electrons in boxes

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Electrons within an orbital always have opposite spins

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Filling the 2p-orbital

Elements with more than 18 electrons 4s and 3d shells have similar energy but the 4s is lower. The 4s fills first. When the 4s is filled the repulsions between the electrons in the shell cause it to move out slightly away from the nucleus. The 3d is then a lower energy shell and starts to fill. As it fills repulsions increase and the 4s is repelled further out. When ionisation occurs the 4s shells are the first to be lost. First in; first out.

______ __ _ ___ _ ___ ______Overlap of 4s- and 3d sub-shells Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original Filling the orbitals in a potassium atom

Sub Shells and the Periodic Table The Periodic Table is structured in blocks which show which sub shells are filling for each element in that block. For example the Transition Metals are better described as ‘ d block elements ’ because the d sub shell is filling as we go across the Periodic Table from e.g. Sc to Zn. An elements position in the block gives a very quick and easy way of working out the electron configuration. (See bottom of page 46)

Week 5 © Pearson Education Ltd 2008 This document may have been altered from the original The Periodic Table, sub-shells and blocks

Transition Metal Configuration For Transition Metals the first in; first out rule is most important. See box p. 47 Shell order gives no information about which electrons are the important ones for the chemistry of Transition Metals whereas energy level order makes it clear that the d sub shells determine the characteristics of these metals.

Noble Gas Core For large atoms, particularly Transition Metals electron configuration can be tedious. Noble Gas core removes the need to write out the configuration for full shells and allows us to focus on the outer shell for a particular atom. So Fe should be 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 But is more easily written [Ar] 4s 2 3d 6 Metal ions form positive ions by the removal of electrons from the highest energy orbitals. Non metal ions form negative ions by the addition of electrons to the highest energy orbitals.