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Electronic Structure of Atoms atomic radii, ionisation energies and trends within groups The size of an atom • It is difficult to measure the size of an atom directly. • The exact position of the furthest electron from the nucleus cannot be found! However.. Scientists can measure the length of a covalent bond between two atoms using various methods The distance is .074 nm Atomic radii The atomic radius is half the distance between two atoms of the same element joined together by a single covalent bond. It only works if it is the same atoms.. And a single covalent bond The atomic radii of the elements A pattern in atomic radii Descrease going across a period Ìncrease going down a group 3 11 19 Potassium As you go down a group… 1. the atoms have extra electrons which have to go into a new shell further away from the nucleus. 2. the atoms have a bigger nuclear charge. You might expect that this would pull the electrons closer to the nucleus but the electrons in the extra shells are being shielded from the extra nuclear charge by the inner shells of electrons so this doesn’t happen! “The shielding effect” The atomic radii increases.. 3 Lithium 4 Beryllium 5 6 Aluminium Carbon As you go across a period.. 1. the nuclear charge in the atom increases. The increasing nuclear charge pulls the electrons closer to the nucleus 2. the atoms have extra electrons but they do not go into new shells! There is therefore no extra shielding effect as you go across a period! The atomic radii decreases.. Higher level only Ionisation energy Ionisation energy is the minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom. Check the ionisation energy • What is the pattern • 1) as you go down a group • 2) as you go across a period Higher level only Ionisation energy – the pattern Decreases as you go across any period Increases as you go down any group Higher level only Ionisation energy – explaining the trend Increases as you go down any group 3 11 19 As you go down any group… 1. The atomic radii increase so the furthermost electron gets further away from the nucleus 2. The atoms have a bigger nuclear charge. But the furthermost electron is shielded from the extra nuclear charge by the inner shells of electrons “The shielding effect” The ionisation energy decreases.. It gets easier to pull the electron away from the nucleus Ionisation energy – explaining the trend Higher level only Decreases as you go across any period 3 Lithium 4 Beryllium 5 6 Aluminium Carbon As you go across a period.. 1. the nuclear charge in the atom increases. The increasing nuclear charge pulls the furthermost electron tighter to the nucleus 2. The atomic radii decrease so the furthermost electron is closer to the nucleus The ionisation energy increases.. It gets harder to pull the electron away from the nucleus Plot of atomic number against ionisation energy The second period Higher level only Higher level only Ionisation energies across the second period Generally they go up - but there are exceptions… Beryllium and Nitrogen have higher ionisation energies than you would expect. Higher level only The electronic configuration of: Lithium It has three electrons. The electronic configuration is 1s2 2s1 Higher level only The electronic configuration of: Beryllium It has 4 electrons. The electronic configuration: 1s2 2s2 It has a full 2s sublevel and this gives it extra stability! It would be extra hard to pull the electron away so the ionisation energy is especially high. Higher level only The electronic configuration of: Boron It has 5 electrons. The electronic configuration: 1s2 2s2 2p1 Higher level only The electronic configuration of: Carbon It has 6 electrons. The electronic configuration: 1s2 2s2 2p2 Higher level only The electronic configuration of: Nitrogen It has 7 electrons. The electronic configuration: 1s2 2s2 2p3 It has a half full 2p sublevel and this gives it extra stability! It would be extra hard to pull the electron away so the ionisation energy is especially high. Higher level only The electronic configuration of: Oxygen It has 8 electrons. The electronic configuration: 1s2 2s2 2p4 Plot of atomic number against ionisation energy The third period Higher level only Higher level only Ionisation energies across the third period Generally they go up - but there are exceptions… Magnesium and Potassium have higher ionisation energies than you would expect. Higher level only The electronic configuration of: Magnesium It has 12 electrons. The electronic configuration: 1s2 2s2 2p6 3s2 It has a full 3s sublevel and this gives it extra stability! It would be extra hard to pull the electron away so the ionisation energy is especially high. Higher level only The electronic configuration of: Potassium It has 15 electrons. The electronic configuration: 1s2 2s2 2p6 3s2 3p3 It has a half full 3p sublevel and this gives it extra stability! It would be extra hard to pull the electron away so the ionisation energy is especially high. Check your learning.. • • • • • What is an energy level? What is an orbital? What is an energy sublevel? Define atomic radius Define ionisation energy Today’s objectives • Second and successive ionisation energies and evidence for energy levels • The chemical properties of the elements based on atomic radii and ionisation energies Higher level only Second ionisation energies • Def the energy needed to remove the second electron from a singly charged positive ion. First ionisation energy needed for: Second ionisation energy needed for: Be Be+ Be+ + e Be2+ +e Higher level only Beryllium: Successive ionisation energies + 4 Beryllium (neutral atom) 4 Beryllium+ (positive ion) The second ionisation energy is always higher than the first because in an ion the atomic radius decreases and the electrons are closer to the nucleus – harder to take away! Higher level only Beryllium: Successive ionisation energies Beryllium There is a dramatic increase in ionisation energy needed to remove the third electron! Higher level only Beryllium: Successive ionisation energies There is a large ionisation energy needed because the electron is removed from a new energy level! The electronic configuration: 1s2 2s2 Higher level only Ionisation values provide evidence for the existence of energy levels: • The ionisation energy required to remove an electron from a new shell jumps dramatically! This is because it is much harder to remove an electron from an inner shell as it is closer to the nucleus and has less shielding from the nuclear charge than before. Chemical Properties of the elements The chemical properties of each element depends on its electronic structure. Elements in the same group in the Periodic Table have similar electronic structures – they all have the same number of electrons in their outermost shell - and so have similar chemical properties. Periodic table Group 1- The Alkali metals Why do they always react to lose an electron? Can you explain why reactivity increases going down the group? They all have one electron in their outermost shell, which they have a tendency to lose. Potassium 3 Group 1 – the Alkali metals • Going down the group the atomic radius increases 11 • Going down the group the nuclear charge also increases, but screening also increases so this effect is cancelled out. 19 Potassium • The result is that as you go down the group the outermost electron becomes easier to remove, and the element is therefore more reactive! Alkali metals reacting with water Sodium Lithium Potassium Periodic table Group 17- The Halogens 9 17 24 Why do these elements always gain an electron? Can you explain why reactivity decreases as you go down this group? They all have 7 electrons in their outermost shell and have a tendency to gain one electron in chemical reactions 9 Group 17- The Halogens • A high nuclear charge and small atomic radius attracts electrons to the halogens 17 • Going down the group the atomic radius increases • Going down the group the nuclear charge also increases, but screening also increases so this effect is cancelled out. 24 • The result is that as you go down the group the ability to gain an electron decreases… and the element is therefore less reactive!