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 X-ray scattering techniques are a family of non-destructive analytical techniques which reveal information about the crystallographic structure, chemical composition, and physical properties of materials and thin films. X ray diffraction - These techniques are based on observing the scattered intensity of an x-ray beam hitting a sample as a function of incident and scattered angle, polarization, and wavelength or energy. Elecron diffraction - Electron diffraction is a technique used to study matter by firing electrons at a sample and observing the resulting interference pattern. This phenomenon occurs due to the wave-particle duality, which states that a particle of matter (in this case the incident electron) can be described as a wave. For this reason, an electron can be regarded as a wave much like sound or water waves. This technique is similar to the X-ray diffraction and neutron diffraction. Electron diffraction is most frequently used in solid state physics and chemistry to study the crystal structure of solids. These experiments are usually performed in a transmission electron microscope (TEM), or a scanning electron microscope (SEM) as electron backscatter diffraction. In these instruments, the electrons are accelerated by an electrostatic potential in order to gain the desired energy and determine their wavelength before they interact with the sample to be studied. 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. The ionic radius of an element is its share of the distance between adjacent ions in an ionic solid.   For example, the distance between the magnesium cation and the oxide anion in magnesium oxide is 211 pm, the radius of the Mg2+ ion is calculated from 205pm - 146 pm = 65 pm. The size of an ion also varies somewhat with its environment so average values for ionic radii are often quoted.

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

As you go down a group… The atomic radii increases.. 3 1. the atoms have extra electrons which have to go into a new shell further away from the nucleus. 11 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” 19 The atomic radii increases.. Potassium

As you go across a period.. 4 5 3 6 Aluminium Carbon Beryllium Lithium 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! Due to more protons. The atomic radii decreases..

Ionisation energy Higher level only Ionisation energy is the minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom. In its ground state and isolated, Not e the position of hydrogen is kind of incorrect

Check the ionisation energy What is the pattern 1) as you go down a group 2) as you go across a period

Decreases as you go across any period Higher level only Ionisation energy – the pattern Decreases as you go across any period Increases as you go down any group Notice the same trend as before!!!!

Ionisation energy – explaining the trend Higher level only Ionisation energy – explaining the trend Increases as you go down any group Notice the same trend as before!!!!

As you go down any group… 3 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” 11 The ionisation energy decreases.. It gets easier to pull the electron away from the nucleus 19

Decreases as you go across any period Ionisation energy – explaining the trend Higher level only Decreases as you go across any period Notice the same trend as before!!!!

As you go across a period.. Lithium Beryllium Aluminium Carbon 3 4 5 6 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 Due to more protons. 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 Maybe an activity where they spot the trends and the exceptions and revise what they have just done! Higher level only

Ionisation energies across the second period 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.

The electronic configuration of: Higher level only The electronic configuration of: Lithium It has three electrons. The electronic configuration is 1s2 2s1

The electronic configuration of: 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. More excercises, all 36 elements!

The electronic configuration of: Higher level only The electronic configuration of: Boron It has 5 electrons. The electronic configuration: 1s2 2s2 2p1 More excercises, all 36 elements!

The electronic configuration of: Higher level only The electronic configuration of: Carbon It has 6 electrons. The electronic configuration: 1s2 2s2 2p2 More excercises, all 36 elements!

The electronic configuration of: 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. More excercises, all 36 elements!

The electronic configuration of: Higher level only The electronic configuration of: Oxygen It has 8 electrons. The electronic configuration: 1s2 2s2 2p4 The trend continues for fluorine and neon!

Plot of atomic number against ionisation energy The third period Maybe an activity where they spot the trends and the exceptions and revise what they have just done! Higher level only

Ionisation energies across the third period 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.

The electronic configuration of: 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. The trend continues for fluorine and neon!

The electronic configuration of: 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. The trend continues for fluorine and neon!

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

Second 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: Be Be+ + e After the first electron has already been moved! Second ionisation energy needed for: Be+ Be2+ +e

Beryllium: Successive ionisation energies Higher level only Beryllium: Successive ionisation energies + 4 4 This cntinues for the third ionisation energies and so on… Beryllium+ (positive ion) Beryllium (neutral atom) 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!

Beryllium: Successive ionisation energies Higher level only Beryllium: Successive ionisation energies Beryllium There is a dramatic increase in ionisation energy needed to remove the third electron!

Beryllium: Successive ionisation energies 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

Ionisation values provide evidence for the existence of energy levels: 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. Inner shell is also more stable!

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. (how they tend to behave in chemical reactions)

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

Group 1 – the Alkali metals 3 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. The result is that as you go down the group the outermost electron becomes easier to remove, and the element is therefore more reactive! 11 19 Potassium

Alkali metals reacting with water Lithium Sodium 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

Group 17- The Halogens 9 A high nuclear charge and small atomic radius attracts electrons to the halogens 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. The result is that as you go down the group the ability to gain an electron decreases… and the element is therefore less reactive! 17 Flourine >chlorine>bromine 24