Chapter 7: Periodicity

• Elements are arranged in vertical columns called groups. Each element in a group has the same number of outer-shell electrons and similar properties. • Elements are arranged in horizontal rows called periods. The number of the period gives the number of shells in the element • Reading from left to right, elements are arranged in order of increasing atomic number

Across each period, there is a repeating trend in properties of the elements.

Each period starts with an electron in a new highest-energy shell. For each period, the s- and p- subshells are filled in the same way - a periodic pattern.

Elements in each group have the same number of electrons in their outer shell. Elements in each group also have the same number of electrons in each sub-shell. The similarity in electron configuration gives elements in the same group their similar chemistry.

• Arranged in order of atomic mass • Elements in groups of similar properties - if it didn’t fit, he swapped around elements and left gaps for elements yet to be discovered • Predicted properties of undiscovered elements, which were relatively accurate • Noble gases hadn’t been discovered yet

The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of 1+ ions.

Na (g) –> Na+ (g) + e-

Al4+ (g) –> Al5+(g) + e-

Successive ionisation energies are a measure of the energy required to remove each electron in turn. e.g. the second ionisation energy of an element is the energy required to remove 1e- from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

They provide important evidence for the different electron energy levels in an atom.

• The number of electrons in the outer shell • The group of the element in the periodic table -The identity of an element

Atomic radius - the greater the distance between the nucleus and the outer electron, the less the nuclear attraction (the force of attraction falls of sharply with increasing distance, so atomic radius has a large effect). Nuclear charge - the more protons there are in the nucleus, the greater the attraction between the nucleus and the outer electron. Electron shielding - Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and outer electron.

• Nuclear charge increases (most important factor) • Same shell: similar shielding • Atomic radius decreases • Nuclear attraction increases • 1st IE increases

• Atomic radius increases • More inner shells so shielding increases • Nuclear attraction on outer electrons decreases • 1st IE increases

• Extra shells • Electron to be lost is further from nucleus - atomic radius increases • Inner shells shield outer electrons from the attraction from the nucleus • large drop in IE

There is a general rise, but there are falls in certain areas.

• In boron the 2p subshell starts to fill • 2p subshell in boron has a higher energy than 2s subshell in beryllium (so easier to lose electron) • In boron the 2p electron is easier to remove (further away from nucleus) than one of the 2s electrons in beryllium • the 1st IE of boron is less than 1st IE of beryllium If you’re stuck on this kind of question, quickly draw the electron configuration for each of these atoms

The fall in 1st IE from Nitrogen to oxygen marks the start of electron pairing in the p orbitals of the 2p subshell. - In nitrogen and oxygen the highest energy electrons are in a 2p subshell - In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom - therefore the 1st IE of oxygen is less than the 1st IE of nitrogen.

It takes more energy to remove the 3rd e- than the 2nd e- because of increasing nuclear attraction since: 1) The number of protons remains the same. 2) The number of electrons decreases by 1. 3) The ratio of p:e increases, increasing the nuclear attraction per e- and reducing the ionic radius.

Metallic bonding is the strong electrostatic attraction between cations (positive ions) and delocalised electrons. The cations are fixed in position, maintaining the structure and shape of the metal. The delocalised electrons are mobile and are able to move throughout the structure. Only the electrons move.

Giant metallic lattice With cations in a fixed position (each atom has donated its negative outer shell electrons to a shared pool of electrons which are delocalised throughout the structure) And mobile delocalised electrons

Ionic radius - greater ionic radius the weaker the attraction, because: - greater ionic radius means same electron cloud has to cover a larger volume (e.g. lithium is stronger than sodium) Electron density - greater electron density, the greater the attraction, because: - greater electron density holds the ions together more strongly (e.g. beryllium is stronger than lithium) Electron density is the number of delocalised electrons.

Most metals have: - strong metallic bonds - attraction between positive ions and delocalised electrons - high electrical conductivity - high melting and boiling points - Insoluble Only mercury is a liquid at room temperature.

Because delocalised electrons can move through the structure, carrying charge So when a voltage is applied, the electrons will move towards the positive terminal, causing an electrical current in response to the potential difference.