A level Chemistry Atomic Structure and the Periodic Table

Structure of an atom

An atom has a nucleus containing protons and neutrons that is surrounded by orbiting electrons.

Protons

have a relative charge of +1 and a relative mass of 1

Neutrons

have a relative charge of 0 and a relative mass of 1

Electrons

have a relative charge of -1 and a relative mass of 1/1840

Atomic Number

The number of protons in an atom.

(A neutral atom will have the same number of electrons as protons, but ions have different numbers of electrons so the proton number determines the element's identity)

Mass number

The sum of the protons and neutrons in an atom's nucleus

Isotope

Atoms of the same element (same number of protons) with different masses due to differing numbers of neutrons

Relative atomic mass

The weighted mean mass of an atom of an element, on a scale in which one atom of 12C weighs 12 units

Relative Isotopic Mass

The mass of an atom of an isotope of an element, on a scale in which one atom of 12C weighs 12 units

Relative Molecular Mass and Relative Formula Mass

Relative molecular mass (Mr) should only be used for molecules (covalently bonded).

RFM can be used for any compound or diatomic species.

Mass Spectra

Mass Spectra will show the masses and how frequently they occur, either as a number or a percentage. This can be used to calculate the relative atomic mass of an element.

Diatomic molecules such as chlorine and bromine have 3 large peaks at the end due to the 3 possible combinations of isotopes.

First ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of 1+ ions

Second ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of 2+ ions

Trends in ionisation energies - down a group

Ionisation energy decreases down the group as although the nuclear charge increases (thereby attracting the electrons more) there are more quantum shells shielding the outer electron from the nucleus and the outer electron is further from the nucleus, thus inhabiting a higher energy quantum shell

Subshells

Quantum shells are made up of subshells.

There are s, p, d and f subshells.

Each new shell contains one more subshell e.g. shell 1 - s only

shell 2 - s and p

shell 3 - s, p and d

shell 4 - s,p,d and f

Orbitals

a region of space/probability function where a maximum of two electrons can be found

s subshell

Shaped like a sphere with one orbital (hence contains 2 electrons)

p subshell

Shaped like 3 dumbbells at right angles to each other. Has 3 orbitals and thus 6 electrons

d subshell

Has 5 orbitals and thus 10 electrons

Electron configuration

This determines an elements chemical properties

Pauli exclusion principle

For electrons to occupy the same orbital they must have an opposite spin

Exceptions to electron configuration - d block

Chromium and Copper are exceptions to the pattern. This is because for both cases the 3d subshell is either one electron away from being half full or completely full. As these states are more stable they pull an electron down from the 4s subshell.

Additionally, when transition metals ionise they lose the electrons in the 4s subshell first, despite gaining electrons in the 3d subshell last.

Periodicity

The repeating of patterns across different periods

s, p, d and f blocks

Evidence for electron configuration

Evidence for electron configuration, shells, subshells and orbitals can be found from:

- Atomic emission spectra

-Successive ionisation energies

-First ionisation energies of successive elements

Atomic emission spectra

When atoms absorb energy (either heat or electrical) the electrons are exited and move up an energy level before returning to the ground state.

The energy is released as electromagnetic radiation. The spectra of this radiation is unique for every atom, because the energy spacing between levels is discrete and unique to individual elements.

If energy levels were not discrete, the emitted radiation would be continuous in frequency.

Successive ionisation energies

The jumps show where an electron is in a new main energy level and so requires more energy to be removed.

This can be used to determine which group the element is in.

First ionisation energies of successive elements

Shows general trends - increases across a period due to increasing nuclear charge (electrons are in the same main energy level but the number of protons in the nucleus increases) and thus greater attraction.

First ionisation energy of successive elements: decrease from group 8 to 1

The outer electron is in a new main energy level. It is therefore further from the nucleus and the nuclear attraction is shielded by another shell of electrons - less attraction and electron is easier to remove.

First ionisation energy of successive elements: decrease from group 2 to 3

The outer electron is now in a p subshell instead of an s subshell and is slightly further from the nucleus, decreasing attraction and so electron is easier to remove.

First ionisation energy of successive elements: decrease from group 5 to 6

The outer electron is no longer in a singly occupied orbital (configuration goes from p3 to p4) but is paired with another electron; the resulting repulsion between electrons makes the electron slightly easier to remove.

Calculating relative atomic mass from mass spectra

sum of the mass x abundance for each isotope, divided by the sum of the abundances (100 if percentages used)