Nuclear atom and electronic configuration

structure of atoms

• atoms contain a positively charged, dense nucleus composed of protons and neutrons

• negatively charged electrons occupy the space outside the nucleus

• electron mass considered negligible

isotopes

• different atoms of the same element containing the same number of protons and electrons but different number of neutrons

calculating relative atomic mass

• Ar: average mass of one atom of an element compared to 1/12 of the mass of a carbon-12 atom

• to calculate:

  • multiply percentage abundance by mass of each isotope and divide by 100

electromagnetic spectra

• frequency: how many waves pass per second

• wavelength: distance between 2 consecutive peaks on the wave

• radio, microwaves, IR, visible light, UV, x-rays, gamma rays

  • energy increases

  • frequency increases

  • wavelength decreases

• continuous spectrum: shows all wavelengths (colors) of light

• line spectrum: only contains emissions at particular wavelengths

emission spectra

• electrons move rapidly in energy shells, if energy is absorbed they get excited and jump to higher energy level

• when electrons drop down from a higher to lower level, they emit photons of energy

  • this energy corresponds to a wavelength and shows up as a line on the emission spectra

hydrogen emission spectra

• lines converge towards the higher energy end, so electron is reaching a max amount of energy

  • max energy corresponds to ionisation energy

• these lines correspond to electron jumping from higher levels to n=2

• electron jumps:

  • to n=3 is IR

  • to n=2 is visible light

  • to n=1 is UV

electron shells

• electrons arranged around nucleus in principal energy levels

  • the lower the number, the closer to nucleus

  • the higher, the greater the energy of electron within the shell

• number of electrons in principal energy level is 2n²

subshells

• principal energy levels split into subshells s,p,d

  • n=1 : 1s

  • n=2 : 2s 2p

  • n=3 : 3s 3p 3d

• subshells contain orbitals which can be occupied by max of 2 electrons

  • s orbital: spherical, size increases with increasing shell number

  • p orbital: dumbbell shaped, lobes become larger and longer with increasing shell number

aufbau principle

• ground state (most stable electron config with lowest amount of energy) is achieved by filling subshells with lowest energy first

• 3d is higher in energy to 4s so is occupied after

  • when ionising electrons are lost from 4s first

• exceptions:

  • Cr is [Ar] 3d⁵ 4s¹ not [Ar] 3d⁴ 4s²

  • Cu is [Ar] 3d¹⁰ 4s¹ not [Ar] 3d⁹ 4s²

    • this is because it is energetically favourable to achieve a full or half full d-subshell

hund’s rule

• electrons with the same spin repel eachother: spin pair repulsion

  • so, electrons occupy separate orbitals in the same subshell to minimise repulsion and have spin in same direction

  • they then pair up, with second electron being added to first p orbital w/ spin in opposite direction

pauli exclusion principle

• orbital can only hold 2 electrons and they must have opposite spin

  • because energy required to jump to higher empty orbital is greater than inter-electron repulsion

ionisation energy

• in an emission spectrum, the limit of convergence at higher frequencies corresponds to the ionisation

• ionisation energy: energy required to remove one electron from an atom in its gaseous state

calculating first ionisation energy

• calculated using frequency/wavelength of convergence limit

• ΔE = h ν

  • E (J)

  • h (J s)

  • v frequency (s^-1)

• c = ν λ

  • λ (m)

• to calculate IE per mole, multiply energy by avogadro’s constant

successive IEs of an element

• successive IE’s increase

  • as more electrons are removed, the pull of the protons holds the remaining electrons more tightly so more energy is required to remove them

successive IE graph

• big jumps: change of shell

• small jumps: change of subshell

• analysing where large jumps appear and number of electrons removed when large jumps occur can deduce electronic config

• to deduce group number:

  • the largest jump in ionisation energy

  • e.g if biggest jump from first to second, means that it’s easier to remove first than second, so the first electron removed is last electron in valence shell, thus group I

factors affecting first IE

• size of nuclear charge: proportional

• distance of outer electrons from nucleus: inverse

• shielding effect of inner electrons: inverse

• spin pair repulsion: inverse

trends in IE in periodic table

• across a period:

  • generally increases

    • nuclear charge increases

    • atomic radius decreases, as outer shell pulled closer to nucleus so distance between nucleus and outer electrons decreases

    • shielding by inner electrons remains constant as electrons added to same shell

  • exceptions:

    • decrease in IE1 between Be and B

      • fifth electron in B is in 2p which is further away from nucleus than 2s subshell of Be

    • decrease in IE1 between N and O due to spin-pair repulsion

      • since 2 electrons in 2p_x orbital of O, there is repulsion so easier for one of them to be removed

  • from one period to next:

    • large decrease

      • increased distance between nucleus and outer electrons

      • increased shielding

  • down a group:

    • decreases

      • atomic radius increases

      • increased shielding