atomic structure

subatomic particles

• atoms are made of 3 fundamental subatomic particles: protons, neutrons electrons

  • protons and neutrons (nucleons) form the dense nucleus (or centre) of atoms

  • electrons surround the nucleus and are arranged in electronic shells

particle	actual mass/kg	relative mass	charge relative to proton	location	deflection in an electric field
proton	$1.67$ x $10^{-27}$	1	+1	in the nucleus of an atom	towards negative electrode
neutron	$1.67$ x $10^{-27}$	1	0	in the nucleus of an atom	not deflected
electron	$9.11$ x $10^{-31}$	1/1840	-1	in electronic shells	towards positive electrode

behaviour of subatomic particles in electric field

• due to relative masses and charges, the three subatomic particles behave differently in an electric field

  • neutrons: not deflected in an electric field as they are electrically neutral

  • positively charged protons: deflected towards the negatively charged plate

  • negatively charged electrons: deflected towards the positively charged plate

• angle of deflection = k(charge/mass)

• deflection of charged particles begins only at the start of the electric field

• extent of deflection for electrons is more than for protons → due to smaller charge/mass

atomic number

number of protons in nucleus of an atom

nucleon number/mass number

total number of protons and neutrons in the nucleus of an atom

isotopes

• atoms of the same element having same number of protons (same atomic number) but different number of neutrons (different mass numbers

  • isotopes of the same element have the same chemical properties

arrangement of electrons

• electrons are arranged around the nucleus of an atom in electronic shells

• shells are numbered starting from the nucleus → numbers are known as principal quantum numbers, n

  • 1: 2 electrons

  • 2: 8 electrons

  • 3: 18 electrons

• electronic configuration: arrangement of electrons in an atom

shells, subshells and atomic orbitals

• electrons in atoms are arranged in a series of electronic shells described by the principal quantum number

• each principal quantum shell comprises of subshell(s) → designated by a letter: s, p, d, f

• each type of subshell contains a specific number of atomic orbitals

• number of different types of subshells in a shell is equal to its principal quantum number

• each orbital can accommodate up to two electrons

  • an electron in an atom behaves like a tiny magnet, spinning on its own axis like a top in a clockwise or anticlockwise direction

• in ground state configuration, two electrons in the same orbital must have opposite spins

  type of subshell	number of orbitals in each subshell	orbitals present
  s	1	s
  p	3	px, py, pz
  d	5	dxy, dyz, dxz, dx2-y2, dz2
  f	7

subshell → definition

a group of orbitals that have the same energy level (degenerate) but different orientation in space

atomic orbital → definition

a region or volume of space near the nucleus where there is a high probability (95%) of finding an electron

energy level of shells and subshells

• as principal quantum number increases, the

  • electrons are further from the nucleus

  • electrons are at a higher energy level

  • size of orbital increases

• energy levels converge as distance from the nucleus increases

• energy difference between shells is much greater than energy difference between subshells

• orbitals within the same subshell are degenerate (same energy) but have different orientation

• energies and distance from nucleus of subshells within each shell differ as follows: s < p < d < f

• vacant 4s orbital is of lower energy than the vacant 3d orbitals despite its higher quantum number

  • once 3d subshell is occupied by electrons (at least 1), 4s orbital is repelled to a higher energy level than 3d

s orbitals

• are spherical in shape and non-directional → probability of finding an electron at a particular distance from the nucleus is the same in all directions

• the size of the s orbitals increases with principal quantum number → size of 1s orbital < 2s orbital < 3s orbital

p orbitals

• three p orbitals of a subshell have dumb-bell shape and are directional → three p orbitals have different orientation in space

  • the probability of finding an electron at a particular distance is greatest along the x, y or z axis

• the size of p orbitals increases with principal quantum number: size of 2p orbital < 3p orbital

d orbitals

• have complex shapes and are directional

  • four of the five d orbitals (d_xy, d_yz, d_xz, d_x²-y²) have four lobes of the same size and shape

    • the d_xy, d_yz, d_xz lie on the xy, yz and xz plane respectively

    • the d_x²-y² orbital consists of four lobes along the x- and y- axes

  • the fifth d orbital (d_z²) has two lobes along the z axis and a ring shaped-region of electron density in the centre

• the size of the d orbitals increases with principal quantum number → size of 3d orbital < 4d orbital

rules for arrangement of electrons in orbitals

1. aufbau building principle states that electrons are accommodated in orbitals of lowest energy first

2. pauli exclusion principle states each orbital can only contain a maximum of two electrons of opposite spin

3. hund’s rule of maximum multiplicity: when filling a set of degenerate energy levels, the electrons will enter the orbitals singly, with spins in the same direction until the set is half-filled before pairing takes place

ground state electronic configuration

• if the above three rules are followed, then every electron is given the lowest possible energy → arrangement of electrons is known as the ground state electronic configuration

• order of filling orbitals: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

• electronic configuration is always written in order of increasing principal quantum number → indicates the order of increasing energy for the subshells

exceptions to the aufbau principle

• chromium: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹

• copper: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

• following the aufbau principle, the 4s orbitals are filled first, followed by the 3d orbitals with the exception of chromium and copper

• 3d and 4s subshells are close in energy → the unpairing of 4s electrons takes place to form a more stable configuration with a lower energy state

excited state electronic configuration

• the outermost electrons of an atom can absorb energy and be excited to a higher energy level → more than one electron can be excited at a time

• excited particles can lose energy emitting radiation

electronic configuration of ions

isoelectronic atoms/ions have the same number of electrons

• the number of electrons in an ion is found from the proton number of the element and the charge of the ion

• to form cations from atoms, electrons must be removed from the subshell with highest energy

• to form anions from atoms, electrons are added to the lowest energy which is not fully filled

first ionisation energy

minimum energy required to completely remove one mole of valence electrons from one mole of ground state atoms in the gaseous state to form one mole of gaseous singly charged cations

• X (g) → X⁺ (g) + e^- ; ΔH: 1st IE > 0

second ionisation energy

minimum energy required to completely remove one mole of valence electrons from one mole of ground-state singly positively charged ions in the gaseous state to form one mole of gaseous doubly charged cations

• X⁺ (g) → X²⁺ + e^-; ΔH: 2nd IE > 0

factors influencing ionisation energies

nuclear charge

• nuclear charge is the attraction of the nucleus for the electrons

• increases with proton number

• the greater the attraction on the valence electrons, the greater the energy required to remove the electron

screening effect or shielding effect

• valence electrons are screened/shielded from attraction of nucleus by inner electrons

• screening effect increase with number of electrons, especially electrons in inner shells

• the greater the screening effect upon the valence electrons, the lower the energy required to remove the valence electron

distance of electron from the nucleus (no. of electronic shells)

• electrons are less strongly attracted when they are further away from the nucleus

• the further away the valence electrons are from the nucleus, the lower the energy required to remove the valence electron

general trend in successive ionisation energies of an atom

trend: there is a general increase in successive ionisation energies: 1st IE < 2nd IE < 3rd IE

• as electrons are being removed, the number of protons remains unchanged hence nuclear charge remains constant

• screening effect decreases (electrons are removed) and remaining electrons are held more strongly and closer to the nucleus

• more energy is required to remove remaining electrons

successive ionisation energies and electronic configuration

• from plotting successive ionisation energies against the order of removal of electrons, the following information can be obtained

  1. predicting group number of an element

  2. number of quantum shells and number of subshells occupied → period number can only be determined if the successive IE graph shows complete ionisation of the atom

increases in ionisation energies

• a large increase in ionisation energy occurs when the 3rd electron is removed (IE2 → IE3)

  • the 3rd electron comes from the inner shell which is nearer to the nucleus and requires more energy to remove it

  • indicates that Mg has 2 valence electrons and is a group 2 element

• a large increase in ionisation energy also occurs when the 11th electron is removed (IE10 → IE11)

  • the 11th electron comes from the inner shell which is closest to the nucleus hence more energy is required to remove it

  • two large increase in ionisation energies indicates that Mg has 3 electronic shells and is a period 3 element

• smaller increase in ionisation energy (IE8 → IE9)

  • the 8th electron is removed is from the 2p subshell while the 9th electron removed is from the 2s subshell

  • since electrons in the 2s subshell are slightly closer to the nucleus, they would require more energy to be removed

predicting group number of an element from successive ionisation energies data

• find the biggest difference in IE values

• determine the number of valence electrons, hence group number

periodic trends in ionisation energies → group

trend: down each group, the first ionisation energies decrease

• down the group, the number of protons increase, hence nuclear charge increase

• there is an additional electronic shell, hence the screening effect increases

• the valence electrons are further away and less strongly attracted by the nucleus

• less energy is required to remove the valence electrons, hence first ionisation energy decreases down a group

periodic trends in ionisation energies → period

trend: across the period, the first ionisation energies generally increase

• group 1 elements have the lowest ionisation energy within each period, and the group 18 elements have the highest

• across the period, number of protons increase hence nuclear charge increases

• additional electrons added to the same shell hence screening effect remains approximately constant

• valence electrons are closer to and more strongly held by the nucleus

• more energy is required to remove the valence electrons, first ionisation energy increases across a period

exceptions/anomalies in general trend

• for period 2 and 3, the exceptions occur between groups 2 and 13, and between groups 15 and 16

1. anomaly between group 2 and 13: IE of Be > IE of B

   • the first electron removed from boron atom is the 2p electron while that from beryllium is the 2s electron

   • the 2p electron is slightly further from the nucleus than the 2s and hence requires less energy to remove

2. anomaly between groups 15 and 16: IE of N > IE of O

   • the first ionisation energies of the two atoms involve the removal of the valence 2p electron

   • the 2p electron removed from oxygen is one of the paired 2p electron which experiences inter-electronic repulsion and hence requires less energy than expected for its removal