Atomic Structure: Mass, Charge, and Numbers

• Mass is concentrated at the nucleus of an atom
• Positively charged protons are in the nucleus
• Negatively charged electrons orbit the nucleus in shells

Mass and Charge in an Atom

• Protons are deflected towards the negative plate
• Electrons are deflected towards the positive plate
• Neutrons have no charge, and continue in a straight path
• With same energy, deflection of protons and electrons is exactly the same
• With same speed, lighter electrons are deflected more strongly than protons

Atomic and Mass Numbers

• Atomic number is the number of protons in an atom, also called the proton number
• Atoms of the same element have the same atomic number and number of protons
• Atoms have no overall charge, so number of electrons = number of protons
• Ions are charged, so number of electrons = atomic number +/- electrons gained/lost
• Mass number = total number of protons and neutrons, also called the nucleon number

Atomic Radius

• Atomic radius decreases across a period; increase in protons → increase in nuclear charge → attract outer electrons more strongly → reduces atomic radius
• Atomic radius increases down a group; nuclear charge increases, energy level increases → outer electrons are further from the nucleus → shielded by inner shell electrons

Ionic Radius

• Ionic radius decreases for positive ions, then increases for negative ions across a period
• Positive ions have same electron configuration; # of protons ↑ → nuclear attraction ↑ → ionic radius decreases
• Negative ions have gained electrons; more electrons than protons → nuclear attraction is weaker → increased ionic radius

The Nucleus of the Atom

• Nucleus contains protons and neutrons (aka nucleons)

Isotopes

• Isotopes: atoms of an element with same # of protons/electrons but a different # of neutrons
• Isotopes of the same element have different mass numbers since mass number = # protons + # neutrons
• Isotopes of the same element have the same chemical properties but differ in physical properties

Electrons

• Orbital: region of space that can contain up to 2 electrons
• Principal quantum number (n) is the shell that electrons occupy
• Higher principal quantum number value means electrons are higher energy + further from the nucleus

Types of Orbitals

• s: spherical shape, one s orbital in each shell from n = 1 upwards (two s electrons per shell), lowest energy
• p: dumbbell shape, three p orbitals in each shell from n = 2 upwards (six p electrons per shell), higher energy than s
• d: five d orbitals in each shell from n = 3 upwards (ten d electrons per shell), higher energy than p

Electron Filling

• Electrons fill orbitals from lowest to highest energy (1s, 2s, 2p, 3s...)
• 4s is filled before 3d, because 4s has a lower energy
• Before electrons start pairing, subshells must be filled with unpaired electrons
• Subshell: specific type of orbitals in a shell (e.g. p subshell has 3 p orbitals)

Exceptions to Standard Electron Configuration

• Two main exceptions due to stability of completely full or half full d sublevels; electron from 4s orbital is excited to 3d
• Chromium: 1s²2s²2p⁶3s²3p⁶3d⁵4s¹
• Copper: 1s²2s²2p⁶3s²3p⁶3d¹⁰4s¹
• When copper and chromium ions form, electrons are removed from 4s beforehand
• Full Configuration of Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
• Short Hand Configuration of Fe: [Ar] 3d⁶ 4s²
• Electron configurations show number of electrons and types of orbitals in each energy level:

Free Radical

• Single unpaired electron, shown as a dot next to the chemical symbol

First Ionisation Energy

• First ionisation energy: energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions, measured in kJ mol⁻¹
• Reaction must have gas state symbol

Factors Affecting Ionisation Energy

• Nuclear Charge: greater charge → stronger attraction to outer shell electrons → higher first ionisation energy
• Atomic Radius: larger radius → weaker attraction (positive nucleus/negative electrons) → lower first ionisation energy
• Electron Shielding: more shells → more shielding → weaker attraction (between outer shell electrons/nucleus) → lower first ionisation
• First ionisation energy of elements increases across a period due to increased nuclear charge/decreased atomic radius

Exceptions in Period 2

• Between Groups 2 and 3: electrons start to be added to a 2p versus 2s orbital
• 2p has a higher energy level → electron is further from nucleus → easier to remove (less energy needed)
• Between Groups 5 and 6: electrons start to pair in the 2p orbitals
• Paired electrons repel each other → easier to remove electrons with less energy

First Ionisation Energy and Atomic Properties

• First ionisation energy decreases down a group because even though nuclear charge increases, the atomic radius and electron shielding also increase

Successive Ionisation Energies

• Successive ionisation energies: removing one mole of electrons from one mole of gaseous ions
• Increase because attraction between outer shell electrons and nucleus is greater, because atomic radius decreases
• Large jump between successive ionisation energies indicates the group number of the element