C1 - Atomic Theory

Structure of an atom

Consist of three subatomic particles: protons, neutrons and electrons

Nucleus of an atom

• Contains protons and neutrons

• Has a positive charge and is very small and dense

  • Contains most of mass

Electrons in an atom

• Negatively charged

• Orbit the nucleus in shells/orbitals

• Most of atom is made of empty space

Properties of subatomic particles 

• Protons

  • RM: 1

  • RC: +1

 

• Neutrons:

  • RM: 1

  • RC: 0

 

• Electrons

  • RM: 1/1840

  • RC: -1

Mass number

• Total number of protons and neutrons in the nucleus

• Bigger number

Atomic number

Total number of protons in the nucleus, and defines the element

Isotope

• Different atoms of the same element which have same number of protons and electrons but a different number of neutrons

  • Means they have a slightly different physical property (mass and density)

  • Same chemical properties (reactivity)

Relative isotopic mass

• Mass of an atom of an isotope compared to 1/12 of the mass of carbon-12

• Rounded to one dp

Relative atomic mass

Mean mass of an atom of an element compared to the mass of 1/12 of carbon-12

Calculated by (sum of (ab x mass)) / sum of ab

Relative molecular mass (Mr)

• Used for simple covalent molecules

• Calculated by adding the Ar values of all atoms in one molecule

Relative Formula Mass

• Used for ionic compounds

• Used for giant structures

• Calculated by adding all the Ar values of all the ions in one formula unit

Emission spectra

 

• Electrons absorb and release energy in discrete amounts

  • Electrons occupy fixed energy levels called quantum shells

  • They can be excited to higher energy levels by absorbing energy or vice versa

Lines in emission spectra

Display the frequencies of light emitted as electrons transition from higher/lower energy levels

Quantum shell model (and how this provides evidence for quantum shells)

• Electrons are confined to fixed shells

• Each shell has a defined energy and electrons cannot exist between shells

• Electrons must absorb/emit electromagnetic radiation to move

• Emitted radiation has fixed frequencies

How does emission spectra support the Quantum Shell Model

• Electrons cannot smoothly transition

• Presence of distinct lines rather an a spectrum indicated quantized energy levels

• Each line represents a specific electronic transition

Mass spectroscopy

• Plots the relative abundances of ions against their mass to charge ratio (m/z)

 

• X axis:

  • m/z values equals relative mass of ion

 

• Y - axis:

  • Relative abundance

Mass spectroscopy: molecular samples

Peak at highest m/z value is the molecular ion (M+)

• This is the molecular ion peak 

• Matches the molecule’s Mr

• Lower peaks come from fragments of molecular ions

Ionisation energy

• Refers to the process of removing one or more electrons

• Endothermic reaction

First Ionisation energy

Energy needed to remove 1 electron from each atom in one mole of gaseous atoms to form 1 mole of gaseous 1+ ions

Factors that affect ionisation energy

• NASA

• Nuclear Charge

• Atomic Radius

• Shielding effect

• How this impacts Nuclear Attraction

Nuclear Attraction

More protons = stronger positive charge = stronger electrostatic force of attraction

Atomic Radius

• Electrostatic attraction drops off steeply with increasing distance

• Electrons in smaller atoms are held closer so attraction is greater

Shielding effect

• Inner electrons create shielding effect for outer electrons

• More electrons means more shielding

Trends in ionisation energy: Group

• Decreases down a group

  • Nuclear charge increases down as more protons are added

    • Increases attraction for electrons

  • Atomic radius increases as more electron shells are added

  • Electron shielding increases down a group as more inner shells reduce nuclear attraction

Trends in ionisation energy: Period

• Increase across a period

  • Nuclear charge increases as more protons are added across a period

  • Atomic radius decreases across a period as more electrons are added

  • Electron shielding stays similar across a period

Successive ionisation energies 

(Energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions) needed to remove each successive electron

Large jumps

• When entering a new shell, there are often large jumps because nuclear attraction increases greatly

• General pattern before is that energy increased

FIE of Na

Na_(g) —> Na⁺_(g) + e^-

5th Ionisation Energy of Na

Na⁴⁺_(g) —> Na⁵⁺_(g) + e^-

Shells

• Hold electrons

• Each one is defined by a principal quantum number, indicating the relative distance

• Is divided into sub-shells

Sub-shells

• Have distinct energy levels labelled s, p, d

• Contain orbitals

Electrons in subshells

• s: can have 1 orbital and therefore 2 electrons

• p: can have 3 orbitals and therefore 6 electrons

• d: can have 5 orbitals and therefore 10 electrons

Orbitals

• Are regions with a high probability of finding an electron found around the nucleus

• Can accommodate up to 2 electrons with opposite spins

S orbital

O shape

P orbital

Dumbbell shape

Electrons in the first four shells

• Electrons in the first four shells

  • Shell 1: 1s, 2 electrons

  • Shell 2: 2s and 2p, 8 electrons

  • Shell 3: 3s and 3p and 3d, 18 electrons

  • Shell 4: 4s, 4p, 4d, 4f and 32 electrons

Subshell notation

• Sub-shell notation

  • Indicated the number of electrons within each subshell

  • 1s² 2s² 2p⁶

Working out electron configuration

• Fill out the lowest energy orbitals first

• Fill subshells singly before pairing up

• Must have opposite spins

• For ions in s/p blocks, electrons are added to/removed from highest occupied subshell

Order of subshells

• 1s

• 2s

• 2p

• 3s

• 3p

• 4s

• 3d

• 4p

Noble gas notation

• Use square brackets of the closest preceding noble gas

• e.g [Ar] 4s²

Chromium notation

1s2 2s2 2p6 3s2 3p6 3d5 4s1

Copper notation

 1s2 2s2 2p6 3s2 3p6 3d10 4s1

Why are copper and chromium irregular

These are irregular because configurations with d subshells are energetically more favourable

Arrangement of periodic table

• In order of increasing atomic number

• Can predict their behaviour

• Divided into groups and periods

s/p/d blocks

Have their valence electron in that orbital

Periodic properties

• Electronic configuration determines the chemical properties

• Type of bonds and their strength will affect the melting/boiling points of elements in period 2 and 3

Metals properties due to e config

• Strength of metallic bonding increases across the period

  • Ions have larger positive charge and smaller ionic radius

  • More delocalised electrons

  • Means greater electrostatic force of attraction

Giant covalent lattices properties due to e config

• Carbon and silicon

• A huge amount of energy is required to break the covalent bonds so very high bp/mp

Simple covalent properties due to e config

Only have weak induced dipole forces existing between molecules which are very easy to overcome

Noble gases properties due to e config

Very low bp/mp as no bonds formed and very weak induced dipole-dipole forces attract

First Ionisation Energy periodicity in the periodic table

• Rules as before apply BUT there is a variation in drops between group 2-3 and 5-6

Why are the rules different for groups 2 - 3

• In group 3 the electron in removed from p orbital rather than s like in group 2

• P orbitals have slightly higher energy than s orbitals so valence electron is an average further

• P orbital also experiences additional shielding from nucleus by s electrons

• Means less energy is required