Chemistry - 1.1 Atomic Structure

Proton

A sub-atomic particle located in the nucleus with a relative mass of 1 and a charge of +1.

Neutron

A sub-atomic particle located in the nucleus with a relative mass of 1 and a charge of 0.

Electron

A sub-atomic particle located in orbitals with a relative mass of 1/1840 and a charge of -1.

Atomic Number (Z)

The number of protons in the nucleus of an atom.

Mass Number (A)

The total number of protons and neutrons in an atom.

Number of Neutrons

Calculated as A - Z.

Isotopes

Atoms with the same number of protons but different numbers of neutrons, resulting in different masses.

Time of Flight Mass Spectrometer

An instrument used to determine the isotopes present in a sample of an element.

Ionisation

The process of converting a sample into ions, which can occur through methods like electron impact or electrospray ionisation.

Electron Impact Ionisation

A method where a vaporised sample is bombarded with high-energy electrons, knocking out outer electrons and forming positive ions.

Electrospray Ionisation

A technique where a sample is dissolved in a volatile solvent and injected through a needle to form a mist, allowing molecules to gain protons.

Acceleration

The process where positive ions are accelerated by an electric field to a constant kinetic energy.

Kinetic Energy (KE)

Given by the formula KE = ½ mv², where m is mass and v is velocity.

Flight Tube

The area where ions travel; lighter ions move faster than heavier ions due to the same kinetic energy.

Detection

The process where ions reach the detector, generating a current proportional to the abundance of the species.

m/z Ratio

The mass-to-charge ratio measured by the mass spectrometer for each isotope.

Abundance

The relative amount of a particular isotope present in a sample.

Example of Isotope Analysis

In a sample of nickel, one of the isotopes found was 59Ni, with ions accelerated to 1.000 x 10-16 J of kinetic energy.

Flight Tube Length

The distance through which ions travel in the flight tube, measured in meters.

Ion Drift Area

The section of the mass spectrometer where ions travel before detection.

Vacuum Requirement

The mass spectrometer must operate under a vacuum to prevent air particles from interfering with ion detection.

Current Generation

The small current produced when ions reach the detector, which is used for analysis.

Avogadro constant

L = 6.022 × 10^23 mol-1

Mass of one ion of 59Ni+

Mass of one mole of 59Ni+ = 59 / 6.022 × 10^23 = 9.797 × 10^-23 g = 9.797 × 10^-26 kg

Time calculation for ion

t = 0.8000 √(9.797 × 10^-26 / (2 × 1.000 × 10^-16)) = 1.771 × 10^-5 s

Relative Atomic Mass (R.A.M)

The relative atomic mass quoted on the periodic table is a weighted average of all the isotopes.

R.A.M calculation formula

R.A.M = Σ(isotopic mass × % abundance) / total relative abundance

Example R.A.M for Magnesium

R.A.M = [(78.7 × 24) + (10.13 × 25) + (11.17 × 26)] / 100 = 24.3

Relative abundance of isotopes

If relative abundance is used instead of percentage abundance, use the equation R.A.M = Σ(isotopic mass × relative abundance).

Example R.A.M calculation for Tellurium

R.A.M = [(124 × 2) + (126 × 4) + (128 × 7) + (130 × 6)] = 127.8

Copper isotopes

Copper has two isotopes 63-Cu and 65-Cu. The relative atomic mass of copper is 63.5.

Percentage abundance calculation for Copper

63.55 = y × 63 + (1 - y) × 65; 2y = 1.45; y = 0.725

Percentage abundance of 63-Cu

% abundance 63-Cu = 72.5%

Percentage abundance of 65-Cu

% abundance 65-Cu = 27.5%

Mass spectra for Cl2

Cl has two isotopes Cl35 (75%) and Cl37 (25%).

Mass spectra for Br2

Br has two isotopes Br79 (50%) and Br81 (50%).

Mass spectrometer function

Mass spectrometers have been included in planetary space probes to identify elements on other planets.

Molecular ion in mass spectrometry

The peak with the largest m/z is due to the complete molecule and is equal to the relative molecular mass, Mr, of the molecule.

Parent ion

The peak called the parent ion or molecular ion is equal to the relative molecular mass of the molecule.

Electro spray ionisation

Fragmentation will not occur; the peak equals the mass of the MH+ ion.

Relative molecular mass calculation

If a peak at 521.1 is for MH+, the relative molecular mass of the molecule is 520.1.

Bohr model of the atom

An early model of the atom with electrons in spherical orbits.

Principle energy levels

Numbered 1, 2, 3, 4... with 1 being closest to the nucleus.

Sub energy levels

Labelled s, p, d, and f with specific electron capacities.

Shapes of orbitals

Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus.

Subshells filling order

An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s): 1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p.

s sublevels

s sublevels are spherical.

p sublevels

p sublevels are shaped like dumbbells.

Electronic structure of oxygen

For oxygen: 1s2 2s2 2p4.

Electronic structure of calcium

For calcium: 1s2 2s2 2p6 3s2 3p6 4s2.

Spin diagrams

An arrow is one electron; the arrows going in the opposite direction represent the different spins of the electrons in the orbital.

Filling orbitals

When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons.

s block element

A s block element is one whose outer electron is filling a s-sub shell, e.g., sodium 1s2 2s2 2p6 3s1.

p block element

A p block element is one whose outer electron is filling a p-sub shell, e.g., chlorine 1s2 2s2 2p6 3s2 3p5.

d block element

A d block element is one whose outer electron is filling a d-sub shell, e.g., vanadium 1s22s22p63s23p6 4s23d3.

Negative ion formation

When a negative ion is formed, electrons are gained; O is 1s2 2s2 2p4 becomes O2- is 1s2 2s2 2p6.

Positive ion formation

When a positive ion is formed, electrons are lost from the outermost shell; Mg is 1s2 2s2 2p6 3s2 becomes Mg2+ is 1s2 2s2 2p6.

d-block electronic structure complications

The electronic structure of the d-block has complications; conventionally, 4s fills before 3d, but there are exceptions, such as chromium and copper.

Ionisation energy definition

The first ionisation energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge.

First ionisation energy equation

This is represented by the equation: H(g) → H+(g) + e-.

Second ionisation energy definition

The second ionisation energy is the enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge.

Second ionisation energy equation

This is represented by the equation: **+(g) → Ti2+(g) + e-.

Factors affecting ionisation energy

There are three main factors: 1. The attraction of the nucleus (the more protons in the nucleus, the greater the attraction); 2. The distance of the electrons from the nucleus (the bigger the atom, the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus).

Shielding of the attraction of the nucleus

An electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus.

Successive ionisation energies

The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element.

First ionisation energy

The first ionisation energy of an element is the energy required to remove the outermost electron from a neutral atom.

Why are successive ionisation energies always larger?

The second ionisation energy of an element is always bigger than the first ionisation energy because the ion formed after the first electron is removed increases the attraction on the remaining electrons.

Big jump between ionisation energies

Notice the big jump between 4 and 5; the fifth electron is in an inner shell closer to the nucleus and is attracted much more strongly by the nucleus.

Example of ionisation energies

The ionisation energies are 590, 1150, 4940, 6480, 8120 kJ mol-1, indicating the element is in group 2 of the periodic table.

Periodic trends in ionisation energy

A repeating pattern across a period is called periodicity.

Why has helium the largest first ionisation energy?

Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells.

Why do first ionisation energies decrease down a group?

As one goes down a group, the outer electrons are found in shells further from the nucleus and are more shielded, so the attraction of the nucleus becomes smaller.

Why is there a general increase in first ionisation energy across a period?

As one goes across a period, the electrons are being added to the same shell, while the number of protons increases, making the effective attraction of the nucleus greater.

Why has Na a much lower first ionisation energy than neon?

Na has its outer electron in a 3s shell further from the nucleus and is more shielded, making it easier to remove.

Why is there a small drop from Mg to Al?

Al starts to fill a 3p subshell, whereas Mg has its outer electrons in the 3s subshell, making Al's outer electron slightly easier to remove.

Why is there a small drop from P to S?

In sulfur, the 4th electron is added to a 3p orbital, causing slight repulsion between the two negatively charged electrons, making the second electron easier to remove.

Patterns in the second ionisation energy

If the graph of second ionisation for each successive element is plotted, a similar pattern to the first ionisation energy is observed but shifted one to the left.

Lithium's second ionisation energy

Lithium would have the second largest ionisation of all elements as its second electron would be removed from the first 1s shell closest to the nucleus.

Comparison of ionisation energies

Li has a bigger second ionisation energy than He as it has more protons.