CHAPTER 1 - ATOMIC STRUCTURE AND THE PT

Atomic number

The number of protons in its nucleus. Sometimes referred to as the ‘proton number’.

Mass number

The number of protons and neutrons in its nucleus. Protons and neutrons are sometimes called nucleons so ‘nucleon number’ is sometimes used.

Mass : charge (m/z)

The ratio of the relative mass, m, of an ion to its charge, z, where z is the number of charges. Spectrometers usually operate so that most ions produced have the value z=1.

Isotopes

Atoms of the same element which have the same number of protons in the nucleus but a different number of neutrons. So isotopes have the same atomic number but different mass numbers.

Relative isotopic mass

The mass of one atom of an isotope relative to 1/12th of the mass of an atom of the isotope carbon - 12. The values are relative so they do not have units.

Relative atomic mass

The average mass of an atom of an element relative to 1/12th of the mass of an isotope of the isotope carbon-12. The values are relative so they do not have units.

Relative molecular mass

The sum of the relative atomic masses of all the atoms in its molecular formula.

Relative formula mass

The sum of the relative atomic masses of all the atoms in its formula.

Ionisation energy

The energy needed to remove the one mole of electrons from one mole of gaseous atoms, or ions, of an element.

Atomic energy levels

The energies of electrons in atoms. According to quantum theory, each electron in an atom has a definite energy. When atoms gain or lose energy, the electrons jump from one energy level to another.

Shielding

An effect of inner electron which reduces the pull of the nucleus on the electrons in the outer shell of an atom. This means that the electrons in the outer shell are attracted by an ‘effective nuclear charge’ which is less than the dull charge on the nucleus.

Atomic Orbitals

The sub-divisions of the electron shells in atoms. The main shells divide into sub-shells labelled s, p, d and f. The sub-shells are further divided into atomic orbitals. An orbital is a region in space around the nucleus of an atom in which there is a 95% chance of finding an electron, or a pair of electrons with opposite spins.

Electron configuration

The number and arrangement of electrons in an atom of the element.

Period

A horizontal row of elements in the periodic table.

Group

A vertical column of elements in the periodic table. Elements in the same group have similar properties because they have the same outer electron configuration.

How to determine the relative molecular mass from mass spectrometry

Look at m/z value furthest to the right.

Successive ionisation energies

The energy required for the removal of the second electron from the ion.

Ionisation

An electron is knocked off each particle by the high energy electrons to form 1+ ions

Acceleration

Positive ions are accelerated by a negative electric plate.

Deflection

Ions are then deflected by a magnetic field according to their masses. The lighter they are, the more they are deflected.

Detection

Positive ions hit negative plate and produce and electric current. Size of current is proportional to the number of ions.

How are ionisation energies influenced

• Number of protons

• Electron shielding

• Electron sub-shell from which electron is removed

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

More protons so higher nuclear charge so the atomic radius decreases as there is a stronger attraction from the nucleus to the outer electron. Shielding is the same.

Why is there a general decrease in first ionisation energy down a group.

More electron shells, so more shielding and a larger atomic radius.

What do the lines in the emission spectrum correspond to in the structure of a helium atom?

Electrons releasing energy of fixed wavelength as they fall back from higher energy levels- shows existence of quantum shells - fixed packets of energy.

Differences between emission and absorption spectra

Absorption lines are where light has been absorbed by the atom thus the dip whereas emission spectra have spikes due to atoms releasing photons at those wavelengths.

Similarities between emission and absorption spectra

Absorption lines correspond to the frequencies of emission spectrum of the same element. The amount of energy absorbed by the electron to move into a higher level is the same as the amount of energy released when returning to the original level.

What causes the lines?

They represent the different shades of light that can be emitted as the electron of a ground state atoms jumps back from different states.

What causes flame colours?

Each colour involves a specific amount of energy released as light energy and each corresponds to a particular wavelength. A spectrum of lines is produced, some of which will be in the visible part of the spectrum.

Formula for number of electron in quantum shells

2n² (n=the shell number)

What is an orbital?

A region within an atom that can hold up to 2 electrons with opposite spins.

How many electrons can occupy s subshells?

2

How many electrons can occupy p subshells

6

How many electrons can occupy d subshells

10

Atomic radius

Decreases down a period

• Outer electrons are in the same shell

• Same amount of shielding

• So stronger attraction between nucleus and outer shell electrons

• Outer shell electrons are pulled closer to nucleus

Group 2-3 dip in ionisation energy

G3 - p orbital G2- s orbital

P orbital is higher energy than s orbital so easier to lose electron

Group 5-6 dip in ionisation energy

G6 element loses electrons from orbital with 2 electrons

G5 element loses electron from orbital with 1 electron.

Extra electron - electron repulsions make it easier to lose electron from p4 than p3

Electronegativity

Increases down period

• Increasing nuclear charge

• Little increase in shielding as electrons are added to same shell

• Decreasing atomic radius

• Increasing attraction for outer shell

Melting and Boiling points - Na, Mg, Al

Increases down period

• High as strong metallic bonding

• Al > Mg > Na as from Na to Al have smaller ions, higher charge on metal ions and more delocalised electrons

Melting and Boiling points - Si

Very high as giant covalent and need to break many strong covalent bonds to melt

Melting and Boiling points - P4, S8, Cl2

• Low as simple molecular with weak van der waal’s forces

• S8 > P4 > Cl2 as S8 has the most electrons so strongest van der waal’s forces

Melting and Boiling points - Ar

Monatomic so very low BPT as very weak van der waal’s forces between atoms.