Chemistry - Chapter 4: Electronic Structure of Atoms and Arrangement of Electrons

gah.

isotope

atoms of the same element with the same atomic number but different mass numbers, due to different numbers of neutrons in the nucleus

mass spectrometer

instrument used to measure the Ar of an element and is used to identify the presence of isotopes and measure their relative abundance.

1. vaporisation

with an oven, samples are heated to ensure they are in a gaseous form

2. ionisation

electron gun with high energy electrons are bombarded at the sample, knocking electrons out of the atoms/molecules to produce positive ions

3. acceleration

done using oppositely charged plates with a hole in the centre to allow ions to pass through. Accelerates ions down the tube

4. separation/deflection

applies a force to the bend in the tube.

3 things can happen:

• ion is too heaving and continues into the wall of the spectrometer

• ion is too light and is bent too far around to the wall

• ion is the correct ratio to the electromagnet and goes straight down

5. detection

electronic detector works out how many ions with corresponding mass charge ratio are hitting it at any one time. By varying charge on electromagnet, we can work out how many of each ion is present in the sample

continuous spectrum

when white light is passed through a prism, it is split into a continuous spectrum of colours corresponding to every energy from red to violet, which appears like a rainbow, which is the visible part of the spectrum

emission spectra

• Refers to the light given out: the light emitted from a heated gaseous element gives rise to a line spectrum.

• Every element has its own unique line spectrum. The fact that the spectrum consisted of a series of lines indicated to him that only certain energy emissions are possible.

• Bohr’s experiment: by using light from a hydrogen discharge tube, he passed it through a prism. He noticed that only lines of a few wavelengths (colours) were present in the resultant spectra. The coloured lines were separated by black regions, which correspond to wavelengths (colours) that were absent in the light. He concluded that H atoms cannot possess all levels of energy but only some, i.e. not all colours but only some will appear

The Bohr Theory

• electrons orbit the nucleus in fixed paths called orbits (which is a 2d structure) and occupy fixed energy levels

• in a normal H atom, electron in its ground state is n=1. When it absorbs enough energy through heat or electricity, it moves up to n=2, then n=3, and so on.

• electrons in excited state are extremely unstable and temporary, therefore it falls back to a lower level.

• when falling back down, it emits a definite amount of energy in the form of photons of light

• each transition has a definite energy and appears as a line of a particular colour in the line emission spectrum. Hence each line has a definite frequency. This indicates that only a limited number of energy changes are possible within the structure of an atom

E2 - E1 = hf

h = planks constant

f = the frequency of light emitted

energy level

the fixed energy value that an electron in an atom may have

excited state

atoms that absorbs enough energy, making it contain higher energy levels than those at ground state

ground state

atoms that occupy the lowest available energy level. When dropped to this state, photons of light are emitted

atomic absorption spectrum (absorption line)

• atoms can also absorb light

• if white light is passed through an element in its gaseous form, the light that comes out has wavelengths missing. These missing wavelengths appear as dark lines and have the same energy as would appear in the emission spectrum

• the amount of light absorbed is proportional to the concentration of the element

• therefore atoms in the ground state can absorb the same amount of energy as they would emit in the excited state

energy level

a fixed amount of energy of an electron in an atom

orbital

region of space around the nucleus of an atom, where there is a high probability of finding an electron. Each orbital can hold 2 electrons

sublevel

a group of atomic orbitals within an atom, all of which have the same energy

Aufbau principle

this states that electrons fill orbitals of lowest available energy level first

Hund’s Rule of Maximum Multiplicity

states that when two or more orbitals of equal energy are available, electrons fill them singly before they fill them into pairs

Pauli’s Exclusion Principle

states that electrons in the same orbital have opposite spins / no more than two electrons can occupy an orbital and this they can only do if they have opposite spins

Heisenberg’s Uncertainty Principle

it is impossible to measure both the velocity and the position of an electron at the same time as electrons move in a wave motion

Erwin Schrodinger

worked out the probability of finding an electron to any particular sublevel using s, p, d, f

transition elements

elements with partially filled d-orbitals in one or more of their common oxidation states

transition metal + their characteristics

one that forms at least one ion with a partially filled d sublevel

characteristics:

• they show variable valency

• they form coloured ions/compounds

• they act as catalysts in many reactions

ion

a charged particle due to the loss of or gain of an electron(s)

discharge tube

used to generate atomic emission spectra from gases. It’s a simple apparatus made with two metal electrodes connected to a high voltage potential difference

the principle of mass spectrometry

the principle involved is that different ions are separated according to their masses (mass/charge ratio) when moving in a magnetic field

light series

• n=1 → Lyman series. These exist in the ultraviolet band

• n=2 → Balmer series (what we get in the atomic emission spectrum of Hydrogen as it is in the visible band

• n=3 → Paschen series. These exist in the infra red band

why do different elements had unique atomic spectra?

each element has a different arrangement of energy levels and a different electron configuration, giving rise to different electron transitions from higher to lower energy levels

absorption vs emission spectrum

absorption: a series of dark lines against a coloured background

emission: a series of coloured lines against a dark background

limitations of Bohr’s theory

• only accounts for the Hydrogen element

• doesn’t take wave-particle duality into account

• does not allow for uncertainty (probability)

• does not explain the discovery of sublevels

Louis de Broglie

discovered wave-particle duality