s1.3 electron configurations

Question 1

what is the electromagnetic spectrum?

Answer 1

the EM spectrum is the range of wavelengths, or frequencies, of EM radiation. higher energy corresponds to higher frequency and shorter wavelength. lower energy corresponds to lower frequency and longer wavelength

Question 2

what equation is related to the EM spectrum?

Answer 2

c = f x λ
the speed of light (‘c’) is constant and has a value of 3.00 x 10^8 m/s
as you can see from the spectrum, frequency (‘f’) is inversely proportional to wavelength (‘λ’)

Question 3

what is line spectra?

Answer 3

white light is made up of a range of colours. these colours can be separated by splitting white light with a glass prism to obtain a spectrum. every element in the periodic table produces a unique line spectrum when heated, consisting of specific colours at specific wavelengths seen as thin bands

Question 4

what are the three types of line spectra?

Answer 4

a continuous spectrum shows all the wavelengths, or frequencies on a coloured background
an absorption spectrum shows black lines on a coloured background
an emission spectrum shows coloured lines on a black background
each element has unique absorption and emission spectra and they can be used to identify unknown elements

Question 5

what does line spectra tell us about electrons?

Answer 5

line spectra tells us that the emitted light from atoms can only be certain fixed frequencies - it is quantised. electrons can only possess certain amounts of energy - they cannot have any energy value

Question 6

how are line spectra produced?

Answer 6

the bohr model of the atom has the protons and neutrons located in the nucleus and the electrons located in energy levels around the nucleus. electrons only exist in stationary orbits of discrete energy levels and electrons in the same energy level have the same amount of energy.
electrons can transition between energy levels by either absorbing or emitting specific amounts of energy. this energy is in the form of small packets of energy called photons.
if an electron absorbs an exact amount of energy, it will transition to a higher energy level. if an electron emits an exact amount of energy, it will transition to a lower energy level - this photo represents the energy difference between the 2 levels.

Question 7

what is the flame test?

Answer 7

the flame test is an analytical technique. it is used to identify the presence of some metals. the flame test is an atomic emission

Question 8

how is a hydrogen emission spectrum produced?

Answer 8

a hydrogen discharge tube is filled with hydrogen gas at low pressure. a spectroscope is used to split the light into its different wavelengths

Question 9

what does the hydrogen emission spectrum show?

Answer 9

electron transitions to the first energy level (n=1) emits the highest amount of energy and corresponds to the UV region of the EM spectrum.
electron transitions to the n=2 energy level emit energy that corresponds to the visible light region of the EM spectrum.
electron transitions to the n=3 energy level emit energy in the infrared region of the EM spectrum.
the longer the arrow, the greater the amount of energy emitted

Question 10

what is the principal quantum number?

Answer 10

a shell or electron level is a group of atomic orbitals with the same principal quantum number (n). the principal quantum number (n) indicated the electron level that the electrons occupy.

Question 11

what are principal energy levels?

Answer 11

electrons are located in principal energy levels (main energy levels). the first main energy level (n=1) has the lowest energy and energy increases as the value of n increases. each main energy level can hold a maximum of 2n² electrons.

Question 12

what are the sub-levels in the atom?

Answer 12

each main energy level is split into sub-levels.
-> n=1 has 1 sub level (1s)
-> n=2 has 2 sub levels (2s, 2p)
-> n=3 has 3 sub levels (3s, 3p, 3d)
-> n=4 has 4 sub levels (4s, 4p, 4d, 4f)

Question 13

what is an orbital?

Answer 13

an orbital is a region in space where one is likely to find an electron (heisenburg uncertainty principle). orbitals can hold up to two electrons as long as they have opposite spin; this is known as pauli’s exclusion principle

Question 14

what shape do orbitals take?

Answer 14

s orbital -> spherical, one in every level
p oribital -> dumb bell, three in every level from 2 up
d orbital -> various, five in levels from 3 up
f orbital -> various, seven in levels from 4 up

Question 15

what is electron configuration?

Answer 15

the arrangement of electrons in orbitals is called electron configuration. only a maximum of 2 electrons can occupy the same atomic orbital. those 2 electrons will have opposite spins - the spin is based on a magnetic property

Question 16

what is the aufbau principle?

Answer 16

the aufbau principle states that electrons are placed into orbitals of lowest energy first

Question 17

what is hund’s rule?

Answer 17

hund’s rule states that if more than one degenerate orbital in a sub-level is available, electrons occupy separate orbitals with parallel spins. always fill orbitals of equal energy with one electron first and then add the second electron once each orbital has one electron in it

Question 18

what is the electron configuration of ions?

Answer 18

metals lose electrons when they form ions. note that the first row d-block elements (sc - zn) lose their 4s electrons first when they form ions

Question 19

what is first ionisation energy?

Answer 19

first ionisation energy is the energy required to remove one electron from each atom (to infinity) in one mole of gaseous atoms to form one mole of gaseous 1+ ions

Question 20

what factors affect first ionisation energy?

Answer 20

1. nuclear charge: the greater the nuclear charge (number of protons), the greater the effective nuclear attraction experienced by the outer electrons so it requires more energy for the electron to be removed
2. atomic radius: the greater the atomic radius, the less effective nuclear attraction experienced by the outer electrons
3. electron shielding (caused by the inner shells of electrons repelling the outer electrons): the more inner shells there are, the larger the shielding effect and the smaller the effective nuclear attraction experienced by the outer electrons

Question 21

why does 1st IE decrease as you go down the group?

Answer 21

despite an increased nuclear charge, the outer shell is easier to remove due to increased shielding and greater atomic radii. the outer electron is held less strongly and is easier to remove, meaning there is less effective nuclear attraction

Question 22

why does 1st IE generally increase as you go across a period?

Answer 22

there is increased nuclear charge, a slightly decreased atomic radii and a similar amount of shielding. the outer electron is held more strongly and harder to remove, meaning there is more effective nuclear attraction

Question 23

why is there a decrease in energy in 1st ionisation energy between mg and al?

Answer 23

mg has the electron configuration 1s2 2s2 2p6 3s2 while al has the electron configuration 1s2 2s2 2p6 3s2 3p1.
the outer electron is removed from the 3p subshell in mg which has subshell shielding from the 3s subshell and is further away from the nucleus. this means there is less effective nuclear attraction

Question 24

why is there a decrease in 1st ionisation energy between p and s?

Answer 24

p has the electron configuration 1s2 2s2 2p6 3s2 3p3. s has the electron configuration 1s2 2s2 2p6 3s2 3p4.
for s, the electron is removed from a doubly occupied p orbital. the outer electron will experience electron pair repulsion in the 3p subshell, so it requires less energy to remove than an electron in a half-filled orbital

Question 25

what are converging energy levels?

Answer 25

as the energy levels increase in energy, they get closer together (the levels converge towards a limit). the highest level is sometimes referred to as the ‘infinite’ level, as the levels get so close together they are impossible to count

Question 26

what is the convergence limit?

Answer 26

the convergence limit is the point at which the spectral lines converge. at the limit of convergence the lines merge, forming a continuum. beyond the continuum, the electrons can have any energy, so the pull of the nucleus is no longer significant - meaning ionisation has occurred. we can use the frequency of radiation in the emission spectra to determine 1st ionisation energy

Question 27

how do you calculate 1st ionisation energy?

Answer 27

the energy difference between n=1 and the energy level of convergence is equivalent to 1st ionisation energy.
ΔE = hf
c = λf
IE₁ = ΔE x Nₐ

Question 28

what do the symbols in the equations for 1st IE mean?

Answer 28

ΔE = change in energy (J)
h = planck’s constant (6.626 x 10⁻³⁴ Js)
f = frequency (Hz)
c = speed of light (2.98 x 10⁸ ms ⁻¹)
λ = wavelength (m)
IE₁ = 1st ionisation energy (kJ mol ⁻¹)
Nₐ = avogadro’s constant (6.022 x 10²³)

Question 29

what is successive ionisation energy?

Answer 29

we can determine which group an element belongs to by looking at a graph of successive ionisation energies. each successive ionisation energies show an increase because as more electrons are removed, the nucleus attracts the remaining electrons more strongly. there is a large increase in ionisation energy after the first electron is removed. this tells us that the element is located in group one of the periodic table