topic 1: atomic structure and the periodic table

define isotope

elements with the same number of protons but different number of neutrons

define relative atomic mass

the weighted mean mass of an atom of an element, compared to 1/12th mass of an atom of carbon 12

define relative isotopic mass

the mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon 12

what must your abundance add up to

100%

formula for relative atomic mass

(abundance A x m/z of A) + (abundance of B x m/z of B) all over total abundance

how do you calculate isotopic mass

do relative atomic mass and then solve for x

why are arrows in electronic configuration in different directions

spin pairing. when 2 electrons occupy one orbital, they ‘spin’ in opposite directions

why do you fill orbitals singly first then pair them up

electron repulsion

why would a transition metal with the electronic configuration of 3d5 4s1 form a unipositive cation with the configuration 3d5 4s1 and not 3d4 and 4s2

lose from the 4s orbital first then 3d as 4s has higher energy when filled but lower energy when empty

what is n=1 in the atomic emission spectra

ground state

why are line spectra used

to identify elements and it is evidence for quantum shells

what is a series in atomic emission spectra

a group of lines

why do lines in atomic emission spectra get closer together

the energy and frequency increases

what does the line spectrum show

the frequency of light in coloured bands

where must an electron fall to for a line to be produced at uv

ground state, n=1

where must an electron fall for a line to appear at visible light in atomic emission spectra

n=2

where will the line in atomic emission spectra appear if n=3

infrared

define first ionisation energy

the minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state to form one mole of unipositive gaseous cations

ionisation energy trends down a group

• decreases down groups

• atomic radius increases

• outer electrons are further from nucleus, so there is a weaker attractive force, therefore less energy is required

• shielding increases so less energy is required to remove electrons

what is successive ionisation

the removal of more than 1 electron form the same atom

what is the general trend of successive ionisation

there is a general increase in energy as removing electrons causes the atom to become increasingly more positive

what happens to the atomic radius across period 3

• decreases

• there is an increased nuclear charge because there is an increase in protons

what happens to the atomic radius down a group

increases due to extra electron shells being added

what is the trend of ionisation energies across periods

• increases due to increasing number of protons and nuclear attraction increases

• shielding is similar and distance marginally decreases

• more energy is required to remove outer electrons

why is there a decrease at the 3rd element in the ionisation energies across periods

element sits in a higher energy sub shell so it is slightly further form the nucleus and there is more shielding from s, the p orbitals are higher in energy so less energy needed to remove an electrons

why is there a decrease at the 6th element in ionisation energies across periods

• electronic configuration p3 to p4 which means the outer electron is paired in its orbital

• means there is more repulsion between the electrons and so they are higher in energy and therefore require less energy to remove

what is the general trend of melting points across period 3 metals

increases as metal ions have a positive charge, increasing the number of delocalised electrons and a smaller ionic radius

why does silicon have the highest melting point in period 3

due to its giant covalent structure

why is argon the lowest melting point in period 3

argon only exists as atoms