T1 atomic structure and the periodic table

topic 1

relative mass of an electron

relative mass of an electron

1/2000

isotope definition

atoms of an element with the same number of protons but diff num of neutrons

relative atomic mass definition

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

relative isotopic mass

it is the mass of an atom of an isotope compared to 1/12th of the mass of a carbon-12 atom

relative atomic mass formula

(abundance A x Mr(A)) + (abundance B x Mr (B))
total abundance

when should term 'relative formula mass' be used instead of 'relative molecular mass'

used for compounds with giant structures

state how relative abundance of two isotopes can be found

compare intensity of signal/ number of particles of each isotope detected
in mass spectrometer

predicting mass spectra for diatomic molecules

1. make abundances decimals out of 1

2. identify the abundances of all possible isotope combinations using a table and multiplying them. all isotopes both across and down to form multiplication table

3. any molecules that are the same, add the abundances together

4. times all abundances by 100 to find abundance as a percentage, or divide all abundances by the smallest abundance to find a whole number ratio

5. plot results using Mr as x axis and abundance as y axis

what is the M+1 peak?

aka molecular ion peak
it is the last peak on the spectra
it shows the Mr unfragmented molecule
(+1 only refers to charge)

explain how atomic emission spectra provides evidence for ideas on electronic configurations

atomic emission spectra provide evidence for the existence of quantum shells
clear lines/ frequencies for different energy levels

sub-shell

s: 1 orbital = can hold 2 electrons
p: 3 orbitals = 6
d: 5 orbitals = 10
f: 7 orbitals = 14
only until 4p⁶ needed

orbital

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

subshell orbital shapes

s: 1 orbital, spherical
p: 3 orbitals, dumbell shape, perpendicular to each other, x,y,z axes

spin pairing

when 2 electrons occupy 1 orbital they spin in opposite directions

energy levels of subshells

we fill from lowest to highest
e⁻s fill subshells singly before pairing up due to e⁻ repulsion

which cells fill up first in a transition metal?

4s
except Cr and Cu

Cr and Cu electronic configuration

4s subshell donates e⁻ to form full/singly filled 3d subshell
Priority: to have a perfect 3d subshell: more stable

In which order do transition metals lose electrons to form ions?

always 4s first, then 3d

first ionisation energy

energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form one mole of gaseous 1+ ions.
(energy is req to overcome electrostatic attr between nucleus and e⁻)

successive ionisation energies

energy required to remove 1 electron from each ion in 1 mole of gaseous (1+) ions to form one mole of gaseous (2+) ions
(second ionisation energy)

explain trend in values of 1st ionisation energies down group

• first ionisation energy decreases down group

• bc although num of protons is increasing

• outer e⁻ is one shell of e⁻s further from the nucleus

• larger atomic radius and more shells of electrons --> more shielding

• this outweighs/ exerts a greater influence than greater positive/nuclear charge
  
  provides evidence for e⁻ shells

explain why the 1st ionisation energy of generally increases across period

• bc atomic number incrs by one

• and e⁻ removed is from same (sub)shell and has similar shielding

  sub necessary if increasing

explain how different factors influence ionisation energies

1. number of protons
   greater nuclear charge = stronger force of attraction from nucleus to (outer) e⁻

2. electron shielding

   • outer e⁻ in same quantuum shell = similair levels of shielding

   • outer e⁻ in higher energy level = larger atomic radius and more shells of electrons --> more shielding, drop in force of attraction between nucleus and (outer) e⁻
     this outweighs/ exerts a greater influence than greater positive/nuclear charge

3. electron subshell from which electron is being removed

explain drop between group 6 after group 5 in 1st ionisation energies

• fully singly filled/ fully filled subshells are more stable than partially filled ones

• so have higher ionisation energies

• in group 6 spin pairing has occured

• resulting in an increase in repulsion between electrons

• so electron lost more easily

explain drop between for group 3 after group 2 in 1st ionisation energies

1. electron removed from 3 is from new subshell, p instead of s

2. extra shielding and distance from nucleus

3. this outweighs increase in nuclear charge
   --> evidence for electron subshells

explain big jump in successive ionisation energies

"new shell broken into":
e⁻ lost from shell closer to nucleus
previous from same shell w similar shielding

why do successive ionisation energies increase

1. same number of protons attracting a decreasing number of electrons

2. decreasing repulsion amongst remaining electrons

what determines the chemical properties of an element

electronic configuration

• same number of electrons in outer shell

• electronic configuration / n of e⁻ in outer shell govern their chemical reactions

explain periodicity in terms of a repeating pattern across different periods

a trend/ pattern of repeating physical and chemical properties with increasing atomic number

explain pattern in atomic radii across a period

• atomic radii decrease across period

• proton number and nuclear charge increases

• attraction increases

• extra electrons gained across period added to outer energy level so do not provide extra shielding

explain reason for trend in melting/ boiling temps for elements in a period

1. at start of period bonding is metallic

2. metallic bonding gets stronger as number of delocalised electrons in a metal increases/ charge on cation increases

3. middle of period (B and C) have giant covalent structure

4. a lot of energy is required to break covalent bonds

5. at the end of the period elements form simple molecules

6. with weak london forces between them

7. noble gases are monoatomic resulting in very weak london forces