topic 1: atomic structure and the periodic table

Question 1

what is the definition of isotopes

Answer 1

Isotopes are atoms of the same element with the same number of protons, but different numbers of neutrons

Question 2

what is the definition of relative isotopic mass

Answer 2

the mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12

Question 3

what is the definition of relative atomic mass

Answer 3

the average weighted mass of an atom of an element compared to one twelfth of the mass of one atom of carbon-12

Question 4

what is the definition of relative molecular mass

Answer 4

the average mass of a molecule
compared to one twelfth of the mass of one atom of carbon-12

Question 5

formula for RAM

Answer 5

R.A.M =  (isotopic mass x % abundance)/100

Question 6

what are the uses of mass spectrometers

Answer 6

• have been included in planetary space probes so that elements on other planets can be identified. elements on other planets can have a different composition of
  isotopes
• drug testing in sport to identify chemicals in the blood and to identify breakdown products from
  drugs in body
• quality control in pharmaceutical industry and to identify molecules from sample with potential
  biological activity
• radioactive dating to determine age of fossils or human remains

Question 7

what is the first ionisation energy

Answer 7

the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge. H(g) –> H+
(g) + e-

Question 8

what is the second ionisation energy

Answer 8

the energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge. Ti+(g) —> Ti2+(g) + e-

Question 9

what are the factors that affect ionisation energy

Answer 9

1.the attraction of the nucleus
(the more protons in the nucleus the greater the attraction)
2. the distance of the electrons from the nucleus (the bigger the atom the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus)
3. shielding of the attraction of the nucleus (an electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus)

Question 10

why has helium the largest first ionisation energy

Answer 10

its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisationenergy than H as it has one more proton

Question 11

why do first ionisation energies decrease down a group

Answer 11

as one goes down a group, the outer electrons are found in shells
further from the nucleus and are more shielded so the attraction of
the nucleus becomes smaller

Question 12

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

Answer 12

as one goes across a period , the number of protons increases making the effective attraction of the nucleus greater. the electrons are being added to the same shell which has the same shielding effect and the electrons are pulled in closer to the nucleus

Question 13

why has Na a much lower first ionisation energy than Neon

Answer 13

this is because Na will have its outer electron in a 3s shell further from the nucleus and is more shielded. so Na’s outer electron is easier to remove and has a lower ionisation energy

Question 14

why is there a small drop from Mg to Al

Answer 14

Al is starting to fill a 3p sub shell, whereas Mg has its outer electrons in the 3s sub shell. the electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s
electrons

Question 15

why is there a small drop from P to S

Answer 15

with sulphur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill the first 3p orbital. when the second electron is added to a 3p orbital there is a slight repulsion between
the two negatively charged electrons which makes the second electron easier to remove

Question 16

what is the order of orbitals

Answer 16

An atom fills up the sub shells in order of increasing energy (note 3d is higher in energy than 4s and so gets filled after the 4s
1s2s2p3s3p 4s3d4p5s4d5p

Question 17

what is the shape of s-subshell

Answer 17

spherical

Question 18

what is the shape of p-subshell

Answer 18

dumbbells

Question 19

how do you draw a spin diagram

Answer 19

2 arrows going in the opposite directions in a box

Question 20

how much electrons does each subshell fill

Answer 20

s - 2 electrons
p - 6 electrons
d - 10 electrons
f - 14 electrons

Question 21

what is the periodicity

Answer 21

the repeating pattern of physical or chemical properties going across the periods

Question 22

what happens to the atomic radii across a period

Answer 22

decreases ,because the increased number of protons create more positive charge attraction for
electrons which are in the same shell with similar shielding

Question 23

what happens to the 1st ionisation energy across a period 3

Answer 23

increases due to increasing number of protons as the electrons are being
added to the same shell

Question 24

what happens to the 1st ionisation energy between Mg and Al

Answer 24

there is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al’s electron is slightly easier to remove because the 3p electrons are higher in energy

Question 25

what happens to the 1st ionisation energy in between P and S

Answer 25

there is a small drop between phosphorous and sulfur. sulfur’s outer electron is being paired up with an another electron in the same 3p orbital. when the second electron is added to an orbital there is a slight repulsion between the two negatively charged
electrons which makes the second electron easier to remove

Question 26

what is the trend of melting and boiling point across period 3

Answer 26

• for Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. a smaller sized ion with a greater positive charge also makes the bonding stronger. higher energy is needed to break metallic bonds.
• Si is Macromolecular: many strong covalent bonds between
  atoms high energy needed to break covalent bonds– very high
  mp +bp
• Cl2 (g), S8 (s), P4 (S)- simple molecular : weak London forces
  between molecules, so little energy is needed to break them –
  low mp+ bp
• S8 has a higher mp than P4 because it has more electrons (S8
  =128)(P4=60) so has stronger London forces between
  molecules
• Ar is monoatomic weak London Forces between atoms

Question 27

describe the trend of melting and boiling points across period 2

Answer 27

similar trend to period 3

Question 28

atomic emission spectroscopy provides evidence for the existence of

Answer 28

quantum shells

Question 29

state why the arrows are all pointing in the same direction in the 3p boxes

Answer 29

three electrons with parallel/ same spin/ direction of rotation (because the electrons are in different orbitals)

Question 30

state what is meant by the 2 arrows in the boxes

Answer 30

up and down arrows represent electrons with opposite spins
OR
2 electrons in the same orbital with opposite spins

Question 31

explain why iodine and chlorine have many similar chemical reactions

Answer 31

iodine has 7 electrons in the outer shells (same as chlorine
electronic configurations/ number of electrons in the outer shell govern their chemical reactions

Question 32

explain why the first ionisation energy of magnesium is higher than that of sodium

Answer 32

Mg has more protons than Na
Mg has similar shielding to Na/ outer electron in Mg in the same sub-shell in a Na atom
so the force of attraction between the nucleus and the electron is greater in Mg than in Na

Question 33

a student suggested that the difference in the rates of reaction of strontium and barium with water is due to the difference in the sum of their first and second ionisation energies.
discuss this suggestion

Answer 33

the sum of the first 2 ionisation energies for barium are lower
barium is a bigger atom
has more shielding
barium has more proton
reacts faster

Question 34

describe how the graph showing the first ionisation energies of period 3 elements change

Answer 34

increase from Na to Mg
Al between Mg and Na
rise from Al to P
dip at S (between Si and P)
rise from S to Ar

Question 35

explain the trend in the values of the first ionisation energies for the group containing
sulfur

Answer 35

although the number of protons is increasing, the electron being removed is further from the nucleus as there is an increase in shells. this gives more shielding from the nucleus

Question 36

explain why the first ionisation energy of sulfur is lower than that of chlorine

Answer 36

in sulfur the proton number is on less than chlorine and the electron being removed is from the same subshell/ similar shielding

Question 37

explain why the first ionisation energy of sulfur is lower than that of phosphorus

Answer 37

in sulfur, spin pairing has occured/ electron being removed from an orbital containing 2 electrons
results in an increase in repulsion between electrons (so the electron is lost more easily)

Question 38

explain why the first ionisation energy of hydrogen is less than that of helium, but greater
than that of lithium

Answer 38

He has more protons than H
in He, the outer electron is in the same shell as H
in Li, the outer electron is in a higher energy level further away from the nucleus and is shielded by inner electrons

Question 39

explain why the first ionisation energy of sodium is greater than that of potassium

Answer 39

outer electron in sodium is closer to the nucleus than that in K and there is less shielding from inner electron shells. these outweigh the greater charge in potassium

Question 40

explain why the difference between the second and third ionisation energies of calcium is much larger than the difference between the first and second ionisation energies

Answer 40

3rd electron is lost from 3p orbital which is closer to the nucleus
1st and 2nd electrons removed from the same orbital

Question 41

give the meaning of the term ‘periodicity’.
illustrate your answer by referring to the atomic radii of the Period 2 and Period 3 elements. specific values of atomic radii are not required

Answer 41

a trend/ pattern of repeating physical and chemical properties
atomic radii decrease across the period
the pattern/ atomic radius trend is repeated in period 3

Question 42

explain the trend in melting temperatures across the elements of Period 2 in terms of their
structure and bonding

Answer 42

at the start of the period Li to Be, the bonding is metallic. the metallic bonding gets stronger as the number of delocalised electrons, radius of cation and charge of cation increases in metal atom
in the middle of the period, C has a giant structure of atoms which requires a lot of energy to break strong covalent bonds, (in graphite and diamond)
at the end of the period, N to Ne are simple molecules which contain weak intermolecular/ London forces (between molecules)