periodicity

Question 1

Define ionisation energy

Answer 1

The energy needed to remove one mole of electrons from one mole of atoms to form one mole of positive ions in the gaseous state

Question 2

What is the first ionisation energy

Answer 2

• Electrons move around the nucleus in particular energy levels
• energy required to remove the outermost electron from an atom = first ionisation energy

Question 3

what is the equation for the first ionisation energy

Answer 3

X(g) —> X+ (g) e-

Energy change = +ve

Question 4

What is the second ionisation energy

Answer 4

• energy needed to remove the second outermost electron

Question 5

What is the equation for second ionisation energy

Answer 5

X+ (g) —> X+2 (g) + e-

Question 6

What does nuclear attraction of an electron depend on?

Answer 6

• atomic radius
• nuclear charge
• electron shielding

Question 7

What is the effect of atomic radius on nuclear attraction

Answer 7

• greater = smaller nuclear attraction experienced by outer electrons

Question 7

What effect does nuclear charge have on nuclear attraction?

Answer 7

greater = greater attractive force = more difficult to remove electron

Question 8

What effect does electron shielding have on nuclear attraction?

Answer 8

• inner shells of electrons repel the outer shell electrons
• more inner shells = larger shielding = smaller nuclear attraction of outer shells = easier to remove electron

Question 9

what is the trend across period 3 in terms of first ionisation energy?

Answer 9

generally increases because the nuclear charge increases but the shielding of the outer electrons remains relatively the same.

Question 10

what is the general trend in 1st IE as you go across a period?

Answer 10

• increase in IE1 across a period
• nuclear charge increases
• atomic radius decreases
• stronger attractive forces between the nucleus and outer electrons (stronger nuclear attraction
• similar shielding
  = harder to remove any electrons
• electron comes from same principle energy level
• size of nuclear charge increases
• more energy needed to remove outer electron

Question 11

what are the deviations in the trend across the periods?

Answer 11

between Mg - Al and P - S

Question 12

why is there a deviation between Mg-Al

Answer 12

• 1st IE of Al is less than Mg
• there is an increase in nuclear charge
• the outer electron is removed from Al from the 3p sub level which is higher in energy than 3s sub level
• less energy is needed to remove the electron from Al
• electrons are less stable in 3p + 3p is further away from 3s so smaller nuclear attraction bc increase nuclear radius

Question 13

why is there a deviation between S and P

Answer 13

• 1st IE of S is less than P
• both remove electron from 3p sub level
• but 3p electron removed from S is a paired electron
• whereas 3p electron from P is unpaired electron
• when electrons are paired there is mutual repulsion bc of like charges
• less energy needed to remove an electron = reduced ionisation energy

Question 14

what is the trend in atomic radius as you go across period 3?

Answer 14

• atomic radii decrease
• number of protons (the nuclear charge) and the number of electrons increases by one every time you go an element to the right
• elements in a period all have the same number of shells = shielding effect is the same)
• as you go across the period the nucleus attracts the electrons more strongly = pulling them closer to the nucleus
• the atomic radius= thus the size of the atoms decreases across the period

Question 15

what is ionic radius?

Answer 15

distance between the nucleus and the outermost electron of an ion

Question 16

what is the size of cations in comparison to their parent atoms?

Answer 16

lost their valence electrons = smaller
fewer electrons, this also means that there is less shielding of the outer electrons

Question 17

what is the general trend in IEs as you go down a group?

Answer 17

• general decrease
• atom gets bigger bc of increase to number of shells
• increased shielding
• increases nuclear radius
• nuclear charge smaller nuclear attraction on electrons repelling them
• easier to lose electron

Question 18

why is the first ionisation energy of sodium is less than neon but the seconds ionisation of sodium is greater than seconds ionisation of neon? EQ

Answer 18

• electron is removed from the third principle level in sodium
• electron is removed from the second principle level
• there is a greater attraction to nucleus = harder to remove
• 1st IE is less for sodium
• after this, they both have remove the electron from the 2nd principle level
• Na ion has a greater nuclear charge so more energy is needed to remove an electron from Na than Ne bc of higher nuclear attraction = greater 2nd IE

Question 19

The change in melting point from Sodium to Aluminium

Answer 19

• all metals and therefore have metallic bonding.
• more delocalised electrons present and the smaller the radius of the atom, the higher the melting point of the metal
• move across period 3 the number of delocalised electrons per metal atom increases and the radius of the elements decreases. = melting point increases.
• greater electrostatic attraction between the positive ions and delocalised electrons and hence the metallic bond is stronger and requires more energy to break.

Question 20

Explain the melting point of Silicon (Si)

Answer 20

• giant covalent structure which consists of covalent bonding, tetrahedral structure
• strong covalent bonds that hold the atoms together.
• require a large amount of energy to break and therefore silicon has a high melting point.

Question 21

The change in melting point from Phosphorous to Chlorine

Answer 21

• simple molecular substances/simple covalent which consist of van der Waals forces.
• van der Waals forces are weak intermolecular forces which require a small amount of energy to break. = melting points are low.
• melting point of these substance depends on the varying strength of van der Waals forces.
• shape of a molecule and the distance between the molecules affects the strength of induced dipole-dipole forces.
  The stronger the induced dipole-dipole forces, the higher the melting point.
• A larger molecule contains more electrons, therefore it consists of larger electron clouds. The greater the number of electron clouds = stronger the induced dipole-dipole forces.
• More energy is required to break stronger induced dipole-dipole forces, therefore the melting point is higher.
• sulfur is the largest molecule out of the three, it contains the most number of electrons and the strongest van der Waals forces.
• Therefore sulfur has the highest melting point, compared to phosphorous and chlorine.

Question 22

Explain why b.p of aluminium is higher than that of sulfur?

Answer 22

• Al is metal = strong matallic bonds holding all the atoms together
• S is a simple molecular substance = weak van der waals forces
• takes less energy to break vaan der waals than metallic bonds = b.p of Al is much higher

Question 23

Explain why the meting point of phosphorus is lower than silicon

Answer 23

• Si is macromolecule = strong covalent bonds linking all atoms together
• P is simple molecular substance = Van der Waals = less energy needed to break them than covalent
• m.p of P is less