trends across period 3

Question 1

what is meant by the atomic radius ?

Answer 1

its the distance between the nucleus and the electrons in the outer energy level

Question 2

under what circumstance would the atomic radius increase ?

Answer 2

atomic radius would increase if there were more electrons and therefore more energy levels

Question 3

in what energy level are the outer electrons in ?

Answer 3

outer electrons are in the same energy level 3

Question 4

if all the elements in period 3 have 3 energy levels, what does that do to the shielding ?

Answer 4

it would mean that they all have the same shielding

Question 5

what happens to the proton number across the period ?

Answer 5

across the period the nucleus gains protons which increases the atomic number causing the positive charge in the nucleus to increase and therefore pull the electrons closer to the nucleus

Question 6

what happens to the atomic radius across period 3 ?

Answer 6

atomic radius decreases across period 3

Question 7

what is meant by electronegativity ?

Answer 7

the power of an atom to attract the 2 electrons in a covalent bond

Question 8

what is the trend of electronegativity across period 3 ?

Answer 8

electronegativity increases across the period

Question 9

why does the electronegativity increase across the period ?

Answer 9

the nuclear charge increases and so atomic radius decreases
- the bonding pair is closer to the nucleus due to a stronger attraction of the electron pair in a covalent bond

Question 10

why is group 0 not shown on the graph ?

Answer 10

group 0 dont form covalent bonds as they have a complete shell and therfore they dont have an electronegativity

Question 11

what is the trend in electronegativity down a group ? and why ?

Answer 11

electronegativity decreases down the group
- larger atoms
- nucleus is further away from the bonding pair of electrons so there is more shielding and therefore a weaker attraction

Question 12

what is meant by ionisation energy ?

Answer 12

the enthalpy change to remove 1 mole of electrons from each atom in a mole of gaseous atoms

Question 13

what state must all the elements be when taking ionisation energy into consideration ?

Answer 13

they must be in a gaseous state

Question 14

what is the rule for ionisation energy ?

Answer 14

the stronger the attraction between nucleus and outer shell electrons, the harder it will be to remove that electron so the larger the ionisation energy

Question 15

what is the general pattern for ionisation energy ?

Answer 15

ionisation energy increases across a period because the outer electron is closer to the nucleus (atomic radius decreases ) and the nuclear charge increases

Question 16

why is there a dip between group 2 and 3 on the graph ?

Answer 16

The Mg electron is from the s orbital but Al electron is from the p orbital
the p orbital has a higher energy than the s orbital and are easier to remove because they are further from the nucleus

Question 17

why is there a dip between group 5 and 6 on the graph ?

Answer 17

phosphorus ends with 3p3 and sulfur ends with 3p4
- the 2 electrons are in the same orbital for sulfur and they are both negatively charged and will repel eachother so 1 electron is easier to remove due to repulsion

Question 18

will the first ionisation energy pattern continue for all groups ?

Answer 18

yes

Question 19

what happens to the ionisation energy when moving from one period to another ?

Answer 19

the ionisation energy decreases because you are going down a group

Question 20

why does the ionisation energy decrease down the group ?

Answer 20

• atomic radius increases so electrons are further away from the nucleus
• weaker attraction between nucleus and outer electron so its easier to remove