Module 3 - Periodicity

Question 1

how is the periodic table arranged

Answer 1

-by atomic (proton) number

Question 2

define periodicity

Answer 2

-the regular repeating patterns in the physical and chemical properties of elements

Question 3

define ionisation

Answer 3

-when an atom loses an electron from its outer shell

Question 4

define first ionisation energy

Answer 4

-the energy required to remove 1 mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

Question 5

describe how to read successive ionisation energy graphs

Answer 5

-where the first big jump is shows the number of electrons in the outer shell
-this determines the group

Question 6

describe how to draw a successive ionisation graph for an element

Answer 6

-ionisation energy on y axis
-number of electrons on x axis
-ionisation energy slowly increases for electrons in the same shell
-big jump when moving into the next shell

Question 7

name the factors affecting ionisation energy

Answer 7

-shielding
-nuclear charge
-atomic radius

Question 8

define shielding

Answer 8

-the effect of the inner electrons protecting the outer electrons from the effect of the nuclear charge

Question 9

describe the trend in shielding across periods (eg period 3)

Answer 9

-the shielding stays the same

Question 10

describe the trend in nuclear charge across periods (eg period 3)

Answer 10

-nuclear charge increases

Question 11

define nuclear charge

Answer 11

-the positive charge on the nucleus

Question 12

define atomic radius

Answer 12

-the distance between two nuclei of atoms, then halving the distance (to measure the radius of an atom)

Question 13

describe the trend in atomic radius across a period (eg period 3)

Answer 13

-decreases

Question 14

explain the trend in atomic radius across a period (eg period 3)

Answer 14

-nuclear charge increases
-shielding stays the same
-nuclear attraction increases
-so atomic radius decreases

Question 15

describe the trend in first ionisation energy across a period (eg period 3)

Answer 15

-increases

Question 16

explain the trend in first ionisation energy across a period (eg period 3)

Answer 16

-nuclear charge increases
-atomic radius decreases
-nuclear attraction increases
-shielding stays the same
-it takes more energy to remove the first electron

Question 17

equation for ionisation energy (eg magnesium)

Answer 17

1st : Mg(g) -> Mg+(g) + e-
2nd : Mg+(g) -> Mg2+(g) + e-

Question 18

explain why the first ionisation energy goes down between magnesium and aluminium

Answer 18

-aluminium has one electron in a higher sub shell (3p)
-that one electron is removed more easily as it is further away from the nucleus
-therefore the first ionisation energy is lower than magnesium

Question 19

explain why the first ionisation energy goes down between phosphorous and sulphur

Answer 19

-sulphur has one 3p orbital that contains a pair of electrons
-these paired electrons repel each other, so one of these electrons is easier to remove
-therefore sulphur has a lower first ionisation energy than phosphorous

Question 20

describe the trend in shielding down a group

Answer 20

-increases

Question 21

describe the trend in nuclear charge down a group

Answer 21

-increases

Question 22

describe the trend in atomic radius down a group

Answer 22

-increases

Question 23

explain the trend in atomic radius down a group

Answer 23

-shielding increases
-nuclear charge increases but it is cancelled out by the extra shielding
-so nuclear attraction decreases
-so atomic radius increases

Question 24

describe the trend in first ionisation energy down a group

Answer 24

-decreases

Question 25

explain why the first ionisation energy decreases down a group

Answer 25

-more shielding -nuclear charge increases -nuclear attraction of electrons to the nucleus decreases (as the increased nuclear attraction is cancelled out by the increased shielding) -atomic radius increases -so less energy is needed to remove the outer electron

Question 26

describe the trend in melting point in period 3

Answer 26

-increases from Na (group 1) to Si (Group 4) -then it decreases

Question 27

describe the types of bonding across period 3

Answer 27

1. metallic bonding - Na, Mg, Al 2. giant covalent - Si 3. simple covalent - P, S, Cl, Ar

Question 28

explain why the melting and boiling points rise across the three metals in period 3

Answer 28

-the charge on the metal ions increases from +1 to +3 -so the number of delocalised electrons increases -so the strength of the metallic bonding increases -so the melting/boiling points increase as more energy is needed to break stronger metallic bonds

Question 29

explain why silicon has the high melting/boiling point in period 3

Answer 29

-silicon is a metalloid -it has a giant covalent structure exactly the same as carbon in diamond (each silicon is bonded to four other silicons) -silicon has a high melting point -you have to break strong covalent bonds to melt it -so a lot of energy is needed to break

Question 30

what does phosphorous exist as

Answer 30

P4

Question 31

what does sulphur exist as

Answer 31

S8

Question 32

what does chlorine exist as

Answer 32

Cl2

Question 33

what does argon exist as

Answer 33

Ar atoms

Question 34

explain why the simple molecular elements in period 3 have low melting points

Answer 34

-when these 4 substances melt/boil, the London force are broken -these are very weak bonds so little energy is needed to overcome them -the strength of the London forces increases with the number of electrons

Question 35

describe the order of melting point decrease in simple molecular elements in period 3

Answer 35

S8 > P4 > Cl2 > Ar

Question 36

define metallic bonding

Answer 36

-the electrostatic force of attraction between the positive metal ions and the delocalised electrons

Question 37

name the factors that affect the strength of metallic bonding

Answer 37

1. number of protons/ strength of nuclear attraction -the more protons the stronger the bond 2. number of delocalised electrons per atom (the outer shell electrons are delocalised) -the more delocalised electrons the stronger the bond 3. size of ion -the smaller the ion, the stronger the bond

Question 38

structure of giant metallic structures

Answer 38

-giant metallic lattice -lattice of positive metal ions -electrons

Question 39

boiling and melting points of giant metallic structures

Answer 39

-high: strong electrostatic forces between positive ions and sea of delocalised electrons

Question 40

solubility in water of giant metallic structures

Answer 40

-insoluble

Question 41

conductivity when solid of giant metallic structures

Answer 41

-good: delocalised electrons can move through structure

Question 42

conductivity when molten of giant metallic structures

Answer 42

-good

Question 43

general description of giant metallic structures

Answer 43

-shiny metal -malleable as the positive ions in the lattice are all identical -so the planes of ions can slide easily over one another -attractive forces in the lattice are the same whichever ions are adjacent