Periodicity, Electron Config, Ionic and Metallic bonding

Question 1

what is an Ionic Bond

Answer 1

a strong electrostatic force of attraction between oppositely charged ions (between metal and non metal atoms)

Question 2

what are the properties of ionic solids

Answer 2

• giant ionic structures called lattices
• high melting point
• solid at room temp
• do not conduct electricity (do when molten or dissolved. ions can move)

Question 3

describe the bonding in MgI2

Answer 3

• Mg reacts with 2 Iodine atoms
• Mg loses 2 electrons to form a positive ion
• I atoms gain 1 electron to form a negative ion
• ions are attracted to all oppositely charges ions in the structure
• a giant structure/lattice is formed

Question 4

sodium chloride has a very high melting point, explain why

Answer 4

• sodium ions and chlorine ions electrostatically attract due to them being oppositely charged
• strong ionic bonds are between all ions as they are all attracted to all oppositely charged ions in the giant structure
• these strong bonds require lots of energy to overcome which provide a high melting point

Question 4

how are sublevels filled when atoms form ions

Answer 4

when forming ions electrons are lost from the 4s before the 3d

**special cases - **
Cu and Cr leave 4s half filled and half/fully fill 3d
Cr = 4s1 3d5
Cu = 4s1 3d10

Question 5

what is periodicity

Answer 5

repetitive patterns across a row in the periodic table

Question 6

what are the 3 factors that affect periodicity

Answer 6

1. nuclear charge (number of protons)
2. shielding (number of energy levels/shells)
3. principal energy level (distance from nucleus to outer shell)

Question 7

what does atomic radius tell us and what is effective nuclear charge

Answer 7

atomic radius tells us about **the size of atoms **- it is measures as half the distance between the centre of 2 atoms

EFC - nuclear charge that is acting on the outer electron

Zeff = Z - S

Z = proton number S = no. of outer electrons

Question 8

how do the 3 factors of periodicity affect atomic radius

nuclear charge, shielding, principal energy levels

Answer 8

• as nuclear charge increases, atomic radius decreases
• higher energy levels means atomic radius increases
• an increase in shielding means atomic radius increases

Question 9

what happens to atomic radius across a period

Answer 9

1. nuclear charge INCREASES
2. principal energy level STAYS THE SAME
3. shielding STAYS THE SAME

overall it decreases

Question 10

what happens to atomic radius as you go down a group

Answer 10

1. nuclear charge INCREASES
2. principal energy level INCREASES
3. shielding INCREASES

overall it increases

Question 11

what happens to ionic radius as you go down a group

Answer 11

1. nuclear charge INCREASES
2. principal energy level INCREASES
3. shielding INCREASES

overall it increases

Question 12

what happens to ionic radius across a period

Answer 12

1. nuclear charge INCREASES
2. principal energy level STAYS THE SAME among positive ions and the same among negative ions
3. shielding STAYS THE SAME among positive ions and the same **among negative ions **

overall it decreases

Question 13

is the positive ion of an element smaller or larger than an atom of it

Answer 13

the positive ion of an atom is smaller due to reducing its number of principal energy levels and shielding whilst having the same nuclear charge

Question 14

is the negative ion of an element smaller or larger than an atom of it

Answer 14

the negative ion of an atom is larger due to the number of principal energy levels and shielding staying constant whilst having the same nuclear charge split among more electrons

Question 15

what is first ionisation energy and show the equation to represent it

Answer 15

it is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state and is measures in KJmol-1

X (g) ———> X+ (g) + e-

Question 16

what is second ionisation energy and show the equation to represent it

Answer 16

it is the energy required to remove one mole of electrons from one mole of 1+ ions in the gaseous state

X+ (g) ———-> X2+ (g) + e-

Question 17

how does the amount of energy change through successive ionisation energies

Answer 17

As successive ionisation energies occur, the amount of energy required increases in general, until you get to a different principal energy level, where a large increase in ionisation energy occurs
e.g when Al goes from losing 3 to 4 electrons

Question 18

how does ionisation energy change as you go down a group

Answer 18

As you go down a group, 1st ionisation energy decreases. Although nuclear charge is increasing, so are the principal energy levels and shielding, making it easier to remove an outer electron

Question 19

how does ionisation energy change as you go across a period

Answer 19

as you go across, first ionisation energy generally increases, however removing the first electron from a sublevel usually requires less energy than removing one from its complete sub level
e.g easier to lose the 2p1 from Al than the 2s2 in Mg

also happens for half filled p and d sub levels:
Phosphorus is half filled, Sulfur has a pair, electrons repel, more easily lost

Question 20

why does ionisation energy increase across period 3

Answer 20

It increases due to an increase in nuclear charge. This means more energy is required to remove electrons are they are more strongly attracted to the nucleus due to more protons being there, so a larger positive charge

Question 21

describe the structure of a metal

Answer 21

positive metal ions arranged in neat regular layers in a lattice, surrounded by a sea of delocalised electrons which they are attracted to

Question 22

what is charge density

Answer 22

the ratio of an ions charge to its volume

e.g a high charge density would be something with a large nuclear charge but a small radius

Question 23

why does magnesium have stronger metallic bonding than sodium

Answer 23

• a magnesium ion has a higher nuclear charge than a sodium ion (forms more delocalised electrons)
• but has a lower ionic radius than sodium
• this gives magnesium a greater charge density, so a higher melting point