PERIODICITY

Question 1

Who created the periodic table that we use today and how is it ordered

Answer 1

Mendeleev ordered the elements in order of atomic mass and grouped the elements in terms of similar chemical properties

today is is ordered by proton number

Question 2

Groups in the periodic table

Answer 2

Elements in the same column will have the same number of electrons in the outer shell (group number relates to number of electrons in the shell)

will have similar properties and will usually increase in reactivity as you move down the group

Question 3

Periods in the periodic table

Answer 3

elements in the same period have the same number of electron shells

Question 4

What is ionisation energy

Answer 4

the minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state

Question 5

1st ionisation energy of sodium equation

Answer 5

Na(g) –> Na+(g) + e-

Question 6

is ionisation energy +ve or -ve

Answer 6

always endothermic (+ve) as energy is required

Question 7

What is electron shielding

Answer 7

the more electron shells between the outer shell electron being removed and the positive nucleus means less energy will be required as there is less nuclear attraction

Question 8

How does atomic size effect ionisation energy

Answer 8

The larger the size of the atom, the further away the electron being removed is from the nucleus, so less attraction between positive and negative charges, less attractive force

Question 9

How does nuclear charge affect ionisation energy

Answer 9

the larger the amount of protons in the nucleus, the bigger the attraction between it and the outer electrons, this means more energy will be required to remove the outer electron

Question 10

ionisation trends in groups explained

Answer 10

as you move down the groups, atomic size and the number of electron shells increase, so this will decrease nuclear attraction as the distance between the outer shell electrons and the nucleus increases and therefore lowers the energy needed to remove the outer shell electron

Question 11

Ionisation moving across a period explained

Answer 11

as you move across a period, the number of protons in the nucleus will increase, while atomic size and electron shielding remains roughly the same so ionisation energy will generally increase as more energy will be required to remove the outer shell electrons

Question 12

Why does Aluminium have a lower ionisation energy than Magnesium and Sulphur than Phosphorous

Answer 12

Aluminium’s most outer electron sits in a higher energy sub shell than magnesium, which is slightly further from the nucleus , so nuclear attraction will be weaker and less energy will be required to remove it.

Sulphur has an outer electron in a p orbital with another electron, so they will repel each other due to like charges and will require less energy to remove than in phosphorous, where all 3 outer electrons sit in their own orbital

Question 13

2nd ionisation of magnesium equation

Answer 13

Mg+(g) –> Mg2+(g) + e-

Question 14

trend of successive ionisations of elements

Answer 14

general increase in ionisation energy, and graph will show jumps upwards as the electron shell number decreases, as the electrons being removed will become closer to the nucleus so attraction will be stronger

Question 15

2 examples of giant covalent structures and their properties

Answer 15

Graphite has each carbon bonded 3 times with every 4th electron delocalised,

-lots of covalent bonds means it has a high melting point and it is insoluble

-layers slide easily as there are weak forces between them

• the delocalised electrons between layers allow graphite to conduct electricity as they can carry charge

Diamond has each carbon bonded 4 times in a tetrahedral shape

• conducts heat well as is arranged very tightly and rigidly
• has a very high melting point as has very strong covalent bonds
• will not conduct electricity as has no delocalised electrons

SILICON ALSO HAS A TETRAHEDRAL COVALENT STRUCTURE

Question 16

Structure of graphene and uses

Answer 16

1 layer thick of graphite

-hexagonal carbon rings
- lots of delocalised free electrons makes it a good conductor of electricity
-is very light weight and transparent as only one atom thick
- used for aircraft shells and super computers and phone screens

Question 17

Describe metallic bonding and the melting points and conductivity

Answer 17

• Metals have giant metallic lattices in which positive metal ions will form as metals will donate electrons to form a sea of delocalised electrons
• the delocalised electrons can transfer kinetic energy and also act as mobile charge carriers so metals are good thermal and electrical conductors
• there’ll be an electrostatic attraction between the positive metal ions and the delocalised electrons meaning the melting points are high
• the more electrons a metal can donate, the higher the melting point , e.g. magnesium is higher than sodium

Question 18

melting point trend of period 3 in detail for each molecule

Answer 18

Magnesium > sodium as will donate 2 electrons into the sea of delocalised electrons in the giant metallic lattice

Aluminium is similar to Magnesium and for the 3 metals there’ll be a general increase due to increasingly positive charge of each ion and number of delocalised electrons

Silicon > Aluminium due to its giant covalent structure which is the same as diamond so lots of strong covalent bonds will need to be broken to melt silicon

Phosphorous is much lower as it has a weaker simple structure, and induced dipole-dipole forces will need to be broken to melt it (P4)

Sulphur will be higher than phosphorous slightly, as even though its still got a weak simple molecular structure, it has a larger structure, so more induced dipole-dipole interactions (S8)

Chlorine is lower than sulphur as its a much smaller molecule (Cl2) so will have a lot less induced dipole-dipole forces and will be easier to melt

Argon molecules only contain one argon atom so it will have tiny induced dipole-dipole forces and the lowest melting point of period 3

TREND IS EXACT SAME FOR PERIOD 2

Question 19

Summaries and examples of giant covalent bonding

State, electrical conductivity, soluble, melting or boiling

Answer 19

Graphite, diamond and SiO2

-solid at room temp

-only graphite will conduct electricity when its a solid

• none conduct as a liquid
• not soluble in water
• high melting point as difficult to break covalent bonds

Question 20

Summary and examples of simple covalent structure

State, electrical conductivity, soluble, melting or boiling

Answer 20

I2 NH3 H2O

• liquid or gas at room temp, iodine is solid
• dont conduct electricity as a liquid or solid
• if molecule is polar it can dissolve in water
• low melting and boiling points as no bonds have to be broken, only weak intermolecular forces

Question 21

Summary and examples of giant ionic structures

State, electrical conductivity, soluble, melting or boiling

Answer 21

NaCl, CaO, MgBr2

• solid at room temp
• will conduct electricity as a liquid as ions are free to move
• are soluble in polar substances
• high melting/ boiling points, stong electrostatic forces have to be broken