Periodicity

Question 1

What is the covalent radius?

Answer 1

It’s the measure of the size of an atom. It’s half the distance between two nuclei in a bond.

Question 2

How are elements arranged on the periodic table?

Answer 2

in order of increasing atomic number

Question 3

How are the groups arranged?

Answer 3

They are arranged in columns that have elements with similar chemical properties due to the same number of outer electrons (number of outer electrons same as group number for groups 1-7)

Question 4

How are the periods arranged?

Answer 4

They are rows of elements ordered by increasing atomic number(number of protons) leading to increasing number of outer electrons in outer shell

Question 5

What are the group names of the periodic table?

Answer 5

Alkali metals(group 1-reactive elements, reactivity increases going down group)
Transition metals(middle section)
Halogens(group 7-reactive elements, reactivity decreases going down group)
Noble gases(group 8/0-unreactive elements)

Question 6

What elements have metallic bonding?

Answer 6

Lithium(Li)
Beryllium(Be)
Sodium(Na)
Magnesium(Mg)
Aluminium(Al)
Potassium(K)
Calcium(Ca)

Question 7

What elements are covalent networks?

Answer 7

Boron
Carbon(graphite, diamond)
Silicon

Question 8

What elements are covalent molecular?

Answer 8

Phosphorus(P4)
Sulphur(S8)
Carbon(C60-fullerenes)
Diatomic-Nitrogen(N2)
Oxygen(O2)
Fluorine(F2)
Chlorine(Cl2)
Hydrogen(H2)

Question 9

Which elements are monatomic(noble gases)?

Answer 9

Helium(He)
Argon(Ar)
Neon(Ne)

Question 10

What is the trend in covalent radius across a period?

Answer 10

The covalent radius decreases

Question 11

Why does the covalent radius decrease across a period?

Answer 11

Because of increasing nuclear charge. As you go across a period a proton is added to the nucleus whilst the electrons are added into the same shell. The increased nuclear charge within the nucleus pulls the outer electron shell in tighter towards the nucleus as the electrons are more strongly attracted to the nucleus so the atoms decrease in atomic size

Question 12

What is the trend in covalent radius down a group?

Answer 12

The covalent radius increases

Question 13

Why does the covalent radius increase down a group?

Answer 13

Because as you go down a group a new electron shell is being occupied for each period so the atom is bigger (increased number of occupied shells).The inner filled electron shells also shield/screen the outer electrons from the increased nuclear charge

Question 14

What is the first ionisation energy?

Answer 14

The first ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state

Question 15

Is the first ionisation energy endothermic or exothermic?

Answer 15

It’s endothermic as energy is absorbed

Question 16

What is the second and subsequent ionisation energies?

Answer 16

The energies required to remove further moles of electrons

Question 17

How do you calculate the total energy required when there’s an equation with the removal of more than one mol of electrons?

Answer 17

The ionisation energies need to be added together to calculate the total energy required

Question 18

Why is the second ionisation energy always bigger than the first?

Answer 18

As the ions already positively charged after the first ionisation energy so the remaining electrons are held more tightly by the nucleus due to an increase in nuclear charge, so more energy is required to overcome the stronger attraction

Question 19

What is the trend in ionisation energy across a period?

Answer 19

The ionisation energy generally increases

Question 20

Why does the ionisation energy increase across a period?

Answer 20

As the outer electrons are closer to the nucleus because of the increased nuclear charge pulling them in tighter. This makes the outer electrons more difficult to remove increasing the ionisation energy. Also, the electrons are added to the same shell across a period

Question 21

What is the trend in ionisation energy down a group?

Answer 21

The ionisation energy decreases

Question 22

Why does the ionisation energy decrease down a group?

Answer 22

Because going down a group atoms get bigger as there’s an additional shell of electrons each time
The outer electrons are then further form the nucleus
Theres also an increased screening/shielding effect from the inner electron shells, so the outer electrons are shielded from the nucleus
This makes the outer electrons easier to remove so the ionisation energy decreases

Question 23

What is the electronegativity?

Answer 23

It’s the measure of the attraction an atom involved in a bond has for the electrons of that bond

Question 24

What does a high electronegativity value mean?

Answer 24

The attraction of the atom for the shared electrons in a covalent bond is strong

Question 25

What’s the trend in electronegativity value across a period?

Answer 25

The electronegativity value increases

Question 26

Why does the electronegativity value increase across a period?

Answer 26

The atoms decrease in size across a period so the bonding electrons are closer to the positive nucleus and are more strongly attracted to the nucleus(increased nuclear charge)

Question 27

What’s the trend in electronegativity value down a group?

Answer 27

The electronegativity value decreases

Question 28

Why does the electronegativity value decrease down a group?

Answer 28

The atoms increase in size as the number of filled electrons shells increases so the bonding electrons in the outer shell are further away from the nucleus and less strongly attracted The bonding electrons are screened from the full effect of the nucleus by the inner shell electrons

Question 29

What is Metallic bonding?

Answer 29

The electrostatic attraction between positive metal ions and a sea of delocalised (negative) electrons

Question 30

What kind of melting points do metals have?

Answer 30

Typically high melting points which reflects energy required to overcome the strong metallic bonds

Question 31

What type of intra and inter molecular forces are in covalent molecular elements?

Answer 31

Intra-covalent bonds Inter-London dispersion forces(very weak)

Question 32

What are the typical melting and boiling points of covalent molecular elements?

Answer 32

Most elements have relatively low melting and boiling points since only the weak LDFs need to be broken to melt and boil them. However molecules like sulfur, phosphorus and carbon fullerenes have more electrons and therefore have stronger LDFs resulting in higher melting and boiling points.

Question 33

What are the two structures of covalent bonding?

Answer 33

Discrete covalent molecular Covalent network(lattice)

Question 34

What are the melting and boiling points of covalent networks?

Answer 34

Very high as strong covalent bonds must be broken for the solid to melt

Question 35

What carbon covalent network can conduct electricity?

Answer 35

-Graphite can as the carbon atoms only form three covalent bonds so the fourth outer electron becomes delocalised between the layers(layers held together by weak LDFs, allowing layers to move easily) -Diamond doesn’t conduct electricity as a ll four outer electrons form covalent bonds

Question 36

What melting points do monatomic elements have?

Answer 36

Low as they’re held together by weak London dispersion forces, as you go down the group the melting point increases due to LDFs increasing as there is an increase in the number of electrons

Question 37

Explain why sodium is a solid at room temp whereas chlorine is a gas at room temp? HINT- same Q as why does sodium have a higher MP than chlorine?

Answer 37

To melt sodium, strong metallic bonds must be brine . Sodium is solid at room temp as there’s not enough energy at room temp to break strong metallic bonds Chlorine molecules are held together by weak LDFs, so at room temp enough energy is supplied to overcome the LDFs allowing chlorine molecules to separate into a gas

Question 38

Explain why the boiling point of fluorine is higher than the boiling point of helium? HINT- compare the LDFs

Answer 38

In both fluorine and helium the force holding the fluorine molecules and helium atoms together is LDFs The fact that fluorine’s boiling point is higher suggests that the LDFs between the fluorine molecules are stronger than the LDFs between the helium atoms. So more energy is required to overcome the LDFs of the fluorine molecules LDF strength increases as the number of electrons increases therefore as there’s more electrons in fluorine molecules compared to helium atoms the LDFs of fluorine molecules are stronger

Question 39

How do you explain trends in reactivity?

Answer 39

Number of outer electrons

Question 40

How do you explain trends in MP and BP?

Answer 40

Structure and bonding of elements

Question 41

Have negative ions gained or lost electrons Have positive ions gained or lost electrons

Answer 41

Negative-gained Positive-lost