1. Periodic Trends and Bonding
Higher Supported Study - Week 2
Periodic Trends and Bonding

2. Key Areas – Periodic Trends
The trends in covalent radius across periods and down groups.
The trends in ionisation energies across periods and down groups.
The trends in electronegativity across periods and down groups.

3. Why is there no covalent radii information for noble gases?
Covalent Radius
The covalent radius is half the distance between the two nuclei in a covalent bond.
There is no information for the noble gases because they do not form covalent bonds.
Why is there no covalent radii information for noble gases?

4. Covalent Radius Trends
Across a period
Covalent radius decreases due to
INCREASING NUCLEAR CHARGE
pulls electrons in closer to the nucleus
Down a group
Covalent Radius increases due to
increasing number of electron shells
Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening to the size of the atom and why.
Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

5. Ionisation Energies
Definition Energy required to remove one mole of electrons from one mole of gaseous atoms
This is a definition which you should be able to recall.
The equation for first ionisation energies and second ionisation energies are given on page 11 of data book.
The elements on this page are arranged in periods.

6. Ionisation Energy Trends
Across a period
Ionisation Energy increases due to
INCREASING NUCLEAR CHARGE
pulls electrons in closer to the nucleus, meaning more energy is required to remove these electrons
Down a group
Ionisation Energy decreases due to
the electron being further away from the
nucleus due to additional energy levels and
the inner electrons screening the outer
electrons from full effect of nuclear charge.
Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why.
Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

7. Ionisation Energy Trends
Always write out the electron arrangements of atoms/Ions.
Compare nuclear charge and no. of electron shells.
Across a period - INCREASES
Increasing nuclear charge
Down a group – DECREASES
additional energy levels and screening
Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why.
Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

8. Ai) K(g)  K+(g) + e- (this should be easiest mark to get in exam as general equation given in data book, STATE SYMBOLS must be included.
Aii) Write out Nuclear charge and electron arrangements for the two things which are being compared.
K (19) 2, 8, 8, 1
Cl (17) 2, 8, 7
K outer electron is further away from nucleus, due to greater number of electron shells, also inner electrons shield the outer electrons from full effect of positive nuclear charge.

9. B C

10. Other Styles of Questions
Explain why potassium has a lower first ionisation energy than lithium
Explain why the second ionisation energy of sodium is much higher than the first ionisation energy of sodium.
Explain why the P3- ion is larger than the Al3+ ion
Explain why the Ca2+ ion is smaller than the P3- ion.
K – 2, 8, 8, 1
Li- 2, 1
Potassium has more electron shells, the electron to be removed is further away from the nucleus and the inner electrons shield the outer electrons from the full effect of the nuclear charge. (Notice I link response to definition of first Ionisation energy)
Na – 2,8,1
The second electron is being removed from a shell closer to the nucleus. This electron has a much greater nuclear charge to overcome, as is closer to nucleus.
P3- – 2, 8, 8.
Al3+ - 2, 8.
Al only has 2 electron shells, while P has 3 making the P ion larger.
Ca2+ - 2, 8, 8
P3- – 2, 8, 8
Have same number of electron shells, so change in size must be due to difference in nuclear charge. Ca has a greater nuclear charge, drawing electrons in closer making the ion smaller.

11. for five elements A, B. C, D and E, in kJ mol-1.
The following table shows the approximate first ionisation energies (IE)
for five elements A, B. C, D and E, in kJ mol-1.
Which of these elements;
(a) could be in group 2 of the periodic table?
(b) could be in the same group of the periodic table?
(c) would require the least amount of energy to convert one mole of
gaseous atoms into ions carrying a three positive charge.
Explain your answers.
3. Write equations corresponding to;
(a) The second ionisation energy of magnesium.
(b) The third ionisation energy of aluminium.
element
first IE
second IE
third IE
fourth IE A
520
7300
11500 B
2100
3900
6100
9400 C
580
1800
2800 D
740
1450
7700
10500 E
420
3050
4500
5900
A – Group 1
B – Group 3
C- Group 3
D – Group 2
E – Group 1
(c) X(g)  X3+(g)
Need to add together (1st IE + 2nd IE + 3rd IE) energy required to remove 3 electrons.
3a) Mg+(g)  Mg2+ + e-
b) Al2+ (g)  Al3+ + e-

12. I.E. Questions State the energy required for the following changes.
a) Al(g)  Al+(g) + e-
b) Al(g)  Al2+(g) + 2e-
c) Al+(g)  Al3+(g) + 2e-
First IE from data book
First IE and Second IE added together, as 2 electrons being removed.
One electron already lost before started, so Second IE and third IE added together.

13. Electronegativity
Atoms of different elements have different attractions for bonding electrons.
Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond.
Another key definition to learn – electronegativity

14. Electronegativity Trends
Across a period
Electronegativity increases due to
INCREASING NUCLEAR CHARGE
causes atoms to attract bonding electrons
more strongly
Down a group
Electronegativity decreases due to
increasing number of electron shells, the inner
electrons shield the outer electrons from the
nuclear charge, meaning the electrons are less
strongly attracted to the atom
Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why.
Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

15. Ionisation Energy and ELECTRONEGATIVITY Trends
Always write out the electron arrangements of atoms/Ions.
Compare nuclear charge and no. of electron shells.
Across a period - INCREASES
Increasing nuclear charge
Down a group – DECREASES
additional energy levels and screening
Tip draw out a few atoms, with nuclear charge in centre, and electron arrangements. This will help you identify what is happening and why.
Remember the covalent radius data is on page 7 of your data books, you don’t actually have to remember the trends, just the reasons why.

16. Key Areas – Bonding Covalent bonding. Polar covalent bonding.
The bonding continuum.
VDWFs
Properties linked to Bonding

17. Recap From Nat 5 Bonding in Elements Bonding in Compounds Metallic
Covalent molecular
Covalent network
Monatomic (group 8)
Bonding in Compounds
Ionic
Covalent

18. Covalent Bonding
This is listed in Higher Arrangement documents, so could be asked at Nat 5.

19. 3 types of Bonding Covered at Nat 5.

20. Intramolecular Covalent molecules
Non polar – no difference in electronegativity
- electrons shared equally
Polar – DIFFERNCE in electronegativity
- electrons NOT shared equally
LARGER THE DIFFERENCE IN ELECTRONEGATIVTY THE MORE POLAR THE BOND WILL BE
The key here is electronegativities.

21. Bonding Continuum
Depending on difference in electronegativities, bonds are classed as Pure covalent, polar covalent or ionic. This scale is called the Bonding continuum.

22. Bonding Continuum
If the electronegativity difference is large then the movement of bonding electrons from one atom to another results in formation of ions.
Compounds formed between metals and non-metals are often, but not always ionic.
If the electronegativity difference is less than 0.4 we class the bonds as pure covalent.
If the electronegativity dofference is between 0.4 and 1.7 it is classed as polar covalent
If the electronegativity difference is greater than 1.7 it is ionic.
The greater the ionic character the greater the difference in electronegativity (common exam question)
Most covalent character the smaller the difference in electronegativity.

23. Intermolecular Bonding
All molecular elements and compounds and monatomic elements condense and freeze at sufficiently low temperatures.
For this to occur, some attractive forces must exist between the molecules or discrete atoms.
Van Der Waals Forces
London Dispersion Forces
Permanent dipole - permanent dipole
Hydrogen Bonding
HINT – Always useful to include diagrams when trying to explain these types of interactions with ∂+ and ∂- signs when appropriate
First two sentences lifted straight from Higher arrangement documents.
Remember these forces only found between molecules and monatomic.
Ionic/metallic/network are a continuous structure of atoms/ions. There is no individual atoms/molecules so forces of attraction can exist “between”.

25. Properties of Molecules
Viscosity
Melting/Boiling Point
Can all be explained by comparing the types of VDWF’s present.
If comparing two molecules state the
Type of VDWF’s in each molecules
The strength of the VDWF
A detailed explanation of how the VDWF arises.

26. ICE
H- Bonding can be used to explain why ice is less dense than water.
Ice contains an OPEN-LATTICE ARRNAGEMENET OF WATER MOELCULES

27. Solubility
Ionic compounds and polar molecules tend to be soluble in polar solvents such as water and are insoluble in nonpolar solvents.
Non-polar molecules tend to be soluble in non-polar solvents and insoluble in polar solvents.
“Like dissolves like”

29. Diagrams can help explain solubility

30. 2012
10. Atoms of nitrogen and element X form a bond in which the electrons are shared equally.
Element X could be
A carbon
B oxygen
C chlorine
D phosphorus.
Answer C

31. 11. Which line in the table represents the solid in
2012
11. Which line in the table represents the solid in
which only van der Waals’ forces are overcome
when the substance melts?
Answer - D

32. 14. In which of the following solvents is lithium
2012
14. In which of the following solvents is lithium
chloride most likely to dissolve?
A Hexane
B Benzene
C Methanol
D Tetrachloromethane
Answer – C
Note the use of the language “most likely”

33. Answer - C