1. Physical Properties

2. Syllabus statements
3.2.1 Define the terms first ionization energy, and electronegativity.
3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities and melting points for the alkali metals, and the halogens.
3.2.3  Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities for the elements across period 3.
3.2.4  Compare the relative electronegativity values of two or more elements  based on their positions on the periodic table.

3. The first ionisation energy of an atom tells us how difficult it is to remove an electron from that atom.
We only talk about first ionisation energy when we make a positive ion – this will be important to remember when we study energetics!
A definition (learn it!)
The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.
Every part of the definition is important!

4. First ionization energies are measured in kJ mol-1
You don’t need to remember the numbers – they are given in the data booklet.
You DO need to remember the trends.

5. Trends in Properties
We will look at the chemical properties later; first we need to think about some of the physical properties:

6. Atomic radius
The atomic radius is the distance from the nucleus of an atom to its outermost electron.

7. BUT
Electrons don’t stand still (and that’s without even mentioning the Heisenberg Uncertainty Principle!)
So atomic radius is sometimes defined as half the distance between neighbouring nuclei

8. For group 1 elements: Element Electron arrangement
Atomic radius (10-12 m)
Lithium Li
2,1
152
Sodium Na
2,8,1
186
Potassium K
2,8,8,1
231
Rubidium Rb
2,8,8, ,1
244
Cesium Cs
2,8, , . . ., 1
262

9. The numbers are in the data booklet – you don’t need to learn them.
The atomic radius increases down a group as the number of occupied electron shells increases.
The data booklet doesn’t give atomic radii for the noble gases.
Why?

10. All the group 1 elements lose a single electron when they are ionized.
As we would expect, positive ions are smaller than the atoms from which they were formed.
Positive ions have fewer occupied electron shells, and greater electrostatic attraction than the atoms from which they were formed.

11. Element Atomic Radius 10-12m
There is a similar trend in atomic radii going down group VII.
Fluorine has more electrons than Li.
So why does it have a smaller atomic radius?
We will answer that question in a minute!
Element
Atomic Radius 10-12m F 64 Cl 99 Br
114 I
133

12. All the group VII elements gain one electron when they form ions.
The ions have more electrons than the atoms from which they were formed.
Hence negative ions are larger than the atoms from which they were formed.

13. Cations contain _______ electrons than protons, so they are ______ than their parent atoms;
Anions contain _______ electrons than protons, so they are ______ than their parent atoms.

14. Cations contain _fewer_ electrons than protons, so they are smaller than their parent atoms;
Anions contain __more_ electrons than protons, so they are bigger than their parent atoms.

15. Now consider period 3
This is the row which contains:
Na Mg Al Si P S Cl Ar
How does atomic radius change between Na and Cl?
It gets smaller, even though Cl has more electrons.
Why?

16. All the elements in period 3 contain electrons in the 3rd shell.
I.e. they all have the same outer shell.
But, as we go across period 3 each element has one more proton than the previous element.
This extra positive charge pulls the outer shell of electrons slightly.
As we go across a period, atomic radius decreases because of increased electrostatic attraction between the nucleus and the outermost electrons.

18. The situation is a little harder for ionic radii.
We need to consider positive ions (cations) and negative ions (anions) separately.
All the ions across a period contain the same number of electrons
They are “isoelectronic” with a noble gas

19. For cations:
As we go across a period, we add protons.
The increased nuclear charge pulls the outer electrons in closer.
Ionic size decreases across a period.

20. When we start forming negative ions (anions)
There is suddenly a large increase in size.
This is because electrons have been added to the outer shell, resulting in increased repulsion with the nucleus.
BUT as we continue across the group we add further protons
(ie the number of protons gets closer to the number of electrons)
Hence the anions get smaller as we add more protons.

22. 1st Ionisation Energy and Electronegativity
Similar trends exist (for similar reasons!) in 1st IE and electronegativity.
As we go down a group, the valence electrons are attracted less strongly to the nucleus
So the 1st IE decreases as we go down a group
Question: What do you think happens to 1st IE across a period? Be prepared to justify your answer!

23. 1st IE increases across a period as the increased nuclear charge makes it harder to remove electrons.
2 small asides:
1st IE is always endothermic (we always have to supply energy)
Sometimes we talk about “effective nuclear charge”. This is the ACTUAL nuclear charge minus the effect of electron-electron repulsion.

24. Notice that 1st IE doesn’t increase smoothly across a period.
HL students need to know much more about this!

26. Just remind me again what 1st ionisation Energy was . . . ?
The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.
I thought I told you to remember that definition!

27. Let’s try another definition:
Electronegativity is a measure of the tendency of an atom in a molecule to attract a pair of shared electrons in a covalent bond towards itself.
A kind of slurpability!

28. Electronegativity is a made up concept!
Aaargh – ToK
It doesn’t have a natural scale, so we just made one up.
The scale the IB use gives each element a value between 0 and 4.
Electronegativity increases as the size of an atom decreases.

29. So just like 1st IE it increases across a period and decreases down a group.
Some attempt has been made to relate this to effective nuclear charge, but the relationship doesn’t always work (for example oxygen has a higher electronegativity than chlorine, but has a lower effective nuclear charge)

30. Question: Derive from first principles a value for the electronegativity of Xenon. Give your answer to 1 decimal place.
Duh!
What’s the definition of electronegativity?

31. Electronegativity is a measure of the tendency of an atom in a molecule to attract a pair of shared electrons in a covalent bond towards itself.
Xenon is a noble gas – it doesn’t form any covalent bonds!
So the answer is 0.0

32. In general the highest values are top left; the lowest values are bottom right.

35. Melting Points That just leaves melting points to discuss.
The melting point of a substance is determined by the size of the forces between particles.
What type of bonding do we expect in metals?
Metallic bonding!
The metal atoms form cations which are held together by a sea of valence electrons

37. As we go down group 1 the distance between the cation and the electrons increases,
So the force holding the particles together decreases,
So melting point decreases down a group of metals

39. As we go from group 1 to group 2
The number of valence electrons increases,
So the forces holding particles together increases
So melting point increases from group 1 to group 2
The same happens going from group 2 to group 3

41. Group 4 elements have even higher melting points.
This is nothing to do with metallic bonding!
They form giant molecular (aka giant covalent) structures.
The particles are VERY strongly held by covalent bonds.

43. a) Diamond
b) Graphite
c) Lonsdaleite (a kind of synthetic diamond)
d) C60 (Buckminsterfullerene)
e) C540 (a Fullerene with 540 carbon atoms)
f) C70 (a Fullerene with 70 carbon atoms)
g) Amorphous carbon
h) Single-walled carbon nanotube

45. Continuing across a period, the meting points get much lower.
Now we have simple molecules with low forces between particles
These are Van der Waal’s forces
(if that’s too hard to remember, you can call them London forces!)
Group 8 have the lowest melting points as the elements exist as single molecules
(Compare them with: N2 O2 Cl2 P4 S8 etc)

46. Melting points increase going down a group of non metals because
Van der Waal’s forces are greater for bigger particles

47. Summary of Trends

48. Syllabus statements
3.2.1 Define the terms first ionization energy, and electronegativity.
3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities and melting points for the alkali metals, and the halogens.
3.2.3  Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativities for the elements across period 3.
3.2.4  Compare the relative electronegativity values of two or more elements  based on their positions on the periodic table.