1. Trends in Periodic Table

3. Atomic Radius
of an element is defined as half the distance between the nuclei of the two atoms of the same element that are joined together by a single covalent bond
Atomic Radius
Bond Length

4. - What happens to the size of the atomic radius as you go down a group
- What happens to the size of the atomic radius as you go down a group? - What happens to the size of the atomic radius as you go across a period?

5. Why does the size of the atomic radius differ?
New shells and sublevels increase the atomic radius as you go down a group
Screening effect of inner electrons cause the outer electrons to experience less nuclear charge so are not as tightly held, increasing the atomic radius
Increasing nuclear charge and no increase in the screening effect from L-R across a period causes a decease in the atomic radius

6. Trends in Atomic and Ionic Size
Metals
Nonmetals
Group 1 Al
143 50 e
Group 13
Group 17 e e
152
186
227 Li Na K 60
Li+ F-
136 F Cl Br 64 99
114 e e 95
Na+
Cl-
181
Al3+ e e
133 K+
Br-
195
Cations are smaller than parent atoms
Anions are larger than parent atoms

7. Why cant you measure the size of a single atom
Why cant you measure the size of a single atom? [2] cant determine where outermost electron is going to be / heisenberg’s Uncertainty principle
What is atomic radius? [4] half distance / between two adjacent nuclei / of same element / joined by single covalent bond
What happens to atomic radius as one goes across period? Why? [3] Gets smaller / increased nuclear charge / same screening effect
What happens to atomic radius as you go down a group? Why? [4] Gets bigger / new main energy level (shell) / increased nuclear charge / cancelled out by increased screening effect
For which group are there no atomic radius values? Explain? [3] noble gases / don’t bond / cant predict position of outermost electron

8. Ionisation Energy
Minimum energy required to completely remove the most loosely bound electron from a mole of gaseous atoms in their ground state
Unit: kJ mol-1
The general symbol is D Hi, and for a first ionisation energy it is D Hi1.
The process may be shown by the example of calcium as:
Ca (g) = Ca+ + e-

9. Page 80 - Log Tables Ionisation Energy Increases L-R
Ionisation Energy decreases going down a group
Increasing nuclear charge/ no increase in screening
Decreasing atomic radius
Increasing atomic radius
Larger screening effect

10. Across a Period Ionisation energy Increases
(i) increasing nuclear charge.
Number of protons increases so attraction increases
electrons in same main energy level so no increase in screening
(ii) decreasing atomic radius
atomic radius decreases so electrons are closer to positive nucleus so held tighter.

11. Down a Group Decreases [gets less]
(i) Increasing atomic radius means electron further from attractive force of the nucleus
(ii) Screening effect of inner electrons. Positive charge of nucleus is increasing but inner e- shells shield the outer electron from this increased charge.

12. Exceptions

13. Explanation Beryllium higher than Boron Nitrogen higher than Oxygen
Removing electron from a full s sublevel
Having a full s sublevel gives Beryllium extra stability making it harder to remove an electron
Nitrogen higher than Oxygen
removing electron from a half filled p sub-level
Having a half full p sublevel gives Nitrogen extra stability making it harder to remove an electron from it than Oxygen

14. Graph of ionisation energies
Why has Aluminium has a lower first ionisation energy than Magnesium?

15. Mg 1s2, 2s2, 2p6, 3s2
Al 1s2, 2s2, 2p6, 3s2 3px1
Aluminium ionisation energy lower than Mg
Removing electron from Mg is from a full s sublevel which requires more energy than removing electron from Al which is from a partially filled p sublevel
Having a full s sublevel gives Mg extra stability making it harder to remove an electron

16. Exceptions to general trend
Look at the general trend.
Argon has the highest value.
Why?
Full sublevel [outer shell]
Are any elements higher than general trend?
Yes
Mg and P
Why?

17. Explanation Magnesium higher Removing electron from a full s sublevel
Phosphorous higher
removing electron from a half filled p sub-level
Having a half full p sublevel gives Phosphorus extra stability making it harder to remove an electron from it

18. Define first ionisation energy
Define first ionisation energy. [5] The energy required / to completely remove / the most loosely bound electron / from a [mole of] neutral gaseous atom [s], / in its [their] ground state
Write an equation for the first ionisation of potassium. [2] K = K+ + e-
What happens to first ionisation energy across a period? Explain [3] Increases / increasing nuclear charge / decreasing atomic radius
What happens to first ionisation energy down a group? Explain [4] Decreases / Increasing atomic radius / increased screening effect / of inner electrons.
Be and N have higher values than the general trend; explain why [4] Be filled / S sublevel / N ½ filled / P sublevel

19. Second and successive Ionisation Energy

20. Second ionisation energy
The energy needed to remove a second electron from each ion in a mole of ions is the second ionisation energy.
e.g. for calcium:
Ca+(g) = Ca2+(g) + e DHi2 = kJ mol-1
The second ionisation energy is much larger than the first.
What are the reasons for this?

21. Reasons
The ionisation energies increase because as each electron is removed from an atom, the remaining ion becomes more positively charged.
Removing the next electron away from an increasing positive charge is more difficult and the ionisation energy is even larger.

22. Further evidence for the existence of Energy Levels
Take a Na atom and remove the electrons one by one.
Plot a graph of the successive ionisation energies.
Use log10 because values so big.

23. Graph of logarithm (log 10) of ionisation energy of sodium against the number of electrons removed.
There are one or more particularly large rises within the set of ionisation energies of each element.
Why??

24. Why?
Why are there large increases between the first and second ionisation energies and again between the ninth and tenth ionisation energies?
Look at the s,p,d configurations to find out

25. Trends within groups
Alkali Metals: What do you think happens to the reactivity as you go down the group?
Reactions with oxygen: 2K + ½O2 = K2O
Reactions with water: Na +H2O = NaOH +½H2

26. Halogens
What happens to the reactivity of halogens as you move down the group?
Reactivity Decreases

27. Trends in Electronegativity
Is the relative attraction of an atom in a molecule has for the shared pairs of electrons in a covalent bond
Page 81 - Log Tables

28. Electronegativity increases
Increasing nuclear charge
Decreasing atomic radius
Increasing atomic radius
Larger screening effect
Electronegativity Decreases

29. Is the relative attraction that an atom in a molecule has for the shared pair of electrons in a single covalent bond between two atoms of the same element.
In a bond between two identical atoms the pair of electrons are shared equally
chemists have found that in many bonds the pair of electrons are attracted to one of the atoms more than to the other.
SCC Science Dept

30. Hydrogen and Chlorine
Electrons attracted to chlorine more than to hydrogen [ bigger , but more +ve nucleus]
therefore the electrons spend more time near the chlorine than near the hydrogen
this gives the chlorine a slightly negative charge [δ- delta minus]
it gives the hydrogen a slightly positive charge [δ+ delta plus]
SCC Science Dept

31. H Cl
SCC Science Dept

32. Linus Pauling measured the electronegativity of each element and put them in a table
Noble gases are not in this table because they do not form bonds H Cl
+
 –
SCC Science Dept

33. Differences and Bond Type
Difference 0 to 0.45
[Pure] Covalent Bond
Difference > 0.45 but < 1.7
Polar Covalent Bond
Difference is = or > 1.7 then
Ionic Bond
SCC Science Dept

34. Trends in Electronegativity
Across a period
Goes up
Bigger nuclear charge –
Smaller atomic radius distance from nucleus
Actually closer as diameter of atom gets smaller –
SCC Science Dept

35. Trends in electronegativity
Down a Group
Goes down
Bigger nuclear charge
But! But! But! But!
Increased Atomic Radius - electron is much further from nucleus
Increased shielding by more inner electron shells
SCC Science Dept

36. Who invented the term? [1] Linus Pauling
Define electronegativity [3] measure of the attraction of an atom / for an electron / in a single covalent bond / between atoms of the same element
Who invented the term? [1] Linus Pauling
What happens to it as you go across a period. Explain.[3] Goes up / bigger nuclear charge / same shell or decreasing atomic radius
What happens to it as you go down group? Explain. [3] Goes down / further from nucleus / more shielding
Give three classes of bond and electronegativity associated with each [6] / pure covalent / 0.4 to 1.7 / polar covalent / over 1.7 / ionic
How can you show that a molecule is polar. [4] Rub pen / pour stream past it / polar bends / non-polar does not bend.
What group of elements is not present in the table of electronegativities. Why? [2] Noble Gases / don’t bond
Total = 22
SCC Science Dept