1. SCH3U Mr. Krstovic Agenda: 1) Atomic and Ionic Trends
2) Trend in Reactivity
3) Ionization Energy
4) Electron Affinity
5) Electronegativity

2. Atomic Radius
Atomic Radius: measurement of the size of an atom in picometers
(1 pm = 10E-12 m)
We will assume that the atomic radius is one-half the distance between the nuclei of two atoms of the same element

3. Trends in Atomic Radius

4. Atomic Radius
How can we explain the trends in atomic radius across each period?
As we move across each period # of p+ increases.
Thus, the nuclear charge increases.
So the strength of the attraction between the nucleus and the valence electrons is stronger.
This nuclear attraction pulls the electrons closer to the nucleus making the atomic radius smaller.

5. Atomic Radius Why does the atomic radius increase down each group?
The number of energy levels increases from top to bottom, therefore the atomic radius increases.
The inner electrons shield the outer electrons from the pull of the nucleus.
With each additional energy level, the attractive force of the nucleus is reduced by shielding effect, so the outer electrons are not as strongly attracted by the nucleus. This results in a larger radius.

6. STOP AND THINK
Which of the following atoms is larger: aluminum or rubidium? Why?

7. Trends in Chemical Reactivity for Metals
As we move DOWN a group, the reactivity of metals INCREASES.
As we moved ACROSS a period, the reactivity of metals DECREASES.
Why do we observe this trend?
We have to look the trends in atomic radius and how close or far away an outer most electron is from the nucleus…

8. Trends in Chemical Reactivity for Metals
…so, metals give up their valence electrons during a chemical reaction.
The further away that electron is from the nucleus, the easier it is for an atom to give it up in a chemical reaction.
Thus, atoms with greater radius are more reactive than those with a smaller radius.
Francium (Fr) is most reactive metal

9. STOP AND THINK
Which of the following elements will be more reactive silver or zinc? Why?

10. IONIC RADIUS
The removal of an electron from at atom results in the formation of an ION.
Example: Remove e- from valance shell of Li, and you get Li+ cation.
Add a electron to Chlorine atom, and you get Cl- anion.
There are differences in the size of ions!

11. Ionic Radius
For example, the radius of Na atom is 186 pm, while Na+ is 102 pm. Explain the difference.
Electron is removed from a larger valence-shell orbital of Na atom.
The valence shell of Na atom is third one.
The valence shell of Na+ is the second one.
Thus, Na+ has a smaller valence shell than Na atom, and is therefore a smaller size.

12. Ionic Radii
Can you explain why O2- has a greater atomic radius than Oxygen atom?

13. Ionic Radius
O2- has a greater atomic radius than oxygen atom because as oxygen accepts 2 more electrons, the repulsive force between the two electrons is stronger than the attractive force between the electrons and the nucleus, thus resulting in a greater ionic radius.

14. STOP AND THINK
Which of the following ions is larger: strontium ion or chlorine ion? Why?

15. Trend in Ionic Radius Remember this: POSITIVE IONS ARE ALWAYS SMALLER
THAN THE NEUTRAL ATOM FROM
WHICH THEY ARE FORMED.
NEGAVTIVE IONS ARE ALWAYS
LARGER THAN THE NEUTRAL ATOMS
FROM WHICH THEY ARE FORMED.

16. Trends in the Periodic Table
Mr.Krstovic
Agenda:
Ionization Energy (IE)
Electron Affinity (EA)
Electron Negativity (EN)

17. Ionization Energy
The ionization energy is the energy needed to remove an electron from an atom or ion. 
For an element X, it is written as:
            X  +  energy    X+  +  e-
The ionization energy is a measure of how tightly the electrons are held by the atom.
cation

18. Ionization Energy (IE)
DO NOT COPY THIS!!! YOU HAVE IT IN YOUR NOTES!
Recall: across a period (ROW) (from L to R) the number of protons increase
AND outer electrons are more strongly attracted to a larger more positive nucleus.
Is it easier or harder to remove electrons? (LEFT side or RIGHT side of the PT). Explain.
The outer electrons in atoms on RIGHT SIDE (non-metals) of the PT will be held more strongly (small radius) then those elements of the LEFT SIDE (metals) of the PT

19. Ionization Energy (cont’)
Therefore as you go across a period from
L to R the IONIZATION ENERGY increases (as the size decreases)
This makes sense b/c we know that atoms of elements from the left are more apparent to lose electrons and become positive ions (i.e. metals)
IE decreases as you go down a group (as the atomic radius increases)

20. - The lower the ionization energy, the easier it is to remove the outer electron.
The higher the ionization energy, the more difficult it is to remove the outer electron.

21. 1st, 2nd, 3rd ionization energies
Atoms with more than one electron have more than one ionization energy.
energies correspond to the stepwise removal of electrons, one after another

22. Example:
two valence electrons
easy to remove first one, twice as hard to remove second one
Remove one more major jump
1st nd rd th
Be          899      1757   |  14,845     21,000

23. Movie

24. Electron Affinity EA

25. Electron Affinity X + e- -----> X- + energy
Under some circumstances, it is possible to get an atom to accept electrons.
EA is the amount of energy released when one electron is added to an atom in the gaseous state.
X  +  e-   ----->  X-   +  energy
anion

26. EA Across a period: (L to R)  INCREASE DOWN a group  DECREASE
The same factors – for EA as for IE
Across a period: (L to R)  INCREASE
a valence shell that holds its electrons tightly will also tend to hold an additional electron tightly.
DOWN a group  DECREASE
A valence shell that loses electrons easily will have little attraction for additional electrons.

28. Electronegativity EN

29. Electronegativity
Electronegativity - Electronegativity is an atom's “desire” to grab another atom's electrons.
EN is high for a nonmetallic element and low for a metallic element. WHY?
Period - EN increases from L to R across a period.
Left side elements  1 -2 valence electrons give up electrons to achieve the octet in a lower energy level
Right side elements only need a few electrons to complete the octet  so they have strong desire to grab another atom's electrons.

30. EN Group - EN decreases as you go down a group.
Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons.
Elements near the bottom (lots of electrons) loosing or acquiring an electron is not as big a deal.
This is due to the shielding affect  outer electrons not being as tightly bound to the atom.

32. summary

33. Summary
Elements with the highest ionization energy, electron affinity and electronegativity are at the top, right on the periodic table (Noble Gases excluded).
Elements with the lowest ionization energy, electron affinity and electronegativity are at the bottom, left of the periodic table.

34. pp.58 Summary Table

35. Summary Atomic Property Situation in which the atom is found
Which electron? IE
Free, isolated atom
Valence electron of the atom EA
An extra electron EN
Atom is chemically bound within a molecule
Electron is being shared with another atom in a chemical bond