1. Warm up for Today!
How many valence electrons are in Titanium (Ti)? Hint – draw the electron configuration.
How many unpaired electrons are in Tellurium (Te)? Hint – draw the orbital diagram.

2. Lecture 6B
Periodic Trends

3. Development of the Periodic Table
Dmitri Mendeleev designed the first periodic table in 1869 by grouping elements with similar chemical & physical properties in rows according to Atomic mass
Henry Moseley rearranged the table in 1913 according to atomic number

4. The Modern Periodic Table
Organized by groups (columns) and periods (rows)

5. Classifying Elements by Electron Configurations
Representative Elements- Outermost s and p are filling
Alkali metals – Outermost s has 1 e-
Alkaline earth metals – Outermost s is full
Halogens – Outermost s is full, outermost p has 5 e-
Noble Gases- Outermost s and p are totally full
Transition Metals -Outermost s is full and d is filling or full
Inner Transition Metals- Outermost s is full and f is filling or full

6. Hence the name periodic table.
Periodic Trends
As you move across the periodic table, you keep periodically running across the same properties...
The location of an element on the Periodic Table gives a lot of information about key properties of that element compared to the other elements
Hence the name periodic table.

7. Trend #1 – Atomic Size
An atom’s size is determined by its atomic radius
Atomic radius is defined as half the distance between the nuclei of two like atoms

8. Atomic Size Trend
Going down a group, atoms get bigger because the number of energy levels increases.
Going across a period, atoms get smaller. No new energy levels are added, but more protons and electrons are added. The increase in protons and electrons result in a greater pull towards the nucleus – like a magnet!

10. Trend #2 – Atomic Size of IONS
Cations (positive ions) are always smaller than their parent atom because they lose electrons and an energy level. A strong attraction forms between the remaining electrons and the nucleus.
Anions (negative ions) are always larger than their parent atom. Gaining electrons causes less attraction to the nucleus – the new electrons aren’t held as tightly and are now freer to move around!

11. Comparison of atomic vs. ionic size

12. Answer these . . . Which is larger: S or S2-. Why?
S2- is larger – Gaining electrons causes less of an attraction to the nucleus – the new electrons are more free to move around.
Which is smaller: Fe or Fe4+. Why?
Fe4+ is smaller – Losing electrons causes more of an attraction to the nucleus – the remaining electrons are pulled in tighter.
Which is smaller: Na1+ or Al3+? Why
They both lose an energy level BUT Al3+ loses MORE e-, nucleus pulls the remaining electrons closer due to the larger nuclear charge.
Which is smaller: Be2+ or Na1+. Why?
They both lose an energy level BUT Be2+ goes down to a smaller energy level.
A neutron walks into a ‘restaurant’ and says, "Hey bartender give me a drink." The bartender gives him one and says, “No charge for you"
An atom walks into a ‘restaurant’ and proclaims, "Hey! Somebody just stole one of my electrons!" The waiter says, "Are you sure?“ The atom replies, "Yes - I'm positive!"

13. Trend #3 – Electronegativity
ELECTRONEGATIVITY: the ability for an atom to attract an electron
Going down a group, electronegativity decreases because the added energy levels ‘shield’ the power of the nucleus to attract electrons.
Going across a period, electronegativity increases because nuclear charge increases, but the electron distance does not! Non-metal atoms want to gain electrons to completely fill their outer shell.

14. Which element would be the most electronegative? Why?
Fluorine – The outermost electrons reside on a low energy level. Therefore, there is a very strong attraction between the electrons and the nucleus. To fill its outer shell, Fluorine needs only one electron!

15. How does ‘Shielding’ work?
SHIELDING EFFECT: The process of the inner electrons shielding/repelling the outer electrons. The inner electrons ‘shield’ the outer electrons from the pull of the nucleus. Therefore the outer electrons are not held as tightly as the inner electrons.
Shielding increases as you go down the periodic table because more energy levels are added.
Shielding remains constant as you go across the periodic table because all electrons added enter the same energy level.

16. Why aren’t noble gases assigned electronegativity numbers?
They are happy the way they are – thank you very much – they don’t need any more electrons 

17. Trend #4 – Ionization Energy
IONIZATION ENERGY - The energy needed to REMOVE an electron from the outer level.
Going down a group, ionization energy decreases because more energy levels are added. This ‘shielding causes the electrons to be less attracted to the nucleus – little energy is required to remove them.
Going across a period, ionization energy increases because the electrons move closer to the nucleus. Therefore more energy is needed to remove them!

18. First vs. Second
First ionization energy is the energy required to remove the first electron
Second ionization energy is the energy required to remove the second electron (and so forth)
First ionization energy is always smaller (compared to the second or third ionization energy) because each successive electron removed is closer to the nucleus and more strongly attracted to it.

19. Graphing First Ionization Energy vs. Atomic Number

20. Trend #5: Metallic Properties First off, what’s meant by ‘metallic’?
Differences between metals and nonmetals
Metals
Nonmetals
Solid at room temp
Most are gases at room temp, some are solid or liquid.
Malleable (Can be hammered into sheets) and ductile (Can be drawn into wires)
Not malleable nor ductile – mostly brittle or powdery if solid
Conduct electricity and heat well
Conduct electricity and heat poorly
Lose electrons to become cations
Gain electrons to become anions

21. Metallic properties increase as you go down a group, and decrease as you go across a period.X
Fr is the most metallic!