The Periodic Table

Force of Attraction:

Q1 & Q2 are charges and d is the distance between them

According to this equation…

1. Electrons close to the nucleus will be held with a greater force than those farther away from the nucleus.

2. A higher the positive charge (effective nuclear charge) will draw electrons closer to the nucleus and hold them with a

greater force.

Valence Electrons (Outer-Shell Electrons)

Electrons that can participate in the formation of chemical bonds.

Electrons in the outermost “s” and “p” orbitals. The number of valence electrons corresponds to the

group number!

Transition & Rare Earth elements are tricky! Sometimes their highest level “d” and “f” electrons

behave like valence electrons and sometimes they behave like shielding electrons.

Shielding ElectronsCore Electrons

(Inner-Shell Electrons)

Electrons included in a noble gas “core” in the electron configuration.

They are between the nucleus and the valence electrons.

The are called shielding electrons because they shield the valence electrons from the attractive force of the nucleus.

Nuclear Charge

In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.

The nuclear charge is equal to the number of protons in the nucleus or to the atomic number (Z)

Effective Nuclear Charge

Net positive charge felt by an electron. Takes into account the core electrons which shield (or

screen or cancel out) some of the nuclear charge.

Effective Nuclear Charge

Zeff = Z − S

where Z is the atomic number and S is a screening constant, usually close to the number of inner

electrons

Example:

Na: 1s22s22p63s1 Z=11, S=10, Zeff = +1

*This is an approximation – there are better methods beyond the scope

of this course – this calculation is close enough for us!

Zeff Across a Series

Li Be B C N O F Ne

Z 3 4 5 6 7 8 9 10

S 2 2 2 2 2 2 2 2

Zeff +1 +2 +3 +4 +5 +6 +7 +8

Zeff increases across a series The electrons are held closer to the nucleus are held more tightly by the nucleus as we move across a series.

Zeff Down a GroupZ S Zeff

Li 3 2 +1

Na 11 10 +1

K 19 18 +1

Rb 37 36 +1

Cs 55 54 +1

We calculate Zeff to be fairly constant going down a group, however it actually increases slightly because the core electrons are spreadout over a larger space and therefore don’tscreen the valence electrons as well.

Size of Atoms: Across a Series Atoms decrease in size as electrons are being added to the same energy level. Zeff increases so the force on the valence electrons gets stronger and they are pulled closer to the nucleus as we move across a series.

ATOM Na Mg Al Si P S Cl Ar

SIZE 1.54 1.36 1.18 1.11 1.06 1.02 .99 .98

(radius in Å)

Size of Atoms: Down a Group Atoms increase in size as the electrons go into a higher energy

level (get farther away from the nucleus)

ATOM SIZE (radius in Å) H .32 Li 1.23

Na 1.54 K 2.03 Rb 2.16 Cs 2.35

Size of Ions

A positive ion (cation) is smaller than its atom – electrons leave from the outermost orbitals,

electron-electron repulsion decreases.

A negative ion (anion) is larger than its atom – adding electrons causes more electron-electron

repulsion causing the electrons to spread out

IONIZATION ENERGY (I.E.) Low ionization energy is good for making ions!

The First Ionization Energy, I1, is the energy required to remove the first electron from the atom

The Second Ionization Energy, I2, is the energy required to remove the 2nd electron

This continues for the successive removal of electrons.

I1 < I2 < I3, etc. because each electron removed is pulled from a more positive ion.

ACROSS A SERIES

I1increases because of the increased attraction between the nucleus and the electrons (higher Zeff).

ATOM Na Mg Al Si P S Cl Ar

I.E. 119 176 138 188 242 239 299 363

(kcal/mole)

ACROSS A SERIES

ATOM Na Mg Al Si P S Cl Ar

I.E. 119 176 138 188 242 239 299 363

(kcal/mole)

It’s not a perfect trend because of more stable filled & half filled orbitals…..

The 3p sublevel has a higher energy than 3s so it’s easier to take an electron from the 3p (Al) than from the 3s (Mg)

Sulfur’s electron is removed from an orbital with paired electrons which have some electron-electron repulsion. Phosphorus’s electron is removed from an orbital with a single electron.

DOWN A GROUP

Ionization energy decreases as the electrons go into higher energy levels which are farther away from the nucleus.

ATOM I.E. (kcal/mole)

H 314

Li 124

Na 119

K 100

Rb 96

Cs 90

Making Ions of Transition Elements

The electrons which are removed to make a positive ion are always removed from the orbitals with the highest

principal quantum number (energy level) first!

Example:

Fe = [Ar]4s23d6

Fe+2 = [Ar]3d6

Fe+3 = [Ar]3d5

ELECTRON AFFINITY (E.A.) The energy change when an electron is gained by an atom The negative on these values means energy is released (it’s

exothermic) when an electron is gained. The more negative the electron affinity, the easier an atom will gain

an electron.

ACROSS A SERIES Electron affinity increases to a maximum in Group

VII It is a minimum in Group VIII because these elements

already have stable electron structures

We see the same anomalies in this trend due to the filled & half filled sublevel stability

DOWN A GROUP

Electron affinity decreases slightly

Fluorine is an exception because it’s extra small size contributes to stronger electron-electron

repulsion.

TRENDS - arrow points in the direction of the INCREASING trend

I.E. E.A.

→ ↑ → ↑

I.E. and E.A. Summary

Small IE is good for makingPositive ions! (Metals)

High EA is good for makingNegative ions! (Non-Metals)

Metals vs. Non-metals ACROSS A SERIES, elements become LESS metallic.

DOWN A GROUP, elements become MORE metallic.

The big “stair-step” line separates metals from nonmetals.

Elements to the left of the line are metals (except H!)

Elements to the right of the line are nonmetals

Some elements right along the line are metalloids.

Metal vs. Nonmetal

Test Note:

Be careful to distinguish trends from explanations!

If a question asks you to explain why an oxygen atom is smaller than a nitrogen atom, “because it’s farther right on the periodic table” is not the answer. That is a trend…the

explanation is that there is a higher effective nuclear charge on oxygen which pulls the electrons closer to the nucleus.