1. Periodic Properties of the Elements
1869: Dmitri Mendeleev and Lothar Meyer
publish identical tables. Mendeleev gets the
credit and becomes “Father of Periodic Table”.
1913: Henry Moseley develops
concept of atomic numbers.
Modern periodic table arranged in order of increasing atomic number.

3. Mendeleev’s Original Periodic Table, 1869

4. Coulomb’s Law of Attraction
F = kQ1Q2 d2
The strength of interaction between two electrical
charges depends on the magnitude of the charges and
the distance between the two.
So…the forces of attraction between the nucleus and the electron depends of the net nuclear charge acting of the electron and the average distance between the nucleus and the electron.

5. Effective Nuclear Charge (Zeff)
In an atom, each electron is simultaneously attracted to the nucleus (+ charge) and repelled by other electrons (- charge).
We can estimate the energy of an individual electron in the average electric field created by the nucleus and the electron density of the other electrons.
We look at the effective nuclear charge (Zeff) located at the nucleus.

6. How Do We Determine Effective Nuclear Charge (Zeff)?
Zeff = Z - S
Zeff = charge acting on electron by nucleus
Z = number of protons in nucleus
S = Average number of electrons between nucleus and specific electron. S can be a non-integer.

7. What is Electron Shielding?
In a multielectron system an e- in
any orbital will partially shield an
e- in any other orbital.
Electron density of core e-’s shield or screen outer e-’s from the full charge of the nucleus.
Electrons in the same shell do little
shielding but do repulse each other.

8. What Determines Effective Nuclear Charge (Zeff)?
The Zeff experienced by outer electrons is determined primarily by the difference between the charge on the nucleus and the charge of the core electrons.
The force of attraction between the electron and the nucleus increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.

9. Estimating Zeff Zeff = Z - S Z = number of protons = 12+
We can roughly estimate Zeff
for an outer electron in a magnesium atom.
Zeff = Z - S
Z = number of protons = 12+
S = avg. # of core electrons = 10
Zeff = = 2+ charge

11. How Rough is our Estimate of Magnesium’s Zeff?
The actual effective nuclear charge for an outer magnesium electron is about 3.3+ compared with our calculated 2+ charge.
Our equation underestimates Zeff because it ignores the times when the outer electron may be inside the core, closer to the nucleus. This penetration causes the Zeff to be higher.

12. Periodic Table Trend 1: Effective Nuclear Charge
Effective nuclear charge steadily increases as we go from left to right across a period.
WHY?
The charge increases because the number of protons increases but the number of core electrons (shielding) stays the same.
Effective nuclear charge increases slightly as we go down a family.
WHY?
Larger cores are less able to shield the outer electrons.
Remember: Electron shielding is constant across a period because the core electrons remain constant!
Note: This trend down a family is much less important that the trend across a period.

13. Periodic Trend 2: Atomic Radius
• Atoms do not have sharply defined boundaries (think electron
cloud) but an atomic radius is known for most elements.
Nonbonding atomic radius (van der
Waals radius): the radius of an atom as
defined by the closest distance separating
its nucleus from the nucleus of another
atom during a collision.
0.5 d
Bonding atomic radius (covalent radius): the radius of an
atom as defined by the distance separating its nucleus from the
nucleus of another atom to which it is chemically bonded.

14. Periodic Trend 2: Atomic Radius
Atomic radius decreases
up a family (bottom to top
of periodic table).
WHY?
Radius decreases with decreasing n because the electrons spend more time closer to the nucleus, thereby decreasing atomic size.
Atomic radius decreases
across period from
left to right.
WHY?
Shielding stays constant across a period so Zeff increases. Greater Zeff draws the electrons in closer to the nucleus, thereby decreasing atomic size.

16. Periodic Trend 3: Ionic Radius
CATIONS
Radius of a cation is smaller than its parent atom.
WHY?
Outermost electrons leave an atom to form a cation.
b. Fewer electrons means less electron-electron repulsion.
ANIONS
Radius of an anion is larger than its parent atom.
WHY?
a. Electrons are added to outer shell to form an anion.
b. More electrons means more electron-electron repulsion

18. Periodic Trend 3: Ionic Radius Trend Up a Family
Cation radii decrease
up a family (bottom to top).
WHY?
Atomic radii decrease up a family, and cations are smaller than their parent ions. Larger atoms at the bottom of the table have larger ions than smaller atoms at the top of the table.
Anion radii decrease
up a family (bottom to top).
WHY?
Atomic radii decrease up a family, and anions are larger than their parent ions. Larger atoms at the bottom of the table have larger ions than smaller atoms at the top of the table.

20. Periodic Trend 3: Ionic Radius Trend Across a Period
The ionic radius trend changes across a period from left to right as we move from metals to nonmetals.
The radii of metallic ions decrease while the radii of nonmetallic ions increase across a period.
This variation between cations and anions causes a wave-like pattern on the table.

21. Isoelectronic Series of Ions
Isoelectronic ions have the same number of electrons.
Example:
O2-, F-, Na+, Mg2+, Al3+ all have 10 electrons.
Increasing nuclear charge 
O2-, F-, Na+, Mg2+, Al3+
Decreasing ionic radius 
In an isoelectronic series of ions, the ion with the most negative charge has the largest radius.

22. Ionization Energy The greater the ionization energy,
Ionization Energy is the minimum amount of energy needed to remove an electron from the ground state of an isolated gaseous atom or ion.
First Ionization Energy (I1) removes the first electron from a neutral atom:
Na(g)  Na+(g) + 1e-
Second Ionization Energy (I2) removes the second electron:
Na+(g)  Na2+(g) + 1e-
Formation of cations
The greater the ionization energy,
the more difficult it is to remove an electron.

23. Variations in Successive Ionization Energies
Ionization energy for removal of successive electrons:
I1 < I2 < I3
There is a sharp increase in
ionization energy as inner shell
electrons are removed.
Every element shows a large
increase in ionization energy
when electrons are removed
from its noble gas core.

25. Periodic Trend 4: Ionization Energy
Ionization energy generally increases as we go up a family (bottom to top).
Ionization energy generally increases as we go from left to right across a period.
WHY?
As atomic radius decreases, it is harder to remove electrons because they are more attracted to the nucleus.
WHY?
As Zeff increases and the atom gets smaller, it is harder to remove electrons because they are more attracted to the nucleus.

28. What Is Electron Affinity?
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It measures the attraction of the atom for the added electron.
Ionization energy measures how easily an atom loses an electron and electron affinity measures how easily an atom gains an electron.

29. Basic Electron Affinity Concepts
For most atoms, energy is released when an electron is added. Therefore, electron affinities are usually negative values.
KABOOM!
Example:
Cl(g) + 1e-  Cl-(g) ∆E = -349 kJ/mol
The greater the attraction between an atom and an added electron, the more negative the atom’s electron affinity will be.

30. What Does It Mean If The Electron Affinity Is Positive?
Some elements such as noble gases have positive electron affinities.
Example:
A(g) + 1e-  Ar- (g) ∆E > 0 kJ/mol
This means that the argon
anion is higher in energy than
the separated atom and
electron. Because ∆E > 0, the
argon ion is unstable and does
not form.
Ar-

31. Periodic Trend 5: Electron Affinity
Electron affinity generally becomes increasingly negative as we go from left to right across a period. Halogens have the most negative electon affinites. Noble gases all have positive electron affinities.
WHY?
Halogens are one electron shy of a very stable noble gas configuration.
Electron affinity does not change greatly as we move up a family (bottom to top).
WHY?
Increasing electron-nucleus attraction is counterbalanced by higher electron-electron repulsions.