Chapter 8: Electron Configuration & Chemical Periodicity
The Original Periodic Table
Mendleev arranged the 65 elements known at the time into a table and summarized their behavior in the periodic law:
when arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties.
The Modern Periodic Table
Arranges the table by atomic number (number of protons) not by atomic mass.
8.1: Characteristics of Many-Electron Atoms
A. The Electron-Spin Quantum Number
The three quantum numbers, n, l and m describe the size (energy), shape and orientation respectively.
An additional quantum number describes the spin, which is a property of an electron not the orbital.
When a beam of atoms that have one or more lone pairs of electrons passes through a magnetic field, it splits into
two beams. Each electron behaves like a spinning charge and generates a tiny magnetic field, which can have one of
two values of spin.
The two electrons have opposing directions, so half of the electrons are attracted by the large external magnetic field
while the other half are repelled by it.
The spin quantum number (m ) has two possible values: +½ or –½.
Each electron in an atom is described completely by a set of four quantum numbers: the first three describe its
orbital, and the fourth describes its spin.
B. The Exclusion Principle
Pauli Exclusion Principle states that: no two electrons in the same atom can have the same four quantum numbers.
The major consequence is that an atomic orbital can hold a maximum of two electrons and they must have opposing
C. Electrostatic Effects and Energy-Level Splitting
Electrostatic Effects: attraction of opposite charges and repulsion of like charges
They play a major role in determining the energy states of many-electron atoms.
The Hydrogen atom is only effected by the n value whereas other elements are also affected by electron-electron
These additional interactions give rise to the splitting of energy levels into sublevels of differing energies: the energy
of an orbital in a many-electron atom depends mostly on its n value (size) and to a lesser extent on its l value
This energy difference arises from three factors:
1. Nuclear Attraction
2. Electron Repulsions
3. Orbital Shape
These three factors lead to shielding and penetration that occur in all elements but H.
The Effect of Nuclear Charge (Z) on Sublevel Energy.
o Higher charges interact more strongly than lower charges (according to Coulomb’s Law)
o Thus, a higher nuclear charge increase nucleus-electron attraction and thus lowers sublevel energy.
o Least Stable = Highest Energy sublevel, least energy required to remove it.
Sheilding: The Effect of Electron Repulsion on Sublevel Energy
o each electron feels not only the attraction to the nucleus but also repulsion from other electrons
o repulsions counteract nuclear attraction, making each electron easier to remove by helping to push it away
o each electron “shields” the other electrons from the nuclear charge somewhat
o Shielding (also called screening) reduces the full nuclear charge to an effective nuclear charge eff), which is
the nuclear charge an electron actually experiences, and this lower nuclear charge makes the electron easier
Shielding by other electrons in a given energy level: electrons in the same energy level shield each
ex. He and He : both have a 2+ nuclear charge (two protons) but He has 2e in the 1s
sublevel and He only has one, thus it takes less than ½ as much energy to remove an e
from He -
Shielding by Electrons in in inner energy levels: b/c inner e spend all their time b/w the outer
electrons and nucleus, they cause a much lower Z effhan do e in the same level. This makes it easier
to remove e from outer levels.
Penetration: The Effect of Orbital Shape on Sublevel Energy
o Penetration has two effects:
it increases the nuclear attraction for a 2s electron over that for a 2p electrons
it decreases the shielding of a 2s electron by the 1s electron
Splitting of Levels into Sublevels
o penetration and the resulting effects on shielding cause an energy level to split into sublevels of differing
o the lower the l value of a sublevel, the more its electrons penetrate, and so the greater their attraction to the
o thus, for a given n value, a lower l value indicates a more stable (lower energy) sublevel
Order of sublevel energies: s < p < d < f
8.2: The Quantum-Mechanical Model and The Periodic Table
A. Building the Periodic Table
the ground state electron configuration = the lowest-energy distribution of electrons in the sublevels of its atoms.
Aufbau Principle: electrons are always placed in the lowest energy sublevel available.
The Pauli Exclusion Principle states that each orbital may contain a maximum of 2 electrons, which must have
There are two common ways to indicate the distribution of electrons:
o Electron configuration is indicated by a shorthand notation which consists of the principle energy level (n
value), the letter designation of the sublevel (l value), and the number of electrons in the sublevel
o Orbital Diagram: Orbital diagrams make use of a box, circle, or line for each orbital in the energy level. An
arrow is used to represent an electron and its spin.
Hund’s Rule: specifies that when orbitals of equal energy are available, the lowest energy electron configuration has
the maximum number of unpaired electrons with parallel spins.
Elements in the same group have similar outer electron configurations and similar patterns of reactivity
Partial orbital diagrams: shows only the sublevels being filled
Condensed Electron Configuration: A condensed electron configuration has the element symbol of the previous
noble gas in square brackets. Al has the condensed configuration [Ne]3s 3p
Elements in the same group of the periodic table have the same outer electron configuration.
Elements in the same group of the periodic table exhibit similar chemical behavior.
Similar outer electron configurations correlate with similar chemical behavior.
Sublevels are filled in order of increasing energy, which leads to chemical properties that recur periodically
B. Building Up the Transition Series:
Effects of Shielding and Penetration on sublevel energy: the 3d sublevel is filled in period 4, but the 4s level is filled
first, this switch in filling order is due to shielding and penetration effects.
o Based on the 3d radial probability distribution, a 3d electron spends most of the time outside the filled inner
n = 1, and n=2 levels, so it is shielded very effectively from nuclear charge.
o The outermost 4s electron penetrates close to the nucleus part of the time, so it is subject to greater
o Thus, the 4s orbital is slightly lower in energy than the 3d and fills first
o In any period the ns sublevel fills before the (n-1)d sublevel
Filling the 4s and 3d sublevels:
o filling the 3d orbitals proceeds at one electron a time with the exception of Cr and Cu
Stability of Half-Filled and filled sublevels
o half-filled and full sublevels are the most stable 2 4 1 5
o Cr – instead of [Ar]4s 3d as expected has [Ar]4s 3d , an electron from the s sublevel moves to the d
sublevel to increase stability
C. Categories of Electrons
Inner (core) electrons are those an atom has in common with the pervious noble gas and any completed transition
series. The fill all the lower energy levels of an atom
Outer electrons are those in the highest energy level (highest n value), they spend most of their time farthest from
Valence Electrons: are those involved in forming compounds
o For main group elements, the valence electrons are the outer electrons
o For transition elements, the valence electrons include the outer electrons and any
Group and Period Numbers
among the main group elements the A number equals the number of outer electrons
the period number is the n value of the highest energy level
for any energy level, n = the number of orbitals, and 2n = maximum number of electrons
D. Intervening Series: Transition and Inner Transition Metals
Transition Series: Periods 4,5,6 and 7 have 3d,4d,5d and 6d orbitals respectively. the order of filling is ns, (n-1)d, np
Inner Transition Series: f orbitals, the f orbitals have l=3, so the possibll m values are [-3,-2,-1,0,+1,+2,+3]
o the period 6 inner transition series are called the lanthanides
o the period 7 inner transition series are called the actinides
o the filling order is: ns, (n-1)d, (n-2)f , np
8.3: Trends in Three Atomic Properties
A. Trends in Atomic Size
Atomic size varies from substance to substance
1. Metallic Radius: it is one-half of the shortest distance b/w nuclei of adjacent, individual atoms in a crystal of the
element (ex. Aluminum)
2. Covalent Radius: used for elements occurring as molecules, it is one-half of the shortest distance b/w nuclei of
bonded atoms (ex. Chlorine)
Main Group Elements
Among the main group(1-8) elements atomic size varies within both groups and periods as a result of two opposing
1. Changes in n: as the principal quantum number ,n, increases, the probability that outer electrons spend most of
their time farther from the nucleus also increases, thus atomic size increases (down the group). As we move
down a group each member has one more level of inner electrons that shields the outer electrons, thus atomic
radius increases down a group.
2. Changes in Zeff : As the effective nuclear charge increase, outer electrons are pulled closer to the nucleus; thus
atomic size decreases (across a period). Across a period, electrons are added to the same outer level, so the
shielding by inner electrons doesn’t change. Thus, atomic radius decreases across a group
1. Down a transition group n increases but shielding by an additional level of inner electrons results in only a
small change in size from period 4 to 5 and none from 5 to 6.
2. Across a transition period, atomic size shrinks through the first two or three elements because of the incr