PERIODIC TABLE TENDENCIES
A. Effective nuclear charge – the attractive positive charge of nuclear protons acting on valence electrons.
1. The effective nuclear charge is always less than the total number of protons present in a nucleus due to shielding effect.
2. Effective nuclear charge is behind all other periodic table tendencies.
B. Shielding effect – the lessening of attractive electrostatic charge difference between nuclear protons and valence electrons by partially or fully filled inner shells.
1. Shielding effect increases with the number of inner shells of electrons.
2. Electrons sharing the same shell do not shield one another from the attractive pull of the nucleus.
C. An estimate of effective nuclear charge can be obtained from
Zeff = Z − S,
where Zeff = effective nuclear charge, Z = atomic number, and, S = the screening constant.
1. The screening constant can be estimated by setting it equal to the number of core electrons present in an atom.
2. It is often used as a measure of the shielding effect present in an atom.
3. Consider aluminum…
[Ne]3s23p1 Z = 13 S = 10 Zeff = Z − S = 13 – 10 = 3+
4. Don’t forget that Zeff is only an estimate. Actual shielding effect is always greater that the screening constant S because core electrons are much closer to the nucleus than are valence electrons.
D. The periodic table tendency for effective nuclear charge:
1. Increase across a period (due to increasing nuclear charge with no accompanying increase in shielding effect).
2. Decrease down a group (although nuclear charge increases down a group, shielding effect more than counters its effect). E. Effective nuclear charge can be used to explain the octet rule.
1. The octet rule is the very strong tendency of an element to adopt the electron configuration of the Noble gas (numerically) closest to it in the periodic table.
2. The octet rule for potassium is to lose one electron.
a) The electron configuration for potassium is [Ar]4s1, so the loss of a single 4s electron leaves it with the configuration of the Noble gas argon.
b) The loss of that electron turns out to be relatively easy to accomplish since the effective nuclear charge on it is actually quite small due to the shielding effect of three inner shells.
3. The octet rule for sulfur is to gain two electrons.
a) The electron configuration for sulfur is [Ne]3s23p4, so the gain of two electrons gives it with the configuration of the Noble gas argon.
b) The gain of those electrons turns out to be relatively easy to accomplish since the effective nuclear charge on them is actually the same as that already felt by the six valence belonging to sulfur and since the addition of the two new electrons will not change the shielding effect in the atom.
II. Atomic Radii.
A. Simply stated, atomic radius an estimate of the size of an atom.
1. Nonbonding atomic radius – the smallest distance separating the nuclei of two atoms of the same chemical element during an elastic (or non-bonding) collision. This distance is also called the van der Waals radius.
2. Bonding atomic radius – the smallest distance separating the nuclei of two atoms of a chemical bond. Because bonded atoms have overlapping electron clouds, the bonding atomic radius is always less that the van der Waals radius.
B. The periodic table tendency for atomic radius:
1. Decrease across a period (due to increasing nuclear charge with no accompanying increase in shielding effect).
2. Increase down a group (although nuclear charge increases down a group, shielding effect more than counters its effect).
III. Ionic Radii.
A. Simply stated, ionic radius is an estimate of the size of an ion.
B. The periodic table tendency for ionic radius:
1. Cations are smaller than their atoms. There are two effects at work here:
a) The loss of valence electrons results in the loss of the outermost shell.
b) The remaining electrons actually experience an increase in nuclear charge with no accompanying increase in shielding effect. 2. Anions are larger than their atoms (additional electrons with no accompanying change in either nuclear charge or shielding effect).
3. Increase down a group for ions carrying the same charge. Why?
4. Decrease across an isoelectronic series.
a) An isoelectronic series is a series of ions all of whom have the same electron configuration.
b) For example; O-2, F-1, Na+1, Mg+2, and Al+3 all have the same electron configuration, namely, [Ne].
c) Arranged in order of increasing atomic number, the ions experience an increase in nuclear charge with no accompanying increase in shielding effect, and decrease in size.
IV. Ionization Energy.
A. The ionization energy of an atom or ion is the minimum amount of energy necessary to remove an electron from an isolated gaseous atom or a ground state ion.
B. Every electron in an atom or ion has a characteristic ionization energy, and in general,
1. Valence electrons have lower ionization than core electrons. Why?
2. First ionization energies are lower than successive ionization energies.
a) The first ionization energy, I1, is the energy associated with the removal of the first electron from a neutral atom:
S → S+1 + 1e-
b) The second ionization energy, I2, is the energy associated with the removal of the second electron:
S+1 → S+2 + 1e-
C. The periodic table tendency for first ionization energy:
1. Generally increase across a period (due to increasing nuclear charge with no accompanying increase in shielding effect).
a) Important exceptions to this tendency occur due to subshell effects.
b) For example, removing a p1 electron is generally easier than breaking up an s2 pair.
c) It is also easier to remove the p4 electron than the p3. Why?
2. Generally decrease down a group (although nuclear charge increases down a group, shielding effect more than counters its effect). V. Electron Affinity.
A. The energy change that occurs when a gaseous atom or ground state ion gains an electron is termed electron affinity.
B. Generally, the addition of an electron to a neutral atom is an exothermic process, so electron affinities are usually reported as negative quantities. For example, the formation of the bromide is generally written as
Br + 1e- → Br-1 ΔE = - 325 kJ mol-1
C. The periodic table tendency for electron affinity:
1. Generally becomes increasingly negative across a period.
a) It is important to keep in mind that the more negative the value for electron affinity becomes, the greater the attraction for an additional electron.
b) It is therefore accurate to state that electron affinity generally increases across a period.
2. Generally becomes decreasingly negative down a group, although this is a very small effect in most groups.
VI. Metallic Character.
A. The stair-step line divides the periodic table into metals, nonmetals, and metalloids.
B. Metals lie below and to the left of the stair-step line. They include the elements aluminum and polonium that border the stair-step line.
C. Properties of metals include:
1. Shiny luster
2. Conductivity of heat and electricity
3. Malleability (can be rolled into thin sheets)
4. Ductility (can be pulled into thin wires)
5. High melting point
6. Low first ionization energy
7. Form ionic compounds with nonmetals
8. Form basic oxides
D. The periodic table tendency for metallic character:
1. Generally decreases across a period.
2. Generally increases down a group.