logo imageChapter 6 Notes
Dr. Floyd Beckford

CHAPTER 6
CHEMICAL PERIODICITY

Remember the Periodic Law: The properties of the elements are periodic functions of their atomic numbers.

PERIODIC PROPERTIES OF THE ELEMENTS
I. Atomic Radii
The radius of an atom is taken as half of the distance between the nuclei in homonuclear molecules.
In an atom, there exist an electrostatic attraction between the electrons and the nucleus. There is also electrostatic repulsion between the electrons. Thus the attraction of an electron to the nucleus is affected by the presence of the other electrons. The nuclear attraction for the other electrons is mitigated by the repulsion of the electrons in place. We say that the outer electrons are screened or shielded by the innermost electrons. The charge felt by the outer electrons is referred to as the effective nuclear charge, Zeff.

Atomic sizes of the representative elements show two periodic trends: sizes increase going from top to bottom of a group and decrease from left to right across a period. This is in agreement with the two major effects that control atomic properties: the outermost orbital of an atom and its Zeff.

Going from Li to F, the increase in nuclear charge pulls the outer electrons closer to the nucleus, thereby decreasing the atomic size. (The increase in Zeff is far greater than the corresponding increase in shielding).

Going from H to Cs, the outermost filled orbital increases in value of n. Coupled with the fact that the effective nuclear charge remains roughly the same down the group, we find that atomic size increases down the group.

II. Ionic Radii
In forming an anion, the number of electrons in the parent atom increases. The increase in electrons creates more shielding, thus decreasing the effective nuclear charge. Each of the electrons is held less tightly (due to increase electron-electron repulsion) and the ion is larger than the atom.

Conversely, when a cation is formed, the number of electrons decreases and the remaining electrons are held more tightly. The ion is therefore smaller than the parent atom.

A general rule is that anions are larger than the atoms from which they are formed and cations are smaller than their parent atoms.

Er > Er+ and Er < Er-
To a limited extent, trends in ionic sizes follow those of atomic sizes. Moving down a group one finds that like-charged ions grow in size, as expected by their larger orbitals. Moving left to right across a period, ion sizes generally decrease from one group to the next except when changing from cations to anions.

Some elements contain the same number of electrons. These species are called isoelectronic. Properties change regularly with atomic number across an isoelectronic series. So within an isoelectronic series, ionic radii decrease with increasing atomic number because of increasing nuclear charge.

III. Ionization Energy
To remove an electron form an atom requires energy. The ionization energy is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form an ion with a +1 charge.


	Mg(g) ---> Mg+(g) + e- 		I.E.(1) = 738 kJ/mol
Each atom can have a series of ionization energies, because more than one electron can be removed:

	Mg+(g) ---> Mg2+(g) + e-	I.E.(2) = 1451 kJ/mol
IE(1) < IE(2) because it is more difficult to remove an electron from an ion than from a neutral atom. This is because in an ion, the remaining electrons are held tighter and so require more energy to pry them loose.

Ionization energy is a measure of how tightly electrons are bound to atoms. Elements with low ionization energies easily form cations.

Ionization energies decrease from top to bottom through a group. As we move downward, the outermost electrons are in orbitals with higher n (farther away from the nucleus) and greater size, so the electrons are held less tightly by the nucleus.

In contrast, moving from left to right across a period, the general trend is one of increasing ionization energy (except Group IIIA and VIA) due to an increase in the effective nuclear charge experienced by the outermost electrons.

The exception at Group IIIA is due to the fact that the electron is being removed from a outer p orbital. This single electron is easier to remove than the filled s orbital which is lower in energy (more stable).

At Group VIA, we are removing a paired electron from a p orbital. This process is a little easier than disrupting the half-filled p shell that is present in Group VA.

IV. Electron Affinity
What ionization energy is to the formation of cations, electron affinity is to the formation of anions. Electron affinity is the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge.

	Cl(g) + e- ---> Cl-(g)		E = Electron Affinity
NOTE: Electron affinity is not the reverse of ionization energy.
	X+(g) + e- ---> X(g)		IE reverse
The EA process starts with a neutral atom while the reverse of IE starts with a cation. The larger the negative EA value of an atom, the greater the affinity the atom has for an electron, and the more likely it is to form a stable anion.

In general electron affinity values become less negative on descending a periodic group and increase on moving from left to right across a period. As we move down the table, entering electrons enter successively higher energy orbitals and EA thereby becomes less negative. Moving left across the table leads to an increase in Zeff, with the concomitant increase in EA (EA becomes more negative). This means that as we go across the period the elements show a greater attraction for an extra electron. As in the case with ionization energies, the trends in electron affinities are not smooth. Adding an electron to Group IIA elements is difficult because we are adding the electron to a filled s orbital. Adding an electron to Group VA elements is adding to relatively stable half-filled p orbitals. Thus EA for Group IIA and VA elements are anomalies in the trends.

V. Electronegativity
The electronegativity (EN) of an element is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atom. This is in contrast to electron affinity that measures the attraction of a free electron to a neutral atom to make an anion.

For the representative elements, electronegativity usually increases from left to right across periods and from bottom to top within groups. The scale of electronegativity is arbitrary. Fluorine has the highest electronegativity and cesium (Cs) has the lowest.

CHEMICAL REACTIONS AND PERIODICITY
Hydrogen and the hydrides
Hydrogen has the lowest atomic and density of all the elements. Compounds of hydrogen are called hydrides.
a. Ionic hydrides contain H- ion and are formed by gaining an electron from a very electropositive metal such as Cs.
b. Molecular hydrides are formed by sharing its electrons with an atom of another nonmetal.

Ionic hydrides are formed with the alkali metals and the more active (starting at Ca) alkaline earth metals.

	2M(l) + H2(g) --->   2[M+,H-](s)
		M = Li, Na, K, Rb, Cs
	M(l) + H2(g) ---> [M2+,2H-](s)
		M = Ca, Sr, Ba

Ionic hydrides are all basic, meaning that they form bases when dissolved in water.

	NaH(s) + H2O(l) ---> NaOH(s) + H2(g)
	SrH2(s) + 2H2O(l) ---> Sr(OH)2(s) + 2H2(g)
Molecular hydrides are formed when hydrogen reacts with other non-metals.
	H2(g) + X2 ---> 2HX(g) 		X = F, Cl, Br, I
	2H2(g) + O2(g) ---> 2H2O
Most of the molecular hydrides are acidic; they produce acids (H+) in water.
	HCl(g) ---> H+(aq) + Cl-(aq)
Oxygen and the Oxides
Oxygen exists in two allotropic forms: O2 and O3 (ozone). Binary compounds of oxygen are called oxides. Oxides are formed by direct combination with virtually all other elements.

Metallic oxides (metal + oxygen) are usually ionic solids. The alkali metals form three kinds of binary compounds with oxygen: oxides, peroxides and superoxides.
a. Oxides contain the O2- group

	4Li(s) + O2(g) ---> 2LiO2(s) 		Lithium oxide
b. Peroxides contain the O22- group 
	2Na(s) + O2(g) ---> Na2O2(g)		Sodium peroxide
c. Superoxides contain the O2- group
	M(s) + O2(g) ---> MO2   (M = K, Rb, Cs)		Metal superoxide
The oxidation number of O is different for each of the above compounds. In oxides it is -2; it is -1 in peroxides and -1/2 in superoxides.

The alkaline earth metals react with oxygen similarly to alkali metals. 

	2M(s) + O2(g) ---> 2[M2+, O2-](s)	(M = Be, Mg, Ca, Sr, Ba)
	M(s) + O2(g) ---> [M2+, O22-]		(M = Ca, Sr, Ba)
Metallic oxides are called basic anhydrides because when they dissolve in water, there is no change in the oxidation state of the metal.
	Na2O(s) + H2O(l) ---> 2NaOH(aq)
Molecular oxides are formed when oxygen reacts non-metals. 
	S8(s) + 8O2(g) ---> 8SO2(g)
Non-metallic oxides are called acid anhydrides because oxidation state of the nonmetal do not change when they are dissolved in water. 
	SO2(g) + H2O(l) ---> H2SO3(aq)		Sulfurous acid
Oxidation number of sulfur in both SO2 and H2SO3 is +4. 
	N2O5(s) + H2O(l) ---> 2HNO3(aq)		Nitric acid
Oxidation number of nitrogen in both N2O5 and HNO3 is +5.

Acid anhydrides also react with basic anhydrides to form salts. 

	MgO(s) + SO3(g) ---> Mg SO4(s)
There is no change in oxidation state.