Chapter 5: Section 3
Electron Configuration and Periodic Properties
atomic radius- may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.
The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus.
*As electrons are added to s and p sublevels in the same main energy level, they are gradually pulled closer to the more highly charged nucleus.
The atomic radii of the main-group elements increase down a group.
*By filling the fourth main-energy level is outweighed by a shrinking of the electron cloud caused by a nuclear charge that is considerably higher than that of aluminum.
(In the picture below, as the shade gets lighter across, the atomic radius is decreasing. As it gets darker going down the atomic radius is increasing)
*An electron con be removed from an atom if enough energy is supplied.A + energy → A⁺ + e⁻
A+ represents an ion of element A with a single positive charge, referred to as 1+ ion.
ion- an atom or group of bonded atoms that has a positive or negative
ionization- any process that results in the formation of an ion
ionization energy, IE-(or first ionization energy, IE1)- The energy required to remove one electron from from a neutral atom of an element
*measurements of ionization energies are made on isolated atoms in the gas phase
Group 1 metals have the lowest first ionization energies in their respective periods and they lose electrons most easily
Group 18 element, noble gases, have the highest ionizations energies, the don’t lose electrons easily.
Ionization energies of the main-group elements increase across each period.
*This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level.
*Increasing nuclear charge is responsible for both increasing ionization energy and decreasing radii across the periods.
*nonmetals have higher ionization energies than metals do.
Among the main-group elements, ionization energies generally decrease down the groups.
*Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus.
*They are removed easily
*Atomic number increases going down a group, more electrons lie between the nucleus and the electrons in the highest occupied energy levels.
Removing Electrons from Positive Ions
*The energies for removal of additional electrons from an atom are referred to as the second ionization energy, third ionization energy, and so on.
*Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge (the nuclear charge minus the electron shielding).
Electron Affinity- the energy that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.
A+ e- -----> A- + energy
On the other hand, some atoms must be forced to gain an electron by the addition of energy
A+ e- + energy ------> A-
The quantity of energy absorbed would be represented by a positive number, but ions produced in this way are very unstable and hence the electron affinity for them is very difficult to determine. An ion produced in this way will be unstable and will lose the added electron spontaneously.
Among the elements of each period, the halogens gain electrons most rapidly. The ease with which halogen atoms gain electrons is a major reason for the high reactivities of the Group 17 elements.
In general, as electrons add to the same p sublevel of atoms with increasing nuclear charge, electron affinities become more negative across each period within the p block.
(An exception to this occurs between Groups 14 and 15.) Compare the electron affinities of carbon and nitrogen.) Adding an electron to a carbon atom gives a half-filled p sublevel. This occurs much more easily than forcing an electron to ait with another electron in an orbital of the already half-filled p sublevel of a nitrogen atom.
Trends in groups are not as as regular as trends for ionization energies. Electrons add with greater difficulty down a group. This pattern is a result of two competing factors. The first increase in effective nuclear charge down a group, which increases electron affilities. The second is an increase in atomic radius down a group, which decreases electron affinities.
In general, the size effect predominates. But there are exceptions, especially among the heavy transition metals, which tend to be the same size or even decrease in radius down a group.
Adding Electrons to Negative Ions
For an isolated ion in the gas phase, it is always more difficult to add a second electron to an already negatively charged ion. Therefore, second electron affinities are all positive.
Certain p-block nonmetals tend to form negative ions that have noble gas configurations. The halogens do so by adding one electron.
cation- a positive ion
*The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius because the removal of the highest-energy-level electrons results in a smaller electron cloud.
anion- a negative ion
*The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius.
*Within each period of the periodic table, the metals at the left tend to form cations and the nonmetals at the upper right tend to form anions.
*anions are more common than cations
*the outer electrons in both cations and anions are in higher energy levels as one reads down a group.
Valence electrons- the electrons available to be lost, gained, or shared in the formation of chemical compounds
*Valence electrons hold atoms together in chemical compounds*chemical compounds form because electrons are lost, gained, or shared between atoms.
*Electronegativity- a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
*Electronegativities tend to increase across each period, although there are exceptions
*The alkali and alkaline-earth metals are the least electronegative elements
-they have low attraction for electrons
*Nitrogen, oxygen, and the halogens are the most electronegative elements.
-attraction is strong in compounds
*Electronegativities tend to either decrease down a group or remain about the same.
*when a noble gas does form a compound, its electronegativity is rather high, similar to the values for the halogens.
Periodic Properties of the d- and f-Block elements
*The d-block elements vary less and with less regularity than those of the main-group elements.
*The atomic radii of the d-block elements generally decrease across the periods.
*As the number of electrons in the d sublevel increases, the radii increase because of repulsion among the electrons.
*the radii of elements following hafnium in the sixth period vary with increasing atomic number in the usual manner
*ionization energies of the d-block and f-block elements generally increase across the periods.
*the first ionization energies of the d-block elements generally increase down each group.*The electrons available for the ionization in the outer s sublevels are less shielded from the increasing nuclear charge by electrons in the incomplete (n-1)d sublevels.
Ion Formation and Ionic Radii
*Electrons in the highest occupied sublevel are always removed first
*The d-block elements all have electronegativities between 1.1 and 2.54.
*The active metals in Groups 1 and 2 have lower electronegativities
*The f-block elements all have similar electronegativities, which range from 1.1 to 1.5.