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Learn about atomic radii trends, ionization energies, electron affinities, cations, anions, and valence electrons. Understand why exceptions occur and how to predict chemical properties. Explore group and period trends in the periodic table.
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Section 5.3 Electron Configuration And Periodic Properties
Trends • Trends are approximations of overall patterns. This section discusses trends that we see in elements in the periodic table. Remember that exceptions can, and do, occur.
Explaining trends… • Properties such as the size of an atom (atomic radius), the energy required to remove an electron from an atom (ionization energy), and the energy associated with the addition of an electron to a gaseous atom (electron affinity) can be understood in terms of the electron configuration of the atom and a competition between electron-nucleus attraction and electron-electron repulsion.
Atomic Radii (see p. 140) • The size of any atom is defined by the edge of its orbital. But since this edge is fuzzy the atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together.
Period Trends of Atomic Radii • Trend: Smaller atoms across a period (from left to right, see p. 141). • Caused by the increasing positive charge of the nucleus (increased electron-nucleus attraction). • Somewhat offset by the repulsion among the number of electrons in the same outer energy level.
Group Trends of Atomic Radii • In general, the atomic radii increase down a group. • n increases and so does the distance from the nucleus • Core electrons shield outer electrons • attractive force b/w nucleus and outer electrons is less • Trend does not always hold (see Al to Ga, it decreases). • Difference is small, for our purposes we will ignore it.
Ionization Energies • Electrons from an element can be removed if enough energy is supplied. • Shown by: A + E → A+ + e– An ion is an atom or group of bonded atoms that has a positive or negative charge.
Ionization • Any process that results in the formation of an ion.
Ionization Energy (see p. 143) • The energy required to remove one electron from a neutral atom of an element (often called the first ionization energy). • To remove a second e– is called the 2nd ionization energy, etc..
Ionization Energy Period Trends (see p. 144) • In general, ionization energies of the main-group elements increase across each period. • Caused by increasing nuclear charge increased attraction b/w nucleus and electrons.
Removing Electrons From Positive Ions • See p. 145 • 2nd, 3rd, 4th, etc. ionization energies are successively higher. Notice the pattern at the stairstep. • Due to the remaining electrons feeling an increasingly stronger effective nuclear charge (the nuclear charge minus the electron shielding).
Electron Affinity (p. 147) • The energy change that occurs when an electron is acquired by a neutral atom.
Exothermic Reaction • Most elements release energy when they gain an electron. The numerical value is negative. A + e-→ A- + E
Endothermic Reaction • Some elements must be “forced” to gain an electron by the addition of energy: • A + e- + E → A- • The energy value is positive. Such ions are unstable, and will lose the electron spontaneously.
Period Trends • Increase in negativity from left to right, as electrons add to the p sublevel. Not true between Groups 14 and 15. 14: ns2 np2 15: ns2 np3
Group Trends – Electron Affinity • Electrons add with greater difficulty down a group (more positive values). Mainly due to the increase in atomic radius.
Cations – positive ions • Loss of e- always leads to a decrease in size because removal leads to a smaller electron cloud. • The remaining e- are drawn closer to the nucleus by the unbalanced positive charge.
Anions – negative ions • Adding electrons leads to larger ionic size because the e- are not drawn as strongly to the nucleus and the e-cloud spreads out due to increasing repulsion between the e-.
Group Trends • In general, because of the increase in atomic radii, both the anions and cations of a group are larger as one goes down a group. Exceptions occur in some of the p-block groups as elements change from non-metals to metals across the staircase. Why?
Period Trends • To the left of the staircase, metals form cations (+ charge). • All metals tend to lose the electrons from the s orbitals first. • To the right and above the staircase, nonmetals tend to form anions (- charge)
Cations • Across a period cationic radii decrease because the electron cloud shrinks due to the increasing nuclear charge.
Anions • Anionic radii decrease across a period for the same reason.
Valence Electrons • Valence electrons are the electrons available to be lost, gained, or shared in the formation of a chemical compound.
Numbers of Valence Electrons • For the main-group elements: Groups 1 and 2; equal to the group number. Groups 13-18; equal to the group number minus 10.
Common Ionic Charge • Atoms gain/lose electrons to achieve a noble gas electron configuration.
Electronegativity • A measure of the ability of an atom in a chemical compound to attract electrons.
Fluorine Reference Point • The most electronegative element was arbitrarily assigned a value of four and all other element values are calculated in relation to this value.
Period & Group Trends • Electronegativities tend to increase across each period and up each group, although there are exceptions.
Metallic Character • Metallic character is the name given to the set of chemical properties associated with elements that are metals. These chemicals properties result from how readily metals lose their electrons.
