CH 7: Atomic Structure and Periodicity

CH 7: Atomic Structure and Periodicity Sections 7.10 -7.13

Periodic Trends • Models explain observed behavior. • The better the model the fewer the exceptions • Consider computer weather models vs. kinetic molecular theory

Periodic Trends • The quantum mechanical model of the atom explains many trends in the properties observed for the elements. • Trends in physical properties • Atomic radius • Size of the ion vs. the “parent” atom • Trends in reactivity: • Charge on the ion formed • Ease of removing or adding an electron to an atom

Atomic Radius • Measuring/defining atomic radius • Metals: atomic radius is half the distance between nuclei in a solid • Nonmetals; atomic radius is half the distance between the nuclei of atoms in a diatomic molecule Cu H H

Atomic radius trends (pg 276) • Atomic radius increases down a group • Valence electrons are in higher (larger) principal quantum levels with increased shielding. • H 1s1 • Li …..2s1 • Na ……......3s1 • K ………………..4s1

Atomic radius trends • Atomic radius decreases across a period of representative elements • Valence shell (PEL) remains the same across a period, same shielding across the period……however… • The # protons increases across a period • The increased nuclear charge “pulls” shells closer to the nucleus

Atomic Radius Consider the 2nd period…filling n = 2 Li Be B C N O F Ne # p 3 4 5 6 7 8 9 10  decreasing atomic radius

Atomic radius • Atomic radius remains ~same across a row of transition metals • Why?

Ionization Energy • Ionization Energy – energy needed to remove the highest energy electron from an atom in its gaseous state. • See page 272/273, IE > 0 Na(g) Na+ (g) + e IE1 = 495 kJ/mole

IE Trends • First IE (IE1 ) becomes less endothermic (less +) down a group • See table 7.5 on page 272 • Why? • As you go down a group, the electron being removed is farther from the nucleus and shielded by more core electrons from the attractive forces of the nucleus. • Therefore, it’s easier to remove.

IE Trends • In general, first IE (IE1 )increases across a period. • See figure 7.31 on page 273 • Why? • Atoms become smaller across a period and the # core electrons (shielding) remains the same while nuclear charge increases. • Electron to be removed is held more tightly to the nucleus across a period.

Exceptions to IE Trends • A dip in IE1 is observed for elements in group 3A and 6A. • 3A elements are all ns2p1 • Hypothesized that the s2 electrons shield the first p electron • 6A elements are all ns2p4 • Hypothesized that the first pairing of p electrons increases repulsions and thus this electron is easier to remove.

Trends in Successive IE • IE increases as additional electrons are removed from a given element • see table 7.5 on page 272 Na(g) Na+ (g) + e IE1 = 495 kJ/mole Na+ (g)  ____ + e IE2 = 4560 kJ/mol

Trends in Successive IE • IE jumps when the first core electron is removed. • Why? Na(g) Na+ (g) + e IE1 = 495 kJ/mole (val. e) Na+ (g)  ____ + e IE2 = 4560 kj/mol (core e)

Electron Affinity • EA – energy change associated with the addition of an electron to a gaseous atom. • In this text, EA < 0 (convention varies) • See page 275 X (g) + e  X-(g)

EA Trends • MANY EXCEPTIONS! • In general, EA becomes less negative down a group. • In general, EA becomes more negative across a period.

Periodic Trends • Atomic radius • Ionization Energy (>0) • First IE and successive IE • 3A and 6A exceptions • Electron Affinity (<0)

CH 7: Atomic Structure and Periodicity