Daniel L. Reger Scott R. Goode David W. Ball Chapter 8 The Periodic Table: Structure and Trends

The periodic table can be divided into four blocks of elements: elements with highest energy electrons in s, p, d, or f subshells. The arrangement of the elements in the periodic table correlates with the subshell that holds the highest energy electron. Electron Configurations and the Periodic Table

Using only the periodic table, determine the electron configurations of Al, Ti, Br, and Sr. Example: Electron Configurations

For anions, the additional electrons fill orbitals following the same rules that applies to atoms. Cl: [Ne] 3s 2 3p 5 Cl - : [Ne] 3s 2 3p 6 As: [Ar] 4s 2 3d 10 4p 3 As 3- : [Ar] 4s 2 3d 10 4p 6 Many stable anions have the same electron configuration as a noble gas atom. Electron Configurations of Anions

For the electron configurations of cations, electrons of highest n value are removed first. For cases of the same n level, electrons are first removed from the subshell having highest. As: [Ar] 4s 2 3d 10 4p 3 As 3+ : [Ar] 4s 2 3d 10 Mn: [Ar] 4s 2 3d 5 Mn 2+ : [Ar] 3d 5 NOTE: For d-block atoms, the ns electrons are removed before the (n-1)d electrons. Electron Configurations of Cations

Test Your Skill Write the electron configurations of the following ions: (a) N 3- (b) Co 3+ (c) K +

An isoelectronic series is a group of atoms and ions that contain the same number of electrons. The species S 2-, Cl -, Ar, K +, and Ca 2+ are isoelectronic – they all have 18 electrons. Isoelectronic Series

An atomic radius is one half the distance between adjacent atoms of the same element in a molecule. Atomic Radii 198/2 = /2 = 114 Sum = 215

Size Trends for an Isoelectronic Series

Sizes of the Atoms and Their Cations Atoms are always larger than their cations.

Sizes of the Atoms and Their Cations If an atom makes more than one cation, the higher-charged ion has a smaller size.

Anions are always larger than their atoms. Atomic and Ionic Radii

Identify the larger species of each pair: (a) Mg or Mg 2+ (b) Se or Se 2- Test Your Skill

Atomic Radii of Main Group Elements

The sizes of atoms are impacted by the effective nuclear charge felt by the outermost electrons. Sizes of Atoms

The sizes of atoms increase going down a group. Effective Nuclear Charge & Size

The increase in effective nuclear charge causes a size decrease across the period. Sizes of Atoms

Test Your Skill Identify the larger species of each pair: (a) Mg or Na (b) Si or C

The ionization energy is the energy required to remove an electron from a gaseous atom or ion in its electronic ground state. Ionization Energy

An atom has as many ionization energies as it has electrons. Example: Mg(g) → Mg + (g) + e - I 1 = first ionization energy Mg + (g) → Mg 2+ (g) + e - I 2 = second ionization energy Ionization Energies

The increase in the effective nuclear charge across a period causes an increase in the ionization energy as you go across that period. Trends in 1 st Ionization Energies

The slight dip in ionization energy for O is because the fourth p electron now pairs with another electron, slightly repelling each other. Trends in 1 st Ionization Energies

Trends in First Ionization Energies

Isoelectronic species with the greatest charge in the nucleus will have the largest ionization energy. For the isoelectronic series S 2-, Cl -, and Ar, Ar has the largest ionization energy because it has the most protons in its nucleus. Ionization Energy Trends in Isoelectronic Series

Predict which species in each pair has the higher ionization energy. (a) Ca or As (b) K + or Ca 2+ (c) N or As Ionization Energy

Successive ionization energies always increase because of the increasing hold the nucleus has on remaining electrons. I 1 I 2 I 3 I 4 Mg Al A much larger increase is seen when an electron comes from a lower-energy subshell. (all values in kJ/mol) Successive Ionization Energies

Which element, magnesium or sodium, has the greater second ionization energy? Test Your Skill

The electron affinity of an element is the energy change that accompanies the addition of an electron to a gaseous atom to form an anion. A(g) + e - → A - (g) Electron affinities are generally favorable (exothermic) for elements on the right side of the periodic table. Electron Affinity

Electron Affinities

Alkali Metals – Group 1A (1) The reactivity of the Group 1A metals increases down the group. Their chemistry is dominated by the formation of M + ions. 2M(s) + H 2 O( ) → 2MOH(aq) + H 2 (g) 2M(s) + H 2 (g) → 2MH(s) 2M(s) + X 2 (g) → 2MX(s) X = F, Cl, Br, I

Only lithium reacts with O 2 to give the expected product, lithium oxide. 4Li(s) + O 2 (g) → 2Li 2 O(s) Sodium reacts mainly to yield sodium peroxide. 2Na(s) + O 2 (g) → Na 2 O 2 (s) Potassium reacts to yield mixtures of the oxide, peroxide, and superoxide. K(s) + O 2 (g) → KO 2 (s) Alkali Metal Reactions with O 2

Flame Colors of the 1A Elements

The Alkaline Earth Metals – Group 2A (2) The Group 2A metals are not as reactive as the Group 1A metals. Reactivity increases down the group, and they all form M 2+ ions. Magnesium alloys are useful in aeronautical applications, where low density and high strength are important.

Flame Colors of 2A Elements CalciumStrontiumBarium

The halogens all exist as diatomic molecules, but they are very reactive. The reactivity decreases as you go down the group. Their chemistry is dominated by the formation of X - ions. The interhalogens are compounds formed from different halogens, like IF 3 and BrCl. The Halogens – Group 7A (17)