Daniel L. Reger Scott R. Goode David W. Ball Lecture 08 (Chapter 8) The Periodic Table: Structure and Trends

The experimental trends in group properties on which the periodic table was based can now be explained by the arrangements of electrons in atoms. The electron configurations of the valence electrons in each group member are similar. Valence electrons: Electrons with highest principle quantum number in atom, and any electrons in an unfilled subshell from a lower shell. Valence orbitals: Orbitals of the highest principle quantum number in atom, and (for d or f electrons), the orbitals of any partially-filled subshells of lower principle quantum numbers. Periodic Trends of the Elements

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

The ordering of orbitals with respect to energy is reflected in the periodic table (i.e., we can count electrons on the periodic table to determine addresses for all of the electrons in the electron configuration). Elements with one electron in a new principal shell (ns 1 ) start a new period in the periodic table. Electron Configurations and the Periodic Table

The 4s orbital is lower in energy than the 3d orbital and fills first, starting the fourth period at potassium. The 3d orbitals fill after the 4s. Similar inversions occur in the remaining periods. Electron Configurations

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. See Pb 2+ as example. Electron Configurations of Cations

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

Test Your Skill Write the electron configurations of the following ions: (a) N 3- (b) Co 3+ (c) K + Answers: (a) 1s 2 2s 2 2p 6 (b) [Ar] 3d 6 (c) [Ar]

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. Would these species be stable, and why? Isoelectronic Series

Since an electron cloud is a “fuzzy” probability function, the atomic radius can be calculated based on the half distance between adjacent atoms of the same element in a molecule. This method does not work as well for metal atoms. Atomic Radii 198/2 = /2 = 114 Sum = 213

Sizes of the Atoms and Their Cations Atoms are always larger than their cations. If the electrons are removed from an orbital, then there is less probability of finding an electron in that orbital. 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

Size Trends for an Isoelectronic Series

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

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

The sizes of atoms are impacted by the effective nuclear charge (Z eff ) 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

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

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

Property trends Atomic radius decreases with addition of protons to nucleus which pulls e- closer to center (Electronegativity). Decreasing radius makes it more difficult to remove e- (ionization energy is energy required to remove e-) Atomic radius increases Ionization energy decreases Electronegativity decreases Small radius e- held tightly Large radius e- held loosely

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. Exceptions: Group 3A. The np 1 electron does not penetrate inner electrons as much as ns 2 electrons. Group 6A. First pairing of electrons in p orbital produces small repulsion between electrons. Trends in 1 st 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 (therefore, the most positive charge) 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. Based on periodic table, do these numbers make sense? (all values in kJ/mol) Successive Ionization Energies

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

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

The electron affinity of an element is the energy change (in kJ/mol) 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 (i.e., these non-metal elements are more likely to gain electrons than lose them). Electron Affinity

Electron Affinities

Alkali Metals – Group 1A (1) Group 1A metals very reactivie, and reactivity increases down the group. Their chemistry is dominated by the formation of M + ions (easier to remove a single electron as atomic radius increases). 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

Reactions of Group 1A metals with molecular oxygen do not necessarily follow assumed trends. 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

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 (loss of 2 electrons results in a noble gas electron configuration).

The halogens all exist as diatomic molecules, but they are very reactive, and tend to gain 1 electron to achieve a noble gas configuration. The reactivity decreases as you go down the group. The interhalogens are compounds formed from different halogens, like IF 3 and BrCl. The Halogens – Group 7A (17)

Summary As Atomic radius increases, electrons are held more distantly from protons, and are easier to remove as the effective nuclear charge is reduced. Metals (with higher atomic radii) tend to lose enough electrons to achieve a noble gas electron configuration. Once they have done so, it is more difficult to remove more electrons (requires higher ionization energy). Non-metals (with smaller atomic radii) tend to gain enough electrons to achieve a noble gas electron configuration. Once they have done so, it is more difficult to add more electrons (electron affinity decreases as they are already in a stable octet)