Chapter 7: Periodic Properties of the Elements

Electronegativity Ionization Energy & Electron Affinity

There are a number of atomic characteristics that either increase or decrease along the periodic table.

Atomic Radius:

As you go down a group, the atomic radius increases with increasing energy level.

Atomic radius

decreases

as you go left to right along a period because the greater nuclear charge pulls the electrons in closer to the nucleus.

s-block p-block

Why ionization energy decreases and atomic/ionic radius increases as go down a group: Shielding effect: The inner electron shells insulate the valence electrons from some of the electrical attraction with the positive charge of the nucleus. + nucleus core electrons valence e-

Why ionization energy increases and atomic/ionic radius decreases as go across a period: Increasing Effective Nuclear charge(Z eff ): Electrons in the outermost energy levels do not effectively screen each other from an increasingly positive nucleus. nucleus + + + + + + +

Z eff Li: Z eff = Z - S = 3 – 2 = 1 N: Z eff = 7 – 2 = 5

A few definitions:

van der Waals radius: The nonbonding radius of atoms Covalent radius: Isoelectronic: The radius of atoms covalently bonded to another atom; is smaller than the van der Waals radius Different ions that have the same number of electrons Ex: O 2 , F , Na + , Mg 2+ , Al 3+ all have 10 electrons Ionization energy: The amount of energy required to remove an electron from an atom or ion.

Electron affinity: The amount of energy released when an electron is added to an atom or ion.

There is a spike in ionization energy whenever a noble gas electron configuration is disrupted.

Q: What is the valence electron configuration for an atom that has the following ionization energies: 1 st : 734 kJ/mol, 2 nd : 1850 kJ/mol, 3 rd : 16,432 kJ/mol A: ns 2 Large spike in IE indicates noble gas core is disrupted

Electron Affinity:

the ability of an atom to gain an electron. This is closely related to ionization energy, and

increases

going left to right, and

decreases

going down.

Electronegativity

: Ability of an atom to attract electrons when in a molecule.

Electronegativity

going down. increases going left to right, and decreases

Electron configurations of ions •If electrons are added to make an anion, they fill the lowest energy levels first. (Auf bau principle) •If electrons are removed to make a cation, the are taken from the highest energy levels first.

Li Atom 1s 2 2s 1 Ion 1s 2 2s 0 Li + Fe [Ar] 4s 2 3d 6 [Ar] 4s 0 3d 6 [Ar] 4s 0 3d 5 Fe 3+ : Fe 2+ Fe 3+ Notice that the Fe has the maximum multiplicity possible. Hence, its greater stability wrt Fe 2+ 3+ ion 4s 3d

Electron configuration exceptions and their ions Cr Atom [Ar] 4s 1 3d 5 Ion [Ar] 4s 0 3d 5 [Ar] 4s 0 3d 3 [Ar] 4s 0 3d 0 Cr + Cr 3+ Cr 6+ Cu [Ar] 4s 1 3d 10 [Ar] 4s 0 3d 10 [Ar] 4s 0 3d 9 Au [Xe] 6s 1 4f 14 5d 10 [Xe] 6s 0 4f 14 5d 10 [Xe] 6s 0 4f 14 5d 8 Cu + Cu 2+ Au + Au 3+

Characteristic Properties of Metals and Nonmetals Metals • Have a shiny luster and are usually silvery in color  Notable exceptions: Metalloids

Cu Au

• Solids are malleable and ductile

Hg is only liquid metal at RT

Nonmetals • Do not have luster; various colors • Solids are usually brittle • Poor conductors of heat and electricity • Good conductors of heat and electricity • Metal oxides are ionic solids and form basic solutions • Form cations in solution • Most nonmetal oxides are molecular solids that form acidic solutions.

• Tend to form anions or oxyanions in solution

Elements with Color Pale yellow gas Light green gas Dark orange liquid/gas Dark violet crystals/gas

Properties of SOLID ionic compounds (M + NM) • Poor conductors of electricity and heat. • Generally high melting (more than 150°C). • Crystalline, hard and brittle. NaCl crystal lattice structure

Properties of ionic compounds (cont.) • Molten ionic compounds form liquids that are an electrical conductors. • Ionic solids that are water soluble, dissolve to form solutions that are electrical conductors.

• The solubility of ionic compounds depends upon the lattice energy. The greater the lattice energy, the lower the solubility. + -

Acid-Base Behavior of Oxides • Most metal oxides form basic solutions • Most nonmetal oxides form acidic solutions • The acidity of the solution increases with oxidation number of the central atom pH: SO 3 (aq) SO 2 (aq) Hint: an easy way to evaluate the acidity is that pH  as the # oxygen atoms  • Amphoteric substances can act as either an acid or a base

Group Trends Alkali metals: low IE  highly reactive, soft silvery metals with a low density and low melting point.

While lithium reacts with oxygen to form lithium oxide 4 Li + O 2  2 Li 2 O the other alkali metals form peroxides. (peroxide = O 2 2 ) 2 Na + O 2  Na 2 O 2 Potassium, rubidium and cesium react with oxygen to form superoxides (superoxide = O 2  ) K + O 2  KO 2

Alkaline earth metals • Harder, more dense and have a higher melting point than Group 1 metals.

• 1 st IE is low, but not as low as Group 1 because disrupting pseudo-noble gas configuration (s 2 ) • Increasing reactivity with increasing atomic number due to increase in nuclear shielding.

• Beryllium will not react with water, but the other alkaline earth metals will to form the metal hydroxide and hydrogen gas.

Ca + 2 H 2 O  Ca(OH) 2 + H 2