Periodicity of Atomic PropertiesElements in the same group have the same number of

valence electrons and related electron configurations; hence have similar chemical properties.

The ground state electron configuration of the elements vary periodically with atomic number; all properties that depend on electron configuration tend to vary periodically with atomic number

Valence electrons - involved in bonding, determine reactivity

Bi (Atomic number 83): [Xe] 4f14 5d10 6s2 6p3

6 valence electrons

Os (Atomic number 76): [Xe] 4f145d6 6s2

8 valence electrons

The variation of effective nuclear charge through the periodic table plays an important role in determining periodic trends

Zeff increases from left to right across a period, but drops every time an outer electron occupies a new shell.

Atomic RadiusElectron clouds do not have sharp boundaries, so do not

really define the radius of an atom.

Atomic radius - half the distance between the nuclei of neighboring atoms

In solids centers of atoms are found at definite distances from one another.

For metals: atomic radius is half the distance between the centers of the atoms in the solid.

For nonmetals the atomic radius is half the distance between the nuclei of atoms joined by a chemical bond; the covalent radius.

van der Waals radius (used for noble gases): half the distance between the centers of the atoms in the solidified gas

Trend: Atomic radius generally decreases from left to right across a period and increases down a group

Across a period: new electron added to the same shell, and nuclear charge is increasing

Ionic radius: effective radius of an ion in an ionic solid

Cations are smaller than the neutral atom; anions are larger

Ionic radii generally increase down a group and decrease from left to right across a period.

Ionization Energies

Ionization energy: energy needed to remove an electron from an atom in the gas phase.

First ionization energy (I1): energy required to remove an electron from the neutral gas phase atom

Cu(g) -> Cu+(g) + e- energy = I1

For Cu I1 = 785 kJ/mol

Second ionization energy (I2): energy required to remove an electron from the singly charged gas phase atom

Cu+(g) -> Cu2+ (g) + e- energy = I2

For Cu I2 = 1955 kJ/mol

Trend: Ionization energy generally decreases down a group, and tends to increase moving across a period from left to right

Departures from these trends can usually be traced to repulsion between electrons, particularly electrons occupying the same orbitals

Elements with low ionization energies can be expected to form cations readily and to conduct electricity in their solid form.

Elements in the lower left of the periodic table tend to have lower ionization energies than those in the upper right.

These are the elements in the s block, d block, f block and the lower left of the p block - metallic solids

Electron AffinityElectron attachment energy: energy change when a gas

phase atom in its ground state gains a single electron

X(g) + e- -> X- (g) E = electron attachment energy

Electron affinity = - E(electron attachment)

Electron affinity higher in the upper right

Main Group Elements (s and p)An s-block element has low ionization energy; outer electrons can easily be lost

Group 1 form +1 ions; Group 2 form +2 ions

An s block element is likely to be a reactive metal

Since ionization energies are lowest at the bottom of each group , these elements are most reactive in the s-block

p- block

Elements on the left have low enough ionization energies to be metallic, but higher than the s-block elements and so are less reactive

Elements on the right have high electron affinities (tend to gain electrons to form closed shell ions)

Transition Metals (d block)

All d-block elements are metals; properties are transitional to s and p block

Many d-block elements exhibit more than one oxidation state since the d-electrons have similar energies (inner shell) and a variable number can be lost

Chemical Bonds

Chemical Bond: arise from sharing or transfer of electrons between two or more atoms

Covalent bond: when electron density is mainly shared between two atoms

Ionic bond: electron density mainly transferred from one atom to another.

Pot

entia

l ene

rgy

(kJ/

mol

)

Electronegativity: ability of an atom to attract electrons shared in a bond toward itself

Elements to the lower left have low ionization energies and low electron affinities; tend to act as electron donors

Elements to the upper right have high ionization energies and high electron affinities; tend to act as electron acceptors

R. Mulliken

= 0.5 (IE1 + EA)

L. Pauling: quantitative measure of electronegativity, | A - B| = 0.102{D(A-B) - 0.5 [ D(A-A) + D(B-B)]}0.5

A polar colavent bond is a bond with partial electric charges arising from their difference in electronegativity.

The partial charges give rise to an electric dipole moment

Pauling Scale for

Bond dissociation energy Ed : energy required to separate the bonded atoms

H-Cl(g) -> H(g) + Cl(g)

Greater the Ed, stronger the bond

Bond enthalpy: enthalpy change (heat absorbed at constant pressure) when bond is broken

Ed = Hd - R T

Since bond enthalpy of a bond X-Y between two atoms X and Y does not vary much between compounds, define an “average” bond enthalpy or “average” bond energy

Variations in atomic nuclei contribute to trends in bond energy

Number of bonds between two atoms also influence bond energy

single < double < triple

Bond length: distance between the centers of two atoms joined by a covalent bond

Bonds between larger atoms tend to be longer

Multiple bonds are shorter than single bonds

A bond length is approximately the sum of the covalent radii of the two atoms

Covalent radius: effective radius of an atom in a covalent bond

Bond Order: number of shared electron pairs

Single bond: bond order = 1

Double bond: bond order = 2

Triple bond: bond order = 3

Bond length decrease with bond order

Bond strength increases with bond order

Chemical Bonding

The Lewis model of the chemical bond assumes that each bonding electron is located between two bonded atoms - localized electron model.

However, location of an electron in an atom cannot be described in terms of a precise position; same is true of electrons in a molecule.

Valence Bond Theory - first quantum mechanical description of distribution of electrons in bonds.

Valence Bond TheoryMinimum requirement for the formation of a covalent bond is

two electrons and two valence orbitals that can overlap.

The overlapping of the two orbitals concentrate the electron probability between the two nuclei, creating a chemical bond.

H2: two H atoms, each in the ground state 1s1

In VB theory: the two H atoms come together and their atomic orbitals merge or overlap

Greater the extent of overlap, greater the strength of the bond

Overlap of the two 1s orbitals form a sigma () bond.

In a sigma bond - electron density is distributed along the bond axis

Note that the electron spin is paired

HF H: 1s1 F: 1s22s22p5

Overlap between the valence orbital of H (1s) and valence orbital of F (2p) to form a bonds

Note: electron spin is paired in the orbital

By definition: z is the direction along the internuclear axis