III. Compound theory

Key focus of this chapter: naming chemical compounds

This chapter reviews chemical nomenclature, bonding, electronegativity, molecular polarity, intermolecular forces, Lewis structures, VSEPR geometry, formal charge, and hybridization.

A. Naming chemical compounds

1. Naming oxyanions (an element bonded to oxygen)

The bromine series below illustrates the naming pattern. Oxygen counts are specific to this series; other elements can have different counts.

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ite
Two		- ite
Three		- ate
Four	Per -	- ate

2. Naming acids

a. Oxyacids (hydrogen-containing acids based on oxyanions)

The chlorine series below illustrates the pattern: -ite becomes -ous acid, and -ate becomes -ic acid; hypo- and per- are retained.

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ous acid
Two		- ous acid
Three		- ic acid
Four	Per -	- ic acid

** H₂SO₃ : Sulfurous acid

H₂SO₄ : Sulfuric acid

H₃PO₄ : Phosphoric acid

b. Hydrogen + anion

Prefix	Suffix	Examples
Hydro -	- ic acid	HCl(aq): hydrochloric acid HBr(aq): hydrobromic acid HI(aq): hydroiodic acid HCN(aq): hydrocyanic acid

3. Table of naming polyatomic anions and acids

Polyatomic anions		Acids
Singly charged anions	Hydrogen carbonate (bicarbonate)		Carbonic acid
	Hydrogen sulfate (bisulfate)		Sulfuric acid
	Dihydrogen phosphate		Phosphoric acid
	Acetate		Acetic acid
	Cyanide		Hydrocyanic acid
	Nitrite		Nitrous acid
	Nitrate		Nitric acid
	Hydroxide
	Hypochlorite		Hypochlorous acid
	Chlorite		Chlorous acid
	Chlorate		Chloric acid
	Perchlorate		Perchloric acid
	Permanganate		Permanganic acid
	Hypobromite		Hypobromous acid
	Bromite		Bromous acid
	Bromate		Bromic acid
	Perbromate		Perbromic acid
Doubly charged anions	Hydrogen phosphate		Phosphoric acid
	Carbonate		Carbonic acid
	Chromate		Chromic acid
	Dichromate		Dichromic acid
	Peroxide		Hydrogen peroxide
	Sulfite		Sulfurous acid
	Sulfate		Sulfuric acid
	Thiosulfate		Thiosulfuric acid
Triply charged anions	Phosphate		Phosphoric acid

** NH₄⁺: Ammonium

4. Naming compounds

Fig. 1 Classification of groups

Combinations	Naming compounds

B. Chemical Bonds

Chemical Bonds	Features
Ionic bonds
Metallic bonds
Covalent bonds (molecules)

C. Covalent bond

1. Electronegativity (EN)

The ability of an atom in a bond to attract the shared electrons toward itself.
An atom with higher electronegativity attracts shared electrons more strongly. Electronegativity describes a bonded atom and is distinct from electron affinity.
A greater electronegativity difference generally means a more polar bond; the polarity of the whole molecule also depends on geometry.
In the same period, electronegativity generally increases as atomic number increases.
In the same group, electronegativity generally decreases as atomic number increases.

Fig. 2 Electronegativity in the periodic table

2. Polar and nonpolar covalent bond

Classification	Features
Polar covalent bond	• A larger electronegativity difference generally produces a more polar bond. Molecular polarity also depends on molecular geometry. • A very large electronegativity difference → greater ionic character → bonding is better described as ionic rather than covalent
Nonpolar covalent bond	• The difference of electronegativity between diatomic elements is zero (H2 ,F2, Cl2) • Usually C-H bond is considered to be nonpolar.

Example: Comparing bond polarity using electronegativity differences:

Therefore, ionic bonds (NaF, LiF) have a higher polarity than covalent bonds (H–F, C–H, F₂)

3. Dipole moments of polar and nonpolar molecules

Classification	Features
Polar molecules
Nonpolar molecules

4. Intermolecular forces

: attractive interactions between particles, including molecules and ions

Classification	Features
London dispersion force
Dipole-dipole force
Hydrogen bond
Ion-dipole

** A common qualitative comparison for molecules of similar size and polarizability (not a universal ranking)

London dispersion force < Dipole – dipole force < Hydrogen bond

** For HCl, HBr, and HI, increasing polarizability strengthens London dispersion forces enough to outweigh the decrease in dipole–dipole attraction. Boiling points follow this trend; melting points also depend strongly on crystal packing.

– the strength order of dipole-dipole force : HI < HBr < HCl

– the strength order of London dispersion force: HCl < HBr < HI

– the order of boiling points: HCl < HBr < HI

** Other covalent bonds

Diatomic molecules
Coordinate covalent
Covalent network solid

5. Octet rule

: many main-group atoms form bonds that give them eight valence electrons, as in a noble-gas configuration. Hydrogen usually follows a duet rule.

a. Valence electron:

For main-group atoms, electrons in the outermost occupied shell
Participate in bonding, including covalent and ionic bonding
Largely determine an element’s chemical behavior
The table shows typical bond-order totals for neutral, closed-shell molecules.

Atom	Number of valence electrons	Forming bonds
C N O F	4 5 6 7	4 3 2 1

b. Lewis structure

The equilibrium bond length balances attractive and repulsive interactions

– repulsion between electrons and between nuclei

– attraction between each nucleus and the electrons

Multiple covalent bonds share electron pairs

– double bonds by sharing 4 electrons (2 pairs): CO₂, O₂

– triple bonds by sharing 6 electrons (3 pairs): HCN, N₂

c. Exception to octet rule

Some compounds of Be, B, and Al have electron-deficient central atoms with fewer than eight electrons. Odd-electron species, such as NO and NO₂, cannot give every atom an octet.
Some central atoms in period 3 or later (such as Si, P, Ge, and Sn) can be represented by expanded-octet Lewis structures.

Fig. 3 Atoms that are exceptions to the octet rule

Expanded-octet Lewis structures assign more than eight electrons to some central atoms in period 3 or later. This electron-counting model should not be interpreted as simple filling of low-energy d orbitals; modern bonding descriptions include delocalized molecular orbitals.

6. VSEPR model (valence-shell electron-pair repulsion model)

: electron domains around a central atom arrange themselves to minimize repulsion. Each single, double, or triple bond counts as one domain; each lone pair also counts as one domain. Molecular geometry describes atom positions, while electron-domain geometry includes lone pairs. Diagrammed angles are idealized or approximate; the ideal tetrahedral angle is 109.5°, and lone pairs often reduce adjacent bond angles.

Molecular shapes	Molecular geometry	Examples	Features
Linear
Linear
Trigonal planar
Bent
Bent
Tetrahedral
Trigonal pyramidal
Trigonal bipyramidal
Seesaw
T-shaped
Octahedral
Square pyramidal
Square planar

7. Formal charges

: bookkeeping charges assigned to atoms in a Lewis structure by treating bonding electrons as equally shared. Formal charge is not the same as an atom’s actual partial charge.
Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons). The bond-order sum in the table counts a single bond as 1, a double bond as 2, and a triple bond as 3. Formal charges sum to the total charge of the molecule or ion.

Atoms	Number of valence electrons	Bond-order sum	Number of non-bonding electrons
N	5	3	2
C	4	4	0
S	6	1	6

Formal charge of N: 5 −3 − 2 = 0

Formal charge of C: 4 − 4 − 0 = 0

Formal charge of S: 6 − 1 − 6 = -1

8. Hybridization

: the mathematical mixing (linear combination) of atomic orbitals on the same atom to form hybrid orbitals.

a. Sigma(σ) bond

Head-on orbital overlap along the internuclear axis
Straight orbital overlap
Electron density is cylindrically symmetric around the bond axis
Each sigma bonding orbital can hold up to two electrons.

b. Pi(π) bond

Sideways overlap
Electron density lies on opposite sides of a nodal plane containing the bond axis
Weaker than sigma(σ) bond

c. Application

Single bond: 1 sigma bond and 0 pi bond
Double bond: 1 sigma bond and 1 pi bond
Triple bond: 1 sigma bond and 2 pi bonds

d. Hybrid orbitals

Count electron domains as bonded atoms plus lone pairs, treating each multiple bond as one domain. The labels sp³d and sp³d² are traditional geometry models for five and six domains; they should not be taken as proof of substantial d-orbital participation in main-group bonding.

Classification	Electron domains (bonded atoms + lone pairs)	Shape of molecules	Examples
		• Linear
		• Bent • Trigonal planar
		• Bent • Trigonal pyramidal • Tetrahedral
		• Linear • Seesaw • T-Shaped • Trigonal bipyramidal
		• Octahedral • Square Pyramidal • Square Planar

Examples:

BF₃: 3 bonds and 0 lone pair → 3+0=3 → sp²hybridization

NH₃: 3 bonds and 1 lone pair → 3+1=4 → sp³ hybridization

9. Resonance

: two or more Lewis structures that differ only in electron placement are needed to describe a molecule or ion. The actual species is a resonance hybrid; it does not switch between the drawings.

Delocalization stabilizes the species; equivalent resonance contributors are especially important
Resonance contributors preserve atom positions, total charge, and the number of unpaired electrons
Examples: O₃, OCN^–, NO₃^–

D. Magnetism

For isolated atoms and molecules, unpaired electrons indicate paramagnetism; species with all electrons paired are diamagnetic. Ground-state Mg ([Ne]3s²) and Ca ([Ar]4s²) atoms have no unpaired electrons. The metallic examples in the table refer to bulk materials, whose magnetic response also depends on conduction electrons and cannot be inferred from isolated-atom configurations alone.

Magnetism	Features
Ferromagnetism
Paramagnetism
Diamagnetism

E. Unit cells

: a repeating three-dimensional portion of a crystal lattice. Repeating the unit cell in space reproduces the crystal structure; lattice points can represent an atom or a group of atoms.

Unit Cell	Number of atom(s) per unit cell	Features
Primitive cubic (Simple cubic)	1	• 1/8 of each atom in unit cell • Ex/ Polonium
Body-centered cubic	2	• 1 atom at center with 1/8 of each corner atom • Ex/ Tungsten, iron, sodium
Face-centered cubic	4	• 1/2 of each face-centered atoms with 1/8 of each corner atom • Ex/ Copper, silver, gold, nickel, lead