III. Compound theory

Key focus of this chapter: naming chemical compounds

This chapter reviews compound naming and the key principles of bonding, molecular geometry, intermolecular forces, magnetism, and crystal structures.

A. Naming chemical compounds

1. Naming oxoanions (atom + oxygen(s))

The hypo-/per- and -ite/-ate naming patterns apply to related oxoanions of the same central element, such as the chlorine oxyanion series. The number of oxygen atoms alone does not universally determine the name of an oxoanion.

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

2. Naming acids

a. Oxoacids (hydrogen + oxoanion)

The corresponding hypo-/per- and -ous/-ic acid patterns apply within an oxoacid series of the same central element, such as the chlorine oxoacids shown below. Oxygen count alone does not universally determine an oxoacid name.

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

The acid formulas correspond to complete protonation of the listed anions. H₂O₂ is hydrogen peroxide, not an ordinary oxoacid named by changing -ate to -ic or -ite to -ous.

Poly atomic anions		Acids
Single 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
Double 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
Triple 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)

Electronegativity is the tendency of an atom in a chemical bond to attract the shared electrons toward itself.
Atoms with high electronegativity strongly attract bonding electrons. Electronegativity is related to, but distinct from, electron affinity.
A greater electronegativity difference generally produces a more polar bond. Molecular polarity also depends on molecular geometry.
In the same period, electronegativity generally increases as the atomic number increases.
In the same group, electronegativity generally decreases as the atomic number increases.

Fig. 2 Electronegativity in the periodic table

2. Polar and nonpolar covalent bond

Classification	Features
Polar covalent bond	• Bond polarity depends primarily on the electronegativity difference between bonded atoms. • Molecular polarity depends on both bond dipoles and molecular geometry. Polar bonds do not necessarily produce a polar molecule.
Nonpolar covalent bond	• Identical atoms have zero electronegativity difference: H₂, F₂, and Cl₂ have nonpolar covalent bonds. • C–H bonds are often treated as nearly nonpolar.

Ex/ The order of polarity between two atoms:

HF has a polar covalent bond, C–H bonds are often treated as nearly nonpolar, and F₂ has a nonpolar covalent bond. NaF and LiF are ionic compounds. Bond polarity depends primarily on electronegativity difference; molecular polarity depends on both bond dipoles and molecular geometry. Polar bonds do not necessarily produce a polar molecule.

3. Dipole moments of polar and nonpolar molecules

Classification	Features
Polar molecules
Nonpolar molecules

4. Molecular forces (van der Waals force)

: attractive forces between molecules

Classification	Features
London dispersion force	London dispersion forces arise from instantaneous and induced dipoles and occur in all atoms and molecules. Strength depends on polarizability and molecular contact area; it often increases with size within related series. Boiling points reflect all intermolecular forces.
Dipole-dipole force	Dipole–dipole interactions occur between permanent molecular dipoles. Strength depends on dipole magnitude, orientation and separation. Comparisons of boiling points must also account for dispersion and other interactions.
Hydrogen bond	Hydrogen bonding occurs when H bonded to N, O or F interacts with a lone pair on a neighboring atom. Its strength depends on geometry and environment; there is no universal N–H / O–H / F–H strength ranking. Examples: H₂O, HF, NH₃.
Ion-dipole

For molecules of comparable size and polarizability, dipole–dipole attractions and hydrogen bonding can strengthen intermolecular attraction beyond dispersion alone. The familiar London dispersion < dipole–dipole < hydrogen-bonding comparison is a useful guideline, not a universal ranking of substances.

Boiling-point examples: CH₄ < SiH₄ < GeH₄ < SnH₄: increasing size and polarizability strengthen dispersion. H₂S < H₂Se < H₂Te: increasing dispersion dominates. All are polar; this is not a ranking of dipole–dipole interactions alone. HF < H₂O: combined forces and hydrogen-bond network connectivity determine boiling points; this is not a universal hydrogen-bond strength order.

Boiling points depend on the combined effects of intermolecular forces, including London dispersion forces, dipole–dipole interactions, and hydrogen bonding when applicable. Molecular size, polarizability, and molecular structure also influence these properties. Melting points additionally depend strongly on crystal packing.

In HCl, HBr, and HI, permanent dipole moments decrease in the order HCl > HBr > HI, while polarizability and dispersion contributions increase in the order HCl < HBr < HI. Their boiling points increase HCl < HBr < HI because the increasing dispersion contribution outweighs the decreasing dipole contribution. Their melting points follow the same order in this series; this does not establish a universal rule for melting points.

** Other covalent bonds

Diatomic molecules
Coordinate covalent
Covalent network solid

5. Octet rule

: atoms in the main group tend to be positioned with 8 electrons in their outmost shell like noble gases.

a. Valence electron:

Electrons of atoms in the outermost shell
Forming covalent bonds
Main factor to distinguish the chemical properties of atoms
Ex/

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

b. Lewis structure

Repulsive and attractive forces make optimum distance between two atoms

– repulsive forces between two different atom’s electrons

– attractive forces between one atom’s electrons and another atom’s nucleus

Indicating covalent bonds

– 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 atoms (Be, B, Al) have less than 8 electrons in their outermost shell.
Some atoms (past the second row in the periodic table: Si, P, Ge, Sn,…), have more than 8 electrons in their outermost shell.

Fig. 3 atoms that are exceptions to the octet rule

Some molecules involving third-period or heavier central atoms are represented by expanded-octet Lewis structures. However, significant participation of low-energy d orbitals is not required to explain their bonding. Modern bonding descriptions use delocalized molecular orbitals and other models.

XeF₄ has 36 valence electrons: four Xe–F bonds, two lone pairs on Xe, and three lone pairs on each F. All formal charges are zero. The electron-domain geometry is octahedral and the molecular geometry is square planar; the sketch below omits the terminal F lone pairs.

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

: electrons will be positioned as far as possible from one another to maintain a stable state of molecular geometry by the repulsive forces.

Count electron domains around the central atom: every single, double or triple bond counts as one bonding domain, and each lone pair counts as one domain. The table distinguishes molecular geometry from electron-domain geometry; ideal angles may be distorted by lone pairs and the identities of surrounding atoms.

Molecular shapes	Molecular geometry	Examples	Features
Linear
Linear
Trigonal planar
Bent			Bond angle <120° (SO₂: about 119°) 2 bonding domains 1 lone pair Example: SO₂
Bent			Bond angle <109.5° 2 bonding domains 2 lone pairs H₂O: 104.5°; H₂S and H₂Te are closer to 90°.
Tetrahedral			Ideal bond angle: 109.5° 4 bonding domains 0 lone pairs Examples: SiH₄, CH₄, NH₄⁺, carbon in diamond
Trigonal pyramidal
Trigonal bipyramidal
Seesaw			Distorted trigonal bipyramidal geometry Angles usually <90° and <120° 4 bonding domains 1 equatorial lone pair Examples: SF₄, IF₄⁺
T-shaped			Angles approximately 90° and 180° (often compressed) 3 bonding domains 2 equatorial lone pairs Example: ClF₃
Octahedral			Ideal angles: 90° and 180° 6 bonding domains 0 lone pairs Example: SF₆
Square pyramidal
Square planar

7. Formal charges

Formal charge is the charge assigned to an atom when bonding electrons are shared equally.

Formal charge = valence electrons − nonbonding electrons − ½(bonding electrons).

In the examples below, “binding number” means bond-order sum, equal to half the number of bonding electrons.

Atoms	Number of valence electrons	Binding number	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

Hybridization is a bonding model in which atomic orbitals on the same atom are mathematically combined to form hybrid orbitals with specific spatial orientations.

a. Sigma(σ) bond

Axis bond by head-on orbital overlap
Straight orbital overlap
Round shape by vertical section
Occupying maximum 2 electrons

b. Pi(π) bond

Sideways overlap
Dumbbell-shape by vertical section
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: each single, double, or triple bond counts as one domain, and each lone pair counts as one. The traditional sp³d and sp³d² labels describe five- and six-domain geometries; they do not imply that substantial d-orbital hybridization is required for hypervalent bonding.

Classification	Total number of bonds & lone pairs	Shape of molecules	Examples
		• Linear
		• Bent • Trigonal planar
		• Bent • Trigonal pyramidal • Tetrahedral
		• Linear • Seesaw • T-Shaped • Trigonal bipyramidal	Xe in XeF₂ S in SF₄ Cl in ClF₃ P in PCl₅
		• Octahedral • Square Pyramidal • Square Planar	S in SF₆ Sb in SbCl₅²⁻ Xe in XeF₄

BF₃: 3 bonds and 0 lone pair → 3+0=3 → sp²types of orbital

NH₃: 3 bonds and 1 lone pair → 3+1=4 → sp³ types of orbital

9. Resonance

Resonance occurs when a single Lewis structure cannot adequately represent the electron distribution in a molecule or ion. Resonance contributors differ in electron placement, not in the positions of atomic nuclei. The actual structure is a resonance hybrid with delocalized electrons.

Electron delocalization can stabilize a molecule or ion. The importance and relative energies of valid contributors matter; merely drawing more structures does not guarantee greater stabilization.
Contributors preserve total charge, electron count, and nuclear positions; major contributors generally favor complete octets and reasonable formal charges.
Examples: O₃, OCN⁻, and NO₃⁻. Ozone has two equivalent major contributors, nitrate has three, and cyanate has nonequivalent contributors with different weights.

D. Magnetism

Magnetism	Features
Ferromagnetism	Ferromagnetism: cooperative alignment of magnetic moments produces spontaneous magnetization within domains. An applied field can align domains, giving a strong positive response. Magnetization can remain after the field is removed. Examples: Fe, Co, Ni.
Paramagnetism	Paramagnetism: unpaired electrons give a weak attraction to an applied magnetic field (positive susceptibility), without ferromagnetic long-range alignment. Induced magnetization disappears when the field is removed. Examples: O₂, Al.
Diamagnetism	Diamagnetism: an applied field induces a weak opposing response (negative susceptibility). Diamagnetic contributions occur in all materials; substances with all electrons paired are usually diamagnetic. Examples: N₂, F₂, Cu.

E. Unit cells

A unit cell is a repeating region that generates a crystal structure by translation in three dimensions. A crystal structure consists of a lattice and its associated atomic basis.

Diamond has a diamond cubic crystal structure consisting of a face-centered cubic Bravais lattice with a two-atom basis. Its conventional cubic unit cell contains eight carbon 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 atom with 1/8 of each corner atom • Ex/ Copper, silver, gold, nickel, lead