Ionic and Molecular Compounds: Structure, Properties, and Naming

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Chapter 6: Ionic & Molecular Compounds

Metals vs Nonmetals

The periodic table is divided into metals, metalloids, and nonmetals, each with distinct chemical behaviors. Metals tend to lose electrons and form cations, while nonmetals gain electrons to form anions. The interaction between these types of elements forms the basis for ionic and molecular compounds.

Metals: Located on the left and center of the periodic table; typically lose electrons.
Nonmetals: Located on the right; typically gain electrons.
Metalloids: Elements with properties intermediate between metals and nonmetals.

Periodic Table highlighting metals, metalloids, and nonmetals

Ionic vs Covalent Bonding

Compounds are formed by either transferring or sharing electrons. Ionic bonds occur between metals and nonmetals, while covalent bonds occur between nonmetals.

Ionic Bond: Formed by transfer of electrons from a metal to a nonmetal, resulting in cations and anions.
Covalent Bond: Formed by sharing electrons between nonmetals.

Diagram showing transfer and sharing of electrons for ionic and covalent bonds

Diatomic Molecular Compounds

Certain elements exist naturally as diatomic molecules, meaning two atoms of the same element are bonded together. These molecules have an oxidation state of zero.

Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2
Oxidation State: For diatomic molecules, the charge is zero.

Periodic table highlighting diatomic elements

Diatomic Molecule	Name
H2	Hydrogen
N2	Nitrogen
O2	Oxygen
F2	Fluorine
Cl2	Chlorine
Br2	Bromine
I2	Iodine

Table of diatomic molecules and their names

Formation of Ionic and Molecular Compounds

Ionic compounds are composed of ions, while molecular compounds are composed of molecules. The type of bond and the nature of the particles determine their properties.

Ionic Compounds: Made of ions, held together by ionic bonds (e.g., NaCl).
Molecular Compounds: Made of molecules, held together by covalent bonds (e.g., H2O, C2H6).

Type	Ionic Compounds	Molecular Compounds
Particles	Ions	Molecules
Bonds	Ionic	Covalent
Examples	Na+ Cl- ions	H2O, C2H6 molecules

Comparison of ionic and molecular compounds

Cations and Anions

Atoms form ions by losing or gaining electrons. Metals form cations (positive), and nonmetals form anions (negative).

Cation: Positively charged ion formed by loss of electrons (e.g., Na+).
Anion: Negatively charged ion formed by gain of electrons (e.g., Cl-).

Table showing common cations and anions

Formation of Sodium and Chloride Ions

Sodium (Na) loses one electron to become Na+, while chlorine (Cl) gains one electron to become Cl-. This transfer results in the formation of sodium chloride (NaCl).

Electron Arrangement: Sodium changes from 2,8,1 to 2,8; Chlorine changes from 2,8,7 to 2,8,8.
Charge Balance: The sum of charges in NaCl is zero.

Formation of sodium ion Formation of chloride ion Diagram showing electron transfer in NaCl formation

Naming Ionic Compounds

Ionic compounds are named by combining the names of the cation and anion. For main group elements, the charge is determined by the group number. Transition metals may have variable charges, indicated by Roman numerals.

Main Group Elements: Charge is consistent within a group.
Transition Metals: Charge varies; oxidation state must be specified.

Compound	Metal Ion	Nonmetal Ion	Name
KI	K+	I-	Potassium iodide
MgBr2	Mg2+	Br-	Magnesium bromide
Al2O3	Al3+	O2-	Aluminum oxide

Periodic table showing charges for main group elements

Variable Charge (Transition Metals)

Transition metals and some other elements (Sn, Pb, Bi) can form ions with different charges. The oxidation state is specified in the compound name.

Element	Possible Ions	Name of Ion
Fe	Fe2+, Fe3+	Iron(II), Iron(III)
Cu	Cu+, Cu2+	Copper(I), Copper(II)
Sn	Sn2+, Sn4+	Tin(II), Tin(IV)
Pb	Pb2+, Pb4+	Lead(II), Lead(IV)

Table of transition metal ions

Compound	Systematic Name
FeCl2	Iron(II) chloride
Fe2O3	Iron(III) oxide
Cu3P	Copper(I) phosphide
CrBr2	Chromium(II) bromide
SnCl2	Tin(II) chloride
PbS2	Lead(IV) sulfide
BiF3	Bismuth(III) fluoride

Table of transition metal compound names

Common Polyatomic Ions

Polyatomic ions are groups of atoms bonded together with an overall charge. They are important in both biological and chemical contexts.

Examples: Hydroxide (OH-), Ammonium (NH4+), Nitrate (NO3-), Carbonate (CO32-), Bicarbonate (HCO3-), Sulfate (SO42-), Phosphate (PO43-).

Nonmetal	Formula of Ion	Name of Ion
Hydrogen	OH-	Hydroxide
Nitrogen	NH4+, NO3-, NO2-	Ammonium, Nitrate, Nitrite
Chlorine	ClO4-, ClO3-, ClO2-, ClO-	Perchlorate, Chlorate, Chlorite, Hypochlorite
Carbon	CO32-, HCO3-, CN-, C2H3O2-	Carbonate, Hydrogen carbonate, Cyanide, Acetate
Sulfur	SO42-, HSO4-, SO32-, HSO3-	Sulfate, Hydrogen sulfate, Sulfite, Hydrogen sulfite
Phosphorus	PO43-, HPO42-, H2PO4-, PO33-	Phosphate, Hydrogen phosphate, Dihydrogen phosphate, Phosphite

Table of common polyatomic ions

Common Polyatomic Compounds and Their Uses

Many polyatomic compounds have important medical and industrial uses.

Formula	Name	Medical Use
AlPO4	Aluminum phosphate	Antacid
Al2(SO4)3	Aluminum sulfate	Antiperspirant, anti-infective
BaSO4	Barium sulfate	Contrast medium for X-rays
CaCO3	Calcium carbonate	Antacid, calcium supplement
Ca3(PO4)2	Calcium phosphate	Calcium dietary supplement
CaSO4	Calcium sulfate	Plaster casts
MgSO4	Magnesium sulfate	Cathartic, Epsom salts
K2CO3	Potassium carbonate	Alkalizer, diuretic
AgNO3	Silver nitrate	Topical anti-infective
NaHCO3	Sodium bicarbonate	Antacid
Zn3(PO4)2	Zinc phosphate	Dental cement

Table of polyatomic compounds and their medical uses

Lewis Structures and the Octet Rule

Lewis structures are used to represent the arrangement of electrons in molecules. The octet rule states that atoms tend to have eight electrons in their valence shell, except for certain exceptions.

Octet Rule: Most atoms achieve stability with eight valence electrons.
Exceptions: H2 (2 electrons), BCl3 (6 electrons), PCl5 (10 electrons), SF6 (12 electrons), ClO4 (14 electrons).

Lewis structure for F2 showing shared and lone pairs

Naming Molecular Compounds

Molecular compounds are named using prefixes to indicate the number of atoms. If the first element is 'mono', the prefix is omitted.

Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.
Example: CO2 is carbon dioxide, not monocarbon dioxide.

Atomic Size and Electronegativity Trends

Atomic size and electronegativity are periodic trends that influence chemical bonding.

Atomic Size: Increases down a group, decreases across a period.
Electronegativity: Increases across a period, decreases down a group.

Bond Polarity and Dipoles

The difference in electronegativity between atoms determines bond polarity. If the difference is 0.5 or greater, the bond is polar; otherwise, it is nonpolar.

Bond Polarity: Polar bonds have unequal sharing of electrons, creating dipoles.
Dipole Moment: A measure of molecular polarity.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory explains the geometry and polarity of covalent compounds. Electron groups (lone pairs or bonded atoms) arrange themselves to minimize repulsion.

Electron Groups: Lone pairs and bonded atoms.
Molecular Shape: Determined by the arrangement of electron groups.
Polarity: Determined by the sum of bond dipoles; equal and opposite dipoles cancel.

Molecular Shapes

Molecules can have various shapes depending on the number of electron groups and lone pairs.

Linear: 2 electron groups, 180° bond angle.
Trigonal Planar: 3 electron groups, 120° bond angle.
Bent: 3 electron groups with a lone pair.
Tetrahedral: 4 electron groups, 109.5° bond angle.
Trigonal Pyramidal: 4 electron groups with a lone pair.

Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties like melting and boiling points.

Dipole-Dipole: Occurs in polar molecules.
Hydrogen Bonding: Extreme dipole-dipole, only for N, O, and F.
Dispersion (London) Forces: Weakest, present in all molecules, strength increases with size.

Melting and Boiling Points

The strength of intermolecular forces determines melting and boiling points.

Ionic Forces: Strongest, highest melting and boiling points.
Dipole-Dipole: Intermediate strength.
Dispersion Forces: Weakest, lowest melting and boiling points.

Example: Mg2+ has stronger attraction than Na+; O-H hydrogen bond is stronger than N-H due to greater electronegativity difference; bromine has stronger forces than chlorine due to larger atomic radius.