Chapter 3: Molecules, Compounds, and Chemical Equations – General Chemistry 1101 Study Notes

Chapter 3: Molecules, Compounds, and Chemical Equations

3.1 Hydrogen, Oxygen, and Water

This section introduces the concept that compounds possess physical and chemical properties distinct from their constituent elements. The comparison of hydrogen, oxygen, and water highlights these differences.

• Compounds have unique properties compared to individual elements.

• Physical properties such as boiling point and state at room temperature differ between hydrogen, oxygen, and water.

• Chemical properties such as flammability also vary significantly.

Selected Properties	Hydrogen	Oxygen	Water
Boiling Point	-253 °C	-183 °C	100 °C
State at Room Temperature	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish flame

3.1 Mixtures and Compounds

Mixtures and compounds differ in how their constituent elements combine.

• Mixtures: Elements can mix in any proportion (e.g., hydrogen H2 and oxygen O2).

• Compounds: Elements combine in fixed, whole-number ratios (e.g., water H2O, hydrogen peroxide H2O2).

• Water molecules always have a 2:1 ratio of hydrogen to oxygen.

3.2 Chemical Bonds

Atoms in compounds are held together by chemical bonds, which can be classified as ionic or covalent.

• Ionic bonds: Atoms held together by opposite charges (cation and anion).

• Covalent bonds: Atoms held together by shared electrons.

3.2 Ionic Bonds

• Formed by electron transfer from a metal to a nonmetal.

• Metal ion becomes a cation (positive charge).

• Nonmetal becomes an anion (negative charge).

• Electrostatic forces attract cations and anions, forming a charged bond.

3.2 Ionic Compounds in Solid Phase

• In the solid phase, ionic compounds form a lattice of alternating cations and anions.

• This structure maximizes attractive forces and minimizes repulsive forces.

3.2 Covalent Bonds

• Occur between two or more nonmetals that share electrons.

• Form discrete molecules called molecular compounds.

3.3 Chemical Formulas and Molecular Models

Chemical formulas indicate the elements present and the relative number of atoms in a compound.

• Water – H2O (covalent bond)

• Sodium Chloride – NaCl (ionic bond)

• Carbon Dioxide – CO2 (covalent bond)

• Carbon Tetrachloride – CCl4 (covalent bond)

Types of Chemical Formulas

• Empirical formula: Simplest whole-number ratio of atoms.

• Molecular formula: Actual number of atoms of each element in a molecule.

• Structural formula: Shows how atoms are connected and the geometry of the molecule.

Type	Example
Empirical formula	HO (for H2O2)
Molecular formula	H2O2
Structural formula	H–O–O–H

Molecular Models

• Ball-and-stick model: Atoms as balls, bonds as sticks; color-coded for elements.

• Space-filling model: Shows relative radii of atoms; more closely represents actual molecular shape.

Color	Element
White	Hydrogen
Black	Carbon
Blue	Nitrogen
Red	Oxygen
Green	Chlorine
Yellow	Sulfur
Purple	Phosphorus

3.4 Compound Representations

Compounds can be represented in various ways, including empirical, molecular, and structural formulas, as well as molecular models.

Name of Compound	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6
Acetylene	CH	C2H2	H–C≡C–H
Glucose	CH2O	C6H12O6
Ammonia	NH3	NH3

Additional info: Images referenced are for illustration; actual study notes would use drawn structures.

Atomic-Level View of Elements and Compounds

• Elements can be atomic (e.g., Ne) or molecular (e.g., O2).

• Compounds can be molecular (e.g., H2O) or ionic (e.g., NaCl).

Molecular Elements

• Some elements exist as molecules rather than single atoms (e.g., O2, N2).

• Polyatomic molecules include P4, S8.

Molecular Compounds

• Composed of two or more covalently bonded nonmetals.

• Molecules are the basic units (e.g., C3H8 for propane).

Ionic Compounds

• Composed of cations and anions bound by ionic bonds.

• Formula unit: Smallest electrically neutral collection of ions (e.g., NaCl).

Polyatomic Ions

• Groups of covalently bonded atoms with an overall charge.

• Examples: NO3- (nitrate), CO32- (carbonate), ClO3- (chlorate).

Classifying Substances

• Substances can be classified as atomic element, molecular element, molecular compound, or ionic compound.

• Examples:

  • Fluorine: Molecular element

  • N2O: Molecular compound

  • Silver: Atomic element

  • K2O: Ionic compound

  • Fe2O3: Ionic compound

3.5 Ionic Compounds: Formulas and Names

Ionic compounds consist of positive and negative ions. Their formulas reflect the smallest whole-number ratio of ions that results in electrical neutrality.

• Chemical formula:

  • Sum of the charges of the cations + sum of the charges of the anions = 0

Naming Ionic Compounds

• Ionic compounds are categorized by the type of metal present:

• Type I: Metal forms only one type of ion (invariant charge).

• Type II: Metal forms more than one type of ion (variable charge).

Naming Type I Ionic Compounds

• Type I metals: Group 1A, 2A, and some others (e.g., Al, Zn, Ag).

• Binary compounds: Name of cation (metal) + base name of anion (nonmetal) + -ide.

• Examples:

  • KCl: Potassium chloride

  • CaO: Calcium oxide

Nonmetal	Symbol for Ion	Base Name	Anion Name
Fluorine	F-	fluor	Fluoride
Chlorine	Cl-	chlor	Chloride
Bromine	Br-	brom	Bromide
Iodine	I-	iod	Iodide
Oxygen	O2-	ox	Oxide
Sulfur	S2-	sulf	Sulfide
Nitrogen	N3-	nitr	Nitride
Phosphorus	P3-	phosph	Phosphide

Naming Type II Ionic Compounds

• Type II metals: Transition metals and some main group metals (e.g., Fe, Cu, Pb, Sn).

• Name of cation (metal) + charge in Roman numerals in parentheses + base name of anion (nonmetal) + -ide.

• Example: CrBr3 is chromium(III) bromide.

Metal	Ion	Name
Chromium	Cr2+	Chromium(II)
Chromium	Cr3+	Chromium(III)
Iron	Fe2+	Iron(II)
Iron	Fe3+	Iron(III)
Cobalt	Co2+	Cobalt(II)
Cobalt	Co3+	Cobalt(III)
Copper	Cu+	Copper(I)
Copper	Cu2+	Copper(II)
Tin	Sn2+	Tin(II)
Sn	Sn4+	Tin(IV)
Mercury	Hg22+	Mercury(I)
Mercury	Hg2+	Mercury(II)
Lead	Pb2+	Lead(II)
Lead	Pb4+	Lead(IV)

Naming Ionic Compounds Containing Polyatomic Ions

• Use the name of the polyatomic ion directly in the compound name.

• Example: NaNO2 is sodium nitrite.

Name	Formula	Name	Formula
Acetate	C2H3O2-	Hypochlorite	ClO-
Carbonate	CO32-	Chlorite	ClO2-
Hydrogen carbonate	HCO3-	Chlorate	ClO3-
Hydroxide	OH-	Perchlorate	ClO4-
Nitrate	NO3-	Permanganate	MnO4-
Nitrite	NO2-	Sulfate	SO42-
Chromate	CrO42-	Hydrogen sulfate	HSO4-
Dichromate	Cr2O72-	Phosphate	PO43-
Hydrogen phosphate	HPO42-	Cyanide	CN-
Dihydrogen phosphate	H2PO4-	Peroxide	O22-
Ammonium	NH4+

3.6 Oxyanions

Oxyanions are anions containing oxygen and another element. Their names depend on the number of oxygen atoms present.

• If two ions in a series:

  • More oxygen: -ate ending

  • Fewer oxygen: -ite ending

• Examples:

  • NO2-: nitrite

  • NO3-: nitrate

  • SO32-: sulfite

  • SO42-: sulfate

• If more than two ions:

  • hypo-: less than

  • per-: more than

  • Example: ClO- (hypochlorite), ClO2- (chlorite), ClO3- (chlorate), ClO4- (perchlorate)

3.5 Hydrated Ionic Compounds

Hydrates are ionic compounds containing a specific number of water molecules per formula unit.

• Example: MgSO4 • 7H2O is magnesium sulfate heptahydrate.

• CoCl2 • 6H2O is cobalt(II) chloride hexahydrate.

Prefix	Number
hemi	1/2
mono	1
di	2
tri	3
tetra	4
penta	5
hexa	6
hepta	7
octa	8

3.6 Molecular Compounds: Formulas and Names

Molecular compounds are composed of two or more nonmetals. Their formulas cannot be determined solely from constituent elements.

• Name the element with the smallest group number first.

• If in the same group, name the element with the greatest row number first.

• Prefixes indicate the number of atoms present:

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

• If only one atom of the first element, the prefix mono- is usually omitted.