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

Chapter 3: Molecules, Compounds, and Chemical Equations

Elements and Compounds

Elements combine with each other to form compounds, resulting in the vast diversity of substances found in nature. When two or more elements combine, a new substance with unique properties is formed.

• Elements: Pure substances consisting of only one type of atom.

• Compounds: Substances composed of two or more different elements chemically bonded in fixed proportions.

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

Definite Proportions: Compounds have fixed ratios of elements. For example, in H2O, the ratio of H to O atoms is 2:1.

Chemical Bonds

Atoms in compounds are held together by chemical bonds, which arise from electromagnetic interactions between charged particles (electrons and nuclei).

• Ionic Bonds: Typically between metals and nonmetals; involve transfer of electrons.

• Covalent Bonds: Typically between nonmetals; involve sharing of electrons.

• Metallic Bonds: Between metal atoms; involve a 'sea' of delocalized electrons.

Electrostatic Interaction

The force between charged particles is described by Coulomb's Law:

$F = \frac{1}{4\pi\varepsilon} \frac{q_1 q_2}{r^2}$

$U = \frac{1}{4\pi\varepsilon} \frac{q_1 q_2}{r}$

• Proportional to the product of the charges.

• Inversely proportional to the square of the distance (force) or the distance (potential energy).

• Attraction occurs between opposite charges; repulsion between like charges.

Chemical Bond (Atomic Perspective)

Electrostatic (electromagnetic) interactions link atoms together to form compounds (ionic and covalent bonds) and hold molecules together in liquids or solids.

• Ionic Bond: Attraction between cations and anions (e.g., Na+ and Cl-).

• Covalent Bond: Shared electrons between atoms (e.g., H2O).

An Atomic-Level View of Elements and Compounds

Elements may be atomic or molecular; compounds may be molecular or ionic.

• All bonds in a molecule are covalent bonds.

• The bond between cations and anions in an ionic compound is an ionic bond.

• Polyatomic ions (e.g., NH4+, CO32-) have covalent bonds within the ion.

Molecular Elements

Certain elements exist naturally as molecules rather than individual atoms. These include diatomic molecules (e.g., H2, O2, N2, F2, Cl2, Br2, I2) and polyatomic molecules (e.g., P4, S8).

Representing Compounds: Chemical Formulas and Molecular Models

A chemical formula indicates the elements present in a compound and the number of atoms or ions of each.

• Water: H2O

• Carbon dioxide: CO2

• Sodium chloride: NaCl

• Calcium perchlorate: Ca(ClO4)2

Types of Chemical Formulas

• Empirical Formula: Shows the simplest whole-number ratio of atoms (used for ionic compounds).

• Molecular Formula: Shows the actual number of atoms of each element in a molecule (used for molecular compounds).

• Structural Formula: Shows how atoms are bonded together in a molecule.

Formula of Ionic Compounds

• Ionic compounds contain positive (cations) and negative (anions) ions.

• The sum of the charges in the formula must be zero (electrically neutral).

• The formula reflects the smallest whole-number ratio of ions.

Writing Formulas of Ionic Compounds

1. The charge on the cation becomes the subscript on the anion.

2. The charge on the anion becomes the subscript on the cation.

3. If necessary, reduce subscripts to the lowest whole-number ratio.

Example: For Mg2+ and N3-, the formula is Mg3N2.

Molecular Formula

The molecular formula gives the actual number of atoms of each element in a molecule. For example, hydrogen peroxide is H2O2 (empirical formula: HO).

Structural Formula

Structural formulas use lines to represent covalent bonds and show how atoms are connected. Isopropanol (C3H8O) and propanol (C3H8O) have the same molecular formula but different structures.

Ways of Representing a Compound

Name	Empirical Formula	Molecular Formula	Structural Formula	Ball-and-Stick Model	Space-Filling Model
Benzene	CH	C6H6	Ring structure	Ball-and-stick	Space-filling
Ethylene	CH2	C2H4	Double bond	Ball-and-stick	Space-filling

Additional info: Table entries inferred for clarity.

Formula Mass and Molar Mass of Compounds

• Formula Mass: Sum of the atomic masses of all atoms in a formula unit (amu).

• Molar Mass: Mass in grams of 1 mole of molecules or formula units (g/mol); numerically equal to formula mass.

Examples:

• NaCl: 23 + 35.5 = 58.5 amu; molar mass = 58.5 g/mol

• H2O2: 2×1 + 2×16 = 34 amu; molar mass = 34 g/mol

Composition of a Compound

To analyze the composition of a compound, use the molar mass to convert between mass and moles, and determine the mass percent of each element.

• Number of moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Mass percent: $\text{mass percent of element X} = \frac{\text{mass of X in 1 mol of compound}}{\text{mass of 1 mol of compound}} \times 100$

Example: In 190 g of MgCl2 (molar mass = 95 g/mol):

• Moles of MgCl2: $\frac{190}{95} = 2$ mol

• Moles of Mg: 2 mol; Moles of Cl: 4 mol

• Mass of Mg: $2 \times 24 = 48$ g; Mass of Cl: $4 \times 35.5 = 142$ g

• Mass % Mg: $\frac{24}{95} = 25\%$; Mass % Cl: $\frac{71}{95} = 75\%$

Determining a Chemical Formula from Experimental Data

Elemental analysis provides the mass percentage of each element in a compound. The empirical formula can be determined as follows:

1. Convert percentages to grams (assume 100 g sample).

2. Convert grams to moles using atomic masses.

3. Write a pseudoformula using moles as subscripts.

4. Divide all by the smallest number of moles.

5. Multiply ratios to obtain whole numbers if necessary.

Example: Hydrogen peroxide contains 5.88% H and 94.12% O.

• H: $\frac{5.88}{1} = 5.88$ mol; O: $\frac{94.12}{16} = 5.88$ mol

• Empirical formula: HO

Molecular Formulas for Compounds

The molecular formula is a whole-number multiple of the empirical formula. To determine it, you need the empirical formula and the molar mass.

• $\text{Molecular formula} = (\text{Empirical formula})_n$, where $n$ is a positive integer.

• $n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$

Example: Hydrogen peroxide: Empirical formula HO (mass = 17), molar mass = 34, so $n = 2$, molecular formula = H2O2.

Inorganic Nomenclature Flow Chart

Inorganic compounds are named according to systematic rules based on the types of elements and ions present. Flow charts help determine the correct naming convention for acids, bases, salts, and molecular compounds.

Organic Compounds

Organic compounds are primarily composed of carbon and hydrogen, sometimes with O, N, P, S, and other elements. The main element in organic chemistry is carbon.

• Early distinction: Organic (from living things) vs. inorganic (from nonliving sources).

• Modern organic compounds can be synthesized in the lab.

Carbon Bonding

Carbon forms four covalent bonds, allowing for a variety of structures:

• 4 single bonds (e.g., methane, CH4)

• 2 double bonds, 1 triple + 1 single, etc.

Examples: Propane (C3H8), Isobutane (C4H10), Cyclohexane (C6H12), Ethene (C2H4), Ethyne (C2H2), Acetic acid (CH3COOH)

Common Hydrocarbons

Name	Molecular Formula	Structural Formula	Space-Filling Model	Common Uses
Methane	CH4	Tetrahedral	Model	Natural gas
Propane	C3H8	Chain	Model	Fuel for lighters
Ethene	C2H4	Double bond	Model	Ripening agent

Additional info: Table entries inferred for clarity.

Families in Organic Compounds

Family	Name Ending	General Formula	Example	Occurrence/Use
Alcohols	-ol	R–OH	CH3CH2OH	Alcohol in beverages
Ethers	ether	R–O–R'	CH3OCH3	Solvent
Aldehydes	-al	R–CHO	CH3CHO	Perfume, flavors
Carboxylic acids	-oic acid	R–COOH	CH3COOH	Vinegar

Additional info: Table entries inferred for clarity.

Chemical Equations

Chemical equations are concise representations of chemical reactions, showing reactants and products, their states, and the stoichiometric coefficients.

• Example: $\mathrm{CH_4(g) + 2\ O_2(g) \rightarrow CO_2(g) + 2\ H_2O(g)}$

• Coefficients are used to balance equations according to the Law of Conservation of Mass.

• Subscripts indicate the number of atoms in a molecule; coefficients indicate the number of molecules.