Chapter 3: Molecules, Compounds, and Nomenclature – Study Notes

Chapter 3: Molecules, Compounds, and Nomenclature

Chemical Bonds

Chemical bonds are the forces that hold atoms together in compounds. They arise from interactions between charged particles (protons and electrons) within atoms, resulting in the formation of stable chemical species.

• Chemical bond: The attractive force that holds atoms or ions together in a compound.

• Result of interactions: Bonds form due to the electrostatic attraction between positively charged nuclei and negatively charged electrons.

• Importance: Chemical bonds determine the structure, properties, and reactivity of substances.

• Example: Water (H2O) is formed by chemical bonds between hydrogen and oxygen atoms.

Types of Chemical Bonds

There are two main types of chemical bonds: ionic and covalent. The type of bond depends on the nature of the atoms involved (metallic or nonmetallic).

Ionic Bonds (Metal/Nonmetal)

Ionic bonds occur when electrons are transferred from one atom (typically a metal) to another (typically a nonmetal), resulting in the formation of oppositely charged ions that are held together by electrostatic forces.

• Electron transfer: Metals lose electrons to form cations (+), nonmetals gain electrons to form anions (−).

• Electrostatic attraction: Oppositely charged ions attract each other, forming a stable ionic compound.

• Lattice structure: In the solid phase, multiple ion pairs group together in a regular three-dimensional array of alternating cations and anions.

• Example: Sodium chloride (NaCl) forms from Na+ and Cl− ions.

$\text{Na} \rightarrow \text{Na}^+ + e^-$ $\text{Cl} + e^- \rightarrow \text{Cl}^-$ $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

Covalent Bonds (Two Nonmetals)

Covalent bonds form when two nonmetal atoms share electrons, resulting in the formation of molecules. This sharing lowers the potential energy of the system through electrostatic interactions.

• Electron sharing: Each atom contributes one or more electrons to a shared pair.

• Molecular compounds: Composed of individual covalently bonded molecules (e.g., H2O, CO2, PCl3).

• Stability: The most stable form is achieved when the potential energy is minimized.

• Example: Water (H2O) and carbon dioxide (CO2) are covalent compounds.

$\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Properties of Ionic and Covalent Compounds

The properties of compounds depend on the type of bonding present.

• Ionic compounds: High melting points, solid at room temperature, conduct electricity when dissolved in water, form crystalline lattices.

• Covalent compounds: Lower melting points, can be gases, liquids, or solids at room temperature, do not conduct electricity in solution.

• Example Table:

Property	Hydrogen	Oxygen	Water
Melting Point	−253°C	−183°C	100°C
State at Room Temp	Gas	Gas	Liquid
Flammability	Explosive	Necessary for combustion	Used to extinguish fire

Chemical Formulas

Chemical formulas represent the composition of compounds, indicating the elements present and the relative number of atoms or ions of each.

• Empirical formula: Shows the simplest whole-number ratio of atoms in a compound.

• Molecular formula: Shows the actual number of atoms of each element in a molecule.

• Structural formula: Shows the arrangement of atoms and the bonds between them.

• Example: For glucose, empirical formula is CH2O, molecular formula is C6H12O6.

Classification of Elements and Compounds

Substances can be classified as atomic elements, molecular elements, molecular compounds, or ionic compounds.

• Atomic element: Exists as single atoms (e.g., Al, Co).

• Molecular element: Exists as molecules made of two or more atoms of the same element (e.g., O2, Cl2).

• Molecular compound: Molecules made of different nonmetals (e.g., CO, C3H6O).

• Ionic compound: Composed of cations and anions (e.g., NaCl, AlCl3).

Chemical Nomenclature (Naming Compounds)

Chemical nomenclature provides systematic names for compounds based on established rules, allowing chemists to communicate clearly about chemical substances.

• Common names: Traditional names (e.g., baking soda, quicklime).

• Systematic names: Based on IUPAC rules, derived from chemical formulas.

• IUPAC: International Union of Pure and Applied Chemistry sets the standards for chemical naming.

Naming Ionic Compounds

• Contain positive (cation) and negative (anion) ions.

• Formula unit is electrically neutral; charges must balance.

• For metals with variable charge, indicate charge with Roman numerals (e.g., Fe2+ is iron(II)).

• Polyatomic ions retain their names in compounds (e.g., sulfate, carbonate).

• Example: Na2SO4 is sodium sulfate.

Naming Molecular Compounds

• Composed of two or more nonmetals.

• Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.).

• First element is more metallic; second element ends with -ide.

• Never use the prefix mono- for the first element.

• Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

Naming Acids

• Acids contain H+ cation and an anion.

• If the anion does not contain oxygen, use the prefix hydro- and suffix -ic (e.g., HCl (aq) is hydrochloric acid).

• If the anion contains oxygen (oxyanion): - If the anion ends in -ate, change to -ic acid (e.g., HNO3 is nitric acid). - If the anion ends in -ite, change to -ous acid (e.g., HNO2 is nitrous acid).

Formula Mass, Molar Mass, and the Mole

Formula mass (or molecular mass) is the sum of the atomic masses of all atoms in a chemical formula. Molar mass is the mass of one mole of a substance, numerically equal to the formula mass in grams per mole.

• Formula mass: $\text{Formula mass} = \sum \text{atomic masses of all atoms in the formula}$

• Molar mass: $\text{Molar mass} = \text{Formula mass (g/mol)}$

• Example: For CHCl3: $\text{Molar mass} = (1 \times 12.01) + (1 \times 1.008) + (3 \times 35.45) = 119.37\ \text{g/mol}$

Percent Composition by Mass

Percent composition expresses the mass percentage of each element in a compound.

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

• Example: For C2H6: $\text{Mass percent C} = \frac{2 \times 12.01}{30.068} \times 100\% = 79.89\%$ $\text{Mass percent H} = \frac{6 \times 1.008}{30.068} \times 100\% = 20.11\%$

Empirical and Molecular Formulas

The empirical formula shows the simplest whole-number ratio of atoms in a compound, while the molecular formula shows the actual number of atoms in a molecule. The molecular formula is a whole-number multiple of the empirical formula.

• Empirical formula: Simplest ratio of elements.

• Molecular formula: Actual number of atoms; $\text{Molecular formula mass} = \text{Empirical formula mass} \times n$

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

• Example: If empirical formula is CH and molar mass is 78 g/mol: $n = \frac{78}{13} = 6$ Molecular formula is C6H6.

Determining Empirical Formula from Experimental Data

Empirical formulas can be determined from mass percent composition or combustion analysis.

• Assume a 100 g sample; convert mass percent to grams.

• Convert grams to moles using atomic masses.

• Write a pseudoformula using mole values as subscripts.

• Divide all subscripts by the smallest value to get whole numbers.

• If necessary, multiply to obtain whole-number subscripts.

• Example: Combustion analysis of a hydrocarbon yields masses of CO2 and H2O; use these to determine moles of C and H, then empirical formula.

$\text{Mass percent} = \frac{\text{mass of element}}{\text{total mass}} \times 100\%$

Conversion Factors in Chemical Formulas

Chemical formulas provide mole-to-mole relationships between atoms and molecules, which can be used to convert between mass, moles, and number of particles.

• Avogadro's number: $6.022 \times 10^{23}$ particles/mol

• Example: 1 mol CaCO3 = 100.09 g = $6.022 \times 10^{23}$ formula units

• Mass to moles: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Moles to particles: $\text{Number of particles} = \text{moles} \times 6.022 \times 10^{23}$

Summary Table: Types of Compounds

Type	Example	Description
Atomic Element	Al, Co	Single atoms
Molecular Element	O2, Cl2	Two or more atoms of same element
Molecular Compound	CO, C3H6O	Two or more nonmetals
Ionic Compound	NaCl, AlCl3	Cations and anions

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