Molecules and Compounds

Composition and Formulas

Understanding the composition and formulas of molecules and compounds is fundamental in chemistry. Compounds are substances formed from two or more elements chemically combined in fixed ratios.

• Molecule: A group of atoms bonded together, representing the smallest fundamental unit of a chemical compound.

• Chemical Formula: A notation that uses chemical symbols and numerical subscripts to convey the elements and their ratios in a compound (e.g., H2O for water).

• Example: Table salt is formed by the reaction .

Law of Constant Composition

The law of constant composition states that a given compound always contains its constituent elements in fixed proportions by mass.

• Definition: All samples of a compound have the same proportions of their constituent elements.

• Example: Water (H2O) always contains 2 hydrogen atoms and 1 oxygen atom. The mass ratio of oxygen to hydrogen in water is always 8.0:1.

• Formula:

Empirical and Molecular Formulas

Chemical formulas can be empirical (simplest whole-number ratio) or molecular (actual number of atoms in a molecule).

• Empirical Formula: Shows the simplest ratio of elements in a compound.

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

• Example: Glucose has a molecular formula of C6H12O6 and an empirical formula of CH2O.

Types of Compounds

Ionic and Molecular Compounds

Compounds are classified based on the types of elements involved and the nature of their bonding.

• Ionic Compounds: Formed from metals and nonmetals; consist of positive and negative ions held together by electrostatic forces.

• Molecular (Covalent) Compounds: Formed from nonmetals; consist of atoms sharing electrons.

• Example: NaCl is ionic; H2O is molecular.

Writing Chemical Formulas

Rules and Examples

Formulas are written by following conventions for element order and ratios.

• Nonmetals: Written in order from left to right as per Table 5.1 (not shown).

• Polyatomic Ions: Groups of atoms that act as a unit (e.g., OH-).

• Examples:

  • Fe2O3: Four oxide ions for every three iron ions.

  • C6H12O6: Six hydrogen atoms, twelve carbon atoms, six oxygen atoms.

  • Al2O3: Two aluminum atoms for every three oxygen atoms.

Periodic Table Trends

Metallic and Nonmetallic Character

The periodic table displays trends in element properties.

• Francium (Fr): Most metallic element.

• Fluorine (F): Most nonmetallic element.

• Trend: Metallic character increases down a group and to the left; nonmetallic character increases up a group and to the right.

Practice Problems and Key Concepts

Unit Conversions

Unit conversions are essential for quantitative chemistry.

• Length:

• Temperature:

• Density:

Conservation of Mass

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction.

• Formula:

• Example: If 24 g of natural gas reacts with 96 g of oxygen to produce 66 g of carbon dioxide, the mass of water formed is g.

Physical vs. Chemical Change

Distinguishing between physical and chemical changes is important in chemistry.

• Physical Change: Change in state or appearance without altering composition (e.g., sugar dissolving in water).

• Chemical Change: Change that alters the chemical composition (e.g., sugar burning).

Pure Substances and Mixtures

Classification of matter is based on composition.

• Pure Substance: Matter with a fixed composition (element or compound).

• Mixture: Combination of two or more substances (homogeneous or heterogeneous).

• Examples:

  • Element: Cu (copper)

  • Compound: H2O (water)

  • Homogeneous mixture: urine

  • Heterogeneous mixture: soil, Snickers bar

Significant Figures

Significant figures reflect the precision of measurements and calculations.

• Rules: The result of a calculation should have the same number of significant figures as the least precise measurement.

• Example:

Endothermic vs. Exothermic Reactions

Chemical reactions can absorb or release energy.

• Endothermic: Absorbs energy; products have higher potential energy than reactants.

• Exothermic: Releases energy; products have lower potential energy than reactants.

Definitions: Ions and Atomic Number

Understanding ions and atomic number is foundational.

• Ion: An atom or molecule with a net electric charge due to loss or gain of electrons.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Atomic Number: Number of protons in the nucleus of an atom.

• Alkaline Earth Metals: Group 2 elements; all are metals.

Kinetic and Potential Energy

Energy in chemistry is classified as kinetic or potential.

• Kinetic Energy: Energy of motion.

• Potential Energy: Energy due to position or composition.

• Example: A book on a shelf has potential energy; a book falling has kinetic energy.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Definition: Atoms with the same number of protons but different numbers of neutrons.

• Example: Chlorine-35 and Chlorine-37 are isotopes of chlorine.

• Chemical Notation: and

• Group: Chlorine belongs to group 17 (halogens).

• Properties: Halogens typically form -1 charge ions.

Prefix Multipliers

Prefix multipliers are used in scientific notation and unit conversions.

Neutral Particles vs. Ions

Atoms are neutral when the number of protons equals the number of electrons; ions have unequal numbers.

• Neutral Atom: Equal numbers of protons and electrons.

• Ion: Unequal numbers of protons and electrons, resulting in a net charge.

• Example: Sodium (Na) with 11 protons and 11 electrons is neutral; with 10 electrons, it is a cation with +1 charge.

Summary Table: Types of Matter

Additional info:

• Some slides and notes referenced textbook images and tables not included; standard chemistry conventions were used to fill in missing details.

• Practice problems covered significant figures, unit conversions, and classification of matter, which are foundational for introductory chemistry.