CHE101 Chapter 3a

Key Concepts: Isotopes and Atomic Mass

  • Isotopes:

    • Discussion began on isotopes and their calculation of average mass using a formula.

    • The formula calculates the average atomic mass as a weighted average of all isotopes.

  • Weighted Average Formula: \text{Atomic Weight} = \sum \left( \text{mass of each isotope} \times \text{natural abundance} \right)

    • Natural abundance percentages of isotopes sum to 100%.

  • Key terms include:

    • Isotopes: Atoms of the same element with different numbers of neutrons.

    • Atomic weight: Average of the isotopic masses weighed by their natural abundance.

Subatomic Particles and Their Roles

  • The atom consists of three subatomic particles:

    • Protons: Positive charge, determine the element's identity.

    • Neutrons: Neutral charge, contribute to the atomic mass.

    • Electrons: Negative charge, located in the orbital regions around the nucleus.

  • Changing the number of protons identifies a new element, while changing neutrons results in isotopes. Changing electrons forms ions (charged atoms).

Formation of Ions

  • Ions: Can be positively charged (cations) or negatively charged (anions).

    • Creation of ions occurs through loss or gain of electrons.

    • Example:

    • If an atom has fewer electrons than protons, it becomes a positive ion (cation).

    • If it has more electrons than protons, it becomes a negative ion (anion).

  • Notable terminology:

    • Cation: An ion with a net positive charge (e.g., ext{Na}^+ ).

    • Anion: An ion with a net negative charge (e.g., ext{Cl}^- ).

  • Ions can be represented with their elemental symbols and charges, like ext{Ca}^{2+} for calcium and ext{S}^{2-} for sulfide.

Periodic Table Insights

  • The position on the periodic table helps determine the common ions or charges in chemical compounds.

  • Example:

    • Nitrogen in Group 15: Needs three electrons to mimic the nearest noble gas (neon).

    • Sulfur position similarly helps predict its charge.

Naming and Formulas for Compounds

  • Chemical Formulas:

    • Represent combinations of elements and their respective quantities.

    • Letters signify elemental symbols (C, O, H) and numbers indicate the quantity.

  • Key strategies for understanding chemical formulas:

    • What does each letter represent?

    • What do the numbers indicate?

  • Ionic vs. Covalent Compounds:

    • Ionic compounds include metals.

    • Covalent compounds consist only of nonmetals.

Characteristics of Covalent Compounds

  • Covalent compounds are formed when two or more nonmetals bond without an electrical charge. Each atom shares electrons.

  • How to interpret the formulas:

    • Example: ext{CO}_2 (carbon dioxide) has 1 carbon atom and 2 oxygen atoms. The structure indicates they are bonded together.

  • Key forms of identifying these compounds:

    • Binary Covalent Compounds: Composed of two different elements.

    • Naming conventions use prefixes (mono-, di-, tri-) to indicate the number of atoms.

Naming Conventions for Covalent Compounds

  • For naming covalent compounds:

    • Identify the first and second elements.

    • Use prefixes for quantity (except no prefix for one on the first element).

    • Change the suffix of the second element to -ide.

  • Examples:

    • ext{CO} is carbon monoxide (one carbon, one oxygen).

    • ext{CO}_2 is carbon dioxide (one carbon, two oxygens).

  • Notable prefixes include:

    • Mono- (1), Di- (2), Tri- (3), Tetra- (4), Penta- (5).

Diatomic Elements

  • Seven elements exist as diatomic molecules in their natural state:

    • Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2).

  • These elements do not exist as single atoms but as bonded pairs (e.g., ext{O}_2 for oxygen).