Introduction to General Chemistry Concepts

These study notes cover foundational topics in General Chemistry, including the periodic table, classification of elements, naming compounds, and balancing chemical equations. Mastery of these concepts is essential for success in introductory college chemistry courses.

Periodic Table and Classification of Elements

Development and Structure of the Periodic Table

The periodic table organizes all known chemical elements in a systematic way, revealing periodic trends in their properties. Russian chemist Dmitri Mendeleev is credited with its development, having noticed that elements arranged by increasing atomic mass showed repeating patterns of properties.

• Periods: Horizontal rows (left to right, L→R) on the periodic table. Properties of elements vary significantly across a period.

• Groups (Families): Vertical columns (top to bottom, T→B). Elements in the same group share similar chemical and physical properties.

• Group Numbering: The modern IUPAC system numbers groups 1–18. The older system uses Roman numerals and letters (e.g., 1A–8A for main group elements).

Classification of Elements

• Metals: Located on the left and center of the periodic table. Properties include:

  • Good conductors of heat and electricity

  • Ductile and malleable

  • Mostly solids at room temperature (except mercury)

  • Tend to lose electrons in chemical reactions (form cations)

• Nonmetals: Found on the right side of the table. Properties include:

  • Poor conductors (good insulators)

  • Can be gases, liquids, or solids at room temperature

  • Tend to gain electrons in chemical reactions (form anions)

• Metalloids (Semiconductors): Elements with properties intermediate between metals and nonmetals. Examples: B, Si, Ge, As, Sb, Te, At. Widely used in electronics (e.g., silicon in computer chips).

Special Groups and Their Names

• Alkali Metals (Group 1A): Highly reactive metals (e.g., Li, Na, K)

• Alkaline Earth Metals (Group 2A): Reactive metals (e.g., Mg, Ca)

• Halogens (Group 7A): Very reactive nonmetals (e.g., F, Cl, Br)

• Noble Gases (Group 8A): Inert gases (e.g., He, Ne, Ar)

Units and Conversion Factors

Length and Mass Units

Understanding SI and common units is essential for chemical calculations. Some conversion factors are exact, while others are approximate and will be provided on exams.

Naming Compounds

Types of Compounds

• Molecular (Covalent) Compounds: Formed by two or more nonmetals sharing electrons.

• Ionic Compounds: Formed by transfer of electrons between metals (cations) and nonmetals (anions).

Naming Molecular Compounds

Use Greek prefixes to indicate the number of each atom. The more metallic element is named first, followed by the less metallic element with an -ide suffix.

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

Naming Ionic Compounds

• Name the cation (metal) first, then the anion (nonmetal) with an -ide suffix.

• Example: NaCl is sodium chloride; CaO is calcium oxide.

• Do not use prefixes for ionic compounds.

Transition Metals and the Stock System

• Transition metals can form more than one cation. Use Roman numerals to indicate the charge.

• Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.

Polyatomic Ions

Polyatomic ions are groups of covalently bonded atoms with an overall charge. You must memorize the names, formulas, and charges of common polyatomic ions.

• Oxyanions: Polyatomic ions containing oxygen. Suffixes indicate the number of oxygens:

  • -ate: more oxygens (e.g., sulfate SO42-)

  • -ite: fewer oxygens (e.g., sulfite SO32-)

  • Prefixes per- (most oxygens) and hypo- (least oxygens) are used for halogen oxyanions (e.g., perchlorate ClO4-, hypochlorite ClO-).

Hydrates

Hydrates are ionic compounds that include water molecules in their structure. The number of water molecules is indicated by a Greek prefix.

• Example: CuSO4·5H2O is copper(II) sulfate pentahydrate.

• Common prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10).

Chemical Formulas: Molecular and Empirical

Molecular vs. Empirical Formulas

• Molecular Formula: Shows the exact number of atoms of each element in a molecule (e.g., H2O2 for hydrogen peroxide).

• Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., HO for hydrogen peroxide).

• Empirical formulas are often determined first in chemical analysis.

Balancing Chemical Equations

Law of Conservation of Mass

Chemical equations must be balanced to reflect the conservation of mass: the number of atoms of each element must be the same on both sides of the equation.

• Reactants are written on the left, products on the right of the reaction arrow.

• Physical states are indicated in parentheses: (s) solid, (l) liquid, (g) gas, (aq) aqueous.

Steps for Balancing Equations

1. Balance atoms in compounds first; leave single elements for last.

2. Use coefficients to balance the number of atoms.

3. If necessary, use fractional coefficients, then multiply all coefficients to obtain whole numbers.

• Example: Formation of ammonia:

• Example: Combustion of methane:

Types of Chemical Reactions

• Combination (Synthesis): Two or more reactants form one product.

• Decomposition: One reactant forms two or more products.

• Combustion: A hydrocarbon reacts with oxygen to form carbon dioxide and water.

Formula Weights and Percent Composition

Formula Weight

The formula weight (also called molecular weight) is the sum of the average atomic masses of all atoms in a chemical formula, expressed in atomic mass units (amu).

• Example: For H2O:

  • 2 × atomic mass of H + 1 × atomic mass of O

  • amu

Percent composition can be calculated by dividing the total mass of each element in a formula by the formula weight and multiplying by 100%.

Summary Table: Common Polyatomic Ions

Additional info: Some content was inferred and expanded for clarity and completeness, including the structure of tables and the explanation of chemical concepts.