Chemistry Chapter: Atomic Theory and Proportions

The Law of Definite Proportions

• In a chemical compound, the elements are always combined in the same proportions by mass.
• For example, water is always composed of 8 parts by mass of oxygen and 1 part by mass of hydrogen.

Dalton’s Atomic Theory

• Proposed in 1803 to explain the law of conservation of mass and the law of definite proportions.
• Matter is composed of atoms, which are tiny, indivisible particles.
• All atoms of a given element are the same, but differ from atoms of other elements.
• Compounds are formed when atoms of different elements unite in fixed proportions.
• Chemical reactions involve the rearrangement of atoms, no atoms are created, destroyed, or broken apart.

The Law of Multiple Proportions

• When two or more compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers.
• Example:
  • In Carbon monoxide (CO), 12.0 g of carbon combines with 16.0 g of oxygen.
  • In Carbon dioxide (CO2), 12.0 g of carbon combines with 32.0 g of oxygen.
  • The ratio of oxygen/carbon in these two compounds is 1:2.

The Modern View of Atomic Structure

• The atom has a structure that includes a nucleus and an electron cloud.
• The nucleus contains protons and neutrons:
  • Protons carry a positive charge.
  • Neutrons carry no charge.
• The electron cloud contains negatively charged electrons.
• Protons and neutrons are much heavier than electrons.

Atomic Notation

• Atomic number (Z) is the number of protons in an atom.
• Mass number (A) is the sum of protons and neutrons in an atom.
• Charge (q) is the difference between protons and electrons (protons - electrons).
• The symbol for an element is written as: ^A_ZQ where Z is the atomic number, A is the mass number, and Q is the charge.

Ions

• An anion is a negatively charged ion formed when an atom gains electrons.
• A cation is a positively charged ion formed when an atom loses electrons.

Molecules and Ions

• Molecular compounds involve shared electrons and consist of electrically neutral, discrete particles (molecules).
• Ionic compounds involve electron transfer between metals and nonmetals and consist of charged particles (ions).

Chemical Bonds

• Covalent bonds involve shared electrons between atoms.
• Ionic bonds involve the complete transfer of electrons from one atom to another.

Types of Formulas

• Empirical formulas represent the lowest whole number ratio of atoms in a compound.
• Molecular formulas show the true number of atoms of each element in the formula of a compound.
• Structural formulas show how atoms are attached together and the order in which atoms are bonded.

The Periodic Table

• Organizes elements based on increasing atomic number.
• Rows are called periods.
• Columns are called groups or families.
• Elements within a group share similar properties due to similar valence electron configurations.

Naming Compounds

• For ionic compounds:
  • Name the cation.
  • Indicate the charge of the cation using Roman numerals in parentheses.
  • Name the anion.
• For covalent compounds:
  • Use prefixes (mono, di, tri, etc.) to indicate the number of atoms of each element.
  • Name the second element using the suffix "-ide".

Identifying Empirical Formulas

• An empirical formula represents the simplest whole number ratio of atoms in a compound.
• Divide all subscripts in a molecular formula by the greatest common factor to get its empirical formula.

Polyatomic Ions

• Consist of multiple atoms with a net charge.
• Have specific names that need to be memorized.
• Some polyatomic ion names change based on the number of oxygen atoms present.
• Use prefixes like "per-", "hypo-", and "-ite" or "-ate" to denote changes in the number of oxygen atoms within the ion.