Atoms, Molecules, and Ions - Summary

Early Ideas in Atomic Theory

• Dalton's Atomic Theory (1803-1807):
  1. Matter is composed of atoms.
  2. Atoms of a given element have identical properties.
  3. Atoms of one element differ from atoms of other elements.
  4. Compounds are combinations of atoms in whole-number ratios.
  5. Atoms are not created or destroyed in chemical reactions.
• Law of Constant Composition/Law of Definite Proportion: Pure compounds contain the same elements in the same proportion by mass (Joseph Proust).
• Law of Multiple Proportions (Dalton’s Law): When two elements react, the mass of one element reacts with masses of a second element in a ratio of small, whole numbers.

Subatomic Particles and Atomic Structure

• J.J. Thomson:
  • Discovered electrons using cathode ray tubes.
  • Determined charge-to-mass ratio of electrons: $1.76 \times 10^8$ C/g.
  • Proposed the "plum pudding" model.
• R.A. Millikan (1911):
  • Determined the charge on an electron: $-1.6022 \times 10^{-19}$ C using the oil drop experiment.
• Ernest Rutherford:
  • Proposed the nuclear model based on the gold foil experiment.
  • Positive charge and mass are concentrated in the nucleus.
• Subatomic Particles:
  • Protons: positively charged, in the nucleus.
  • Neutrons: neutral, in the nucleus.
  • Electrons: negatively charged, distributed around the nucleus.

Atomic Number, Mass Number, and Isotopes

• Atomic Number (Z):
  • Number of protons in the nucleus; determines the identity of an element.
• Mass Number (A):
  • Total number of protons and neutrons (nucleons) in the nucleus.
    $^{\text{A}}_{\text{Z}}X$
• Isotopes:
  • Atoms of the same element with different mass numbers (different numbers of neutrons).

Isotopes and Atomic Weight

• Atomic Mass:
  • Mass of an atom in atomic mass units (amu).
• Average Atomic Mass:
  • The weighted average of the masses of naturally occurring isotopes.
  • Calculated using isotopic masses and natural abundances.
• Mass Spectrometry:
  • A method for determining atomic and molecular masses.

Covalent Bonding and Molecules

• Molecule:
  • A combination of at least two atoms held together by chemical bonds.
• Diatomic Molecules:
  • Contain two atoms; can be homonuclear or heteronuclear (BrINClHOF).
• Polyatomic Molecules:
  • Contain more than two atoms.
• Chemical Formula:
  • Denotes the composition of a substance.
• Allotropes are different forms of the same element.

Molecular and Empirical Formulas

• Molecular Formula:
  • Shows the exact number of atoms of each element in a molecule.
• Empirical Formula:
  • The simplest whole-number ratio of elements in a molecule.

Molecular and Formula Mass

• Molecular Mass:
  • The mass in atomic mass units (amu) of an individual molecule.
• Formula Mass:
  • The mass of a formula unit of an ionic compound.

The Mole and Molar Mass

• Mole:
  • The amount of a substance that contains as many elementary entities as there are atoms in exactly 12 g of carbon-12.
  • Avogadro’s Number (NA): $6.022 \times 10^{23}$
• Molar Mass:
  • The mass (grams) in 1 mole of a substance (g/mol).
    $moles = \frac{grams}{Molar\,Mass}$

Molecular Mass and Molar Mass

• Molecular mass (amu) = molar mass (grams)

Conversion between Mass, Moles, Number of Atoms

• Molar mass converts between mass and moles.
• Avogadro’s constant converts between moles and number of atoms.

Molar Mass

• Molar mass (M) of a substance is the mass in grams of one mole of the substance.
• Molar mass of a compound is the sum of molar masses of the elements it contains.
  1 mol H2O = $2 \times 1.008$ g + 16.00 g = 18.02 g