Lesson 3: Atoms and Molecules

Each element is represented by a unique name and symbol.
Symbols are based on elemental properties or derived from scientists, places, astronomical bodies, or mythological characters.
Elemental symbols: One capital letter or a capital letter followed by a lowercase letter.

Consists of the symbols of the atoms in the molecule.
Each elemental symbol represents one atom of the element.
Subscripts indicate the number of atoms of an element present in the compound.
At least 2 atom types in compound (ex. triatomic - 3 types of atoms)
Example:
- Nitrogen dioxide: $NO_2$ (one nitrogen atom, two oxygen atoms).
- Sulfuric acid: H2SO4 (two hydrogen atoms, one sulfur atom, four oxygen atoms).

Atoms are composed of three subatomic particles: protons, neutrons, and electrons.
Protons(+, defines element) and neutrons(+- is a filler):
- Located in the nucleus.
- Tightly bound together.
Electrons(-):
- Located outside the nucleus.
- Move rapidly throughout a large volume of space surrounding the nucleus.

Atomic number:
- Equal to the number of protons in the nucleus.
- Symbol: Z
- Located above the element symbol in the periodic table.
Mass number:
- Equal to the sum of protons and neutrons in the nucleus.
- Symbol: A

Atoms with the same number of protons but different numbers of neutrons.
Same atomic number (Z) but different mass numbers (A).
Example: Chlorine-35 and Chlorine-37.
All isotopes of the same element have:
- Same number of electrons outside the nucleus.
- Same number of protons in the nucleus.

Atomic Mass:
- Numbers beneath the element symbol in the periodic table.
- Provide a means of comparing the masses of atoms.
- Atomic weight of elements is the average relative mass of the atoms in the isotope mixture.
Atomic Mass Unit (u):
- Used to express the relative masses of atoms.
- $1 u = 1/12$ the mass of a carbon-12 atom.
- One carbon-12 atom has a relative mass of 12 u.
- An atom with a mass equal to twice the mass of a carbon-12 atom would have a relative mass of 24 u.
Molecular Weight:
- Relative mass of a molecule expressed in atomic mass units.
- Calculated by adding together the atomic weights of the atoms in the molecule.
- Example: Water ( $H_2O$ )
  - MW = (2 × atomic weight of H) + (1 × atomic weight of O)
  - MW = $(2 × 1.01 u) + (1 × 16.00 u) = 18.02 u$
Example Problem: Molecular weight of urea, $CH<em>4N</em>2O$

Avogadro’s number: Number of atoms or molecules in a specific sample of an element or compound.
Mole (mol): Number of particles in a sample with a mass in grams equal to the atomic or molecular weight.
- 1 mol = $6.022 × 10^{23}$
- Example: 1 mol S atoms = $6.02 × 10^{23}$ S atoms = 32.1 g S
Relationships:
- 1 mol S atoms = $6.02 × 10^{23}$ S atoms
- $6.02 × 10^{23}$ S atoms = 32.1 g S
- 1 mol S atoms = 32.1 g S

Chemical formulas represent numerical relationships among atoms in a compound.
Example: $H_2O$ represents a 2:1 ratio of hydrogen to oxygen atoms.
Relationships:
- $6.02 × 10^{23} H_2O$ molecules contain $12.04 × 10^{23}$ H atoms and $6.02 × 10^{23}$ O atoms
- 1 mol $H_2O$ contains 2 mol H atoms and 1 mol O atoms
- 18.0 g $H_2O$ contains 2.0 g H and 16.0 g O
Mole Calculations:

Calculate the number of moles of Ca in a 15.84 g sample of Ca.
- $15.84 g Ca × (1 mole Ca / 40.08 g Ca) = 0.3952 moles Ca$

One mole of any compound is a sample with a mass in grams equal to the molecular weight of the compound.
- 1 mol CO2 $=$ 6.02 × 10^{23} CO2 $molecules = 44.0 g$ CO2
Relationships for factor-unit calculations:
- 1 mol CO2 $=$ 6.02 × 10^{23} CO2molecules
- 6.02 × 10^{23} CO2 $molecules = 44.0 g$ CO2
- 1 mol CO2 $= 44.0 g$ CO2
- 1 mol $CO_2$ contains 1 mol C atoms
- 1 mol $CO_2$ contains 2 mol O atoms
- 1 mol C atoms = 12.01 g C