Fundamental Concepts in General Chemistry: Units, Atomic Structure, and Chemical Compounds

Conversion Factors and Units

Temperature Conversions

Temperature is a fundamental physical property and is commonly measured in Celsius, Fahrenheit, or Kelvin. Conversion between these units is essential in chemistry calculations.

• Celsius to Fahrenheit: $\text{Celsius} = \frac{5}{9} \times (\text{Fahrenheit} - 32)$

• Fahrenheit to Celsius: $\text{Fahrenheit} = \frac{9}{5} \times (\text{Celsius}) + 32$

• Kelvin to Celsius: $\text{Kelvin} = 273 + \text{Celsius}$

Conversion Factor Formula:

• $\text{Given Quantity} \times \frac{\text{Desired Quantity}}{\text{Conversion Factor}}$

Density, Volume, and Mass

Density is a physical property defined as mass per unit volume. It is useful for identifying substances and converting between mass and volume.

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Volume: $\text{Volume} = \frac{\text{Mass}}{\text{Density}}$

• Mass: $\text{Mass} = \text{Density} \times \text{Volume}$

Law of Definite and Multiple Proportions

Law of Definite Proportion

This law states that a chemical compound always contains the same proportion of elements by mass.

• Formula: $\text{Element mass} / \text{Compound mass} \times 100$

Law of Multiple Proportions

When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Example: Carbon monoxide (CO) and carbon dioxide (CO2)

Stepwise Calculation:

1. Determine atomic mass of 1st and 2nd element for each substance. (e.g., C = 12, O = 16 for CO; C = 12, O = 32 for CO2)

2. Divide both numbers by the atomic mass unit (AMU) of the 1st element. $\frac{16}{12} = 1.33$ for CO; $\frac{32}{12} = 2.66$ for CO2

3. Divide the solutions for each compound by each other. $\frac{2.66}{1.33} = 2$

This shows the ratio is a small whole number (2).

Atomic Structure

Atomic Number and Isotopes

The atomic number is the number of protons in an atom and defines the element. Isotopes are atoms of the same element with different numbers of neutrons, resulting in different masses.

• Atomic Number (Z): Number of protons; determines element identity.

• Isotopes: Same number of protons, different number of neutrons.

• Example: $^{12}\text{C}$, $^{13}\text{C}$, $^{14}\text{C}$ (all carbon isotopes with 6 protons, but different neutrons and masses)

Average Atomic Mass

The average atomic mass (atomic weight) is the weighted average of the masses of an element's isotopes as they occur naturally.

• Formula: $\text{Atomic weight} = \sum (\text{isotope mass} \times \text{fractional isotope abundance})$

• Step 1: Find the abundance of each isotope.

• Step 2: Multiply each isotope's mass by its fractional abundance and sum the results.

Ions and Ionic Compounds

Cations and Anions

Ions are charged particles formed when atoms gain or lose electrons.

• Cation: Positively charged ion (loss of electron)

• Anion: Negatively charged ion (gain of electron)

Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations (usually metals) and anions (usually nonmetals). Their naming depends on the type of metal involved.

• Binary Ionic Compounds (No Transition Metal):

  • Composed of only two elements

  • Name ends in -ide

• Ternary Ionic Compounds (Polyatomic Ions):

  • Name ends in -ate or -ite (e.g., cyanide, hydroxide, peroxide)

• Transition Metal Ionic Compounds:

  • Use Roman numerals to indicate the charge of the metal (e.g., Iron(III) chloride)

Molecular (Covalent) Compounds

Definition and Naming

Molecular compounds are composed of nonmetals only, and electrons are shared between atoms. Naming uses numerical prefixes to indicate the number of each atom present, except for 'mono-' on the first element.

• No Oxygen:

  • Start name with 'hydro'

  • End in -ic

• With Oxygen:

  • If the anion ends in -ate, acid name ends in -ic

  • If the anion ends in -ite, acid name ends in -ous

Molecular and Empirical Formulas

• Molecular Formula: Shows the actual number of atoms of each element in a molecule (e.g., H2O, C6H12O6).

• Empirical Formula: Shows the simplest whole-number ratio of atoms in a compound (e.g., CH2O for glucose).

Summary Table: Types of Compounds and Naming Rules

Type of Compound	Constituents	Naming Rule	Example
Ionic (Binary, no transition metal)	Metal + Nonmetal	End in -ide	NaCl (Sodium chloride)
Ionic (with transition metal)	Transition metal + Nonmetal	Use Roman numeral for metal charge	FeCl3 (Iron(III) chloride)
Ionic (Polyatomic ion)	Metal + Polyatomic ion	End in -ate or -ite	NaNO3 (Sodium nitrate)
Molecular (Covalent)	Nonmetal + Nonmetal	Use prefixes (mono-, di-, tri-, etc.)	CO2 (Carbon dioxide)