General Chemistry: Atomic Structure, Isotopes, and Laws of Chemical Combination – Study Notes

Atomic Structure and Isotopes

Definition and Identification of Isotopes

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. They are commonly represented by the element's name followed by the mass number.

• Isotope Notation: Element name-mass number (e.g., carbon-13, uranium-238).

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Example: Bromine exists as a mixture of bromine-79 and bromine-81 isotopes.

Calculating Subatomic Particles:

• Number of protons = Atomic number (Z)

• Number of neutrons = Mass number (A) – Atomic number (Z)

• Number of electrons = Number of protons (for a neutral atom)

Examples of Isotopes

• Chlorine-35: 17 protons, 18 neutrons, 17 electrons

• Chlorine-37: 17 protons, 20 neutrons, 17 electrons

• Carbon-12: 6 protons, 6 neutrons, 6 electrons

• Carbon-14: 6 protons, 8 neutrons, 6 electrons

Atomic Ions and Electron Counting

Formation of Ions

Atoms can gain or lose electrons to form ions. The charge of an ion is determined by the difference between the number of protons and electrons.

• Cation: Positively charged ion (fewer electrons than protons)

• Anion: Negatively charged ion (more electrons than protons)

• Example: An atom with 15 protons and a mass number of 31 (phosphorus-31):

  • Neutral atom: 15 electrons

  • As3– ion: 18 electrons (gained 3 electrons) why is it As

Determining Subatomic Particles in Ions

• Number of protons = Atomic number

• Number of neutrons = Mass number – Atomic number

• Number of electrons = Number of protons – (charge on ion)

• Example: For an anion formed by one atom of 75As3–:

  • Protons: 33

  • Neutrons: 75 – 33 = 42

  • Electrons: 33 + 3 = 36

Laws of Chemical Combination

Law of Multiple Proportions

The law of multiple proportions states that when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers.

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

• Application: Used to distinguish between compounds formed from the same elements in different ratios.

Law of Definite Proportions

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

• Example: Water (H2O) always contains hydrogen and oxygen in a mass ratio of 1:8.

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

• Antoine Lavoisier is credited with this law and is known as the father of modern chemistry.

Identifying Compounds and Elements

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of atoms of each element in a molecule.

• Example: The empirical formula of hydrogen peroxide is HO, while the molecular formula is H2O2.

Determining the Identity of Elements

• The number of protons (atomic number) uniquely identifies an element.

• The number of electrons in a neutral atom equals the number of protons.

• Example: An element with 26 electrons (neutral atom) is iron (Fe), since its atomic number is 26.

Practice Problems and Applications

Sample Calculations

• Calculating Mass from Proportions: If 5.0 g of oxygen, 10.0 g of carbon, and 20.0 g of nitrogen are present in a sample, and the law of definite proportions applies, then a 70 g sample would contain:

  • Oxygen: $\frac{5.0}{35.0} \times 70 = 10.0$ g

  • Carbon: $\frac{10.0}{35.0} \times 70 = 20.0$ g

  • Nitrogen: $\frac{20.0}{35.0} \times 70 = 40.0$ g

Table: Comparison of Law of Definite and Multiple Proportions

Law	Description	Example
Law of Definite Proportions	Each compound has a fixed ratio of elements by mass	Water (H2O): 1 g H : 8 g O
Law of Multiple Proportions	Elements can combine in different ratios to form different compounds	CO and CO2: 12 g C : 16 g O (CO), 12 g C : 32 g O (CO2)

Key Terms

• Isotope: Atoms of the same element with different numbers of neutrons

• Ion: Atom or molecule with a net electric charge due to loss or gain of electrons

• Atomic Number (Z): Number of protons in the nucleus

• Mass Number (A): Total number of protons and neutrons

• Empirical Formula: Simplest whole-number ratio of atoms in a compound

• Molecular Formula: Actual number of atoms of each element in a molecule