General Chemistry - Chapter 1

1. Introduction to Chemistry and Matter

1.1 Atoms, Molecules, and Chemistry

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. At the core of chemistry are atoms and molecules, which make up all substances.

• Atom: The smallest unit of an element that retains its chemical properties.

• Molecule: A group of two or more atoms bonded together.

• Chemistry: The science that studies the composition, structure, properties, and changes of matter.

1.2 Mixtures, Pure Substances, Elements, and Compounds

Matter can be classified based on its composition and uniformity.

• Mixture: A physical combination of two or more substances where each retains its own properties. Mixtures can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

• Pure Substance: Matter with a fixed composition and distinct properties; includes elements and compounds.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions.

Example: Salt water is a homogeneous mixture, while sand and iron filings form a heterogeneous mixture.

2. Modern Atomic Theory

2.1 Laws of Conservation and Proportions

Modern atomic theory is based on several fundamental laws:

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A chemical compound always contains the same elements in the same proportions by mass.

• Dalton's Atomic Theory: John Dalton proposed that all matter is made of indivisible atoms, atoms of the same element are identical, and chemical reactions involve rearrangement of atoms.

3. Discovery and Structure of the Atom

3.1 Discovery of the Electron

Experiments with cathode ray tubes by J.J. Thomson provided evidence for the existence of the electron, a negatively charged subatomic particle.

• Cathode Ray Tube Experiment: Showed that cathode rays were composed of negatively charged particles (electrons).

Additional info: Thomson measured the charge-to-mass ratio of the electron, leading to further discoveries about atomic structure.

3.2 Structure of the Atom

Atoms consist of a dense nucleus containing protons and neutrons, surrounded by electrons.

• Radioactivity: The spontaneous emission of radiation from an unstable atomic nucleus.

• Nucleus: The central core of the atom, containing protons and neutrons.

• Proton: Positively charged subatomic particle found in the nucleus.

• Neutron: Electrically neutral subatomic particle found in the nucleus.

Example: Rutherford's gold-foil experiment demonstrated the existence of a small, dense, positively charged nucleus.

4. Subatomic Particles: Protons, Neutrons, and Electrons

4.1 Atomic Mass Unit, Atomic Number, and Chemical Symbol

Atoms are characterized by their numbers of protons, neutrons, and electrons.

• Atomic Mass Unit (amu): A unit of mass used to express atomic and molecular weights; 1 amu is defined as one-twelfth the mass of a carbon-12 atom.

• Atomic Number (Z): The number of protons in the nucleus of an atom; determines the element.

• Chemical Symbol: One- or two-letter abbreviation for an element (e.g., H for hydrogen, O for oxygen).

4.2 Isotopes and Relative Abundance

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

• Mass Number (A): The total number of protons and neutrons in an atom's nucleus.

• Relative Abundance: The percentage of a particular isotope in a natural sample of the element.

Example: Carbon-12 and Carbon-14 are isotopes of carbon, with mass numbers 12 and 14, respectively.

4.3 Ions: Anions and Cations

• Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.

• Anion: A negatively charged ion (gains electrons).

• Cation: A positively charged ion (loses electrons).

5. Atomic Mass and the Mole

5.1 Average Atomic Mass

The average atomic mass of an element is calculated from the masses and relative abundances of its isotopes.

• Formula:

$\text{Average Atomic Mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Mass Spectrometry: An analytical technique used to measure the masses and relative abundances of isotopes.

5.2 The Mole and Avogadro's Number

The mole is a fundamental unit in chemistry for counting particles such as atoms, molecules, or ions.

• Mole (mol): The amount of substance containing as many entities as there are atoms in 12 grams of carbon-12.

• Avogadro's Number ($N_A$): The number of particles in one mole, given by:

$N_A = 6.022 \times 10^{23}$ particles/mol

• Relates mass, number of moles, and number of particles.

5.3 Converting Between Moles, Mass, and Particles

• To convert between mass, moles, and number of particles, use the following relationships:

$\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

$\text{Number of Particles} = \text{Moles} \times N_A$

Example: Calculate the number of atoms in 10 grams of aluminum (Al):

• Molar mass of Al = 26.98 g/mol

• Moles of Al = $\frac{10}{26.98} = 0.370$ mol

• Number of atoms = $0.370 \times 6.022 \times 10^{23} = 2.23 \times 10^{23}$ atoms