Atoms, Molecules, and Ions – Chapter 2 Study Notes

Atoms, Molecules, and Ions

2.1 Atomic Theory of Matter

The atomic theory of matter describes the fundamental nature of substances and their composition. Early Greek philosophers, such as Democritus, proposed that all matter is made of tiny, indivisible particles called atoms. In the 1800s, John Dalton developed the first scientific atomic theory, which explained and supported several fundamental chemical laws.

• Law of Constant Composition: Compounds have a definite composition; the relative number of atoms of each element in a compound is the same in any sample.

• Law of Conservation of Mass: The total mass of substances present after a chemical process is the same as before the process.

• 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 small whole-number ratios.

Dalton's Atomic Theory – Postulates

1. Each element is composed of extremely small particles called atoms.

2. All atoms of a given element are identical in mass and properties, but atoms of different elements differ.

3. Atoms of one element cannot be changed into atoms of another element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

4. Atoms of more than one element combine to form compounds; a given compound always has the same relative number and kind of atoms.

2.2 Discovery of Subatomic Particles

Dalton considered atoms indivisible, but later experiments revealed that atoms are made of smaller particles: electrons, protons, and neutrons.

• Cathode Rays: Streams of electrons move from the cathode (–) to the anode (+) in a vacuum tube, making a fluorescent screen glow and being deflected by magnets, proving they are charged particles. J. J. Thomson discovered the electron in 1897.

• The Electron: Thomson measured the charge-to-mass ratio ($e/m$) of the electron: $1.76 \times 10^8 \ \text{C/g}$.

• Millikan Oil-Drop Experiment: Robert Millikan measured the charge of a single electron ($e = 1.602 \times 10^{-19} \ \text{C}$) and, using Thomson's $e/m$, determined the electron's mass.

• Early Atomic Models: Thomson's 'plum pudding' model (electrons in a positive matrix) and Nagaoka's 'Saturn-like' model (electrons orbiting a positive sphere).

• Discovery of the Nucleus: Ernest Rutherford's gold foil experiment showed that atoms have a small, dense, positively charged nucleus at the center, surrounded by electrons.

2.3 Modern View of Atomic Structure

Modern atomic theory describes atoms as mostly empty space, with a tiny, dense nucleus containing protons and neutrons, and electrons moving around the nucleus.

• Atoms are very small: $1$–$5 \ \text{Å}$ or $100$–$500 \ \text{pm}$.

• The nucleus is about 100,000 times smaller than the atom.

• Subatomic Particles:

  • Proton: Charge $+1$, mass $\approx 1.0073$ amu

  • Neutron: Charge $0$, mass$\approx 1.0087$ amu

  • Electron: Charge $-1$, mass $\approx 5.486 \times 10^{-4}$ amu

Particle	Charge	Mass (amu)
Proton	+1	1.0073
Neutron	0	1.0087
Electron	-1	5.486 × 10-4

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; determines the element's identity.

• Mass Number (A): Number of protons plus neutrons in the nucleus.

• Atoms are represented by their symbol, atomic number, and mass number (e.g., $^{12}_{6}\text{C}$ for carbon-12).

Symbols of Elements

Each element is represented by a one- or two-letter symbol (e.g., H for hydrogen, O for oxygen, Na for sodium).

Element	Symbol
Hydrogen	H
Carbon	C
Oxygen	O
Sodium	Na
Iron	Fe
Chlorine	Cl
Magnesium	Mg
Lead	Pb

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons (and thus different masses).

• Same element: same number of protons

• Different mass: different number of neutrons

Symbol	Number of Protons	Number of Electrons	Number of Neutrons
$^{12}$C	6	6	6
$^{13}$C	6	6	7
$^{14}$C	6	6	8

Ions

Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Cations: Positively charged ions (lost electrons)

• Anions: Negatively charged ions (gained electrons)

• In a neutral atom: number of protons = number of electrons

• In an ion: number of protons ≠ number of electrons

Example Calculations

• For $^{23}$Na:

  • Protons: 11

  • Electrons: 11

  • Neutrons: $23 - 11 = 12$

• For $^{43}$Ca$^{2+}$:

  • Protons: 20

  • Electrons: $20 - 2 = 18$

  • Neutrons: $43 - 20 = 23$

• For $^{17}_{8}$O$^{2-}$:

  • Protons: 8

  • Electrons: 10

  • Neutrons: $17 - 8 = 9$

2.4 Atomic Mass Unit (amu)

The atomic mass unit (amu) is a standard unit for measuring atomic and molecular masses. It is defined as 1/12 the mass of a carbon-12 atom.

• 1 amu = $1.66054 \times 10^{-24}$ g

• Each proton and neutron ≈ 1 amu

• Atomic mass of an atom ≈ mass number

Atomic Weight

Atomic weight is the average mass of all isotopes of an element, weighted by their natural abundance.

• Atomic Weight = $\sum [\text{isotope mass} \times \text{fractional natural abundance}]$

• Example (chlorine): $\text{Atomic weight} = (35 \times 0.75) + (37 \times 0.25) \approx 35.5$ amu

Additional info:

• These notes cover the foundational concepts of atomic theory, subatomic particles, isotopes, ions, and atomic mass, which are essential for understanding chemical reactions and the structure of matter in General Chemistry.