Chapter 2: Atoms and Elements

Einstein and Jean Perrin confirmed that particle motion is due to collisions with invisible molecules, providing evidence for the existence of molecules and atoms.
Robert Brown (1773–1858), a Scottish botanist, observed pollen grains suspended in water under a microscope.
Brown's observation showed a random, jittery motion of the pollen grains.
Initial interpretation: motion might indicate the grains were alive, but later explanation attributed it to continual collisions with surrounding molecules in the water.
This phenomenon offered direct evidence for the molecular/atomic nature of matter and the real motion of atoms/molecules in fluids.

Idea of fundamental particles dates back to ancient civilizations (Greek, Chinese, Hindu, Buddhist traditions around the 6th century BCE).
Democritus and Leucippus proposed an atomistic view: matter can be divided until reaching small, indivisible particles called atoms; atoms differ in shape and size; move randomly through empty space; various atoms combine in different ways to form all matter.
Key quote attributed to atomists: “Nothing exists except atoms and empty space.”
Plato and Aristotle did not embrace atomism; they argued that matter is made of four elements: Water, Fire, Earth, and Air.
They believed matter is infinitely divisible (no smallest unit).
John Dalton (1766–1844) later offered convincing evidence that supported the early atomic ideas of Leucippus and Democritus, bridging ancient ideas to modern chemistry.
Dalton’s work laid groundwork for turning atomistic ideas into a scientific theory compatible with experimental data.

The modern atomic theory emerged from observations and guiding laws that established consistent patterns in matter.
The three most important laws:
- Law of Conservation of Mass
- Law of Definite Proportions (also called the Law of Constant Composition)
- Law of Multiple Proportions

Formulated by Antoine Lavoisier.
States that in a chemical reaction, matter is neither created nor destroyed.
Mathematically: mass of the reactant = mass of products
This implies that mass is preserved across chemical processes, forming a foundational principle for balancing chemical equations and understanding reactions.

Observed by French chemist Joseph Proust in 1797.
All samples of a given compound, regardless of source or preparation, contain the same proportion of constituent elements.
Also called the law of constant composition.
Core idea: the composition of a compound is fixed by its identity, not by how it was formed.

Water, \mathrm{H_2O}
- Sample 1 (18 g water): H = 2 g, O = 16 g → \text{H:O} = 1:8
- Sample 2 (36 g water): H = 4 g, O = 32 g → \text{H:O} = 1:8
- Sample 3 (54 g water): H = 6 g, O = 48 g → \text{H:O} = 1:8
Ammonia, \mathrm{NH_3}
- Sample 1 (17 g ammonia): N = 14 g, H = 3 g → \text{N:H} = 14:3 \approx 4.67:1
- Sample 2 (34 g ammonia): N = 28 g, H = 6 g → \text{N:H} = 14:3 \approx 4.67:1
- Sample 3 (51 g ammonia): N = 42 g, H = 9 g → \text{N:H} = 14:3 \approx 4.67:1

Example problem: Determine the mass ratio of nitrogen to oxygen in dinitrogen pentoxide. (Atomic masses: N = 14.01\,\text{amu},\quad O = 16.00\,\text{amu})
For \mathrm{N2O5}, the masses are:
- N mass: 2 \times 14.01 = 28.02
- O mass: 5 \times 16.00 = 80.00
Mass ratio of N to O: \text{N:O} = 28.02:80.00 \approx 0.350:1 (or equivalently, 28.02:80.00\approx 0.350:1)

Published in 1804 by John Dalton.
If two elements (A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.
Conceptually, an atom of A can combine with one, two, three, or more atoms of B (AB1, AB2, AB3, …).
This law explains why different compounds made from the same elements have simple, related mass ratios.

Example: Carbon monoxide (CO) and carbon dioxide (CO₂) are two compounds formed from the same elements carbon (C) and oxygen (O).
Determine the mass ratio of carbon to oxygen in each compound using given atomic masses: C = 12.01, O = 16.00.
CO₂:
- Masses: C = 12.01, O (2 atoms) = 2 × 16.00 = 32.00
- Divide by the smallest mass (12.01): C → 1, O → 32.00/12.01 ≈ 2.67
- Ratio: \text{C:O} = 1:2.67
CO:
- Masses: C = 12.01, O (1 atom) = 16.00
- Divide by the smallest mass (12.01): C → 1, O → 16.00/12.01 ≈ 1.33
- Ratio: \text{C:O} = 1:1.33

The ratio of these two mass ratios is itself a small whole number:
- \frac{2.67}{1.33} \approx 2
This numerical relationship corroborates Dalton’s law that simple whole-number ratios govern compound formation.

Dalton’s atomic theory connected the laws to a coherent model of matter:
- Each element is composed of tiny, indestructible particles called atoms.
- All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements.
- Atoms combine in simple, whole-number ratios to form compounds.
- Atoms of one element cannot change into atoms of another element.
- In a chemical reaction, atoms only change the way they are bound together with other atoms.
This framework provided a consistent explanation for the laws and established the foundation for modern chemistry.