Chapter 2: Atoms and Elements

Imaging and Moving Individual Atoms

Atoms are the fundamental building blocks of matter, connecting the macroscopic and microscopic worlds. Despite their small size, they are the smallest identifiable unit of an element.

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

• There are about 91 naturally occurring elements and over 20 synthetic elements created by scientists.

Early Ideas About the Building Blocks of Matter

Ancient Philosophies

Early Greek philosophers such as Leucippus and Democritus (5th century B.C.) proposed that matter was composed of small, indestructible particles called atoms.

• Democritus stated: "Nothing exists except atoms and empty space; everything else is opinion."

• They suggested that atoms differed in shape and size and moved randomly through empty space.

Rejection and Later Acceptance

• Plato and Aristotle did not accept the atomic theory, believing instead that matter was continuous and composed of earth, air, fire, and water in various proportions.

• Scientific approaches later validated the atomic theory, especially through the work of John Dalton (1766–1844), who provided evidence supporting the existence of atoms.

Modern Atomic Theory and the Laws That Led to It

The modern atomic theory is based on experimental observations and several fundamental laws:

• Law of Conservation of Mass

• Law of Definite Proportions

• Law of Multiple Proportions

The Law of Conservation of Mass

Formulated by Antoine Lavoisier, this law states:

• In a chemical reaction, matter is neither created nor destroyed.

• The total mass of substances involved in a chemical reaction remains constant.

The Law of Definite Proportions

Proposed by Joseph Proust in 1797, this law states:

• All samples of a given compound, regardless of their source or preparation, have the same proportions of their constituent elements.

Example: Decomposition of 18.0 g of water yields 16.0 g of oxygen and 2.0 g of hydrogen. The mass ratio is:

Sample Calculation (from Example 2.1):

• Sample 1:

• Sample 2:

• Both samples have the same mass ratio, consistent with the law of definite proportions.

The Law of Multiple Proportions

Formulated by John Dalton in 1804, this law states:

• When two elements (A and B) form more than one compound, the masses of B that combine with a fixed mass of A are in the ratio of small whole numbers.

Example: The mass ratio of oxygen to carbon in carbon dioxide is 2.67:1, and in carbon monoxide it is 1.33:1. The ratio of these two ratios is:

This is a small whole number, consistent with the law.

Sample Calculation (from Example 2.2):

• Mass oxygen to 1 g nitrogen in nitrogen dioxide: 2.28

• Mass oxygen to 1 g nitrogen in dinitrogen monoxide: 0.570

• Ratio:

John Dalton and the Atomic Theory

Dalton's atomic theory explained the above laws as follows:

1. Each element is composed of tiny, indestructible particles called atoms.

2. All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

3. Atoms combine in simple, whole-number ratios to form compounds.

4. Atoms of one element cannot change into atoms of another element in a chemical reaction; they only change the way they are bound together.

The Discovery of the Electron

J. J. Thomson (1856–1940) discovered the electron through cathode ray experiments:

• He used a cathode ray tube to show that a beam of negatively charged particles (electrons) traveled from the cathode (negative electrode) to the anode (positive electrode).

• Electrons are negatively charged, low-mass particles present in all atoms.

• The charge-to-mass ratio of the electron is coulomb per gram.

Properties of Electrical Charge

• Like charges repel; opposite charges attract.

• Positive and negative charges of equal magnitude sum to zero when combined.

Millikan’s Oil Drop Experiment: The Charge of the Electron

Robert Millikan (1868–1953) measured the charge of a single electron using the oil drop experiment:

• The measured charge on any oil drop was always a whole-number multiple of C, the fundamental charge of a single electron.

Knowing the charge and Thomson’s charge-to-mass ratio, the mass of the electron can be calculated:

The Structure of the Atom

J. J. Thomson proposed the "plum-pudding model" of the atom:

• Negatively charged electrons are embedded within a positively charged sphere.

• This model was widely accepted until later experiments revealed the nuclear structure of the atom.

Additional info: Later experiments by Ernest Rutherford (not shown in these slides) led to the nuclear model of the atom, where a dense, positively charged nucleus is surrounded by electrons.