The Nature of Scientific Theories

Scientific Theories and Models

Scientific theories are comprehensive explanations of natural phenomena, built from well-established hypotheses and validated by experimental evidence. They aim to explain not only what happens in nature, but also why and how it occurs.

• Theory Formation: Theories arise from repeated experimental validation of hypotheses.

• Predictive Power: Good theories reliably predict future observations under various circumstances.

• Provisional Status: Theories are never absolutely proven; new observations may reveal limitations.

Early Ideas About Matter

Ancient Philosophical Concepts

Ancient Greek philosophers such as Democritus proposed that matter is composed of small, indestructible particles called atoms. They believed different types of atoms existed, each with unique shapes and sizes, moving randomly through empty space.

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

• Plato and Aristotle: Rejected atomic theory, proposing instead that matter was continuous and composed of four elements: fire, air, earth, and water.

Transition to Scientific Approach

The scientific method eventually replaced philosophical speculation. John Dalton provided experimental evidence supporting the atomic theory, laying the groundwork for modern chemistry.

Modern Atomic Theory and Foundational Laws

Key Laws Leading to Atomic Theory

The acceptance of atomic theory was driven by three fundamental laws:

• Law of Conservation of Mass: In a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products.

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

• 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 ratios of small whole numbers.

Dalton's Atomic Theory

Dalton's atomic theory provided a systematic explanation for these laws:

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

• All atoms of a given element are identical in mass and properties, but differ from atoms of other elements.

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

• Atoms are not changed into other elements in chemical reactions; they only rearrange how they are bonded.

Atomic Structure: Subatomic Particles

Discovery of the Electron

J. J. Thomson's cathode ray experiments revealed the existence of electrons—negatively charged particles present in all atoms.

• Electrons travel in straight lines, are independent of the cathode material, and carry a negative charge.

Charge-to-Mass Ratio of the Electron

Thomson measured the charge-to-mass ratio of the electron as using electric and magnetic fields.

Millikan's Oil Drop Experiment

Robert Millikan determined the charge of a single electron ( C) by balancing gravitational and electrical forces on tiny oil droplets.

Thomson's Plum-Pudding Model

Thomson proposed that electrons were embedded in a sphere of positive charge, like "plums in a pudding." This model was later disproven.

Radioactivity and the Discovery of Subatomic Particles

Radioactivity, discovered by Henri Becquerel and further studied by Ernest Rutherford, revealed that atoms emit high-energy radiation and contain additional subatomic particles.

Rutherford's Gold Foil Experiment and the Nuclear Model

Rutherford's experiment showed that most alpha particles passed through gold foil, but some were deflected, indicating a small, dense, positively charged nucleus at the center of the atom.

The Nuclear Atom

Rutherford's model described the atom as mostly empty space, with electrons orbiting a tiny, dense nucleus containing protons and neutrons.

Subatomic Particles: Properties and Comparison

Atoms are composed of three main subatomic particles:

• Protons: Positive charge (+1), mass ≈ 1 amu, located in the nucleus.

• Neutrons: No charge, mass ≈ 1 amu, located in the nucleus.

• Electrons: Negative charge (–1), mass ≈ 0.0005 amu, orbit the nucleus.

Atomic Mass and Isotopes

Atomic Mass Unit (amu)

The atomic mass unit (amu) is defined as one-twelfth the mass of a carbon-12 atom: .

Mass Spectrometry

Mass spectrometers measure atomic and molecular masses with high precision by separating ions based on their mass-to-charge ratio.

Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons. The atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.

• Atomic Weight Formula:

The Periodic Table and Periodic Law

Development of the Periodic Table

Dmitri Mendeleev organized elements by increasing atomic mass, observing that elements with similar properties recur periodically. This led to the modern periodic table, where elements are arranged by atomic number.

Structure of the Periodic Table

• Groups (Families): Vertical columns (numbered 1–18 or with A/B notation).

• Periods: Horizontal rows (numbered 1–7).

• Main-Group Elements: Groups 1, 2, and 13–18 (1A–8A).

• Transition Metals: Groups 3–12 (B groups).

Classification of Elements

• Metals: Good conductors, malleable, ductile, shiny, tend to lose electrons.

• Nonmetals: Poor conductors, not malleable or ductile, tend to gain electrons.

• Metalloids: Exhibit mixed properties, some are semiconductors.

Special Groups

• Alkali Metals (Group 1A): Highly reactive metals (e.g., Li, Na, K).

• Alkaline Earth Metals (Group 2A): Fairly reactive metals (e.g., Mg, Ca).

• Halogens (Group 7A): Very reactive nonmetals (e.g., F, Cl, Br, I).

• Noble Gases (Group 8A): Unreactive gases (e.g., He, Ne, Ar).

Atoms, Ions, and Chemical Formulas

Atomic Number and Mass Number

The atomic number (Z) is the number of protons in an atom and defines the element. The mass number (A) is the total number of protons and neutrons.

Ions and Ionic Compounds

Main-group metals lose electrons to form cations; nonmetals gain electrons to form anions. Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Alkali metals: Form 1+ ions.

• Alkaline earth metals: Form 2+ ions.

• Halogens: Form 1– ions.

• Oxygen family: Form 2– ions.

Chemical Formulas

• Empirical Formula: Lowest whole-number ratio of elements.

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

• Structural Formula: Shows the connectivity of atoms.

Summary Table: Subatomic Particles

Key Equations

• Law of Conservation of Mass:

• Atomic Weight: