Notes: Atomic Structure and Matter Classification
Classifying Matter
Matter: Anything with mass and occupying space; consists of atoms and molecules in motion. The physical universe is entirely comprised of matter and energy.
Atoms: Submicroscopic particles, building blocks of elements and molecules. Each atom has a unique identity based on its atomic number.
Molecules: Two or more atoms combined through chemical bonds, forming a discrete unit.
States of Matter:
Solid: Possesses a fixed volume and rigid shape. Particles are tightly packed in a fixed, ordered arrangement, vibrating in place. This is the most ordered state of matter.
Liquid: Has a distinct volume but takes the shape of its container. Particles are closely spaced but are free to move past one another, allowing liquids to flow.
Gas: Neither a fixed volume nor a fixed shape; it expands to fill its container. Particles are far apart and move randomly and rapidly. This is the least ordered state.
Composition Classification:
Pure Substances: Matter with distinct properties and a constant composition throughout. They cannot be separated into simpler substances by physical means.
Elements: Fundamental substances that cannot be chemically decomposed into simpler substances. Each element is defined by its atomic number (number of protons). Examples: Oxygen (O), Gold (Au).
Compounds: Substances formed when two or more different elements are chemically combined in fixed, definite proportions. Their properties are distinct from those of their constituent elements. Examples: Water (H{2}O), Carbon Dioxide (CO{2}).
Mixtures: Combinations of two or more pure substances that are physically mixed but not chemically combined. Their proportions can vary, and they retain the properties of their components.
Homogeneous: Mixtures with a uniform composition and appearance throughout. All parts of a homogeneous mixture are identical (e.g., salt water, air). Also known as solutions.
Heterogeneous: Mixtures where the composition is not uniform; different regions have different properties or appearances. Components are visibly distinguishable (e.g., sand and water, oil and vinegar).
Properties of Matter:
Intrinsic Properties: Independent of the amount of substance present (e.g., density, melting point, boiling point, color).
Extrinsic Properties: Dependent on the amount of substance present (e.g., mass, volume, length).
The Scientific Method
Empirical: Based on observation and experimentation.
Terminology:
Observation: Phenomenon competent observers agree upon, often quantitative measurements.
Scientific Law: General statement about natural quantities, tested and uncontradicted, describing what happens but not why.
Hypothesis: Initial attempt to explain the underlying cause of an observation; a testable, tentative explanation.
Experiment: Controlled observation designed to test a hypothesis by manipulating variables.
Theory: One or more confirmed hypotheses forming a comprehensive model that explains why natural phenomena occur and predicts behavior. Theories are extensively tested and widely accepted.
Atomic Theory and Its Laws
Early Ideas: Greek concept of "atomos" (indivisible particles) by Democritus and Leucippus, who postulated that matter is composed of finite, indivisible particles.
Law of Conservation of Mass (Antoine Lavoisier, 1789): In any closed system, during a chemical reaction or physical transformation, matter is neither created nor destroyed. The total mass of the reactants must equal the total mass of the products. For example, if 10 g of A react with 5 g of B, 15 g of product C will be formed.
Law of Definite Proportions (Joseph Proust, 1799): All samples of a given chemical compound, regardless of their source or method of preparation, contain the same elements in the exact same proportions by mass. For instance, water (H_{2}O) always consists of hydrogen and oxygen in an 8:1 mass ratio of oxygen to hydrogen.
Law of Multiple Proportions (John Dalton, 1804): When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other element are in a ratio of small whole numbers. For example, carbon and oxygen can form carbon monoxide (CO) and carbon dioxide (CO{2}). For a fixed mass of carbon, the mass of oxygen in CO{2} is exactly double the mass of oxygen in CO, a 2:1 ratio.
Dalton's Atomic Theory Postulates (1808):
Each element is composed of extremely small, indivisible particles called atoms.
All atoms of a given element are identical in mass and properties; atoms of different elements are distinct and possess different masses and properties.
Atoms are neither created nor destroyed, nor changed into atoms of another element, through chemical reactions (conservation of mass).
Compounds are formed when atoms of more than one element combine in a fixed ratio of small whole numbers (explains definite and multiple proportions).
The Discovery and Modern View of Atomic Structure
Cathode Rays (J.J. Thomson, 1897): Experiments with cathode ray tubes demonstrated that electricity was transmitted by negatively charged particles, electrons. Thomson measured the charge-to-mass ratio (e/m) of these particles as -1.76 imes 10^{8} C/g, which was independent of the cathode material, indicating electrons are fundamental to all matter.
Electron Charge and Mass (Robert Millikan, 1909): Millikan's oil drop experiment precisely determined the elementary charge of an electron as -1.60 imes 10^{-19} C. Combining this with Thomson's charge/mass ratio allowed for the calculation of the electron's mass: 9.10 imes 10^{-28} g.
Thomson's "Plum Pudding" Model: Proposed that the atom was a positively charged sphere with negatively charged electrons embedded within it, like plums in a pudding. This model was short-lived.
Radioactivity (Becquerel, Marie and Pierre Curie): The spontaneous emission of radiation by certain elements like Uranium and Radium revealed that atoms are not indivisible. Three types of radiation were identified:
\alpha particles: Positively charged (equivalent to a helium nucleus).
\beta particles: Negatively charged (high-energy electrons).
\gamma rays: Electromagnetic radiation with no charge or mass.
Rutherford's Gold Foil Experiment (Ernest Rutherford, 1911): By bombarding a thin gold foil with \\alpha\ particles, Rutherford observed that most particles passed straight through, but a small fraction were deflected at large angles, some even bouncing back. This led to the discovery of the nucleus: a tiny, incredibly dense, positively charged center within the atom, with electrons orbiting in a much larger, mostly empty space.
Subatomic Particles:
Electrons (e^{-}): Carry a fundamental negative charge (-1 or -1.602 imes 10^{-19} C) and have a mass of 9.10 imes 10^{-28} g (effectively zero relative to protons/neutrons). Located in the electron cloud outside the nucleus.
Protons (p^{+}): Carry a fundamental positive charge (+1 or +1.602 imes 10^{-19} C) and have a mass of 1.672 imes 10^{-24} g (approximately 1 atomic mass unit, amu). Located in the nucleus.
Neutrons (n$^{0}$): Carry no electrical charge (neutral) and have a mass of 1.674 imes 10^{-24} g (approximately 1 amu, slightly heavier than a proton). Located in the nucleus.
Atomic Number (Z): This is the characteristic number of protons in the nucleus of an atom. It defines the element and its chemical identity. For a neutral atom, the number of protons equals the number of electrons.
Mass Number (A): The sum of the number of protons (p^{+}) and neutrons (n^{0}) in an atom's nucleus (A = p^{+} + n^{0}). This value gives the approximate mass of an atom in atomic mass units (amu).
Isotopes: Atoms of the same element (meaning they have the same atomic number, Z) but possess different numbers of neutrons and therefore different mass numbers (A). For example, Carbon-12 and Carbon-14 are isotopes of carbon (Z=6), with 6 and 8 neutrons, respectively.
Ions (C$^{C}$): Atoms that have either lost or gained electrons, resulting in a net electrical charge.
Cations: Positively charged ions formed when an atom loses one or more electrons. For example, Na^{+} has lost one electron.
Anions: Negatively charged ions formed when an atom gains one or more electrons. For example, Cl^{-} has gained one electron. The charge (C) is determined by (number of protons - number of electrons).
Atomic Mass (Weight): The weighted average of the masses of all naturally occurring isotopes of an element. This is the value commonly found on the periodic table. It is calculated by:
\text{Atomic Mass} = \sum (\text{isotope molar mass} \times \text{natural fractional abundance})
Where natural fractional abundance is the percentage abundance expressed as a decimal (e.g., 10\% = 0.10). The notation for an isotope is \prescript{A}{Z}{}X^{C} where X is the element symbol, A is the mass number, Z is the atomic number, and C is the charge.
The Mole and Molar Mass
Mole Concept: A fundamental unit in chemistry that provides a bridge between the macroscopic world (grams) and the microscopic world (atoms, molecules). One mole (mol) represents a specific number of particles.
Avogadro's Number (N_{A}): Exactly 6.022 \times 10^{23} particles (atoms, molecules, ions, or formula units) per mole. This constant allows for conversion between the number of moles and the number of individual particles.
Molar Mass: The mass in grams of one mole of a substance (g/mol). It is a crucial conversion factor. For elements, the molar mass in grams per mole is numerically equal to the average atomic mass of that element in atomic mass units (amu) as found on the periodic table. For example, 12.011 amu for Carbon means its molar mass is 12.011 g/mol C. For compounds, molar mass is the sum of the atomic masses of all atoms in its chemical formula.