Atoms, Molecules, and Ions

What is Matter?

Matter is anything that has mass and occupies space. According to atomic theory, matter is composed of discrete units called atoms. Atoms are the smallest identifiable unit of an element and are much smaller than the wavelength of visible light, making them impossible to see optically but possible to image with advanced techniques.

• Atom: The fundamental unit of chemical elements.

• Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Atoms combine to form molecules and compounds.

• Example: Imaging techniques such as atomic force microscopy can visualize atoms in a molecule.

Dalton’s Atomic Theory of Matter

John Dalton proposed a foundational theory describing the nature of atoms and their role in chemical reactions.

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

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

• Postulate 3: Atoms combine in simple, whole-number ratios to form molecules or compounds.

• Postulate 4: In chemical reactions, atoms are rearranged but not changed into atoms of another element.

• Example: Water (H2O) is formed by combining two hydrogen atoms and one oxygen atom.

Discovery of Subatomic Particles

Subatomic Structure

Dalton originally viewed the atom as the smallest particle possible, but later discoveries revealed that atoms are made up of smaller particles.

• Proton (p+): Positively charged particle found in the nucleus.

• Neutron (n0): Neutral particle found in the nucleus.

• Electron (e-): Negatively charged particle found outside the nucleus.

Notes on Charge

Electric charge is a fundamental property of subatomic particles.

• There are two kinds of charge: positive (+) and negative (−).

• Opposite charges attract; like charges repel.

• Charges of equal magnitude but opposite sign sum to zero when combined.

• Example:

Discovery of the Electron

J.J. Thomson (1900-1909) discovered that cathode rays are made of tiny, negatively charged particles called electrons. Robert Millikan (1909) determined the charge and mass of the electron.

• Electrons have the same amount of charge as a hydrogen ion, but with opposite sign.

• The mass of an electron is approximately 2000 times lighter than a hydrogen ion.

• Electrons are fundamental components of atoms.

• Key Point: The atom is not unbreakable!

Early Atomic Models

The "plum pudding" model, proposed by Thomson, depicted the atom as a positive sphere of matter with negative electrons embedded within it.

• Positive charge is spread throughout the sphere.

• Electrons are scattered within the positive matrix.

• Example: The plum pudding model was later replaced by the nuclear model.

Summary Table: Comparison of Subatomic Particles

Atomic Mass and Isotopes

Atomic Mass Unit (amu)

Atoms have extremely small masses, so the atomic mass unit (amu) is used for convenience.

• 1 amu = g

• Atomic masses are compared to carbon-12 (C), which is defined as exactly 12 amu.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons.

• Isotope Notation:

• Example: , ,

Isotope Abundance and Atomic Weight

The atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.

• Formula:

• Example: Neon has three isotopes: , , .

The Periodic Table

Periodic Law and Classification

The periodic table organizes elements by increasing atomic number, revealing recurring patterns in their properties.

• Elements with similar properties are grouped in columns (groups or families).

• Periodic law: Properties of elements recur periodically when arranged by atomic number.

• Metals, nonmetals, and metalloids are classified based on their properties.

Important Groups

• Alkali metals (Group 1A): Li, Na, K, Rb, Cs, Fr – highly reactive, form +1 cations.

• Alkaline earth metals (Group 2A): Be, Mg, Ca, Sr, Ba – less reactive, form +2 cations.

• Halogens (Group 7A): F, Cl, Br, I – reactive nonmetals, form -1 anions.

• Noble gases (Group 8A): He, Ne, Ar, Kr, Xe, Rn – inert gases.

Chemical Formulas and Nomenclature

Chemical Formulas

Chemical formulas use subscripts to indicate the number of atoms of each element in a molecule or compound.

• Empirical formula: Shows the simplest whole-number ratio of atoms.

• Molecular formula: Shows the exact number of atoms of each element.

• Structural formula: Shows the connectivity of atoms.

• Example: Glucose: Empirical formula CH2O, Molecular formula C6H12O6

Molecular and Ionic Compounds

• Molecular compounds: Composed of molecules, usually only nonmetals.

• Ionic compounds: Composed of ions, typically formed between metals and nonmetals.

Ions

Atoms gain or lose electrons to form ions.

• Cation: Positively charged ion (loss of electrons), usually formed by metals.

• Anion: Negatively charged ion (gain of electrons), usually formed by nonmetals.

• Example: Na loses one electron to form Na+; Cl gains one electron to form Cl-.

Polyatomic Ions

Polyatomic ions are groups of atoms that carry a charge.

Writing Formulas for Ionic Compounds

Ionic compounds are electrically neutral. The charge on the cation becomes the subscript for the anion and vice versa. Subscripts are reduced to the lowest whole-number ratio.

• Example: Magnesium chloride: Mg2+ and Cl- combine to form MgCl2.

• Formula:

Naming Ionic Compounds

• Name the cation first; if it has variable charge, indicate the charge with Roman numerals.

• If the anion is an element, change its ending to -ide.

• If the anion is polyatomic, use its name directly.

• Example: Fe2O3 is iron(III) oxide.

• Example: KClO4 is potassium perchlorate.

Oxyanion Nomenclature

Oxyanions are polyatomic ions containing oxygen.

• Fewer oxygens: -ite (e.g., NO2- is nitrite)

• More oxygens: -ate (e.g., NO3- is nitrate)

• Prefix hypo- for least oxygens, per- for most oxygens (e.g., ClO- is hypochlorite, ClO4- is perchlorate)

Naming Binary Molecular Compounds

• Formed between two nonmetals.

• Name the element farther left (or lower) in the periodic table first.

• Use prefixes to denote the number of atoms (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-).

• Never use the prefix mono- for the first element.

• Change the ending of the second element to -ide.

• Example: SF6 is sulfur hexafluoride; P4O10 is tetraphosphorus decoxide.

Naming Acids

• Acids contain H+ cation and an anion.

• If the anion ends in -ide, use the prefix hydro- and suffix -ic acid (e.g., HCl: hydrochloric acid).

• If the anion ends in -ite, use the suffix -ous acid (e.g., HClO: hypochlorous acid).

• If the anion ends in -ate, use the suffix -ic acid (e.g., HClO4: perchloric acid).

Summary Table: Prefixes for Binary Molecular Compounds

Key Equations and Concepts

• Atomic Number:

• Mass Number:

• Charge of Ion:

• Atomic Weight:

Additional info: Some content was inferred and expanded for clarity and completeness, including standard definitions, examples, and tables for subatomic particles, isotopes, and nomenclature.