Atoms, Molecules, and Ions: Foundations of General Chemistry

Atoms, Molecules, and Ions

Historical Development of Atomic Theory

The concept of the atom has evolved over centuries, beginning with philosophical ideas and culminating in scientific theories supported by experimental evidence.

• Democritus (460–370 BC): Proposed that matter is composed of tiny, indivisible particles called atoms.

• Plato & Aristotle: Argued against the existence of ultimately indivisible particles, believing matter was continuous.

• John Dalton (17th century): Developed the first modern atomic theory based on experimental evidence.

Dalton’s Atomic Theory

John Dalton formulated a set of postulates that laid the foundation for modern chemistry. These postulates explain the nature of atoms and how they combine to form compounds.

• Postulate 1: Each element is composed of extremely small particles called atoms.

• Postulate 2: All atoms of a given element are identical in mass and other properties, but atoms of different elements differ in these properties.

• Postulate 3: Atoms of one element cannot be changed into atoms of another element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

• Postulate 4: Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

Example: Water (H2O) always consists of two hydrogen atoms and one oxygen atom.

Key Laws Based on Dalton’s Theory

• Law of Conservation of Mass: The total mass of substances present after a chemical reaction is the same as before the reaction. Equation: $\text{Total mass before reaction} = \text{Total mass after reaction}$

• Law of Constant Composition: In a given compound, the relative numbers and kinds of atoms are constant. Example: Carbon dioxide (CO2) always contains one carbon atom and two oxygen atoms.

• Law of Multiple Proportions: If 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.

Subatomic Particles

Atoms are composed of smaller particles known as subatomic particles: electrons, protons, and neutrons.

• Electron: Negatively charged particle discovered by J.J. Thomson in 1897 using cathode ray experiments. Charge/mass ratio: $1.76 \times 10^8$ C/g

• Proton: Positively charged particle found in the nucleus. Charge: $+1.602 \times 10^{-19}$ C

• Neutron: Neutral particle found in the nucleus. Discovered by: James Chadwick (1932)

Comparison of Subatomic Particles

Particle	Charge	Relative Mass
Proton	+1	1
Neutron	0	1
Electron	-1	~0 (1/1836 of a proton)

Atomic Structure and Isotopes

Atoms are characterized by their atomic number (number of protons) and mass number (number of protons plus neutrons).

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

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

• Isotopes: Atoms of the same element with different numbers of neutrons and thus different masses.

Example: Carbon has three isotopes: $^{12}\text{C}$, $^{13}\text{C}$, and $^{14}\text{C}$.

Atomic Mass and Atomic Weight

Atomic mass is the mass of a single atom, while atomic weight is the weighted average mass of all naturally occurring isotopes of an element.

• Atomic Mass Unit (amu): $1\ \text{amu} = 1.66054 \times 10^{-24}\ \text{g}$

• Atomic Weight Calculation: $\text{Atomic Weight} = \sum [(\text{isotope mass}) \times (\text{fractional natural abundance})]$

The Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties together.

• Periods: Horizontal rows

• Groups: Vertical columns; elements in the same group have similar properties

• Metals: Left side; lustrous, conductive, malleable

• Nonmetals: Right side; varied states, poor conductors

• Metalloids: Border between metals and nonmetals; properties intermediate between metals and nonmetals

Group Name	Elements
Alkali Metals	Li, Na, K, Rb, Cs, Fr
Alkaline Earth Metals	Be, Mg, Ca, Sr, Ba, Ra
Chalcogens	O, S, Se, Te, Po
Halogens	F, Cl, Br, I, At
Noble Gases	He, Ne, Ar, Kr, Xe, Rn

Chemical Formulas and Molecular Compounds

Chemical formulas represent the composition of compounds using element symbols and subscripts.

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

• Molecular Formula: Shows the exact number of atoms of each element in a molecule.

• Structural Formula: Shows the arrangement of atoms within the molecule.

Example: Glucose: Empirical formula CH2O, Molecular formula C6H12O6

Diatomic Molecules

Some elements exist naturally as molecules containing two atoms.

• Hydrogen (H2)

• Nitrogen (N2)

• Oxygen (O2)

• Fluorine (F2)

• Chlorine (Cl2)

• Bromine (Br2)

• Iodine (I2)

Ions and Ionic Compounds

Ions are formed when atoms gain or lose electrons. Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Cation: Positively charged ion (loss of electrons)

• Anion: Negatively charged ion (gain of electrons)

• Ionic Compound: Formed between metals and nonmetals; only empirical formulas are written

Example: Sodium chloride (NaCl) is formed from Na+ and Cl-.

Naming Inorganic Compounds

Chemical nomenclature provides systematic rules for naming compounds.

• Ionic Compounds: Name the cation first, then the anion. If the cation has multiple charges, indicate the charge with Roman numerals.

• Acids: If the anion ends in -ide, use the prefix hydro- and the suffix -ic acid. If the anion ends in -ate, use -ic acid; if -ite, use -ous acid.

• Binary Molecular Compounds: Use prefixes to indicate the number of atoms (mono-, di-, tri-, etc.). The element farther left or lower in the periodic table is named first.

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

Example: CO2 is carbon dioxide, N2O5 is dinitrogen pentoxide.

Organic Compounds and Isomerism

Organic chemistry focuses on compounds containing carbon. The simplest hydrocarbons are alkanes, named according to the number of carbon atoms.

• Alkanes: Methane (CH4), Ethane (C2H6), Propane (C3H8), etc.

• Alcohols: Replace a hydrogen in an alkane with an -OH group; name ends in -ol (e.g., ethanol).

• Isomers: Molecules with the same molecular formula but different structural arrangements.

Example: 1-propanol and 2-propanol both have the formula C3H8O but differ in the position of the -OH group.

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