chapter 2 lec 1

Atoms, Molecules, and Ions

Early Ideas in Atomic Theory

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

• Democritus (5th century BC): Proposed that matter consists of indivisible particles called atomos.

• John Dalton (1808): Formulated the first modern atomic theory, defining atoms as indivisible units of elements.

Dalton's Atomic Theory:

1. Elements are composed of extremely small particles called atoms.

2. All atoms of a given element are identical in size, mass, and chemical properties; atoms of different elements differ.

3. Compounds are formed from atoms of more than one element, with ratios of atoms as integers or simple fractions.

4. Chemical reactions involve the rearrangement of atoms; atoms are neither created nor destroyed (Law of Conservation of Mass).

Laws of Chemical Combination

Experimental observations led to fundamental laws describing how elements combine to form compounds.

• Law of Definite Proportions (Joseph Proust, 1799): A given compound always contains the same elements in the same proportion by mass.

• 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.

Examples:

• CO and CO2: The ratio of oxygen in CO to CO2 is 1:2.

• SO2 and SO3: The ratio of oxygen in SO2 to SO3 is 2:3.

Atomic Structure and Symbolism

Atoms are composed of subatomic particles: electrons, protons, and neutrons. The discovery of these particles revolutionized our understanding of atomic structure.

• Electrons: Negatively charged particles discovered through cathode ray experiments. J.J. Thomson determined the charge-to-mass ratio; R.A. Millikan measured the charge, allowing calculation of the electron's mass.

• Protons: Positively charged particles located in the nucleus. Their mass is about 1840 times that of an electron.

• Neutrons: Electrically neutral particles in the nucleus, with a mass slightly greater than that of a proton. Discovered by James Chadwick in 1932.

Key Equations:

• Electron charge-to-mass ratio: $\frac{\text{charge}}{\text{mass}} = -1.76 \times 10^8 \ \text{C/g}$

• Electron charge: $-1.6022 \times 10^{-19} \ \text{C}$

• Electron mass: $\frac{-1.6022 \times 10^{-19} \ \text{C}}{-1.76 \times 10^8 \ \text{C/g}} = 9.10 \times 10^{-28} \ \text{g}$

Radioactivity

Some elements emit radiation spontaneously, a phenomenon called radioactivity. There are three main types of radioactive decay:

• Alpha (α) particles: Positively charged, deflected by positive plates.

• Beta (β) rays: Electrons, deflected by negative plates.

• Gamma (γ) rays: High-energy, uncharged, not deflected by electric or magnetic fields.

Rutherford's Nuclear Model

Ernest Rutherford's gold foil experiment demonstrated that atoms have a dense, positively charged nucleus surrounded by electrons. Most of the atom is empty space.

• Protons: Located in the nucleus, charge +1, mass $1.67262 \times 10^{-24}$ g.

• Nucleus: Contains nearly all the atom's mass, but occupies a tiny fraction of its volume.

Atomic and Nuclear Dimensions:

• Atomic radius: ~100 pm (1 Å = 100 pm)

• Nuclear radius: ~5 x 10-3 pm

Subatomic Particle Properties

Particle	Mass (g)	Charge (C)	Charge	Mass (amu)
Electron	9.10939 x 10-28	-1.6022 x 10-19	-1	5.4858 x 10-4 (≈0)
Proton	1.67262 x 10-24	+1.6022 x 10-19	+1	1.00728 (≈1)
Neutron	1.67493 x 10-24	0	0	1.00867 (≈1)

Isotopes

Atoms of the same element with different numbers of neutrons are called isotopes. Isotopes have identical chemical properties but different masses.

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

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

Examples:

• Hydrogen isotopes: $^1$H (protium, 0 neutrons), $^2$H (deuterium, 1 neutron), $^3$H (tritium, 2 neutrons).

• Chlorine isotopes: $^{35}$Cl (17 protons, 18 neutrons, 77.4%), $^{37}$Cl (17 protons, 20 neutrons, 22.6%).

• Neon isotopes: $^{20}$Ne (90.4838%), $^{21}$Ne (0.2696%), $^{22}$Ne (9.2465%).

Calculating Average Atomic Mass:

• Weighted average: $\text{Average mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example for Neon: $\text{Average mass} = (19.9924 \times 0.904838) + (20.9940 \times 0.002696) + (21.9914 \times 0.092465) = 20.1799$ amu

Example for Silver:

• Let $X$ = abundance of $^{107}$Ag, $(1-X)$ = abundance of $^{109}$Ag

• Equation: $106.90X + 108.90(1-X) = 107.87$

• Solve for $X$: $X = \frac{-1.03}{-2.00} = 0.515$ (51.5%)

• Abundances: $^{107}$Ag = 51.5%, $^{109}$Ag = 48.5%

The Periodic Table

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

• Periods: Horizontal rows

• Groups: Vertical columns

Element Categories:

• Metals: Good conductors of heat and electricity

• Non-metals: Poor conductors

• Metalloids: Properties intermediate between metals and non-metals

Special Groups:

• Alkali metals (Group 1A): Li, Na, K, Rb, Cs, Fr; form +1 ions

• Alkaline earth metals (Group 2A): Be, Mg, Ca, Sr, Ba, Ra; form +2 ions

• Halogens (Group 7A): F, Cl, Br, I, At; form -1 ions

• Chalcogens (Group 6A): O, S, Se, Te; form -2 ions

• Noble gases (Group 8A): He, Ne, Ar, Kr, Xe, Rn; generally unreactive

• Transition metals: Groups 3B–8B, 1B, 2B; can form multiple ions

• Lanthanides: Elements 58–71

• Actinides: Elements 90–107

Molecules and Ions

Most elements exist as molecules or ions, except for noble gases, which are monatomic. Molecules are aggregates of atoms held together by chemical bonds.

• Molecule: Aggregate of at least two atoms in a definite arrangement.

• Diatomic molecules: Two atoms, either same (homonuclear: H2, N2, O2, F2, Cl2, Br2, I2) or different (heteronuclear: CO, HF, HCl, HBr, HI).

• Polyatomic molecules: More than two atoms (SO2, SO3, C8H14O3).

Chemical Formulas

Chemical formulas indicate the elements present and their ratios in molecules and ionic compounds.

• Molecular formula: Shows exact number of atoms (e.g., H2O vs. H2O2).

• Structural formula: Shows arrangement of atoms (e.g., H-O-H for water).

• Ball-and-stick and space-filling models: Visualize 3D structure.

Bonding: Hydrogen forms only one bond; structural formulas must reflect correct bonding.

Example: Water is H-O-H, not H-H-O; hydrogen peroxide is H-O-O-H.

Bond angle in water: 109.45° (not linear).

Naming Molecular Compounds

Binary molecular compounds consist of two non-metallic elements. Naming conventions indicate the elements and their ratios.

• First element is named as is; second element's name ends with "-ide".

• Prefixes indicate the number of atoms: mono-, di-, tri-, tetra-, penta-, hexa-, etc.

Examples:

• CO: carbon monoxide

• CO2: carbon dioxide

• SO2: sulfur dioxide

• SO3: sulfur trioxide

• P2O3: diphosphorus trioxide

• P2O5: diphosphorus pentoxide

• N2O4: dinitrogen tetroxide

• PCl3: phosphorus trichloride

• P4S10: tetraphosphorus decasulfide

Special Compounds: Some compounds have common names:

• B2H6: diborane

• SiH4: silane

• NH3: ammonia

• PH3: phosphine

• H2O: water

• H2S: hydrogen sulfide

Acids

Acids are substances that produce hydrogen ions (H+) when dissolved in water. The name changes when the compound is dissolved in water.

• HCl (g): hydrogen chloride; in water: hydrochloric acid

• HBr: hydrobromic acid

• HI: hydroiodic acid

• HF: hydrofluoric acid

• HCN: hydrocyanic acid