Atoms, Molecules, and Ions: Foundations of Atomic Theory and the Periodic Table

Chapter 2: Atoms, Molecules, and Ions

Announcements & Reminders

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• Bring a calculator and periodic table to class for problem-solving activities.

Atomic Theory of Matter

Historical Development

The concept of the atom has evolved from philosophical ideas to a scientific theory based on experimental evidence.

• Democritus (460–370 BC): Proposed that matter is composed of indivisible particles called "atomos."

• Dalton (1803–1807): Formulated the atomic theory, stating that elements are made of tiny, indivisible particles (atoms) and that atoms of the same element are identical.

Key Laws Derived from Atomic Theory

• Law of Constant Composition: Compounds have a definite composition; the ratio of atoms of each element is always the same. (Joseph Proust)

• Law of Conservation of Mass: Mass is conserved in chemical reactions; total mass before and after a reaction is unchanged. (Antoine Lavoisier)

• 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 small whole number ratios. (John Dalton)

Discovery of Subatomic Particles

Subatomic Structure

Atoms are not indivisible; they are composed of smaller particles:

• Electrons: Discovered via cathode ray experiments; negatively charged.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles also located in the nucleus.

• Radioactivity: Showed that atoms can emit particles, further supporting the existence of subatomic structure.

Modern Atomic Model

• Atoms consist of a dense nucleus (protons and neutrons) surrounded by a cloud of electrons.

• Most of the atom's volume is empty space.

• Typical atomic size: 1–5 Å (100–500 pm).

Subatomic Particles: Properties and Comparison

• Proton (p+): Charge +1, mass ≈ 1 amu.

• Neutron (n0): Charge 0, mass ≈ 1 amu.

• Electron (e-): Charge –1, mass ≈ 5.486 × 10–4 amu (much smaller than protons/neutrons).

Particle	Charge	Mass (amu)
Proton	+1	1.0073
Neutron	0	1.0087
Electron	–1	5.486 × 10–4

Atomic Numbers, Mass Numbers, and Isotopes

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

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

• Element notation: $^{A}_{Z}X$ (e.g., $^{12}_{6}C$ for carbon-12).

• For neutral atoms: Number of electrons = Number of protons.

• Neutrons: $n^0 = A - Z$

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons (and thus different masses), but the same number of protons.

Symbol	Number of Protons	Number of Electrons	Number of Neutrons
$^{11}C$	6	6	5
$^{12}C$	6	6	6
$^{13}C$	6	6	7
$^{14}C$	6	6	8

Examples

• $^{138}Ba$: 56 protons, 82 neutrons, 56 electrons

• $^{59}Ni$: 28 protons, 31 neutrons, 28 electrons

• $^{197}Au$: 79 protons, 118 neutrons, 79 electrons

Atomic Mass Unit (amu) and Atomic Weight

Definition and Use

• 1 amu = $1.66054 \times 10^{-24}$ g

• 1 g = $6.02214 \times 10^{23}$ amu

• By definition, $^{12}C$ has a mass of exactly 12 amu.

• Atomic weights on the periodic table are average values based on isotopic abundance.

Calculating Atomic Weight

• Atomic Weight (AW) is the weighted average of all naturally occurring isotopes of an element.

Formula:

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

Example: For carbon:

• 98.93% $^{12}C$ (12 amu), 1.07% $^{13}C$ (13.00335 amu)

$AW = (0.9893)(12) + (0.0107)(13.00335) = 12.01$ amu

Measurement of Atomic and Molecular Weights

• Mass spectrometry is used to measure atomic and molecular weights and determine isotopic abundances.

The Periodic Table

Organization and Structure

• The periodic table arranges elements by increasing atomic number (Z).

• Rows are called periods; columns are called groups.

• Elements in the same group have similar chemical properties.

Reading the Periodic Table

• Each box lists the atomic number (above the symbol), atomic symbol, and atomic weight (below the symbol).

• Example: 19 (atomic number), K (symbol), 39.0983 (atomic weight)

Special Groups

Group	Name	Elements
1A	Alkali metals	Li, Na, K, Rb, Cs, Fr
2A	Alkaline earth metals	Be, Mg, Ca, Sr, Ba, Ra
6A	Chalcogens	O, S, Se, Te, Po
7A	Halogens	F, Cl, Br, I, At
8A	Noble gases	He, Ne, Ar, Kr, Xe, Rn

Metals, Nonmetals, and Metalloids

• Metals: Mostly on the left side; majority of elements.

• Nonmetals: Mostly on the right side (including H).

• Metalloids: Elements along the step-like line (except Al, Po, At); have properties intermediate between metals and nonmetals.

Example Application

Elements in the same group (vertical column) show the greatest similarity in chemical and physical properties. For example, Na and K (both in group 1A) are more similar to each other than to elements in other groups.

Additional info: This summary covers the foundational concepts of atomic theory, subatomic particles, isotopes, atomic mass, and the organization of the periodic table, which are essential for understanding chemical behavior and periodic trends.