Atomic Structure & Periodicity: Foundations and Calculations

Atomic Structure & Periodicity

Introduction

The study of atomic structure and periodicity forms the basis of understanding chemical properties and behavior. This section covers the historical development of atomic theory, the structure of atoms, and the calculation of atomic mass using isotopic abundances.

Atomic Structure History

Democritus’ “Atomo”

• Democritus (460 B.C. – 370 B.C.) was a Greek philosopher who first proposed the concept of the atom.

• He described atoms as indivisible and indestructible particles.

• His ideas were philosophical and not based on experimental evidence or the scientific method.

Dalton’s Atomic Theory

• John Dalton (1766 – 1844) formulated the first modern atomic theory.

• Key points of Dalton’s theory:

  • All elements are composed of tiny, indivisible particles called atoms.

  • Atoms of the same element are identical; atoms of different elements are different.

  • Atoms of different elements combine in simple whole-number ratios to form compounds.

  • In chemical reactions, atoms are separated, rearranged, and combined, but atoms of one element are never changed into atoms of another element.

J.J. Thomson’s Discovery of the Electron

• In 1897, J.J. Thomson used a cathode ray tube to discover the electron, a negatively charged subatomic particle.

• This experiment demonstrated that atoms are divisible and contain smaller particles.

Robert Millikan’s Measurement of Electron Mass

• In 1916, Robert Millikan determined the mass of the electron to be $9.11 \times 10^{-28}$ g.

• The electron has a charge of -1 and a mass approximately 1/1840 that of a hydrogen atom.

Sub-Atomic Particles

Types, Properties, and Locations

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

Particle	Charge	Mass (g)	Location
Electron (e-)	-1	9.11 × 10-28	Electron cloud
Proton (p+)	+1	1.67 × 10-24	Nucleus
Neutron (n0)	0	1.67 × 10-24	Nucleus

Atomic Structure Models

Thomson’s “Plum Pudding Model”

• Thomson proposed that electrons were embedded in a positively charged "pudding," like plums in a pudding.

• This model was later disproven by further experiments.

Rutherford’s “Gold Foil Experiment”

• Rutherford fired alpha particles (helium nuclei) at a thin sheet of gold foil.

• Most particles passed through, but some were deflected, and very few were greatly deflected.

Rutherford’s Findings & Conclusions

• Most of the atom is empty space.

• The nucleus is small, dense, and positively charged.

• All positive charge and most mass are concentrated in the nucleus, which contains protons and neutrons.

• Electrons occupy most of the atom’s volume, distributed around the nucleus.

• This is known as the nuclear model of the atom.

Atomic Number and Mass Number

Atomic Number (Z)

• The atomic number ($Z$) is the number of protons in the nucleus of an atom.

• It uniquely identifies each element.

Element	# of protons	Atomic # (Z)
Carbon	6	6
Phosphorus	15	15
Gold	79	79

Mass Number (A)

• The mass number ($A$) is the sum of protons and neutrons in the nucleus.

• Formula: $A = p^+ + n^0$

Nuclide	p+	n0	e-	Mass #
Oxygen-18	8	10	8	18
Arsenic-75	33	42	33	75
Phosphorus-31	15	16	15	31

Atomic Symbols (Nuclear Symbols)

Notation

• The nuclear symbol contains the element symbol, mass number (superscript), and atomic number (subscript).

• Format: Mass numberAtomic numberElement symbol

• Example: 8035Br

Isotopes

Definition and Naming

• Isotopes are atoms of the same element (same number of protons) but different numbers of neutrons (different mass numbers).

• Isotopes are named by placing the mass number after the element name, e.g., carbon-12, carbon-14, uranium-235.

Isotopes in Nature

• Elements exist in nature as mixtures of isotopes.

• Examples:

  • Hydrogen: Protium (H-1), Deuterium (H-2), Tritium (H-3)

  • Helium: 3He, 4He

  • Lithium: 6Li, 7Li

Measuring Atomic Mass

Average Relative Atomic Mass

• The average relative atomic mass of an element is based on the abundance (percentage) of each isotope in nature.

• Atomic mass is not measured in grams due to the small size; instead, the atomic mass unit (amu) is used.

• 1 amu is defined as one-twelfth the mass of a carbon-12 atom: $1 \text{ amu} = 1.67377 \times 10^{-24}$ grams

Calculating Atomic Mass

• The atomic mass of an element is the weighted average of all its naturally occurring isotopes.

• Example for Carbon:

Isotope	Symbol	Composition of the nucleus	Abundance in nature
Carbon-12	12C	6 protons, 6 neutrons	98.89%
Carbon-13	13C	6 protons, 7 neutrons	1.11%
Carbon-14	14C	6 protons, 8 neutrons	<0.01%

• Calculation for Carbon atomic mass:

  • $\text{Atomic Mass (C)} = 12 \times 0.9889 + 13 \times 0.0111 + 14 \times 0.0001 = 12.011$

Practice Example: Chlorine

• Chlorine has two major isotopes:

  • One with 18 neutrons (mass number 35, 75.77% abundance)

  • One with 20 neutrons (mass number 37, 24.23% abundance)

  • Atomic number of chlorine is 17

• Calculation:

  • $\text{Atomic Mass (Cl)} = 35 \times 0.7577 + 37 \times 0.2423 = 35.48$

Summary Table: Key Atomic Structure Concepts

Concept	Definition	Example
Atomic Number (Z)	Number of protons in nucleus	Carbon: Z = 6
Mass Number (A)	Number of protons + neutrons	Phosphorus-31: A = 31
Isotope	Same Z, different A	Carbon-12, Carbon-14
Atomic Mass Unit (amu)	1/12 mass of carbon-12 atom	1 amu = $1.67377 \times 10^{-24}$ g

Additional info: These notes provide foundational knowledge for understanding atomic structure, isotopes, and atomic mass calculations, which are essential for further study in general chemistry and periodicity.