Chap 2 Atoms, Atomic Structure, and the Periodic Table: Foundations of Modern Atomic Theory

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Modern Atomic Theory

Foundations of Atomic Theory

The modern atomic theory describes the fundamental nature of matter and the structure of atoms. It builds upon Dalton's atomic theory and incorporates discoveries about subatomic particles and atomic structure.

Atoms are the smallest particles of an element that retain the chemical properties of that element and participate in chemical reactions.
All atoms of a given element are alike, meaning they have the same properties and structure. Atoms of different elements are different.
Compounds are combinations of atoms of more than one element, always in specific whole-number ratios.
Chemical reactions involve the rearrangement of atoms; atoms themselves are unchanged and cannot be created or destroyed in ordinary chemical reactions.

Atoms and Their Electronic Structure

Subatomic Particles

Atoms are composed of three main subatomic particles, each with distinct properties and locations within the atom.

Proton (p): A positively charged particle located in the nucleus.
Neutron (n): An electrically neutral particle also located in the nucleus.
Electron (e-): A negatively charged particle that moves through a large volume of space around the nucleus.

Key Interactions:

Like charges repel each other (proton-proton or electron-electron).
Opposite charges attract (proton-electron).

Structure of the Atom

Most of the mass of an atom is concentrated in its nucleus, which contains protons and neutrons. Electrons occupy the space around the nucleus and are attracted to the positively charged nucleus.

Atomic mass unit (amu): Standard unit for atomic mass; .
Mass of 1 proton ≈ 1 amu; mass of 1 neutron ≈ 1 amu.

Comparison of Subatomic Particles

The following table summarizes the properties of the three main subatomic particles:

Particle	Symbol	Mass (g)	Mass (amu)	Charge
Proton	p	1.672622 × 10-24	1.007276	+1
Neutron	n	1.674927 × 10-24	1.008665	0
Electron	e-	9.109328 × 10-28	5.485799 × 10-4	-1

Atomic Number and Mass Number

Atomic number (Z): The number of protons in the nucleus of an atom. Determines the identity of the element. For example, for hydrogen.
Mass number (A): The sum of the number of protons and neutrons in an atom.

Each element on the periodic table is represented by its symbol, atomic number, and atomic mass.

Isotopes

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

Example: Hydrogen has three isotopes—protium (), deuterium (), and tritium ().
Isotopes are represented as , where A is the mass number, Z is the atomic number, and X is the element symbol.
The atomic weights listed on the periodic table reflect the weighted average of all naturally occurring isotopes of each element.

Arrangement of the Periodic Table

Organization of the Periodic Table

The periodic table arranges elements by increasing atomic number into rows and columns based on recurring chemical properties.

Periods: Seven horizontal rows.
Groups (families): Eighteen vertical columns. Elements in the same group have similar chemical properties.

Main Categories of Groups:

Main Groups: Groups 1-2 (far left) and 13-18 (far right).
Transition Metals: Groups 3-12 (center block).
Inner Transition Metals: Lanthanides and actinides (bottom two rows).

Properties of Selected Groups

Group 1 (Alkali Metals): Shiny, low-melting, highly reactive metals. React vigorously with water. Not found in pure form in nature.
Group 2 (Alkaline Earth Metals): Reactive metals, less reactive than alkali metals. React with water to form basic solutions. Not found in pure form in nature.
Group 17 (Halogens): Mostly nonmetals, exist in all three physical states (solid, liquid, gas). Colorful and very reactive. Found in nature only in compounds.
Group 18 (Noble Gases): Nonmetals, colorless gases, very unreactive. Found in pure form in nature. Have stable electron configurations.

Trends in Atomic Radius

Atomic radius is defined as one-half the distance between the nuclei of two adjacent atoms. Trends in the periodic table:

Atomic radius decreases from left to right across a period.
Atomic radius increases from top to bottom down a group.

Electronic Structure of Atoms

Quantum Mechanical Model

The quantum mechanical model describes how electrons are arranged around the nucleus in regions of space called orbitals. This arrangement determines the chemical reactivity of atoms.

Electrons are restricted to certain energy levels (quantized energies).
Electrons are grouped into shells (numbered 1, 2, 3, ...), with each shell further divided into subshells (s, p, d, f).
Each subshell contains a specific number of orbitals (s: 1, p: 3, d: 5, f: 7).
Each orbital can hold up to two electrons with opposite spins.

Summary Table: Shells, Subshells, and Orbitals

Shell Number (n)	Subshells	Number of Orbitals	Maximum Electrons
1	s	1	2
2	s, p	1, 3	8
3	s, p, d	1, 3, 5	18
4	s, p, d, f	1, 3, 5, 7	32

Shapes of Orbitals

s orbitals: Spherical shape.
p orbitals: Dumbbell-shaped, oriented in three directions (x, y, z).

Order of Orbital Energy Levels

Electrons fill orbitals in order of increasing energy, following the Aufbau principle. The general order is:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s, etc.

Electron Configuration Rules

Aufbau Principle: Electrons occupy the lowest energy orbitals available.
Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins.
Hund's Rule: Orbitals of equal energy are each occupied by one electron before any is occupied by a second electron.

Electron Configuration Notation

Electron configurations are written by listing the shell number and subshell letter, with the number of electrons in each subshell as a superscript.
Example for carbon:
Noble gas shorthand: Use the symbol of the previous noble gas in brackets to abbreviate inner-shell electrons. Example: for carbon.

Valence Electrons and Lewis Dot Symbols

Valence shell: The outermost, highest energy shell of an atom.
Valence electrons: Electrons in the outermost shell; these determine the chemical properties and reactivity of the element.
Lewis dot symbol: An atomic symbol surrounded by dots representing the number of valence electrons.

Elements in the same group have the same number of valence electrons, explaining their similar chemical behavior. Noble gases have completely filled valence shells, making them especially stable and unreactive.

Example Table: Valence Electrons by Group

Group	Number of Valence Electrons	Example Element
1 (Alkali Metals)	1	Na
2 (Alkaline Earth Metals)	2	Mg
17 (Halogens)	7	Cl
18 (Noble Gases)	8	Ne

Summary: Understanding atomic structure, electron configuration, and the organization of the periodic table is essential for predicting the chemical behavior of elements and compounds.