Module 1: Atomic Theory and the Periodic Table

Atomic Theory of Matter

• Conceptual Origins: The term "atomos" was suggested as the fundamental component of all matter in the universe. The word is derived from Greek where "a" means "not" and "tomos" means "cut," implying an indivisible unit.

• Modern Atomic Theory: Dalton, an English meteorologist and chemist, formulated the first modern description of an atom. He characterized it as the fundamental building block of all chemical structures.

• Law of Conservation of Mass: Formulated by Antoine Lavoisier, this law states that the atoms of an object cannot be created or destroyed. However, they can be moved around and changed into different particles.

  • Combustion Example: When wood is burnt in oxygen, it yields ashes, carbon dioxide gas ($CO_2$), and water vapor ($H_2O$). The total mass of the wood and oxygen consumed is equal to the total mass of the ashes, carbon dioxide, and water produced.

  • Methane Combustion: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

• Law of Definite Proportions: If a compound is broken down into its constituent elements, the masses of those constituents will always have the same proportions, regardless of the quantity of the substance or its original source.

• Law of Multiple Proportions: If two elements form more than one compound between them, the masses of one element that combine with a fixed mass of the second element form ratios of small integers.

  • Nitrogen and Oxygen Example: Nitrogen ($N$) and Oxygen ($O$) can form various compounds including $NO$, $NO_2$, $N_2O$, $N_2O_4$, and $N_2O_5$.

  • Ratio of molar masses ($N:O$):

    • $NO$: $14:16$ (Ratio: $2$)

    • $NO_2$: $14:32$ (Ratio: $4$)

    • $N_2O$: $28:16$ (Ratio: $1$)

    • $N_2O_4$: $28:64$ (Ratio: $3$)

    • $N_2O_5$: $28:80$ (Ratio: $5$)

Dalton's Postulates and Their Limitations

• Postulates of Dalton's Theory:

  • All matter consists of indivisible particles called atoms.

  • Atoms of the same element are similar in shape and mass but differ from the atoms of other elements.

  • Atoms cannot be created or destroyed.

  • Atoms of different elements may combine with each other in fixed, simple, whole-number ratios to form compound atoms.

  • Atoms of the same element can combine in more than one ratio to form two or more compounds.

  • The atom is the smallest unit of matter capable of taking part in a chemical reaction.

• Drawbacks on Dalton's Postulates:

  • Subdivision: An atom is not indivisible; it can be further subdivided into protons, neutrons, and electrons.

  • Isotopes: Atoms of some elements vary in their masses and densities. These variations are known as isotopes.

  • Isobars: Atoms of different elements may share certain physical traits; for example, Argon ($Ar$) and Calcium ($Ca$) share the same atomic mass.

  • Allotropes: The theory fails to explain the existence of allotropes, such as the differences in properties between charcoal, graphite, and diamond, all of which are forms of carbon.

Sub-atomic Particles and Atomic Identification

• Primary Sub-atomic Particles: The three main components of an atom are the Electron, Proton, and Neutron.

• Atomic Identification Numbers:

  • Atomic Number ($Z$): The number of protons located in the nucleus of each atom of the element.

  • Mass Number ($A$): The total number of protons and neutrons in an atom of the element.

• Chemical Notation: Elements are represented by a chemical symbol with the mass number ($A$) as a superscript and the atomic number ($Z$) as a subscript to the left of the symbol.

  • Hydrogen: ${}_1^1 \text{H}$

  • Carbon: ${}_6^{12} \text{C}$

  • Fluorine: ${}_9^{19} \text{F}$

Quantum Numbers

Quantum numbers are used to completely describe the movement and trajectories of each electron within an atom. They designate specific shells, subshells, orbitals, and spins.

• Principal Quantum Number ($n$):

  • Designates the principal electron shell.

  • Describes the energy level and the most probable distance of the electron from the nucleus.

  • Valid values are any positive integer starting at $1$ (where $n=1$ designates the first principal shell, also called the ground state or lowest energy state).

  • A higher value of "$n$" corresponds to a higher average distance from the nucleus and a higher energy level. For example, $n=3$ has a greater energy level than $n=2$.

• Orbital Angular Momentum Quantum Number ($l$):

  • Identifies the subshell.

  • If $l=2$, the subshell is named "d".

  • If $l=3$, the subshell is named "f".

• Magnetic Quantum Number ($m_l$):

  • Describes the orientation of the orbital.

  • If $l=3$, the possible values for $m_l$ range from $-3$ to $+3$ (i.e., $-3, -2, -1, 0, 1, 2, 3$).

• Electron Spin Quantum Number ($m_s$):

  • Describes the direction of the electron's spin.

Electron Configuration and Governing Principles

• Definition: Electron configuration is the arrangement of atoms in energy levels around an atomic nucleus. In the shell atomic model, it is expressed by indicating the number of electrons in each shell, beginning with the first.

• Governing Rules:

  • Pauli Exclusion Principle: In an atom or molecule, no two electrons can have the same four electronic quantum numbers.

  • Aufbau Principle: In the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher energy levels.

  • Hund's Rule: Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied. Furthermore, all electrons in singly occupied orbitals possess the same spin.

• Shorthand Configuration: An example is the shorthand electron configuration of Sulfur ($S$), which utilizes the preceding noble gas to simplify the notation.

The Periodic Table of Elements

• Historical Development: Mendeleev organized the table into columns, assigning them numerals and the letters "A" for some and "B" for others.

  • Group A Elements: Known as main-group elements (Groups $1A$ through $8A$).

  • Group B Elements: Called transition elements.

  • Inner Transition Elements: Elements $58 - 71$ (Lanthanides) and $90 - 103$ (Actinides) are shown separately at the bottom and are not included in the main body of the table.

• Classification of Elements:

  • Metals: Generally solids at room temperature (except Mercury ($Hg$), which is liquid), shiny, good conductors of electricity, ductile, and malleable. Chemically, they tend to give up electrons.

    • Alloys: Solutions of metals.

      • Brass: $Cu$ and $Zn$

      • Bronze: $Cu$ and $Sn$

      • Pewter: $Sn$, $Sb$, and $Pb$

  • Non-metals: Located on the right side of the Periodic Table (with the exception of Hydrogen). They generally do not conduct electricity (graphite is an exception). At room temperature, Phosphorus ($P$) and Iodine ($I$) are solids, Bromine ($Br$) is a liquid, and Group $8A$ contains gases.

  • Metalloids: Six elements (Boron, Silicon, Germanium, Arsenic, Antimony, and Tellurium) that share properties of both metals and non-metals. Silicon is a semiconductor, meaning it conducts electricity only under specific applied voltages.

Periodicity and Trends

• Periods: Each horizontal row is called a period. Each row represents the number of energy levels present in an atom of the element.

• Periodic Trends:

  • Atomic Size/Radius: Increases down a column as the valence shell moves farther from the nucleus due to a larger principal quantum number. Atomic size decreases across a period from left to right.

  • Ionization Energy (IE): The amount of energy required to remove an electron from an atom in the gas phase (usually in $kJ/mol$). This process is endothermic (positive energy). IE decreases down a group (electrons are farther from the nucleus) and generally increases across a period.

  • Electron Affinity: The energy involved when an electron is added to a gaseous atom or ion. It generally increases moving up a group and from left to right across a period. Non-metals typically have the highest electron affinities.

  • Metallic Property: Metals tend to lose electrons (forming cations), while non-metals tend to gain electrons (forming anions). Metallic character decreases across the table from left to right.

Molecularity of Elements

• Monoatomic Elements: Consist of single atoms not connected to each other, such as Helium ($He$) and Neon ($Ne$).

• Diatomic Elements: Elements that occur naturally as molecules containing two atoms connected by a chemical bond. These seven elements do not exist as free atoms:

  • Hydrogen ($H_2$)

  • Nitrogen ($N_2$)

  • Fluorine ($F_2$)

  • Oxygen ($O_2$)

  • Iodine ($I_2$)

  • Chlorine ($Cl_2$)

  • Bromine ($Br_2$)

  • Mnemonic: "Have No Fear Of Ice Cold Beer"

• Polyatomic Elements: Elements with more than two atoms per molecule.

  • Ozone: $O_3$

  • Phosphorus: $P_4$

  • Sulfur: $S_8$

  • Diamond: Consists of millions of carbon atoms bonded in a gigantic cluster.