Atomic Structure and Electron Configuration

Atoms and Isotopes

Atoms are the fundamental units of matter, consisting of protons, neutrons, and electrons. The atomic number defines the element and equals the number of protons. The mass number is the sum of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Ions: Atoms that have gained or lost electrons, resulting in a net charge.

• Proton: Positively charged subatomic particle.

• Neutron: Neutral subatomic particle.

• Electron: Negatively charged subatomic particle.

Example: Hydrogen Isotopes: Protium (1 proton), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons).

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals, following specific principles:

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing up.

Ground State Electron Configuration: The lowest energy arrangement of electrons for an atom.

Condensed Electron Configuration: Uses the previous noble gas to abbreviate the configuration.

Example: Phosphorus (Z = 15): Ground state: Condensed:

Periodic Trends

Electronegativity

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases from left to right across a period and increases going up a group.

• Most electronegative element: Fluorine (F).

Example: Among Group 7A elements, Cl is more electronegative than Br or I.

Chemical Bonding and the Octet Rule

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, resembling the electron configuration of noble gases.

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

• Shared Electrons: Electrons shared between atoms in a covalent bond.

Example: In H3COH, oxygen has 6 valence electrons and 2 shared electrons, achieving an octet.

Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule.

• Formula:

• Sum of formal charges in a molecule equals the overall charge.

Example: For the thiocyanate ion (NCS-), calculate formal charges for each atom using the formula above.

Lewis Dot Structures

Lewis structures represent the arrangement of valence electrons among atoms in a molecule.

1. Count total valence electrons.

2. Place the least electronegative atom in the center (except hydrogen).

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then central atom.

5. Use double/triple bonds if needed to satisfy octet rule.

6. Check formal charges for stability.

Example: Draw the Lewis structure for COCl2.

Resonance Structures

Some molecules have multiple valid Lewis structures, called resonance structures, differing only in the placement of electrons.

• Resonance structures are connected by double-headed arrows.

• The actual structure is a resonance hybrid, a blend of all resonance forms.

• Resonance involves the movement of pi electrons or lone pairs.

Example: Nitrate ion (NO3-) has three resonance structures, each with a different oxygen atom double-bonded to nitrogen.

Hybridization and Molecular Geometry

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Electron Groups: Number of atoms bonded to the central atom plus lone pairs.

Example: HCN has 2 electron groups, so the central atom is sp hybridized.

Molecular Polarity

Molecular polarity arises from the distribution of electron density in a molecule.

• Nonpolar Molecule: Symmetrical shape, even distribution of charge.

• Polar Molecule: Asymmetrical shape or presence of lone pairs, uneven charge distribution.

Example: Nitrogen trifluoride (NF3) is polar due to the presence of a lone pair on nitrogen.

Organic Chemistry Fundamentals

Functional Groups

Functional groups are specific groups of atoms within molecules responsible for characteristic chemical reactions.

• Hydrocarbons: Compounds containing only carbon and hydrogen (alkanes, alkenes, alkynes, aromatic compounds).

• With Carbonyls: Aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides.

• Without Carbonyls: Alcohols, ethers, amines, alkyl halides, thiols.

Example: Alcohols contain an -OH group; carboxylic acids contain a -COOH group.

Organic Molecules and Biological Relevance

Organic chemistry studies molecules containing carbon, often with hydrogen, oxygen, nitrogen, and other elements. Organic molecules are the basis of life and are found in biological systems and everyday products.

• Organic molecule: Contains both carbon and hydrogen.

• Hydrocarbon: Contains only carbon and hydrogen.

Example: Identifying organic molecules in product labels and their functional groups.

Summary Table: Key Concepts