BackGeneral Chemistry Review: Atoms, Bonding, and Organic Structure
Study Guide - Smart Notes
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Atoms and Subatomic Particles
Structure of the Atom
Atoms are the fundamental units of matter, composed of three main subatomic particles: protons, neutrons, and electrons. Understanding their properties is essential for studying chemical bonding and molecular structure.
Protons: Mass ≈ 1 atomic mass unit (amu), charge = +1. Located in the nucleus.
Neutrons: Mass ≈ 1 amu, charge = 0. Located in the nucleus.
Electrons: Mass ≈ 0.0005 amu (often approximated as 0), charge = -1. Found in the electron cloud surrounding the nucleus.
Key atomic numbers:
Atomic number (Z): Number of protons in the nucleus.
Mass number (A): Total number of protons and neutrons.
Charge: Number of protons minus number of electrons.
Isotopes
Definition and Examples
Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons, resulting in different mass numbers. Changing the number of protons, neutrons, or electrons alters the identity, isotope, or charge of the atom, respectively.
Different number of protons: New element (different atomic number).
Different number of neutrons: New isotope (different mass number).
Different number of electrons: New ion (different charge).
Isotope | Protons | Neutrons | Electrons |
|---|---|---|---|
12C | 6 | 6 | 6 |
13C | 6 | 7 | 6 |
14N | 7 | 7 | 7 |
15N | 7 | 8 | 7 |
1H: Hydrogen
2H: Deuterium
3H: Tritium
Electron Distribution in Atoms
Atomic Orbitals and Energy Levels
Electrons occupy specific regions around the nucleus called atomic orbitals. Each orbital has a characteristic energy, and electrons fill these orbitals according to several principles:
Aufbau Principle: Electrons fill orbitals starting with the lowest energy available.
Hund's Rule: Electrons occupy degenerate (equal energy) orbitals singly before pairing up.
Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins.
General energy ordering for orbitals:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s ...
Core vs. Valence Electrons
Definitions and Periodic Trends
Electrons in an atom are classified as core electrons or valence electrons:
Core electrons: Occupy inner shells (lower principal quantum number, n).
Valence electrons: Occupy the outermost shell; responsible for chemical reactivity.
Elements in the same group (column) of the periodic table have the same number of valence electrons and similar chemical properties.
Element | Valence Electrons |
|---|---|
H | 1 |
C | 4 |
N | 5 |
O | 6 |
F | 7 |
Ne | 8 |
Atoms tend to gain, lose, or share electrons to achieve a stable octet (eight valence electrons).
Chemical Bonding
Ionic and Covalent Bonds
Atoms form chemical bonds by transferring or sharing electrons to achieve stable electron configurations.
Ionic bonding: One atom transfers electrons to another, forming oppositely charged ions.
Covalent bonding: Two atoms share one or more pairs of electrons.
Preferred Bonding Patterns and Formal Charge
Atoms have preferred numbers of bonds and lone pairs to minimize formal charge. The formal charge is calculated as:
Element | Uncharged | Positive Formal Charge | Negative Formal Charge |
|---|---|---|---|
H | 1 bond, 0 lone pairs | 0 bonds, 0 lone pairs | 0 bonds, 1 lone pair |
C | 4 bonds, 0 lone pairs | 3 bonds, 0 lone pairs | 3 bonds, 1 lone pair |
N | 3 bonds, 1 lone pair | 4 bonds, 0 lone pairs | 2 bonds, 2 lone pairs |
O | 2 bonds, 2 lone pairs | 3 bonds, 1 lone pair | 1 bond, 3 lone pairs |
F | 1 bond, 3 lone pairs | 2 bonds, 2 lone pairs | 0 bonds, 4 lone pairs |
Bond Polarity
Electronegativity and Dipoles
Covalent bonds can be nonpolar or polar depending on the difference in electronegativity between the bonded atoms:
Nonpolar bonds: Atoms have similar electronegativity; electron density is shared equally.
Polar bonds: Atoms have different electronegativity; electron density is unequally shared, creating a dipole.
Electrostatic Potential Maps
Visualizing Electron Density
Electrostatic potential maps are graphical representations showing regions of high and low electron density in molecules. These maps help visualize areas of partial positive and negative charge, which are important for understanding reactivity and intermolecular interactions.
Red/orange: Regions of high electron density (attract positive charge).
Blue/green: Regions of low electron density (attract negative charge).
Example: In water (H2O), the oxygen atom appears red (electron-rich), while the hydrogen atoms appear blue (electron-poor).
Skeletal Structures
Drawing Organic Molecules
Skeletal (line-angle) structures are a simplified way to represent organic molecules:
Each line represents a carbon-carbon bond; the ends and bends are carbon atoms.
Hydrogens bonded to carbon are not shown explicitly; each carbon is assumed to have enough hydrogens to complete four bonds.
All other atoms (e.g., O, N, halogens) are shown explicitly.
Lone pairs may or may not be shown, but are important for understanding reactivity.
Example: The condensed formula CH3CH2OH is drawn as a two-carbon chain with an -OH group attached to the second carbon.
Practice: Converting and Interpreting Skeletal Structures
Convert condensed structural formulas to skeletal structures by identifying carbon chains and functional groups.
Determine the number of implicit hydrogens on each carbon by ensuring each carbon has four bonds.
Example: In a skeletal structure, a carbon at the end of a line with no other atoms attached has three implicit hydrogens (methyl group).
Additional info: Skeletal structures are essential for quickly visualizing large organic molecules and are the standard in organic chemistry literature.
