General Chemistry Translated: Finding the Electrons (Chapter 2 Study Notes)

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General Chemistry Translated: Finding the Electrons

Structure of the Atom

Atoms are the fundamental building blocks of matter, composed of three subatomic particles: protons (p+), neutrons (n0), and electrons (e-). The arrangement and number of these particles determine the identity and properties of each element.

Atomic number (Z): Number of protons in the nucleus; distinguishes one element from another.
Isotopes: Atoms of the same element with different numbers of neutrons (different mass numbers).
Average atomic mass: Weighted average of all isotopes, as shown on the periodic table.

Shells and Orbitals

Electrons in atoms are arranged in shells and orbitals, described by quantum numbers. The principal quantum number (n) corresponds to the shell or energy level.

n = 1: Holds 2 electrons
n = 2: Holds 8 electrons
n = 3: Holds 18 electrons
Organic chemistry primarily focuses on n = 1 and n = 2 shells.
Each orbital can hold a maximum of 2 electrons.

Energy levels of orbitals: s-orbitals are lower in energy than p-orbitals within the same shell.

Electron Configuration

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

Pauli Exclusion Principle: Each orbital holds a maximum of 2 electrons with opposite spins.
Aufbau Principle: Electrons fill orbitals in order of increasing energy.
Hund's Rule: Electrons occupy degenerate (equal energy) orbitals singly before pairing up.

Atoms to Molecules

Atoms combine to form molecules through chemical bonds. The two main types of bonds in organic chemistry are:

Covalent bonds: Electrons are shared between atoms.
Ionic bonds: Electrons are transferred from one atom to another.

Lewis dot structures are simplified representations showing only valence electrons around the elemental symbol. Atoms with similar valence configurations exhibit similar reactivity.

Lewis's Octet Rule

The octet rule states that atoms tend to bond in ways that give them eight electrons in their valence shell, achieving a noble gas configuration.

Ionic bonding: Electrostatic attraction between oppositely charged ions (e.g., Na+ and Cl-).

Covalent Versus Ionic Bonds

The type of bond formed depends on the electronegativity difference (ED) between atoms:

Ionic bonds: ED between 2.0 and 4.0; electrons are not shared.
Polar covalent bonds: ED between 0.4 and 2.0; electrons are shared unequally.
Pure covalent bonds: ED below 0.4; electrons are shared equally.

Ionic Bonding

Metals lose electrons to become cations.
Non-metals gain electrons to become anions.
Cations and anions combine so the overall charge of the ionic compound is zero.

Covalent Bonding

Covalent bonds form by sharing electrons between two non-metal atoms.
The type of covalent bond is governed by the electronegativity difference between the atoms.

Drawing Lewis Structures (LS)

Lewis dot structures predict how many bonds an atom will form by showing electron sharing. For example, methane (CH4) consists of one carbon atom (needs four bonds) and four hydrogen atoms (each needs one bond).

Formal Charge

Formal charge helps determine the most stable Lewis structure. It is calculated as:

Formula:
The most stable Lewis structure minimizes formal charges on atoms.

Lewis Structures and Reactivity

Lewis structures are foundational for understanding molecular reactivity. Key rules include:

Negatives attack positives: Negative charges (electron-rich sites) donate electrons to positive charges (electron-poor sites).
Arrows in mechanisms show the movement of electron pairs.

VSEPR and Molecular Shape

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes by assuming electron pairs repel each other and arrange as far apart as possible.

Areas of electron density: Includes lone pairs, single, double, and triple bonds (each counts as one area).

Shapes Around a Central Atom

The number of electron density areas around a central atom determines molecular geometry and bond angles:

Areas of Electron Density	Predicted Shape	Average Bond Angle	Example Molecule
4	Tetrahedral	109.5°	Methane (CH4)
3	Trigonal planar	120°	Ethene (C2H4)
2	Linear	180°	Ethyne (C2H2)

Polarity of Molecules

Polar covalent bonds within a molecule can combine to give an overall dipole moment, making the molecule polar. The dipole moment depends on both bond polarity and molecular shape.

C–Cl bonds are polar; C–H bonds are non-polar.
Molecular dipoles result from the vector sum of all bond dipoles.

Valence Bond Theory

Valence bond theory describes covalent bonding as the overlap of atomic orbitals, using only valence electrons to form bonds. For example, two hydrogen atoms form H2 by sharing their 1s electrons.

Orbital Hybridization (sp3, sp2, sp)

Hybridization explains the observed shapes of molecules by mixing atomic orbitals:

sp3 hybridization: Mix 1 s-orbital with 3 p-orbitals to get 4 sp3 orbitals (tetrahedral geometry, e.g., methane).
sp2 hybridization: Mix 1 s-orbital with 2 p-orbitals to get 3 sp2 orbitals (trigonal planar, e.g., ethene).
sp hybridization: Mix 1 s-orbital with 1 p-orbital to get 2 sp orbitals (linear, e.g., ethyne).

Bond Length and Bond Strength

Bond length and strength depend on the hybridization of the carbon atom:

Molecule	Hybridization of Carbon	C–H Bond Length (Å)
Methane	sp3	1.09
Ethene	sp2	1.08
Ethyne	sp	1.06

The more s-character in the hybrid orbital, the shorter and stronger the bond.

s Orbitals	p Orbitals	Hybridization	% s Character	% p Character
1	3	sp3	25%	75%
1	2	sp2	33%	67%
1	1	sp	50%	50%

Molecular Orbital Theory

Molecular orbital (MO) theory describes bonding by combining atomic orbitals to form molecular orbitals. For H2:

Two atomic orbitals combine to form one bonding (σ) and one antibonding (σ*) molecular orbital.

MO Pictures of Neutral Molecules

MO diagrams can be drawn for molecules like methane, ethene, and ethyne, showing the overlap of atomic orbitals and the location of bonding and antibonding orbitals.

Resonance

When more than one valid Lewis structure exists for a molecule, the actual structure is a resonance hybrid of these forms. Resonance is depicted using curved arrows to show electron movement.

Contributing Resonance Structures

Not all resonance structures contribute equally. The major contributor is determined by:

Maximum number of covalent bonds.
Fewest atoms with incomplete octets.
Least separation of unlike charges.
Negative charges on the more electronegative atom.

Note: The first two criteria are more important than the last two.