Atomic Structure

• Atoms consist of protons, neutrons, and electrons.
• Protons are positively charged and reside in the nucleus.
• Neutrons have no charge and reside in the nucleus.
• Electrons are negatively charged and surround the nucleus.
• The number of neutrons determines isotopes.
• Atomic number is the number of protons, giving an element its identity.
• The number of electrons determines the charge of an atom.
• For carbon, the atomic number is 6 (6 protons), usually with 6 neutrons, making the atomic weight 12.
• Opposite charges attract.

Atomic Orbitals

• Protons and neutrons exist in the nucleus
• Electrons are found outside the nucleus in atomic orbitals.
• Atomic orbitals exist in concentric shells, energy increases further from the nucleus.
• Each row of the periodic table represents a new orbital shell.
• Position on the periodic table dictates the number and type of orbitals in an atom.
• First-row elements (H, He) have one 1s orbital.
• Second-row elements (Li, Be, B, C, N, O, F, Ne) have one 1s orbital, one 2s orbital, and three 2p orbitals which are degenerate

Electron Configuration and Valence Shell

• An atom's electron configuration describes which orbitals have electrons.
• Neutral atoms have the same number of electrons as protons.
• Aufbau principle: fill the lowest energy orbital first.
• Pauli exclusion principle: only 2 electrons per orbital, with opposite spin.
• Hund's rule: put one electron into each degenerate orbital before pairing.
• Valence electrons are electrons in the outer shell.
• The most important electrons are in the valence shell.
• Count from left to right on the periodic table to find the number of valence electrons.

Ionic Bonding

• Atoms strive to achieve a full valence shell, which is done by transferring or sharing electrons to make chemical bonds.
• For second-period atoms, a full valence shell has 8 electrons (octet).
• Two types of chemical bonding are ionic and covalent bonding.
• Occurs through electrostatic attraction between charged ions (likes attract).
• Requires full transfer of an electron (no sharing).
• Occurs between elements with large differences in electronegativity, typically far right and far left elements on the periodic table.
• Forms a salt, not a molecule.
• Atoms on the left of the periodic table readily lose an electron becoming a positively charged cation.
• Atoms on the right of the periodic table readily gain an electron becoming a negatively charged anion.

Covalent Bonding

• Electrons are shared to fill valence shells for atoms closer in electronegativity.
• Creates discrete bonds rather than a crystal lattice.
• Forms molecules with a distinct shape.
• Involves the sharing of electrons between two nuclei to complete the octet of both atoms.
• Unequal sharing of electrons between two nuclei creates a polar covalent bond with a bond dipole.
• Bond is created between two atoms with a large electronegativity difference.

Covalent Bonds and Organic Structures

• Each element forms a predictable number of covalent bonds when incorporated into a molecule.
• For atoms with 1-4 valence electrons, # bonds = # of valence electrons.
• For atoms with 5 or more valence electrons, # of bonds = 8 - # of valence electrons; they also have lone pairs.
• Carbon generally forms 4 bonds which are neutral.
• Nitrogen generally forms 3 bonds and has 1 lone pair which are neutral.
• Oxygen generally forms 2 bonds and has 2 lone pairs.
• Hydrogen generally forms 1 bond and has no lone pairs which is neutral.
• Halogens (F, Cl, Br, I) typically form 1 bond and have 3 lone pairs.
• Second-period elements (C's row) never exceed 8 electrons in the valence shell, and can never have more than 4 shells.

Covalent Bonds and Formal Charge

• When an atom deviates from its predicted number of bonds, a formal charge is created.
• Formal charge is the charge assigned to an atom, assuming electrons are shared equally.
• Formal charge = valence electrons - (# of bonds + # of unshared electrons).

Drawing Lewis Structures

• Need to know where electrons are and how atoms are connected, which is represented in Lewis structures
• Determine valence electrons, keeping overall charge in mind (add an electron for negative charge, subtract for a positive charge).
• Place the atom that makes the most bonds in the middle.
• Draw a skeleton that places atoms that form more than one bond in the center.
• Add hydrogen and halogens around the outside, and use electrons to form bonds between the atoms
• Avoid O-O bonds as they are reactive and unstable.
• 2nd-period elements follow the octet rule.
• Add lone pairs to complete octets, starting with the most electronegative atoms, until running out of valence electrons.
• If there are atoms that do not have full octets, fill the remaining octets by making multiple bonds (using lone pairs of electrons).
• Add any formal charges, if required.
• The first rule of Lewis structures states 2nd period elements cannot have more than eight electrons.

Beyond Lewis Structures

• Kekule structures are similar to Lewis structures, but lone pairs are not drawn in.
• Number of lone pairs are implied by formal charge.
• Condensed structures are different in that bonds are not drawn in, only atoms are shown.
• Knowing bonding patterns of atoms help interpret condensed structures
• Atoms connected to each carbon are written beside it in a CXn ​​pattern
• Double and triple bonds are implied to fulfill octets.

How Atoms Form Covalent Bonds

• Atoms form covalent bonds when two orbitals physically overlap in space
• Each electron then has access to both nuclei
• Forms a sigma (σ) bond

Bonding and Shape of Molecules: Orbital Hybridization

• The covalent bonding of organic molecules determines its shape
• Methane has four identical bonds to hydrogen but does not have four identical orbitals.
• Four identical bonds are created from s and p orbitals via orbital hybridization.

Molecules with a C-C Bond

• Ethane molecule has a C-C bond where each Carbon is tetrahedral and sp3 hybridized
• Sigma bond is formed through end-on overlap and allows for rotation around the C-C bond

Orbital Hybridization w/ a C-C Double Bond

• Carbon atoms involved in a double bond in ethylene are trigonal planar.
• Carbon only has 3 atoms surrounding it.
• Each Carbon is Mix 1s+2p =3, and has one unhybridized p orbital.
• Double bond consists of 1 sigma bond + 1 pi bond.
• A Pi bond cannot rotate, the Pi bond is side-on overlap of p-orbital.

Orbital Hybridization w/ a C-C Triple Bond:

• Molecules with a C-C triple bond have hybridized orbitals connected with geometry that is linear
• Each carbon atom has only two bonds, needs a hybridized orbital
• Triple bond consists of 1 sigma bond and 2 pi bonds
• An atom can only hace 1 sigma bond + 2 pi bonds

Atoms with Lone Pairs:

• Lone pairs still use orbitals and take up space around the central atom
• All 3 atoms have"4 things" around them
• The 3 sigma bonds and 1 lone pair need 4 hybrid orbitals in ammonia.
• Two sigma bonds and 2 lone pairs need 4 orbitals in water.
• Count sigma bonds, atoms, and lone pairs.

Orbital Hybridization: A Summary

• The number of atomic orbitals in equals the number of hybridized orbitals out.
• All hybridized orbitals are identical.
• Not all atomic orbitals have to hybridize, which still becomes P orbitals For Pi bonds.
• 4 "things" has sp3 Hybridization
• 3 "things" has sp2 Hybridization
• 2 "things" has sp Hybridization

Hybridization Summary:

• All single bonds are sigma bonds.
• All double bonds are one sigma and one pi bond.
• All triple bonds are one sigma and two pi bonds.
• To determine hybridization of C, Nor O, count the number of pi bonds.
• No pi bonds formed mean sp3 hybridization.
• One pi bond formed means sp2 hybridization.
• Two pi bonds formed means sp hybridization.
• Carbocations and carbon radicals are an exception, and are sp2 hybridization without pi bonds because of its partially filled or empty p-orbital.

Determining Hybridization:

• To determine hybridization and geometery
• Determine the "things" equal on the compound

Module 1 Learning Objectives

• Determine the electron configuration for an element.
• Determine the number of valence electrons for an atom.
• Identify an ionic or covalent bond.
• Identify bond dipoles and predict the direction of the dipole with a dipole arrow or partial charges.
• Draw a Lewis structure for an organic molecule, including formal charges and lone pairs.
• Determine the formal charge or the number of lone pairs on an atom.
• Interpret condensed structures of molecules.
• State the hybridization of atoms in organic molecules.
• Predict molecular shapes and bond angles based on hybridization.
• Describe the bonding in a molecule in terms of the overlapping orbitals and the types of orbitals holding lone pairs.