Resonance and Carbonate Ion Structures

Resonance

• Resonance structures are multiple valid Lewis diagrams for a molecule or ion, where only the positions of bonding and non-bonding electrons differ (connectivity of atoms stays the same)
• Each resonance structure contributes to the actual structure (resonance hybrid), which is the overall representation of the molecule
• Resonance only occurs within the pi system, which involves pi bonds, lone pairs, and atoms with incomplete octets
• Sigma bonds are not involved in resonance, as they form the backbone of the molecule and cannot be broken without breaking the molecule itself

Carbonate Ion Resonance

• The carbonate ion (CO3^2-) has three equivalent resonance structures
• Each structure depicts the carbon atom double-bonded to one oxygen atom and singly bonded to the other two
• The resonance hybrid shows the carbon atom with partial double-bond character to all three oxygen atoms, giving them equal bond lengths and an overall structure that's more stable than any single resonance contributor

Resonance Energy

• Resonance lowers the energy of a molecule or ion
• The resonance hybrid actually has lower energy than any individual resonance contributor would suggest
• This lower energy is called resonance energy and makes the molecule or ion more stable
• Resonance energy can be visualized as the "extra" stability that results from delocalized electrons

Common Resonance Patterns

• Two-center systems are resonance patterns involving the movement of only pi electrons and lone pairs
• Three-center systems involve the movement of electrons across more than two atoms, without breaking any sigma bonds

Chapter Objectives

• Understand the use of hybrid orbitals to describe bonding in molecules
• Write hybridization schemes for sp, sp2, and sp3 hybridized atoms and determine their hybridization in molecules, radicals, and ions
• Determine the shape of a molecule and bond angles using hybridization
• Identify lone pairs and charges on atoms
• Identify the presence and direction of dipoles in molecules
• Distinguish various types of bonds in compounds including ionic and covalent bonds
• Draw Lewis structures
• Understand the concept of delocalized electrons
• Explain the significance of resonance contributors and resonance energy
• Draw resonance contributors and the resonance hybrid of a structure

Structure of an Atom

• Atoms are composed of a positively charged nucleus (made of protons and neutrons) surrounded by a cloud of negatively charged electrons
• The nucleus is very dense and small (about 10^-15 meters in diameter)
• The electron cloud has a larger diameter (about 10^-10 meters or 200 picometers)

Atomic Number and Atomic Mass

• The atomic number (Z) represents the number of protons in an atom's nucleus
• The mass number (A) is the sum of protons and neutrons
• All atoms of the same element share the same atomic number
• The number of electrons in an atom equals the number of protons
• Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers
• The atomic mass (atomic weight) is the weighted average mass of an element's naturally occurring isotopes, expressed in atomic mass units (amu)
• The molecular weight of a compound is the sum of the atomic weights of all the atoms in the molecule

Atomic Orbitals

• Atomic orbitals describe the regions where electrons are most likely to be found

• Each electron in an atom is defined by four quantum numbers:

  • Principal quantum number (n): Indicates the electron shell (n=1, 2, 3,...), influencing size and energy
  • Angular momentum quantum number (l): Describes the shape of the orbital (l= 0, 1, 2,...n-1), corresponding to s, p, d orbitals
  • Magnetic quantum number (ml): Determines the orientation of the orbital in space (ml=-l, -l+1, ..., l-1, l)
  • Electron spin quantum number (ms): Specifies the electron's spin (ms=+1/2, -1/2), which determines its magnetic moment

• No two electrons can have the same set of all four quantum numbers (Pauli Exclusion Principle)

s Orbitals

• All s orbitals are spherical in shape
• As the principal quantum number (n) increases, the size of the s orbital also increases
• The number of nodes (regions of zero electron probability) in an s orbital is n-1 (1s has 0 nodes, 2s has 1 node, 3s has 2 nodes).

p Orbitals

• There are three p orbitals (px, py, pz) oriented along the x, y, and z axes respectively
• Each p orbital has a dumbbell shape with a node at the nucleus
• As n increases, the p orbitals get larger

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory explains the shapes of molecules based on the repulsions among electron pairs in the valence shell of an atom
• Electron pairs repel each other, whether they are in bonds (bond pairs) or lone pairs (non-bonding pairs)
• Electron pairs adopt orientations around the atom to minimize repulsions

Electron Group and Molecular Geometry

• Electron group geometry refers to the spatial arrangement of all electron pairs around the central atom
• Molecular geometry describes the shape formed by the atomic nuclei only
• Electron group geometry can be linear (2 electron pairs), planar (3 electron pairs), or tetrahedral (4 electron pairs)
• Molecular shape can vary depending on the number and position of lone pairs on the central atom

Hybridization

• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies that better account for the observed bond angles and molecular geometry
• sp3 Hybridization (s + 3p orbitals combined): Forms four equivalent sp3 hybrid orbitals, directed towards the corners of a tetrahedron, resulting in tetrahedral geometry, e.g., methane (CH4).
• sp2 Hybridization (s + 2p orbitals combined): Forms three equivalent sp2 hybrid orbitals, oriented in a trigonal planar arrangement, e.g., ethene (C2H4).
• sp Hybridization (s + 1p orbital combined): Forms two equivalent sp hybrid orbitals, oriented in a linear arrangement, e.g., acetylene (C2H2).

Valence Bond Theory

• Valence Bond (VB) theory explains bonding in terms of the overlap of atomic orbitals
• The overlap of atomic orbitals leads to the formation of covalent bonds
• Hybridization helps explain the observed bond angles and shapes of molecules using VB theory

Delocalized Electrons

• Delocalized electrons are electrons that are not associated with a single atom or bond, but are spread out over multiple atoms
• Delocalization contributes to the stability of molecules and ions, such as in resonance structures

Covalent Bonding

• Covalent bonds form when atoms share electrons.
• Atoms achieve a noble gas configuration by sharing electrons.
• Carbon, nitrogen, oxygen, sulfur, and phosphorus readily form multiple bonds.

Lewis Structures

• Lewis structures show the arrangement of valence electrons in a molecule or ion.
• The number of valence electrons for an atom equals its group number on the periodic table.
  • For main group elements.
• For negative ions, add one electron for each negative charge.
• For positive ions, subtract one electron for each positive charge.

Formal Charges

• Formal charge = number of valence electrons - (number of lone pair electrons + 1/2 number of bonding electrons).

Polar Covalent Bonds

• Polar covalent bonds have a dipole moment, which is a measure of the separation of charge in the bond.
• Dipole moment (µ) = Q (charge on atom) × r (distance between two charges).
• The net dipole moment of a molecule is the vector sum of all its bond dipoles.

Valence Bond (VB) Theory

• VB theory explains the formation of covalent bonds in terms of the overlap of atomic orbitals.
• Sigma (σ) bonds are formed by the end-on overlap of two atomic orbitals.
• Pi (π) bonds are formed by the sideways overlap of two parallel atomic orbitals.

Hybridization

• Hybridization is the mixing of atomic orbitals to form new hybrid orbitals that have different energies and shapes.
• Hybridization allows for the formation of multiple bonds between atoms.
• Hybridized orbitals can be used to explain the geometry of molecules.

Delocalized Electrons

• Delocalized electrons are electrons that are not associated with a single atom or bond but are spread out over multiple atoms or bonds.
• Delocalized electrons are found in molecules with multiple bonds, such as alkenes and aromatic compounds.
• Delocalized electrons contribute to the stability of the molecules.
• Localized electrons are restricted to a specific atom or bond.

Condensed Structures

• Condensed structures provide a shorthand way of writing Lewis structures.
• Carbon and hydrogen atoms are not explicitly shown in condensed structures.
• The number of hydrogen atoms attached to each carbon atom is implied.
• Carbon atoms connected by a single bond are shown as a single letter.

Skeletal or Line Structures

• Skeletal or line structures are even more simplified representations of molecules.
• Each vertex and end of each line represents a carbon atom.
• Hydrogen atoms are omitted, except when they are attached to atoms other than carbon.
• Atoms other than carbon or hydrogen are explicitly drawn.

Key concepts

• An atom's atomic number equals the number of protons.
• An atom's mass number equals the sum of protons and neutrons.
• Isotopes of an element have the same atomic number but different mass numbers.
• Electrons are assigned to orbitals according to the Aufbau principle, the Pauli exclusion principle, and Hund's rule.

This quiz explores the concept of resonance in chemistry, focusing on the carbonate ion and its multiple valid Lewis diagrams. Discover how each resonance structure contributes to the stability and representation of the molecule, highlighting the significance of pi bonds and electron arrangement.