Organic Chemistry Chapter 1

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Valence Shell
outermost electron shell
Valence Electrons
electrons on the outermost energy level of an atom
Valence electrons on the periodic table
The number of valence electrons in an A group of the periodic table = its group number
Octet Rule Theory
Atoms lose, gain, or share electrons to have a full eight valence electrons to reach noble gas configuration
Chemical Bond
The force that holds atoms together within molecules
Ionic compound
a compound that consists of positive and negative ions
Electrostatic attraction
A stabilizing interaction between opposite charges
Ionic Bond
The electrostatic attraction that binds oppositely charged ions together through transfer of electrons
Covalent Bond
Covalent Bond
A chemical bond formed when two atoms share electrons
Lewis Structures
Molecular structures represented by dots and – bonds
Lone pairs (Unshared pairs)
Pairs of valence electrons that are not shared between atoms
Octet Rule
Octet Rule
The sum of all shared and unshared valence electrons around each atom in many stable covalent compounds is eight
Double bond
Double bond
a covalent bond in which two pairs of electrons are shared between two atoms
Triple bond
a covalent bond in which three pairs of electrons are shared between two atoms
Formal charge
A positive or negative charge on an individual atom. # of valence electrons – ( # dots + # lines)
Steps to formal charge
1. Group number = # of valence electrons in the neutral atom
2. Valence electron count for the atom (unshared V e-‘s + covalent bonds)
3. Valence electron count – group number
Polar bond
a covalent bond between atoms in which the electrons are shared unequally
Electronegativity
The ability of an atom to attract electrons when the atom is in a compound
High electronegativity
Top right, the atom attracts more electrons
Low electronegativity (electropositive)
Bottom left, the atom attracts less electrons
Polar molecules
Molecules that have negative and positive ends (Ex. HCl)
Non-polar Molecules
Molecules that have no charged poles. No dipole movement. (Ex. H)
Bond dipole
The dipole moment that is due to unequal electron sharing in a covalent bond.
Molecular Shape
The geometric shape formed by atoms bonded to the central atom in a molecule
Linear
Linear
(CO2) Bond dipoles oriented in opposite directions. 180 degrees, line.
Atomic connectivity
the specification of how atoms in a molecule are connected
Molecular geometry
The specification of how far apart the atoms are and how they appear in space. 3D shape, determined by lone pairs and the # of bonds.
Bond length
The distance between the nuclei of two bonded atoms
Bond angle
The angle formed by two bonds to the same atom.
Bond order
Number of electron pairs (covalent bonds) shared by two bonded atoms. (Bond lengths decrease with bonding order)
Bond order formula
bond order = 1/2 (number of electron in bonding orbitals – number of electrons in antibonding orbitals)
VSEPR Theory
Valence-shell electron-pair repulsion theory: Bonds and electrons repel to arrange around the central atom so bonds are as far apart as possible.
Tetrahedron
A 3D object with 4 triangular faces (Ex. Methane, CH4)
Trigonal planar
an arrangement of atoms where the three pairs of electrons are placed 120 degrees apart on a flat plane.
Trigonal planar
an arrangement of atoms where the three pairs of electrons are placed 120 degrees apart on a flat plane.
Trigonal pyramidal
The molecular geometry of a compound with 3 shared pairs and 1 lone pair of electrons. 107 degree angles.
Trigonal pyramidal
The molecular geometry of a compound with 3 shared pairs and 1 lone pair of electrons. 107 degree angles.
Dihedral angle
An angle formed by two half planes with a common edge.
Dihedral angle
An angle formed by two half planes with a common edge.
Torsion angle
Spatial relationship of the bonds on adjacent atoms
Torsion angle
Spatial relationship of the bonds on adjacent atoms
Resonance Hybrid
Structures that arise from the possibility to draw a multiple bond in different positions equivalently. Delocalization of pi bonds.
Resonance Hybrid
Structures that arise from the possibility to draw a multiple bond in different positions equivalently. Delocalization of pi bonds.
Heisenberg uncertainty principle
it is impossible to know exactly both the velocity and the position of a particle at the same time
Heisenberg uncertainty principle
it is impossible to know exactly both the velocity and the position of a particle at the same time
Quantum Numbers
The four numbers that define each particular electron of an atom.
Quantum Numbers
The four numbers that define each particular electron of an atom.
Principle Quantum Number
n: describes the energy of the electron and distance from the nucleus. (1, 2, 3, 4, 5)
Principle Quantum Number
n: describes the energy of the electron and distance from the nucleus. (1, 2, 3, 4, 5)
Angular momentum quantum number
symbolized by l, indicates the shape of the orbital (n-1) (1=s, 2=p, 3=d, 4=f)
Angular momentum quantum number
symbolized by l, indicates the shape of the orbital (n-1) (1=s, 2=p, 3=d, 4=f)
Magnetic Quantum number
symbolized by m, indicates the orientation of an orbital around the nucleus (-L/+L)
Magnetic Quantum number
symbolized by m, indicates the orientation of an orbital around the nucleus (-L/+L)
Spin Quantum number
The two fundamental spin states of an electron in an orbital (+1/2, -1/2)
Spin Quantum number
The two fundamental spin states of an electron in an orbital (+1/2, -1/2)
Node
A point in a 3D wave where the wave amplitude = 0
Node
A point in a 3D wave where the wave amplitude = 0
Trough
Lowest point on a wave
Trough
Lowest point on a wave
Peak
Highest point on a wave
Peak
Highest point on a wave
Electronic configurations
The arrangements of electrons indicated by a specific notation. (1s2 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6…)
Electronic configurations
The arrangements of electrons indicated by a specific notation. (1s2 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6…)
Electron spin
2 electrons occupying an orbital must have opposite spins (magnetic property of electrons)
Electron spin
2 electrons occupying an orbital must have opposite spins (magnetic property of electrons)
Aufbau principle
Rule that electrons occupy the orbitals of lowest energy first
Aufbau principle
Rule that electrons occupy the orbitals of lowest energy first
Pauli Exclusion principle
No two electrons in the same atom can have the same four quantum numbers
Pauli Exclusion principle
No two electrons in the same atom can have the same four quantum numbers
Hund’s Rule
Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. Spins of the unpaired electrons must be the same.
Hund’s Rule
Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. Spins of the unpaired electrons must be the same.
Valence orbitals
orbitals that contain the outer-shell electrons of an atom
Valence orbitals
orbitals that contain the outer-shell electrons of an atom
Energy in orbitals
Energy increases up the orbitals (Low- 1s2, 2s2, 2p6 Higher)
Energy in orbitals
Energy increases up the orbitals (Low- 1s2, 2s2, 2p6 Higher)
Molecular Orbital theory
Molecular orbitals having zero values in regions between
nuclei.
– Obtained by subtracting atomic orbitals.
Molecular Orbital theory
Molecular orbitals having zero values in regions between
nuclei.
– Obtained by subtracting atomic orbitals.
Antibonding molecular orbitals
Sigma bonds. Have their “electron density” concentrated “outside” the 2 atoms. Higher in energy than bonding orbitals. (Destructive) Out-of-phase interactions
Antibonding molecular orbitals
Sigma bonds. Have their “electron density” concentrated “outside” the 2 atoms. Higher in energy than bonding orbitals. (Destructive) Out-of-phase interactions
Bonding molecular orbital
Has lower energy and greater stability than the atomic orbitals from which it was formed. (Constructive) In-phase interactions.
Bonding molecular orbital
Has lower energy and greater stability than the atomic orbitals from which it was formed. (Constructive) In-phase interactions.
Cylindrical symmetry
electron density looks the same no matter how a molecule is turned about the line joining two nuclei.
Cylindrical symmetry
electron density looks the same no matter how a molecule is turned about the line joining two nuclei.
Sigma bonds
Single bonds, overlap of two S orbitals, 2 P orbitals (end-to end), or a S and P orbital
-represents the sharing of one pair of electrons
Sigma bonds
Single bonds, overlap of two S orbitals, 2 P orbitals (end-to end), or a S and P orbital
-represents the sharing of one pair of electrons
Pi bonds
-Overlap of two P orbitals side by side
-Double bonds = 1 π bond
-double bond = 1 π bond and 1 sigma bond
– Weaker than σ bonds (less energy to break bond).
Pi bonds
-Overlap of two P orbitals side by side
-Double bonds = 1 π bond
-double bond = 1 π bond and 1 sigma bond
– Weaker than σ bonds (less energy to break bond).
Total electron density
the probability of finding electrons in a molecule
Total electron density
the probability of finding electrons in a molecule
Hybridization
the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies
Hybridization
the mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies
Hybrid orbitals
orbitals of equal energy produced by the combination of two or more orbitals on the same atom
Hybrid orbitals
orbitals of equal energy produced by the combination of two or more orbitals on the same atom
Categories: Organic Chemistry