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