CHEM 1111: BONDS AND STRUCTURES
49 Cards in this Set
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Bond
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Strong attractive force between atoms
Ionic and Covalent
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Ionic Bonds
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Electrostatic attraction between a POSITIVE and NEGATIVE ion
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Cation
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Positively charged ion
Made from loss of e-
Generally smaller than neutral atom
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Anion
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Negatively charged ion
Made from gain of e-
Generally larger than neutral atom
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Lattice Energy
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Change in energy that occurs when an ionic solid is separated into isolated ions in the gas phase (strength of attraction between ions)
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Ionic Bond Properties
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High melting points
Generally conduct electricity in the liquid or aqueous phase
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Monoatomic
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Consisting of one ion
Main group
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Main Group
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Lose/gain electrons to "pseudo" noble gas configuration
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Monoatomic Metals
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Lose electrons
Positive/Cations
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Monoatomic Non-Metals
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Gain electrons
Negative/Anions
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Transition
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May form more than one ion by losing/gaining electrons to a stable configuration
Ag+
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Polyatomic
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Contains more than one atom
Memorized from sheet
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Ionic Radii
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Measurement of spherical region around nucleus in which electrons are most likely present (use Zeff)
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Isoelectronic
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Two atoms, ions or molecules that have the same electronic structure and same number of valence electrons
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Covalent Bond
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Sharing of electrons between two atoms
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Coordinate Covalent
Dipolar
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Covalent bond between two atoms where one of the atoms provides both electrons that form the bond
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Lewis Dot Structures
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Pictorial representation of electrons in an atom or molecule
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Bond
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Sharing of two electrons that can be written as electrons or as a dash
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Lone Pair
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Pair of electrons not participating in a bond
Electrons must be paired (no loners)
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Octet Rule
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Tendency for atoms to have eight electrons in their valence shell
C, N, O, F, Ne, Si
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Hydrogen Octet
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Two electrons
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Boron Octet
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Six Electrons
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Expanded Octet
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Anything that has access to d subshells can expand to more than eight
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Multiple Bonds
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Sharing of more than two electrons between atoms
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Single Bond
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Sharing two electrons
H3C − CH3
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Double Bond
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Sharing four electrons
H2C = CH2
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Triple Bond
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Sharing six electrons
HC ≡ CH
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Resonance Structures
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All possible structures of multiple bonds
Structures drawn with a ↔ between them
No single structure represents actual structure it is a hybrid of all possibilities
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Formal Charges
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Hypothetical charge difference between original valence electrons and electrons in a bond
FC = VE − 1/2 BE − LPE
FC = VE − N − B/2
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Bond Length
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Distance between nuclei in a bond
Triple < Double < Single
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Bond Order
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Number of pairs of electrons in a bond
BO = 1 (2 e-) single
BO = 2 (4 e-) double
BO = 3 (6 e-) triple
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Bond Energy (Strength)
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The average energy required to break a bond in the gas phase
Triple > Double > Single
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Molecular Geometry
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General shape of a molecule determined by the three dimensional positions of the atomic nuclei
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VSEPR (Valence Shell Electron Pair Repulsion) Model
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Predicts shapes of molecules by assuming that the valence shell electron pairs are arranged about each atom so that electron pairs repulsion are minimized
Shapes fall into five categories / "families"
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Polarity
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Separation of charge resulting in poles that can be oriented in a electric field
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Dipole Movement
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Quantitative measurement of the degree of charge separation in a molecule
Vector Quantity
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Non-Polar
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Molecule does not have a "net" dipole
Generally diatomics and symmetrical molecules
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Polar
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Molecule that has a "net" dipole
Larger difference in electronegativity the larger the dipole
Generally any molecule has an asymmetry
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Atomic Orbital Theory
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Shows electrons in an atom
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Molecular Orbital Theory
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Shows electrons in molecule (Why some exist and some don't)
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Hybrid Orbital Theory
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Shows electrons in terms of molecular shape and direction - uses overlap of orbitals
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VSEPR Orbital Theory
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Valence Shell Electron Pair Repulsion
Predict shapes of molecules
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Valence Bond Theory/Hybrid
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A bond forms when both of the following are met
1. An orbital on one atom (atomic) comes to occupy a portion of the same region of space as an orbital on the other atom. They are said to overlap
2. The total number of electrons in both orbitals is no more than two
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Hybrid Orbitals
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Orbitals that formed by the "combination" of atomic orbitals
One hybrid orbital formed for every atomic orbital used
Identical hybrid orbitals have the same energy
Any unused atomic orbitals are still present and available for multiple bonding
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Sigma (σ) Bonds
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From overlapping hybrid orbitals
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Pi (π) Bonds
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From overlapping/"interacting" atomic orbitals other than s
Must also have an σ bond
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Molecular Orbital
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Orbital formed when two atomic orbitals overlap creating two molecular orbitals
One molecular orbital formed for every atomic orbital used
Helps predict the number of bonds using the bond order
Helps predict magnetic properties
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Paramagnetic
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Attracted to a magnetic field
Has unpaired electrons
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Diamagnetic
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Not attracted (or slightly repelled) by a magnetic field (no unpaired electrons)
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CHEM 1111: CHAPTER 1