ch 3 - Bonding and Chemical Interactions Flashcards
exceptions to octet rule
hydrogen can only have two valence electrons achieving the configuration of helium; lithium and beryllium bond to attain two and four valence electrons respectively; boron bonds to attain six electrons; and all elements in period 3 and greater which can expand the valence shell to include more than eight electrons by incorporating d-orbitals
Incomplete octet
hydrogen, helium, lithium, beryllium, and boron
expanded octet
any element in period 3 and greater
odd numbers of electrons
any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom
ionic bonding
one or more electrons from an atom with low ionization energy, typically a metal, are transferred to an atom with a high electron affinity, typically, a nonmetal; resulting electrostatic attraction between opposite charges holds the ions together; form crystal lattices; high melting and boiling points, dissolve in water and other polar solvents, good conductors of electricity in molten or aqueous state
covalent bonding
an electron pair is shared between two atoms, typically nonmetals, that have relatively similar values of electronegativity. the degree to which the pair of electrons is shared equally or unequally between the two atoms determines the degree of polarity in the covalent bond - if equal, nonpolar; unequal, polar; consist of individually bonded molecules
coordinate covalent
if both of the shared electrons are contributed by only one of the two atoms
cation
atom in ionic bonding that loses an electron
anion
atom in ionic bonding that gains an electron
bond order of covalent bonds
number of shared electron pairs between two atoms (single, double, or triple covalent bonds have bond orders of one, two, or three respectively)
bond length of covalent bonds
the average distance between the two nuclei of atoms in a bond; as number of shared electron pairs increases, the two atoms are pulled closer together, resulting in a decrease in bond length; thus a triple bond is shorter than a double which is shorter than a single
Bond energy of covalent bonds
the energy required to break a bond by separating its components into their isolated, gaseous atomic states; the greater the number of pairs of electrons shared between atomic nuclei, the more energy is required to break the bonds holding atoms together which means greater bond energy (with triple bond having the greatest)
Polarity of covalent bonds
occurs when two atoms have a relative difference in electronegativities (between 0.5 and 1.7). higher electronegativity atoms get the larger share of the electron density; polar bond creates a dipole, with the positive end of the dipole at the less electronegative atom and the negative a the more electronegative atom
nonpolar covalent bonding
no separation of charge across the bond; only bonds between atoms of the same element will have exactly the same electronegativity and exhibit a purely equal distribution of electrons; also occurs with nearly identical electronegativities but not perfectly. The seven common diatomic molecules are H2, N2, O2, F2, Cl2, Br, and I2; any bond with a difference in electronegativity of less than 0.5 is generally considered nonpolar
dipole moment equation
p = qd, where p is the dipole moment, q is the magnitude of the charge, d is the displacement vector separating the two partial charges; measured in Debye units (coulomb-meters)
bonding electrons
electrons involved in a covalent bond that are in the valence shell
nonbonding electrons
electrons in the valence shell that are not involved in covalent bonds
Lewis structure system
system of notation developed to keep track of the bonded and nonbonded electron pairs; the number of valence electrons attributed to a particular atom in the Lewis structure of a molecule is not necessarily the same as the number of valence electrons in the neutral atom
formal charge
difference between neutral atom number of electrons and number of electrons attributed to atom of a molecule by Lewis structure system; in instance of more than one arrangement of electron pairs for a molecule, one can assess the likelihood of each arrangement by checking these on the atoms; the one that minimizes the number and magnitude is usually the most stable arrangement of the compound
equation for formal charge
formal charge= V - N sub nonbonding - 1/2N sub bonding; where V is the normal number of electrons in the atom’s valence shell, N sub nonbonding is the number of nonbonding electrons in the atoms valence shell, N sub bonding is the number of bonding electrons (double the number of bonds because each bond has two electrons)
difference between formal charge and oxidation number
formal charge underestimates effect of electronegativity differences, whereas oxidation numbers overestimate its effect, assuming that the more electronegative atom has 100% share of the bonding electron pair
resonance structures
two or more Lewis structures that demonstrate the same arrangement of atoms but that differ in the specific placement of electrons
resonance hybrid
the nature of the bonds within the actual compound of different resonance structures is a hybrid; this is the actual structure of the compound; the more stable the structure of one resonance structure the more it contributes to the resonance hybrid
rules for determining stability via formal charges
Lewis structure with small/no formal charges is preferred over that with large ones; one with less separation bt opposite charges is preferred over one with large separation of opposite charges; one in which negative formal charges are placed on more electronegative atoms is more stable than one in which the negative formal charges are placed on less electronegative atoms
valence shell electron pair repulsion theory (VSEPR)
uses Lewis dot structures to predict the molecular geometry of covalently bonded molecules; states that the three-dimensional arrangement of atoms surrounding the central atom is determined by the repulsions between bonding and nonbonding electron pairs in the valence shell of the central atom
steps to predict the geometrical structure of a molecule using the VSEPR theory
-draw Lewis dot structure; -count total number of bonding and nonbonding electron pairs in valence shell of central atom; -position electron pairs around central atom so they are as far apart as possible
geometry and angles of pairs in VSEPR theory
2 regions of electron density: linear with 180 degrees between electron pairs; 3 regions: trigonal planar with 120 degrees between; 4 regions: tetrahedral with 109.5 degrees between; 5 regions: trigonal bipyramidal with 90, 120, and 180 degrees between; 6 regions: octahedral with 90, and 180 degrees between
electronic geometry
describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs; molecules with certain number of electron pairs (whether bonded or not) around them have the same electronic geometry as each other
molecular geometry
describes the spatial arrangement of only the bonding pairs of electrons
ideal bond angle
associated with electronic geometry; if all electron pairs (bonded or not) exerted same force and created equal angles to each other
polarity of molecules
a molecule containing only nonpolar bonds must be non
molecular orbital
formed when two atoms bond to form a compound, describes the probability of finding the bonding electrons in a given space; obtained by combining the wave functions of the atomic orbitals; the overlap of two atomic orbitals describes the molecular orbital
bonding orbital
formed when the signs of two atomic orbitals that overlap to form molecular orbitals are the same
antibonding orbital
formed when the signs of the two atomic orbitals that overlap to form molecular orbital are different
sigma bond
when orbitals overlap head-to-head this bond is formed; allows for free rotation about axis because electron density of the bonding orbital is a single linear accumulation between the atomic nuclei
pi bonds
when orbitals overlap in a parallel fashion this bond forms, does not allow free rotation because the electron densities of the orbitals cannot be twisted in such a way that allows continuous overlapping of the clouds
intermolecular forces
defines the strength of weak electrostatic interactions that atoms and compounds participate in; can impact some physical properties; weakest are London (dispersion) forces, then dipole-dipole interactions, and strongest are hydrogen bonds (there is no actual sharing or transfer of electrons); note that even the strongest is only about 10% as strong as covalent
London dispersion forces
a type of van der Waals force: result of rapid polarization and counterpolarization of the electron cloud and formation of short-lived dipole moments; the dipoles react with neighboring compounds inducing formation of more dipoles; end will be temporarily negative causing end of near molecule to be temporarily positive and other end of same molecule to be briefly negative and causes other molecules to be polarized; only relevant at small distances
Dipole-Dipole interactions
different from London dispersion forces only in time frame. Last longer. same concept; dipole lines up with opposite dipole
Hydrogen bond
dipole-dipole interaction that may be intra- or intermolecular; hydrogen essentially acts as a naked proton that does not carry almost any of the electron cloud; high boiling points compared to molecules of similar weights
