electronegativity: increases across a period, decreases down a group
electronic structure of atoms
shell model
Coulomb’s law
nonpolar covalent bond: valence electrons are shared between atoms of similar electronegativity
e.g. carbon and hydrogen (even though carbon is slightly more electronegative than hydrogen)
polar covalent bond: valence electrons are shared between atoms of unequal electronegativity
atom with higher electronegativity will develop a partial negative charge relative to the other atom in the bond
in single bonds, greater differences in electronegativity lead to greater bond dipoles
all polar bonds have some ionic character
difference between ionic and covalent bonding is not distinct but rather a continuum
generally, bonds between a metal and nonmetal are ionic, and bonds between two nonmetals are covalent
the best way to characterize the type of bonding is by examining the properties of a compound
in a metallic solid, the valence electrons from the metal atoms are considered to be delocalized(not associated with any individual atom)
intramolecular force and potential energy
graph of potential energy versus the intranuclear distance (distance between atoms) describes interactions between atoms
equilibrium bond length: separation between atoms at which the potential energy is loewst
bond energy: energy required to separate the atoms
in a covalent bond, the bond length is influenced by:
the size of the atom’s core
bond order (i.e. single, double, triple)
bonds with a higher order:
shorter
larger bond energies
Coulomb’s law can be used to understand the strength of interactions between cations (positive ions) and anions (negative ions)
larger charges lead to stronger interactions (interaction strength is proportional to the charge of each ion)
smaller ions lead to stronger interactions (interaction strength increases as the distance between the nuclei decreases)
ionic crystal: the cations and anions in an ion are arranged in a systematic, periodic 3-D array
maximizes the attractive forces
minimizes the repulsive forces
structure of metals and alloys
metallic bonding can be represented as an array of positive metal ions surrounded by delocalized valence electrons (i.e. a “sea of electrons”)
interstitial alloys: form between atoms of significantly different radii
smaller atoms fill the interstitial spaces between the larger atoms
e.g. steel: carbon occupies the interstices in iron
substitutional alloys: form between atoms of comparable radius
one atom substitutes for another in the lattice
e.g. brass: zinc substitutes for copper
Lewis diagrams
construction of Lewis structures
1 dot represents 1 valence electron
the element with more lone pairs goes in the middle
add atoms to fulfill octet/duet rule
types of bonds
single bond: two dots or one line
double bond: four dots or two lines
triple bond: six dots or three lines
resonance: when more than one equivalent Lewis structure can be constructed
needed to provide qualitatively accurate predictions of molecular structure and properties
to determine which possible valid Lewis diagram provides the best model:
octet/duet rule: elements prefer to have filled valence shells
formal charge: difference between the number of valence electrons on free neutral atom and the number of valence electrons assigned to the atom in the molecule
there are limitations to the use of the Lewis structure model, particularly in cases with an odd number of valence electrons
VSEPR theory: uses Coulombic repulsion between electrons as a basis for predicting the arrangement of electron pairs around a central atom
molecular geometry
bonding groups
lone pairs
bond angle
linear
2
2
0
3
180°
trigonal planar
3
0
120°
tetrahedral
4
0
109.5°
trigonal pyramidal
3
1
107°
bent
2
2
1
2
< 120°
104.5°
trigonal bipyramidal
5
0
120° (equatorial), 90° (axial)
seesaw
4
1
< 120° (equatorial), < 90° (axial)
T-shaped
3
2
< 90°
octahedral
6
0
90°
square pyramidal
5
1
< 90°
square planar
4
2
90°
There are no rows in this table
relative bond energies based on bond order
relative bond lengths (multiple bonds, effects of atomic radius)
presence of a dipole moment
hybridization/hybrid atomic orbital: arrangement of electrons around a central atom and ideal bond angles
sp: 180°
sp²: 120°
sp³: 109.5°
bond formation is associated with overlap between atomic orbitals
multiple bonds: leads to formation of sigma bondsandpi bonds
overlap is stronger in sigma bonds than pi bonds, so sigma bonds have greater bond energy
the presence of a pi bond prevents the rotation of the bond and leads to geometric isomers
