Unit 1 - Chapter 4 Flashcards
Drawing Lewis structures
determine the total number of valence electrons, taking into account charge if it’s an ion
place the least electronegative as a central atom surrounded by the other atoms
determine the electron distribution by arranging the electrons between the central atom and the outer atom, then completing the octets for the outer atoms, then as lone pairs around the central atom
Molecules of second row elements
second row elements have the second energy level as their valence shell
with only s and p orbitals, their bonding capacity is limited by the octet rule
Resonance structures
when drawing Lewis structures, sometimes there are two or more acceptable diagrams for one molecule
these molecules are classified as resonance structures, and the real structure cannot be represented by a single diagram
the true molecule is a hybrid of all of the possible diagrams
Lewis structures are therefore the best approximation
Measurable properties of covalent bonds
bond length = average distance between two nuclei of bonded atoms (shorter = stronger)
bond energy = energy released as bond forms (higher = stronger)
by measuring these properties on resonance structures, all bonds have the same length and energy
Molecules of third row elements and beyond
elements with a valence shell beyond the second energy level have access to s, p, and d orbitals
their valence shell is therefore not limited by the octet rule
with these elements, the concept of formal charge allows us to draw the best structure
Formal charge
the apparent charge on atoms in a Lewis structure that arises when the atoms have not made equal contributions of electrons to the covalent bonds
once formal charges have been assigned to each atom in the Lewis structure, they can be reduced by shifting lone pairs to form double bonds
minimizing formal charge in this way does not charge the fact that a structure is a resonance structure
minimizing formal charge may also expose the presence of a resonance structure
The best Lewis structure
all atoms have a formal charge of zero
if formal charge can’t be reduced to zero, remaining charges are as small as possible
formal charges add to zero if molecule is neutral, or add to charge of ion
any negative charges are on most electronegative ion
VSEPR theory
valence shell electron pair repulsion theory
valence shell electron pairs around central atom determine the shape of a molecule due to optimal arrangement
repulsion forces between electron pairs will position the valence electron pairs into a geometric arrangement that minimizes repulsion forces
AXE formula
A = the central atom X = the number of atoms bonded to A E = the number of lone electron pairs around A
Molecular polarity
a molecule is polar when there is an asymmetric distribution of electron density towards the more electronegative atom, resulting in a permanent electrical dipole
this means that one side of the molecule has a partial positive charge, the other a partial negative
this polarity is caused by the asymmetric arrangement of polar covalent bonds
if bonds are nonpolar or polar bonds are symmetrically arranged, no electrical dipole exists and molecule is nonpolar
Valence bond theory
proposed that a covalent bond forms when two half-filled valence orbitals from two atoms overlap
the two overlapping orbitals form one bonding orbital = a volume of space between two atomic nuclei in which there is a high probability of finding the now spin paired bonding electrons
electron promotion = amount of energy used to unpair electrons, this increases the number of half-filled valence orbitals that can now bond
Problems with valence bond theory
a good bonding theory will explain bond length, energy, angles, and shape
for most molecules, the VB theory does not match experimental observations
eg. based on VB theory, the water molecule should have a bond angle of 90, however in reality, the bond angle is 104.5
Linus Pauling and orbital hybridization
when bonding, the pure atomic orbitals of the central atom are replaced with new, hybrid orbitals
hybridizing electrons from one s orbital and three p orbitals results in four identical hybrid orbitals called sp3 orbitals that are arranged in a tetrahedral pattern
hybrid orbitals are intermediate in energy to the original pure orbitals
Sigma bonds
a sigma bond is formed when two orbitals overlap in a direct head-to-head fashion
the bonding orbital formed by this overlap lies directly between and in the same plane as the two atomic nuclei
this very effective overlap results in a very strong bond
Pi bonds
pi bonds are formed by the less effective side-by-side overlap of two p orbitals
this results in a bonding orbital forming above and below the plane of the nuclei
this less effective orbital overlap results in the pi bond being weaker than the sigma bond
double bond = one sigma bond and one pi bond
triple bond = one sigma bond and two pi bonds
