Chapter 9 Flashcards
__________ __________ predict chemical bonding in nature.
Lewis structures
Step 1 for writing a Lewis structure for a molecular compound:
- Write the correct skeletal structure for the molecule
Step 2 for writing a Lewis structure for a molecular compound:
- Calculate the total # of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule.
Step 3 for writing a Lewis structure for a molecular compound:
Distribute the electrons among the atoms, giving octets to as many atoms as possible.
Step 4 for writing a Lewis structure for a molecular compound:
If any atoms lack an octet, form double or triple bonds as necessary to give them octets.
“At the ends”.
Terminal
The only way to determine the skeletal structure of a molecule with absolute certainty is to __________ _____ _________ ________.
Examine it’s structure experimentally
__________ atoms are always terminal.
Hydrogen
__________ atoms must form at least two bonds.
Central
Put the more __________ __________ in terminal positions.
Electronegative elements
Put the __________ electronegative elements (other than hydrogen) in the central position.
Less
If you are writing a Lewis structure for a __________ _____, the charge of the ion must be considered when calculating the total # of electrons.
Polyatomic ion
_____ one electron for each negative charge.
Add
__________ one electron for each positive charge.
Subtract
The Lewis structure for a __________ _____ is usually written within brackets.
Polyatomic ion
For the polyatomic ion NH(subscript 4)^+ , there are 9 electrons. Would you add an electron to make the total 10, or subtract and make it 8?
Subtract and make it 8
This is used when two or more valid Lewis structures can be drawn for the same compound.
Resonance
An electron book-keeping system that allows us to discriminate between alternative Lewis structures.
Formal charge
Where there are two or more valid Lewis structures for the same molecule, we find that, in nature, the molecule exists as an _________ of the structures.
Average
Two or more valid Lewis structures that are shown with double-headed arrows between them to indicate that the actual structure of the molecule is intermediate between them:
Resonance structures
The actual structure of a molecule that is intermediate between two or more resonance structures:
Resonance hybrid
The only structure that actually exists is the __________ structure-the individual resonance structures do not exist and are merely a convenient way to describe the actual structure.
Hybrid
The concept of __________ is an adaptation of Lewis theory that helps account for the complexity of actual molecules.
Resonance
In Lewis theory, electrons are ________ either on one atom or between atoms.
Localized
The _________ of electrons lowers their energy; it stabilizes them.
Delocalization
__________ __________ is a fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures.
Formal charge
The charge that an atom would have if all bonding electrons were shared equally between the bonded atoms.
Formal charge
The calculated charge for an atom if we completely ignore the effects of electronegativity.
Formal charge
Formal charge =
(# of valence electrons) - [# of nonbonding electrons for element +1/2(number of bonding electrons)]
The sum of all formal charges in a neutral molecule must be _____.
Zero
The sum of all formal charges in an ion must equal _______.
The charge of the ion
Small (or zero) formal charges on individual atoms are __________ than large ones.
Better
When formal charge cannot be avoided, negative formal charge should reside on the __________ __________ atom.
Most electronegative
A technique in which X-rays are scattered from crystals of the molecule of interest.
X-ray crystallography
In 1989, researchers used X-ray crystallography to determine the structure of a molecule called what?
HIV-protease
Without HIV-protease, HIV cannot spread in the human body because the virus cannot __________.
Replicate
The working part of a molecule.
The active site
Models that predict how atoms bond together to form molecules.
Bonding theories
Many AIDS patients are still alive today because of:
Protease inhibitors
Without chemical bonding, there would be only how many different kinds of substance?
91 (# of elements on the periodic table)
Chemical bonds form because they lower the __________ __________ between the charged particles that compose atoms.
Potential energy
If atomic/electron interactions result in an overall net reduction of energy between charged particles, what forms?
A chemical bond
Metals tend to have low:
Ionization energies
Nonmetals tend to have negative:
Electron affinities
Nonmetals readily gain __________.
Electrons
Nonmetals tend to have high:
Ionization energies
Lowest potential energy = most __________.
Stable
Since metals have low ionization energies, they tend to lose electrons __________.
Easily
In the __________ __________ model, all of the atoms in a metallic lattice pool their valence electrons.
Electron sea model
Lewis theory focuses on:
Valence electrons
The easiest way to calculate lattice energy is with the:
Born-Haber cycle
The magnitude of lattice energy __________ as we move down a column.
Decreases
Lone pair electrons are also called:
Nonbonding electrons
In general, double bonds are __________ and __________ than single bonds.
Shorter and stronger
Triple bonds are even __________ and __________ than double bonds.
Shorter and stronger
__________ bonds are directional.
Covalent
The fundamental units of colvalently bonded compounds are __________ __________.
Individual molecules
In covalently bonded molecular compounds, the interactions __________ molecules is much weaker than the bonding interactions __________ a molecule.
Between, within
__________ compounds tend to have lower melting and boiling points than _________ compounds.
Molecular, ionic
Intermediate in nature between a pure covalent bond and an ionic bond.
Polar covalent bond
Having a positive pole and a negative pole.
Polar
The ability of an atom to attract electrons to itself in a chemical bond (which results in polar and ionic bonds).
Electronegativity
Elecgronegativity was quantified by American chemist:
Linus Pauling
Energy required to break a bond.
Bond energy
Electrons are shared equally and exactly.
Pure covalent bond
Non-polar.
Purely covalent
