Ch 10 Flashcards
Lewis structures
show us how atoms are bonded to each other. They can be used to predict molecular shapes, bond strength and bond length.
Covalent Bonds
Formed when two (or more) nonmetals share valence electrons in order to obey the octet rule.
Single bond
two atoms share a pair of electrons between them (2 shared electrons = 1 bond)
Electron-dot formulas show
- The order of bonded atoms in a covalent compound.
- The bonding pairs of electrons between atoms.
- The unshared (lone) valence electrons.
- A central atom with an octet.
Double bond
two atoms share 2 pairs of electrons between them (4 shared electrons = 2 bonds) in order to obey the octet rule
Triple bond
two atoms share 3 pairs of electrons between them (6 shared electrons = 3 bonds)
In a single bond
One pair of electrons is shared.
In a double bond
Two pairs of electrons are shared
In a triple bond
Three pairs of electrons are shared
Bond Length
- Distance between 2 nuclei
- Triple bond < Double Bond < Single Bond
Bond Strength
- Energy required to separate 2 nuclei
- Triple bond > Double Bond > Single Bond
Coordinate Covalent Bonds
bond in which there is an uneven contribution of the shared electron pairs.
As long as all the atoms in your molecule are satisfied under the octet rule (exception is H with only 2 electrons),
then you have an acceptable Lewis Dot Structure
Guide to Writing Electron-Dot Formulas
STEP 1 Determine the arrangement of atoms. (usually the first one in the formula unless the first atom is H)
STEP 2 Add the valence electrons from all the atoms.
STEP 3 Attach the central atom to each bonded atom using one pair of electrons.
STEP 4 Add remaining electrons as lone pairs to complete octets (2 for H atoms).
STEP 5 If octets are not complete, form one or more multiple bonds.
Resonance structures are
- Two or more electron-dot formulas for the same arrangement of atoms
- Related by a double-headed arrow
- Written by changing location of a double bond from the central atom to a different attached atom
- Sometimes written as a hybrid resonance structure
Molecular Shape
- Lewis Structures tell us how atoms are connected in a covalent molecule or ion
- They can also be used to predict molecular shape
Molecular shape
- determines how molecules function.
- is essential to drug design and other biochemical processes
In order to predict molecular shape, we assume that electron groups repel each other (because they are all negatively charged).
Therefore, the molecule adopts whichever 3D geometry minimized this repulsion
VSEPR
Valence Shell Electron Pair Repulsion theory
Molecules adopt shapes that
minimize the repulsion between electron groups/domains
What is an electron group/domain?
A collection of valence electrons present in a localized region about the central atom in a molecule
There are two kinds of electron groups/domains
- A pair of nonbonding electrons
- An area where there are bonding electrons (it doesn’t matter if the bond is a single, double or triple bond)
Determining Molecular Shape
First, draw the Lewis Structure for the compound
There are two important features of Lewis Structures that you must recognize:
a. The number of ATOMS bonded to the central atom (this is the number of bonding electron groups/domains)
b. The number of PAIRS of nonbonding electrons on the central atom (this is the number of nonbonding electron groups/domains
Electronegativity
A measure of the relative attraction that an atom has for the shared electrons in a bond
Trends in electronegativity
Electronegativity values increase from left to right on the periodic table
Electronegativity values increase from the bottom to the top of the periodic table
Electronegativity difference: 0.0 to 0.4
nonpolar covalent.
Electronegativity difference: 1.8 or greater
bond considered ionic.
Electronegativity difference:greater than 0.4 or less than 1.8
bond considered polar covalent
Polar covalent bonds
A covalent bond in which there is unequal sharing of electrons between two atoms
Nonpolar covalent bonds
A covalent bond in which there is equal sharing of electrons between two atoms
Bond Polarity
a measure of the degree of inequality in the sharing of electrons between two atoms in a chemical bond
If a bond is between two of the same kinds of atoms
It will be NONpolar
If a bond is between two different atoms
It will be polar to some degree (more so the further apart the two atoms are on the periodic table)
Polar Molecule
A molecule in which there is an unsymmetrical distribution of charge
Molecular Polarity
A measure of the degree of inequality in the attraction of bonding electrons to various locations within a molecule
Why do we care if the molecule is polar?
Polar molecules dissolve in polar solvents
Nonpolar molecules dissolve in nonpolar solvents
Molecules with the following geometries are nonpolar (if there is at least one polar bond in the molecule):
Linear (with 2 of the same atom bonded to a central atom) Trigonal Planar (with 3 of the same atom bonded to a central atom)
Tetrahedral (with 4 of the same atom bonded to a central atom)
Trigonal bi-pyramidal (with 5 of the same atom bonded to a central atom)
Octahedral (with 6 of the same atom bonded to a central atom)
Square planar (with 4 of the same atom bonded to a central atom)
Covalent or ionic bonds that hold a molecule together are
intramolecular forces
The attraction between molecules is an
intermolecular force
Intermolecular forces are much ___than intramolecular forces
weaker
There are different intermolecular forces
Dipole-Dipole Forces
Hydrogen “Bonds”
London Dispersion Forces
London Dispersion Forces
Formed between adjacent neutral molecules
Weak attractions between nonpolar molecules.
Caused by temporary dipoles that develop when electrons are not distributed equally.
Increase with size of atoms or molecular weight of molecules
Intramolecular Bonds
Ionic Bonds
Intermolecular Forces
Hydrogen “Bonds”
Hydrogen Bonds
Unusually strong dipole-dipole interactions that occur between molecules with H-N, H-O, or H-F covalent bonds and a lone pair of electrons on another small, electronegative atom
