Ch 6 Section 5 Flashcards
Molecular geometry is the
3-dimensional arrangement of a molecule’s atoms in space
The polarity of each bond along with the geometry of the molecule determines
Molecular polarity
Molecular polarity is the
Uneven distribution of molecular charge
Molecular polarity strongly influenced the forces that act
Between molecules in liquids and solids
A chemical formula reveals little information about
A molecules geometry
After performing many tests designed to reveal the shapes of various molecules chemists developed two different
Equally successful theories to explain certain aspects of their findings
One theory accounts for
Molecular bond angles
The other theory is used to describe the orbitals that
Contain the valence electrons of a molecules atoms
Diatomic molecules like h2 and HCl must be
Linear because they consist of only two atoms
To predict the geometries of more complicated molecules one must consider the
Locations of all electron pairs surrounding the bonded atoms
The abbreviation VSEPR stands for
valence-shell, electron-pair repulsion
The abbreviation for VSEPR refers to the repulsion between
Pairs of valence electrons of the atoms in a molecule
VSEPR theory states that repulsion between the sets of valence level electrons surrounding an atom causes these sets to be
Oriented as far apart as possible
According to the VSEPR theory the shared pairs in BeF2 will be as far
Away from each other as possible
The distance between electron pairs is maximized if the bonds to fluorine are on
Opposite sides of the beryllium afl , 180 degrees apart. Thus the molecule is linear
If we represent the central atom in s molecule by the letter A and we represent the atoms bonded to the central atom by the letter B then according to VSEPR theory BeF2 is an example of an
AB2 molecule which is linear
In an AB3 molecule the three A-B bonds stay farthest apart by pointing to the corners of an
Equilateral triangle giving 120 degree angles between the bonds
The central atoms in AB4 molecules follow the octet rule by sharing
Four electron pairs with B atoms
In AB4 molecules the distance between electron pairs is maximized if each A-B bond points to
One of four corners of a tetrahedron
Ammonia (NH3) and water (H2O) are examples of molecules in which the central atom has both
Shared and unshared electron pairs
VSEPR theory postulates that the line pair of electrons occupies space around the (ammonia atom)
Nitrogen atom just as the bonding pairs do
Thus in ammonia the electron pairs maximize their separation by assuming the
Four corners of a tetrahedron
Lone pairs do occupy space but our description of the observed space of s molecules refers to the
Positions of atoms only
The general VSEPR formula for molecules such as ammonia is
AB3E where E represents the unshared electron pair
The VSEPR formula for water is
AB2E2
For H2O VSEPR theory states that the lone pairs occupy space around the central atom but that the actual shape of the molecule is determined by the position of the
Atoms only, resulting in a bent molecule
The bond angles in ammonia and water are smaller because the unshared electron pairs
Repeal electrons more strongly than do bonding electron pairs
In VSEPR theory double and triple bonds are treated in the same way as
Single bonds
In VSEPR theory poly atomic ions are treated similarly to
Molecules
Lewis structures and VSEPR theory and molecular geometry can be used together to predict the… Of poly atomic ions as well as… With
Shapes; molecules with double or triple bonds
VSEPR troth does not reveal the relaid shop between a molecules
Geometry and the orbitals occupied by its bonding electrons
To explain the orbitals of an atom become rearranged when the atom forms covalent bonds, a different
Model, hybridization is used
Hybridization is the mixing of two or more atomic orbitals of similar energies on the same atom to
Produce new hybrid atomic orbitals of equal energies
The sp3 orbitals all have the same
Energy which is greater than that of the 2s orbitals but less than that of the 2p orbitals
Hybrid orbitals are orbitals of
Equal energy produced by the combination of two or more orbitals on the same atom
The number of hybrid orbitals produced equals the number of
Orbitals that have combined
Hybridization also explains the
Bonding and geometry of many molecules formed by group 15 and 16 elements
The linear geometry of molecules such as beryllium fluoride is made possible by hybridization involving the
S orbital and one available empty p orbital To yield sp hybrid orbitals
The trogonal-planar geometry of molecules such as boron fluoride is made possible by hybridization involving the
S orbital one singly occupied p orbital and one empty p orbital to yield sp2 hybrid orbitals
Atomic orbitals: s, p
Type of hybridization:sp
Number of hybrid orbitals: 2
Geometry:
180 degrees; linear
Atomic orbitals:s,p,p
Type of hybridization: sp2
Number of hybrid orbitals: 3
Geometry:
120 degrees: trigonal planar
The properties of molecules depend not only on the bonding of atoms but also on
Molecular geometry
Atomic orbitals:s,p,p,p
Type of hybridization: sp3
Number of hybrid orbitals: 4
Geometry:
109.5 degrees; tetrahedral
Atoms bonded to central atom: 2
Lone pairs of electrons: 0
Type of molecule: AB2
Molecular shape:
Linear
Atoms bonded to central atom: 3
Lone pairs of electrons: 0
Type of molecule: AB3
Molecular shape:
Trigonal-planar
Atoms bonded to central atom: 2
Lone pairs of electrons: 1
Type of molecule: AB2E
Molecular shape:
Bent or angular
Atoms bonded to central atom: 4
Lone pairs of electrons: 0
Type of molecule: AB4
Molecular shape:
Tetrahedral
Atoms bonded to central atom: 3
Lone pairs of electrons: 1
Type of molecule: AB3E
Molecular shape:
Trigonal-pyramidal
Atoms bonded to central atom: 2
Lone pairs of electrons: 2
Type of molecule: AB2E2
Molecular shape:
Bent or angular
Atoms bonded to central atom: 5
Lone pairs of electrons: 0
Type of molecule: AB5
Molecular shape:
Trigonal-bipyramidal
Atoms bonded to central atom: 6
Lone pairs of electrons: 0
Type of molecule: AB6
Molecular shape:
Octahedral
As a liquid is heated the kinetic energy of its particles
Increases
At the boiling point the needy is sufficient to overcome the force of
Attraction between the liquids particles
The particles pull away from each other and enter the
Gas phase
Boiling point therefore is a good measure of the force of
Attraction between particles of a liquid
The higher the boiling point the stronger the
Forces between particles
The forces of attraction between molecules are known as
Intermolecular forces
Intermolecular forces. Art in strength but are generally weaker than bonds that join
Atoms in molecules, ions in ionic compounds, or metal atoms in solid metals
The values for ionic compounds and metals are much higher than those for
Molecular substances
The strongest intermolecular forces exist between
Polar molecules
Polar molecules act as tiny dipoles because of their
Uneven charge distribution
A dipole is created by equal but opposite charges that are
Separated by a short distance
The direction of a dipole is from the dipoles
Positive pole to its negative pole
A dipole is represented by an arrow with a head pointing toward the
Negative pole and a crossed tail situated at the positive pole
The negative region in one polar molecule attracts the
Positive region in adjacent molecules
The forces of attraction between polar molecules are known as
Dipole-dipole forces
The forces of attraction between polar molecules are known as
Dipole-dipole forces
Dipole-dipole forces are short range forces, acting only between
Nearby molecules
The polarity of diatonic molecules such as ICl is determined by just
One bond
For molecules containing more than two atoms molecular polarity depends on both the
Polarity and the orientation of each bond
Because the molecule is bent the polarities of these two bonds combine to make the molecule
Highly polar
In some molecules individual bond dipoles
Cancel one another causing the molecular polarity to be zero
A polar molecule can induce a dipole in an Nonpolar molecule by
Temporarily attracting its electrons
The result is a short range intermolecular force that is somewhat
Weaker than the dipole-dipole force
Some hydrogen containing compounds have unusually high
Boiling points
These high boiling points is explained by the presence of a particularly strong type of
Dipole-dipole force
In compounds containing H-F, H-O, or H-N bonds, the large electronegativity differences between hydrogen atoms and fluorine, oxygen, or nitrogen atoms make the bonds connecting them
Highly polar
The high polarity gives the hydrogen atom a positive charge that is almost half as large as that of a
Proton
The small size of the hydrogen atom allows the atom to come very close to an
Unshared pair of electrons on an adjacent molecule
The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an electronegative atom in a nearby molecule is known as
Hydrogen bonding
Hydrogen bonds are usually represented by dotted lines connecting the hydrogen bonded hydrogen to the unshared electron pair of the
Electronegative atom to which it is attracted
Even noble gas atoms and molecules that are Nonpolar experience a
Weak intermolecular attraction
In any atom or molecule the electrons are in
Continuous motion
Thus at any instant the electron distribution may be slightly
Uneven
The momentary uneven charge creates a positive pole in one part of the atom or molecule and a
Negative pole in another
This temporary dipole can then induce a dipole in an
Adjacent atom or molecule
The two are held together for an instant by the weak attraction. Between the
Temporary dipoles
The intermolecular attract drinks resulting from the constant motion of electrons and the creation of instantaneous dipoles are called
London dispersion forces
London dispersion forces are named after
Fritz London who first proposed their existence in 1930
London forces act between all
Atoms and molecules
London forces are the only intermolecular forces acting among
Noble gas atoms and Nonpolar molecules
Because London forces are dependent on he motion of electrons their strength increases with the number of
Electrons in the interacting atoms or molecules
London forces increase with increasing
Atomic or molar mass
