Chapter 6 Study Guide - Chemical Bonding
(INTRAMOLECULAR
FORCES)
Chemical Bond – forces holding atoms together to form a molecule (Intramolecular) By examining the forces (electronegativity
difference – see p. 151) between
atoms in a compound, you can determine if a molecule is covalent or ionic: The three types of bonds are:
Non-polar covalent bonds - zero – 0.3 electronegativity difference (p. 163) (H2, N2,
O2)
Polar covalent bonds - between 0.3 and 1.7
electronegativity difference (p. 163)
(HCl, H2O, CO2)
Ionic bonds - greater than 1.7
electronegativity difference (p. 162) (NaCl, MgCl2)
Ionic Bonds - Cations are Positive ions, Anions are Negative ions – Wow,
great general question!
-Cathodes are
negative electrodes, Anodes are positive electrodes – great general
question!
-Ionic compounds exist as crystal lattice structures. One atom donates its
electron to another atom and they “hang out
together” because of the positive/negative attraction.
-Metals and non-metals often react to form ionic
compounds. If the
two elements are from opposite sides of the periodic
table, it is most likely an ionic
compound (NaCl is a good example)
1. The first element is positive and the second one is negative. Look at the column the first element is in and put a plus charge next to it as needed. (see examples to the right)
2. The second element is negative and it has a
full octet of electrons surrounding it.
Put it in brackets and put a negative charge next to it as needed. (see
examples to the right)
Covalent Bonds - A sharing of electrons is
called a covalent bond. If the
electronegaitvity difference between two atoms is 1.7 or less, then it is
probably a covalent bond. Look on p.
151 for a table of electronegativities.
There are two types of covalent bonds:
Non-polar covalent and polar covalent.
Use the electronegativity chart to determine which one you have.
Polar vs. Non
Polar covalent bonds: A polar bond is a bond in
which the electron(s) tend to favor one atom over another. For example, in HCl, the electron would
prefer to be with Cl than with H.
Therefore, the Cl is slightly more negative and the H is slightly more
positive.
This is called a dipole moment and is indicated by a
partial positive or a partial negative sign.
For example, you could indicate the dipole moment on HCl as
follows: HCl (fig 6-3 p. 163)
A non-polar covalent compound would be like H2,
N2, O2…etc because they have no net electronegativity
difference. (fig 6-3 p. 163)
Rules/Tips for Drawing Lewis
Structures: (COVALENT MOLECULES ONLY) 1.7 or less
1. Count up the TOTAL valence electrons for your molecule or polyatomic ion (remember to ADD electrons if it has a negative sign and to SUBTRACT electrons if it has a positive sign).
2. Draw a single bond between the central
atom and all of the atoms to which it is attached. The central atom is USUALLY the first
atom in a compound, but it is NEVER Hydrogen.
(Remember that each single bond counts as TWO electrons.)
3. After you have drawn your single bonds, put
the rest of the electrons (from the total number of valence electrons) around
your atoms in your molecule. Remember
that each atom must have 8 electrons around it or shared between it and another
atom. (There are, however, exceptions
to the octet rule. H, for instance,
needs only 2 electrons. See below for
other exceptions)
Exceptions to Octet Rule: H
(2e), Be (4e), Al
(6e), B (6e) , Ga (6e)
4. Remember that if you see a double bond, that
means that 4 electrons are being shared between those two atoms. If you see a triple bond, that means there
is 6 electrons being shared.
5. No atom may have 1 electron on any one of
its sides. Electrons travel in pairs,
you will never find an orbital with only one electron in a molecule, it must
always have a completely filled or an empty orbital.
Resonance (p. 175): If a covalent molecule requires a double bond on one side, but
not on the other of a central atom, that double bond may “resonate” throughout
the molecule. Take the example of SO2.
Molecular shapes (see table
6-5 on page 186) - Memorize this chart
In Summary of INTRAMOLECULAR FORCES:
INTRAMOLECULAR FORCE: The force holding individual atoms together to make
up one molecule. Molecular shapes are based on this force. Whether a molecule is ionic
or covalent. Whether a molecule is polar or
non-polar.
Bond polarity
and Molecular polarity: Any molecule which has a
central atom and some unbonded electron pairs around it. This provides UN-equal pulling of electrons
and thus, polarity. (example: Bent - H2O;
Trigonal pyramid – NH3; and Trigonal planar – SO2).
Bond polarity and No Molecular polarity: Any
molecule which has a central
atom and NO unbonded
electron pairs around it. This provides
equal pulling of
electrons and thus, no polarity. (example: Linear – CO2; Trigonal planar – BF3;
Tetrahedral – CH4).
No Bond polarity and No Molecular
polarity: Any Diatomic (N2,
H2, O2…)
A
Molecule cannot have molecular polarity unless it also has bond polarity. But a molecule with bond polarity does not necessarily have to
have molecular polarity. (see p. 190-191 if you don’t understand this
comment)
INTERMOLECULAR
FORCE: The force holding molecules together. The phase of
matter is based on this force. Boiling point/Freezing point. When you boil or freeze a substance, the
boiling point and freezing point are dependent upon the substance’s
INTERMOLECUAR forces only. For Example,
you would not break the individual bonds between hydrogen and oxygen when you
melt ice. You would merely break the
bonds between water molecules.
There are three basic INTERMOLECULAR FORCES (also
called Van der Walls Forces). They are
described on pages 189-193 and below.
London
Dispersion forces p. 193: depending upon where an electron is in its
orbit around the nucleus, a “pole” may be set up. Wherever the electron is would be more negative than where it
isn’t. Example: Hydrogen is a diatomic atom and it has 2
electrons. Hydrogen also has an
electronegativity difference of 0 (non-polar covalent). If the electrons go to one side of the
molecule, then the molecule has a positive and a negative end which would then
be attracted to other positive and negative ends of other hydrogen molecules
floating by.
Dispersion forces occur between two
otherwise nonpolar molecules. Dispersion
forces are the weakest of all three van der Waals forces. All molecules, including polar ones, have
dispersion forces. As a rule: As molar mass increases, so do the
dispersion forces (because there are more electrons) and so do the boiling
points.
Bottom line: Large dispersion forces make molecules want to “hang-on” to each other so it takes more energy to boil them away from each other. Similarly, low dispersion forces would make molecules not want to get together so it would take a good deal of COLD to get fluorine to freeze as opposed to say Iodine. Dispersion forces are the only way non-polar molecules stick.
Noble gases only have dispersion forces
and this is why they must get so COLD to condense or freeze.
Dipole-Dipole forces: Like dispersion forces, these are created by
electrons being on one end of a molecule.
Unlike dispersion forces, dipole-dipole forces are on polar molecules. This means that these molecules all ready have
those electrons shoved to one end of the molecule. An example is PCl3.
This Trigonal Pyramidal molecule has a lone pair of electrons on the P
atom and is therefore a polar molecule.
Hydrogen Bonding: Hydrogen has an electronegativty number of 2.1. When it is attached to Nitrogen (3.0),
Oxygen (3.5) or Fluorine (4.0), the electron is drawn more to those atoms and
thus Hydrogen is left with a partial positive charge and they with a partial
negative charge. A good example is our
old polar friend dihydrogen
monoxide: H2O.
Since hydrogen is so small also, it tends
to hug up next to the negative charged N, O, or F of another molecule. This is called hydrogen bonding (see page
192 fig 6-28.)
Carbon has a electronegativity of 2.5 and
is therefore close to that of hydrogen so the hydrogen bonds are not as strong.
Hydrogen bonding is about 10 times a
strong as regular dipole-dipole forces.
Therefore, a molecule of HF has an incredibly high boiling point, and so
does water. Molecules with hydrogen
bonding also have dipole-dipole and dispersion forces.
NOTE:
H2 does not exhibit hydrogen bonding. Only Dispersion forces. Why?
Because Hydrogen-Hydrogen bonds have equal electronegativity pulls so
there is no dipole effect on diatomic hydrogen. WATCH OUT, THIS IS OFTEN A TRICK QUESTION ON TESTS! Just because it is hydrogen doesn’t mean it
will form an intermolecular hydrogen bond!
Instructions: I will assign each homework as we go along. Make sure that you indicate which homework part you are doing on your paper (sub-headings!!!!!)
I will collect homework 1-4 on the day of the quiz. Be sure to do ALL of the problems – both the ones from the book and the ALSO ones I wrote.
Homework Part 1: Is
it ionic or covalent bonding? Electron
dots for elements
Section review p. 163 #3; Problems #33 & #37 p. 196
Practice #1-2 p. 172, Practice #1-2 p. 174, Section Review #4 p. 175 (these are all covalent) Problem #38 p. 197 (ionic), Problems #39, 41 & 42 p. 197 (all covalent)
Also: Draw the Lewis dot structure for the following ionic compounds:
Aluminum Chloride
Potassium Chloride
Sodium Oxide
Aluminum Sulfide
Homework Part 3: Predicting the Shape of a molecule using
VSEPR
Practice p. 185 #1, Practice p. 187 #1, p. 196 #24, P. 197-198 #43, 45, 46, 48, 49, 51, 57
ALSO draw the structure
and name the shapes for the following molecules:
H2,
H2O, HBr, CH3NH2 (Where C is bonded to N and 3
hydrogens. N is bonded to C and 2
hydrogens), CO2, H2CO, C2H2, C6H6,
CH3Cl, NH3, CH4, O2, AlH3,
CH3OH
Homework Part 4: Intermolecular forces and determining
polar/non-polar molecules.
p.
197 # 47. Also do problem #48 only tell
me if the bonds of those are polar or non-polar.
ALSO Look at the following molecules: , CH3Cl,
O2 NH3 H2O PCl3 BF3 H2
-Which
ones have polar bonds?
-Which
ones have molecular polarity?
-Which
ones have ONLY dispersion forces?
-Which
ones have dipole-dipole forces?
-Which
ones have hydrogen bonding forces?