CHEM 1151K
September 28, 2004
Molecules:
H2
O2
N2
F2
single bonds
double bond
triple bond
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Naming Binary Molecular (covalent)Compounds
Binary means it's composed of only 2 elements. For molecular compounds this means
two non-metals.
Names need prefixes to tell how many of atoms of each element are in a molecule.
mono
di
tri
tetra
penta
1
2
3
4
5
hexa
hepta
octa
nona
deca
6
7
8
9
10
When you list the two elements in the name, if there is only one atom of the first named
element, you leave off the prefix "mono".
For binary covalent compounds, both in the names and in the chemical formulas the least
electronegative element comes first, followed by the more electronegative element.
For the second element in the name, change the suffix to –ide.
Add prefixes to both elements, except as noted above for the first element.
Electronegativity – the ability of an atom to attract electrons in a covalent bond.
Fluorine is the most electronegative element.
The closer and element is to fluorine on the Periodic Table, the more
electronegative it is.
The general trends are that electronegativity increases as you go across any given
row and that it decreases as you go down any given column.
Naming examples:
CO
carbon monoxide
notice than one of the o’s is dropped
so it’s not monoxide
CO2
SF6
N2O5
carbon dioxide
sulfur hexafluoride
dinitrogen pentaoxide
Put a space between the two words.
Summary: To know which naming rules to use look for the following:
Two non-metals = molecular
use molecular rules, prefixes will be used
Using the Ionic compounds rules when you see:
Metal + nonmetal
Metal + polyatomic anion
Ammonium + nonmetal
Ammonium + polyatomic anion
NaCl
NaNO3
NH4Cl
NH4NO3
You HAVE to learn to recognize the names and formulas of the polyatomic ions.
Polarity in bonds – uneven sharing of electrons in a bond
100% non-polar means equal sharing of electrons. This only happens when identical
(like) atoms are bonded.
Example: the H2 molecule
H–H
This is comparable to identical twins playing tug-of-war.
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When unlike atoms are bonded they will not share the bonding electrons exactly equally.
The more electronegative atom will tend to pull the bonding electrons closer to itself and
away (but not disconnected from) the other atom. Another way of saying this is that the
“electron cloud” for the bonding electrons is “skewed” toward the more electronegative
atom.
This gives the more electronegative atom a tiny bit more negative charge in its vicinity.
It takes away negative charge from the other atom and leaves it a little bit positively
charged.
In a bond between oxygen (more electronegative) and hydrogen, the electrons in the bond
will be skewed toward the oxygen. This gives oxygen a partial negative charge (-) and
the hydrogen has a partial positive charge (+).
By “partial” charge, we mean fractional. The absolute value of the charge is between
zero and 1.
On a polar bond you can indicate the polarity with an arrow pointing toward the - end.
You also put a + on the other end. A hydrogen fluoride molecule is shown below.
Fluorine is the most electronegative element.
When atoms of similar electronegativity are bonded they almost share the electrons
equally, so chemists will still classify the bond as “non-polar”. The most important
example is the bond between carbon and hydrogen.
C-H
this bond is classified as non-polar
There is a mathematical way to find out if the electronegativities are close enough to do
this kind of classification, but we aren’t going to do it.
Lewis Dot Structures for Molecules
Lewis structures are a way of drawing molecules to show which atoms are bonded to
each other and also where the valence electrons are. Valence will appear as either
bonding electrons or lone pairs (non-bonding electrons).
Below is a step-by-step sequence to help you draw Lewis structures. It is essentially the
same as what’s in your text, in you lab manual, and what your lab instructor went over.
You just need to PRACTICE and find which wording of the steps you best understand
and then stick to that set.
1. Draw the atoms connected by single bonds with the correct connectivity.
You need to know the connectivity of the atoms (how they are joined). For our purposes
we will have one atom in the center (central atom) and any other joined to this one (the
outside atoms).
The connectivity will either be given or be simply deduced. For “simply
deduced” I mean that there will be only one way to connect the atoms.
2. Count the total number of valence electrons. This is how many will be displayed one
way or another on the structure.
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add up the number of valence electrons each atom has
for anions add one to account for each negative charge
for cations subtract one to account for each positive charge
(remember, we’re counting electrons and they are negative)
3. Subtract two electrons from the grand total in step 2 for each single bond drawn in step
1. That’s because you’ve already shown two valence electrons each time you drew a
bond. The result here will tell you how many valence electrons are left to be displayed
on the structure.
4. Distribute the remaining electrons to be shown found in step 3, first pair-wise to the
outside atoms to a maximum of 8 valence electrons on any given atom. Then distribute
to the central atom if any are left.
Notes: Distribute the electrons pair-wise as non-bonding electrons.
Don’t put any on hydrogen. A single bond connecting hydrogen to
another atom already gives the hydrogen the most valence electrons it can
have (two).
5. Move non-bonding pairs to be bonding pairs to try to give all atoms (except H) 8
valence electrons.
Note: As before, the number of valence electrons any given atom “feels” is the sum of all
the non-bonding electrons on that atom plus all the electrons it’s sharing via bonding.
We’ll add a few “exceptions” later but for now look at an example.
Water.
1. Connectivity: Oxygen is the central atom and both the hydrogens are connected to it.
H–O–H
2. Count the valence electrons. Each hydrogen has 1 and the oxygen has 6, for a grand
total of 8 valence electrons to be displayed on the structure.
3. Adjust the number to account for bonds. In step 1 we drew two single bonds which
already put 4 electrons on the picture, so subtract 4 from the grand total found in step 3.
8-4= 4 electrons left to show on the picture
4. Distribute electrons. You can’t put any more on the hydrogens, so all the remaining
electrons go to the oxygen. This also gives oxygen a nice total of 8 valence electrons.
****The desire for many atoms to have 8 valence electrons is called the Octet Rule.