Transcript Chapter 6

Chapter 4
Chemists have found it convenient to represent elements,
especially when discussing chemical bonding, using a system
devised by G.N. Lewis, called Lewis Dot Symbols
Let’s review the names of families in the Periodic Table
Element Family Names:
•Alkali Metals – First Column
•Alkaline Earth Metals – Second Column
•Halogens – Next to last column
•Noble Gases – Last Column
Chemical Bonds – Force that holds atoms together in a
molecule.
Types of Chemical Bonds:
1. Ionic Bonds – Assume that every atom wants to have a
filled valence level (2 electrons for the first and 8 for each
of the others).
•
Then each atom will try to add or lose electrons to
achieve this. This is called the Octet Rule.
•Loss or gain depends on which is easier. (The number of
valence electrons in an atom is the # on top of the column
(before the A or B).
•Except for H and He, if the # of valence electrons is less
than 4, it will lose that number of electrons.
•If it is greater than 4, it will gain enough to make it eight.
•H either gains or loses 1 electron, depending on the other
atom. More about this later. He does not do anything. It is
chemically inert (it does not bond to any other atom)
A) Ion – Any charged atom (i.e., one that has gained or lost
electrons
B) Cation – A + charged ion
C) Anion – A - charged ion
An atom cannot gain or lose electrons without other atoms
near to accept or give these electrons. Hence when a
cation is formed, an anion is also formed (or perhaps more
than one of each type). These then attract each other
(opposite charges attract) and form the force we call an
Ionic Bond. Ionic bonds between atoms can only occur
between a metal and a non-metal atom.
•Covalent Bonds – Chemical Bonds in which 2 atoms share
exactly 2 electrons. An example is F2.
•Sometimes 2 atoms will share 4 electrons. This is 2
covalent bonds between the same atoms, always referred to
as a double bond (the word covalent is not necessary,
because there is no such thing as a double ionic bond).
• Sometimes 6 electrons are shared, which forms a triple
bond. 6 is the maximum # of electrons that can be shared
by 2 atoms. Covalent bonds occur between 2 non-metal
atoms or sometimes between a metal and a non-metal. Two
metal atoms never bond together. Why?
G.N. Lewis also developed symbols to illustrate chemical
bonds. A dash is used to represent a single bond (i.e. 2
electrons) and dots to represent the remaining valence
electrons not involved in bonding). ═ is used to represent a
double bond and ≡ is used to represent a triple bond.
Two Types of Covalent Bonds:
1. Non-polar – When 2 identical atoms bond
together, each has an equal attraction for the shared
electrons, therefore there is no net gain or loss of
electrical charge by either atom. One end of the bond
is identical to the other.
2. Polar – Whenever 2 different atoms covalently
bond together, one atom will have a stronger attraction
for the shared electrons than the other, thus one side
appears to gain some negative charge and the other
side seems to lose negative charge (become positive).
The 2 ends of the bond are different in charge and,
similar to a magnet, we say that each end is a pole (in
this case a positive and negative pole (like a battery),
hence the name polar covalent bond.
Sometimes molecules that have polar bonds are, as a
whole non-polar compounds, because of symmetry. We
won’t discuss how to identify these except in a couple of
very important examples. Any compound in which all
bonds are non-polar, will be non-polar as a molecule.
There is a special class of compounds that we will discuss
in more detail later in the semester, called hydrocarbons.
These are compounds that contain C and H and nothing
else. All bonds are either non-polar CC bonds or
slightly polar CH bonds. Nevertheless, all these
compounds (and there are thousands of them) are nonpolar, because of symmetry.
•Compounds composed of only covalent bonds are
considered to be molecular compounds and compounds
composed of ionic bonds are considered to be ionic
compounds.
•Ionic compounds can either be simple binary compounds
composed of ionic bonds between atoms
•or more complicated molecules, where one or both ions is
actually composed of 2 or more atoms covalently bonded but
forming an ion. These are called polyatomic ions. Some of
the more common polyatomic ions are listed in table 4.4 on
page 106:
04_T04.JPG
We will learn a few of the most common
( hydronium, ammonium, hydroxide, nitrate,
sulfate, cyanide, carbonate, bicarbonate and
phosphate). Note that all, except two, are
anions.
•A compound between 2 elements is called a binary
compound. For compounds between a metal and non-metal,
we can predict the formula of the compound that will form.
We simply assume each will form the most likely ion
(remember the Octet Rule).
•Then the value of the charge on the cation (without the sign)
becomes the subscript for the anion in the compound
formula, while the value of the charge on the anion becomes
the subscript for the cation.
•We simply cross-exchange.
•If both subscripts can be divided evenly by the same
number, we do that.
•Anytime the subscript is one, we don’t write it.
Let’s do some examples:
Na and Cl - Na has 1 valence electron which it loses and
becomes Na+1 while Cl has 7 valence electrons, so it gains 1
and becomes Cl-1. The 1 on Na becomes the subscript for Cl
while the 1 on Cl becomes the subscript for Na. The formula
for the compound is NaCl (1 is never written).
Ca and Cl – Ca has 2 valence electron which it loses and
becomes Ca+2 while Cl has 7 valence electrons, so it gains 1
and becomes Cl-1 (gains 1 electron to get 8). The 2 on Ca
becomes the subscript for Cl and the 1 on Cl becomes the
subscript for Ca. Thus the formula for the compound is CaCl2
1
2
Mg and SO4-2 – Mg has 2 valence electrons which it loses
and becomes Mg+2. the sulfate already has a charge of –2.
The 2 from Mg becomes the subscript for SO4-2 and the 2
from the sulfate ion becomes the subscript for Mg. Thus the
formula for the compound becomes Mg2(SO4)2. Note that
when there are more than one of a particular polyatomic ion,
its formula is placed inside parentheses. But we are not
through yet. The subscript on both the sulfate and the Mg
can be evenly divided by 2, so we do so. In each case the
result is 1. Thus the final formula is MgSO4. NOTE: No
parentheses.
Let’s try one more:
Al and O
2
3
Naming Binary Compounds or ionic compounds
with Polyatomic ions.
Binary Compounds – Name the metal first followed by
the non-metal name (dropping endings such as, “ygen”,
ogen, “ur”, “ine”, “ic” or “orous” and replacing with “ide”.
Examples – Sodium chloride, beryllium oxide, aluminum
nitride.
Ionic compounds with polyatomic ions. Simply name
the metal followed by the anion polyatomic name
unchanged. If the compound contains the ammonium ion
with a non-metal element, follow the non-metal rule
above.
Examples – sodium sulfate, magnesium carbonate,
ammonium phosphate.
NOTE: It makes no difference how many of each atom or
ion is present in the compound, for these cases. For now,
we will not worry about binary compounds involving
transition metals.
The following is discussed on pages 154 and 155 in Chapter 6 of
your book.
When molecules are mixed together, they tend to attract
each other, sometimes strongly and sometimes weakly.
With polar and ionic compounds it is easy to understand.
The positive end of one molecule attracts the negative end
of its neighbor. These attractions are what cause solids
and liquids. With ionic compounds this is very strong.
Ionic compounds tend to be solid at room temperature and
at high temperatures.
•In gases, the attractions are extremely weak. The stronger
the attractions, the more likely it is that the substance will be
a solid at room temperature and the higher the melting and
boiling points will be.
•With polar molecular substances, these attractions are
much weaker (called Dipole Forces), than in ionic
compounds, hence lower melting and boiling points, with
one exception. We will return to this in a minute.
•With non-polar compounds these attractions are
extremely weak, (called Dispersion Forces or
London Forces) hence many of these are gases or
liquids at room temperature.
The exception to the polar molecular situation arises whenever
H is directly bonded covalently to N, O or F in the compound.
Then the polarity is so strong that the intermolecular attraction
is much stronger than would be expected. This is called
Hydrogen Bonding.
Intermolecular Forces and the
States of Matter
Hydrogen bonds: When a
hydrogen atom is
covalently bonded to a
highly electronegative atom
like nitrogen, oxygen, or
fluorine (N,O,F), it can
exhibit an additional polar
attraction. This attraction is
called a hydrogen bond.
© 2010 Pearson Prentice Hall,
Inc.
6/29
The best example is H2O. Very strong attractions:
•Much higher MP and BP than expected
•Solid is less dense than liquid (ice floats)
•DNA
Without these special properties, life on earth as we know
it, could not exist. We will discuss this more at the end of
the semester.
The following can be found on pages 150-152 in your
text.
We have mentioned that substances can change state
from solid to liquid (melting) or vice versa (freezing) or
from liquid to gas (vaporization) or vice versa
(condensation). These are phase changes. They are
physical changes. One other physical change is worthy
of mention. Sublimation occurs when a solid vaporizes
directly without ever becoming a liquid. Best example is
dry ice.