Transcript of Chapter 4. Chemists have found it convenient to represent elements, especially when discussing...
- Slide 1
- Chapter 4
- Slide 2
- 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
- Slide 3
- 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.
- Slide 4
- 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)
- Slide 5
- A) Ion Any charged atom (i.e., one that has gained or lost
electrons B) Cation A + charged ion C) Anion A - charged ion
- Slide 6
- 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.
- Slide 7
- Covalent Bonds Chemical Bonds in which 2 atoms share exactly 2
electrons. An example is F 2. 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?
- Slide 8
- 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.
- Slide 9
- 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.
- Slide 10
- 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.
- Slide 11
- Slide 12
- Sometimes molecules that have polar bonds are, as a whole
non-polar compounds, because of symmetry. We wont 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.
- Slide 13
- 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 non- polar, because of symmetry.
- Slide 14
- 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 107: 2
elements
- Slide 15
- 04_T04.JPG
- Slide 16
- 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.
- Slide 17
- 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 dont
write it.
- Slide 18
- 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). Lets do some examples:
- Slide 19
- 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 CaCl 2 1 2
- Slide 20
- Mg and SO 4 -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 SO 4 -2 and the 2 from the sulfate ion
becomes the subscript for Mg. Thus the formula for the compound
becomes Mg 2 (SO 4 ) 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 MgSO 4. NOTE: No
parentheses.
- Slide 21
- Al and O Lets try one more: 2 3
- Slide 22
- 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.
- Slide 23
- 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.
- Slide 24
- 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. The following is discussed on
pages 156 and 159 in Chapter 6 of your book.
- Slide 25
- 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.
- Slide 26
- 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. Ionic Compounds Ionic bonds
& dispersion forces between molecules Polar Covalent Compounds
Dipole forces (attractions) & dispersion forces Non-polar
Covalent Compounds Dispersion forces only
- Slide 27
- 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.
- Slide 28
- 2010 Pearson Prentice Hall, Inc. 6/28 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.
- Slide 29
- Much higher MP and BP than expected Solid is less dense than
liquid (ice floats) DNA The best example is H 2 O. Very strong
attractions:
- Slide 30
- 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.
- Slide 31
- 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.
The following can be found on page 154 in your text.