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Chapter 4 - Compounds and Bonding
(Chapter 4 in 1st Edition)
In their natural state, most elements are not found in
form.
Instead, elements are almost always found as combinations
of atoms from the same or different elements. These
. Within each
combinations are called
together.
molecule the elements are
Do you know of any substances that contain only pure
elements?
Are they ATOMS or MOLECULES?
Give some examples of compounds around you:
Name
Name
Name
Name
Name
Formula_______________
Formula_______________
Formula_______________
Formula_______________
Formula_______________
How do the elements bond together?
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Bonding
Elements
electrons.
by exchanging/sharing valence
The valence electrons occupy the valence shell which is
energy level (last shell).
the
Electrons found in the inner shells are called
electrons.
Valence electrons are primarily responsible for the
chemical behaviour of the element
Electron Dot Structures (G.N. Lewis)
An Electron dot structure is a convenient way to
represent the valence electrons.
When writing electron dot structures, the valence
electrons are placed as single dots around the symbol of
the element.
e.g. Write the electron dot structures for the following
elements:
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N
O
Ar
S
Octet Rule.
Consider the electron configurations of the noble gasses.
He (#e = 2)
E. C. =
Ne (#e = 10)
E. C. =
Ar (#e = 18)
E. C. =
Kr (#e = 36)
E. C. =
Xe (#e = 54)
E. C. =
Rn (#e = 86)
E. C. =
Considering that atoms of most of the elements react to
form compounds except the noble gasses, it would seem
that the noble gasses have a particular stability associated
with them. This is attributed to fact that their valence
shell contains an octet of electrons (except for helium).
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This is known as the Octet Rule.
Compounds are formed when two or more different
elements are bonded together.
Ionic bonds are formed when elements lose or gain
valence electrons.
Covalent bonds are formed when elements share
electrons.
By losing, gaining, or sharing electrons, atoms in
compounds can achieve an octet of valence electrons.
The octet rule gives us a good understanding as to why
certain compounds form and others don’t.
e.g.
NaCl
F2
Ionization Energy.
In the previous example, why does sodium give up an
electron to chlorine? Whether or not an atom gains or
loses electrons is dependent on its ionization energy.
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The ionization energy of an atom is the amount of energy
required to remove an electron from the valence shell of
the atom.
Because metals (ie. Na) have lower ionization energies
than non-metals (ie. Cl), metals will lose valence
electrons more easily than non-metals.
Ions and Ionic Compounds
When atoms lose or gain electrons they form ions.
Consider both the Na and Cl atoms.
Na
Na+
Cl
Cl-
Sodium loses an electron and becomes a positively
charged ion called a cation.
Chlorine gains an electron and becomes a negatively
charged ion called an anion.
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Ionic compounds are formed from cations and anions.
The ions are held together by a strong electrical attraction
between the opposite charges, called an ionic bond.
As with all chemical equations, they must always be
balanced in terms of number of atoms and charge.
Na + Cl
Na Cl
Mg + Cl
Mg Cl2
Write the correct ionic formula for compound that formed
between the following pairs of ions:
Ca2+ and BrAl3+ and P3Al3+ and O2Naming Ionic Compounds.
Of the Representative elements, the cations are named
like the element and the anions are named by changing
the end of the element name to –ide.
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e.g.
Na ion =__________
O ion =__________
The ionic compound formed between them is called
Ba ion =__________
N ion =__________
The ionic compound formed between them is called
Unlike the metals of the representative elements, the
transition metals can form two or more kinds of cations.
In order to distinguish between the different cations, a
roman numeral matching the charge is placed in
parentheses after the element name. An older system
changes the ending of the latin name of the element to
-ous for the lowest charged cation and –ic for the higher
charged cation.
eg.
Fe (II) or
Au (I) or
Fe (III) or
Au (III) or
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Covalent Bonds
When two or more non-metals combine, covalent
molecules result. Unlike in ionic compounds, covalent
compounds are held together by sharing electrons. This
sharing of electrons is known as a covalent bond. Atoms
held together by covalent bonds are called molecules.
Recall F2:
Fluorine is called a diatomic molecule. Hydrogen,
nitrogen, oxygen, chlorine, bromine, and iodine are also
diatomic molecules. Why do these exists as diatomic
molecules?
Typically, the number of covalent bonds a non-metal
forms is equal to the number of electrons needed to
acquire a noble gas arrangement.
e.g.
Carbon has four valence electrons. It needs four more
electrons to achieve an octet. Therefore it can share four
electrons from another non-metal to form four bonds.
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H
H
C
H
H
This molecule is known as methane.
Be aware that there are exceptions to the octet rule. In the
hydrogen molecule (H2), each hydrogen atom is sharing
one electron for a total of two electrons. Some atoms
(typically P and S) can have more than an octet of
electrons while some elements, notably Be, B, and Al,
have less than 8 electrons around them. This is not
surprising considering their electron dot structures.
BeCl2
AlCl3
BF3
Multiple Bonds.
In the molecule F2, the two electrons shared between the
fluorine atoms represent a single bond. The pair of
electrons of a covalent bond is usually represented as a
line joining the two atoms.
e.g.
F F
or
F F
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If a two atoms share two or three pairs of electrons, a
double bond and triple bond respectively, will result.
Atoms that usually from multiple bonds are C, O, N, P,
and S. Generally, hydrogen and the halogens only form
single bonds. Multiple bonds form when single bonds
fail to complete the octet for the atoms. Electron pairs
that are not shared between atoms are called lone pairs.
Steps for Writing Electron-Dot Structures for
Molecules (Very Important).
1.
Determine the central atom. Typically this is the one
atom that is different from the others. Unless you are
given no choice, oxygen is never the central atom.
Conversely, carbon is almost always the central atom.
Hydrogen is never the central atom.
2.
Calculate the total number of valence electrons.
Remember: the number of valence electrons for a
given element is equal to its group number.
3.
Connect the atoms to the central atom using single
bonds (pairs of electrons).
4.
Subtract the number of bonding electrons from the
total number of valence electrons. Distribute these
remaining electrons around the atoms and try to give
each atom an octet of electrons.
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5. If an octet of electrons is not accomplished for each
atom, form a multiple bond by moving a lone pair from an
adjacent atom to the atom that needs an octet.
Write the electron dot structure for each of the following
molecules
SiH4
N2
HCN
C2H4
Polar Covalent Bonds.
You have seen how a pair of electrons is shared between
two atoms in a covalent single bond (e.g. F2). In the case
of F2, because the two atoms are identical (F), the
electrons are shared equally. However, if the two atoms
are different (e.g. HCl) the electron pair will not be shared
equally. This will give rise to a polar covalent bond.
Polar covalent bonds are represented by the symbols δ–
).
and δ+ or by dipole vectors (
δ δ
+
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-
H-Cl
H-Cl
The electrons are shared unequally because different
atoms have different electronegativities. The
electronegativity of an atom is its ability to attract the
shared electrons and
this is quantified by a number called an electronegativity
value.
As you go from left to right in the Periodic Table, the
electronegativity values increase with fluorine having the
highest value. As you go from top to bottom in the
Periodic Table, the electronegativity values decrease. By
comparing the electronegativity values of two atoms, we
can determine if the bond between them is ionic, polar
covalent, or covalent.
covalent
polar covalent
ionic
0-0.3
0.4-1.7
>1.7
electronegativity difference
Determine whether the bonds in the following molecules
are ionic, polar covalent, or covalent.
KBr
HCN
N2
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In general:
1. A bond between the same atoms or atoms with the
similar electronegativities is a covalent (nonpolar)
bond.
2. A bond between two different nonmetals having
different electronegativities is a polar covalent
bond.
3. Typically, a bond between a metal and non-metal is
ionic.
Naming Covalent Compounds
Name the following covalent compounds.
a) NCl3
b) SiBr4
c) PCl5
d) SO3
Write the formulas for the following covalent compounds.
a) Oxygen difluoride
b) Boron trifluoride
c) Dinitrogen trioxide
d) Sulfur hexafluoride
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Polyatomic Ions.
Polyatomic ions are a charged (cationic or anionic) group
of atoms usually made up of nonmetals such as C, Cl, S,
P, or N bonded to oxygen.
IMPORTANT: You must know (i.e. memorize) the
formulas and names of the polyatomic ions.
prefix/suffix
Carbon
hydrogen or bi-
Nitrogen
Phosphorus
Sulfur
ClO4-
per-
-ate
Chlorine
perchlorate
CO32-
ClO3-
NO3-
PO43-
SO42-
carbonate
chlorate
nitrate
phosphate
sulphate
HPO42-
HSO4-
HCO3-
bicarbonate
-ite
ClO2-
chlorite
NO2nitrite
hydrogen sulphate
or bisulphate
PO33-
SO32-
phospite
sulphite
HSO3-
hydrogen or bi-
hypo-
hydrogen phosphate
or biphosphate
hydrogen sulphite
or bisulphite
ClOhypochlorite
Other common polyatomic ions:
NH4+
ammonium
OHhydroxide
CNcyanide
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Write the formula for the polyatomic ion in each
compound and name the compound.
a) KOH
b) NaNO3
c)
CuCO3
d) NaHCO3
e)
BaSO4
Polyatomic ions are charged, but where is the charge? To
find out where the charge resides, we need to assign
formal charges to each atom in the polyatomic ion.
Chemist use formal charges as a means of ‘bookkeeping’
for the valence electrons. Consider the cyanide CN- ion.
Does the negative charge reside on the carbon or nitrogen
atom? To assign a formal charge, follow these rules:
1.
2.
3.
All the unshared electrons (lone pairs) are assigned
to the atom on which they are found.
Half the bonding electrons are assigned to each
atom in the bond
The formal charge of an atom equals the number of
valence electrons in the isolated atom, minus the
number of electrons assigned to the atom in the
electron-dot structure (steps 1 and 2).
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Using these rules, we can now find out where the negative
charge resides in the cyanide ion.
Sometimes there is more than one way to write the
electron-dot structure. This gives rise to resonance
structures. Resonance structures are drawn by moving
electrons and not atoms. Consider the ozone molecule,
O3. Draw the electron-dot structure and assign formal
charges.
Draw as many electron dot structures as you can for the
sulphate anion, SO42-
It is important to remember that molecules do not
oscillate
between resonance structures but are a weighted average
of the resonance structures.
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In the case of the sulphate anion, more than one resonance
can be drawn. Are they all equal contributors to the
overall structure of the anion? Is one structure better than
the others? In order to find the best structure, follow
these rules:
1.
2.
3.
The best resonance structures are those in which all
atoms have eight electrons in its valence shell. This
rule is INFALLIBLE for the first row elements (Be,
B, C, N, O, F). Under NO circumstance should any
of these atoms ever have more than eight electrons in
its valence shell.
Resonance structures where atoms are neutral are
better than those where atoms carry a formal charge.
When atoms do carry a charge, electronegative atoms
should carry the negative charge and electropositive
atoms should carry the positive charge.
Shapes of Molecules
The three dimensional shape of molecules is very
important in our understanding of how molecules interact
with biological entities such as enzymes and how our
sense of taste and smell function. A very simple and
effective way to determine the shape of a molecule is by
the valence shell electronpair repulsion (VSEPR) theory.
VSEPR theory arranges the electron pairs (bonding and
nonbonding) around the central atom in a way that
minimizes electric repulsion.
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Using VSEPR to predict molecular shape.
1) draw the electron dot structure
2) count the electron pairs (bonding and
nonbonding) around the central atom.
3) arrange the electron pairs to minimize
electrostatic repulsion.
4) determine the shape of the molecule from its
coordination number.
Coordination number (C.N.)= #bonded atoms +
#nonbonding electron pairs (lone pairs).
bonding
pairs
2
3
nonbonding
pairs
0
0
C.N.
Shape
2
3
2
4
3
1
0
1
3
4
4
2
5
2
0
4
5
6
0
6
linear
trigonal
planar
bent
tetrahedral
trigonal
pyramidal
bent
trigonal
bipyramidal
octahedral
Angle
θ
180
120
120
109
109
109
N/A
N/A
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Use VSEPR to predict shapes of the following molecules
and ions:
a) CF4
b) CO32-
c)
NCl3
d) NO2-
Polar and Nonpolar Molecules.
Similar to covalent bonds, molecules can also be polar or
nonpolar. When a covalent bond is polar, the more
electronegative atom has a partial negative charge and the
less electronegative atom has a partial positive charge (δor δ+ respectively). This separation of charge is called a
dipole. If the dipoles over the whole molecule cancel out,
the molecule is nonpolar. If the dipoles do not cancel out,
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then the molecule is polar. Consider the bonds in the
following molecules. Are they polar or non-polar? Is the
molecule polar? Important – just because a molecule
contains polar bonds does not mean that the molecule is
polar.
F2
CO2
H2O
CCl4
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Important Concepts from Chapter 4
• Atoms and Molecules
• Valence Electrons and Core Electrons
• Electron Dot Structures (for Atoms and Molecules)
• Octet Rule
• Ionic Bonds and Ionic Compounds (Nomenclature)
• Covalent
Bonds
(Nomenclature)
and
Covalent
• Ionization Energy and Electronegativity
• Polyatomic Ions
• Formal Charge
• Molecular Shapes
• Polarity of Bonds and Molecules
Molecules