Title: Chemical Bonding
1Chapter 6
2Chapter Sections
- Introduction to chemical bonding
- Covalent bonding molecular compounds
- Ionic bonding ionic compounds
- Metallic bonding
- Molecular geometry
3Section 1
- Introduction to chemical bonding
4Introduction to chemical bonding
- What is a chemical bond???
- A mutual electrical attraction between the
nuclei and valence electrons of different atoms
that binds the atoms together
5Introduction to chemical bonding
- Why do atoms bond?
-
- They are working to achieve more stable
arrangements where the bonded atoms will have
lower potential energy than they do when existing
as individual atoms.
6Introduction to chemical bonding
- Types of Chemical Bonding
- 1. Ionic an electrical attraction that forms
between cations () and anions (-) - 2. Covalent are formed when electrons are
shared between atoms - 3. Metallic formed by many atoms sharing many
electrons
7Introduction to chemical bonding
- However.
- Bonds are never purely covalent or purely ionic.
- The degree of ionic-ness or covalent-ness depends
on property of electronegativity.
8Degree of Ionic/Covalent Character in Chemical
Bonds
100 50 5 0
Ionic Polar-Covalent Nonpolar-Covalent
9Introduction to chemical bonding
- Recall what electronegativity is
-
- The degree of attraction that an atom has to
electrons that are within a bonded compound. - (see page 161)
10Introduction to chemical bonding
- To determine the degree of ionic-ness or
covalent-ness you must take each of the
electronegativities for the elements in the
compound and subtract them.
11Introduction to chemical bonding
- If difference is 0-0.3 nonpolar covalent
- If difference is 0.3 1.7 polar covalent
- 1.7 and above Ionic
12Ionic/Covalent Character Due to Electronegativity
Differences
100 50 5 0
3.3 1.7 0.3 0
Ionic Polar-Covalent Nonpolar-Covalent
13Introduction to chemical bonding
2.5 - 2.1 0.4 Polar Covalent 2.5 - 0.7
1.8 Ionic 2.5 3.0 0.5 Polar Covalent
- Sulfur Hydrogen
- Sulfur Cesium
- Sulfur Chlorine
14Introduction to chemical bonding
- In general however
- If bonding elements are on opposite sides of the
periodic table then they tend to be ionic. - If elements are close together, then they tend
to be covalent.
15Assignment
- Page 177 3, 4, 5
- Page 209 6
16Section 2
- Covalent Bonding Molecular Compounds
17Covalent Bonding
- What is a molecule?
- A neutral group of atoms that are held together
by covalent bonds. - May be different atoms such as H2O or C6H12O6
- May be the same atoms such as O2
18Covalent Bonding
- Molecular compounds are made of molecules .. Not
ions! - We represent molecular compounds by chemical
formulas that show numbers of atoms of each kind
of element in the compound. CH4 - methane
19Covalent Bonding
- Diatomic molecules are those elements that exist
in pairs of like atoms that are bonded together. - There are 7 diatomic molecules
- H2 N2 O2 F2 Cl2 I2 Br2
20Covalent Bonding
- Formation of a covalent bond
- When atoms are far apart they do not attract
potential energy is zero. - As they come closer the electrons are attracted
to protons but electrons and electrons repel
but e- to p attraction is stronger!
21Covalent Bonding
- The electron clouds of the bonded atoms are
overlapped and form a bond length.
22Covalent Bonding
- Energy is released when these atoms join together
with a bond. - Energy must be added to separate these atoms
called bond energies. - Bond energy is expressed in kilojoules per mole.
23Covalent Bonding
- Octet Rule Atoms will either gain, lose, or
share electrons so that their outer energy levels
will contain eight electrons (H is an exception
since it can only have 2 in the outer level). - These electrons that are being gained, lost, or
shared are represented by using the electron dot
diagrams.
24Examples of electron dot notations
- 1 valence electron
- 3 valence electrons
- 5 valence electrons
- 7 valance electrons
X
X
X
X
25Covalent Bonding
- Shared electron pairs and unshared pairs
- ClCl Shared pair
- Unshared pairs
26Covalent Bonding
- These electron dot representations are called
Lewis structures. - Dots represent the valence electrons
27 28Covalent Bonding
- Lewis structures can also be represented using
structural formulas. - Dashes indicate bonds of shared electrons
(unshared e- are not shown - Cl - Cl
- One pair (2 e-) is shared here.
29Covalent Bonding
- Lewis structure for ammonia (NH3)
-
30Covalent Bonding
- Practice
- Draw Lewis structure for methane CH4
- Ammonia NH3
- Hydrogen Sulfide H2S
- Phosphorus trifluoride PF3
31Covalent Bonding
- Some atoms can form multiple bonds especially
C, O, N. - Double bonds are bonds that share 2 pair of
electrons - CC means CC
- Triple bonds share 3 pair
- CC means CC
32Covalent Bonding
- Resonance
- Some substances cannot be drawn correctly with
Lewis structure diagrams - Some electrons share time with other atoms ex.
Ozone O3
33Covalent Bonding
- Electrons in ozone may be represented as O
OO - Other times it may be represented as OOO
- Actually these structures are shared electrons
resonate (go back forth) between them
34Covalent Bonding
35Section 3
- Ionic Bonding and Ionic Compounds
36Section 3 Ionic Bonding Compounds
- Ionic compounds are formed of positive and
negative ions - When combined these charges equal zero
- Ex Na 1
- Cl 1-
0 charge
37Section 3 Ionic Bonding Compounds
- Ionic substances are usually solids
- Ionic solids are generally crystalline in shape
- An ionic compound is a 3-D network of and
ions that are attracted to each other
38Section 3 Ionic Bonding Compounds
- Crystals in ionic compounds exist in orderly
arrangements known as a crystal lattice.
39Section 3 Ionic Bonding Compounds
- Ionic substances are not referred to as
molecules - Ionic substances are referred to as formula
units - A formula unit is the simplest ratio of the ions
that are bonded together.
40Section 3 Ionic Bonding Compounds
- The ratio of ions depends on the charges.
- What would result when F- combines with Ca2?
- CaF2
41Section 3 Ionic Bonding Compounds
- When ions are written using electron dot
structures the dots are written and symbols for
their charges. - Na. ? Na
- Cl ? -
42Compared to molecular compounds, ionic compounds
- Have very strong attractions
- Are hard, but brittle
- Have higher melting points and boiling points
- When dissolved or in the molten state they will
conduct electricity
43Polyatomic Ions
- A group of atoms covalently bonded together but
with a charge. - Sulfate SO42-
- Carbonate CO32-
- Nitrate NO3-
- Ammonium NH4
44Section 4
45Metallic Bonding
- Metals are excellent electrical conductors in the
solid state. - This is due to highly mobile valence electrons
that travel from atom to atom.
e-
46Metallic Bonding
- Generally metals have either 1 or 2 s electrons
- p orbitals are vacant
- Many are filling in the d level
- Electrons become delocalized and move between
atoms
47Metallic Bonding
- A metallic bond is the mutual sharing of many
electrons among many atoms. - Electrons travel in what is known as the zone of
conduction.
48Metallic Properties
- High electrical conductivity
- High thermal conductivity
- High luster
- Malleable (can be hammered or pressed into shape)
- Ductile (capable of being drawn or extruded
through small openings to produce a wire)
49Metallic Bond Strength
- Varies with nuclear charge and number of
electrons shared. - High bond strengths result in high heats of
vaporization (when metals are changed into
gaseous phase)
50Section 5
51Molecular geometry
- A molecules properties depend on bonding of
atoms, but also the molecular geometry.
52Molecular geometry
- Is the three dimensional arrangement of a
molecules atoms in space.
53VSEPR Theory
- Valence Shell Electron Pair Repulsion
- Electrons around a nucleus repel each other to be
as far away from each other as possible.
54VSEPR Theory
- AB2 forms linear molecule as with beryllium
- However, water (H2O) is bent due to electrons
repulsion!
55VSEPR Theory
- AB3 forms trigonal planar molecule-ex. ammonia
- AB4 forms tetrahedral molecule ex. methane
- See page 200 for other shapes
56Hybridization
- Explains how atoms orbitals become rearranged to
form covalent bonds. - Hybridization is the mixing of 2 or more orbitals
of similar energies on the same atom to produce
new orbitals of equal energies.
57Hybridization
- Methane (CH4) is an example of hybridization
- Carbons normal configuration is 2s22p2
- In methane all the electrons in the 2nd energy
level become equal in energy and is referred to
as sp3
58Intermolecular Forces
- What happens to liquid molecules when they are
heated? - As energy is added particles overcome their
attraction to each other. - IM Forces are the forces of attraction between
molecules not within the molecule. - IM forces vary in strength but are weaker than
bonds that join atoms
59Intermolecular Forces
- Strongest IM forces exist in polar molecules.
- Polar molecules act as tiny dipoles (equal
opposite charges separated by short distances)
60Intermolecular Forces
- Dipole dipole forces attract between molecules
such as between two water molecules. - Positive H region is attracted to negative O
region of a different molecule.
61Intermolecular Forces
- Another IM force is Hydrogen bonding.
- Is a strong type of dipole-dipole force
- Explains high boiling points of H-containing
substances such as water and ammonia
62Intermolecular Forces
- In hydrogen bonding, a hydrogen atom is attracted
to an unshared pair of electrons of an
electronegative atom in a nearby molecule. - The double helix of DNA is held together by
hydrogen bonding.
63Intermolecular Forces
- London Dispersion forces
- Are very weak bonds
- Occur due to the fact that since electrons are in
constant motion that briefly there are moments
where electrons are unevenly distributed and thus
the molecule briefly has a charged area.