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Chapter 6
Chemical Bonding
Section 1:
Introduction to chemical bonding
Introduction 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
Introduction 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- increase stability.
Introduction 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
Introduction 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.
Degree of Ionic/Covalent Character in Chemical Bonds
Ionic
Polar-Covalent
Nonpolar-Covalent
100%
50%
5%
0%
Introduction to chemical bonding
Recall what electronegativity is:
The ability or degree of attraction that an atom has to electrons that are within a bonded compound.
(see page 161)
Introduction 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.
Introduction to chemical bonding
If difference is 0-0.3 = nonpolar covalent
If difference is 0.4 – 1.7 = polar covalent
Above 1.8 = Ionic
Ionic/Covalent Character Due to Electronegativity Differences
Ionic
Polar-Covalent
Nonpolar-Covalent
100%
50%
5%
0%
3.3
1.7
0.3
0
Introduction to chemical bonding
Sulfur + Hydrogen
Sulfur + Cesium
Sulfur + Chlorine
2.5 - 2.1 = 0.4 Polar Covalent
2.5 - 0.7 = 1.8 Ionic
2.5 – 3.0 = 0.5
Polar Covalent
Introduction to chemical bondingIn general however…
If bonding elements are on opposite sides of the periodic table (metal with a nonmetal) then they tend to be ionic.
If elements are close together (nonmetal to nonmetal), then they tend to be covalent.
Section 2:
Covalent Bonding & Molecular Compounds
Covalent 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
Covalent Bonding
Molecular compounds are made of molecules ….. Not ions!
We represent covalent or molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH4 - methane
Covalent 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
Big 7
Covalent 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!
Covalent Bonding
The electron clouds of the bonded atoms are overlapped and form a “bond length.”
Covalent Bonding
Energy is released when these atoms join together with a bond.
Energy must be added to separate these atoms into neutral isolated atoms – called bond energies.
Bond energy is expressed in kilojoules per mole.
Covalent 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.
Examples of electron dot notations
1 valence electron
3 valence electrons
5 valence electrons
7 valance electrons
X
X
X
X
Covalent Bonding
Shared electron pairs and unshared pairs:
Cl:ClShared pair
Unshared pairs
Covalent Bonding
These electron dot representations are called Lewis structures.
Dots represent the valence electrons
Covalent 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.
Steps To Drawing Lewis Structures
Calculate the number of valance electrons.
Arrange atoms. Compare number of electrons used with
number of electrons available. Check octet rule. Change dots to dashes where
appropriate.
Covalent Bonding
Lewis structure for ammonia (NH3)
Covalent Bonding
Practice: Draw Lewis structure for
methane CH4
Ammonia NH3
Hydrogen Sulfide H2S
Phosphorus trifluoride PF3
More Guidelines
H and halogen atoms usually bond to only one other atom in a molecule and are usually on the outside or end of a molecule (each only need 1 electron to form stable octet and electronegativity)
More Guidelines
The atom with the smallest electro-negativity is often the central atom
When a molecule contains more atoms of 1 element than the other, these atoms often surround the central atom
Covalent Bonding
Some atoms can form multiple bonds – especially C, O, & N.
Double bonds are bonds that share 2 pair of electrons
C=C means C::CTriple bonds share 3 pair
C≡C means C:::C
Covalent Bonding
Resonance:Some substances cannot be
drawn correctly with Lewis structure diagrams
Some electrons share time with other atoms – ex. Ozone – O3
Covalent Bonding
Electrons in ozone may be represented as: O = O–O
Other times it may be represented as O–O=O
Actually these structures are shared – electrons “resonate” (go back & forth) between them
Section 3:
Ionic Bonding and Ionic Compounds
Section 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
Section 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
Section 3: Ionic Bonding & Compounds
Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.
Section 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.
Section 3: Ionic Bonding & Compounds
The ratio of ions depends on the charges.
What would result when F-
combines with Ca2+?
CaF2
Section 3: Ionic Bonding & Compounds
When ions are written using electron dot structures the dots are written and symbols for their charges.
Na. Na+
Cl -
Compared 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
Polyatomic Ions:
A group of atoms covalently bonded together but with a charge.
Sulfate SO42-
Carbonate CO32-
Nitrate NO3-
Ammonium NH4+
Section 4:
Metallic Bonding
Metallic 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-
Metallic Bonding
Generally metals have either 1 or 2 s electrons
p orbitals are vacantMany are filling in the d levelElectrons become delocalized
and move between atoms (sea of electrons)
Metallic Bonding
A metallic bond is the mutual sharing of many electrons among many atoms.
Metallic Properties
High electrical conductivityHigh thermal conductivityHigh lusterMalleable (can be hammered or
pressed into shape)Ductile (capable of being drawn or
extruded through small openings to produce a wire)
Metallic 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)
Section 5:
Molecular Geometry
Molecular geometry…
A molecule’s properties depend on bonding of atoms, but also the molecular geometry.
Molecular geometry…
Is the three dimensional arrangement of a molecule’s atoms in space.
VSEPR Theory
Valence Shell Electron Pair Repulsion
Electrons around a nucleus repel each other to be as far away from each other as possible.
VSEPR Theory
Types of e- Pairs Bonding pairs - form bonds Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
Draw the Lewis Diagram. Tally up e- pairs on central atom.
double/triple bonds = ONE pair
Shape is determined by the # of bonding pairs and lone pairs.
Know the common shapes & their bond angles!
Common Molecular Shapes
2 total
2 bond
0 lone
LINEAR180°BeH2
3 total
3 bond
0 lone
TRIGONAL PLANAR
120°
BF3
Common Molecular Shapes
4 total
4 bond
0 lone
TETRAHEDRAL
109.5°
CH4
Common Molecular Shapes
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
NH3
Common Molecular Shapes
4 total
2 bond
2 lone
BENT
104.5°
H2O
Common Molecular Shapes
PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°
F P FF
Examples
CO2
O C O2 total
2 bond
0 lone LINEAR
180°
Examples
Hybridization
Explains how atom’s 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.
Hybridization
Methane (CH4) is an example of hybridization: Carbon’s normal configuration is
2s22p2
In methane all the electrons in the 2nd energy level become equal in energy and is referred to as sp3
Intermolecular Forces
11.2
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
• 41 kJ to vaporize 1 mole of water (inter)
• 930 kJ to break all O-H bonds in 1 mole of water (intra)
Generally, intermolecular forces are much weaker than intramolecular forces.
“Measure” of intermolecular force
boiling point
melting point
Intermolecular Forces:
Strong IM forces exist in polar molecules.
Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)
Types of Intermolecular Forces
Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2
Types of IMF
Dipole-Dipole Forces
+ -
View animation online.
Types of Intermolecular Forces
Dipole ForcesAttractive forces between an ion and a polar molecule
11.2
Ion-Dipole Interaction
Intermolecular Forces:
Another IM force is Hydrogen bonding.
Is the strongest type of dipole-dipole force
Explains high boiling points of H-containing substances such as water and ammonia
Intermolecular Forces:
In hydrogen bonding, a hydrogen atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.
Types of Intermolecular ForcesHydrogen Bond (strongest)
11.2
The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. IT IS NOT A BOND.
A H…B A H…Aor
A & B are N, O, or F
Types of IMF
Hydrogen Bonding
Intermolecular Forces:London/ Dispersion/VanDerWaals
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.
Types of IMF
London Dispersion Forces
View animation online.