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Chapter 6 . Chemical Bonding. Chapter Sections:. Introduction to chemical bonding Covalent bonding & molecular compounds Ionic bonding & ionic compounds Metallic bonding Molecular geometry. Section 1:. Introduction to chemical bonding. Introduction to chemical bonding.

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chapter 6

Chapter 6

Chemical Bonding

chapter sections
Chapter Sections:
  • Introduction to chemical bonding
  • Covalent bonding & molecular compounds
  • Ionic bonding & ionic compounds
  • Metallic bonding
  • Molecular geometry
section 1
Section 1:
  • Introduction to chemical bonding
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 bonding5
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.

introduction to chemical bonding6
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 bonding7
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
Degree of Ionic/Covalent Character in Chemical Bonds

100%

50%

5%

0%

Ionic

Polar-Covalent

Nonpolar-Covalent

introduction to chemical bonding9
Introduction 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)

introduction to chemical bonding10
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 bonding11
Introduction 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
ionic covalent character due to electronegativity differences
Ionic/Covalent Character Due to Electronegativity Differences

3.3

1.7

0.3

0

100%

50%

5%

0%

Ionic

Polar-Covalent

Nonpolar-Covalent

introduction to chemical bonding13
Introduction 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
introduction to chemical bonding14
Introduction 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.

assignment
Assignment:
  • Page 177 #3, 4, & 5
  • Page 209 #6
section 2
Section 2:
  • Covalent Bonding & Molecular Compounds
covalent bonding
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 bonding18
Covalent 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
covalent bonding19
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

covalent bonding20
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 bonding21
Covalent Bonding
  • The electron clouds of the bonded atoms are overlapped and form a “bond length.”
covalent bonding22
Covalent Bonding
  • Energy is released when these atoms join together with a bond.
  • Energy must be added to separatethese atoms – called bond energies.
  • Bond energy is expressed in kilojoules per mole.
covalent bonding23
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
Examples of electron dot notations
  • 1 valence electron
  • 3 valence electrons
  • 5 valence electrons
  • 7 valance electrons

X

X

X

X

covalent bonding25
Covalent Bonding
  • Shared electron pairs and unshared pairs:

Cl:Cl Shared pair

Unshared pairs

covalent bonding26
Covalent Bonding
  • These electron dot representations are called Lewis structures.
  • Dots represent the valence electrons
covalent bonding28
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.
covalent bonding29
Covalent Bonding
  • Lewis structure for ammonia (NH3)
covalent bonding30
Covalent Bonding
  • Practice:
    • Draw Lewis structure for methane CH4
    • Ammonia NH3
    • Hydrogen Sulfide H2S
    • Phosphorus trifluoride PF3
covalent bonding31
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::C

  • Triple bonds share 3 pair

C≡C means C:::C

covalent bonding32
Covalent Bonding
  • Resonance:
  • Some substances cannot be drawn correctly with Lewis structure diagrams
  • Some electrons share time with other atoms – ex. Ozone – O3
covalent bonding33
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
covalent bonding34
Covalent Bonding
  • Assignment:

p. 189 #4 a – e

section 3
Section 3:
  • Ionic Bonding and Ionic Compounds
section 3 ionic bonding 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 compounds37
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 compounds38
Section 3: Ionic Bonding & Compounds
  • Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.
section 3 ionic bonding compounds39
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 compounds40
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 compounds41
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
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
Polyatomic Ions:
  • A group of atoms covalently bonded together but with a charge.
  • Sulfate SO42-
  • Carbonate CO32-
  • Nitrate NO3-
  • Ammonium NH4+
section 4
Section 4:
  • Metallic Bonding
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 bonding46
Metallic 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
metallic bonding47
Metallic 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.
metallic properties
Metallic 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)
metallic bond strength
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
Section 5:
  • Molecular Geometry
molecular geometry
Molecular geometry…
  • A molecule’s properties depend on bonding of atoms, but also the molecular geometry.
molecular geometry52
Molecular geometry…
  • Is the three dimensional arrangement of a molecule’s atoms in space.
vsepr theory
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 theory54
VSEPR Theory
  • AB2 forms linear molecule as with beryllium

However, water (H2O) is bent due to electrons repulsion!

vsepr theory55
VSEPR Theory
  • AB3 forms trigonal planar molecule-ex. ammonia
  • AB4 forms tetrahedral molecule ex. methane
  • See page 200 for other shapes
hybridization
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.
hybridization57
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
Intermolecular 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
intermolecular forces59
Intermolecular Forces:
  • Strongest IM forces exist in polar molecules.
  • Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)
intermolecular forces60
Intermolecular 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.
intermolecular forces61
Intermolecular 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
intermolecular forces62
Intermolecular 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.
intermolecular forces63
Intermolecular 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.
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