1. Chapter 6
Chemical Bonding

2. Introduction Chemical Bonding Ionic Bonding Covalent Bonding
- mutual electrical attraction between the nuclei and valence electrons of different atoms that bind atoms together.
- Why bond?
Stability! Decreased potential energy.
Two types:
Ionic Bonding
- the electrical attraction between cations and anions.
- involves the transfer of electrons.
Covalent Bonding
- the sharing of electron pairs between atoms.
- nonmetals and nonmetals or nonmetals and metalloids

3. Chemical Bond

4. Types of Bonds
- The difference in electronegativity can be used to determine the type of bond.
Ionic bonds
- difference is >1.7 or 50%
Covalent bonds
Two Types:
Nonpolar (equal sharing of electrons) – 0.3 or <5%
Polar (unequal sharing of electrons)
or 5 – 50%

5. Using Electronegativity Difference to Classify Bonding

6. Ionic Bonding

7. Covalent Bonding

8. What type of Bond? H - Cl Na – Cl Mg – O H – H C – H Covalent nonpolar
2.1
0.0
2.3
0.4
Covalent
nonpolar
Covalent
polar
Ionic
Ionic
H - Cl
This tells us the direction that
electrons are attracted.
δ+ δ-

9. The Octet Rule Electron Dot Notation
When a chemical compound is formed, each atom will have an octet (8) of electrons in its highest energy level.
Achieved by gaining, losing or sharing electrons.
Exceptions:
- Boron: 6 electrons and Hydrogen: 2 electrons
Electron Dot Notation
Shows the valence electrons of an element.
Indicated by using dots around the symbol of the element.

10. Lewis Structures
Show the arrangement of electrons among atoms in a molecule.
Rules for writing Lewis Structures:
1. Find the total number of valence electrons.
2. Arrange symbols to show how elements are bonded.
- Carbon is central when present
- Least electronegative atom is central.
- Exception is hydrogen which can never be central!
3. Arrange electrons to form stable octets.
4. Compare number of electrons used to number from step 1.
5. Change dots to dashes to indicate shared pairs of electrons.

11. Structural Formula Resonance Structures
Shows the bonding (sharing of electrons) of atoms in a molecule without illustrating unshared pairs of electrons.
Resonance Structures
Bonding in ions or molecules that cannot be correctly represented by a single Lewis structure.

12. Ionic Compounds
- (+) and (-) ions that are combined so charges are equal and opposite.
- Most exist as crystalline solids.
- Chemical formula is the simplest ratio of ions present in a compound.
- Formula unit is the simplest collection of atoms from which an ionic compound’s formula can be derived.
Examples:
Na+ and Cl-
Ca2+ and F1-

13. Characteristics of Ionic Compounds
Crystal Lattice
- orderly arrangement of ions.
- designed to minimize potential energy.

14. Characteristics of Ion Bonding in a Crystal Lattice

15. Characteristics of Ionic Compounds
Lattice Energy
- Energy released (-) when ions bond to form a lattice structure.
- The more negative the number, the higher the lattice energy.

16. Comparing Ionic and Molecular Compounds

17. Comparing Ionic and Molecular Compounds

18. Polyatomic ions
- a group of atoms that are covalently bonded that carry a charge.
- Function as a unit, usually do not break apart during chemical reactions.

19. Molecular Geometry Linear molecules VSEPR theory
- valence shell electron pair repulsion
Repulsion between sets of valence electrons surrounding an atom causes these sets to be oriented as far apart as possible.

20. Electrons want to orient as far apart as possible.
How is the shape of a molecule determined?
Draw the Lewis structure.

21. VSEPR and Lone Electron Pairs

27. Hybridization
- Mixture of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energy.
The result is a hybrid orbital.
sp hybrid orbital s and p
sp2 hybrid orbital -- s and p, p
sp3 hybrid orbital -- s and p, p, p

28. Geometry of Hybrid Orbitals

29. Hybrid Orbitals

30. Intermolecular Forces
Forces of attraction between molecules.
Molecular polarity and dipole-dipole forces -
dipole – created by opposite charges that are separated by a short distance.
H - Cl
- dipole-dipole force -
- forces of attraction between
polar molecules.
- short range forces between nearby
molecules.
δ+ δ-

31. Dipole-Dipole Forces

32. Intermolecular Forces continued
Hydrogen Bonding -
- hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.
- specific type of dipole-dipole force.
London dispersion forces -
- intermolecular attraction from the constant motion of electrons and the creation of instantaneous dipoles.
- ONLY forces acting among NOBLE GASES and nonpolar molecules.

33. London Dispersion Force