1. Bonding
Chapter 6

2. Chemical Bonds
All types of chemical bonds form due to the mutual attraction between the nucleus of one atom and valence electrons of another atom.

3. Ionic Bonds
Ionic bonds form when electrons are transferred from one atom to another.
A metal will lose electrons and form a cation
A nonmetal will gain electrons and form an anion

4. Ion charges
When an ion forms it usually has an octet of electrons in its valence shell.
An octet (8 electrons) would result from a full s and p sublevel in the outermost energy level.
What is the charge on the fluorine ion?
What is the charge on the strontium ion? F-
Sr2+

5. Ionic Compounds
Ionic compounds form when the cation and anion stick together due to the electrostatic attraction between the oppositely charged ions.
The overall compound will be neutral.
cation
anion
compound
Na+
Br-
Al3+
Cl-
Ca2+
N3-
NaBr
AlCl3
Ca3N2

6. Ionic Crystals
Ionic solids contain cations and anions that are arranged in an organized three dimensional crystal lattice.
Within the lattice, each cation is surrounded by anions and vice-versa.

7. Ionic Bond Strength
Lattice energy is the energy released when one mole of an ionic compound is formed from gaseous ions.
The greater the lattice energy the stronger the ionic bond.

8. (Honors)
Lattice energy is affected by two variables; the ions’ charges and the ions’ radii.
The greater the charges, the greater the LE. LENaCl LEMgCl2
The smaller the radii, the greater the LE. LENaCl LELiCl
<
<

9. Covalent Bonds
Covalent bonds result from the sharing of electrons between two atoms
Covalent bonding occurs when nonmetal atoms are bonded together
There are two types of covalent bonds; and nonpolar covalent bonds polar covalent bonds

10. Nonpolar Covalent Bonds
In a nonpolar covalent bond, electrons are shared equally between the atoms.
This results in a balanced electron distribution

11. Polar Covalent Bonds
In a polar covalent bond, electrons are shared unequally
This results in an uneven charge distribution (one end is positive and the other is negative)

12. Electronegativity
Def: the ability of an atom in a compound to attract electrons from another atom
Values given on page 161 in text

13. Electronegativity differences between atoms in a compound can help us predict the type of bond present
Nonpolar ~ Polar~ ionic ~ 1.7 +

14. Lewis Dot Diagrams Show how atoms are bonded together in a molecule.
A dot represents a valence electron
A pair of dots between two atom symbols represents a bond between the atoms

15. Drawing Lewis Diagrams
Step 1: determine the total number of valence electrons in the molecule
Step 2: Arrange element symbols to form a skeleton structure
If carbon is present, it should go in the center
Hydrogen atoms should always be on the edge of the molecule. Halogens usually are on an edge.
Otherwise, molecules tend to be somewhat symmetrical

16. Step 3: Put one pair of electrons between adjacent atoms
Step 4: Add unshared pairs (lone pairs) to the outside of each atom until each atom has a total of 8 electrons (an octet), except hydrogen will have 2 (a duet).

17. Step 5: Count all the electrons you have drawn
If that equals the amount that you had available from Step 1, you are done!
If you have drawn too many dots, you must make double and/or triple bonds until you have the proper number of electrons present. For every additional bond made, you will remove two lone pairs of electrons

18. HONORS – Exceptions to the Octet Rule
Some molecules will not obey the octet rule. (They often have a noble gas in them.)
If you complete Step 4 and in Step 5 you have fewer electrons shown in your diagram that you counted from Step 1, you will need to place all extra electrons (in pairs) around the central atom.

19. VSEPR Valence Shell Electron Pair Repulsion Theory
States that regions of electrons around the central atom of a molecule will be arranged as far apart as possible
Determines the shapes of molecules

20. Hybridization
When a molecule is formed, the atomic orbitals (s, p, etc) from the individual atoms combine with each other to make new orbitals called hybrid orbitals
Hybridization must occur because the shapes of atomic orbitals does NOT explain the molecular shapes.

21. sp3 hybridization
If there are four regions of high electron density around the central atom, then
1 s-orbital has hybridized with 3 p-orbitals resulting in 4 sp3 orbitals arranged tetrahedrally

22. sp2 hybridization
If there are three regions of high electron density around the central atom, then
1 s-orbital has hybridized with 2 p-orbitals resulting in 3 sp2 orbitals arranged in a trigonal pyramid

23. sp hybridization
If there are two regions of high electron density around the central atom, then
1 s-orbital has hybridized with 1 p-orbital resulting in 2 sp orbitals arranged linearly

24. Sigma vs Pi Bonding
A sigma (s) bond occurs when electron density lies between the nuclei of the bonded atoms.
All single bonds are s-bonds.

25. Pi (p) bonding occurs when orbitals overlap above and below the line between the nuclei

26. All double bonds contain 1 s and 1 p bond
All triple bonds contain 1 s and 2 p bonds
The s-bond forms from the overlap of the hybrid orbitals while the p-bonds form from the overlap of the unhybridized p-orbitals

27. Molecular Polarity Dipoles are polar molecules.Cl H
Dipoles are polar molecules.
A dipole exists when there is an asymmetrical charge distribution in the molecule
For a molecule to be a dipole it must contain at least one polar bond and must have an asymmetrical shape O H H

28. Examples of Dipoles

29. Examples of Nonpolar Molecules

30. Intermolecular Forces
def: The forces of attraction that exist between molecules.
Intermolecular forces are weaker than a covalent or ionic bond
There are three types of intermolecular forces:
Dipole-Dipole Forces
Hydrogen Bonding
London Dispersion Forces

31. Dipole-Dipole Forces Exists between two dipoles (polar molecules)
The positive end of one molecule is attracted to the negative end of another molecule
The greater the polarity of the molecule, the stronger the dipole-dipole force

32. Hydrogen Bonding
Def: a hydrogen atom bonded to a highly electronegative atom (F, O, or N) is attracted to an unshared pair of electrons on the F, O, or N of a nearby molecule.
H-bonding is a very
strong type of
dipole-dipole force.

33. London Dispersion Forces
A weak intermolecular force that results from a momentary (temporary) dipole.
All molecules both polar and nonpolar experience LDFs, but this is the ONLY type of intermolecular force that nonpolar molecules experience.
The larger the molecular mass, the stronger the LDFs.
A temporary uneven electron distribution results in a momentary dipole
one momentary dipole can cause another nearby molecule to also have a momentary dipole

34. Metallic Bonding
The bonding in metals consists of a crystal lattice of metal ions immersed in a sea of delocalized, mobile valence electrons
These mobile electrons
are able to flow
throughout the sample
allowing metals to
conduct electricity