1. CHEMICAL BONDS
SC2. Obtain, evaluate, and communicate information about the chemical and physical properties of matter resulting from the ability of atoms to form bonds. 
a. Plan and carry out an investigation to gather evidence to compare the physical and chemical properties at the macroscopic scale to infer the strength of intermolecular and intramolecular forces. 
b. Construct an argument by applying principles of inter- and intramolecular forces to identify substances based on chemical and physical properties.
c. Construct an explanation about the importance of molecular level structure in the functioning of designed materials.
d. Develop and use models to evaluate bonding configurations from nonpolar covalent to ionic bonding. 

2. Intermolecular vs. Intramolecular forces
Intramolecular forces are the forces that hold atoms together within a molecule.
These are the strong forces of chemical bonding (ionic/covalent/metallic)
Intermolecular forces are forces that exist between molecules.
These are much weaker interactions (dipole-dipole/hydrogen bonding/van der Waals), but they influence physical properties

3. What is a chemical bond?
Remember that although atoms are neutral (# of protons = # of electrons), they are not all stable (unreactive).
The most stable atoms are those that have a full outer energy level of valence electrons.
Which elements have the most stable atoms?

4. What is a chemical bond?
So, for an atom to achieve that “noble gas-like” stability, it must gain or lose valence electrons. Atoms must, therefore, interact with other atoms to exchange or share electrons. This interaction of atoms is called bonding. Lewis structures are commonly used to “track” electron exchange/sharing.

5. What is a chemical bond?
Bond formation releases energy, while bond cleavage (breaking) absorbs energy.
The amount of energy required to break a bond is called bond energy, and it is a measure of the strength of a bond.
Higher bond energy = stronger bond
Bond length is inversely related to bond energy
Higher bond energy = shorter bond length

6. Types of Bonding
If atoms exchange electrons, it results in the formation of ions – either positive (cations) or negative (anions). This interaction is called ionic bonding. If atoms share electrons, it is called covalent bonding. If metal atoms share their delocalized "sea" of electrons, it is called metallic bonding.

7. Ionic bonding
The ion formed by an atom depends on its current number of valence electrons and how close the atom is to noble gas configuration. Remember that valence electron count is predicted by the group number. Generally, atoms will only exchange up to 3 (rarely 4) electrons to gain the “octet.”

8. IONIC BONDING
Ionic bonds are essentially electrostatic attractions between atoms.
Which types of elements will tend to form ionic bonds between their atoms (ions)?
Why?
Ions will bond so that the resulting compound is neutral.

9. bond

10. IONIC BONDING Properties resulting from ionic bonding include:
Crystalline solid
Very high melting point
Soluble in H2O
Nonconductor of heat and electricity
Conducts electricity in aqueous solutions
Examples: NaCl, CaCO3

11. COVALENT bonding
Atoms that participate in covalent bonding exchange two electrons per bond.
The electrons come from overlapping orbitals – typically one from each atom
Which kind of elements tend to form covalent bonds?
Why?
A covalently bonded compound is called a molecule.

13. COVALENT bonding
Multiple covalent bonds may form if more than one pair of electrons is shared.
This is common with C, N, O, P, S

14. COVALENT bonding
An alternate form (coordinate covalent bond) occurs when one atom contributes both electrons to the bond

15. COVALENT bonding
Atoms that covalently bond can either result in a charged ion or a neutral compound.

16. COVALENT bonding
For several compounds, multiple valid Lewis structures may be drawn. These are called resonance structures.

17. COVALENT bonding
Some elements that participate in covalent bonding also undergo a process called hybridization
Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties.
Hybridized orbitals are very useful in the explanation of the shape of molecules.

18. COVALENT bonding There are 3 types of hybrid orbitals:
sp            s orbital + 1 p orbital
Ex: C2H2

19. COVALENT bonding There are 3 types of hybrid orbitals:
sp2           s orbital + 2 p orbitals
Ex. C2H4

20. COVALENT bonding
sp3           s orbital + 3 p orbitals
Ex. CH4

21. COVALENT bonding
When orbitals hybridize, additional orbital bonds form:
Sigma bonds (σ) form when orbitals overlap end-to-end
Pi bonds (π) form when orbitals overlap side-by-side
Sigma bonds
are stronger than
pi bonds

22. COVALENT bonding

23. COVALENT BONDING How to describe the shape of a molecule:
VSEPR (valence shell electron pair repulsion) theory states that the valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion, thus determining the molecule's geometry.

24. COVALENT BONDING Three basic electron geometries: Linear
Trigonal planar
Tetrahedral
There are others, but they involve greater degrees of hybridization

25. COVALENT BONDING Not all atoms share equally:
Polar covalent bonds occur among atoms that share unequally, producing a dipole
Often occurs in compounds containing halogens, N and O

27. COVALENT BONDING Not all atoms share equally:
Nonpolar covalent bonds occur among atoms that share nearly equally
Generally among same kinds of atoms
Polar covalent bond strength > nonpolar c.b.s.

28. COVALENT BONDING Properties resulting from covalent bonding include:
Gas, liquid, or a soft solid
Low melting point and low boiling point
Insoluble in H2O
Soluble in nonpolar solvents
Nonconductor of heat and electricity
Nonlustrous
Examples: CO2, CH3COOH

29. Predicting Bonds
Calculating the difference in electronegativity can predict ionic vs. polar covalent vs. nonpolar covalent bonds
Difference  1.9  ionic
Difference 0.5 – 1.9  polar covalent
Difference < 0.5 nonpolar covalent

30. metallic BONDING
Metallic bonding occurs when transition metals bond to either themselves or mixed with other metals in alloys
Having generally low electronegativity they tend to lose their valence electrons easily 
The valence electrons detach from the atoms but are not held by any one of the other atoms – they are only held collectively by the entire assemblage of atoms

31. metallic BONDING
In a metal the ion cores are held in an ordered, or crystal, lattice. The valence electrons are free to move about under applied stimulation, such as electric fields or heat.

32. metallic BONDING Properties resulting from metallic bonding include:
Malleable solid
High melting point and boiling point
Insoluble in H2O
Insoluble in nonpolar solvents
Conducts heat and electricity
Lustrous
Examples: gold, copper

33. Intermolecular forces
Intermolecular forces are generally referred to as van der Waals forces. These are distance-dependent interactions between atoms or molecules. These forces play a fundamental role in polymer science, nanotechnology, and surface science, and define many properties of compounds, including solubility in polar and nonpolar solvents.

34. Intermolecular forces
There are 2 types:
Dipole-dipole
Electrostatic interactions between molecules which have permanent dipole(s)

35. Intermolecular forces
There are 2 main types of intermolecular forces:
Dipole-dipole
Hydrogen bonding is a form of dipole-dipole interaction

36. Intermolecular forces
London Dispersion Forces
Arise from interaction between uncharged atoms or molecules