1. Chemical Bonds

2. Forming Chemical Bonds
The force that holds two atoms together is called a chemical bond .
The valence electrons are the electrons involved in forming chemical bonds.
Elements tend to react to acquire eight electrons. This is called a stable octet.
Noble gases (group VIIIA/18) have this structure (octet) and are inert (does not form bonds).
Atoms can gain, lose, or share electrons to reach an octet.

3. Forming Ions
Positive ions (cations) are formed when atoms lose one or more valence electrons. [Usually these atoms are metals.]
Reactivity of metals (cations) are based on how easily they lose electrons. [ Ionization energy]
Negative ions (anions) are formed when atoms gain one or more valence electrons. [These atoms tend to be non-metals.]
Electronegativity is the ability to attract or gain electrons.

5. Forming Ionic Bonds
Ionic bonds: Complete transfer of valence electrons between atoms (difference of electronegativity of 1.6 or greater).
Bond between a metal and a nonmetal
Two neutral atoms will form ions. The resulting compound is called an ionic compound.

6. Properties of Ionic Compounds
Form crystal lattice bonding structure.
Example: NaCl (sodium chloride)
High melting point and boiling point due to strong electrostatic charge between the atoms (cation is attracted to anion after the transfer of electrons).
Hard, rigid and brittle solids.
Conducts electricity in liquid state or dissolved in water only.

7. Forming Covalent Bonds
Covalent bonds: Sharing of valence electrons between atoms (difference of electronegativity of less than 1.6).
Occurs usually between elements close to each other on the periodic table (mostly 2 non-metals).
The resulting compound is called a molecule.

8. Properties of Covalent Compounds
Have definite and predictable shapes.
Low melting and boiling points.
Relatively soft solids.
Can exist as solids, liquids or gases.

9. Metallic bonds
This weak bond is formed due to the attraction between kernels and the mobile electrons in a metal lattice.
Properties:
High thermal and electrical conductivity.
Luster and high reflectivity. (shiny)
Malleability and ductility. (They can be beaten or shaped without fracture.)
Variability of mechanical strengths

10. Complete your Venn diagrams
Requirements:
5 for ionic only
5 for covalent only
1 similarity

11. Comparison of Bonding Types
Ionic
Covalent
Metal and nonmetal
ions
2 nonmetals
molten salts
conductive
molecules
non-conductive
valence e-
transfer of electrons
sharing of electrons
high mp & bp
low bp & mp
DEN < 1.6
DEN > 1.6

12. Ionic and Covalent Song

13. Diatomic Elements
These atoms are never found alone in nature. They are always bonded to something. (that can mean another atom of their same element)
Memorize: H2, N2, O2, F2, Cl2, Br2, I2
To help memorize remember at 7 make a 7 + H.

14. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together. Subtract to find the difference.

15. Examples in calculating the difference of electronegativity
HF, we would subtract the electronegativity of hydrogen (2.1) from fluorine (4.0).
= 1.9
O2 is a diatomic element. Since the two oxygen's have the same electronegativity, the difference between them is 0.
MgCl2 is an ionic compound. We would subtract the electronegativity of magnesium (1.2) from chlorine (3.0).
= 1.8

16. Electronegativity and Bond Character
The character and type of a chemical bond can be predicted using the electronegativity difference of the elements that are bonded.
Electronegativity is a measurement of how strong the atom is pulling on electrons that it is sharing in a bond with another atom.
For identical atoms, the electronegativity difference is zero, the electrons are shared equally and the bond is considered nonpolar covalent, which is a pure covalent bond.

17. Lewis Dot Structures Ionic
A system of “drawing” bonds
Shows how valence electrons are arranged
Dots represent valence electrons
Pairs of dots represent bonding pairs of electrons.
Show the reaction between sodium and chlorine:
Show the reaction between calcium and bromine:

18. Determining Number of Covalent Bonds
To determine how many bonds exist in a molecule, use the following formula
N - A = # bonds 2
-- Where N is the # of needed electrons, which is 8 for all elements but H, which is 2.
-- Where A is the # of available electrons, which is the number of valence electrons.

19. Drawing Lewis Structures Covalent
1.Determine the # of bonds needed
2.Draw a skeleton structure.
3.Connect the atoms with the # of bonds. (determined in step 1)
4.Finish by making sure all atoms in the structure have an octet of electrons.
5.Remember, if the structure has a positive charge, it has fewer available electrons, and if it has a negative charge it has more available electrons.

20. Lewis Structure Practice
Draw Lewis Dot Structures for the following molecules and polyatomics:
CO N O Cl H2O SO4 2-

21. Exceptions to the Octet Rule
There are two elements which don’t want an octet of electrons and those are Be (wants 4) and B (wants 6).
Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule using their empty valence d orbitals, especially when surrounded by highly electronegative atoms such as chlorine, fluorine, bromine, and oxygen.

22. Exceptions to the Octet Rule
The exceptions are easy to recognize-- when you calculate the # of bonds, the formula will give you an answer that makes no sense. If this happens, do the following:
Connect all of the atoms to the central element with one bond.
Give all of the surrounding elements an octet of electrons.
Count the number of electrons you have put in the structure so far.
- If it equals the available # of electrons you calculated, you’re done.
- If it is less than the available # of electrons you calculates, place
the needed number of electrons around the central element.

23. Exceptions to the Octet Rule
Draw Lewis Dot Structures for the following molecules:
BF PCl SF XeF2

24. Carbon Bonding
Carbon follows the rules previously discussed, but it has a few unique properties.
Carbon is always the central atom in its compounds and it tends to bond to itself.
Carbon always has four bonds.
Anytime you have a choice for a skeleton structure, always go for the most symmetrical option.

25. Carbon Bonding Draw Lewis Dot Structures for the following molecules:
C2H C3H C2H2

26. Bond Strength and Length
C C C C C C
Bond Type single double triple
Bond Length
Bond Strength
Bond lengths are in picometers and bond energies are in kJ
Stronger bonds = higher boiling & melting points
Weaker bonds = lower boiling & melting points

27. Resonance
Resonance occurs when you have a double bond which can be placed in more than one location.
More than one valid Lewis structure can be written for a molecule. Examples: O3, NO3-, NO2-, SO22-, CO32-

28. Let’s Practice! H2
SO2
CCl4
SiH4
PCl5
XeCl6
H2O
SO4-2
NO3-
C2H3O2-

29. Molecular Shape
Valence Share Electron-Pair Repulsion (VSEPR) model allows us to predict the molecular shape by assuming that the repulsive forces of electron pairs cause them to be as far apart as possible from each other.
Only the valence electron pairs are considered in determining the geometry.

30. Effect of the number of electron pairs around the central atom
4 charge clouds, tetrahedral
2 charge clouds, linear
3 charge clouds, trigonal planar

31. PREDICTING EXPECTED GEOMETRY ACCORDING TO VSEPR THEORY
Lewis dot structure determines the total # of electrons around the central atom.
Multiple bonds (double and triple) count as one.
The number of bonding and nonbonding electron pairs around the central atom determines the geometry of electron pairs and the molecular geometry.
Lone e pairs affect geometry more than bonding pairs.
The shape is referred to as bent if there are lone pairs on the central atom.

32. Molecular Shapes 2,3,4 Electron Pairs

33. Molecular Shape Practice
Example:
H2CO
HCN

34. Molecular Shape Practice
Example:
CO2
PF3

35. Molecular Shape Practice
Example:
SO42-