1. A chemical bond is the simultaneous attraction for the same electrons by two different nuclei. Remember protons attract electrons. Opposites attract.
2. Bond formation is exothermic and stability increases. Energy is released, lower energy state is more stable.
3. Bond breaking is endothermic. It takes energy to break things apart.
4. The stronger the bond the more energy is released during formation and the more energy it takes to break it.
5. There are 3 bond types broken down into more specific categories know them all.
a. Ionic bonds are formed when a metal transfers an electron to a nonmetal. The bond is actually electrostatic attractions between a positive and negative ion.
b. Covalent bonds are formed when two non-metallic atoms share electrons.
i. Non polar covalent bonds exist between atoms of the same element. The diatomic molecules have N.P.C. bonds.
ii. Polar covalent bonds exist between two different non- metallic atoms. Ex. HCl, H2O, NH3 all have P.C. bonds.
c. Metallic bonds exist in solid metals. They consist of “fixed” nuclei and mobile electrons. The mobile electrons explain why metals conduct electricity.
6. Types of compounds:
a. Ionic compounds: Contain Ionic bonds. Metal with Non-Metal, or positive polyatomic ion with negative polyatomic ion examples: KCl, NH4NO3
i. Ionic compounds exist as solids with a crystal lattice structure
ii. Ionic compounds have high melting and boiling points
iii. Ionic compounds conduct electricity when melted or dissolved in water (aq) because the ions are free to move.
iv. Ionic compounds are commonly referred to as “Salts”
b. Covalent compounds are also called molecular compounds. They exist as molecules. Molecules only contain non-metals.
i. Molecular compounds have low melting points and boiling points. Most non-organic molecules will be gases at room temp. Why is water the exception?
ii. Molecular compounds do not conduct electricity in any state of matter.
iii. Molecular compounds contain covalent bonds.
7. Drawing Lewis Structures for Ionic compounds and Molecules: The following webpage gives an excellent overview of writing Lewis Structures: http://www.ausetute.com.au/lewisstr.html
a. Ionic compounds: draw the positive ion with no electrons, [ ] and charge. Draw the negative ion with 8 valence electrons inside [ ] with the correct charge. Subscripts in the chemical formula become coefficients outside the [ ] . The brackets show the transfer of electrons. Here is LiF
b. Covalent compouds: two electrons are shared to form 1 bond. All atoms must have a share of eight valence electrons. For help go to this webpage: http://www2.gasou.edu/chemdept/general/molecule/lewis.htm
c. You should know by heart the common molecules: HCl, H2O, NH3, CH4, N2, O2, H2, and CO2.
H:H :N:::N: (triple bond)
d. Remember Group 17 elements form 1 covalent bond, group 16 form 2 covalent bonds and Group 15 elements form 3 covalent bonds. Carbon always forms 4 covalent bonds.
8. Intermolecular Forces: These are not bonds! IMF’s hold molecules together not atoms.
a. Responsible for physical properties such as: boiling pt., melting pt., vapor pressure, and solubility.
9. Types of IMF’s: http://207.10.97.102/chemzone/lessons/03bonding/mleebonding/moleculeion_attractions.htm
a. Hydrogen Bonds: strong IMF’s explain why water has such a high boiling pt.. Hydrogen bonds only occur in molecules when there are H atoms and F, O, or N. Small atoms with high electronegativity. HF, H2O, NH3
b. Van der Waals’ Forces aka London Dispersion Forces: all molecules have them but for Non Polar molecules these are the only IMF. They are weak. The more electrons a molecule has the stronger the Van der Waals’ Forces. Big atoms = more electrons. All the diatomics, CH4, CCl4 all have Van der Waals.
c. Molecule ion attraction explains why salts dissolve in water.
10. Polarity of Molecules: determined by electronegativity differences. The more electronegative element will be the negative “pole”. If electronegativities are the same like in the diatomic molecules the molecule must be non-polar.
11. Non-polar molecules are symmetrical. Examples: diatomics, CO2, CH4, CCl4
12. Polar molecules are asymmetrical. Examples: HF, H2O, NH3
13. Like dissolves like. Polar molecules and ionic compounds dissolve in polar solvents. Salts dissolve in water. HCl dissolves in water. The positive pole or ion is always attracted to the negative oxygen pole of a water molecule. See the website in #9.
14. Non-polar molecules only dissolve in non-polar solvents. I2 dissolves better in CCl4, than H2O because I2 and CCl4 are both non polar while water is polar.
Here’s a picture of Hydrogen Bonding: dashed lines are H – bonds (imf’s), solid lines are polar covalent bonds. Why are the H’s attracted to the O of a different molecule?
15. Network solids are covalent compounds that exist as a network of covalently bonded atoms. In a large sample of a network solid every atom is covalently bonded to another atom. Diamonds, silicon dioxide and silicon carbide are examples. They are very hard with very high melting points due to the strength of the covalent bonds.