1. I. Electrons and Bonding
Valence electrons: the electrons in an atom’s outermost s and p orbitals; determine the chemical properties of an atom.
Octet rule: states that atoms lose, gain, or share electrons in order to acquire a full set of eight valence electrons

2. I. Electrons and Bonding
Electron dot diagram: uses the symbol of the element and dots to represent the electrons in the outer energy level
Chemical bond: the force that holds atoms together in a compound
Ion: a charged atom that has either more or fewer electrons than protons
Ionic bond: attraction formed between oppositely charged ions in an ionic compound

3. I. Electrons and Bonding
Ionic bonds form between a metal and a nonmetal.
Covalent bond: attraction formed between atoms when they share electrons
Covalent bonds form between nonmetals.
Molecule: a neutral particle that forms as a result of electron sharing
Formula unit: the simplest ratio of ions represented in an ionic compound

4. II. Ionic Compounds Cation: a positively charged ion
Anion: a negatively charged ion
The formation of ionic compounds from positive and negative ions is always exothermic (energy is released). The attraction of the positive ion for the negative ion close to it forms a more stable system that is lower in energy than the individual ions.

5. II. Ionic Compounds
Monatomic ion: a one-atom ion (such as Mg2+ or Br -)
Binary ionic compounds are composed of positively charged monatomic ions of a metal and negatively charged monatomic ions of a nonmetal.
Oxidation number: the charge of a monatomic ion
Most transition metals and group 3A and 4A metals have more than one oxidation number.

6. II. Ionic Compounds Naming ionic compounds:
Name the cation first and the anion second.
Monatomic cations use the element name.
Monatomic anions take their name from the root of the element name plus the suffix –ide.
For elements that have more than one oxidation number, the oxidation number is written as a Roman numeral in parentheses after the name of the cation.
If the compound contains a polyatomic ion, simply name the ion.

7. III. Covalent Compounds
For an atom to be stable – whether in an ionic or covalent compound – it must have a full outer energy level.
Sometimes two atoms that both need to gain valence electrons to become stable have a similar attraction (electronegativity) for electrons.
Sharing of electrons is another way that these atoms can acquire the electron configuration of eight valence electrons, or an octet.

8. III. Covalent Compounds
Group 7A elements form a single covalent bond because they need one more electron to attain an octet.
Group 6A elements share two electrons to form two covalent bonds.
Group 5A elements form three covalent bonds with atoms of nonmetals.
Group 4A elements will form 4 covalent bonds.

9. III. Covalent Compounds
Naming Binary Covalent Compounds:
The first element in the formula is always named first, using the entire element name.
The second element in the formula is named using the root of the element and adding the suffix –ide.
Prefixes are used to indicate the number of atoms of each type that are present in the compound.
Example: N2O – dinitrogen monoxide
The first element in the compound never uses the prefix mono-.
Drop the final letter in the prefix when the name begins with a vowel.

10. III. Covalent Compounds
Prefixes for Naming Covalent Compounds:
mono hexa-
di hepta-
tri octa-
tetra nona-
penta deca-

11. IV. Molecular Structures
Structural formula – a molecular model that uses element symbols and bonds to show relative positions of atoms
One type of structural formula is a Lewis diagram.

12. IV. Molecular Structures
How to draw Lewis diagrams:
Determine the total number of valence electrons the atoms in the molecule need to be “happy.”
H, He, B, and Be are satisfied with 2.
All other elements are satisfied with 8.
Determine the total number of valence electrons all of the atoms in the molecule actually “own.”

13. IV. Molecular Structures
Subtract the number of “owned” electrons from the number of “happy” electrons, and divide the difference by two. This is the number of bonds that should connect the terminal atoms to the central atom. (If this number is less than the number of bonds needed to attach all of the terminal atoms, then use one bond to attach each of the atoms.)
The central atom is the one listed first in the chemical formula (except hydrogen).
Attach each of the terminal atoms to the central atom using the total number of bonds specified in part C.

14. IV. Molecular Structures
Add dots around each of the terminal atoms to represent the number of valence electrons they need to be satisfied in addition to the ones in the bonds. Each bond represents two valence electrons.
Remember that, in general, carbon, nitrogen, oxygen, and sulfur can form double or triple bonds with the same element or with another element.

15. IV. Molecular Structures
Exceptions to the Octet Rule:
Some molecules have an odd number of valence electrons and cannot form an octet around each atom. For example, NO2 has five valence electrons from nitrogen and 12 from oxygen, totaling 17, which cannot form an exact number of electron pairs.

16. IV. Molecular Structures
Some compounds form with fewer than eight electrons present around an atom. BH3 is an example. A total of six electrons is shared by the boron atom, which is two less than an octet. BeCl2 is another example.

17. IV. Molecular Structures
Some compounds have central atoms that contain more than eight valence electrons, which is called an expanded octet. PCl5 and SF6 are examples. When you draw the Lewis structures for these compounds, extra lone pairs are added to the central atom or more than four bonding atoms are present in the molecule.

18. IV. Molecular Structures
Resonance: condition that occurs when more than one valid Lewis structure exists for the same molecule

19. V. Determining Bond Type
Two ways electrons can be shared:
Sigma (σ) bond – occurs when the valence atomic orbital of one atom overlaps end-to-end with the valence atomic orbital of another atom, concentrating the shared electrons in an area centered between the two atoms
Pi (π) bond – occurs when two lobes of one orbital overlap two lobes of a parallel orbital to share electrons

20. V. Determining Bond Type
Single covalent bonds are sigma bonds.
In addition to one sigma bond, double and triple bonds have one or two pi bonds, respectively.
Pi bonds are usually weaker than sigma bonds.
As the number of bonds increases, they become shorter and stronger.

21. V. Determining Bond Type
Bonds are also classified by the degree to which electrons are shared.
Depends on the electronegativities of the atoms involved (ability to attract electrons)
Ionic Bond – electrons are removed from one atom and received by another atom
Difference in electronegativities is greater than 1.7

22. V. Determining Bond Type
Polar Covalent Bond – electrons are not shared equally
Difference in electronegativities of bonded atoms is between 0.4 and 1.7
Partial charges occur at the ends of a polar covalent bond:
partial positive (ς +) beside the less electronegative atom
partial negative (ς -) beside the more electronegative atom
This is called a dipole.

23. V. Determining Bond Type
Nonpolar Covalent Bond – electrons are shared equally
Difference in electronegativities is less than 0.4
Not only are individual bonds polar or nonpolar, but entire molecules are also either polar or nonpolar.
Geometry of nonpolar molecules is symmetric, while polar is asymmetric.

24. V. Determining Bond Type
Polar substances will dissolve in water, which is also polar.
Nonpolar substances will not dissolve in water, but rather only in other nonpolar liquids.
Oil is nonpolar and doesn’t mix with water.
Melting Points:
Nonpolar < Polar < Ionic

25. VI. Molecular Geometry
Linear

26. VI. Molecular Geometry
Trigonal planar

27. VI. Molecular Geometry
Tetrahedral

28. VI. Molecular Geometry
Trigonal pyramidal

29. VI. Molecular Geometry
Bent

30. VI. Molecular Geometry
Trigonal bipyramidal

31. VI. Molecular Geometry
Octahedral