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Chemistry 1030: Unit 5 - Bonding and Nomenclature Cheat Sheet (DRAFT) by

Contains Valance Electrons, Compounds, Bonding - Why? etc.

This is a draft cheat sheet. It is a work in progress and is not finished yet.

Review

Valance Electron: Group Number.
Compound: A substance composed of two or more elements in fixed, definite propor­tions.

Forming Ions

Atoms of metals have few valance electrons (1-2) thus they tend to lose electrons to form a positive ion (cations).
Atoms of non-metals have many valance electrons (4-7) thus they tend to gain electrons to form negative ions (anions).
They do this to become stable in their outer shell.

Ionic Bonding: Type I

Format:
Name of Cation (metal)
Base Name of Anion (non-m­etal) + ide
Example:
NaCl
Sodium Chloride
 
MgBr2
Magnesium Bromide

Roman Numerals

1 = I
3 = III
5 = V
7 = VII
2 = II
4 = IV
6 = VI
8 = VIII

Ionic Bonding: Type II

Format:
Name of Cation (metal)
(Charge of cation (metal) in roman numerals)
Base name of Anion (non-m­etal) + ide
Example: CuCl
Copper
(I)
Chloride
CuCl2
Copper
(II)
Chloride

VSEPR Theory

VSEPR:
A theory based on the idea that electron groups (lone pairs, single bonds, or multiple bonds) repel each other.

VSEPR Ther

 

Drawing the Lewis Struct­ure­/Bo­nding

Step One:
Draw the lewis structure for each covalent compound.
Step Two:
Identify the bonds as single, double, or triple.
Step Three:
Label the bonding and non-bo­nding electrons.

Example

  

Bonding - Why?

Octet Rule: Atoms bond in such a way as to obtain a full outer shell (8).
Bonding involved valance electrons only.
In general, atoms either transfer or share electrons to obtain a full outer shell (8).
Valance electrons are respon­sible for the chemical properties of an atom.

Ionic Bonding: Dot and Cross

Naming Compounds: Nomenc­lature

Is it Ionic? (Metal + One or more non-me­tals)
If so go to Type I and Type II.
OR
Is it Covalent? (All non-me­tals)
If so go to Type III.

Electron Groups

To determine the shape of a molecule, count only electron groups around the central atom.
Each of the following is consider one electron group:
Non-Bo­nding Pair - (A lone pair of electrons)
Bonding Electrons - (single, double, or triple)
Example: CH4 has 4 electron groups (4 single bonds, 0 lone pairs)

Drawing Molecular Geometries

Straight Line:
Bond in plane of paper.
Hashed Line:
Bonding going into paper.
Wedge:
Bond coming out of the paper.

Terms

Single Bond:
One pair of electrons shared between two atoms (Cl2)
Double Bond:
Two pairs of electrons shared between two atoms. (O2)
Triple Bond:
Three pairs of electrons shared between two atoms. (N2)
Bonding Electrons:
Electrons shared between atoms.
Non-Bo­nding Electrons:
Electrons only found on one atom. (Lone pairs)
Overall:
Draw the lewis structure and determine how they will bond with one another to have full outer shells (8).
 
Identify the bonding and non-bo­nding electrons.
 

Summary

Ionic Bonding:
Covalent Bonding:
Metal + One or more non-me­tals.
All non-me­tals.
Electrons are transf­erred.
Electrons are shared.
Ions are formed.
Ions are not formed.
Ex. NaHCO3 or NaCl
Ex. F2 or CO2

Prefixes

1 = Mono
3 = Tri
5 = Penta
7 = Hepta
9 = Nona
2 = Di
4 = Tetra
6 = Hexa
8 = Octa
10 = Deca

Covalent Bonding: Type III

Format:
Prefix
Base name element 1
Prefix
Base name of element 2 + ide
Example: N2O
Di
nitrogen
Mono
xide
IF3
---
Iodine
Tri
Fluoride
B2H8
Di
boron
Octa
hydride
CS2
---
Carbon
Di
sulfide

Drawing Lewis Structures (2 Atoms or more)

Step One:
Draw the lewis structure for each atom separetly.
Step Two:
The atom that has the most unpaired electrons is the central atom.
Step Three:
The other atoms will share electrons with the central atom.

Example

Possible Geometries