CHEM 1405 Concept Review: Ionic and Molecular Compounds

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Octet rule: The rule that states that all atoms will attempt to gain, lose, or share electrons in order to have the same number of electrons as the closest noble gas.  (Noble gases, except He, always have 8 valence electrons, hence the name “octet”.)

Example:  Sodium (Na) has 11 electrons.  The closest noble gas is Neon (Ne), which has 10 electrons.  Sodium (Na) will lose one electron to become Na+.  Since electrons are negative, when sodium loses an electron (a negative charge), it becomes more positive.  For a summary see chart below:

1406 Ch 6 img 1.jpg 

(Elements in the shaded area are mostly transition metals and can have multiple different charges.)

 

 

Ionic Compounds: Bonding and Naming Molecular Compounds: Bonding and Naming

Ionic bonding:  This type of bonding involves the complete giving up of 1 or more electrons from one atom to another, resulting in a positively charged ion called a “cation” and a negatively charged ion called an “anion”.  This giving and receiving of electrons allows each atom to obtain a “noble gas” electron configuration, making each ion more stable than the original atoms.  These oppositely charged ions then stick together like magnets in a repeating pattern.

 

Naming Ionic Compounds:  When naming ionic compounds,  follow the following guidelines:

  1. Name the cation first. Then, name the anion.

 

  1. For a metal cation, it is named the same as the element. If the metal can have multiple charges (i.e. a transition metal), you must indicate its charge with a Roman numeral.  (See above chart.  The shaded area indicates where transition elements are.)

 Example:  Al3+ is simply called Aluminum, but Pb4+                      would be Lead (IV) and Cu+ would be Copper (I)

 

  1. For non-metal monatomic ions (one-atom ions), the name of the anion is simply the first syllable of the non-metal followed by an “-ide” ending.

      Example:  O2- is oxide, Cl- is chloride, P3- is                                          phosphide.

 

4.  For polyatomic ions, there is no change in the name of the ion.

 

 

Molecular (or Covalent) bonding:  This type of bonding involves the sharing of electrons between atoms so that each atom can have 8 electrons in its outermost (or valence) shell.  These electrons are shared in pairs, with the two nuclei of the “bonded” atoms pulling on them in a sort of tug-of-war fashion.  Atoms that are more electronegative have a stronger “pull”.  (See more about electronegativity on the next page.)

 

Covalent Compound Naming:  When naming covalent compounds, follow the following guidelines:

  1. For covalent compounds, the same two elements can have numerous combinations due to the fact that there are many ways to share electrons. As such, covalent naming uses the following prefixes:

Prefix

Meaning

Prefix

Meaning

Mono-

1

Hexa-

6

Di-

2

Hepta-

7

Tri-

3

Octa-

8

Tetra-

4

Nona-

9

Penta-

5

Deca-

 10

 

  1. Name the first element with its appropriate prefix (unless the prefix is mono. No “mono-“ prefix for the first element.)  Then, name the second element with its appropriate prefix AND change the ending to an “-ide”.
      Example:  N2O is Dinitrogen monoxide, CO2 is Carbon                          dioxide.

 

 

Determining the Charge of a Transition Metal

  1. Calculate the total negative charge from all the anions. Then, recognize that there will be an equal amount of positive charge.
    Example Problem:  Name the compound with the formula Fe2(SO4)3.
    1406 Ch 6 img 2.jpg

  2. Divide the total positive charge by the number of transition metal ions to get the charge per transition metal.

                  LaTeX: \frac{+6}{2\:Fe\:ions}=+3\:charge\:per\:Fe\:ion+62Feions=+3chargeperFeion                    

The name of the above compound is therefore “Iron(III) sulfate”.

 

 

Writing Ionic Formulas

  1. Identify the ions present along with their charges. 
    Example:  Magnesium Phosphate would have the ions Mg2+  and PO43-.

 

  1. Find the lowest common multiple of the two different charges.

      Example:  1406 Ch 6 img 3.jpg The lowest common multiple is 6.

 

  1. Determine how many cations would be needed to give you a total positive charge equal to the LCM (lowest common multiple) and how many anions would be needed to give you a total negative charge equal to the LCM.
    Example:         3 Mg2+ =+6                and                  2 PO43- = -6

 

  1. Write the formula for the compound using chemical symbols and indicate how many of each ion with subscripts. If there is more than 1 polyatomic ion needed, use parentheses.
    Example:  Mg3(PO4)2

 

 

Lewis Structures

Lewis Symbol:  A symbol that represents the valence electrons of an element as dots which are placed on the four sides of that element’s chemical symbol.           

Example:

 Chapter 4 img 1.png

 

 

Lewis Structures:  Drawings that show how atoms are bonded together and electron placement in a molecule using lines (dashes) to represent shared pairs (i.e. bonds) and dots to represent unshared electrons.  For a systematic method of drawing Lewis Structures for covalent compounds, please see the handout “How to Draw Lewis Structures for Covalent Compounds” available online at the TCC Learning Commons.

Lewis Structures showing electron transfer: 

Chapter 4 img 2.png

 

 

 

Polarity of Bonds and Molecules

Bond: A shared pair of electrons between two atoms that holds those atoms together.

Electronegativity:  A measure of the ability of an atom to attract the shared electrons of a bond to itself.

Polar Covalent Bonds:  Bonds that involve the unequal sharing of electrons between the two atoms.  This occurs when the atoms have different electronegativities.  The atom with the higher electronegativity will have a slight negative charge whereas the atom with the lower electronegativity will have a slight positive charge.  A few examples would be the H-F bond, the C-H bond, and the S-O bond.

Non-Polar Covalent Bonds:  Bonds that involve equal sharing of electrons between the two atoms.  This occurs when the atoms have the same (or nearly the same) electronegativity.  Some examples would be the bonds of the diatomic elements (Br2, I2, N2, Cl2, H2, O2, and F2)

 

Chapter 4 img 3.png

Chapter 4 img 4.png

 

 

VSEPR Theory

      The VSEPR (Valence Shell Electron Pair Repulsion) theory states that molecules will arrange themselves in such a way as to maximize the distance between electron domains in each atom.  This allows us to predict electron domain geometry and molecular geometry.  Electron domain geometry is the arrangement of electron domains around the central atom.  Molecular Geometry is the arrangement of ONLY the atoms in a molecule or ion.

 

To predict the shapes of molecules or ions using the VSEPR model, use the following steps:

  1. Draw the Lewis structure of the molecule or ion and count the number of electron domains around the central atom.  (See above definition of Electron Domain)
  2. Predict electron domain geometry by arranging electron domains around the central atom so that the repulsions among them are minimized (i.e. maximum distance between domains).  See table below:
  3. Use the arrangement of bonded atoms to determine the molecular geometry.

Chapter 4 img 5.png

 

 

Simplified Rules for Determining Molecule Polarity (Requires Lewis Structure)

  1. If a molecule has one lone pair of electrons on the central atom, the molecule is polar. If a molecule has 2 or more lone pairs on the central atom, the molecule will be polar UNLESS those lone pairs are in geometrically opposite positions AND all the atoms bonded to the central atom are the same.
  2. If a molecule has no lone pairs and the atoms bonded to the central atom are all the same, the molecule will be non-polar. If a molecule has no lone pairs and the atoms bonded to the central atom are NOT the same, the molecule will usually be polar.
  3. For any diatomic (2 atom) molecule, the polarity of the bond is the same as the polarity of the molecule.