Molecular (Covalent Compounds) Flashcards
Ionic vs. molecular compounds: ratios
The compound NaCl includes sodium and chloride ions in a 1:1 ratio
We could build a crystal from 8 ions of each element or 8 million of each and it would form the same structure and would have the same properties
This is not true of the bonding of nonmetals with each other!
For example
C2H2 is NOT the same as C6H6
The Covalent Bond
To get full valence shells, each element shares its electron with the other element
Molecular Elements
There are 7 elements that never exist as a single neutral atom
Instead, they form covalent bond(s) with themselves and exist as diatomic molecules
H, O, F, Br, I, N, Cl ALWAYS exist as H2, O2, F2, Br2, I2, N2, Cl2
Say “HOFBrINCl” to help you remember!
Lewis Structures of Molecular Compounds
1 pair of shared electrons is represented with a dash (−) that connects the two element symbols
You must show all other non-shared valence electrons, and write them in pairs.
Don’t draw the square brackets or charges.
A lone pair of electrons
Non-bonded pair of electrons, is called a lone pair. Usually drawn on the side of the element.
Shared Electrons
1 pair of shared/bonded electrons (1 covalent bond)
Bonding Capacity
Bonding capacity indicates how many bonds an atom tends to make.
As we will see, an atom can be part of a molecular compound and make a number of bonds not equal to its bonding capacity if other criteria are met!
Structural Formulas
Chemists often simplify their diagrams even further by only including the covalent bonds in their diagrams
This representation, without the lone pairs, is called a structural formula
Drawing Lewis Structures for more complicated molecular compounds
NO3-
Rule 1: The central atom is usually the least electronegative atom (or largest bonding capacity)
Rule 2: Other atoms surround the least electronegative atom
Rule 3: Count the total number of valence electrons, including charges
-1 ionic charge = gained 1 electron = +1 electron
+2 ionic charge = lost 2 electrons = -2 electrons
Rule 4: Place electron pairs between each atom to represent a single covalent bond, then distribute the remainder of the electrons around the surrounding atoms (except hydrogen) to satisfy the octet rule. [Any remaining electrons that have not been distributed can be placed on the central atom as lone pairs]
[Should an octet not be satisfied on central atom, non-bonding electrons on the outer atoms can be used to create a double bond]
Rule 5: Replace electrons with lines to represent bonds and place square brackets around the final structure. Indicate the charge on the structure by placing the charge on the outside of the right square bracket.
polyatomics are never found alone, they always combine.
Your Lewis structure is correct if…
Every atom has a full valence energy level (8 electrons for every atom except hydrogen, which is full with 2 electrons)
The total number of valence electrons on the structure is correct (you determined this total in step 2)
The structure has the proper overall shape (everything is spread out as far apart as possible around the central atom – remember that lone pairs and bonds take up the same amount of space!)
If the central atom has two lone pairs, remember to put them adjacent to each other!
Exceptions to the octet rule
Consider nitrogen dioxide, NO2
No matter how you draw it, there’s always one atom that doesn’t have a full octet!
To understand why, we need a more advanced model of the atom (Grade 12!)
Molecular compound properties: State at room temperature
Solid, liquid, or gas
Molecular compound properties: Physical Properties
Solids can be soft, waxy, flexible, or crystalline
Molecular compound properties: Relative Melting point/boiling point
Lower than ionic
Molecular compound properties: Solubility in Water
Some good, some poor
(Can be answered with polar/non-polar compounds)
non-polar compounds can’t dissolve in water (from the beaker lab)
