Principles of Physical Science (POPS) - Teacher Notes
Bonding - Molecules and Compounds – Chemistry Unit 3
What’s the difference between a molecule and compound?
*
A molecule is formed when two or more atoms join together chemically.
*
A compound is a molecule that contains at least two different elements.
*
All compounds are molecules but not all molecules are compounds.
Who’s involved?
*
In the last unit we learned that an atom is the smallest unit of matter “indivisible”
*
And that an atom has electrons, protons, and neutrons. The protons and neutrons are packed
together in hard center called the nucleus.
*
The electrons are energy bundles that rotate around the nucleus. This is where we will focus
this unit.
Electrons – What make it all happen
*
Atomic number = number of electrons if not an ion
*
Electrons vary in the amount of energy they possess, and they occur at certain energy levels or
electron shells.
*
Electron shells determine how an atom behaves when it encounters other atoms
*
The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.
*
The electrons in the outer shell are called valance electrons
What is a Bond?
*
Most atoms don’t exist by themselves
*
Most atoms are attracted to other elements and they exist together through bonding.
*
Group 18 elements do not bond.
Why do atoms bond?
*
Atoms are most stable when they have a full shell of electrons
*
Most atoms do not have this so they are naturally attracted to other atoms to gain a full shell
of electrons.
*
When atoms share or borrow electrons to fill their outer electron shell this is called bonding
*
Bonding is between more than one atom and can often be between different elements
Valence Electrons
*
These are electrons in the outer energy level of the atom.
*
These are the available electrons for bonding.
How do you know how many valance electrons elements have?

The number of valance electrons can be determined by looking at the group number on the
periodic table. For example In group 1 elements all atoms have 1 valance electron.

Group 2 – 2 valance electrons

Group 3 – 3 valance electrons

Group 14 – 4 valance electrons

Group 15 – 5 valance electrons

Group 16 – 6 valance electrons

Group 17 – 7 valance electrons

Group 18 – 8 valence electrons
How do I show bonding?
*
Scientists use Lewis dot structures to show the number of valance electrons and the process of
bonding.
Lewis Dot Structures
Symbols of atoms with dots to represent the valence-shell electrons - They can be used to
show bonding. How many more does each of them need to get to eight?
How do you know who bonds with who?
*
Octet Rule - atoms tend to gain, lose or share electrons so as to have 8 electrons in their
valence shell (outermost shell of electrons)
*
This means that if an atom has 3 valence electrons it will lost those three in order to have a full
inner orbital
*
This also means that if an atom has 6 valance electrons the atom will gain two more in order to
have a total of eight for the whole atom.
Octet Rule
For Example
*
Hydrogen gains 1
*
Carbon gains 4
*
Nitrogen gains 3
*
Oxygen gains 2
Oxidation number

Oxidation Number (oxidation state) - Electric charge of an atom or group of atoms in
becoming stable.
-Example – oxygen has an oxidation number of -2 because it has six valance electrons and takes two
more to get to eight.
-Example Sodium has an oxidation number of +1 because it is easier to give up one because
underneath there are eight electrons.

You try – what would be the oxidation number of potassium?
Ionic Bonding
*
bond formed between two ions by the
transfer of electrons
*
When one atom donates electrons to another so that the other atom has a full shell of
electrons. This creates a positive and negative charge which pulls the two atoms together.
What is an Ion?
*
Ions - An atom that has become charged by gaining or losing an e-
*
There are two ions in an ionic bond.
*
Cations - positive ions lose electrons
*
Example: Na+ is a cation because it tends to lose one electron during an ionic
bond
*
Anions - Negative ions gain electrons
*
Example: Cl– is an anion because it tends to gain one electron during an ionic
bond.
Example of Ionic Bonding
Another Example of Ionic Bonding
Other Ionic Bonding Examples
*
Combine
Ca+2 with Br-1 = CaBr2
*
Ba+2
with
S-2 =
BaS
*
Mg+2 with
P-3 =
Mg3P2
*
Al+3
with
Br-1 = AlBr3
*
K+1
with
S-2 = K2S
*
Fe+3
with
Cl-1 = FeCl3
Polyatomic Ions
*
Polyatomic ions - group of atoms with a single charge Ionic bonding is the
attraction force between two ions of the opposite charge
Why do polyatomic ions have a charge?
Examples of Polyatomic Ions
Covalent Bonding
Covalent Bonds
*
When two atoms share electrons (e-) so that their outer shell of electrons is full
*
Between two nonmetals
*
Examples CO2
Covalent Bonding
*
If you count, you’ll see that the Lewis structure shows they both have 8 electrons in their
valance shell
*
This force holds them together –
*
*
The Energy of Bonds: ATP-ADP
Most substances in living organisms
have covalent bonds
*
Covalent Bond – Example
*
*
Carbon dioxide
*
Another Example of Covalent Bonding
An electronegativity difference dictates how much sharing
What are diatomic Molecules? Diatomic molecules - two similar atoms to form a neutral molecule - have the same name as the
element- these are covalent bond.
Examples:
H2 , O2 , N2 , F2 , Cl2 , Br2 , I2
Chemical Formula
*
Chemical Formulas - symbolic representation of a compound
*
Subscript - the terms subscript is given to the number that is written beneath each chemical
element. It indicates the number of atoms of this element in the chemical compound.
Chemical Formulas
*
Coefficient - a coefficient is the number that normally appears at the beginning of a term in a
chemical equation. It indicates the number of molecules or atoms that were involved in the
chemical reaction.
Naming Ionic Compounds Guidelines for Naming Ionic Compounds-NOT TRANSITION METALS
*
The cation is always written first
*
The full name of the metal is given
*
The non-metallic ion is written after the metal. The ending on the non-metal is dropped and –
ide is added
*
A compound is not capitalized unless it is at the beginning of a sentence.
*
***There are three exceptions for the transition metals. Silver, Cadmium and Zinc. These three
metals follow the basic naming guidelines outlined above.
Examples and You Try
*
1. KI _ ___potassium iodide________
2. Ag2O___silver oxide______
*
3. MgCl2___magnesium chloride_____
*
4. Li2O ________________________
*
5. NaCl ________________________
*
6. Al2S3 ________________________
*
7. CaCl2________________________
*
8. MgO ________________________
*
9. NaF ________________________
*
10. ZnCl2________________________
Naming Ionic Bonds with Polyatomic Ions
*
These are ionic bonds. The polyatomic ion is the anion (or in rare cases the cation) and the
metal is still the cation. The bonds within the polyatomic ion are covalent, but the bond
between the cation and anion are ionic.
Guidelines for naming…
*
The cation is always written first.
*
The full name of the cation is named (most often this will be just a metal)
*
The anion is written after the cation (most often this will be the polyatomic ion). The ending on
the polyatomic ion will be –ate or –ite. This depends on the number of oxygens in the ion itself.
One more oxygen means “ate” and less oxygens means “ite”
*
Examples and You Try!
*
Practice:
*
CaCO3 _calcium carbonate___________
*
(NH4)2(SO4)_ammonium sulfate______________
*
AgNO3 _silver nitrate____________________
*
Al2(CO3)3___________________________
*
NaNO3 ___________________________
*
NaOH ___________________________
Guidelines for Naming Ionic Compound - TRANSITION METALS
*
The cation is always written first
*
The full name of the metal is given
*
The number of electrons lost is listed after the metal
*
The non-metallic ion is written after the metal. The ending on the non-metal is dropped and the
appropriate suffix is added. It will be ide or ite, ate if it’s a polyatomic ion.
*
If using the old fashioned classical method, instead of listing the number of electrons lost after
the metal, an ic or ous is added. Ic stands for higher electron loss than ous so we can say that it
is “ic”ky to lose electrons.
*
***There are three exceptions for the transition metals. Silver(1+), Cadmium (2+) and Zinc(2+).
These three metals follow the basic naming guidelines outlined above.
Examples and You Try
*
Zn3P2____Zinc II phosphide_________
*
Cu2O____Copper (IV) oxide_____
*
Ag3PO4_____________________________
*
CdSO3______________________________
*
Fe(HCO3)2____________________________
*
Sc(OH)3_____________________________
*
V2(SO4)3_________________________________
Guidelines for naming covalent compound
*
1
2
3
4
5
6
7
8
9
10
Look at the symbol for the first written element in the compound. If a number follows the
symbol write the prefix that corresponds to that number and then add the elements name after
the prefix.
monoditritetrapentahexaheptaoctanonadeca*
***If the number for the first element is one you do not need to use a prefix.
*
Now look at the second element in the compound. Write the prefix corresponding to the
number of times the second element appears. You need to write mono if there is only one of
the element.
*
Change the ending of the second element to –ide.
*
***If the element begins with a vowel drop the “O” in mono
*
Prefixes for Covalent Compounds
Examples and You try
*
CO2
____carbon dioxide___________
*
ClO2
____chlorine dioxide____________
*
NF3
___nitrogen trifluoride__________
*
PI3
________________________
*
Si2Br6
________________________
*
CCl4
________________________
*
P4S5
________________________
*
NO3
________________________
*
SbBr3
________________________
*
CO
________________________
*
Some important chemistry vocabulary
A. Binary: Compound made of only 2 elements
B. Diatomic: two similar atoms to form a neutral molecule - have the same name as the element
C. Empirical Formulas: simplest ratio of atoms of a compound
D. Molecular formulas: actual ratio of atoms of a compound
E. Coefficients: number used to represent the number of formula units of a compound
F. Electronegativity: Relative tendency for an atom to gain an electron.