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Introduction to Bonding
Most of the matter on Earth is in the form of compounds. Even when a substance exists as a pure
substance, it tends to eventually combine with other elements. For example, if you leave an iron nail
out in the rain it will combine with the oxygen in the air to form iron oxide (rust). In this investigation,
you will build models of atoms and discover one of the fundamental ideas of chemistry: how electrons
are involved in the formation of chemical bonds.
You have already learned:



A neutral atom has the same number of electrons and protons.
The electrons occupy energy levels surrounding the nucleus.
The electrons occupy the lowest energy level first. Once an energy level is full, electrons will
start to fill the next level.
Part I: The Valence Electrons
Directions: Using the atom building game, build each element in the following table. You do not have to
add protons and neutrons to the models, only the electrons, since they are the only subatomic particles
involved in bonding. The first one is done for you as an example.
Element
Atomic
Number
B
He
Li
F
Ne
Na
Cl
Be
K
N
O
5
Electrons in
outermost
energy level
3
How does it become
stable?
Ion
formed
Lose 3 electrons
B3+
Use your table to answer the following questions:
A. What do lithium, sodium, and potassium have in common?
B. What do fluorine and chlorine have in common?
C. The electrons in the outermost energy level of an atom are called valence electrons. These are
the electrons involved in chemical bonds. How is this number related to the group number of
the element?
Part II: Modeling a Chemical Bond
Atoms that have a complete outermost energy level are stable. We have been using the yellow
marbles and either adding more or taking away marbles to show the atom gaining or losing electrons,
but what really supplies those electrons? Another atom either gives or takes away electrons to make
the atom stable. This is how chemical bonds are formed.
Directions:
1. Take out two atom building game boards and the red, blue and yellow marbles.
2. Find sodium on the periodic table.
a. How many protons does sodium contain? ____ Add this many red marbles to the first
board.
b. How many neutrons does sodium contain? ____ Add this many blue marbles to the first
board.
c. How many electrons does sodium contain? ____ Add this many yellow marbles to the
first board.
3. Find chlorine on the periodic table.
a. How many protons does chlorine contain? ____ Add this many red marbles to the second
board.
b. How many neutrons does chlorine contain? ____ Add this many blue marbles to the
second board.
c. How many electrons does chlorine contain? ____ Add this many yellow marbles to the
second board.
Questions:
A. What will sodium do to become stable? _______________________________
B. What will chlorine do to become stable? ______________________________
C. Why might these two atoms bond together to form a molecule?
An element’s oxidation number is the same as the charge an atom has when it ionizes (in other words,
when it gains or loses electrons to become stable).
Move one electron from the outer ring of sodium to the outer ring of chlorine.
D. Are the sodium and chlorine now stable? _________________
E. What is sodium’s oxidation number? (Hint: what is the charge on it?) ______
F. What is chlorine’s oxidation number? ______
G. Suppose the sodium ion and chlorine ion stayed together as a molecule.
How many total positive charges does this molecule have? _______
How many total negative charges does this molecule have? _______
What is the molecule’s overall charge? _____
Compounds overall are electrically neutral. To show this, we add together the oxidation numbers:
Oxidation number of sodium
________
+
Oxidation number of chlorine ________
Equals:
________
In a compound, the sum of the oxidation numbers is always __________.
Part III: Oxidation Numbers and the periodic table
Remember, the oxidation number is the charge an atom has when it loses or gain electrons to become
stable. It indicates how many electrons are lost or gained. If the atom loses one or more electrons,
the number is positive. If the atom gains one or more electrons, the number is negative. We can’t build
a model every time we want to find an element’s oxidation number. Luckily, the periodic table shows
some trends that we can remember to help us make things easier.
Look at the following partial periodic table. We are just going to be working with the “A” groups of
elements.
1. On the first line, write the correct Roman numeral for that group.
2. On the second line, write the number of valence electrons that group has.
3. On the third line, write what the elements must do to become stable. (i.e. gain 2 e- lose 3 e- etc.)
4. On the fourth line, write the charge the elements will have when they become stable. (This is
the oxidation number.)
Part IV: Naming Ionic Compounds
Compounds that are formed from ions are called ionic compounds.
1. Using the periodic table on the previous page, write down the ion formed by each element (symbol
and oxidation number).
Sodium ______
Sulfur ______
Chlorine ______
Oxygen ______
Magnesium ______
Aluminum ______
Beryllium ______
Nitrogen ______
2. The cards on the card stock page represent the ions. A peg represents a negative charge (extra
electrons that have been gained); a notch represents a positive charge (missing electrons –they
have been given away).
4. For each compound below, put the cards together so that each notch is matched by a peg. This will
provide equal numbers of positive and negative charges, as in real ionic compounds.
5. Write down the number of each element that you need to create a neutral molecule. Add the
oxidation numbers of the elements represented to make sure it is zero.
6. Write the chemical formula in the box provided. The first one has been done for you.
Elements
Cation
Anion
sodium and oxygen
Na1+
O2sodium and chlorine
beryllium and sulfur
magnesium and oxygen
aluminum and chlorine
sodium and nitrogen
beryllium and oxygen
aluminum and oxygen
sodium and sulfur
beryllium and nitrogen
magnesium and chlorine
magnesium and sulfur
aluminum and sulfur
beryllium and chlorine
magnesium and nitrogen
aluminum and nitrogen
Questions:
A. Can you join together sodium and aluminum? Explain.
Number
of Cation
needed
2
Number
of Anion
needed
1
Chemical Formula
Na2O
B. Look at your cation, anion, and your chemical formula. Is there a “shortcut” way to get to the
formula from the ions? Explain.
Writing Formulas for Binary Ionic Compounds
Things to remember:
 Ion – an atom that has gained or lost electrons, so has a positive or negative charge.
 Cation – an ion with a positive charge; Anion – an ion with a negative charge
 Oxidation number – the number that indicates how many electrons are lost or gained
during bonding. This is the same as the charge on the stable ion of that element.
 In order to gain stability, the sum of the oxidation numbers for a compound must equal
zero.
 Binary compound – a compound formed between a metal and nonmetal
Writing formulas for binary compounds:
1. Write the ion (symbol and oxidation number) of the cation first.
2. Write the ion (symbol and oxidation number) of the anion second.
3. Write subscripts (small numbers near the bottom) is to indicate how many ions of
each element is needed to make the sum of the oxidation numbers zero. Do not
write the number 1 as a subscript.
Examples:
Magnesium and Oxygen
1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable.
Mg2+
2. The negative ion will be oxygen (Group VI), since it gains 2 electrons to become stable.
Mg2+ O23. The sum of the oxidation numbers is already zero, so we only need one of each ion.
Mg1O1 --- we don’t write the ones, so it is MgO
Magnesium and Bromine
1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable.
Mg2+
2. The negative ion will be bromine (Group VII), since it gains 1 electron to become stable.
Mg2+ Br13. The sum of the oxidation numbers is not zero, so we only need to add more ions until it
becomes zero.
Mg2+ Br1Br1By adding another bromine ion, we have +2 + -2, which = 0.
So, our formula is MgBr2
Practice
1. Sodium and oxygen
2. Calcium and fluorine
3. Iron(II) and sulfur
Crisscross Method
A shortcut to figure out the number of each ion needed to form a neutral compound is called
the crisscross method:
1. Write each ion – cation first, followed by the anion.
2. Cross the number that is a superscript of one to be the subscript of the other. (Just
cross the number, not the positive or negative sign. The sign is dropped.)
3. Reduce your subscripts if they are not in the lowest possible ratio.
Examples:
Magnesium and Bromine
1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable.
Mg2+
2. The negative ion will be bromine (Group VII), since it gains 1 electron to become stable.
Mg2+ Br13. Cross the two numbers and drop the positive or negative sign
Mg2+
Br1-
So, our formula is MgBr2
Magnesium and Oxygen
1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable.
Mg2+
2. The negative ion will be oxygen (Group VI), since it gains 2 electrons to become stable.
Mg2+ O23. Cross the two numbers and drop the positive or negative sign
Mg2+
O2-
So, our formula is Mg2O2
But you can reduce 2:2 to 1:1, so the formula becomes MgO.
Practice
1. Sodium and sulfur
2. Lithium and iodine
3. Copper(IV) and oxygen
Binary Compounds with Polyatomic Ions
Polyatomic Ions are covalently-bonded groups of atoms that behave as a unit and carry a charge
EX:


NO3Nitrate ion
C2H3O2acetate ion
the charge that is shown is for the whole group of atoms together, not just the last
element
these ions have very specific names – you will use the following list to find the name and
formula for the polyatomic ions you will use in this class.
Name
Ammonium
Acetate
Chlorate
Hydroxide
Nitrate
Carbonate
Sulfate
Sulfite
Phosphate
Nitrite
Chlorite
Cyanide

Formula
NH41+
C2H3O21ClO31OH1NO31CO32SO42SO32PO43NO21ClO21CN1-
Use the crisscross method for polyatomic ions, just as you did for regular ions. Never
change anything within the polyatomic ion – if you need to add a subscript, put it
outside of a parenthesis.
Examples:
Magnesium Phosphate
Mg2+ PO43Mg3(PO4)2
Practice
1. magnesium carbonate
2. calcium hydroxide
3. manganese(IV) sulfite
Naming Binary Compounds
Ionic compounds have two word names:
1. Write the name of the cation first – this is just the name of the element
2. If the cation is a transition metal, add a Roman numeral in parentheses to indicate the
oxidation number of the ion in this compound.
a. to find the oxidation number do the crisscross method backwards.
b. check the negative ion to see if it has the correct oxidation number. If not, multiply
both charges by the appropriate factor until it does.
c. remember: the Roman numeral only goes in the name, not into the formula.
3. Write the name of the anion – take the root of the element name and end with –ide.
For example:
1. NaF is __________________________________________
2. MgO is _________________________________________
3. CuCl2 is _________________________________________
4. PbO is __________________________________________
Practice
1. Name the following compounds
a. ZnS
b. K3N
c. BaO
d. CaO
e. AlF3
f. CuI2
Naming Polyatomic Compounds
 It is important that you recognize the polyatomic. If there are more than 2 elements in
the formula, then you know it contains a polyatomic. Note: only one polyatomic is a cation,
ammonium (NH4+) All the rest will come after the metal ion.

You name polyatomic compounds just like you would other ionic compounds, but do not
change the ending of the polyatomic ion.
EX: Name ZnSO4 ____________________________
EX: Name NaOH ____________________________
Practice
1. Name these compounds
a. CaCO3
b. KClO2
d. Al(OH)3
e. Mg3(PO4)2
Notes: Bonding
I. Bonding
 Remember, a compound consists of 2 or more elements chemically bonded together.
Examples: H20, NaCl, Sb3(PO4)5
 smallest unit of a compound is called a molecule
 Only electrons move during bonding.
 it is the movement of electrons in the outer energy level (valence electrons) that
determines how one atom will bond with another.
 Octet rule – all atoms need 8 valence electrons to be stable (the exception is those
atoms that are stable by filling the first energy level, they only need 2)
 4 types of chemical bonds:
1. Ionic bonds
2. Covalent bonds
3. Polar covalent bonds
4. Metallic bonds
II. Ionic bonds
 Electrons are transferred from one atom to another.
 When an atom loses an electron it becomes a positive ion (cation).
 When an atom gains an electron it becomes a negative ion (anion).
 Opposites attract, so the cation is attracted to the anion
 This is the strongest type of bond
 Ionic bonds form between metals and nonmetals
 Where are the nonmetals found on the periodic table? ____________
 Where are the metals found on the periodic table? _______________
Example: NaCl
Lewis dot:
Energy levels:
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III. Covalent Bonds
 Electrons are shared between atoms, but not transferred
 The electrons are in the outermost energy level of both atoms at the same time
 Covalent bonds form between nonmetals
Example: NH3
Lewis dot:
Energy levels:
Ionic vs. Covalent Properties
type of elements
type of bonding
structure
melting point
hardness
conduct electricity?
Covalent Bonds
Nonmetals only
share electrons
not rigid – solid, liquid or gas
low
soft  hard
no
Ionic Bonds
Metals and Nonmetals
transfer electrons
crystalline
high
very hard
yes (when dissolved in water)
Practice
1. Determine whether the following are ionic or molecular covalent compounds
a. N2O5
b. a liquid at room temperature
c. PbNO3
d. a salt that conducts electricity when in solution
e. KF
f. AgCl
g. gasoline
h. PCl3
2
IV. Polar Covalent Bonds
 A special type of covalent bond in which the electrons are shared unequally
 The electrons spend a little more time with one atom than with the other.
 This causes the “stronger” atom in the molecule to be slightly negative, while the
other atom is slightly positive.
Example: H2O
Lewis dot:
Energy levels:
V. Metallic Bonds


Bonds between positive ions and free-roaming electrons.
These are flexible and are good conductors as solids (electrons are free to move)
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