Quarter 2 Unit 4
Compounds and Atomic Stability:
Learning Objectives
 Understand why atoms form compounds
 Understand chemical stability and the octet rule
 Understand how Ionic bonds and Ionic compounds
are formed
How Elements Form
Compounds
 Compounds: a chemical combination of two or more
different elements joined together in a fixed proportion
 Collisions between particles of the atom cause reactions
 Reactions between atoms Involve Only The Electron Cloud!


Remember, chemical properties of elements on the periodic table
repeat because the pattern of valence electrons repeat in each period
Valence electrons of colliding atoms react to form compounds
 Compounds form when electrons in atoms rearrange to
achieve stable electron configurations.
Chemical Stability
 Nobel Gasses: they are almost
completely un-reactive
 None of these have ever been found
naturally in the environment as a
compound.
 They are extremely un-reactive or stable
 Group 18 (except He) all have 8 valence
electrons

Electron arrangement determines chemical
properties
 The electron arrangement of the noble
gasses is the cause of their stability
(extremely un-reactive)
Check Your Understanding
Turn and discuss the following with the person on your right,
then answer each of the questions in your notebook.
1.
Identify each of the following as a compound or not a
compound
a.
b.
c.
d.
Water
Nitrogen
Carbon dioxide
Deuterium
What subatomic particles are involved in forming
compounds?
3. Why do elements form compounds?
4. Which group of elements rarely forms compounds?
Explain why.
2.
Answers
1.
Identify each of the following as a compound or not a
compound
a.
b.
c.
d.
Water--- compound
Nitrogen---- no, an element
Carbon dioxide-----compound
Deuterium-----no , an element
What subatomic particles are involved in forming
compounds?--- valence electrons
3. Why do elements form compounds? --To rearrange
electrons in order to achieve a stable electron
configuration
4. Which group of elements rarely forms compounds?
Explain why. ---Noble gasses, because they already have a
stable electron configuration
2.
The Octet Rule
 Atoms combine to become
more stable
 Atoms become stable by having eight electrons in
their outer energy levels (He is an exception w/ 2
because it is so small)
 Atoms become stable by achieving a noble gas
configuration
Achieving Chemical Stability
 Collisions between atoms, which involve enough
energy, can cause valence electron rearrangements
 Forming a stable octet
 Noble gas configuration
 Total number of electrons never change
Establishing Stable Octets
 2 options
Transfer of electrons
between atoms
2. Sharing electrons
between atoms
1.
Electrons are Transferred
Example
 Na- sodium is a shiny gray
metal which quickly
oxidizes upon exposure
to the atmosphere
 Cl - Chlorine gas is a dense,
pale yellowish-green,
poisonous, gas
 Na and Cl react to form–salt
Electron Transfer On a Subatomic
Level
 Na- a group 1 element, with 1
valence electron
 Cl - a group 17 element, with
7 valence electrons
 How can the valence
electrons be rearranged to
provide a stable octet for
each?
Chlorine Gets a Stable Octet
 If the Cl atom gains an




electron from the Na atom,
it will achieve a noble gas
configuration
It will be stable
It will have a complete octet
It will also have a negative
charge, because it has an
extra electron
It will be an ion
Na Gets a Stable Octet
 If Na gives away its 1 valence
electron, what will it have for a
valence number?

11 electrons minus 1 electron =
10 electrons
 1s 2s 2p
- 8 valence electrons
It will be stable
It will have a complete octet
It will also have a positive
charge, because it has one less
electron
It will be an ion
2




2
6
Transfer of the Electron Forms an
Ionic Bond
An Ionic Bond
 The exchange of the electron creates 2 ions,
 1 positively and 1 negatively charged
 Ionic bond: A strong attractive force between atoms
with opposite charges which is formed by an electron
transfer between atoms.
 Ionic bonds form ionic compounds
 A compound that is made up of ions
 Na+ Ions are attracted to all nearby Cl- Ions and visa
versa thereby forming crystal structures

Cl
Na+
The Results of Ionic Attraction
 Affects properties of the compound
 Example; sodium chloride (salt) is a crystal because of
intermolecular forces of attraction between ions, it is a
solid at room temperature
 Melting Ionic Compounds: Breaking the strong crystal
structure requires a lot of energy, therefore the melting
point of NaCl is more than 800°C
 Hardness and Brittleness: It takes a great deal of force
to break the structure of an ionic crystal
Naming Binary Ionic Compounds
 Binary compounds are compounds with only two
different elements.
1. First write the name of the positively charged ion
(usually a metal)
2. Then add the name of the negatively charged ion
(nonmetal)
3. Modify the negatively charged ion name to end in
-ide
1.
Example NaCl is called sodium chloride
NOT sodium chlorine
Check Your Understanding
 Name compounds formed by the following ions:
Mg+ Cl2. Cl- Ca+
3. Cl- K+
Name these compounds:
1. AlN
2. KI
3. ZnO
1.
Answers
 Name compounds formed by the following ions:
Mg+ Cl- Magnesium chloride
2. Cl- Ca+
calcium chloride
3. Cl- K+
potassium chloride
Name these compounds
1. AlN
aluminum nitride
2. KI
potassium iodide
3. ZnO
zinc oxide
1.
Oxidation Numbers
 Charged atoms or compounds are called Ions
 The total charge on the ion is known as the Oxidation
Number of the atom
 Examples:
 Mg+2 is magnesium ion, the charge is
positive 2, the oxidation number is 2+
 F- has an oxidation number of 1-
 Some metals have the same oxidation number in
all compounds (memorize this)
 Group 1 elements , oxidation number = 1+
 Group 2 elements , oxidation number = 2+
 Aluminum, oxidation number = 3+
The Charge of Ionic Compounds
 In Ionic Compounds, the total positive charge is equal
to the total negative charge
 One Mg2+ ion will combine with 2 Cl- ions
 Forming MgCl2, The total positive charge is 2+, the total
negative charge is 2 +2+-2 = 0
 In a correctly written formula, the sum of the total
positive charges and the total negative charges = 0
Predicting Oxidation Numbers
 Oxidation numbers for most elements can be predicted from




their position on the periodic table.
Groups 3-12 Transition metals are difficult as many of these
elements Have more than one oxidation number depending
on the reaction
Group 13 elements have 3 valence electrons, so will lose 3 and
have an oxidation number of 3+
Group 14 may have 2+ or 4+ oxidation number
Groups 15, 16 and 17 tend to gain electrons to complete the
octet since they are already ½ full
 Their oxidation numbers are 3-, 2-, and 1- respectively
 They can also lose electrons and have positive oxidation numbers
 The tendency to lose electrons increases as you move down the column
Check for Understanding
 Predict the oxidation numbers for the following elements:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
Al
N
Cl
Mg
S
Na
K
O
Ga
P
Se
Br
Answers
 Predict the oxidation numbers for the following elements:
1. Al 3+
2. N 33. Cl 14. Mg 2+
5. S 26. Na 1+
7. K 1+
8. O 29. Ga 3+
10. P 311. Se 212. Br 1-
Writing Chemical Formulas For
Ionic Compounds
 The key to writing formulas is to make the oxidation
numbers add to zero, making a neutral compound.
 Example:
 Ca 2+ , located in group 2
 F 1-, located in group 17
 The formula for a compound of these elements is
 CaF2
1(2+) + 2( 1-) =0
The compound is neutral
Representing Compounds as Formulas
 The formula of a compound tells:
 what elements make up the compound
 and how many of each element are present in one unit
of the compound
 Example : H2O two H for each O
Progress Check
 Write formulas for the following compounds:
1.
2.
3.
4.
5.
6.
7.
8.
Sodium Fluoride
Potassium Chloride
Rubidium bromide
Sodium selenide
Potassium oxide
Lithium sulfide
Strontium fluoride
Calcium Chloride
Answers
 Write formulas for the following compounds:
1.
2.
3.
4.
5.
6.
7.
8.
Sodium Fluoride (Na+1 F-1) NaF
Potassium Chloride (K+1 Cl -1) KCl
Rubidium bromide (Rb+1 Br-1) RbBr
Sodium selenide
(Na +1 Se -2) Na2Se
Potassium oxide
(K+1 O-2) K2O
Lithium sulfide
(Li+1 S -2) Li2S
Strontium fluoride (Sr 2+ F 1-) SrF2
Calcium Chloride
(Ca 2+ Cl 1-) CaCl2
The Formation Of Ionic Compounds
 Lab Learning Objective: To model the transfer of
electrons thereby achieving noble gas configurations and
the formation of stable ionic compounds
 Apply what you have learned!
 The Formation Of Ionic Compounds (MiniLab 4.2)
 Question: What other atoms give up and gain electrons
(creating ions, forming ionic bonds) to form ionic
compounds?
Pre-Lab Procedure
You and your lab partner must locate the following atoms
on the periodic table: Li, S, Mg, O, Ca, N, Al and I.
2. Using the information on the Periodic Table, and what you
have learned, complete the table with the Valence
electrons, class of element, Lewis dot diagram, and electron
configuration for the given atoms
3. Based on what you have discovered about the atoms
1.
above, construct a hypothesis predicting which of
the atoms will give up electron(s) and which will
receive electron(s) when forming compounds.
 REMEMBER a hypothesis is an if… then… statement which
answers the lab question
4. Have your hypothesis checked before you begin the lab.
The Formation Of Ionic Compounds
 Learning Objectives:
 To model the transfer of electrons thereby achieving
noble gas configurations and the formation of stable
ionic compounds
 To name ionic compounds
 To identify formulas of ionic compounds
 Identify a pattern which can be used to predict which
atoms will form Ionic compounds
 Follow the Lab instructions to complete your
investigation.
 Prepare for a poster presentation
Polyatomic Ions
 Some Ions are formed by more than one or two types
of atoms
 They are called polyatomic ions
 Some of these are very common, so recognizing them
will be very helpful to you.
 Let’s look at some
 Some do not follow simple naming rules, often because
they were named before rules for naming were written


These we need to use a reference sheet for, or memorize!
Use your reference sheets to help name complex polyatomic
ionic compounds
Conventions for Naming Polyatomic Ionic
Compounds When Given the Formula
 Steps
Name positive ion (cation) first, then negative ion (anion) second.
1.
If one of these ions has more than 1 atom in it, look up the name on the chart
A.
You should plan on memorizing the ones I have starred!
i.
Determine if you need a Roman numeral in the name
2.
If cation is NOT a transition metal, then NO Roman numeral
If the cation is a transition metal then see if it can have more than 1 oxidation
number, if not then go to step 4
If yes then go to step 3
A.
B.
C.
Determining the Roman numeral, usually the number of anions = the charge
on the cation and the number of cations =anions.
3.
Example: Fe2[SO4}3
1.
1.
2.
1.
2.
There are three SO4 sulfate ions -- charge of 2-= 6 There are two Fe atoms – charge of ? = must total 6
Fe has to be 3+ so use Roman Numeral III Fe(III)
Name is Iron (III) sulfate!
Check to see that the sum of oxidation numbers= 0
4.
If yes then correct!
1.
1.
1.
2.
3.
Example :
3, SO4 ions = 3 x 2- = 62 Fe(III) ions = 2 x 3+ = 6+
6- + 6+ = 0 !!!!!
Let’s Try Another One
 NH4Cl
 Look at your list of common polyatomic ions, do you see
any here?
 Yes NH4 is ammonium it is a cation – 1+
 Follow step 1-- Ammonium is the first part of the name
 Chlorine is the second atom, the anion- 1-, becomes chloride
(just like before)
 So far we have ammonium chloride
 Follow step 2-- Is ammonium a transition metal? No!
 Follow step 4
 1+ + 1- = 0 correct!  ammonium chloride
Now you try it alone!
 Name the following :
 Fe(NO3)3
Solution
 Iron(III)nitrate
 Iron  cation = 2+ or 3+
 NO3  nitrate = 1 Three anions, so the charge needs to be 3+ so Roman
numeral is III
 Check
 1 Fe(III) has charge of 3+
 3 NO3 ions have a charge of 3 x 1- = 3 3+ + 3- = 0 correct!  Iron(III)nitrate
Let’s Try Another
 Al(CN)3
Answer
 Al(CN)3
 Aluminum Cyanide
Establishing Stable Octets
 2 Options
1. Transfer of electrons between atoms – Form Ionic Bonds
2. Sharing electrons between atoms
Obtaining a Stable Octet by Sharing
 Example:
 H (hydrogen)


A gas at room temperature
Has one valence Electron
 O (oxygen)
 A gas at room temperature
 A group 16 element
 Has 6 valence electrons
 Could they achieve a stable
octet by transferring
electrons?
 Could H give up its only
electron and have a noble gas
configuration?

No electrons!
Sharing Electrons Between Atoms
 When atoms collide with enough energy to cause a
reaction
 And neither atom attracts electrons strongly enough to
take electrons from the other atom (small difference in
electronegativity)
 The atoms combine by sharing valence electrons
Reaction of Hydrogen and Oxygen
•Hydrogen needs one
more electron to have the
same electron
configuration as He
•Oxygen needs two more
electrons to have the same
electron configuration as
Ne
•Hydrogen and Oxygen can share one electron from each
atom.
•This makes Hydrogen stable
•But Oxygen still has only 7 valence electrons
Sharing to Make Oxygen Stable
 Oxygen gets a complete octet by sharing an
electron with another hydrogen
 This explains the formula of H2O
Electrons always rearrange in a
chemical reaction
Covalent Bond
 The attraction of two atoms for a shared pair of
electrons is called a covalent bond
 Notice, neither atom has an ionic charge
 A compound whose atoms are held together by
covalent bonds is called a covalent compound
 A molecule is an uncharged group of two or more
atoms held together by covalent bonds
Sharing More Than Two Electrons
 More than two electrons can be shared.
 The reaction Between Carbon and Oxygen for example
 You can arrange these 16 valence electrons to produce a
molecule in which all three atoms have a noble gas
configuration (a complete octet)
Sharing two pairs of Electrons:
Double Covalent bonds are formed
Atoms can also form Triple bonds
Molecular Elements
 Molecular Elements: a molecule that forms when atoms of the
same element bond together
 They are not compounds (not two elements)
7 nonmetal elements naturally found as diatomic molecules:
1. Hydrogen
2. Nitrogen
3. Oxygen
Gases
4. Fluorine
5. Chlorine
6. Bromine -liquid
7. Iodine- solid
Allotropes: molecules of an element that form different crystalline
structures.

The properties of allotropes are usually different, because structure is
important
Formulas and Names of Covalent
Compounds
 Binary inorganic compounds (2 elements, not carbon)
 The Suffix –ide
 Write first nonmetal followed by the name of the second
nonmetal with its ending changed to – ide


The element closest to the left of the periodic table is written first
( some exceptions with H)
If both are in the same group, name the one lower on the column
1st.
 Indicating the number of atoms

Add a prefix to the name of each element indicating the number
of each element
Add prefix to indicate number of atoms
•If only one atom of the first element is listed, the mono is usually left out
•If adding the prefix creates double vowels, the first is usually omitted
Example of Naming Protocol
FORMULAS
NAMES
NO
NO2
N2O
N2O5
Nitrogen monoxide
Nitrogen dioxide
Dinitrogen monoxide
Dinitrogen pentoxide
Practice Naming
 Name the following compounds
1.
2.
3.
4.
5.
6.
CO2
NO2
SO3
PCl3
NO
P2O5
ANSWERS
 Name the following compounds
1.
2.
3.
4.
5.
6.
CO2
NO2
SO3
PCL3
NO
P2O5
CARBON DIOXIDE
NITROGEN DIOXIDE
SULFUR TRIOXIDE
PHOSPHORUS TRICHLORIDE
NITROGEN MONOXIDE
DIPHOSPHORUS PENTOXIDE
General Atomic Bonding Trends
 Two nonmetallic elements usually achieve stability
by sharing electrons to form a covalent compound
 Reacting atoms, when one is a metal and one a
nonmetal, are much more likely to transfer electrons
and form an ionic compound.
Comparing Ionic and Covalent bonds
 When elements combine they form either ions or
molecules
 This changes them dramatically, this is why
compounds have different properties from the
elements that make them up
Properties of Ionic Compounds
 Composed of well-organized, tightly bound ions
 They form strong three-dimensional crystal structures
 They are crystalline solids at room temperature
 They are generally hard, rough and brittle
 They have high melting points
 They are generally soluble in water and form a solution
which conducts electricity (called an electrolyte)
 In liquid state they also conduct electricity
Properties of Covalent Compounds
 Relatively weak Interparticle forces: forces between
particles that make up a substance
 The molecules have no ionic charge so the attractive force
between them is weak
 Many are liquid or gas at room temperature
 If they are solid at room temp, they have low melting point
 Ex. Sugar
 They do not conduct electricity
 Many do not dissolve in water (ex vegetable oil), but some
do (sugar)
 Generally Less soluble in water
Interparticle Forces Make the Difference
 Interparticle Forces are the key to determining a
substances state of matter at room temperature as
well as many other properties
Compare and Contrast Properties
Ionic Compounds
1.
Ions held tightly in solid
state by strong
interparticular forces- solid
at room temp
2. Good conductors in liquid
form and when dissolved in
water
3. Water soluble- ions are
attracted by water
Covalent Compounds
Much weaker
interparticular forces hold
them together more looselyliquid or gas at room temp
(few solids)
2. Not electrical conductors
1.
3.
Less soluble or insoluble in
water-not attracted by water
Check your Understanding
Describe two processes by which elements can form
stable compounds? What is the type of bonding that
occurs from each process?
2. An unknown compound dissolves in water but does
not conduct electricity. Is the compound more likely
ionic or covalent? Explain
3. Why does NaCl have to be heated to 800 degreesC
before melting, and candle wax starts to melt at 50
degrees C.
4. Explain why ionic compounds conduct an electric
current in solution but covalent compounds do not?
1.
Answers
Two processes by which elements can form stable
compounds are transferring valence electrons (Ionic
Bonds) or sharing valence electrons (Covalent Bonds)
2. It is more likely to be covalent because ionic compounds in
solution would be made of mobile ions that would conduct
an electrical current
3. NaCl is an ionic compound and is held together by strong
interparticle forces which have to be broken in order to
melt, candle wax is a covalent compound with very weak
interparticle forces so melting is easy.
4. To conduct electricity there must be ions that are free to
move so that they can pass the electrons along, this does
not happen with covalent compounds but does with ionic
compounds in solution.
1.
Predicting the Type of Bond
The type of bond formed depends on two properties
Electron configuration
Attraction for electrons: electronegativity
1.
2.

In some covalent bonds the sharing is not even
These bonds are called polar covalent bonds




They have some ionic character
The unequal sharing creates 2 poles (- end and + end)
The less electronegative atom will be positive, less attraction for
the electron
Differences in electronegativity can determine bond type



Differences greater than 2.0 makes ionic bond
Differences between 0.8 and 2.0 makes polar covalent
Differences between 0.0 and 0.5 covalent
Lets See how This Works!
 The bond between calcium and oxygen, is it ionic,
covalent or polar covalent?
 Ca electonegtaivity is 1.0
 O electonegtaivity is 3.5
 3.5 - 1.0 = 2.5
 2.5 is greater than 2.0 so it is an ionic bond!
Applying Our Knowledge
 Using figure 9.2 0n page 302, determine the
electronegativity difference to classify the bonds
between the following atoms as ionic, covalent or polar
covalent.
1. NiO
2. BN
3. CaCl2
4. FeSi
5. NaF
6. Zn3P2
Answers
1.
2.
3.
4.
5.
6.
NiO
BN
CaCl2
FeSi
NaF
Zn3P2
Ni 1.8, O 3.5, 1.8 – 3.5 = 1.7 polar covalent
B 2.0, N 3.0, 2.0 – 3.0 = 1.0 polar covalent
Ca 1.0, Cl 3.0, 1.0 – 3.0 = 2.0 ionic
Fe 1.8, Si 1.8, 1.8 – 1.8 = 0.0 covalent
Na 0.9, F 4.0, 4.0 – 0.9 = 3.1 ionic
Zn 1.6, P 2.1, 1.6 – 2.1 = o.5, covalent
Bonding in Metals
 Bonding in metals DOES NOT RESULT IN
COMPOUNDS
 It results in an interaction that holds metal atoms
together and accounts for some of their common
properties.
 Remember, metals are Malleable (thin sheets) and ductile
(drawn into wires) and conduct electricity
 Conductivity is a measure of how easily electrons can flow
through a material- electrical current.
THESE PROPERTIES ARE THE RESULT OF THEIR BONDS
How Do They Bond?
 Metal to Metal bonding does not involve valence electron
transfer, as it does in metal-nonmetal bonding
 Metal atoms release their valence electrons in a sea of
electrons, shared by all of the metal atoms
 This is called a metallic bond
 Positively charged nucleus surrounded by loosely associated
electrons shared with all surrounding nucleus
Properties Associated with Metallic
Bond
 Electrons are bonded to a large network, not to a single
atom.
 More freedom of movement
 Easily reorganize
 (bendable, stretchable etc.)
 Easy, free movement of electrons makes conductivity
easy to understand
Comparing and Contrasting Ionic, Covalent and Metallic
Bonds
Ionic Bonds
Covalent Bonds
Metallic Bonds
Transferred
Shared, evenly or
unevenly
Electron Sea
Metal to Nonmetal
Nonmetal to
Nonmetal
Metal to Metal
Electronegativity
Differences
Differences
greater than 2.0
Differences betwn
0.0 and 2.0
N/A
Make Compounds
Yes, by attraction
of opposite
charged ions
Molecules, or
molecular elements
No
State (STP)
Crystalline solid
Liquid, gas, or solid
Malleable and
ductile solid
Melting Pt
high
low
low
Liquid and
aqueous state, yes
no
yes
high
low
no
Electrons
Properties
Bond
Conductivity
Water
solubility