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Assignment #1: Review (Part 1)
1. Sodium and chlorine become chemically bonded together to form sodium chloride. Sodium chloride must
be
(1) a compound, which is a pure substance (3) a compound, which is not a pure substance
(2) a mixture, which is a pure substance
(4) a mixture, which is not a pure substance
2. Which substance cannot be decomposed by a chemical change?
(1) AlCl3
(2) H2O
(3) HI
(4) Cu
3. Two solid samples each contain sulfur, oxygen, and sodium, only. These samples have the same color,
melting point, density, and reaction with an aqueous barium chloride solution. It can be concluded that the
two samples are the same
(1) compound
(2) element
(3) mixture
(4) solution
4. Tetrachloromethane, CCl4, is classified as a
(1) compound because the atoms of the elements are combined in a fixed proportion
(2) compound because the atoms of the elements are combined in a proportion that varies
(3) mixture because the atoms of the elements are combined in a fixed proportion
(4) mixture because the atoms of the elements are combined in a proportion that varies
 Questions 1-4 are reviewed in section 1 of the review assignment.
5. Which equation represents an exothermic reaction?
(1) H2O(s) + heat → H2O(l)
(2) 2H2(g) + O2(g) → 2H2O(g) + heat
(3) H2(g) + I2(g) + heat → 2HI(g)
(4) N2(g) + 2O2(g) + heat → 2NO2(g)
 Question 5 is reviewed in section 2 of the review assignment.
6. How many of each type of atom are given in 5 Ca(NO3)2?
Score = ______/1
5 Ca atoms 10 N atoms 30 O atoms
 Question 6 is reviewed in section 3 of the review assignment.
7. Which equation represents a physical change?
(1) H2(g) + I2(g) + heat → 2HI(g)
(2) 2H2(g) + O2(g) → 2H2O(g) + heat
Score = ______/4
Score = ______/1
(3) H2O(s) + heat → H2O(l)
(4) N2(g) + 2O2(g) + heat → 2NO2(g)
8. Sodium metal will react with chlorine gas in a spectacular, sparking reaction that produces the compound
NaCl. This interaction can be described by the equation below:
Equation:
Particle Picture:
2 Na (s) + Cl2 (g)  2 NaCl (s)
+

Explain why the above equation and particle picture demonstrate the law of conservation of matter.
This shows the Conservation of Matter because particles have not been created or destroyed, they stay the
same. There are 2 black (Na) and 2 white (Cl) particles on the reactant (left) side and 2 black (Na) and 2 white
(Cl) particles on the product (right) side.
9. Which is a possible description of the sample of matter below?
(1) 4 N2 and 3 H2
(2) 4 NO and 3 H2O
(3) 4 N2 and 3 NO2
(4) 4 N2 and 3 H2O
Use the particle diagrams below to answer Questions #11– 12:
(1)
(2)
(3)
1
10. Which particle diagram contains only a pure compound?
3
11. Which particle diagram contains a diatomic element?
2
12. Which diagram contains a mixture of two monatomic elements?
 Question 9-12 is reviewed in section 5of the review assignment.
(4)
Score = ______/4
Assignment #1: Part 2-Improving Your Understanding of Unit 1
Review Section #1-The Matter Flow Chart:
MATTER
SUBSTANCE
MIXTURES
COMPOUND
ELEMENT
HOMOGENEOUS
HETEROGENEOUS
**This Unit will focus on COMPOUNDS, how and why they are formed, how they are
named, how to write their formulas and how they react with each other**
a) Elements are the basic building block of matter that cannot be broken down by a chemical reaction.
(can / cannot)
b) A compound is formed when atoms
combine
in a
(combine / come apart)
fixed
ratio.
(fixed / variable)
An example: Formation of water from hydrogen and oxygen gases.
c) Use Table S to compare some basic physical properties of O2, H2 and H2O and answer the question:
Substance
Melting Point (K)
Boiling Point (K)
Density (g/cm3)
O2
55
90
0.001429
H2
14
20
0.00009
H2O
273
373
1
Based on this chart, compounds have
unique
chemical and physical properties.
(the same / unique)
Note:
H2O has its own unique properties (not the same as O2 or H2)!!!
This indicates that a compound is a new substance with its own properties!!!
Review Section #2- Exothermic and Endothermic Changes
Exothermic means heat is
released
by the system and is written as a
(released / absorbed)
Endothermic means heat is
absorbed
product
.
(reactant / product)
by the system and is written as a
(released / absorbed)
reactant .
(reactant / product)
Identify each of these as either exothermic or endothermic.
a.) Heating a chunk of iron until it melts.
b.) 2 Na + Cl2  2 NaCl + Light
c.) CaCO3(s) + heat  CaO(s) + CO2 (g)
d.) 2MgO + 288 kjoules  Mg + O2
ENDO
EXO
ENDO
ENDO
Review Section #3-Working with Formulas (Topics 3, 4, & 5)
 A subscript in a formula tells how many atoms of each kind are in one unit of that compound.
(no subscript = 1 atom)
Example:
Na2SO4
2 Na, 1 S, and 4 O in each unit
 Subscripts are distributed when they are behind a parentheses, but only to the atoms in the parentheses.
Example:
Ca(NO3)2
1 Ca, 2 N, and 6 O in each unit
 A number before the formula is called a coefficient. This number tells how many molecules we are
dealing with or describing. That number is ALWAYS distributed to all atoms.
Example:
7 Na2CO3
7 molecules of Na2CO3 containing a total of 14 Na, 7 C and 21 O
How many molecules are shown in the following?
a) 2 H2O = 2
b) CO2 = 1
c) 25 NaCl = 25
d) a mixture of 2 H2O and 3 NaCl = 5
How many atoms of each type are shown by:
a) NH4Cl
#N : _1_
#H : _4__
# Cl : _1__
b) (NH4)3PO4
#N : _3_
#H : _12_
# P : _1__
c) 7 (NH4)2S
#N : _14_
#H : _56__
# S : _7__
# O : _4_
Review Section #4- Chemical Equations: (Topic 6)
Common Notations in Chemical Equations
Symbol
Meaning
+

(s)
(l)
(g)
(aq)
Separates two reactants or two products
Separates reactants from products; read or written as yields or
produces
Identifies the substance as a solid
Identifies the substance as a liquid
Identifies the substance as a gas
Identifies the substance as being in aqueous solution
(i.e. –dissolved in water)
Review #4- Chemical Equations (cont’d): (Topic 6)
Label the chemical equation using PRODUCT, REACTANTS, SUBSCRIPT, COEFFICIENT, and YIELDS.
Yields
Coefficient
Product
2 Mg +
O2 ⟶
2 MgO
Reactants
Subscript
When substances interact, they often end up combining their set of atoms into new and different substances.
For example:
_ 1__ Ca + _ 2__ H2O  _ 1__ Ca(OH)2 + _ 1__ H2
a) The chemical equation describes a chemical change or reaction.
We know this shows a chemical change, because…
different substances (different compositions) are being formed!
b) Reactants (the substances that react) are on the ____left_____side of the equation.
c) Products (the substances that are produced) are on the ____right____side of the equation.
d) Reactants and products are separated by an ___arrow____. It shows the direction of the chemical change.
e) In a chemical change, the arrangement of the atoms has changed in going from reactants to products, while
in a physical change it is the distance between the particles that changes.
f) Remember that the number and kind of atoms on the left side of the equation must equal the number
and kind of atoms on the right side of the equation. The process of adjusting the coefficients in the
chemical equation to show equality of atoms or matter on either side of a chemical equation is called
balancing the equation. Doing so illustrates the principle of the Conservation of Matter.
g) Given the balanced equation representing a reaction:
CaO(s) + CO2(g)  CaCO3(s) + heat
X
+ 88g = 200g
What is the total mass of CaO(s) that reacts completely with 88 grams of CO2(g) to produce 200 grams of
CaCO3(s)?
(1) 56 g
(2) 88 g
(3) 112 g
(4) 288 g
h) Which equation shows conservation of atoms?
(1) H2 + O2 → H2O
(2) H2 + O2 → 2H2O
(3) 2H2 + O2 → 2H2O
(4) 2H2 + 2O2 → 2H2O
Review Section #5-Particle Drawings: (Topic 6)
Use particle diagrams to illustrate each of the reactions from the previous page.
Make sure to make a key!
a.)
__2___NaCl
+
__1___Br2 
__2____NaBr
+ __1___Cl2
Draw your reactants in the “Reactants” box and your products in the “Products” box.
Key:
=Na
=Cl
= Br
Reactants Box
(“BEFORE” reaction)
b.)
__2___Al
+
__3___ Cl2

Products Box
(“AFTER” reaction)
__2___AlCl3
Draw your reactants in the “Reactants” box and your products in the “Products” box.
Key:
=Al
= Cl
Reactants Box
(“BEFORE” reaction)
c.)
__4___Al(s)
+
__3___ O2 (g)
Products Box
(“AFTER” reaction)

__2___Al2O3 (s)
Draw your reactants in the “Reactants” box and your products in the “Products” box.
Key:
=Al
=O
Reactants Box
(“BEFORE” reaction)
REVIEW: ISOTOPES
Products Box
(“AFTER” reaction)
Isotopes = atoms of the same element (same number of protons) with a different number of
neutrons
Ex. 12C and 14C, Magnesium-24 and Magneisum-25
Mass Number = The total number of protons and neutrons in an atom. It is a whole number.
Atomic Mass = The weighted average of the naturally occurring isotopes of an element. It is a
decimal found on the Periodic Table.
Helium has two Isotopes:
He-3 (0.0001% abundance)
He-4 (99.999% abundance)
p+ = 2
p+ = 2
n0 = 1
n0 = 2
e- = 2
e- = 2
Which is most abundant?
He-4
What number is the atomic mass going to be closest to?
4
Base your answers to questions 1 through 3 on the data table below, which shows three isotopes of neon.
1. In terms of subatomic particles, state one difference between these three isotopes of neon.
The isotopes of neon have a different number of neutrons.
2. In terms of subatomic particles, state one similarity between these three isotopes of neon.
The isotopes of neon have the same number of protons.
3. Based on natural abundances, the average atomic mass of neon is closest to which whole number?
20
Honors only
Answers to average atomic mass.
1.) 178.55 g
2.) 126.86 g
3.) 197.50 g
4.) 6.94 g
5.) 1.012 g
6.) 24.31 g
7.) 192.95
8.) Lithium – 7 is closest to 6.94 thus the most abundant isotope.
Ion Mini Lesson Notes
Ions = charged particles
14
ex.
N
ex. cation = (+) positively charged particle
anion = (-) negatively charged particle
Lewis
p+ = 7
n0 = 7
e- = 7
electron configuration = 2-5
 atom
 neutral
N -3
14
ex.
 anion
 lost 3 e-
p+ = 7
n0 = 7
e- = 10
electron configuration = 2-8
Bohr
p=7
n=7
Lewis
2
5
Bohr
-3
[ N ]-3
p=7
n=7
2
8
(Larger radius)
23
ex.
Na
Lewis
p+ = 11
n0 = 12
e- = 11
electron configuration = 2-8-1
 atom
 neutral
23
ex.
Na +1
p+ = 11
n0 = 12
e- = 10
electron configuration = 2-8
 cation
 lost 1 e-
Bohr
p = 11
n = 12
Lewis
[Na]+1
2
Bohr
p = 11
n = 12
8
+1
2
8
(Smaller radius)
Question - In terms of subatomic particles, what is the difference between an atom and an ion?
_An atom has an equal number of protons and electrons while an ion’s electrons will vary.
Assignment #3: Ions Tutorial
Go to www.chemthink.com. Then…
1
A.) Use your username and password from earlier in this unit.
B.) Once you have logged on to the website, select “The Atom”  Select “Ions” Select “Tutorial”
Answer the following questions as you go through the tutorial:
1) What are ions?
Ions are individual atoms or groups of atoms that have a charge. (This charge comes from an
imbalance, or deviation from the normal number of positive protons and negative electrons.)
2) What happens to a lithium atom’s charge when it loses an electron?
When Li atom loses and electron, the total charge changes from zero to positive 1 (0  +1).
3) In terms of subatomic particles, why does the lithium ion have this charge?
There are still three positive protons in the nucleus, creating a net charge of +3, but there are
now only two negative electrons, creating a net charge of -2. So overall, the ion has one more
positive than it does negative, and is left with a +1 charge.
4) What is the symbol for the lithium ion?
Li+1
5) What happens to oxygen’s charge when it gains two electrons?
It changes from zero to -2.
6) In terms of subatomic particles, why does the oxygen ion have this charge?
It has eight positive protons, creating a net charge of +8, but now there are ten negative
electrons, with a net charge of -10. So overall, the ion has two more negatives than it does
positives, and is left with a -2 charge.
7) What is the symbol for the oxygen ion?
O-2
8) Do an ion and a neutral atom react the same chemically?
No! An ion has a different number of electrons and the electrons determine the chemical
properties or reactivity of an element, so it will behave much different than the atom it came
from.
9) MULTIPLE CHOICE: What particle(s) is (are) gained or lost in forming ions?
(1) protons
10)
(2) neutrons
(3) electrons
(4) either protons or electrons
Negative
ions can only be formed by gaining electrons &
can only be formed by losing electrons.
Assignment #4: Atoms & Their Ions
positive
ions
Directions: Read the following information (gray boxes) and answer any questions that follow.
Why Learn About This?
The universe tends toward stability and thus low energy. At the atomic level, atoms achieve stability
by having an energetically stable number of electrons. Some atoms achieve this by attracting or
pulling electrons in from another atom. Those that lose the electron achieve stability as a result as
well. This competition for electrons is the essence of chemical interaction.
Learning Objectives:
 Associate metallic behavior with the likelihood to lose electrons, thus forming a positively
charged ion.
 Associate non-metallic behavior with the likelihood to gain electrons, thus forming a
negatively charged ion.
Success Criteria:
 Evaluate the electronic structure of an atom to predict metallic or non-metallic behavior.
 Assign electrical charge to an ion based on the number of protons and electrons it contains
after forming an ion.
Prerequisites:
 Atomic Structure
 Electron Configuration, valence electrons
 Electrical Charge
Resources:
 Periodic Table
Information:
Depending on electron structure, some atoms will lose or gain electrons in order to achieve a
complete valence electron shell which is more energetically stable. Atomic stability is the driving
force behind chemical interactions.
Model: Structure of Atoms and Their Ions
Number of Protons
(p+)
Number of Electrons
(e-)
Electron
Configuration
Number of Valence
Electrons (val. e-)
Mg ATOM
Mg ION
S ATOM
S ION
12
12
16
16
12
10
16
18
2-8-2
2-8
2-8-6
2-8-8
2
8
6
8
[Mg ]
Metal
Metal
Nonmetal
Nonmetal
N/A*
N/A*
Lost
2
N/A*
N/A*
Gained
2
0
+2
0
-2
N/A*
Cation
N/A*
Anion
N/A*
Neon
(Ne)
N/A*
Argon
(Ar)
Lewis Electron Dot
Diagram
Mg
Metal, Nonmetal, or
Metalloid?
e- lost or gained?
Number of elost/gained
Electrical Charge
-2
S
+2
S
Type of Ion
Becomes like which
noble gas (in terms
of electron
configuration)?
*The atom is included in the model for comparison purposes; therefore some of the information does not apply.
Additional Info:
Chemists have noticed that ALL elements will interact chemically with other substances, EXCEPT the
Noble Gases. Since we know that chemical interaction is a result of the number of valence electrons,
must be there is something unique, or as a chemist would say, “stable” about being a noble gas. With the
exception of helium (He) all noble gases have 8 valence electrons. We also know that all elements end
up with 8 valence electrons when they interact with other substances.
Key Questions:
(1) A) Compare the number of protons and electrons in a metallic atom.
They are the same.
B) Compare the number of protons and electrons in a metallic ion.
There are more protons than electrons.
(2) Would it be easier to gain electrons or lose electrons for a metallic atom? Explain in terms of valence
electrons.
Metallic atoms usually have only 1 or 2 valence electrons (some have 3 or 4). Since every element
reacts in order to achieve 8 valence electrons, it will be easier for metals to lose valence electrons
to achieve this stable arrangement.
(3) Why does a metallic element form a cation? Explain in terms of number of charged subatomic
particles.
A cation is a positively charged ion. Metallic elements lose their valence electrons, which means
that since they started with equal numbers of protons and electrons, and they just lost some, that
they now have more positively charged subatomic particles than they do negatively charged
subatomic particles, creating an overall positive charge on the ion.
(4) A) Compare the number of protons and electrons in a nonmetallic atom.
They are the same.
B) Compare the number of protons and electrons in a nonmetallic ion.
There are more electrons than protons.
(5) Would it be easier to gain electrons or lose electrons for a nonmetallic atom? Explain in terms of
valence electrons.
Nonmetallic atoms have 4, 5, 6, or 7 valence electrons. Since every element reacts in order to
achieve 8 valence electrons, it will be easier for non-metals to gain valence electrons to achieve
this stable arrangement.
(6) Why does a nonmetallic element form an anion? Explain in terms of number of charged subatomic
particles.
An anion is a negatively charged ion. Nonmetallic atoms gain valence electrons, which means
that since they started with equal numbers of protons and electrons, and they just gained some,
that they now have more negatively charged subatomic particles than they do positively charged
subatomic particles, creating an overall negative charge on the ion.
(7) When atoms form ions, the electron configuration of the ion resembles the electron configuration of a
CERTAIN noble gas. Which Noble Gas will each of these atom’s ions be like?
a. Sr
Kr
c. Br Kr
e. Al Ne
g. Li He
b. O
Ne
d. Si Ar
f. P
h. Cs Xe
Putting It All Together:
Ar
1.) In terms of subatomic particles, what is the difference between an atom and an ion?
An atom has the same number of protons and electrons, ions have different numbers of
protons and electrons.
2.) In terms of subatomic particles, what is the difference between an ion and an isotope of an
atom?
An ion has a different number of electrons than the original atom, whereas an isotope
has a different number of neutrons.
3.) How can you determine the electrical charge on an ion?
Find the number of protons and the number of electrons and see which one there is
more of. If there are more electrons, the ion has a negative charge, if there are more
protons, the ion has a positive charge. Either of these charges is equal to the number of
particles lost or gained.
4.) When an atom becomes an ion, does the element’s nucleus change?
No!
Assignment #5: Atoms vs. Ions Chart
Fill in the Table below:
Symbol
Mass
Number
Charge
# of
protons
# of
neutrons
# of
electrons
# of
valence
electrons
40
2+
20
20
18
8
20
0
10
10
10
8
Na
23
1+
11
12
10
8
Li1
6
1+
3
3
2
2
12
6
C
12
0
6
6
6
4
C
16
8
O 2
16
2-
8
8
10
8
[ O ]2-
26
2+
12
14
10
8
[Mg] 2+
68
1-
35
33
36
8
[
40
20
Ca 2
20
10
23
11
6
3
26
12
Ne
1
Mg
68
35
2
Br1
Assignment #6: Depicting and Describing Ions
Lewis Dot
structure
[Ca] 2+
Ne
[Na] 1+
[Li] 1+
Br
]
1-
Directions: Do the first 2 elements. Do the rest if you need more practice.
Element
Metal or
Nonmetal?
Lewis Dot
Structure
as an ATOM
Gain or
lose
electrons
to become
stable?
How
many
are
gained
or lost?
Lewis Dot Structure
of Stable ION
Becomes
like which
noble gas?
Gain
1
Lose
1
[ Li ]+1
He
Lose
3
[ Al ]+3
Ne
Gain
2
Lose
2
Gain
1
Fluorine
Nonmetal
Lithium
Metal
F
Li
Aluminum
Metal
Al
[ F ]
-1
Ne
Sulfur
Nonmetal
Radium
Iodine
Metal
Nonmetal
S
Ra
I
When ions are formed the charge is called the
oxidation
[ S ]
[ Ra ]+2
[ I ]
state
(Hint: See Periodic Table KEY)
Assignment #7: Atoms vs. Ions Guided Practice
-2
-1
Ar
Rn
Xe
of the element.
For each of the following:
a. Tell whether the given atom more likely to lose or gain electrons
b. What would the ionic Lewis structure look like?
c. Tell whether the radius increases or decreases as the atom becomes an ion.
Ex) Phosphorus atom (becomes phosphide ion)
a.
Gain
b.
[ P ] -3
c.
Increase
8 valence electron dots should be placed around the P symbol
1.) Bromine atom (becomes bromide ion)
a.
Gain
b.
[ Br ] -1 8 valence electron dots should be placed around the P symbol
c.
Increase
2.) Magnesium atom (becomes magnesium ion)
a.
Lose
b.
c.
[ Mg ] + 2
Decrease
3.) Sulfur atom (becomes sulfide ion)
a.
Gain
b.
[ S ] -2
c.
Increase
8 valence electron dots should be placed around the P symbol
4.) Lithium atom (becomes lithium ion)
a.
Lose
b.
[ Li ] + 1
c.
Decrease
Ions Drill and Regents Questions
1. Find the number of p+, no, and e- in the following species.
a.) 48Ti2+
b.) 39K+
p+:
22
p+:
o
n:
26
no:
e-:
20
e-:
19
20
18
2. Which statement best describes electrons?
(1) They are positive subatomic particles and are found in the nucleus.
(2) They are positive subatomic particles and are found surrounding the nucleus.
(3) They are negative subatomic particles and are found in the nucleus.
(4) They are negative subatomic particles and are found surrounding the nucleus.
3. The following equation represents the formation of a
(1) fluoride ion, which is smaller in radius than a fluorine atom
(2) fluorine atom, which is smaller in radius than a fluoride ion
(3) fluoride ion, which is larger in radius than a fluorine atom
(4) fluorine atom, which is larger is radius than a fluoride ion
4. As an atom becomes an ion, its mass number
(1) decreases
(2) increases
5. What is the total number of electrons in a Cu1+ ion?
(1) 28
(2) 29
(3) 30
(3) remains the same
(4) 36
6. After a neutral sulfur atom gains two electrons, what is the resulting charge of the ion?
(1) 2+
(2) 2(3) 3+
(4) 3-
7. Which statement is true about the charges assigned to an electron and a proton?
(1) Both an electron and a proton are positive.
(2) An electron is negative and a proton is positive.
(3) An electron is positive and a proton is negative.
(4) Both an electron and a proton are negative.
8. What is the total number of electrons in a Cr3+ ion?
(1) 18
(2) 21
(3) 24
(4) 27
9. How many electrons are contained in an Au3+ ion?
(1) 76
(2) 79
(3) 82
(4) 197
c.) 14N3p+: 7
no: 7
e-: 10
10. Which changes occur as a cadmium atom, Cd, becomes a cadmium ion, Cd2+?
(1) The Cd atom gains two electrons and its radius decreases.
(2) The Cd atom gains two electrons and its radius increases.
(3) The Cd atom loses two electrons and its radius decreases.
(4) The Cd atom loses two electrons and its radius increases.
11. What is the total number of electrons in a S2– ion?
(1) 10
(2) 14
(3) 16
(4) 18
12. Compared to a calcium atom, the calcium ion Ca2+ has
(1) more protons
(3) more electrons
(2) fewer protons
(4) fewer electrons
13. An atom of an element has a total of 12 electrons. An ion of the same element has a total of 10 electrons.
Which statement describes the charge and radius of the ion?
(1) The ion is positively charged and its radius is smaller than the radius of the atom.
(2) The ion is positively charged and its radius is larger than the radius of the atom.
(3) The ion is negatively charged and its radius is smaller than the radius of the atom.
(4) The ion is negatively charged and its radius is larger than the radius of the atom.
14. An atom of an element forms a 2+ ion. In which group on the Periodic Table could this element be
located?
(1) 1
(3) 13
(2) 2
(4) 17
15. What can be concluded if an ion of an element is smaller than an atom of the same element?
(1) The ion is negatively charged because it has fewer electrons than the atom.
(2) The ion is negatively charged because it has more electrons than the atom.
(3) The ion is positively charged because it has fewer electrons than the atom.
(4) The ion is positively charged because it has more electrons than the atom.
16. What is the net charge on an ion that has 9 protons, 11 neutrons, and 10 electrons?
(1) 1+
(3) 1–
(2) 2+
(4) 2–
17. What is the overall charge of an ion that has 12 protons, 10 electrons, and 14 neutrons?
(1) 2(3) 4(2) 2+
(4) 4+
Assignment #9: Ionic Formulas ChemThink
Go to www.chemthink.com. Then…
A.) Use your username and password from the last unit.
B.) Once you have logged on to the website, select “Ionic Bonding”  Select “Ionic
Formulas” Select “Question Set”
Complete the questions asked. You must get 10 questions correct before missing 3 to get
credit. If you do not succeed the first or second time, consider viewing the tutorial before
trying the questions again.
MINI-LESSON Lewis Dot Diagrams for Ionic Compounds
1. One atom that has the ability to lose electrons gets near one that needs to take electrons and the transfer occurs.
a. Draw a Lewis Structure for a potassium atom:
b. Draw a Lewis Structure for a bromine atom:
K
Br
c. Draw a Lewis Structure for a potassium ion:
d. Draw a Lewis Structure for a bromide ion:
[ Br ]-1
[ K ]+1
e. Draw a Lewis Structure for the compound potassium bromide:
[ K ]+1[ Br ]-1
2. The electron from K was transferred to the Br and now they are both stable and “bond” as a compound.
The chemical formula of the compound will be: _____KBr______
3. Lithium and sulfur form lithium sulfide
[Li ]+1[
S ]-2[Li ]+1
Li2S
Practice time! NOTE: Sometimes the number of electrons the metal atom needs to lose is not the same as the
number of electrons the non-metal atom needs to gain. The ratio is not always 1:1.
4. Aluminum and nitrogen form aluminum nitride
[Al ]+3 [ N ]-3
AlN
5. Potassium and sulfur form potassium sulfide
[ K ]+1 [ S ]-2[ K ]+1
K2S
6. Lithium and nitrogen form lithium nitride
[ Li ]+1 [ N ]-3[ Li ]+1
Li3N
[ Li ]+1
**Note that the positive ions and negative ions attract (are next to one another), while like charges repel.
7. Calcium and nitrogen form calcium nitride
[ Ca]+2 [ N ]-3[ Ca ]+2 [ N ]-3[ Ca ]+2
Ca3N2
What is in common between each of the above formulas/names?
1.) Valence electrons are
transferred
.
(shared / transferred)
2.) Bonding occurs between
metals
and
(metals/nonmetals)
3.) The combination of ion charges are equal to
nonmetals
types of elements.
(metals/nonmetals)
ZERO
.
(insert #)
4.) Naming involves metal listed first and
nonmetal
(metal/nonmetal)
5.) Chemical formula is written with
second, with “-ide” ending.
(metal/nonmetal)
positive
ion (or cation) first.
(positive / negative)
NEW-PHYSICAL PROPERTIES OF IONIC COMPOUNDS:
Think about Table Salt=–Sodium Chloride (NaCl)
1.)
High
Info: Mp= 801°C
Bp= 1413°C
melting/boiling points
(High / low)
2.)
Hard
(this has to do with crystalline structure)
(Hard / soft)
3.)
Strong
conductor in liquid or dissolved (aqueous) form.
(Strong / weak)
4.)
Is NOT
conductor in solid form.
(Is / Is NOT)
Assignment #10: Using Lewis Dot Diagrams to Write Formulas
Draw the Lewis Dot Structure for the following ionic compounds formed from a cation and an anion to
give a neutral formula. Remember: total (+) charges and total (-) charges are equal! Then write the
chemical formula and number of ions present in the compound.
Chemical Name
1.) potassium chloride
2.) sodium sulfide
3.) aluminum bromide
4.) magnesium phosphide
Lewis Dot Diagram
Chemical
Formula
Number of
Ions Present
KCl
2
Na2S
3
[ Br ]-1[Al]+3[ Br ]-1
[ Br ]-1
AlBr3
4
[Mg]+2 [ P ]-3[Mg]+2
[ P ]-3[ Mg ]+2
Mg3P2
5
[ K ]+1[ Cl ]-1
[Na]+1 [ S ]-2[Na]+1
See Review Materials Assignment #32 for additional practice if needed!
Assignment #11: Ionic Formula Details
Summary: So…IONIC BONDING is…
1.) A
transferring
of valence e- between
metals
(Sharing/transferring)
2.) Made of
and _____nonmetals_____.
(metals/nonmetals)
positive
and
(positive / negative)
negative
(metals/nonmetals)
, but are electrically
neutral
.
(positive / negative)
3.) Bond occurs because of VERY STRONG
attraction
between opposite charges.
(attraction / repulsion)
Naming with Transition Metals:
Recap: Metallic elements in Groups ___3____-___12____ are TRANSITION METALS!
Because there is typically more than one possible oxidation state we can use, we must SPECIFY which
transition metal ion we are using! We do this by using
Roman numerals
.
Roman Numerals: 1= __I__, 2= __ II___, 3= __ III ___, 4=__ IV ___, 5= ___ V____
For Example:
Compound:
Ox. # of Co:
CoCl2
CoCl3
+2
+3
Ox. # of Cl:
-1
-1
Cobalt (II) chloride
Cobalt (III) chloride
Name:
*Remember- The overall charges have to add up to zero!!!!
Name these:
Compound:
CuCl
CuCl2
Ox. # of Cu:
+1
+2
Ox. # of Cl:
-1
-1
Copper (I) chloride
Copper (II) chloride
Compound:
Fe2O3
FeO
Ox. # of Fe:
+3
+2
Ox. # of O:
-2
-2
iron (III) oxide
iron (II) oxide
Name:
Determine the formula:
Name:
Assignment #12: Ionic Compounds – Naming & Writing Formulas
When you are asked to name a compound, it will often tell you to name it according to IUPAC guidelines. This is
just something the Regents throws in to confuse you… don’t let it! IUPAC just stands for International Union of
Pure and Applied Chemistry, and basically, these are the people who make all the rules for naming. So, if a
question asks you to name something according to IUPAC, just name it the same way you always do!
Hint: Use your periodic table in your reference table to see if the metal has more than one oxidation number.
Part A: Ionic Compound Formulas  Name
1. CaI2
calcium iodide
2. LiBr
lithium bromide
3. CsF
cesium fluoride
4. AlBr3
aluminum bromide
5. RbBr
rubidium bromide
6. NiBr3
nickel (III) bromide
7. AuCl3
gold (III) chloride
8. BaF2
barium fluoride
9. Fe2O3
iron (III) oxide
10. PbS
lead (II) sulfide
Part B: Ionic Compound Formulas from a Name
1. Potassium Iodide
KI
2. Barium Chloride
BaCl2
3. Sodium Phosphide
Na3P
4. Tin (IV) Oxide
SnO2 (don’t forget to reduce! Sn2O4 has a simpler ratio!)
5 Gold (III) Bromide
AuBr3
6. Copper (II) Iodide
CuI2
7. Silver Oxide
Ag2O
8. Chromium (III) Sulfide
Cr2S3
Assignment #13: What Does the Charge Tell Us?
Ionic Bonding Regents Questions
1.) A compound is made up of iron and oxygen, only. The ratio of iron ions to oxide ions is 2:3 in this
compound. The IUPAC name for this compound is
(1) triiron dioxide
(3) iron(III) oxide
(2) iron(II) oxide
(4) iron trioxide
2.) What is the IUPAC name for the compound FeS?
(1) iron(II) sulfate
(3) iron(II) sulfide
(2) iron(III) sulfate
(4) iron(III) sulfide
3.) Which type of bond results when one or more valence electrons are transferred from one atom to another?
(1) a hydrogen bond
(3) a nonpolar covalent bond
(2) an ionic bond
(4) a polar covalent bond
4.) A metal, M, forms an oxide compound with the general formula M2O. In which group on the Periodic
Table could metal M be found?
(1) Group 1
(3) Group 16
(2) Group 2
(4) Group 17
5.) Which formula correctly represents the composition of iron (III) oxide?
(1) FeO3
(3) Fe3O
(2) Fe2O3
(4) Fe3O2
6.) In the formula X2O5, the symbol X could represent an element in Group
(1) 1
(3) 15
(2) 2
(4) 18
7.) What is the chemical formula for iron(III) oxide?
(1) FeO
(3) Fe3O
(2) Fe2O3
(4) Fe3O2
8.) Solid MgCl2 and liquid MgCl2 have different
(1) empirical formulas
(3) ion ratios
(2) formula masses
(4) physical properties
9.) Which statement describes the composition of potassium chlorate, KClO3?
(1) The proportion by mass of elements combined in potassium chlorate is fixed.
(2) The proportion by mass of elements combined in potassium chlorate varies.
(3) Potassium chlorate is composed of four elements.
(4) Potassium chlorate is composed of five elements.
10.) Magnesium and calcium have similar chemical properties because an atom of each element has the same
total number of
(1) electron shells
(3) neutrons
(2) valence electrons
(4) protons
11.) Given the balanced equation representing a reaction:
Which statement explains why the energy term is written to the right of the arrow?
(1) The compound CuS is composed of two metals.
(2) The compound CuS is composed of two nonmetals.
(3) Energy is absorbed as the bonds in CuS form.
(4) Energy is released as the bonds in CuS form.
12.) When sodium and fluorine combine to produce the compound NaF, the ions formed have the same
electron configuration as atoms of
(1) argon, only
(3) both argon and neon
(2) neon, only
(4) neither argon nor neon
13.) In which compound is the ratio of metal ions to nonmetal ions 1 to 2?
(1) calcium bromide
(3) calcium phosphide
(2) calcium oxide
(4) calcium sulfide
14.) Which group on the Periodic Table of the Elements contains elements that react with oxygen to form
compounds with the general formula X2O?
(1) Group 1
(3) Group 14
(2) Group 2
(4) Group 18
Base you answer to question 15 on the information below.
One process used to manufacture sulfuric acid is called the contact process. One step in this process, the reaction
between sulfur dioxide and oxygen, is represented by the forward reaction in the system at equilibrium shown
below.
2SO2(g) + O2(g)  2SO3(g) + 394 kJ
A mixture of platinum and vanadium (V) oxide may be used as a catalyst for this reaction. The sulfur trioxide
produced is then used to make sulfuric acid.
15.) Write the chemical formula for vanadium (V) oxide.
V2O5
Base you answer to question 16 on the information below.
When a person perspires (sweats), the body loses many sodium ions and potassium ions. The evaporation of
sweat cools the skin. After a strenuous workout, people often quench their thirst with sports drinks that
contain NaCl and KCl. A single 250.-gram serving of one sports drink contains 0.055 gram of sodium ions.
16.) Draw a Lewis electron-dot diagram for one of the positive ions lost by the body as a person perspires.
Assignment #14: Covalent Bonding Tutorial
Go to www.chemthink.com. Then…
A.) Use your username and password from earlier in this unit.
B.) Once you have logged on to the website, select “Covalent Bonding”  Select “Covalent
Bonding” Select “Tutorial”
Answer the following questions as you go through the tutorial. You will only go through slides 1
through 24 (of 35 slides).
1.) What is the definition of a covalent bond?
A bond that forms when atoms are sharing electrons.
2.) When you move 2 atoms close together…
a.) How is movement of electrons different when atoms are close?
When the atoms are apart, the electrons travel all around the atom. When the
atoms are close together, the electrons move just between the two atoms.
b.) What happens if you try to move the atoms very close to each other?
You can’t; the atoms have the electrons between them.
3.) Explain what a covalent bond really is. Make sure to mention the electrons of one atom
and the nucleus of a different atom in your response.
A covalent bond is really two atoms fighting over electrons. One atom’s nucleus
attracts its own electrons as well as the electrons from the other atom. The other
nucleus does the same but neither atom actually takes the electrons from the other.
4.) MULTIPLE CHOICE: Which type of elements participate in covalent bonding?
(1) metals
(2) nonmetals
(3) noble gases
5.) Explain your answer to Question #4.
Nonmetals want to gain electrons and in order to participate in covalent bonding the
atom must be able to attract another atom’s electron while still holding onto its own.
Metals don’t hold onto their own electrons very well, and noble gases don’t form
bonds.
6.) Is hydrogen a metal or a nonmetal?
Nonmetal
7.) When are atoms considered “most stable”?
When the atoms are arranged so that the potential energy is lowest, which occurs
when the atoms are bonded.
8.) A single solid line (a single bond) =
2
shared valence electrons
(how many?)
A double bond
=
4
shared valence electrons
(how many?)
A triple bond
=
6
shared valence electrons
(how many?)
9.) Which type of covalent bond is the strongest? Explain why you think this is so.
Triple bonds because they have the most shared valence electrons.
You do NOT need to do the “Question Set.”
Obtain a checkpoint from the teacher.
Mini-Lesson: Covalent Bonding
Covalent Bonding:
1.)
Sharing
of valence electrons between
(sharing/transferring)
nonmetals
.
(metals/nonmetals)
2.) Molecules are electrically
neutral
.
(charged / neutral)
3.) Ratio depends on how many electrons are needed complete the
octet
.
(Chemistry word for 8)
Examples: Do these with your teacher...
Note that arrows are no longer used to show bonding, because we are showing a different type of bond!
Cl and Cl=
Cl2
H and N= NH3
Cl - Cl
C and H=
CH4
C and O= CO2
PHYSICAL PROPERTIES OF COVALENT (MOLECULAR) COMPOUNDS:
Think about Chlorine (Cl2)- Info: Mp= -101°C
Bp= -34°C
1.) Most are in the
liquid
(Solid/ Liquid/ Gas)
or
gas
state.
(Solid/ Liquid/ Gas)
**Not as strong of a particle attraction as ionic substances!
2.)
Soft
(Due to weaker particle attractions!)
(Hard/Soft)
3.)
Low
melting/boiling points
(High/Low)
4.)
Does not
conduct electricity.
(Does / Does not)
Assignment #15: Lewis Dot Diagrams for Covalent Compounds
Use Lewis Dot structures to determine the chemical formula for compounds formed between the
given pairs of non-metal elements.
Non-Metals
Bonding
Lewis Dot Diagram (Rough)
Lewis Dot Diagram (Final)
Chemical
Formula
Description of
Bonds
H and H
H2
1 Single Bond
HCl
1 Single Bond
HF
1 Single Bond
F2
1 Single Bond
H2O
2 Single Bonds
H and Cl
or
H and F
F and F
H and O
or
H and N
NH3
3 Single Bonds
or
Note: You can double check your work by counting that H has 2e- and everything else has 8e-.
Sometimes atoms share more than one pair of electrons. See if you can work these out:
Non-Metals
Bonding
Lewis Dot Diagram (Rough)
One atom
of O bonds
to only one
other atom
of O
One atom
of N bonds
to only one
other atom
of N
Lewis Dot Diagram (Final)
Chemical
Formula
Description of Bonds
O2
1 Double Bond
N2
1 Triple Bond
OR
OR
Check for Understanding:
1) How many electrons are represented by each line in a Lewis Dot Diagram of a molecular
compound?
2 electrons
1) a) How many pairs of electrons are shared between elements in a single bond?
b) How many total electrons are shared between elements in a single bond?
2) a) How many pairs of electrons are shared between elements in a double bond?
b) How many total electrons are shared between elements in a double bond?
3) a) How many pairs of electrons are shared between elements in a triple bond?
b) How many total electrons are shared between elements in a triple bond?
1 pair
2 e2 pairs
4 e3 pairs
6 e-
Get teacher initials to make sure you did these right: _________
See Review Materials Assignment #33 and 34 for additional practice if needed!
Covalent Bonding Regents Questions
1.) What is the total number of pairs of electrons shared in a molecule of N2?
(1) one pair
(3) three pairs
(2) two pairs
(4) four pairs
2.) Which type of chemical bond is formed between two atoms of bromine?
(1) metallic
(3) ionic
(2) hydrogen
(4) covalent
3.) What is the total number of electrons shared in the bonds between the two carbon atoms in a molecule of
(1) 6
(2) 2
(3) 3
(4) 8
4.) Which formula represents a molecular compound?
(1) HI
(3) KCl
(2) KI
(4) LiCl
5.) Which element is composed of molecules that each contain a multiple covalent bond?
(1) chlorine
(3) hydrogen
(2) fluorine
(4) nitrogen
6.) Which formula represents a molecular compound?
(1) Kr
(3) N2O4
(2) LiOH
(4) NaI
7.) Which statement describes what occurs as two atoms of bromine combine to become a molecule of
bromine?
(1) Energy is absorbed as a bond is formed.
(2) Energy is absorbed as a bond is broken.
(3) Energy is released as a bond is formed.
(4) Energy is released as a bond is broken.
8.) Given a formula for oxygen:
What is the total number of electrons shared between the atoms represented in this formula?
(1) 1
(3) 8
(2) 2
(4) 4
9.) As a bond between a hydrogen atom and a sulfur atom is formed, electrons are
(1) shared to form an ionic bond
(2) shared to form a covalent bond
(3) transferred to form an ionic bond
(4) transferred to form a covalent bond
10.) Given the balanced equation representing a reaction:
What occurs during this change?
(1) Energy is absorbed and a bond is broken.
(2) Energy is absorbed and a bond is formed.
(3) Energy is released and a bond is broken.
(4) Energy is released and a bond is formed.
11.) Given the formula of a substance:
What is the total number of shared electrons in a molecule of this substance?
(1) 22
(3) 9
(2) 11
(4) 6
12.) A sample of CO2(s) and a sample of CO2(g) differ in their
(1) chemical compositions
(3) molecular structures
(2) empirical formulas
(4) physical properties
13.) Given the formula for hydrazine:
How many pairs of electrons are shared between the two nitrogen atoms?
(1) 1
(3) 3
(2) 2
(4) 4
Base your answers to questions 14 through 16 on the information below and on your knowledge of chemistry.
The balanced equation below represents a reaction.
O2(g) + energy  O(g) + O(g)
14.) Identify the type of chemical bond in a molecule of the reactant.
Covalent
15.) Draw a Lewis electron-dot diagram of one oxygen atom.
16.) Explain, in terms of bonds, why energy is absorbed during this reaction.
When bonds are broken energy is absorbed.
Assignment #16: Mixing it Up!
Appendix A naming covalent compounds.
1. Carbon dioxide
2. Carbon monoxide
3. Sulfur dioxide
4. Sulfur trioxide
5. Dinitrogen monoxide
6. Nitrogen monoxide
7. Dinitrogen trioxide
8. Nitrogen dioxide
9. Dinitrogen tetraoxide
10. Dinitrogen pentoxide
11. Phosphorous trichloride
12. Phosphorous pentachloride
13. Ammonia – nitrogen trihydride
14. Sulfur hexachloride
15. Diphosphorous pentaoxide
16. Carbon tetrachloride
17. Silicon dioxide
18. Carbon disulfide
19. Oxygen diflouride
20. Phosphorous tribromide
Appendix A Con’t –
1. SiF4
2. IF5
3. ClO2
4. P4S3
5. N2O5
6. NH3
7. BF3
8. CS2
9. SF6
10. As2O5
Honors only
Appendix A (cont’d)
Honors only
1. Ionic – Sodium Carbonate
2. Covalnet – diphosphorous pentoxide
3. Covalent – Nitrogen trihydride
4. Ionic – Iron (II) Sulfate
5. Covalent – Silicon Dioxide
6. Ionic – gallium chloride
7. Ionic – cobalt (II) bromide
8. Covalent – diboron tetrahydride
9. Covalent – carbon monoxide
10. Ionic – lithium sulfate
11. Covalent - N2O3
12. Ionic – NiSO4
13. Covalent – CCl4
14. Ionic – LiC2H3O2
15. Covalent – PF3
16. Ionic – V2O5
17. Ionic – Al(NO3) 3
18. Ionic – ZnS
19. Covalent – SBr6
20. Ionic – Ag3PO4
Assignment #17: Comparing and Contrasting Ionic & Covalent Bonding
1.) Fill out the following chart.
Ionic
Molecular (Covalent)
Metal & Nonmetal
2 Nonmetals
Valence electrons are
transferred
Valence electrons are
shared
Arrows
Loops
Strong attractions between
units
Weak attractions between
molecules
Hardness
(Hard or Soft)
Hard
Soft
Relative Melting/Boiling
Points
(Low or High)
High
Low
Conductive as a solid?
(Yes or No)
No
No
Conductive as a liquid?
(Yes or No)
Yes
No
Conductive in aqueous?
(Yes or No)
Yes
No
Which type(s) of elements
form these bonds?
(Metals/Nonmetals)
How are these bonds
formed?
(Valence electrons are…
Shared or Transferred?)
Bonding in Lewis
Structures Shown using…
(Arrows/Loops)
Relative Strength of
Attractions between
Units/Molecules
(Strong or Weak)
4.) Examine the data collected below. Use this information to classify compounds A-F as
either ionic or molecular compounds.
Compound
Physical
State at
Room Temp.
Conductivity
as a pure
substance
Conductivity
when
“Aqueous”
Melting Point
solid
no
yes
1049oC
solid
no
no
223oC
liquid
no
no
20oC
solid
no
yes
378oC
liquid
no
no
-94oC
solid
no
yes
650oC
A
B
C
D
E
F
a.) Categorize the compounds A – F as Ionic or Molecular.
IONIC
A
D
F
MOLECULAR (Covalent)
B
C
E
b.) Pick one of each type of compound and explain how you came to your decision.
Compounds A, D, and F can be classified as ionic because they do not conduct as a
solid but they do when dissolved in water. They also have very high melting points.
Compounds B, C, and E can be classified as molecular because they do not conduct
electricity as a solid or when dissolved in water. They also have very low melting
points.
Assignment #18: Discovering Polyatomic Ions
Polyatomic ions are ions that are made from more than one type of element. They are particles
that behave like both an ion and a molecule because they both have a charge and share
electrons.
Look at your reference table. How many valence electrons do four oxygen atoms and a sulfur
atom have between them, total?
30
Examine the Lewis diagram for the “sulfate” polyatomic ion
 Why is this a “polyatomic ion?”
It is made of more than one atom and has a charge.
 How many valence electrons are shown in the Lewis
diagram for a sulfate ion? 32
 Why does sulfate have an electrical charge of −2 ?
Two electrons were gained.
In what way is a polyatomic ion like a molecule?

Atoms are sharing electrons.
What is it that makes sulfate behave like an ion?

It is not neutral; it has a charge.
Since this has three elements, it is possible to have
both ____ionic___ and ___covalent____ bonds.
 If sulfate bonds with calcium ion, the formula is: Ca2(SO4)2, which reduces to
become Ca(SO4). So… the polyatomic ion is always the same… it never
changes.
Look at Table E in your Reference Table
Find the formula for the following polyatomic ions…
a. chromate
CrO42-
charge on a chromate ion?
2-
b. ammonium
NH4+
charge on an ammonium ion?
+1
c. carbonate
CO32-
charge on a carbonate ion?
-2
d. nitrite
NO2-
charge on a nitrite ion?
-1
Assignment #19: Polyatomic Ions and Predicting Chemical Formulas
1.) Circle the polyatomic ion in each of these formulas:
NaOH
(NH4)2S
Al2(SO4)3
CaCO3
NaC2H3O2
2.) Identify each of the following as either a compound or a polyatomic ion.
a.) NO2
compound
b.) NO21-
polyatomic ion
c.) SO32-
polyatomic ion
d.) SO3
compound
e.) HOH
compound
f.) OH1-
polyatomic ion
COMPOUNDS INVOLVING POLYATOMIC IONS CONTAIN
BOTH IONIC AND COVALENT BONDING!!
A look at Reference Table E…
Naming with Polyatomic Ions
This is simple! Take the name right off of Table E. Names of polyatomic ions never change!
Examples:
Sodium Hydroxide
Iron (II) Phosphate
______________
______________
Calcium Hydroxide
Cobalt (I) Phosphate
______________
______________
Assignment #20: Naming Compounds with Polyatomic Ions
(Do every other one, then check your answers. Go back and do the rest if you need more practice.)
1. CaCO3
calcium carbonate
2. FeSO4
iron (II) sulfate
3. Zn3(PO4)2
zinc phosphate
4. NH4NO3
ammonium nitrate
5. Al(OH)3
aluminum hydroxide
6. CuC2H3O2
copper (I) acetate
7. PbSO3
lead (II) sulfite
8. NaClO3
sodium chlorate
9. CaC2O4
calcium oxalate
10. (NH4)3PO4
ammonium phosphate
11. NaHSO4
sodium hydrogen sulfate
12. Mg(NO2)2
magnesium nitrite
13. CuSO4
copper (II) sulfate
14. NaHCO3
sodium hydrogen carbonate
15. Be(NO3)2
beryllium nitrate
16. ZnSO4
zinc sulfate
17. KMnO4
potassium permanganate
Assignment #21: Understanding Formulas
(Do every other one, then check your answers. Go back and do the rest if you need more practice.)
Compound
Formula
Ratio of ions within compound
ammonium phosphate
(NH4)3(PO4)
3:1
iron (II) oxide
FeO
1:1
iron (III) oxide
Fe2O3
2:3
calcium chloride
CaCl2
1:2
potassium nitrate
K(NO3)
1:1
magnesium hydroxide
Mg(OH)2
1:2
aluminum sulfate
Al2(SO4)3
2:3
copper (II) sulfate
Cu(SO4)
1:1
lead (IV) chromate
Pb(CrO4)2
1:2
potassium permanganate
K(MnO4)
1:1
sodium hydrogen carbonate
NaHCO3
1:1
zinc nitrate
Zn(NO3)2
1:2
aluminum sulfite
Al2(SO3)3
2:3
See Assignment #35 for additional formula/naming practice if needed
Assignment #22: Classifying Ionic and Covalent Compounds
(Do every other one, then check your answers. Go back and do the rest if you need more practice.)
1. CaCl2
ionic
11. MgO
ionic
2. CO2
covalent
12. NH4Cl
both
3. H2O
covalent
13. HCl
covalent
4. BaSO4
both
14. KI
ionic
5. K2O
ionic
15. NaOH
both
6. NaF
ionic
16. NO2
covalent
7. Na2CO3
both
17. AlPO4
both
8. CH4
covalent
18. FeCl3
ionic
9. SO3
10. LiBr
covalent
19. P2O5
ionic
Appendix B: Drawing Polatomic ions
20. N2O3
Honors only
covalent
covalent
Hydronium:
Carbonate:
Nitrate:
Peroxide:
Dichromate:
Hydrogen Carbonate:
Assignment #23: What Does a Chemical Equation Represent?
Here’s an example of a Chemical Equation…
2 H2 (g)
+
O2 (g)
2 H2O (l)

 The equation describes a chemical reaction in terms of
atoms/elements
and
heat
needed or produced.
 It shows
by
reactants
(H2 and O2) and
(reactants/products)
products
(H2O) separated
(reactants/products)
an arrow.
 It is
balanced
. The number (and type) of atoms on the left
equals
the
number (and type) of atoms on the right.
CONSERVATION OF MASS
 REMINDER:

Co-efficients
are the numbers in front of a compound. It means we
(coefficients/subscripts)
have that number of the entire formula. They tell us the ratio of
within an equation.

Subscripts
compounds
tell us that there are that number of atoms of that type in
(coefficients/subscripts)
the formula. WE CANNOT CHANGE THESE WHEN BALANCING!!!
They tell us the ratio of elements within a formula.
 Sometimes it shows energy changes. We can expand the equation! (DO IT!!)
2 H2 (g)
+
O2 (g)
 The energy content of the reactants is

2 H2O (l) + heat
equal to
(less than, greater than, equal to)
products, plus the extra energy produced.
CONSERVATION OF ENERGY
the energy content of the
What about ENERGY when bonds are formed? Bonds are formed-energy is
released
 Bonds allow atoms to go from unstable (
high
energy) to stable ( low
(low / high)
 Therefore….energy must be
released
energy).
(low / high)
! This is an example of an
(absorbed / released)
exothermic
reaction.
(endothermic / exothermic)
 Put “energy” in its proper place (either reactant or product) in the reaction below:
2 H2 +
O2

2 H2O + heat
What about ENERGY when bonds are broken? Bonds broken – energy
absorbed
 Bonds allow atoms to go from stable (
low
energy) to unstable (
(low / high)
 Therefore….energy must be
high
energy).
(low / high)
absorbed
! This is an example of an
(absorbed / released)
endothermic
reaction.
(endothermic / exothermic)
 Put “energy” in its proper place (either reactant or product) in the reaction below:
2 H2O
+ heat

2 H2
+
O2
Practice Questions:
1.) During all chemical reactions, mass and energy are
(1) absorbed
(2) conserved
(3) formed
(4) released
2.) Which equation represents an exothermic reaction?
(1) H2O(s) + 6.01 kJ → H2O(l)
(3) H2(g) + I2(g) + 53.0 kJ → 2HI(g)
(2) 2H2(g) + O2(g) → 2H2O(g) + 483.6 kJ
(4) N2(g) + 2O2(g) + 66.4 kJ → 2NO2(g)
3.) Given the balanced equation representing a reaction:
H2(g) + Cl2(g) →2HCl(g) + energy
Which statement describes the energy changes in this reaction?
(1) Energy is absorbed as bonds are formed, only.
(2) Energy is released as bonds are broken only.
(3) Energy is absorbed as bonds are broken, and energy is released as bonds are formed.
(4) Energy is absorbed as bonds are formed, and energy is released as bonds are broken.
4.) Which statement describes what occurs as two atoms of bromine combine to become a molecule of
bromine?
(1) Energy is absorbed as a bond is formed.
(3) Energy is released as a bond is formed.
(2) Energy is absorbed as a bond is broken.
(4) Energy is released as a bond is broken.
Balancing Equations Mini Lesson
When you get good at this, you should be able to:
 Explain how the “Conservation of Matter” is demonstrated by a properly balanced
equation.
 Balance equations, which includes being able to tell whether or not an equation is
balanced.
 Compare a balanced equation to the particle picture showing it.
Models:
1)
2 K
2)
3 NaBr
You Try:
+
_____ MgBr2  2 KBr
+
2 Al
_____ GaCl3  3 NaCl
+
3 Br2  2 AlBr3
+
_____ Mg
+
_____GaBr3
Level 1: Try these. Try not to rely too much on others. Try to solve on your own.
1)
_____ N2
+
2)
2 KClO3  2 KCl
3)
2 NaCl
4)
2 H2
+
+
3 H2  2 NH3
+
3 O2
_____ F2  2 NaF
_____ O2  2 H2O
+
_____ Cl2
Level 2: These are little more challenging, but YOU CAN DO IT!
5)
2 Na
+
6)
2 Ag2O  4 Ag
7)
4P
8)
____ CH4 + __2__ O2  ____ CO2 + __2__ H2O
+
2 H2O  2 NaOH
+
+
_____ H2
_____ O2
5 O2  2 P2O5
Level 3:
These are some worthy of a challenge. Ask if you need help using the
(Parentheses).
9) _____ Pb(OH)2
+
2 HCl  _____ PbCl2
+
2 H2O
10) 2__ AlBr3 + __3__ K2(SO4)  __6__ KBr + ____ Al2(SO4)3
11) ____ C3H8 + __5__ O2  __3__ CO2 + __4__ H2O
SUPER CHALLENGE
_____ C6H12O6
+
6 O2  6 H2O
+
6 CO2
JUST ABOUT AS TOUGH AS THE REGENTS WILL EVER GET:
12)
____ S8 + __12__O2  __8__ SO3
13)
__2__ H2O + ____ O2  __2__ H2O2
14)
__2__ NaBr + ____ CaF2  __2__ NaF + ____ CaBr2
CHALLENGE!  SEE IF YOU CAN GET THESE:
15)
____ FeCl3 + __3__ Na(OH)  ____ Fe(OH)3 + __3__NaCl
16)
____ H2(SO4) + __2__ Na(NO2) __2__ H(NO2) + ____ Na2(SO4)
17)
__2__ C8H18 + __25__ O2  __16__ CO2 + __18__ H2O
Assignment #25A: Balancing Equations Regents Questions
1.) Which chemical equation is correctly balanced?
(1) H2(g) + O2(g) → H2O(g)
(2) N2(g) + H2(g) → NH3(g)
(3) 2NaCl(s) → Na(s) + Cl2(g)
(4) 2KCl(s) → 2K(s) + Cl2(g)
2.) Given the unbalanced equation:
When the equation is correctly balanced using the smallest whole-number coefficients, what is the
coefficient of CO?
(1) 1
(3) 3
(2) 2
(4) 4
3.) During all chemical reactions, mass, energy, and charge are
(1) absorbed
(3) formed
(2) conserved
(4) released
4.) Given the balanced equation representing a reaction:
What is the total mass of water formed when 8 grams of hydrogen reacts completely with 64 grams of
oxygen?
(1) 18 g
(3) 56 g
(2) 36 g
(4) 72 g
5.) Which equation shows conservation of atoms?
(1) H2 + O2 → H2O
(2) H2 + O2 → 2H2O
(3) 2H2 + O2 → 2H2O
(4) 2H2 + 2O2 → 2H2O
Base your answer to question 6 on the information below.
When magnesium is ignited in air, the magnesium reacts with oxygen and nitrogen. The reaction between
magnesium and nitrogen is represented by the unbalanced equation below.
3 Mg(s) +
1
N2(g) →
1
Mg3N2(s)
6.) Balance the equation on the lines above for the reaction between magnesium and nitrogen, using the
smallest whole-number coefficients.
Base your answers to questions 7 through 8 on the information below.
A piece of magnesium ribbon is reacted with excess hydrochloric acid to produce aqueous magnesium chloride
and hydrogen gas. The volume of the dry hydrogen gas produced is 45.6 milliliters. The temperature of the gas
is 293 K, and the pressure is 99.5 kilopascals.
7.) Balance the equation in your answer booklet, using the smallest whole-number coefficients.
1 Mg
+
2 HCl
→
1 MgCl2
+
1 H2
8.) Identify the type of bond between the atoms in a molecule of the gas produced in this laboratory
investigation. The gas produced is H2 so the bond would be covalent since they are all nonmetals
See Review Materials Assignment #36 & 37 for additional practice if needed!
Mole Notes
What about those MOLES?
Because an atom is so small, and no one has been able to see or hold one atom by itself, chemists use a larger, more
manageable amount. This amount is the MOLE
Definition: A mole is defined as 6.02 x 1023 particles.
1 mole of atoms = 6.02 x 10 23 atoms.
1 mole of formula unit = 6.02 x 10 23 formula units.
1 mole of ions = 6.02 x 10 23 ions.
1 mole of molecules = 6.02 x 10 23 molecules.
The Mass of a MOLE:
The gram atomic mass⟶ the mass of one mole of atoms=the amount which contains 6.02 x 10 23 atoms=the atomic mass (g).
Example- The mass of 1 mole of Sulfur (S) = the mass which contains 6.02 x 10 23 atoms of sulfur = 32.1 grams
Practice- The mass of 1 mole of Sodium (Na) = the mass which contains 6.02 x 10 23 atoms of sodium =22.99 g.
The gram molecular mass⟶the mass of one mole of molecules = the amount which contains 6.02 x 10 23 molecules = the molecular mass in grams.
Example- The mass of 1 mole of water (H2O) = the mass which contains 6.02 x 1023 molecules of water = 18.0 grams (the
mass of one molecule of water)
2 H = 2 (1g) = 2 g
1 O = 1 (16g) = 16 g
18 grams/mole
Practice- The mass of 1 mole of Carbon Dioxide (CO 2) = the mass which contains 6.02 x 1023 molecules of carbon dioxide
=
44 g/mole (the mass of one molecule of CO2)
1 C = 1 (12g) = 12 g
2 O = 2 (16g) = 32 g
44 g/mole
The gram formula mass⟶the mass of one mole of formula units = the amount which contains 6.02 x 10 23 formula units = the formula mass in
grams.
Example- The mass of 1 mole of sodium chloride (NaCl) –an ionic compound , therefore no molecules = the amount
which contains 6.02 x 1023 formula units = 58.5 grams (the mass of one formula unit in grams)
1 Na = 1 (23.0 g) = 23.0 g
1 Cl = 1 (35.5 g) = 35.5 g
58.5 grams/mole
Practice- The mass of 1 mole of magnesium fluoride (MgF 2) = the amount which contains 6.02 x 10 23 formula units =
64 grams (the mass of one formula unit in grams)
1 Mg = 1 (24 g) = 24 g
2 F = 2 ( 19 g) = 38 g
64grams/mole
Assignment #26: Molar Mass Practice Worksheet
Find the molar masses of the following compounds:
1)
NaBr
23 + 80 = 103 g/mol
2)
AgF
107.9 + 19.0 = 126.9 g/mol
3)
PbSO4
207.2 + 32.1 + (16 x 4) = 303.3 g/mol
4)
Ca(OH)2
40.08 + (16 x 2) + (1.01 x 2) = 74.1 g/mol
5)
Na3PO4
(23
6)
x 3) + 31 + (16 x 4) = 164 g/mol
(NH4)2CO3
(14 x 2) + (1 x 8) + 12 + (16 x 3) = 96 g/mol
7)
C6H12O6
(12 x 6) + (1 x 12) + (16 x 6) = 180 g/mol
8)
Fe3(PO4)2
(55.8 x 3) + (31 x 2) + (16 x 8) = 357.4 g/mol
9)
(NH4)2S
(14 x 2) + (1 x 8) + 32.1 = 68.1 g/mol
10)
Zn(C2H3O2)2
65.4 + (12 x 4) + (1 x 6) + (16 x 4) = 183.4 g/mol
Assignment #27: Grams/Moles Calculations KEY
1)
30 grams of H3PO4
0.31 moles
2)
25 grams of HF
1.25 moles
3)
4 moles of Cu(CN)2
462 grams
4)
110 grams of NaHCO3
1.31 moles
5)
88.4 moles of NI3
34,893 grams
6)
1.1 grams of FeCl3
0.0068 moles
7)
6.6 moles of ZnO
537.2 grams
8)
5.4 moles of K2SO4
941.2 grams
9)
987 grams of Ra(OH)2
3.80 moles
10)
564 grams of copper
8.87 moles
11)
1.2 moles of (NH4)3PO3 159.6 grams
12)
12.3 grams of CO2
0.28 moles
13)
21.3 moles of BaCO3
4202.5 grams
14)
89 grams of Pb(CH3COO)4 0.20 moles
15)
5.6 moles of C6H6 436.8 grams
Assignment #28: Mole Calculation Practice Worksheet
Answer the following questions:
1)
How many moles are in 25 grams of water?
H2O: 18 g/mol
Moles =
2)
Given Mass
Gram formula mass
Moles =
25 g
18 g/mol
= 1.4 mol H2O
How many grams are in 4.50 moles of Li2O?
Li2O: 29.8 g/mol
Moles =
3)
Given Mass
Gram Formula mass
4.50 mol = X g
1
29.8 g/mol
= 134.1 g Li2O
How many grams are in 23.0 moles of oxygen?
O2: 32 g/mol
Moles =
4)
Given Mass
Gram Formula Mass
23.0 mol
1
=
X g
= 736 g O2
32 g/mol
How many moles are in 0.00576 grams of H2SO4?
H2SO4: 98 g/mol
Moles =
5)
Given Mass
Moles =
Gram Formula Mass
.00576 g = 0.0000588 mol H2SO
98 g/mol
How many moles are in 25.0 grams of NH3?
NH3: 17 g/mol
Moles =
6)
Given Mass
Gram Formula Mass
Moles =
X g
17 g/mol
=1.47 mol NH3
How many grams are in 0.136 moles of N2I6?
N2I6: 789.4 g/mol
Moles =
Given Mass
Gram formula Mass
.136 mol
1
= Xg
= 107.36 g N2I6
789.4 g/mol
Assignment #29: Moles Regents Questions
1.) A 1.0-mole sample of krypton gas has a mass of
(1) 19 g
(2) 36 g
(3) 39 g
(4) 84 g
2.) What is the gram-formula mass of Ca3(PO4)2?
(1) 248 g/mol
(2) 263 g/mol
(3) 279 g/mol
(4) 310. g/mol
Base your answer to questions 3 and 4 on the information below.
Rust on an automobile door contains Fe2O3(s). The balanced equation representing of the reactions between iron
in the door of the automobile and oxygen in the atmosphere is given below.
4Fe(s) + 3O2(g) → 2Fe2O3(s)
3.) Determine the gram-formula mass of the product of this reaction.
159.6 g/mol
4.) Write the IUPAC name for Fe2O3.
Iron (III) Oxide
Base your answer to question 5 on the information below.
The compound 1,2-ethanediol can be mixed with water. This mixture is added to automobile radiators as an
engine coolant. The cooling system of a small van contains 6690 grams of 1,2-ethanediol. Some properties of
water and 1,2-ethanediol are given in the table below.
5.) Calculate the total number of moles of 1, 2-ethanediol in the small van’s cooling system. Your response
must include both a correct numerical setup and the calculated result.
Moles =
Given Mass
Gram formula mass
Moles =
( From Table T)
6690
62
= 107.9 moles
Base your answer to question 6 on the information below.
An unsaturated solution is made by completely dissolving 20.0 grams of NaNO3 in 100.0 grams of water at
20.0°C.
6.) Calculate the number of moles of NaNO3 (gram-formula mass = 85.0 grams per mole) used to make this
unsaturated solution.
Moles =
20
= .23 moles
85
Assignment # 30: More Practice with moles with your teacher
Propane reacts with (burns) oxygen to give carbon dioxide and water as products.
Written unbalanced:
____C3H8
+
____O2

_____ CO2
+ _____ H2O
⟶We balance the equation in order to be able to tell how many propane molecules react with how many
oxygen molecules and how many molecules of products are formed.
(Note- When it’s balanced it means that the number and kind of atoms on the left side = the number and kind of atoms
on the right side)
C3H8
+
5 O2
 3 CO2
+ 4 H2 O
1. How many molecules of O2 are needed to react with 1 molecule of propane?
____1_____
2. How many moles of molecules of O2 are needed to react with 1 mole of propane molecules?
_____1____
3. How many grams of C3H8 are needed to react with 5 moles of O2?
_________44 g_______
4. How many grams of H2O are produced? HINT- there are 5 moles produced, not 1! _______72 g______
Assignment #31: Working with Moles
For the reaction:
2H2S + 3O2 ----------> 2SO2
+ 2 H2O
1.) If you start with 6.0 moles of O2 how many moles of H2S will react and how many
moles
of SO2 and H2O will be formed?
Given:
2 moles
3 moles
2 moles
2 moles
__4 mol____
6 moles
_4 mol__
_4 mol__
2.) If you start with 8.0 moles of H2S how many moles of O2 will react and how many
moles
of SO2 and H2O will be formed?
Given:
2 moles
3 moles
2 moles
2 moles
8 moles
_12 mol__
_8 mol__
_8 mol_
3.) If you end up with 7 moles of H2O, how many moles of H2S and how many moles of O2
did you start with and how much SO2 will be formed
Given:
2 moles
__7 mol__
3 moles
2 moles
2 moles
10.5 mol__
_7 mol__
7 moles
4.) If you end up with 6.3 moles of SO2, how many moles of H2S and how many moles of O2
did you start with and how much H2O will be formed
Given:
2 moles
_6.3 mol_
3 moles
2 moles
2 moles
_9.45 mol__
6.3 moles
_6.3 mol_
Consider the reaction of N2 and H2 to form ammonia (NH3)
N2 + 3 H2

2 NH3
5.) If you start out with 0.40 moles of H2
a. How many moles of N2 are required to react completely?
= 0.13 mol N2
.4 = x
3
1
b. How many moles of NH3 are formed?
.4
3
=
x
= 0.27 mol NH3
1
6.) If you start out with 6.3 moles of N2
a. How many moles of H2 are required to react completely?
1 = 3 = 18.9 mol H2
6.3 X
b. How many moles of NH3 are formed?
1 = 2 = 12.6 mol NH3
6.3 X
CHALLENGE:
c. How many grams of NH3 are formed?
gfm = (14.01 x 1) + (1.01 x 3) = 17.04 g/mol
Moles
= Given Mass
Gram Formula Mass
12.6 =
X
1
17.04
Appendix E Honors Only
= 214.7 g NH3
Questions
No
No
No
18 g
55.8 g
32 g
44 g
Assignment #32
1.)
2.)
3.)
4.)
5.)
6 moles of hydrogen
9 moles of oxygen
3 moles of hydrogen
20 moles of oxygen
6 moles of potassium nitrate
Ionic bonding and covalent bonding Lewis dot diagrams will be drawn on the
white board/Smartboard for assignment #32 and #33.
Assignment #33: Covalent Bonding Activity
Purpose: To determine how atoms of non-metals work out the problem of getting
all atoms stabilized.
Preparation: Answer the following…
When atoms bond using an “ionic bond,” the metal atoms will lose valence electrons to nonmetals atoms.
1.) What type of atoms are involved in making “Covalent Bonds”?
All nonmetal atoms.
2.) Why can’t a “transfer” of valence electrons (like what happens in ionic bonding) from one
atom to another happen in some situations?
Because electrons are not being lost to be transferred they are being shared.
3.) What DOES happen in terms of valence electrons, when atoms form a “covalent bond”?
Valence electrons are shared.
4.) What does the word “Co-Valent” mean?
Sharing of valence electrons between nonmetal atoms.
5.) The smallest units in ionic compounds are the oppositely charged ions. These ions can exist
independent of one another (as in solutions). What are the smallest units in covalent
compounds?
Molecules
Assignment #34: Naming—Additional Practice
Name the following compounds
1. KCl
Potassium Chloride
2. MgBr2
Magnesium Chloride
3. Ca(CO3)
Calcium Carbonate
4. FeCl3
Iron (III) Chloride
5. Pb(SO3)
Lead (II) Sulfite
6. Be(NO3)2
Beryllium Nitrate
7. Al2O3
Aluminum Oxide
8. NaF
Sodium Fluoride
9. Zn3(PO4)2
Zinc Phosphate
10. Fe2O3
Iron (II) Oxide
Write the formula for the following compounds
1. iron (II) sulfite
FeSO3
2. magnesium chloride
MgCl2
3. chromium (IV) bromide
CrBr4
4. potassium permanganate
KMnO4
5. sodium phosphide
Na3P
6. beryllium nitride
Be3N2
7. nickel (III) oxide
Ni2O3
8. manganese (IV) sulfate
Mn(SO4)2
9. barium carbonate
BaCO3
10. aluminum hydroxide
Al(OH)3
Balancing equations
These are the coefficient for the balancing.
1.
2 :
1
→
2 : 1
2.
1 : 2
→
1 : 1
3.
1 : 2
→
2 : 1
4.
2
→
2 : 1
5.
4 : 1
→
2
6.
2 : 2
→
1 : 2
7.
2 : 1
→
2
6 : 6
→
1 : 6
Challenge