PRACTICE PACKET LEVEL 5: BONDING
Regents Chemistry: Mr. Palermo
Practice Packet
Level 5: Bonding
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PRACTICE PACKET LEVEL 5: BONDING
LESSON 1: INTRO TO BONDING & TYPES OF BONDS
1. For each phrase, check either “bond breaking” or “bond forming”
Bond
Breaking
1.
2.
b.
c.
d.
e.
f.
g.
Bond
Forming
Stability of the chemical system increases
Energy is released
Cl + Cl  Cl2
exothermic
endothermic
N2  N + N
Energy is absorbed
Stability of the chemical system decreases
2 Identify which bond type is described by each statement below. Choices: ionic, covalent,
metallic
Substance
a.
b.
c.
d.
e.
g.
h.
i.
j.
k.
Bond type
NaCl(s)
CO2 (g)
NO (g)
Cu (s)
MgBr2 (g)
HCl (aq)
SO2 (g)
AlCl3 (s)
Ag (s)
NaBr (s)
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PRACTICE PACKET LEVEL 5: BONDING
3. For each example provide the molecule, bond and determine when and if it conducts
electricity
Type of Molecule
(metallic, ionic,
molecular)
Type of Bond
(Metallic, ionic, covalent,
both ionic and covalent)
Conducts
electricity?
(check all that apply)
No
a.
b.
c.
d.
e.
Li2O
AlCl3
F2
CH4
f.
g.
h.
Fe
Na3PO4
i.
C (diamond)
j.
k.
l.
C (graphite)
H2
m.
n.
o.
p.
q.
r.
s.
t.
(s)
(aq)
(l)
HI
CaO
Na
NH4Br
KNO3
O3
SiO2
NH3
FeBr2
Hg
CO2
Indicate which type of substance is described by each statement.
Choices: covalent (molecular), ionic, metallic
Type of substance
a.
Can conduct electricity in the solid and liquid phases
A soft substance whose atoms are held together by
b.
covalent bonds
c. Low melting point and poor electrical conductor
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PRACTICE PACKET LEVEL 5: BONDING
d. Can conduct electricity when aqueous or molten (liquid)
Lesson 2: Bond Polarity
1. Electronegativity values generally _______________ down a group and ____________ across a
period.
2. Metals tend to have ____________ electronegativity values and nonmetals are _____________
values.
Using your table above find the electronegativity difference for each substance.
Then, check which bonds are present. If it’s a metal and a nonmetal it is automatically
Ionic.
Electronegativity
difference(s)
Substance
Ionic
Covalent
Polar
Nonpolar
I2
PCl3
SiO2
Br2
CO2
NaCl
CH4
N2O5
NH3
KCl
3. Indicate which atom will have the positive charge and which will have the negative
charge in the following polar bonds:
H-Cl
H-F
S-F
N-O
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PRACTICE PACKET LEVEL 5: BONDING
4. What is electronegativity?
5. What factor causes some combinations or atoms to form ions, and other combinations
of atoms to form covalent bonds. Explain in detail.
6. What is a nonpolar covalent bond? Explain the electronegativity differences attributed
to this type of bond.
7. What is a polar bond? Explain the electronegativity differences attributed to this type of
bond.
8. Explain the relationship between electronegativity and polarity.
9. What is a dipole?
10. What symbol indicates a partial charge? ___________________
11. How do you determine which atom gets the partial negative charge?
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PRACTICE PACKET LEVEL 5: BONDING
12. Given the following indicate which atom will receive the partial negative charge and
which atom will receive the partial positive charge. Place the partial charges in the
upper right hand corner of the atom symbol:
a.
H – Cl
b.
H–F
c.
S –F
d.
N–O
13. Compare the degree (which compound is most polar, which is least polar) of polarity in
HF, HBr, HCl, and HI.
14. Classify the type of molecule the diagrams below represent (Ionic, Polar Covalent, or
Nonpolar Covalent), and explain your reasoning.
Electron Distribution
Diagram
Type of Compound
Reason for Classification of Compound
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PRACTICE PACKET LEVEL 5: BONDING
LESSON 3: LEWIS (ELECTRON) DOT DIAGRAMS FOR IONIC
COMPOUNDS
1. Complete the table below (electron dot diagrams for ions)
Ion
Electrondot
Diagram
Electron
Configuration
Ion
a.
sodium
Na+
e.
oxide
O2
b.
aluminum
Al3+
f.
bromide
Br
c.
calcium
Ca2+
g.
phosphide
P3
d.
magnesium
Mg2+
h.
sulfide
S2
Ionic
compound
(name & formula)
a.
b.
c.
d.
e.
sodium
fluoride
NaF
potassium
chloride
KCl
calcium
iodide
CaI2
magnesium
oxide
MgO
rubidium
oxide
Electron-dot
Diagram
Total
# of
ions
Electrondot
structure
Ionic
compound
Electron
Configuration
Electron-dot
structure
Total
# of
ions
(name & formula)
f.
aluminum
chloride
g.
sodium sulfide
h.
lithium
hydride
i.
aluminum
oxide
j.
calcium
phosphide
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PRACTICE PACKET LEVEL 5: BONDING
Ru2O
2. Complete the table below (electron dot diagrams for ionic compounds)
LESSON 4: WRITING FORMULAS FOR IONIC COMPOUNDS
1. Complete the table for the following ionic substances
Name
Criss-Cross,
Reduce
Formula
Electron-dot diagram
potassium fluoride
lithium bromide
strontium chloride
barium iodide
gallium nitride
zinc sulfide
2. Write the formula for the ionic substances below containing polyatomics (use table E)
Name
Criss-Cross,
Reduce
Formula
Name
sodium sulfate
barium
phosphate
aluminum
chromate
calcium
hydroxide
magnesium
hydrogen
carbonate
potassium
hydrogen
sulfate
lithium
permanganate
ammonium
chloride
rubidium oxalate
sodium
acetate
Criss-Cross,
Reduce
Formula
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PRACTICE PACKET LEVEL 5: BONDING
More Practice Writing Ionic Formulas
1. Write the formula for each ionic compound below. Remember to Criss cross and reduce the
oxidation states. The first problem is done for you.
Name
Cation (+)
Anion (-)
1
Sodium Chloride
Na1+
Cl1-
2
Aluminum
Chloride
Al3+
Cl1-
3
Aluminum
Phosphide
4
Magnesium Oxide
5
Cesium Fluoride
6
Strontium Nitride
7
Lithium Sulfide
8
Calcium Chloride
9
Sodium Bromide
10
Beryllium Iodide
11
Strontium Fluoride
12
Aluminum
Fluoride
13
Potassium Nitride
Criss Cross &
Reduce
Formula
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PRACTICE PACKET LEVEL 5: BONDING
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Sodium Sulfide
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Lithium Oxide
2. Writing Formulas containing Polyatomic Ions
1. Ionic Compound
Criss Cross and Reduce
Formula
2. 1. calcium carbonate
3. barium nitrate
4. ammonium sulfate
5. aluminum hydroxide
6. calcium phosphate
7. cesium nitrate
8. sodium nitrite
9. 8. calcium sulfate
10. beryllium sulfate
11. sodium carbonate
12. magnesium phosphate
14. ammonium nitrate
3. What is the sum of the charges on all ionic compounds?
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PRACTICE PACKET LEVEL 5: BONDING
LESSON 5: NAMING IONIC COMPOUNDS
Formula
How MANY
oxidation states
listed for the
Metal?
One
Two or
more*
Name
*Use charge of
NONmetal to figure
out charge of metal
(reverse criss cross)
Show work here:
LiBr
Ag2O
SnO
tin (II) oxide
+2 –2
SnO
Ba3N2
AgBr
Cu3P
Mg(NO3)2
Co2O3
NaNO3
KI
NaClO
Fe(OH)3
PbSO4
NaHCO3
Ni2(SO4)3
Ti2O3
Al2(SO3)3
Al(CN)3
NH4Cl
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PRACTICE PACKET LEVEL 5: BONDING
KNO3
CaCO3
(NH4)2CO3
MORE NAMING PRACTICE:
Formula
How MANY
oxidation states
listed for the
Metal?
One
Two or
more*
Name
*Use charge of
NONmetal to figure
out charge of metal
(reverse criss cross)
Show work here:
1. KOH
2. LiI
3. AlF3
4. FeCl2
5. MgO
6. Co(NO3)2
7. MgSO4
8. NH4Cl
9. CrPO4
10. Ba(OH)2
11. PbS
12. Na2CO3
13. BaF2
14. Cu(NO3)2
15. AgI
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PRACTICE PACKET LEVEL 5: BONDING
16. NiSO4
17. Zn3(PO4)2
18. Na3N
LESSON 6: NAMING AND FORMULA WRITING: COVALENT
(MOLECULAR) COMPOUNDS
1. Write formulas for the following molecular substances.
Name
Formula
Name
dinitrogen trioxide
silicon tetrafluoride
diphosphorus
pentoxide
carbon tetrachloride
sulfur dioxide
boron triiodide
silicon dioxide
carbon disulfide
xenon pentafluoride
phosphorus
pentabromide
dihydrogen monoxide
boron trihydride
Name
Formula
Name
Formula
Formula
N2O5
H2S
SF5
BF3
PBr3
PH3
SO3
H2O
B2H4
Cl2
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PRACTICE PACKET LEVEL 5: BONDING
PCl5
PCl3
P2O5
SCl6
CS2
CO2
CO
NO
BCl3
NO2
2. Write IUPAC Names for the following molecular (covalent) substances
LESSON 7: LEWIS (ELECTRON) DOT DIAGRAMS FOR COVALENT
SUBSTANCES
1. Complete the chart.
Molecule
(name &
formula)
a.
b.
methane
CH4
nitrogen
N2
Total #
of
valence
e-’s
Electron-dot
structure
Molecule
(name &
formula)
g.
carbon
tetrachloride
CCl4
h.
carbon
dioxide
CO2
i.
phosphorus
trichloride
PCl3
c.
ammonia
NH3
d.
water
H2O
j.
e.
oxygen
O2
k.
Total #
of
valence
e-‘s
Electron-dot
structure
dihydrogen
monosulfide
H2S
carbon
monoxide
CO
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PRACTICE PACKET LEVEL 5: BONDING
f.
fluorine
F2
l.
hydrogen
H2
MORE PRACTICE:
Draw the lewis (electron) dot structures for the following molecular (covalent) substances.
Cl2
SH2
H2
CF4
CS2
SiO2
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PRACTICE PACKET LEVEL 5: BONDING
SF2
HF
LESSON 8: MOLECULAR POLARITY
1. Fill in the chart below.
2. In terms of lone pair electrons, how can you determine if a molecule is polar?
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PRACTICE PACKET LEVEL 5: BONDING
3. What molecular shapes are always polar?
4. What molecular shapes are always nonpolar?
5. How can a molecule be nonpolar if it contains polar bonds?
2. Fill in the chart below
Molecule
Dot Diagram
Distribution of
Charge
(symmetrical or
asymmetrical)
Molecular
Polarity (polar
or nonpolar
molecule)
Molecular
Shape (linear,
pyramidal,
tetrahedral or
bent)
CCl4
NF3
Br2
CS2
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PRACTICE PACKET LEVEL 5: BONDING
SiO2
LESSON 9: INTERMOLECULAR FORCES (IMF’S)
1. Which of the following will have the higher boiling point? Explain your answer using
intermolecular forces. NH3 or N2
2. Why does dry ice (solid CO2) evaporate before sodium chloride?
3. Why does gasoline (C8H18) exist in the liquid form while methane (CH4), the gas we use to
power out bunsen burners, exists in the gas form even though both compounds are
nonpolar?
4. Identify the intermolecular forces that exist in the following molecules.
Compound
Type of IMF
H2O
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PRACTICE PACKET LEVEL 5: BONDING
N2
HCl
LiCl
5. Of the compounds in question 4, which has the strongest surface tension?
6.
In terms of the forces of attraction holding them together, explain why a NaCl crystal
has a melting point of 800C while an ice cube of pure water has a melting point of 0C.
7.
List the noble gases from highest to lowest boiling point. Explain your answer based on
intermolecular forces of attraction.
8.
Explain why I2 is a solid, Br2 is a liquid but Cl2 and F2 are gases even though they are all
Halogens.
9.
List the following substances from highest to lowest melting point; use attractive force
to justify your answers. KCl, Cl2, CH4, H2O, PCl3
10.
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PRACTICE PACKET LEVEL 5: BONDING
11.
Network Solids
The best example of a network solid is a “diamond”. Look at the model of a diamond below. Note
that the carbon atoms are bonded together with covalent bonds. The basic building unit is an
atom of carbon. The structure has a very definite tetrahedral crystal shape, because these atoms
are arranged and held rigidly in a fixed pattern. A diamond is very hard (a “10” on the Moh’s Scale
of Hardness…the highest value possible). In order to scratch a diamond, you must break 1000’s of
very strong covalent bonds! Similarly, to melt (or boil) a network solid, like a diamond, you must
break 1000’s of these covalent bonds. This involves considerable energy and is the reason for their
high melting points. It is because of these high temperatures and their hardness that network
solids are frequently used in industry as “abrasives” (on sandpaper and on the tips of drills for
cutting tools). You don’t have to worry about them melting if it gets too hot from friction or being
scratched and dulled when contacting most other surfaces. Network solids have the type of
properties you would expect from atoms being held together via strong covalent bonds, e.g.
diamonds. They have very high melting points and are practically insoluble; are mostly
nonconductors (no free electrons or ions); and they are very brittle (atoms must maintain a fixed
crystal structure, if they are pushed too close together …. they repel).
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PRACTICE PACKET LEVEL 5: BONDING
Graphite is also shown below. Note that it is also pure carbon, like a diamond. However, the
covalent bonds only attach carbon atoms in 2 directions, not 3 like diamonds. The dashed lines
between the layers of covalently bonded carbon atoms represent weak Van der Waal forces.
Graphite is a 2 dimensional network solid. The strong covalent bonds only go in “plates”, in 2
directions. The “plates” are connected via weak VDW forces. Graphite STILL has a high melting
point. – You must weaken/break all of its bonds (VDW and Covalent) to melt it. But since the
weaker VDW forces are present and break easily, graphite is often used as a “dry lubricant”. If you
squirt graphite dust into a lock, it will lodge between the lock’s moving metallic parts. When you
put in a key and turn, the graphite structure will break apart between its “plates.” VDW’s break
and make it turn more easily. Graphite also has free, “delocalized” electrons (it is a resonant
structure… there really are no double bonds present, but free electrons) thus…graphite is a
network solid that is capable of conducting electricity. This is not characteristic of most network
solids. Of course, pencil lead is graphite. What bonds break when you write??? Silicon bonds like
carbon to form a network structure. The computer industry depends on “silicon chips” ,which are
made conductive by placing impurities in their structure. These then provide for free electrons and
allow the chip to do its job. But, pure silicon does not conduct.
Many network solids are composed of various combinations of relatively few elements on the
periodic table. The elements B, C, Al, Si are found in many network solids. They can be pure or
combine with one another or combine with elements near them. For example, SiO 2, quartz is an
example of a network solid. Corundum, Al2O3, is a network and a common abrasive used on
“sandpaper”. Many gem stones, like diamond, are network solids. Emerald is made of the mineral
“beryl”. Its formula is Be3Al2 (Si6O18) . Ruby is a form of corundum.
1. What is a Network solid? Give an example
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PRACTICE PACKET LEVEL 5: BONDING
2. What are some physical properties of Network solids?
Unit Review/Study Guide
INTRODUCTION TO BONDING
Elements are the simplest form of matter and cannot be decomposed. Compounds can be formed
between two or more elements. They can be decomposed chemically.
a. Which of the following is a compound?
Ne
H2O
Be
F
b. Which of the following cannot be decomposed by chemical means?
C12H24
NH3
Li
CS2
Atoms bond in order to obtain a stable electron configuration, like noble gases, called the octet.
Most atoms will gain or lose electrons in order to have eight valence electrons. However, small
elements such as H, Li, and Be will settle for two valence electrons. Obtaining an octet makes the
atoms more stable and they can release energy. The electrons obtain the octet by sharing or
transferring electrons.
a. Draw the Lewis dot diagram of the following elements:
Na
Mg
Al
Si
P
S
Cl
Ar
b. Draw the Lewis dot diagram of the following ions:
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PRACTICE PACKET LEVEL 5: BONDING
Na+
Mg+2
Al+3
P-3
S-2
Cl-
c. Explain why the metals lost electrons but the nonmetals gained electrons.
__________________________________________________________________________
d. Fill the blanks with release or absorb: “When atoms bond they ____________ energy. In order
to break a bond, energy must be __________________.
IONIC BONDING
Compounds that form between a metal and a nonmetal contain ionic bonds, transferring
electrons. Ionic bonds are strong. Ionic compounds have high melting points, are generally solids
at room temperature, and conduct in the liquid phase.
a. Which of the following has ionic bonds?
NaCl
NH3
Mg
b. Which of the following transfers electrons?
MgBr2
Li
CO2
c. Which of the following has a higher melting point?
Cu
C6H12
LiF
KI
Ne
d. Which of the following can conduct in the aqueous phase? NO
COVALENT BONDING
Compounds that form between two nonmetals have covalent bonds, sharing electrons. Covalent
bonds are weaker than ionic bonds. Covalent compounds have low melting points, are generally
gases, liquids, or powdery solids at room temperature, and never conduct. These are also known
as molecular compounds.
a. Which of the following has covalent bonds?
HF
LiCl
Rb
b. Which of the following shares electrons?
H2O
Ag
CaCl2
c. Which of the following can never conduct electricity?
Kr
Rb2O
H2O
NH3
CaCO3
d. Which of the following has both ionic and covalent bonds? Li
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PRACTICE PACKET LEVEL 5: BONDING
e. Which of the following is a molecular compound?
H2O
Mg
LiBr
METALLIC BONDING
Metallic Bonds form when a metal loses their valence electrons and a “sea of mobile electrons”
form that allows the metal to conduct electricity in the solid or liquid phase.
a. Which of the following is metallic?
NaCl
NH3
Mg
b. Which of the following has a sea of mobile electrons? Cu
C6H12
LiF
c. Which of the following can conduct in the solid phase? Ne
NAMING COMPOUNDS/FORMULA WRITING
Ag
CaCl2
When Ionic Compounds, always name the positive, cation first and then the negative, anion
last. The elements are named in the same order they appear on the periodic table. When
compounds have more than 2 elements, it contains a polyatomic ion. Use Table E on page 2 of
your reference tables. Transition Metals are in the middle group of the periodic table.
Nonmetals are on the right side of the staircase. They have multiple charges or oxidation
numbers and so you must show which charge you are using with roman numerals. Polyatomic
ions are a group of 2 or more atoms that are bonded very strongly and act as one ion. Name
the following:
CaCl2
NaF
LiOH
KNO3
CuBr2
CuBr3
Ni(OH)2
NiCl3
To write a formula of an ionic compound, write the two ions separately showing their charges.
Charges are on the periodic table. Then, swap the two numbers and drop the sign. Write the
formula for the following:
Sodium fluoride
Cesium oxide
Strontium acetate
Aluminum phosphate
Iron(III) iodide
Manganese (VII) oxide
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PRACTICE PACKET LEVEL 5: BONDING
When naming Covalent Compounds, use prefixes to indicate the number of each atom present in
the compound. Determine the prefix of each element using the subscript #. Remember, if only
1 atom is present for the first element do not use the prefix mono for that atom. Name the
following:
HCl
PCl5
N2O2
NH3
To write a formula of a covalent compound, write the least electronegative element first.
Determine the prefix of each element using the subscript #. Write the formula for the
following:
Sulfur hexafluoride
Nitrogen dioxide
Carbon Dioxide
Nitrogen monoxide
LEWIS STRUCTURES/GEOMETRY
Ionic Lewis diagrams show the ions involve in the bond, but no arrangement. Covalent Lewis
diagrams show the sharing of electrons with lines representing two electrons. They form shapes
such as linear, bent, pyramidal, and tetrahedral.
a. Draw the following and give the number of shared pairs, unshared pairs, and the shape if
applicable.
LiF
NH3
MgF2
CH4
Cl2
H2O
POLARITY
Bonds are polar when two atoms have different electronegativities and share unevenly. The more
electronegative atom has the electrons more of the time. Nonpolar bonds form when two atoms
have the same electronegativity values and share equally.
a. Label the bonds as polar or nonpolar:
NH3
CH4
Cl2
H2O
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PRACTICE PACKET LEVEL 5: BONDING
Molecules are polar when the molecule is asymmetrical. They are nonpolar if the molecule is
symmetrical.
b. Label the bonds as polar or nonpolar (Use your drawing to help you):
NH3
CH4
Cl2
H2O
INTERMOLECULAR FORCES
Intermolecular forces are what keeps molecules together (not atoms-that’s bonds) and are
responsible for phases, phase changes, surface tension and various other properties. Nonpolar
molecules have the weakest attractive forces dependent on their size (the bigger the stronger).
Polar molecules have stronger forces dependent on their polarity. Hydrogen bonds are a special
case of polar forces between H and either F,O, or N. Molecules that are hydrogen bonded have
high melting and boiling points, strong surface tension, and have closely packed particles.
a. Which of the following has the highest melting point?
HF
HCl
HBr
HI
b. Which of the above has the lowest boiling point?
______
______
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