Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 1 of 18
Honors Chemistry Unit 5B: Covalent Bonding, Nomenclature, VSEPR and Polarity
Vocabulary
Problem Set Due:
Test :
VOCABULARY:
covalent bonds
molecular compound
resonance structure
molecule
empirical formula
intermolecular forces
polar covalent
molecular formula
Van der Waals forces
nonpolar
VSEPR
(London) dispersion forces
OBJECTIVES:
 Be able to describe and identify the three types of bonds: ionic, molecular, and metallic.
 Be able to describe characteristics of ionic, molecular and metallic compounds.
 Be able to describe the difference between polar and nonpolar molecular bonds.
 Be able to determine bond type using the periodic table or electronegativity difference
Unit 5B EQs:
Day 1 – Review
 If X represents an element from group 2A of the periodic table, what would the formula of its fluoride be?
 If element M forms the compounds MBr3 and M2S3, element M would be found in which group of the periodic table? Why?
 According to the graph below, which letter indicates the noble gas? The alkali metals?
 Why is there such a large difference in ionization potential between them?
Day 2 - New Material:
 What type of bonds are found in the following:
a. CH4
c. KCl
b. Na3PO4
d. gold bar
 List the above in order of strongest bond to weakest bond.
Day 3
 Write the molecular formula for the following:
a. diantomony trioxide
d. oxygen difluorids
b. sulfur hexachloride
e. carbon tetrachloride
c. tetraphosphorus decoxide
f. fluorine monosulfide
 Write the name for each of the following molecular formulas:
a. CO2
d. CO
b. NH3
e. BI3
c. P2O5
f. SO3
Day 4
 Illustrate the correct structure formula for C2H4.
 What is a resonance structure? Draw an example of one from your notes.
 List the 6 main VSEPR shapes and draw an example of each.
Day 5
 What is the correct molecular shape of CH4? How was this determined?
 Compare and contrast polar covalent and nonpolar covalent molecules. Give an example of each.
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 2 of 18
BOND
IONIC – transfer
of electrons
(∆ EN) is >1.7
COVALENTSharing of
electrons
(∆ EN) is <1.7
METALLICFree flow of
Electrons
Atoms involved
Metal &
Nonmetal or
Metal &
polyatomic ion
Force
Attraction
between ions,
opposite charges
attract
Two nonmetals
Polar = unequal
sharing = partial
charge
Nonpolar = equal
sharing = no
charge
Two metals
Sharing of
electrons
Sharing of
electrons
between all
atoms
Properties
High melting point.
High boiling point
Water soluble
Crystalline
Aq solutions conduct a
current
Low melting & boiling
point
Brittle
Nonconductors
Examples
NaCl
MgO
CaS
K2SO4 (also
includes cov
bond)
Water
CO2
NH3
Good conductors.
Malleable = shapeable,
Ductile = able to be
drawn into wire
Copper wire
Iron bar
Relative strength
Generally the strongest
bond type = the larger
the ion charges the
stronger the ionic bond.
Weakest bond type(exception network
solids like diamonds)
Strongest: H bond
Dipole-dipole
Weakest: dispersion
forces
VSEPR & Molecular Geometry
Molecular Shape
Linear
Bent
Type of
Molecule
Atoms Bonded to Central
Atom
Lone Pairs of e-s on
Central Atom
AB2
2
0
2
1
3
0
4
0
3
1
2
2
AB2E
2 surrounding atoms – I
pair of lone e-s
Trigonal Planar
AB3
Tetrahedral
AB4
Trigonal
Pyramidal
AB3E
3 surrounding atoms – I pair
of lone e-
Bent
AB2E2
2 surrounding atoms – 2
pairs of lone e-
Covalent Bonding Guided Notes
Covalent Bonds: between two (or more) nonmetals
A little review…Ionic Bonds

Between a _____________________ and a _____________________________________

The metal __________________________________ loses electrons that the nonmetal ____________________.

The elements are held together (bonded) by an electrostatic charge, ___________________ attracted to negative.
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 3 of 18
Lewis Dot Diagrams
We use dots to represent the number of __________________________ electrons an atom has.
Can have up to __________ electrons (dots).
A new kind of bond…(Example slides: one fluorine atom bonding with another)
By ________________________electrons, both fluorine atoms end up with a full octet of electrons and stable orbitals.
Covalent Bonds: A bond in which two atoms _________________ pairs of electrons so that each can achieve stability.
Properties of Covalent Compounds

Made from 2 or more _______________________________________

Are usually liquids and gases (or solids with ____________________ melting points)
Covalent vs. Ionic Compounds
Covalent
Ionic
Chemical formulas called __________________________.
Chemical formulas called __________________________.
Nonmetal bonds with _____________________________.
________________ bonds with nonmetal.
Chemical formula begins with a _____________________.
Chemical formula begins with a _____________________.
Bonds formed by _____________________ electron pairs.
Bonds formed by attraction of positive ________________ to
negative ____________________.
Chemical formula represents _________________ number of
atoms in the compound.
Chemical formula represents the lowest whole number
________________ of elements in the compound.
Lewis Dot Diagrams
Writing Lewis Structures Using the NASL Method:
There are no “perfect” methods for writing Lewis structures. The NASL method makes counting electrons and number of bonds (shared
electron pairs) easier. You always have to think about te molecule or ion that you are working on. Use the NASL method as a guide.
You must makek decisions along the way when a molecule is an exception to the rule of eight.
Each letter in NASL designates a specific quantity. I present these to your sequentially:
N stands for the number of electrons that will satisfy the rule of eight.
For representative elements,
EXCEPT,
N = 8,
H and He N = 2
Be
N=4
B
N=6
To get N for a molecule or molecular ion, add the N values for all atoms in the molecule.
A stands for the number of available electrons, ie, electrons in the valence shell. This is the same as the group no. : IA, IIA, IIIA, IVA,
VA, VIA, VIIA, and VIIIA. If you do not know the group number for some reason, write the electron configuration and count up all the
electrons in the valence shell.
To get A for a molecule or molecular ion, count up the valence electrons for each atom in the molecule and add the charge if it is
negative, but subtract the charge if it is positive.
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 4 of 18
S stands for the number of electrons that are shared.
S=N–A
Divide this number by 2 to the get the number of covalent bonds in the molecule.
L stands for the number of electrons that are left over.
L=A–S
Divide this number by 2 to get the number of lone pairs in the structure.
When using this method, always keep the molecule before you. There are exceptions to the rule of eight and you must be conscious of
this when using any method for writing Lewis structures.

Practice NASL Method : Lewis Structures of Covalent Compounds worksheet
Remember: atoms form compounds to create a full ___________________of valence electrons
(Example slides: bonding of hydrogen and oxygen to form water)
We can use Lewis dot structures to represent covalent compounds!
1.
Add up the valence electrons of ______________the atoms in the compound.
2.
Write the symbols for the elements and _______________________them with a line(s).
a.
Least electronegative element goes in the ________________ (for example: carbon). H is __________________ on the
outside.
b.
3.
The other elements usually ___________________the center element.
______________ 2 from the total for each line drawn and arrange the rest of the electrons (dots) in pairs around the
____________________ elements first so that each atom has an octet. (H likes 2. B is happy with ____________.)
4. Check for happiness (_________________________)!
Example…NH3, Ammonia
N is in group 5A, so it has ________ valence electrons. H is in group 1A, so it has ________ valence electron.
The compound NH3 has 1 N and 3 Hs, so 5 + (3 x 1) gives a total of _______ valence electrons for the whole compound.
••
H–N–H
Subtracting 2 from the total for each line drawn:
8 – 6 = 2 __________________________.
H
You try…
HF
H2O
H–N–H
H
N is surrounded by 8 electrons, and
each H has 2 electrons. (Each
_____________ in the dot
structure represents 2 electrons.)
F2
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 5 of 18
Double and Triple Bonds

DOUBLE and TRIPLE bonds are formed when atoms share 2 or 3 pairs of electrons to attain a ___________________, noble gas
configuration.

These bonds are represented with ______________________________ lines in a dot diagram.
Example: Nitrogen molecule, N2
1.
Count the total number of _________________________ electrons, 5 for each nitrogen atom giving a total of 10.
2.
Draw the atoms _________________________ with a line.
3.
Then ___________________ 2 for each bond (line). 10 – 2 = 8 remaining. Place the 8 remaining electrons in
________________around the 2 nitrogens.
4.
But, the Ns aren’t happy. They only have _________ electrons each. What can we do??
We can share ________________ than one pair.
Draw the final N2 structure with a triple bond here:
Now each N has _______ electrons and is happy.
Let’s try a couple together…
(Show your work.)
CO2
C2H2
Coordinate Covalent Bonds
Formed when one atom contributes _________ bonding electrons so that each atom can achieve a noble gas configuration
To represent this type of bond, we use an arrow going ____________ the donor TO the atom who is receiving the electrons.
Let’s look at the compound _______________________________________________.
O donates a _______________ pair of electrons for sharing, forming a COORDINATE COVALENT BOND and allowing each atom to
achieve noble gas configuration.
Draw the final structure for CO here:
Remember to include the arrow.
Resonance Structures
Structures that occur when it is possible to write 2 or more _____________________electron dot structures that have the same # of
electron pairs for a molecule or an ion.
Example O3, Ozone
Try NO2 The structures of ions are enclosed in _______________________ with the ________________ written on the outside.
 Negative ions _____________ electrons to the total count; positive ions take away electrons.
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 6 of 18
Can also classify bonds based on electronegativity - if the electronegativity difference (∆ EN) is <1.7 = covlent bond
 Find the ∆ EN for the bond formed between O and N
Example: Show the bonding that occurs between Cl and F.
Example two: H and F – hydrogen is an exception to the octet rule!
You try: using Lewis dot diagrams illustrate the bonding that occurs between Br and I.
Structural Formula Practice
Draw the Lewis dot structure for the following compounds with single bonds:
1) HBr
2) PCl3
3) H2S
4) SCl2
5) PH3
6) CCl2F2
Draw the Lewis dot structures for the following compounds with double or triple bonds:
1) SO3
2) HCN
3) SO2
4) N2
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 7 of 18
Structural Formulas
CS2
BF3
NI3
H2Se
Cl2
C2H4
VSEPR Theory
V___________________ S____________ E________________ P__________ R__________________
Definition: Because electrons repel, _________________________________________ adjusts so the valence-electron pairs are as far
apart as possible.
Molecular Shape
The shape of a molecule is determined by 2 factors:
1. The number of atoms bonded to the _______________________ atom.
Double and triple bonds ________________ change the general shape.
2. The number of __________________________ pairs of electrons on the ___________________ atom.
Unshared pairs of electrons on ______________ atoms in the compound don't affect the shape.
VSEPR Theory: Predicts 3D shape of molecules.
1. _________________ 2. ___________ 3. trigonal _______________ 4. trigonal__________________
5. ______________________ 6. trigonal ________________________ 7. _______________________
Linear
 2 atoms bonded to the _______________ atom

0 lone _____________ on the central atom.

All 2-atom molecules are considered linear.
Bent
 2 atoms bonded to the central atom

1 or 2 lone pairs on the central atom

Notice that the Lewis structure does _______ have
to show the bent shape.
2 Examples
of Linear
Example of
Bent
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 8 of 18
Trigonal Planar
 3 atoms bonded to the ________________ atom

0 lone pairs on the central atom.

Since there are no lone pairs on the central atom,
Example of
Trigonal Planar
the compound is _____________.
Trigonal Pyramidal
 3 atoms bonded to the __________________ atom

1 lone pair on the central atom

The lone pair on the central atom ______________
Example of
Trigonal
Pyramidal
down on the other 3 atoms causing a pyramidal shape.
Tetrahedral
 4 atoms bonded to the central atom

Example of
Tetrahedral
0 lone pairs on central atom
Trigonal Bipyramidal
 _______ atoms bonded to the central atom

o
An _____________________ octet
o
Possible for p-group elements from ______ or later periods
0 _______ pairs on central atom
Octahedral
 _______ atoms bonded to the central atom
o

Example of
Trigonal
Bipyramidal
Also an ___________________ octet
Example of
Octahedral
0 lone pairs on central atom
Summary
Shape
Linear
Bent
Trigonal Planar
Trigonal Pyramidal
Tetrahedral
Trigonal Bipyramidal
Octahedral
# of atoms bound to
central atom
# of lone pairs on
central atom
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 9 of 18
Model Building Lab
Chemical Formula
H2O
NH3
CH4
H 2S
CCl4
CCl2F2
N2
CO2
Color-coded key:
Lewis Structure
Sketch of Molecule
Shape of Molecule
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 10 of 18
Bond Types
3 Types of Bonds
Ionic, Polar Covalent (Molecular), & ___________________________________ (Molecular)
Remember Electronegativity?
The tendency for an atom to attract electrons to itself in a _____________________.
The higher the value, the ______________________ it is at attracting electrons.
The difference b/w electronegativity values determines what type of ______________ will be formed.
Electronegativity is a scale from __________, Cs, to __________, F.
It generally ______________________ across a period and decreases ________________ a group.
Why don’t noble gases have a value? _________________________________________________________
Ionic Bonds
If the electronegativity difference is ___________________________________, one atom will
pull the electron completely ______________ from the other atom.
The electrons are _______________ shared.
An __________________ bond is formed as + and – attract.
Electronegativity of Na is _____________; Cl is _____________.
3.0 – 0.9 = 2.1; difference __________________________, so…Ionic.
Polar Covalent Bonds
Covalent bonds _____________________ electrons.
The shared pairs are pulled, similar to a ____________________________, between the nuclei of the atoms sharing the electrons.
If the electronegativity difference is between _____________________________, one side of the bond becomes
_______________________more negative and the other side becomes slightly more _____________________.
This is a _______________________________________ bond.
The electronegativity of H is ___________; Cl is ____________
o
Oxygen
3.0 – 2.1 = __________; difference is b/w 0.4 – 2.0,
so ________________ covalent.
The electronegativity of O is ____________; H is_____________
o
3.5 – 2.1 = ________; difference is b/w 0.4 – 2.0, so _______________ covalent.
Nonpolar Covalent Bonds
When the atoms have equal pull, causing the electrons to be equally shared, the bond
is _________________________________ covalent.
Neither side of the bond is even slightly _________________ or negative.
The electronegativity difference is b/w ______________________.
This is the type of bond that occurs between 2 atoms of the _______________ element. (H 2, O2, Cl2, etc.)
The electronegativity of H is 2.1.
o
2.1 – 2.1 = _________; difference is b/w 0.0 – 0.4, so ________________________ covalent.
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 11 of 18
Attractions Between Molecules: Van der Waals Forces
Weaker than either the ionic or _________________________ bonds that form between atoms in a compound.
Responsible for determining if a compound is a liquid, gas, or ________________.
3 basic types from weakest to ______________________________
o
(London) _________________________________ forces
o
Dipole __________________________________
o
_________________________________________
London Dispersion Forces
________________________ of all molecular attractions
Caused by the motion of electrons producing a temporary ____________
Strength of dispersion forces generally increases as # of electrons in the molecule _______________________.
___________ molecules have these weak attractions
Dipole Interaction
Occurs when ______________ molecules are attracted to one another.
Hydrogen Bonds
Occurs b/w molecules in which H is covalently bonded to O, N, or ________ which are very electronegative
o
Causes __________ polar molecules that are _________________ attracted to each other
o
Still only has about __________ of the strength of a covalent bond
Very Important
o
Reason ice is less ____________ than water
o
Reason for the relatively ____________ b.p. of water
o
Responsible for the double helix of the _________ molecule.
Nonpolar or Polar Molecules
We now know how to determine if the bond b/w atom and atom in a ____________________ is polar or nonpolar.
But…what about the whole _________________________?
Draw the Lewis __________________. If the central atom has any unshared pairs, the molecule is _______________.
If there are no unshared pairs on the _________________ atom, look at the atoms around the central atom.
If they are all the _________________, the molecule is nonpolar.
If any one of them is different, the molecule is ________________.
(In a 2-atom molecule, if the bond between the 2 atoms is _______________, then the whole molecule is polar.)
Dot Structure
Polar /
Nonpolar
Dot Structure
H20
HCN
CO2
N2
HCl
CH3Cl
Polar /
Nonpolar
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 12 of 18
Polarity
 Involves sharing electrons (to get an octet = 8 valence e-s)
o equal sharing = nonpolar - no charge; this occurs when both elements have the same electronegativity (the
attraction to an electron) so this occurs when there is a bond between 2 of the same nonmetal for example N2 – both
N have an equal pull on the shared electron pair – so there is no partial charge.
o
unequal sharing = polar – slight charges; this occurs when the two elements have different electronegativities ie you
have two different nonmetals in a covalent bond; the more electronegative element (the one further to the right and up
on the PT) pulls more on the shared electrons and has a slightly negative charge – the other element will be slightly
positive.
Ex: Classify the following as polar or nonpolar bonds: HCl, I 2
You try: classify as polar or nonpolar: F2, NO
* There are some special molecular cmpds – called network solids (or network crystals) that contain extensive covalent bonding
throughout a network of atoms. These compounds are very hard and brittle and have very high melting points ex: graphite and
diamonds These are the strongest bonds!
Polarity Practice
Draw the Lewis Structure.
 If the central atom has any unshared pairs, the molecule is polar.
 If there are no unshared pairs on the central atom, look at the atoms around the central atom.
o If they are all the same, the molecule is nonpolar.
o If any one of them is different, the molecule is polar.
 In a 2-atom molecule, if the bond between the 2 atoms is polar (electronegativity difference is greater than 0.4), then the whole
molecule is considered polar.
Remember, a molecule with more than 2 atoms can have polar bonds between its atoms (based on differences in electronegativity) and
still be an overall nonpolar molecule.
Molecule
H2
BH3
Lewis Structure
Geometry
(Shape Name)
Polar /
Nonpolar
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 13 of 18
Molecule
Lewis Structure
Geometry
(Shape Name)
Polar /
Nonpolar
N2
SO2
NH3
CO2
BOND STRENGTHS:
o
Ionic: Generally the strongest bond type = the larger the ion charges the stronger the ionic bond ie Calcium Chloride, CaCl2 is
stronger than potassium chloride KCl because Ca has 2+ charge involves 2 e- transferring and K has a 1+; only 1 e- transferring
o
Metallic: medium strength between ionic and covalent
o
Covalent: Weakest bond type- (exception network solids like diamonds- strongest bond type)
Ex problem: Classify the following from strongest to weakest bond:
Copper wire, carbon dioxide, lithium chloride (LiCl) and aluminum nitride (Al3N2).
You try: Rank the following from strongest to weakest bond:
Water, potassium iodide (KI), pure gold necklace, and calcium fluoride (CaF2)
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 14 of 18
Review Sheet & Naming Practice
Lewis Dot Structures
1. Write the symbols for the elements and connect them with a line. H is always on the outside. The other elements usually surround the
center one.
2. Add up the valence electrons of all atoms in the compound. Subtract 2 for each line drawn.
3. Arrange the remaining electrons in pairs around the outer elements first so that each atom has an octet.
(H likes 2. B is happy with 6.)
4. Check for happiness (stability)!
5. Change unshared pairs of electrons into extra bonds if necessary.
VSEPR Shapes
Shape
Linear
Bent
Trigonal Planar
Trigonal Pyramidal
Tetrahedral
Trigonal Bipyramidal
Octahedral
Atoms Bound to
Central Atom
2
2
3
Unshared Pairs on Central
Atom
0
1 or 2
0
3
4
5
6
1
0
0
0
Polar or Nonpolar Bonds
Type of Bond
Nonpolar Covalent (Molecular)
Polar Covalent (Molecular)
Ionic
Difference in Electronegativity Values
0.0 – 0.4
0.5 – 2.0
Greater than 2.0
Polar or Nonpolar Molecule
1. Draw the Lewis Structure.
2. If the central atom has any unshared pairs, the molecule is polar.
3. If there are no unshared pairs on the central atom, look at the atoms around the central atom.
a. If they are all the same, the molecule is nonpolar.
b. If any one of them is different, the molecule is polar.
4. In a 2-atom molecule, if the bond between the 2 atoms is polar (electronegativity difference is greater than 0.4) then the whole molecule
is considered polar.
Naming Molecular Compounds
 First, is the compound molecular?
o If the first element is a nonmetal, then the compound is molecular.
 DO NOT USE CHARGES!!! DO NOT SIMPLIFY!!!
 Use PREFIXES to show how many of both elements!
 Don’t write mono- on first element.
 First element keeps its name; 2nd element ends in –ide.
1- mono
2- di
3- tri
4- tetra
5- penta
6- hexa
7- hepta
8- octa
9- nona
10- deca
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 15 of 18
Molecular Naming Practice
Write the Formula
Write the Name
1) diantimony trioxide ________________________
1) CO2 __________________________________
2) carbon monoxide _________________________
2) CI4 __________________________________
3) oxygen difluoride _________________________
3) P2O3 _________________________________
4) sulfur hexachloride _______________________
4) N2O3 _________________________________
5) sulfur tetraiodide _________________________
5) Cl2O7 ________________________________
6) dicarbon octafluoride ______________________
6) BI3 ___________________________________
7) dichlorine heptaoxide _____________________
7) S2Cl4 _________________________________
8) tetraphosphorus decaoxide _________________
8) NO ___________________________________
9) pentaphosphorus nonachloride ______________
9) CS2 ___________________________________
10) sulfur trioxide __________________________
10) I2O5 _________________________________
11) dihydrogen tribromide ___________________
11) N2O4 ________________________________
12) nitrogen trisulfide _______________________
12) SiBr4 ________________________________
13) heptachlorine monophosphide _____________
13) P6F9 _________________________________
14) fluorine monosulfide ____________________
14) Br2Cl _________________________________
15) hexabromine dinitride ___________________
15) I8P ___________________________________
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 16 of 18
Bonding Basics Review
Element
Atomic
symbol
Total # of
Electrons
# of Valence
Electrons
# of Electrons
Gained or Lost
Oxidation Number
(charge)
Bromine
Lithium
Calcium
Sulfur
Boron
Silicon
Phosphorus
2. Ionic Bonds - Draw the Lewis dot diagrams for each atom, draw arrows to show the transfer of electrons, write the charge for
each ion, and then write the chemical formula. We will do A together!
(A) Potassium + sulfur
(B) Magnesium + Oxygen
(C) Lithium + Nitrogen
3. Covalent Bonds – Draw the Lewis structures for each atom, draw lines to show the electrons that are shared, and then write
the chemical formula.
(A) Fluorine + Fluorine
(B) 3 Hydrogen + 1 Phosphorus
(C) 2 Hydrogen + 1 Sulfur
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 17 of 18
Properties of ionic, covalent and metallic bonds:
Strongest bond type: __________________ remember this involves a transfer of electrons.
Because this is the strongest bond type: ionic compounds tend to have:
_______________ melting points and boiling points (it takes more energy for these compounds to change states of matter).
Are ______________________ - they will break down into ions in water and will conduct an electric current.
Ionic compounds tend to be brittle.
Weakest bond type: __________________ remember this involves sharing electrons.
Because this is the weakest bond type: covalent bonds tend to have:
_______________________ melting points and boiling points (it requires less energy for these compounds to change states of matter).
Are ______________ - they will not conduct an electric current in water.
Covalent bonds may be polar (unequal sharing – partial charges) or nonpolar (equal sharing – no charges).
Review on Bonding
Fill in the table to the
right.
Bond Type
Type of atoms
One Property
One example
2. Explain why the oxidation number of Magnesium is 2+ in ionic compounds (It is recommended you use an electron configuration
diagram in your explanation)
For the following use:
3. KBr
C) Covalent Bond
I) Ionic Bond
M) Metallic Bond
4. Copper wire
5. NaCl
6. NH3
7. Br2
8. Which of the above would have the highest boiling point(s)?
9. Which of the above have the weakest bonds?
10. which of the above is soluble in water?
11. Which of the above is nonpolar covalent?
12. Which of the above is polar covalent?
13. Which of the above would be malleable and ductile?
Answers:
1. Refer to Pg 1 with exact same table!
2. Mg electron configuration = 1s2 2s2 2p6 3s2 – has 2 valence electrons – can gain 6 e-s or lose 2 e-s to have 8 (octet rule = 8 e-s
full outer level = very stable) Mg will lose its 2 val electrons resulting in a positive 2 charge.
3. I
4. M
5. I
6C
7C
8 KBr and NaCl ionic
9. NH3 and Br2 – covalent
10. KBr and NaCl – ionic
11. Br2
12. NH3
13. Cu wire
Covalent Bonding/Nomenclature/VSEPR/Polarity Packet Page 18 of 18
Problem Set- Show all work
1. What is the wavelength and energy type when a Hydrogen atom jumps from n = 5 to n = 3?
Is Energy released or absorbed?
Wavelength:_________________
Energy type:________________
E released or absorbed:________________
2. We are taking a class field trip to Europe to practice the factor-label conversion method. For each question show your work and
include all units.
a. You are filling out the passport application using a pencil containing lead 0.7 mm in diameter. What is the diameter in cm?
____________
b.
The class is flying from GSO to London. If the trip is 3200 miles, what is the distance in km?
____________
c.
In meters?
____________
3. Describe the 3 intermolecular forces (IMFs), include their relative strengths.
4. Why would sugar (C12H22O11) melt before salt (NaCl)?
5. Write the Lewis Structures for each of the following. Identify as polar or nonpolar.
a.
water
c. sulfuric acid
b.
ammonia
d. carbon tetrachloride