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Chemistry Unit 5: Bonding – Ch. 5, 6, and 11(pp. 385-392)
Vocab Quiz/Polyatomic Ion Quiz: Thurs, March 20
VOCABULARY:
ionic bonds
malleable
nonelectrolytes
molecular formula
covalent bonds
ductile
anion
hydrate
Problem Set/Pre-lab Due: Thurs, March 27
molecule
polar covalent
cation
VSEPR theory
metallic bond
nonpolar
polyatomic ion
anhydrous substances
Test: Thurs. March 27
octet rule
electrolytes
empirical formula
FORMULAS/COSTANTS:
Memorize the VSEPR Chart – first 6 shapes!
Polyatomic Ions – p. 3 of Unit Pack
OBJECTIVES: Students will…
 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 describe single, double and triple bonds.
 Be able to describe the three intermolecular forces (Van der Waal’s forces): dipole-dipole, London dispersion and
hydrogen bonding.
 Be able to draw Lewis structures for various compounds and polyatomic ions.
 Be able to apply the VSEPR theory in determining molecular geometry.
 Be able to describe hybridization and give an example.
Unit 5 EQs:
Day 1- Review
 Write the noble gas configurations and dot diagram for Se.
 ID the following as a metal or nonmetal: Ca, O, Sn, H, K, Cl and Li
Day 2 - Review
 Given: O, F and Cl put them in order of
a. Decreasing atomic radius
b. increasing electronegativity
 If Candy cost 3.99 per pound - how many kilograms of candy can you buy with $10.00? (Conversion factor for pounds to
grams in first unit pack.)
Day 3  What is the octet rule?
 When electron pairs are shared equally between two atoms, what type of bond is formed?
Day 4  What type of bond explains water’s abnormally high boiling point? Explain.
Day 5  What type of bonds exist in the following compounds?
a. CH4
c. KCl
b. Na3PO4
d. gold bar
 List the above compounds in order of strongest bond to weakest bond.
Day 6 –
 If the sharing of an electron pair is unequal, what is this sharing called?
 A charged group of covalently bonded atoms is called a(n) ________________________. Give an example of 3.
Day 7 –
 Use the VSEPR model to predict the molecular shape of the following molecules.
a. carbon dioxide (CO2)
b. methane (CH4)
 The VSEPR model of the BF3 molecule results in the same trigonal planar structure as that of the hybridized form represented by
which electron configuration?
** Problem Set located on last page of packet!
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VSEPR & Molecular Geometry
(You must memorize the first six (6) shapes in this table!!)
Molecular Shape
Type of Molecule
AByEz
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
5
0
6
0
Linear
Bent
AB2E
Trigonal Planar
AB3
AB4
Tetrahedral
Trigonal Pyramidal
AB3E
Bent
AB2E2
Trigonal
bipyramidal
AB5
Octahedral
AB6
Bonding notes!
When bonds are formed atoms are following the octet rule – trying to get 8 electrons in its outer level.
This is achieved by losing, gaining or sharing electrons.
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
Bonding Packet
POLYATOMIC IONS – MUST MEMORIZE FOR QUIZ ON THURSDAY!!!
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Acetate
C2H3O2-
Ammonium
NH4+
Carbonate
CO32-
Cyanide
CN-
Hydroxide
OH-
Nitrate
NO3-
Peroxide
O22-
Nitrite
NO2-
Phosphate
PO43-
Sulfate
SO42-
CHAPTER 5 & 6 - CHEMICAL BONDING AND MOLECULAR SHAPES
Answer the following questions as you read Chapters 5 & 6, to learn more about bonding.
1. What are the 3 types of bonds?
2.
Describe each type of bond.
3.
How do you determine whether you have an ionic or covalent bond?
4.
Define polar and nonpolar.
5.
Give at least 3 properties of ionic compounds.
6.
Give at least 3 properties of covalent compounds.
7.
What is a molecule?
8. Define:
a. bond energy
b. bond length
9. Give at least 3 properties of a metallic compound.
10. What is the octet rule?
11. What kind of bonding is found in polyatomic ions?
INTERMOLECULAR FORCES
Fill in the following worksheet.
_____________________________ is the force that holds atoms (molecules) together (i.e. bonds).
_____________________________ are the forces between molecules that hold the molecules together. For example, keeping all the
water molecules in a glass of water “stuck together”. These forces are also called Van der Waals forces. They are generally weaker
forces than bonds.
There are three types of Van der Waals forces. The first type, _____________________________
is present when two polar molecules are attracted to each other. A ____________________________ is created by equal but opposite
charges that are separated by a short distance. An example would be two Hydrochloric acid molecules. The partially negative chlorine of
one HCl molecule is attracted to the partially positive hydrogen of another HCl molecule. The partial charges are represented by a lower
case delta () and the sign.
Hydrogen bonding is the second type of intermolecular force. Hydrogen bonding is a special form of a dipole-dipole force (it is
stronger than a regular dipole-dipole force). Hydrogen bonding is only found in molecules where hydrogen is bonded to highly
electronegative elements (F, O and N). Water is an example of a substance with hydrogen bonding.
The third type of intermolecular force is found in nonpolar molecules. _______________________ occurs when a temporary
instantaneous dipole is formed due to the orbiting motion of the electrons in an atom.
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Ionic Bonds: between a metal and a nonmetal
Can also classify based on electronegativity if the electronegativity difference (∆ EN) is >1.7 = ionic bond
Example:
Using the chart on page 24, Find the ∆ EN for the bond formed between Ca and N
You try: find the ∆ EN for the bond formed between Li and F
Ionic Bonding
 Involves a transfer of electrons: the metal loses e’s to get a + charge and the nonmetal gains e’s to get a – charge
 Held together by electrostatic force (opposites attract!)
 When bonds are formed atoms are following the octet rule – trying to get 8 electrons in its outer level. (exception: H stable with
just 2 valence electrons)
Examples:
Ex 1: Mg and S
Ex 2: Ca and I
You try Al and F
Covalent Bonds: between two (or more) nonmetals
Can also classify based on electronegativity if the electronegativity difference (∆ EN) is <1.7 = covlent bond
Example:
Using the chart on page 241, Find the ∆ EN for the bond formed between O and N.
Example: Show the bonding that occurs between Cl and F.
Example 2: 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.
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Covalent Bonding:
 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 (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 compounds – 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!
Metallic bonds: involves two metals – the electrons are “delocalized” – metal ions plus a ‘sea’ of mobile electrons
MOLECULAR STRUCTURES:
Review: Covalent bond =? Octet rule = ?
H and He are exceptions to octet rule – He is stable with 2 e-s, b/c of their 1s2 shell.
LEWIS STRUCTURE:
Element Symbol = nuclei and inner-shell electrons
Dashes = shared electron pairs in covalent bond
Dots = unshared electrons
Single bond =
Steps for Lewis Structure Problem (and an example)
Example: Draw the Lewis Structure of iodomethane - CH3I
1.
2.
3.
4.
5.
Determine the number and type of atoms.
Determine the total number of valence electrons in the atoms to be combined. (use periodic table)
Arrange the atoms to form a skeletal structure. (First atom is central unless otherwise specified or H is first). Add up the electrons
used to form the bonds (each bond counts as 2 e-s). Subtract the number of electrons used from the total in step 2. Use these
electrons in step 4.
Add electron pairs so that each atom is surrounded by 8 electrons (octet rule) (except H – it can only hold 2 e-s) starting with the
outer atoms and saving the central atom until last.
a. Count the number of e-s in the structure to be sure the number of valence electrons used equals the number available and the octet
(duet) rule is satisfied for all atoms. **See step 5b below if not satisfied.**
Practice: together: ammonia and silicon tetrafluoride
You try: BrI, CH3Br
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Bond Strength? Bond Length?
Need to use multiple bonds when there are not enough valence electrons to complete octets by adding unshared electron pairs (step 5 =
numbers do not match)
5. b. If the octet rule has not been satisfied for all atoms, form multiple bonds (double or triple) by moving electron pair until all atoms
have been satisfied.
Example: (together) CH2O
Practice: carbon dioxide and HCN
POLYATOMIC IONS:
6. Additional Step for Completing Lewis Structures:
when counting valence electrons must include the charge:
 If positive subtract electrons
 If negative add electrons
 Also, brackets must be placed around the structure with the charge of the ion outside the brackets to indicate the gain or loss of e-s in
the total.
Example: nitrate
Lewis Structures Practice Worksheet
Draw Lewis Structures for the following molecules or ions:
1. TeCl2
2. PCl3
3. phosphate
4. I2
5. ICl
6. H2S
7. nitrite
8. carbonate
9. water
10. P2
11. boron trifluoride
12. SF2
13. bromine monochloride
14. ammonium
Practice: sulfate
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VSEPR Theory
valence-shell, electron pair repulsion
Way to predict molecular geometry (shape)
There is a repulsion between valence e- pairs
Steps to Work Problems:
(Be sure to refer to and memorize VSEPR table on page 3 of the packet – first six shapes)
1) Draw Lewis Structure of the Molecule
2) Put molecule in AByEz form, where:
A represents the central atom
B represents the atoms bonded to A (y is the # of B atoms; B may different elements.)
E represents the lone pair e-s on A (z is the # of lone pairs.)
3) Predict the shape based on the AByEz form
Example 1: Use VSEPR Theory to predict the shape of CBr4.
AB4
shape - tetrahedral
Example 2: sulfur difluoride (SF2)
AB2E2
shape – bent
AB6
octahedral
Example 3: SF6
BOND STRENGTHS:
Ionic:
 Generally the strongest bond type = the larger the ion charges the stronger the ionic bond ie Calcium Chloride, CaCl 2 is stronger
than potassium chloride KCl because Ca has 2+ charge involves 2 e- transferring and K has a 1+; only 1 e- transferring
Metallic:
 medium strength between ionic and covalent
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 (Al 3N2).
You try: Rank the following from strongest to weakest bond:
Water, potassium iodide (KI), pure gold necklace, and calcium fluoride (CaF2)
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1. 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
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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).
Metallic bonds involve delocalized electrons or a metal ion plus a “sea” of mobile electrons.
Metallic bonds have high melting points and high boiling points, the solid compound is highly conductive (metal compounds are not
water soluble so they do not make electrolytic solutions with water – but they are the only compounds that conduct a current (and
heat) as a solid.
They are malleable:_______________________________________________________________ and
Ductile :_____________________________________________________________________
Review on Bonding & Lewis Structures
1.
Fill in the following table:
Bond Type
Type(s) of Atoms
involved
Force
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.).
3.
What type of bond(s) is found in lead (II) phosphate { Pb3(PO4)2}?
Properties
One Example
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4. Draw the Lewis structures and predict the molecular geometries for:
a. PCl3
b. Carbonate
Identify the types of bonding using the following: (C) Covalent Bond
5. KBr
(I) Ionic Bond
(M) Metallic Bond
6. Copper wire
7. NaCl
8. NH3
9. Br2
10. Which of the above would have the highest boiling point(s)?
11. Which of the above have the weakest bonds?
12. which of the above is soluble in water?
13. Which of the above is nonpolar covalent?
14. Which of the above is polar covalent?
15. Which of the above would be malleable and ductile?
Answers:
1. Refer to page 2of the packet with the 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 valence electrons, resulting in a positive 2 charged cation.
3. The phosphate (between the Os and P) contains covalent bonds – between the lead and the phosphate is an ionic bond.
4. a) P in the middle, 3 single bonded Cls, 2 dots on P, 6 dots on each Cl = AB 3E1 = trigonal pyramidal
b) C in the middle, 2 single bonded O’s with 6 dots and a double bonded O with 4 dots brackets with a negative 2 at the top
right = AB3 = trigonal planar
5. I
6.
M
7.
I
8.
C
9.
C
10. KBr and NaCl ionic
11. NH3 and Br2 – covalent
12. KBr and NaCl – ionic
13. Br2
14. NH3
15. Cu wire
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Silly Putty Lab
Purpose:
This experiment is a demonstration of a type of chemical reaction called polymerization. The compound formed is a polymer. This type
of reaction takes a small chemical compound (called a monomer) and reacts it with a cross-linking agent to yield a chemical compound
composed of many of the smaller monomers which are now chemically joined together as a larger single molecule. A polymer may be
visualized as a necklace or fishing net. A necklace has links going in two directions. Afishing net has links in all four directions.
Background:
There are many types of polymerization reactions. Some examples of materials made by polymerization reactions include paints, plastics,
foam rubber, insulation, polyester, and nylon among many others. By controlling and modifying variables such as temperature, pressure,
amount of reagents, reaction time and the cross-linking agent used, the properties (hardness, brittleness, strength, etc.) can be changed.
Procedure:
To a ziploc bag (this makes clean up easy) add 2 parts (2 tablespoons) glue to 1 part (1 tablespoon) water stir and then add a tablespoon
(15 mL) of the borax solution. CLOSE THE BAG TIGHTLY. Squish the mixture and mix the contents for ~20 seconds. A gelatinous
blob will form and excess liquid may be present. Remove the blob (this is the silly putty) and knead it 1 to 2 minutes to get rid of the
stickiness. Wash and dry your hands. The putty is now ready to play with. The silly putty can be kept in an air-tight container, such as a
ziploc bag.
Analysis Questions:
1.
2.
3.
How do the physical properties of the glue, water mixture change as a result of adding the sodium borate?
What would be the effect (your thoughts) of adding more sodium borate (Borax) solution?
In your own words: Define a polymer (hint: read the info below!)
Review sheet – Bonding
1. What type of bond is formed between the following?
a. Na and F
c. copper atoms in a wire
b.
S and O
d. K and sulfate ions
2. Describe the following bond types (ionic, covalent, metallic) include the
a. atoms involved
b.
a description of the bond
c.
properties
d.
relative strength
3. Describe polar and nonpolar covalent bonds.
4. Use electron configurations to explain the bond formed between potassium and chlorine.
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5. Use electron configurations to explain the bond formed between Br and F.
6. Identify the bond type in the following and then list from strongest to weakest:
a. Iron bar
b.
hydrochloric acid HCl
c.
table salt NaCl
7. Use electron configurations to explain the bond formed between sodium and oxygen.
8. Fill in the following table:
Element
# of valence
electrons
# of electrons
Lost
OR # of e-s gained
charge
S
N
Li
Sr
Answers:
1. Remember M & NM (or polyatomic ion) = ionic M & M = metallic
NM & NM = covalent
2. SEE the table on page 2 on your packet with this info on it!!!!
3. polar= unequal sharing of electrons in a covalent bond results in a partial charge
Nonpolar= equal sharing of e- in a covalent bond results in no charges
4. K = [Ar]4s1 has 1 valence electron that needs to be lost to result in an octet (8 e-s) in its outer level – so it will lose 1 e- to form K+1. Cl
= [Ne] 3s23p5 has 7 valence electrons, needs to gain 1 to get an octet; so it gains 1 e- to from Cl-1. In the ionic bond that is formed K will
lose 1 electron to Cl to form KCl (their + and – charges will be the force that holds them together).
5. Br = [Ar] 4s2 3d10 4p5 has 7 valence e- needs to gain 1 more for an octet. F =[He] 2s22p5 has 7 valence e- needs to gain 1 more for an
octet. Since both need to gain they will form a covalent bond and share 1 e- to each reach an octet.
6. strongest to weakest: NaCl (ionic)
Iron bar (metallic)
HCl (covalent)
7. use electron configurations to explain the bond formed between sodium and oxygen.
O = [He]2s22 p4 6 val e- 9 2 pairs , 2 singles) needs to gain 2 more es – will gain 1 from 2 Nas (one from each Na atom) to form O -2
Na = [Ne] 3s1 one valence electron, one dot, needs to lose one e-s for an octet – will transfer 1 electron to form Na+. 2 Nas will be
necessary to transfer the 2 electrons that O needs = Na 2O
Bonding Packet
Problem Set- Show all work
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Review:
1. What is the wavelength and energy type when a Hydrogen atom jumps from n = 4 to n = 2?
Is Energy released or absorbed?
Wavelength:______________
Energy type:____________
E released or absorbed:________________
2. Given: S
I Sr
O
A. Place them in order of increasing electronegativity
B. Place them in order of decreasing ionization energy
C. Place them in order of increasing atomic size
3. Explain why the ionization energy is larger for F than for K (– do not just write because it is further up/right)
4. Write the complete electron configuration for Rf. Circle the valence electrons
5. Write the noble gas electron configurations, and circle & count the valence e-s for:
A. Ir
# of valance electrons:_______
B. Ge
# of valance electrons:_______
C. Tl
# of valance electrons:_______
6. Write the noble gas configuration and draw a Bohr diagram and a Lewis dot diagram for:
A. Cs
B. As
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7. Find the average atomic mass of a sample of isotopes where 95 % is N-14, 1% is N -15 and 4 % is N-16. Must show all work to
receive credit!
8. A typical person has about 200 mg of cholesterol per 100 ml of blood. If the total blood volume of an individual is 5.0 liters, how
many grams of total blood cholesterol does the individual contain?
Current Unit Material
9. Draw the Lewis structures for the following and identify its molecular geometry and VSEPR Type (A xByEz):
A. Ozone, O3
B. phosphate ion
10. There are some exceptions to the octet rule; one is when the central atom contains more than an octet. An example is the molecule
SF6, where the formation of sp3d2 hybrids occur. Using the notion of hybridization and diagrams explain how this occurs.
(extra credit)