Chapter 8: Molecular Shapes
• Draw the Lewis Structure for Carbon tetrachloride
Lone pairs of electrons: ones not involved in bonding
Bonded pairs of electrons: ones involved in bonding
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html# Go to Bond Types
Get into groups of 2 using your Venn Diagram write down all that you know about ionic and covalent bonds. Be sure to include: ­Difference in the naming system
­Elements on periodic table involved
­How electrons behave
­Differences in Lewis Structures
­Properties of each
Review: Orbital shapes will play a role in the molecular shape.
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Single Covalent Bonds­ sigma bonds ( )
­occurs when the pair of electrons is centered between the two atoms.
­Sigma bonds from when an s orbital overlaps with another s orbital or a p orbital or 2 p orbitals overlap.
How do you know if you have an s or p orbital?
­orbital diagrams/electron configurations in outer energy levels
Some molecules have a full octet if they share more than 1 pair of electrons with one or more atoms­this requires multiple covalent bonds.
­Carbon, Nitrogen, Oxygen, and Sulfur often form multiple bonds with other nonmetals
How do you know if 2 atoms will form a multiple bond?
­The number of valence electrons needed to get to an octet = the number of covalent bonds that can form.
Double bonds ­ 2 pairs of electrons are shared between 2 atoms. Only considered one bond in determining the shape.
Ex: Oxygen O 2
Triple bonds­ 3 pairs of electrons are shared between 2 atoms. Only considered one bond in determining the shape.
Ex: Nitrogen N
2
3
Pi bond ( )­ a multiple covalent bond that consists of one sigma and at least one pi bond.
­Forms when parallel orbitals overlap and share electrons. ­The shared electron pair of a pi pond occupies the space above or below the line that represents where 2 atoms are joined together.
Double Covalent Bond:
Triple Covalent bond:
1 sigma and 1 pi bond
1 sigma, 2 pi bonds
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# Go to Sigma Pi Bonding­Animations
Have you ever rubbed two balloons in your hair to create a static electric charge on them? If you brought the balloons together their like charges would cause them to repel each other. Molecular shapes are also affected by the forces of electric repulsion.
­The shape of a molecule determines many of its physical and chemical properties.
­Electron densities created by overlap of orbitals of shared electrons determines the molecular shape.
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Factors that affect the strength of covalent bonds.
1. Bond Length ­distance between 2 bonded nuclei at maximum attraction.
­as the number of shared electron pairs increases, the bond length decreases
­the shorter the bond length, the stronger the bond
*single bonds are weaker than double bonds
*double bonds are weaker than tripple bonds
2. Energy­ an energy change occurs when a bond between atoms in a molecule form or breaks.
­Energy is released when a bond breaks.
­bond dissociation energies­
amount of Energy required to break a specific covalent bond, always a positive value.
­The sum total of bond dissociation energies for all bonds in a molecule is the amount of chemical potential energy in a molecule of that compound.
*Inverse relationship between bond energy and bond length: the smaller the bond length the greater the bond dissociation energy.
*The total energy change of a chemical reaction is determined from the energy of the bonds broken and formed.
­Endothermic
­ greater energy requried to break bonds in reactants than what's released when new bonds from in products.
­Exothermic­ greater energy released when new bonds form in the products than the amount of energy required to break the bonds in the reactants.
HOMEWORK
*pg 247 #7­9, 12 no graph, 13 *Complete pg 100/101 study guide for 8.1/8.
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*Can determine the molecular geometry (shape) once a Lewis structure is drawn.
­use V alence S hell E lectron P air R epulsion model or VSEPR
­based on an arrangement that minimizes the repulsion of shared/unshared electron pairs around the central atom.
Balloon Demo: each balloon is an electron dense region...they will always position themselves as far away from each other as they can get.
­Repulsion between electrons cause the atoms in a molecule to be positioned at fixed angles relative to one another.
Bond Angle­ the angle between 2 terminal atoms and the central atom.
­predicted/supported by VSEPR
­unpaired electrons are also important in determining the shape
­larger orbital
­shared bonding orbitals are pushed together by unshared pairs
Quick Demo:
Using small balloons to represent shared pairs of electrons and larger balloons to represent unshared pairs of electrons, make models to represent HCl, H
2 O, NH 3 , and CH
4 . Tie the balloons together and ask students to notice the shape of the resultant molecules.Use the followig balloon combinations for the models:
HCl: one small, three large
H 2 O: two small, two large
NH 3 : three small and one large
CH 4 : four small
Hybridization
hybrid­ 2 things combined that now has characteristics of both
­During chemical bonding different atomic orbitals undergo hybridization
Ex: CH 4 [He] 2s 2 2p 2 ­hybridize the orbitals
Hybridization­
a process in which atomic orbitals ix and form new identical hybrid orbitals.
Figure 8.19
Ex: Hybrid of CH
4
is sp 3 (1s orbital and 3p orbitals)
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Table 8.6: Re­create it using model kits.
Note: Total pairs of electrons is ONLY considered around the central atom.
Note: Single, Double, and Triple bonds contain only ONE hybrid orbital.
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ml Go to molecular shapes
Snowball Fight!!!
Write the formula of a binary covalent molecule on a sheet of paper. Crumple it up and throw them around the room. Then choose one, draw a model of the molecule, predict the geometry, the hybrid, and bond angle(s).
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Homework
page 264 #61­67
Bellwork: ­Find someone who is about the same height as you. ­Complete pg 92/93 titled "VSEPR Model and Molecular Shape" Classwork: Complete 'Shapes Worksheet' Part 1 with partner
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Bellwork:
Complete 8.4 Study guide
on VSEPR theory
8.5 Electronegativity & Polarity
The type of bond formed during a reaction is related to
each atom's attraction for electrons.
Review:
Electron affinity- a measure of the tendency of an
atom to accept an electron.
TREND:
Electronegativity- (assigned value) the ability of an
atom to attract electrons in a chemical bond.
TREND:
Main Idea Activity: Have two students help with a quick demo. Ask both students to pull on a rope with
equal strength. Tell the class the rope represents a shared pair of electrons. Ask the class what this
represents when atoms share electrons. (equal sharing of electrons) Ask one student to pull harder than
the other student. the second student should be pulled toward the first student. Ask the class what
this represents when atoms share electrons. (unequal sharing of electrons) Have students identify which
atoms have a greater tendency to gain electrons. (the ones with the greatest pull on the electrons) Ask
students what type of bond is represented if the electron is completely pulled away from one atom. (ionic
bond)
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A chemical bond is NEVER completely ionic or covalent.
-Electronegativity is used to determine bond character
(type of bond) based on the difference between two
atom's electronegativity values.
-Identical atoms will have the same values in
electronegativity due to the electrons being shared
equally so it will be a nonpolar covalent bond, aka a
covalent bond or aka a pure covalent bond.
-If a difference in electronegativity values exist,
electron pairs in a covalent bond are not shared equallythese will be polar covalent bonds.
-If there is a really large difference between
electronegativity values (>1.7) then it will be ionic because
the electron(s) will be relocated to another atom.
*a difference of 1.7 means 50% covalent, 50% ionic...more
than this means the bond is ionic
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concepts_in_motion.html# Bond Types
Practice! Let's complete the
electronegativity transparency
worksheet-pg 98/99 with your neighbor.
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Note: Covalently bonded molecules are either polar or nonpolar.
-Nonpolar: not attracted by electric field (no overall charges)
-Polar: attracted by an electric field, uneven electron density
Polar Covalent Molecules: tug of war where the 2 teams do NOT
have equal strength
-shared electron pair(s) are pulled more towards one atom (more
electronegative one)
( ) Delta symbol is used to represent a partial charge
- (partial negative) side where electrons spend more
time/
pulled farther to that side
-more electronegative atom
- (partial positive) side where electrons are not at as often
-less electronegative element
*We add partial charges to molecular models to show polarity
-the resulting polar bond is referred to as a dipole (two poles
with partial pos/neg charge)
Polarity & Molecular Shape
Ex: H2O & CCl4
*Note: Both have polar covalent bonds
*Note: symmetrical molecules are usually nonpolar, asymmetrical
are polar as long as bond type is polar.
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Solubility- physical property, defined as the ability of a substance
to dissolve in another substance.
-Bond type/shape of molecules present determine the solubility
*Polar molecules/Ionic compounds are soluble in polar substances.
*Nonpolar molecules dissolve only in nonpolar substances.
'Like Dissolves Like'
Ex: Oil/Water
**Let's review the properties of Ionic compounds­pull out your venn diagram.
The differences in propertes are a result of the differences in attractive forces.
­Covalent bonds between atoms in molecules are strong BUT attractive forces between molecules are weak.
­Weak attraction forces are called Intermolecular Forces aka Van der Waals forces.
­vary in strength
­weaker than bonds that join atoms in a molecule or ions in ionic compounds.
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Types of Intermolecular Forces
1. Between nonpolar molecules ­called dispersion forces (London) or induced dipoles
*weakest type 2. Between oppositely charged ends of 2 polar molecules (pos end aligns with neg end) ­dipole­
dipole
*the more polar the stronger the dipole/dipole
3. Hydrogen end on one dipole and F, O, N atom on another dipole ­hydrogen bonds (STRONG forces)
Properties relate to weak intermolecular forces.
­Low melting points. (sugar melts, salt doesn't)
­Many molecular substances exist as gases or vaporize easily.
­Soft solids (paraffin in candles) ­Form crystal lattices in solid phase like ionics but with less attraction between particles.
Covalent Network Solids
­solids composed of only atoms interconnected by a network of covalent bonds.
­typically brittle, nonconductors of heat/electricity, extremely hard.
*Ex: quartz & diamond
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HOMEWORK
*pg 270 #68­77 *and Read the ‘how it works’ on page 271
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Bellwork:
complete pages 94/95
‘Electronegativity and Polarity’
*Pull out your 'Shapes Worksheet'
from Friday/Monday.
*Complete Part 2
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HOMEWORK:
complete Study guide 8.5 pages 104/105 on electronegativity and polarity
Bellwork: Work with someone you usually don't. Get into a group of 4 and compare answers to your homework.
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Try one yourself!
1. BH
3
2. Nitrogen trifluoride
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Try one yourself!
1. C 2 H 4
2. Carbon Disulfied
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Try one yourself!
1. NH
2. ClO
4
+
4
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Resonance & Exceptions to the Octet Rule
* It is possible to have more than one correct Lewis structure when a molecule or polyatomic ion has both a double and a single bond.
See example:
Resonance­ a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.
*Resonance Structures differ only in the position of the electron pairs, atom position does not change.
­Each molecule or ion that undergoes resonance one structure.
behaves as if it has only ­Bond lengths show that the bonds are identical to each other. Shorter than single bonds, but longer than double bonds. Actual bond length is an average of the bonds in the resonance structures.
­O 3 , NO 3 ­ , NO
resonance.
2
­
, SO
3
2­
and CO
3
2­
commonly form 20
Octet Rule Excepons
(do not obey octet rule)
Coordinate covalent bond­ forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons to form a stable electron arrangement with lower potential energy. 1. A small group of molecules might have an odd number of valence electrons.
2. A few compounds form sub octets or stable configurations with fewer than 8 electrons, rare
Example: BH 3
3. Central atoms contain more than 8 electrons, we call this expanded octets.
HOMEWORK
Resonance Practice Problems: 43­45 pg 258
Octet Rule Exceptions Practice Problems: 47­48 pg 260
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Bellwork:
Complete 8.3 Study guide
'Molecular Structure'
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