chapter 3. section 3.1 (pg. 78-84) homework: lewis symbols pg. 82 #2 – 4 pg. 84 # 2, 4, 5, 7-10...
TRANSCRIPT
Chemical BondingChapter 3
Bonding Theory & Lewis FormulasSection 3.1 (pg. 78-84)
Homework: Lewis Symbols Pg. 82 #2 – 4 Pg. 84 # 2, 4, 5, 7-10
Objectives:1) Define valence electron, electronegativity, and ionic bond
2) Use the Periodic Table and Lewis structures to support and explain ionic bonding
3) Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions.
Bonding Theory: Valence Electrons & Orbitals
To describe where electrons exist in the atom, chemists created the concept of an orbital.
Orbital – region of space around an atom’s nucleus where an electron may exist
An “orbital” is not a definite race track, it is a 3-D space that defines where an electron may be (like a rain drop in a cloud)
For bonding study we are only concerned with an atom’s valence orbitals (the volume of space that can be occupied by electrons in an atom’s highest energy level)
WHY? Bonding only involves valence e-’s because lower energy levels are held so strongly by their positively charged nucleus
FYI Read pg. 78-79 for the history on Bonding Theory
According to bonding theory, valence electrons are classified in terms of orbital occupancy.
(0 = empty, 1 = half filled , 2 = full)
An atom with a valence orbital that has a single electron can theoretically share that electron with another atom Such an electron is called a BONDING
ELECTRON
An atom with a full valence orbital (2 e-’s), repels nearby orbitals and wants to be alone Such a pairing is called a LONE PAIR
Bonding Theory: Valence Electrons & Orbitals
1. The first energy level has room for only one orbit
- can only hold 2 e-’s max
2. Energy levels above the first have room for four orbitals = 8 electrons max
- Noble gases have this structure; their lack or reactivity indicates that eight electrons filling a valence orbital is very stable
(Remember the OCTET RULE)
The Four Rules of Bonding Theory 2e-
2 p+
He
8e-
8e-
2e-
18 p+
Ar
2e-
2e- 2e- 2e- 2e-
FYI: Only C, N, O, and F atoms always obey the octet rule when bonding
EXCEPTIONS:B = stable with 6 valence e- (3 orbitals)P = stable with 10 valence e- (5 orbitals)S = stable with 12 valence e- (6 orbitals)
The Four Rules of Bonding Theory 3. An orbital can be unoccupied, or it may contain one or two electrons – but never more than two (Pauli Exclusion Principle)
4. Electrons “spread out” to occupy any empty valence orbitals before forming electron pairs
Never more than 2e- in an orbital
“Aluminum has three half-filled orbitals and one vacant orbital.” How would you describe Sulfur?
Atomic Models: LEWIS SYMBOLS(aka Lewis Dot Diagrams, Electron Dot Diagrams, LDD, Lewis Models
• Named after Gilbert Lewis who is responsible for the Octet Rule. He
reasoned that all atoms strive to be like the nearest noble gas.
• Used dots or ‘x’ to represent the valence electrons
• The inner electrons and the nucleus are represented by the element symbolHow to draw Lewis Symbols:
1. Write the element symbol2. Add a dot to represent each valence e-
3. Start by placing valence e-’s singly into each of the four valence orbitals (4 sides)
4. If additional e-’s need to be placed, start filling each of the orbitals with a second e- up to 8
Q: Which element has 4 bonding e-’s? Which has 3 lone pairs and 1 bonding e-?
PracticeDraw the Lewis Symbols for the elements in Period 3
For each one indicate how many lone pairs or bonding electrons are present
It is important to remember that the Lewis symbols do not mean that electrons are dots or that they are stationary.
The four sides represent the four orbitals that may be occupied by electrons; it is a simplistic 2-D diagram of a complex 3-D structure
Electronegativity- A measure of the force that an atom exerts on electrons of other atoms; (the “pull” on bonding
electrons)
- Each atom is assigned a value between 0.0 – 4.0; the larger the number
the greater the “pulling” force
- Example: Fluorine has an EN = 4.0 and francium has an
EN = 0.7- This means fluorine wants to pull on other electrons very
strongly- This means francium doesn’t want to pull on other
electrons
Q: Does lithium (EN = 1.0) want to lose or gain an electron to be stable?
Q: Does fluoride (EN = 4.0) want to lose or gain an electron to be stable?
Do you see any relation to their electronegativity numbers?
So how do we assign each atom an electronegativity number?
a) The farther away from the nucleus that electrons are, the weaker their attraction to the nucleus
b) Inner electrons shield valence electrons from the attraction of the positive nucleus
c) The greater the number of protons in the nucleus, the greater the attraction for more electrons
EN = 0.8
EN = 2.6
Cesium's valence electrons are not held as tightly by its nucleus because the atom is larger
EN = 0.8
EN = 3.0
Potassium’s valence electrons are not attracted to its nucleus as much as Nitrogen’s valence electrons because their are more inner electrons present in K
1 e-8e-8e-2e-19p+K
5e-2e-7p+N
EN =1.9
EN = 3.0
14p+Si
35p+Br
Bromine has more protons (+ charge) which attracts the negative charge of electrons more so than silicon’s 14 protons
.
In this 3-D image, the electronegativity scale is vertical.
Q: Which element has the highest EN? Give three reasons why?
Electronegativity
Q: What is the EN trend within a period and a group?
Why do we care about electronegativity??
Imagine that two atoms, each with an orbital containing one bonding electron, collide in such a way that these half-filled orbitals overlap.
As the two atoms collide, the nucleus of each atom attracts and attempts to “capture” the bonding electrons of the other atoms
A “Tug of War” over the bonding electrons occurs
Which atom wins?By comparing the electronegativities of the two atoms
we can predict the result of the contest = 3 different types of bonds result
Covalent BondingBoth atoms have a high EN so neither atom “wins”
The simultaneous attraction of two nuclei for a shared pair of bonding electrons = covalent bond
EN difference can be zero = Cl – Cl EN = 3.2 EN = 3.2
EN difference can be small = H - Cl EN = 2.2 EN = 3.2
This is called a polar covalent bond – because one side pulls on the electrons more but we will learn more about this in Section 3.3
Cl2 = diatomic
Ionic BondingThe EN of the two atoms are quite different
The atom with the higher EN will remove the bonding e- from the other atom electron transfer occurs
Positive and negative ions are formed which electrically attract each other
EN = 0.9
EN = 3.2
Metallic Bonding Both atoms have a relatively low EN so atoms
share valence electrons, but no actual chemical reaction takes place
In metallic bonding:a) e-’s are not held very strongly by their
atomsb) the atoms have vacant valence orbitals
- This means the electrons are free to move around between
the atoms and the (+) nuclei on either side will attract them
Analogy: The positive nuclei are held together by a glue of
negative e-’s
Metallic bonding visual This diagram represents Mg atoms that have released their electrons and are embedded in a sea (or glue) of electrons.
Note: These metal atoms don`t have to be in a particular arrangement to attract each other therefore they are flexible, malleable and ductile = useful alloys (Brass, Stainless Steel, etc.)
Summary of Bonding Theory:Chemical Bond = competition for bonding electrons
1) Atoms with equal EN = electrons shared equally
If both have high EN = covalent bond (equal = non-polar)
If both have a low EN = metallic bond
2) Atoms with unequal EN = covalent bond (unequal = polar)
3) Atoms with unequal EN = ionic bond
metallic
The nature of chemical bonds changes in a continuous way,
creating a broad range of characteristics.
PRACTICE
Copy pg. 84 – Bonding Theory Summary into your Notes
Pg. 82 #2 - 4
Pg. 84 # 2, 4, 5, 7-10
Section 3.2 (pg. 85-90)
Explaining Molecular Formulas
Pg. 89 #5 (a-f), 6 (a-e)
Pg. 90 #1-4, 6
Section 3.2 (pg. 85-90)
Explaining Molecular Formulas
Objectives:
1) Draw electron dot diagrams of atoms and molecules, writing structural formulas for molecular substances using Lewis structures to predict bonding in simple molecules
2) Illustrate, by drawing or building models, the structure of simple molecular substances
3) Explain why the formulas for molecular substances refer to the number of atoms of each constituent element
Molecular ElementsMany molecular elements are
diatomic and some are polyatomic
You will need to memorize the formulas of the 9 molecular elements as they will not be given to you:
Name Symbol
hydrogen H2(g)
nitrogen N2(g)
oxygen O2(g)
fluorine F2(g)
chlorine Cl2(g)
iodine I2(g)
bromine Br2(g)
phosphorous
P4(g)
sulfur S8(g)
Why are they diatomic?Remember fluorine has 7 valence e-’s and
needs 1 more e- to be stable?Well 2 fluorine atoms could obtain a stable
octet of e-’s if they shared a pair with each other
F - F F2
Remember: This is a simplified 2-D version, not where the electrons actually are
In structural formulas, lone pairs
are not shown
Diatomic ElementsWhat about oxygen and nitrogen?
Each oxygen atom only has 6 valence electrons. So by sharing 2 electrons each, the two oxygen atoms can create a full octet. This creates a double bond
O = O
Each nitrogen atom only has 5 valence electrons. So by sharing 3 electrons each, the two nitrogen atoms can create a full octet. This creates a triple bond
N Ξ N
Would sharing only 1 electron each work?
Molecular CompoundsBackground:
Molecular compounds have covalent bonds (shared electrons) between non-metals and non-metals
Can be solid, liquid or gas as SATP
May or may not be soluble in water (more later)
Don’t ever conduct electricity - even when (aq)
Generally have lower m.p. and b.p than ionic compounds
Review from Section 1.5 Notes
Molecular CompoundsBackground:
Empirical Formulas – show the simplest whole number ratios of atoms in a compound Very useful for IONIC compounds
Formula Unit – the ratio of ions that repeats in a pattern within the crystal; the chemical formula of ionic compounds represents the formula unit
Not useful for MOLECULAR compounds
Na242Cl242 Na16Cl16 NaCl
CH C2H2 acetylene C6H6 benzene C8H8 octene
All are extremely different compounds but the empirical formula would be the same
Molecular CompoundsBackground:
Molecular Formulas – a molecular formula shows the actual number of atoms that are covalently bonded to make-up each molecule
We use this because chemical formulas for molecular compounds result from sharing electrons, therefore a variety of compounds are possible (which we determine empirically through experiments) Often the symbols are written in a sequence that helps you determine how
the atoms are bonded
C2H4O2 CH3COOH
Empirical formula: CH2O incorrect for molecular
compounds
Review of Molecular Compound Formulas
SeePg. 88
Q: How do we know how molecular compounds bond?
(Aka: How do we draw Lewis Formulas?)
Where do these come from???
Determining Lewis Formulas Bonding Capacity: the maximum number of single
covalent bonds that an atom can form
REMINDER: How many e-’s does an atom want in its valence energy level to be satisfied?
H = 2 e- C, N, O, F, P, S, Cl, etc. = 8 e-
REMINDER: What types of covalent bonds are possible?
F – F single = sharing one e- pairO = O double = sharing two e- pairsN Ξ N triple = sharing three e- pairs
So why do we care about bonding capacity?If we know how many bonding e-’s an atom has, we can
predict what structure a molecular compound will have
Determining Lewis Formulas
Atom Number ofvalence electrons
Number ofbonding electrons
Bonding capacity
carbon 4 4 4
nitrogen 5 3 3
oxygen 6 2 2
halogens 7 1 1
hydrogen 1 1 1H
I.e. Carbon can form 4 single bonds, 2 double bonds, 1 triple and 1 single, or 1 double and 2 singles
Determine the Lewis formula and structural formula for sulfur trioxide, SO3(g)
1. Count the number of valence electrons there are in total? (If polyatomic ions are included, subtract or add electrons to account for the net charge)
Oxygen = 3 atoms x 6 valence e-’s each = 18 valence e-’s Sulfur = 1 atom x 6 valence e-’s each = 6 valence e-’s
24 valence e-’s
Lewis Formulas- Guided Ex. #1
2. Choose your central atom In our course, we will limit our formulas to ones with one
central atom (unless extra info is provided)
So how do you know which is the central atom? Usually the one in lesser quantity (SO3(g))
OR The one with the higher bonding capacity
Carbon usually – because 4 is the highest bonding capacity
So which is the central atom?
Lewis Formulas – Guided Ex. #1
3. Arrange peripheral atoms around central atom and place one pair of valence e-’s between them
4. Place lone pairs on all peripheral atoms to complete their octet
5. Place any remaining valence e-’s on the central atom as lone pairs.
For this example, all 24 have been assigned
Lewis Formulas – Guided Ex. #1
6. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new position between the peripheral and central atom.
7. Show the structural formula but omit lone pairs and replace every bond with a line(Count the electrons around each atom to confirm the octet rule. Each atom should have 8 e-’s around it; exception: H)
Lewis Formulas – Guided Ex. #1
Determine the Lewis formula & structural formula for the nitrate ion, NO3
-
1. Count the valence electrons (*look for a net charge if an ion).
nitrogen = 1 x 5 valence e-’s = 5
oxygen = 3 x 6 valence e-’s = 18 23 + 1 (b/c net charge is -1) = 24
Lewis Formulas – Guided Ex. #2
N
2. Which is the central atom? Nitrogen (in lesser quantity)
3. Arrange peripheral atoms around central atom and place 1 pair of valence e-’s between them
4. Place lone pairs on all peripheral atoms to complete their octet
5. Place any remaining valence e-’s on the central atom as lone pairs.
6. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new position between the peripheral and central atom.
7. If the entity is a polyatomic ion, place square brackets around the entire Lewis formula and then write the net charge outside the bracket on the upper left.
Lewis Formulas – Guided Ex. #2
N N
N
PracticePg. 89 #5 (a-f), 6 (a-e)
Watch 5 (f) there is an exception noted. The central atom does not follow the octet rule.
We will go through these answers as a class.
Pg. 90 #1-4, 6
The theory presented today is not absolute – there are exceptions. But rather than presenting a more detailed theory, your textbook will always note such exceptions.
Tomorrow...Molecular Model InvestigationThought Lab InvestigationMorse Code Assignment
VSEPR TheorySection 3.3 – Part A
Pg. 91-96
Objective:
1) Apply VSEPR theory to predict molecular shapes
Stereochemistry – is the study of the 3-D spatial configuration of molecules and how this affects their reactions.
The shape of molecules is determined by the repulsion that happens between electron pairs
The theory behind molecular shapes is called VSEPR Theory (Valence Shell Electron Pair Repulsion)
Molecular Shapes
Solid = in plane of page Dashed = behind (away) Wedge = ahead (toward)
General Rule: Pairs of electrons in the valence shell of an atom
stay as far apart as possible because of the repulsion of their negative charges
The type, number and direction of bonds to the central atom of a molecule determine the shape of the resulting molecule.
So how do we predict these molecular shapes?
VSEPR
We will be using the following compounds to analyze the 6 shapes possible
BeH2(s), BH3(g), CH4(g), NH3(g), H2O(l), HF(g)
To start, draw a Lewis formula for each of the molecules and then consider the arrangement of all pairs of valence electrons.
(Remember – all pairs of valence e-’s repel each other and want to get as far apart as possible)
Using VSEPR to Predict Molecular Shapes
Shape #1 = LinearLewis Formula
Bond Pairs
Lone
Pairs
Total Pair
s
General Formula
Electron Pair Arrangement
Stereochemical Formula
2 0 2 AX2 linearX – A – X
linear
Be
• This Lewis formula indicates that BeH2(s) has two bonds and no lone pairs on the central atom.
• VSPER theory suggests that the two bond pairs will be farthest apart by moving to opposite sides to a bond angle of 180°
• This gives the molecule a linear orientation
* A is the central atom; X is another atom
*Exception* Beryllium does not follow OCTET RULE
Shape #2 = Trigonal PlanarLewis Formula
Bond Pairs
Lone
Pairs
Total Pair
s
General Formula
Electron Pair Arrangement
Stereochemical Formula
3 0 3 AX3trigonal planar
• This Lewis formula indicates that BH3(g) has three bonds and no lone pairs on the central atom.
• VSPER theory suggests that the three bond pairs will be farthest apart by moving to a bond angle of 120° to each other.
• This gives the molecule a trigonal planar orientation.
* A is the central atom; X is another atom
B
*Exception* - Boron Does not follow OCTET RULE
Draw the Lewis Formula for BF3
Practice
Does not obey the octet rule
Trigonal Planar
F
F
F
Shape #3 =TetrahedralLewis
FormulaBond Pairs
Lone
Pairs
Total Pair
s
General Formul
a
Electron Pair Arrangement
Stereochemical Formula
4 0 4 AX4tetrahedr
al
• This Lewis formula indicates that CH4(g) has four bonds and no lone pairs on the central atom.
• VSPER theory suggests that the four bond pairs will be farthest apart by arranging in three dimensions so that every bond makes an angle of 109.5° with each other.
• This gives the molecule a tetrahedral orientation.
* A is the central atom; X is another atom
Draw the Lewis Formula for SiH4
Practice
H
H
H
H
Tetrahedral
Shape #4 =Trigonal PyramidalLewis
FormulaBond Pairs
Lone
Pairs
Total Pair
s
General Formul
a
Electron Pair Arrangement
Stereochemical Formula
3 1 4AX3
Etetrahedra
lTrigonal
pyramidal
• This Lewis formula indicates that NH3(g) has three bonds and one lone pair on the central atom.
• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)
• But the lone pair is very repulsive, thus pushes the atoms more to a 107.3° bond angle
• This gives the molecule a trigonal pyramidal orientation.
* A is the central atom; X is another atom, E is a lone pair of electrons
Draw the Lewis Formula for PCl3
Practice
Cl
Cl
Cl
Trigonal pyramidal
Shape #5 =Angular (Bent)Lewis
FormulaBond Pairs
Lone
Pairs
Total Pair
s
General Formul
a
Electron Pair Arrangement
Stereochemical Formula
2 2 4AX2E
2
tetrahedral Angular
(Bent)
• This Lewis formula indicates that H2O(l) has two bonds and two lone pairs on the central atom.
• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)
• But the TWO lone pairs are very repulsive, thus pushes the atoms more to a 105° bond angle
• This gives the molecule an angular (bent) orientation.
* A is the central atom; X is another atom, E is a lone pair of electrons
Draw the Lewis Formula for OCl2
Practice
Angular (bent)
Shape #6 =Linear (Tetrahedral)Lewis
FormulaBond Pairs
Lone
Pairs
Total Pair
s
General Formul
a
Electron Pair Arrangement
Stereochemical Formula
1 3 4 AXE3
Linear(Tetrahedra
l)
• This Lewis formula indicates that H2O(l) has two bonds and two lone pairs on the central atom.
• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)
• But since there are only two atoms with one covalent bond holding them together, by definition, the shape is linear, as is the shape of every other diatomic molecule.
* A is the central atom; X is another atom, E is a lone pair of electrons
FH
Draw the Lewis Formula for HCl
Practice
VSEPR theory describes, explains, and predicts the geometry of molecules by counting pairs of electrons that repel each other to minimize repulsion. The process for predicting the shape of a molecule is summarized below:
Step 1: Draw the Lewis formula for the molecule, including the electron pairs around the central atom.
Step 2: Count the total number of bonding pairs (bonded atoms) and lone pairs of electrons around the central atom.
Step 3: Refer to Table 7, and use the number of pairs of electrons to predict the shape of the molecule.
Summary
Pg. 95
Draw the Lewis and stereochemical formulas for a sulfate ion, SO4
2- and predict the shapeSee pg. 95
Draw the Lewis and stereochemical formulas for a chlorate ion, ClO3
- and predict the shapeSee pg. 96
On your own: Pg. 96 #3
Practice
It is important to remember that a double or triple bond is one bond, and to treat it as such, when predicting the VSEPR shapes of molecules.
Example: Predict the shape of C2H4(g)
Draw the Lewis formula for the molecule
Count the # of pairs of e-’s around the central carbon atoms. The carbon atoms have 3 bonds (2 single, 1 double) and no
lone pairs. This is the same as a trigonal planar configuration.
Practice: Predict the shape for C2H2(g).
Multiple Bonds in VSEPR Models
H
H
H
H
Answer: See pg. 97
1) Finish pg. 96 #1-3
2) Pg. 98 #6-7 (Multiple Bond Practice)For 7 c, d, e - If there is more than one central atom
involved, tell me the shape around each of the central atoms Example:
3) Pg. 104 #1, 2, 3 #2: If there is more than one central atom involved, tell
me the shape around each of the central atoms
Homework
trigonal planar—first two carbonstetrahedral—third carbon
Draw the Lewis Formula for PCl3
Practice
Molecular PolarityDipole Theory
Section 3.3 – Part BPg. 98 - 104
1) Determine the polarity of a molecule based on simple structural shapes and unequal charge distribution
2) Describe bonding as a continuum ranging from complete electron transfer to equal sharing of electrons.
PolarityChemists believe that molecules are made up of charged
particles (electrons and nuclei).
A polar molecule is one in which the negative (electron) charge is not distributed symmetrically among the atoms making up the molecule. Thus, it will have partial positive and negative charges on
opposite sides of the molecule.
A molecule with symmetrical electron distribution is a nonpolar molecule.
The existence of polar molecules can be demonstrated by running a stream of water past a charged object. Demo: See Figure 9
TESTING A LIQUID WITH A CHARGED OBJECT:
In a liquid, molecules are able to rotate freely.
Polar molecules in a liquid will rotate so that their positive sides are closer to a negatively charged material.
Near a positively charged material they become oriented in the opposite direction.
EMPIRICAL RULES FOR POLAR AND NONPOLAR MOLECULES
Type Description of molecule
Examples
Polar AB diatomic with different atoms
HCl(g), CO(g)
NxAy containing nitrogen and other atoms
NH3(g), NF3(g)
OxAy containing oxygen and other atoms
H2O(l), OCl2(g)
CxAyBz containing carbon and two other kinds of atoms
CHCl3(l), C2H5OH(l)
Nonpolar
Ax all elements Cl2(g), N2(g)
CxAy containing carbon and only one other kind of atom (except CO(g))
CO2(g), CH4(g)
When the water test was repeated with a large number of pure liquids, this provided the set of empirical rules above.
Pg. 99
PREDICTING AND EXPLAINING POLARITY
Linus Pauling explained polarity by creating the concept of electronegativity. Introduced in Section 3.1
Electronegativity increases as you go up or to the right on the periodic table.
Pauling explained the polarity of a covalent bond as the difference in electronegativity of the bonded atoms.
If the bonded atoms have the same electronegativity, they will attract any shared electrons equally and form a nonpolar covalent bond.
If the atoms have different electronegativities, they will form a polar covalent bond.
The greater the electronegativity difference, the more polar the bond will be.
For a very large electronegativity difference, the difference in attraction may transfer one or more electrons resulting in ionic bonding.
PREDICTING AND EXPLAINING POLARITY
Cl2(g)
We use the Greek symbol delta to show partial charges
PREDICTING AND EXPLAINING POLARITY Pauling liked to think of chemical bonds as being different
in degree rather than different in kind.
According to him, all chemical bonds involve a sharing of electrons, with ionic bonds and nonpolar covalent bonds being just the two extreme cases
The bonding in substances therefore ranges anywhere along a continuum from nonpolar covalent to polar covalent to ionic.
For polar covalent bonds, the greater the electronegativity difference of the atoms, the more polar the bond.
~EN difference:nonpolar (<
0.4)polar (0.4+)ionic (m +
nm)
PRACTICESee pg. 100 Sample Problem 3.4
Try on your own pg. 100 #9 (a-c), 10, 11(a only)
DOES BOND POLARITY = MOLECULAR POLARITY?? NO! Chemists have found that the existence of polar bonds
in a molecule does not necessarily mean that you have a polar molecule.
Example: Carbon dioxide is found to be a nonpolar molecule, although
each of the CO bonds is a polar bond. WHY??
According to VSEPR (two bonds, no lone pairs) = linear arrangement
We will start showing bond polarity as arrows pointing in the negative direction (where e-’s want to go) = bond dipole Points from lower to higher electronegativity
The arrows are vectors and when added together, the equal but opposite bond dipoles equal zero.
Non-polar molecules are ones where the bond dipoles balance each other; producing a molecular dipole (vector sum) of zero
δ– δ+ δ–
O = C = O 3.4 2.6 3.4
Prediction Molecular Polarity Step 1: Draw a Lewis formula for the molecule.
Step 2: Use the number of electron pairs and VSEPR rules to determine the shape around each central atom.
Step 3: Use electronegativities to determine the polarity of each bond.
Step 4: Add the bond dipole vectors to determine whether the final result is zero (nonpolar molecule) or nonzero (polar molecule).
Guided Practice #1Predict the polarity of the water molecule.
1) Draw the Lewis formula
2) VSEPR: Draw the stereochemical formula
3) Assign the EN of the atoms, assign δ– and δ+ to the bonds
4) Draw in the bond dipoles
O
H HAngular (bent)
• The bond dipoles (vectors) do not balance. • Instead, they add together to produce a
nonzero molecular dipole (shown in red).• This results in a polar molecule (explains
bending water)
Go to Learning Tip
pg. 102
Guided Practice #2Predict the polarity of the methane
molecule.
1) Draw the Lewis formula
2) VSEPR: Draw the stereochemical formula
3) Assign the EN of the atoms, assign δ– and δ+ to the bonds
4) Draw in the bond dipoles• Notice how all the bond dipoles point into the central atom.
• There are no positive or negative areas on the outer part of
the molecule. • A tetrahedral molecule is symmetrical in 3-D and
four equal tetrahedral bond diploes always sum to zero
Tetrahedral
PracticePredict the bond polarity of the
ammonia, NH3(g) molecule. Include your reasoning.
Answer: See pg. 102
FYI: DO YOU REMEMBER “LIKE DISSOLVES LIKE”?
Means:“Polar substances are soluble in polar
substances; Non-polar substances are soluble in non-polar substances”
Mixing non-polar and polar substances results in them forming layers, with the least dense one on top.This occurs because polar molecules attract
each other more strongly; thus they stay close together excluding nonpolar molecules
Two clear liquids formed layers in this tube: nonpolar hexane C6H14(l)on top, and polar water, H2O(l) below.
Nonpolar dark orange liquid bromine, Br2(l) was then added. The bromine dissolves much more readily in the nonpolar hexane.
WHY DO WE CARE ABOUT POLAR MOLECULES?
Cleaning!! Water (polar) is useless at removing oil (nonpolar) so detergents are artificially created molecules that overcome this problem
Detergents have long, nonpolar sections which are attracted to (dissolve in) a tiny oil droplet.
The polar end of each of these detergent molecules helps form a polar “layer” around the droplet, which attracts polar water molecules.
This allows them to pull the oil droplet away from a stained area of fabric and hold it suspended in the wash water.
Homework: Pg. 102-103 #13-16
Pg. 104 # 4, 5, 10
Intermolecular Forces
Section 3.4 Pg. 105-117
1) Explain intermolecular forces, London (dispersion) forces, dipole-dipole
attractions and hydrogen bonding
2) Relate properties of substances to the predicted intermolecular bonding in the
substance.
BACKGROUND• All chemical changes (reactions) are accompanied by energy
changes
▫ Energy is mostly heat, light, or electrical energy
▫ Energy can be released slowly (battery) or quickly (fireworks)
▫ Two types of energy changes are possible:
EXOTHERMIC – energy is released into the surroundings - the product’s bonds have less energy than the reactant’s bonds
ENDOTHERMIC – energy is absorbed from the surroundings - the product’s bonds have more energy than the reactant’s
bonds
▫ Bond Energy – the energy required to break a chemical bond or the energy released when a bond is formed
BACKGROUND• There are three types of forces in matter:
1) Intranuclear force (bond) – bonds within the nucleus between protons and neutrons (very strong)
2) Intramolecular force (bond) – bonds between atoms within the molecule or between ions within the crystal lattice (quite strong)
3) Intermolecular force (bond) – bonds between molecules (quite weak); are electrostatic (involve positive and negative charges)
There are 3 types of intermolecular bonds:
a) Dipole-Dipole Forces (a.k.a. Polar Forces)
b) London Force (a.k.a. London Dispersion Force, Dispersion Force)
c) Hydrogen BondingNote: “Van der Walls force” – includes London and dipole-dipole forces
Weakest
Medium
Strongest
1) Dipole-Dipole Force• The simultaneous attraction between
oppositely charged ends of polar molecules.
▫ Simply put, the attraction between diploes
▫ Dipole-dipole forces are among the weakest intermolecular forces, but still control important properties (i.e. Solubility because water is polar))
Dipole: a partial separation of positive and negative charges within a molecule, due to electronegativity differences
1) Dipole-Dipole ForceIn a liquid, polar molecules can move and rotate to maximize attractions and minimize repulsions. The net effect is greater overall attraction.
The strength of the dipole-dipole force is dependent on the overall polarity of the molecule
Note: If a molecule is polar it will be soluble in water? Why?
1) Dipole-Dipole Forces In a liquid: In a solid:
2) London ForceSimultaneous attraction between a momentary dipole
in a molecule and the momentary dipoles in surrounding molecules
momentary dipole: an uneven distribution of electrons around a molecule, resulting in a temporary charge difference between its ends
They last for just the instant
that the electrons are
not distributed perfectly even.
2) London Force• Fritz London also showed that momentary dipoles
occurring in adjacent molecules would result in an overall attraction
• The strength of the London force is directly related to the number of electrons in the molecule, and inversely related to the distance between the molecules.
▫ Increase electrons = Increase force (directly related)..▫ Increase distance = Decrease force (inversely related)
2) London Force• The key point is that:
▫ the more electrons a molecule has, the more easily momentary dipoles will form, and the greater the effect of the London force will be.
• London forces are present between all molecules, whether any other type of attraction is present.
Why do we care about intermolecular forces?We can use Dipole-Dipole and London Forces to predict Boiling
PointsCompound (at SATP) Electrons Boiling Point (°C)
CH4(g) 10 -164
SiH4(g) 18 -112
GeH4(g) 36 -89
SnH4(g) 54 -52
Remember (if all other factors are equal):
1) The more polar the molecule = The stronger the dipole-dipole force
2) Increase the number of electrons = Increase the strength of London Force
A higher boiling point temperature means more energy has to be added, thus we assume the intermolecular
forces are stronger. (see Learning Tip pg. 109)
Example #1• Use Intermolecular force theory to predict which of the
following hydrocarbons has the highest boiling point:
▫ methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10)
1) Are the molecules polar or non-polar? non-polar (no dipole-dipole force)
2) Which has more electrons? butane: greatest # of e-’s = greatest London force
Check:Alkane Boiling Point (°C)
methane -162
ethane -89
propane -42
butane -0.5
Example #2• Use Intermolecular force theory to predict which of the following
has the highest boiling point:
▫ bromine (Br2 ) or iodine monochloride (ICl)
1) Which has more electrons?
They are isoelectronic: have the same number of electrons (70 e-’s)
-Therefore the London force is the same (or nearly the same)
2) Are the molecules polar or non-polar?
-Bromine is non-polar (has no dipole-dipole force; only London forces)
- Iodine monochloride is polar (has dipole-dipole forces and London forces)
- This extra attraction between ICl molecules produces a higher boiling point
Check: Substance Electrons Boiling Point (°C)
bromine 70 59
iodine monochloride
70 97
You cannot predict boiling points when:One molecule has a stronger dipole-dipole force and
the other has a stronger London force
The two molecules differ significantly in shape
The central atom of either molecule has an incomplete octet
PracticePg. 109 #1-4
4) a) Boron is stable with 6 valence e-
b) chloromethane (CH3Cl)
3) Hydrogen Bonding• Occurs when a hydrogen atom bonded to a
strongly electronegative atom, (N, O and F) is attracted to a lone pair of electrons in an adjacent molecule.▫ Hydrogen nucleus (proton) is simultaneously
attracted to two pairs of electrons; one closer (in the same molecule) and one further away (on the next molecule)
Why do you need a strongly
electronegative atom?
It pulls the hydrogen’s electron away making it
“unshielded”, so the lone pair on the other side can come much
closer
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3) Hydrogen Bonding• Hydrogen bonds are momentary attractive
forces between passing mobile molecules but are the strongest of the intermolecular forces.
• Hydrogen bonds only act as continuous bonds between molecules in solids, where the molecules are moving slowly enough to be locked into position.
• Hydrogen force would have been a better name.
3) Hydrogen BondingIn ice, hydrogen
bonds between the molecules result in a regular hexagonal crystal structure.
The ···H– represents a hydrogen nucleus (proton) being shared unequally between two pairs of electrons
3) Hydrogen Bonding• Do lakes freeze from the
bottom-up or the top-down?
• Top–down, because water is unique in that its solid form (ice) is less dense than its liquid form. Why??
• The hydrogen bonds hold water molecules in a hexagonal lattice with open space in the center, which explains the low density (mass/volume) of ice.
Hydrogen Bonding in DNA• FYI: The double helix of the
DNA molecule owes its unique structure largely to hydrogen bonding.
• The red bonds are hydrogen bonds.
• If the helix were held together by covalent bonds, the DNA molecule would not be able to unravel and replicate and life could not continue!!
• Explains surface tension, shape of a meniscus, volatility and capillary action
1) Surface Tension
▫ Molecules within a liquid are attracted by other molecules in all directions equally, but right at the surface, molecules are only attracted downwards and sideways. This means the net pull is downward so the surface tends to stay intact
▫ The stronger the intermolecular force the stronger the surface tension.
Why do we care about intermolecular forces?This shows water adhering to the faucet gaining mass until it is stretched to a point where the surface tension can no longer bind it to the faucet.
It then separates and surface tension forms the drop into a sphere.
2) Capillary Action – due to adhesion (attraction between unlike molecules) and cohesion (attraction of like molecules)
▫ The adhesion between water and glass is greater than the cohesion between water molecules.
▫ The cohesion between mercury molecules is greater than the adhesion between mercury and glass
Why do we care about intermolecular forces?
In a sense, water is pulled up the tube by
the intermolecular forces between
water and glass
Hg clip
Meniscus
PracticePg. 117 # 1, 4, 5
#1 – use pg. 99 table to determine polarity#1 – look for NH2, NH, OH2, OH, to determine if hydrogen
bonding is possible
Ex. CH3CHOHCH3 will it have hydrogen bonding?