Electron DotsElectron Dotsand and
VSEPR TheoryVSEPR Theory(mostly Chapter 9)
Metallic Bonding
• In metallic bonding the valence electrons are shared between all the atoms in a positive metal crystal. delocalized “sea” of electrons
metallic bonded materials have good thermal and electrical conduction.
Ionic BondsOccur when the nonmetal takestakes one or more electrons away from a metal.
The nonmetal becomes a negative ion metal becomes a positive ion.
The atoms are held together by their opposite charges.
Ionic Bond Strength
Strength of crystal lattice depends on two factors, sizesize and charge transferredcharge transferred.
Smaller atoms have stronger ionic bonds.
Ex: NaF is stronger than NaCl
Atoms transferring more electrons are stronger. Ex: MgCl2 is stronger than NaCl.
2 e- transferred 1 e- transferred
Covalent Bond“the bonds of nature”
• Shared valence electrons
• Complete outer energy levels
• Molecule has 2 or more nonmetal atoms covalently bond– Carbohydrates, proteins, fats, DNA,
stupendous seven (H2, N2, O2, F2, Cl2, Br2, I2)
How do covalent bonds form?
Attractive forces balancebalance the repulsive forces
e- & e- repulsive
p+ & e- attractive
distance is too great repul. = attract
p+ & p+ repulsive
e- & e- repulsive
Electronegativity
• Electronegativity is the ability of atoms in a molecule to attract electrons to themselves.
• On the periodic table, electronegativity increases as you go…– from left to right across a
row.– from the bottom to the top
of a column.
What types of bonds are they?
MgO, water, Calcium Carbide, Potassium Oxide, Nitrogen trihydride
Electronegativity and Bond Type
Find the difference in 2 atoms’ electronegativies to predict bond type…
• Ionic Bonds: 1.7 or greater
• Polar Covalent Bonds: <1.7 and >0.2
• Pure or Nonpolar Covalent bonds: <0.2
Atom Number of Valence Electrons
Number of Bonding Electrons
Bonding Capacity
Carbon
Nitrogen
Oxygen
Halogens
Hydrogen
Bonding Capacity
Electronegativity Table
Drawing Lewis Dot Structures1. Count the valence electrons.2. Predict the location of the atoms
a. Hydrogen is a terminal atomb. The central atom has the smallest electronegativity.
3. Draw a pair of electrons between the central atom and the surrounding atoms.
4. Use the remaining electrons to complete the octets of each atom. If there are electrons left over, place them on the central atom.
5. If the central atom does not have a complete octet then try double or triple bonds.
a. If the atom has 1, 2, or 3 valence electrons, it doesn’t require an octet.
STEP 1: count the total # of valence e- for all atoms involved in the bonding
Carbon: 1 carbon with 4 valence electrons (1x4) = 4
Chlorine: 4 chlorine with 7 valence electrons (4x7) = 28
CCl4
4+28
=32
CCl4CCl4
STEP 2–place the single atom in thecenter and other atoms around it evenly spaced
CClCl
ClClCCl4
4+28 =32 e-
STEP 3: place the electrons in pairs between the central atom and each non-central atom
C ClClCl
Cl
CCl4
4+28
=32 -8
=24
STEP 4: place the remaining electrons around the non-central atom until each has 8 electrons (H atoms have only 2e-)
CCl
Cl
Cl
Cl
CCl4
4+28 =32
-8
=24 -24
=0
Step 5: If you run out of electrons before the central atom has an octet, form multiple bonds until it does. Example: HCNExample: HCN
Hydrogen- 1 electron
Carbon- 4 electrons
Nitrogen- 5 electrons
TOTAL is 1+4+5 = 10 e-
H:C:N
H:C:N:..
..
H:C:::N:
Drawing Lewis Dot Structures
Draw Lewis Dot Structures for:
PH3
H2S
HCl
CCl4
SiH4
CH2Cl2
Draw Lewis Dot Structures
Cl2
NF3
CS2
BH3
CH4
SCl2
C2H6
BF3
(stop)
Covalent Bond StrengthCovalent Bond Strength• Based on proximity (closeness), also called
“bond length”
Influenced by atom size and number of shared electrons
Smaller is stronger
F2 is stronger than Cl2 is stronger than Br2
F2: 1.43 x 10-10 m single bondO2 1.21 x 10-10 m double bondN2 1.10 x 10-10 m triple bond
Bonding OrbitalsBonding Orbitals
• When atoms bond together, their valence shell electron orbitals overlap
• Overlapping electron orbitals create a bonding orbital an area with a high probability of finding an electron
–Sigma Bonds (Sigma Bonds (σσ))•Orbitals overlap head-to-head•Form first, there’s only 1
–Pi Bonds (π)Pi Bonds (π)•Orbitals overlap side-to-side•Form after sigma bonds
Types of BondsTypes of Bonds
• When atoms form a molecule, their orbitals can form different types of bonds:
Every molecule has one sigma bond, but all subsequent bonds between the same two atoms must have a different
way of connecting so they use pi bonds!
Multiple Covalent Bonds – DoubleMultiple Covalent Bonds – Double
6 valence electrons
6 valence electrons
12 valence electrons
Octet satisfied
More stable and stronger
1 sigma bond1 sigma bond
1 pi bond1 pi bond(lines represent
bonded pairs of e-)
5 valence electrons
5 valence electrons
10 valence electrons
Octet satisfied
More stable and stronger
1 sigma bond1 sigma bond
2 pi bonds2 pi bonds
Multiple Covalent Bonds – TripleMultiple Covalent Bonds – Triple
Molecular ShapesMolecular Shapes
• The shape of a molecule plays an important role in its reactivity.
• Look at bonding and non-bonding electron pairs– You can predict the
shape of the molecule!
What Determines the Shape What Determines the Shape of a Molecule?of a Molecule?
• Electron pairs repel each other.• Assuming electron pairs are placed as far as
possible from each other, we can predict the shape of the molecule.
Valence Shell Electron Pair Valence Shell Electron Pair Repulsion Theory (VSEPR)Repulsion Theory (VSEPR)Valence Shell Electron Pair Valence Shell Electron Pair Repulsion Theory (VSEPR)Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”
Molecular Shape ChartMolecular Shape ChartFormula Dot
Structure
Name of
Shape
Nonbonding e- pairs
Bonding electrons
Polarity Hybridization Bond Angle
BeH2
BF3
CH4
NH3
H2O
Molecular PolarityMolecular Polarity• Molecules can be polar and non-polar.• Imagine you are turning over the 3D models on the table.
Are they still the same when you flip them over?– If yes, then the molecule is non-polar (symmetrical)
Molecular PolarityMolecular Polarity
• Non-bonding electron pair = polar– The free pair pushes
the other atoms away
• Non-polar molecule has equal pull from the same atoms
Non-bonding electron pair
(stop)
Bonding Orbital HybridizationBonding Orbital Hybridization• Electron orbitals mix to
make a new set of bonding orbitals (hybrids)– These have different shapes
than regular atomic orbitals– Requires energy but the
energy is returned during bond formation
2s 2p
2sp3
This occurs to allow more bonds!
new hybridized orbital
Bonds can form here
Hybrid OrbitalsHybrid Orbitals
Consider beryllium:• In its ground state, it
would not be able to form bonds because it has no singly-occupied orbitals.
Hybrid OrbitalsHybrid Orbitals
But by promoting an electron from the 2s to the 2p orbital, it can now form two bonds.
This new hybridized orbital is called 2sp2sp
2s 2sp orbitals
2sp
2s 2pGroup 3A elements make spsp22 hybridized
orbitals
Group 2A elements make spsp
hybridizedorbitals
2s 2p
2sp2
2s2 2p2
Group 4A elements have 4 valence electrons
- need 4 bonds to make an octet
- they will have sp3 hybridization.
2sp3
Endothermic and Endothermic and Exothermic ReactionsExothermic Reactions
• Endothermic ReactionsEndothermic Reactions – the energy – the energy needed to break the bonds is greater than needed to break the bonds is greater than the energy that is released, energy is usedthe energy that is released, energy is used– They feel coolThey feel cool
• Exothermic ReactionsExothermic Reactions – the energy – the energy needed to break the bonds is less than the needed to break the bonds is less than the energy released, energy given offenergy released, energy given off– They feel warmThey feel warm
• Why do some solids dissolve in water but others do not?
• Why are some substances gases at room temperature, but others are liquid or solid?
• What gives metals the ability to conduct electricity, what makes non-metals brittle?
• The answers have to do with …
Intermolecular forcesIntermolecular forces
QuestionsQuestions
2 types of attraction in molecules:
IntIntraramolecular bondsmolecular bonds: (Covalent and ionic) attraction between atoms in a molecule
IntInterermolecular forces molecular forces (IMF): the attraction between molecules
– 1) dipole-dipole – 2) hydrogen bonding – 3) London forces
Intermolecular forcesIntermolecular forces (also called Van der Waal’s forces)
Dipole - Dipole attractionsDipole - Dipole attractions•Dipoles: a separation of charge
•This happens in both ionic and polar covalent bonds
H Cl
+ –
• Oppositely charged dipoles (+δ and –δ) are attracted to each other in a molecule
+ –
+ –
+ –
+ –
Hydrogen BondingHydrogen Bonding
H-bondingH-bonding is a special type of dipole - dipole attraction that is very strong (5x stronger)
– Happens when N, O, or F are bonded to H
– Due to the high electronegativity difference between the H and the other atom
– Compounds containing these bonds are important in biological systems (special!)
London forcesLondon forces• Named after Fritz London, sometimes called
dispersion forces
• London forces are due to small dipoles that exist in non-polar molecules
• Random movement of electrons can sometimes form temporary dipoles
• The resulting tiny dipoles cause attractions between atoms/molecules
This is how non-polar molecules This is how non-polar molecules can form solids and liquids!can form solids and liquids!
London forcesLondon forcesInstantaneous dipole: Induced dipole:
Sometimes the random arrangement of electrons
forms tiny dipoles
A random dipole forms in one atom or molecule, inducing a
dipole in the other
(stop)
IMF Strength and Molar MassIMF Strength and Molar Mass• The sizesize of a molecule (molar mass) affects
the strength of intermolecular forces (IMFs)
• Larger size = stronger forcesLarger size = stronger forces– Because the large molecule has more area and
electrons available for intermolecular attractions such as London Forces
– (this is opposite of covalent bond strength)
Stronger IMFs Weaker IMFs
• Consider the halogens (group 7A) as an example
• F2 and Cl2 are gases, Br2 is liquid, I2 is a solid
– Liquids and solids form when IMFs are stronger
– Since they are further down the group, the atoms are bigger
– Larger mass = stronger IMFsLarger mass = stronger IMFs
IMF Strength and Molar MassIMF Strength and Molar Mass
• Boiling (liquid gas) occurs when there is enough energy to overcome intermolecular attractions
• Boiling point tends to increase down a group, as size of atoms in molecules increases
Boiling Point and IMFsBoiling Point and IMFs
Predicted and actual boiling points
-200
-150
-100
-50
0
50
100
Period
Bo
ilin
g p
oin
t
Group 4
Group 5
Group 6
Group 7
2 3 4 5
This is because the largerlarger atoms/molecules have strongerstronger IMFs
so it takes moremore energy to
break those attractions
higher boiling point!higher boiling point!
What about these? (such as H2O)
Hydrogen Bonds and Boiling PointHydrogen Bonds and Boiling Point
• H2O, HF, and NH3 have particularly high boiling points
• This is because of hydrogen bonds!hydrogen bonds!
• Because they are the strongest IMF, they require more heat energy to break the attraction higher boiling points higher boiling points
Predicted and actual boiling points
-200
-150
-100
-50
0
50
100
Period
Bo
ilin
g p
oin
t Group 4
Group 5
Group 6
Group 7
2 3 4 5
(end)
***Hints for IMF Lab***
• Activity 1, Question 3 asks you to draw the Lewis Dot structures for acetone and ethanol.
• Here are their shapes to help you…
ethanol acetone