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