Chemistry 1A Other Notes

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Guidelines for Drawing Hybridized Structures

Guidelines for Drawing Hybridized Structures

Find the total # of valance e-

Find # of bonds in each atom

Find the electronegativity of each atom

Draw the Lewis dot structure

Find electron groups / effective pairs

Lone pairs = 1 e-group

1 single bond = 1 e-group

1 double bond = 1 e-group

1 triple bond = 1 e-group

Match electron groups to hybridization (TABLE 9.24, PAGE 426 Zumdahl textbook)

2 e-group = sp → Linear

3 e-group = sp² → Trigonal Planar

4 e-group = sp³ → Tetrahedral

5 e-group = sp³d → Trigonal Bipyramidal

6 e-group = sp³d² → Octahedral

Using the Lewis structure, count the:

# of σ sigma bonds (single bonds)

# of π pie bonds (double bonds)

# of π pie bonds (triple bonds)

# of lone pairs (if any) on each atom

Find electron configuration of each atom in the molecule

Draw the energy diagrams (from ground state to hybridized state) Show # of orbitals involved in hybridization and those remaining pure

sp hybridization = 2 orbitals mix -s and pₓ only and pᵧ and p? remain pure

sp² hybridization = 3 orbitals mix –s,pₓ, and pᵧ and p? remains pure

sp³ hybridization = 4 orbitals mix –s,pₓ, pᵧ , and p?

sp³d hybridization = 5 orbitals mix –s,pₓ, pᵧ , p? and 1d and remaining 4 do not mix

sp³d² hybridization = 6 orbitals mix –s, pₓ, pᵧ , p? and 2d and remaining 3 do not mix

Make lobes for each atom separately and place electrons in each lobe (single e-and lone pairs) (This will allow you to see which lobes will overlap and which will not)

sp = 2 lobes

sp²= 3 lobes

sp³= 4 lobes

sp³d = 5 lobes

sp³d²= 6 lobes

SIGMA σ BONDS –FIRST, Superimpose those lobes that make sigma bonds–Lateral

PIE π BONDS –SECOND, Show Sideways overlap of p orbitals for pie bonds-Parallel IMPORTANT: Leave the ORIENTATIONS of Pᵧ and P? orbitals AS IS

COMPLETE STRUCTURE – (combine 11 and 12 above) Label atoms, sigma, pie bonds, lone pairs. Total number of electrons in the lobes should be equal to total valence electrons

SHAPE -Electron Geometry & Molecular Geometry

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4/13/2020 Lecture Notes

Bond energy: Energy required to break a chemical bond

Indicates the strength of a bonding interaction

Ionic bonding: Occurs when an atom that loses

electrons easily reacts with an atom that has high affinity for electrons

Ionic compound: forms when a metal reacts with a nonmetal like NaCl.

Coulomb's Law Determines the energy of interaction between a pair of ions using a equation where:

E = units of joules

r = distance between ion centers in nanometers

Q1 and Q2 = numerical ion charges

Used to determine repulsive energy where two like-charged ions are brought together

A bond will form if the energy of the aggregate is lower than that of the separated atoms.

Distance between two atoms when potential energy is minimal

Covalent bonding: equal sharing of electrons between two identical atoms.

Caused by the mutual attraction of nuclei for shared electrons.

Polar covalent bond: bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. Ex: HF

Electronegativity: ability of an atom in a molecule to attract shared electrons to itself.

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4/20/2020 Lecture notes

Molecular orbital theory ch 9

Lewis structure of oxygen (each e- in O2 is paired)

Liquid o2 > ::O=O::

Liquid o2

Take oxygen gas and cool it to turn it into a liquid.

Oxygen is attracted to the magnetic field

Paramagnetism occurs only when there is an unpaired electron.

There has to be an unpaired electron present in O2, but according to lewis structure, there isn’t one.

Magnetic susceptibility: The force experienced by a substance in the magnetic field.

Weight of a sample in a magnetic field in(A) and out(B)

A-B=# of unpaired electrons

Paramagnetism: there is an unpaired electron present, there’s attraction for magnetic field, sample feels heavier when in the magnetic field.

Diamagnetism: paired electrons, repulsion in magnetic field, lighter weight, (opposite of paramagnetism)

N2 (diamagnetism, 1 unpaired electron, repulsion in magnetic field)

:N=-N:

What is happening? The lewis structure doesn’t show unpaired electron for N2. Answer lies in…

Molecular orbital theory:

Postulates of molecular orbital theory

Electrons

are always moving in a spinning/ rotating around the nucleus

They exist as waves and particles

Resonance: SO4-2 (radio label) the charges keep moving

Also in CO3-2 and NO3-

MOT also explains why some substances are good conductors of electricity or why some are semiconductors or insulators.

Molecular orbital theory is like VSEPR theory (lewis structures)

Probability > molecular orbital (si squared)

MO(molecular orbitals) become fully filled with 2e-

Electronic configuration is the same for both

Wave and particle nature of electrons are observed. (si functions apply)

Energy levels are present and quantized.

Comparison between the VSEPT and MOT

VSEPT: Bonds are localized between 2 atoms

Bonds are created by the overlap of orbitals (s,p,d,f, hybridized orbitals)

Forms only bonds (sigma and pi)

Predict shapes based on a number of regions of electron density

Resonance structure>multiple structures

2d presentation

MOT: Bonds are delocalized (bonds are spread out all over) throughout the molecule

Combine atomic orbitals> molecular orbitals

Forms bonding (pi and sigma) and anti bonding (pi star and sigma star)molecular orbitals

Predicts the arrangement of electrons

Resonance by itself

3d presentation

Wave nature of electrons (si)

LCAO: linear combination of atomic orbitals

When the waves meet in phase: both waves are in sync and merge to produce a bigger wave. This leads to a higher probability of finding the density of the electron. (Constructive interference)

Out of phase: Waves that meet are not in the same motion, it means nodes are now there, no electrons are present, no electron density (destructive interference)

Graphical representation

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4/22/2020 Lecture notes

Bond order relates to the strength relates to bond length and number of bonds.

If there are no bonds then bond order = 0

Bond order= 1 = single bond

The less bonds are weaker but less stable

The more bonds the stronger and more stable

More energy present in triple bonds and less in single bonds

Bond order and bond strength have a direct relationship

Bond order and bond length have a inverse relationship

Border and bond energy(bond dissociation energy) have a direct relationship EXCEPT: depending on atoms. b2 is more f2 is less?

Bond strength is directly proportional to bond energy

Electron configuration for molecular orbital theory

Use pi and sigma, and raised to the power.

Molecular orbital energy diagram for N2 and is it para or di magnetic, and what is the bond order.

First identify valence electrons

Draw the orbital diagrams

N:1s22s22p3 look at valence electrons (2s22p3)

No unpaired electrons means diamagnetic

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5/3/2020 Lecture notes

Chapter 17 spontaniety, entropy and free energy

Rate of chemical reaction: what is it? What are the factors involved?

Speed reaction>rate

Reaction> time>product

How are new bonds formed?

-Activation energy

-proximity/distance(close range)

-Orientation / looking

-Catalyst/ speeds ( lowers the energy of activation, alternate reaction pathways[enzyme])

Concentration: amount

Reactants-> limiting reactant

Temperatures: Kinetic energy increases

Exothermic: G= Gibbs free energy

Free: easily available to molecules>internal energy/bond energy

Gibs free energy: something that’s maintaining bond energy

G(reaction) – G(product) = delta G

^ change in free energy

Spontaniety: when new bonds are formed

Exergonic(spontaneous): heat given off

Endergonic(non spontaneous): heat absorbed from outside

Humans by default are lazy so less work for humans is favorable.

DeltaH(reaction)(enthalpy/change in heat of reaction)=H(product)-H(reactant)

Standard heat of formation= 1 mole of product (found in textbook table)

DeltaH=0

Diatomic molecules (H2, O2, Cl2)

Free metals (uncombined)

DeltaH>

Favorable: exothermic(heat is released to surroundings) [less energy of activation]

Not favorable: endothermic(heat absorbed from surroundings)[large energy of activation]

Entropy: DeltaS (change in entropy)

Chaos(theory of chaos) > create newer bonds

Chaos is basically randomness, and nature prefers that. It is a good chance to produce new stuff.

Example: Leaves on ground> decomposes (cellular beakdown) Abcd> ab+cd>a+b+c+b

Building blocks build and rearrange.

When entropy increases or positive, it is good. We want chaos?

DeltaS= S(product) – S(reactant)

If S S(product)^-S(reactant)v=+ (good)

If S S(product)v-s(reactant)^=- (bad)

Ea=Energy of activation

DeltaH(external heat/ bond disassociation energy)

DeltaG(internal/bond energy

^^ All of those at constant temperature

G(products, lower energy)-G(reactants higher energy) Exothermic/negative/spontaneous

H(products, lower energy) – H(reactants, higher energy) exothermic/negative/favorable

DeltaS= generally for exothermic reactions, entropy increases, it is favorable and positive

Endothermic: N2(g)+O2(g)+180kj > 2NO(g)

DeltaG= positive/nonspontaneous/ Endergonic

DeltaH= positive/non fav/ endothermic

DeltaS= Decreasing/negative

Spontaneous > ball rolls down hill

Nonspontaneous > push ball up

What are the driving forces behind spontaneous processes (coffee loses heat)

Exothermicty > it will be a spontaneous except: melting of ice,

Absorbing heat> endothermic

Entropy:all spontaneous reactions, the entropy increases

Larger probability> more randomness/chaos

Postulates of entropy

Chemical reaction = reactants> products

In what case does entropy increases?

More types of products

More types of states (g>>>>L>>>>s)

Temperature higher(exothermic)

More microstates (different types of molecular arrangements, also includes if ions are formed)

Total # of particles increase

When solutions are formed, salvation of ions, h bonding, dipole-dipole, London dispersion

When the volume increases/ when gas expands

# of moles increase

When larger molecules are formed

Question: list all factors:

H2O(l) >100c> H2O(g)

Less entropy >>More entropy this is favorable

2. H2O(s)>25c>H2O(L)

Spontaneous, favorable and positive entropy

3.NaCl(s)+H2O(l)>Na + Cl

Favorable and positive entropy because entopy of ion is higher

4. CaCO3(s)>>>>CaO(s)+CO2(g)

> higher entropy, more things, change of state, more bonds, more microstates, more moles(total of 2 on products side)

5. 2CO(g) + O2(g)>2CO2(g)

Negative entropy because 3mol>2mol, O2 is in more microstates on reactant side but less microstates in product side.

6.Ag(aq)+Cl(aq)>AgCl(s)

Entropy decrease, aq(particles chaotic) to s(still, vibrating particles), 2 mols reactant 1 mol product, less microstates

Ch 10 phase diagrams

Heat of fusion (solid to liquid) heat of vaporization (liquid to gas) , kj/mol

Polarizability/e- cloud distortion

More viscous(greasy)= more hydrogen carbon bonds