Chem Unit 1

Black body radiation

• Light emitted by a perfect black object when heated

• Study by Max Planck

Electromagnetic Spectrum

A range of electromagnetic waves

Wavelength

The length of a wave (λ)

Frequency

The number of cycles per unit of time (f)

Wave Equation

c = λf

λ = 3.0 × 10^8 m/s

f = s^-1 (Hz)

Photoelectric Effect

EM waves can give electrons enough energy to leave a substance. The shorter the wavelength, the greater the energy, therefore more electrons are able to leave.

Plancks equation

E = hf

E = energy of a quantum

h = Planck’s constant (6.63x10^-34)

f = frequency

Continuous Spectrum/Visible Spectrum

Splitting white light into its constituents using a prism (ROYGBV).

Bright Line Spectrum/Emission Spectrum

Shows specific quanta an element emits when heated up using a spectrophotometer.

Absorption Spectrum

Shows specific quanta an element absorbs when in cold gas form.

Bohr’s conclusions

Elements can only absorb and emit certain wavelengths of energy.

#1) Electrons do not emit energy when grounded

#2) Each orbit is an energy level - the further away the orbit, the higher the energy - electrons can jump to higher energy levels by absorbing more energy

What is happening here?

What did Balmer study?

He calculated the amount of photon energy released when electrons jump to ground state.

Balmer Series

Visible light - electrons jumps from any shell down to 2nd, producing visible light - violet is produced when there is a larger jump

Lynmann Series

(UV) - electron jumps down from any orbit to 1st

Paschen Series

(Infrared) - electron jumps down from any orbit to 3rd

How do you calculate the energy of an electron?

E_n = (-2.18x10^-18)/n²

n = shell number

How do you calculate the amount of energy released when an electron jumps down an orbit?

E = E_n - E_l

E = equivalent to E=hf

E_n = higher shell

E_l = lower shell

What formula relates wavelength and frequency?

c = λf

c = 3×10^8 m/s

What was Louis de Broglie’s Atomic theory?

Electrons do not move in circular orbits but in wave patterns

What was Heisenberg’s Atomic theory?

It is impossible to know the position and speed of an electron simultaneously.

What was Schrodinger’s Atomic theory?

He determined electron probability density (orbitals).

What are the main differences between the bohr model versus the quantum mechanical model?

Bohr Model: Protons and neutrons are in the nucleus, electrons orbit the nucleus in energy levels

Quantum Mechanical Model: Protons and neutrons in nucleus, electrons have translational movement around the nucleus in an energy level in a space called an orbital or probability density.

What is the Principal Quantum Number?

n = energy level

How do you determine the total number of:

a) orientations in an energy level

b) electrons in an energy level

a) n²

b) 2n²

What do each of these stand for?

1s²

1 = shell number

s = type of orbital

2 = number of electrons

What is the electron configuration for shell #3?

3s²3p⁶3d¹⁰

What is the electron configuration for shell #6?

6s²6p⁶6d¹⁰6f¹⁴

What is the electron configuration for Na?

3s1

What is the electron configuration for Cr?

4s2 3d4

How do you draw orbitial diagrams for transitional metals?

1) remove electrons from s and p-orbitals first

2) don’t remove any electrons from d and f orbital if full

3) remove all electrons from s and p-orbitals all at once

How do you determine the following quantum numbers?

a) principal quantum number

b) second quantum number

c) magnetic quantum number

d) magnetic spin number

a) n = energy level/shell #

b) l = type of subshell orbital

s = 0, p = 1, d = 2, f = 3

c) m_l = represents the number of orientations of each subshell

m_l is a number from -l to +l

to determine, look at the position of the last electron in an orbital diagram

d) m_s = the last electron in the orbital diagram

m_s = +1/2 if unpaired

m_s = -1/2 if paired

Intramolecular forces

The bonds between atoms in a molecule

Ionic Bonding

• metal + non-metal

• metal loses valence electron to become stable → + ions (cations)

• non-metal gains valence electrons to become stable → - ions (anions)

What is the Lewis Dot Diagram for this Compound?

LiF

What is the Lewis Dot Diagram for H₂O?

What is the Lewis Dot Structure for H₂O?

Which is the center atom from the following molecules?

a) CH₄

b) C₂H₃O₂

c) H₂O₂

a) C - least number of atoms

b) C - can bond the most

c) O - can bond the most

What is a Co-ordinate Bond?

Also known as a Dative Bond, it is a bond where one atom shares both electrons.

One atom is full while the other requires 2 electrons to be full. Represented by an arrow that points to the atom that needs 2 electrons.

When is a co-ordinate bond not needed?

1) If the center atom is on the 3rd period or lower, they may have more than 8 electrons - they possess empty d-orbitals that allow them to take on more electrons. E.g. P and S

2) Certain center atoms can have less than 8 electrons. E.g. B and Be

How do you draw the Lewis structure for the following polyatomic ion?

H₃O⁺

How do you draw the Lewis structure for the following polyatomic ion?

CO₃^2-

Resonance

Delocalized electrons in pi bonds can move around freely. This means that pi bonds can be switched around between atoms. E.g. NO₃^- has 2 resonance structures

What kind of bond are co-ordinate bonds?

Sigma bonds

Hybridization

A theory developed to explain the bonding and shape of molecular structures. It only occurs in molecules with center atoms.

Valence Bond Theory

Covalent bonds share electrons between atoms as orbitals overlap. All covalent bonds have a sigma bond. In multiple bonds, one bond is a sigma bond and the extras are all pi bonds.

How are pi bonds created?

When unhybridized p-orbitals that are vertical to the sigma bond axial overlap.

Bond Length

The distance between the two nuclei of atoms in a bond. Multiple bonds have more electrons, therefore more electrostatic attraction of the electron to the nuclei, which shortens the bond length. The longer the bond length, the weaker the bond strength.

Bond Strength

The amount of energy needed to break a bond. Pi bonds have less bond strength than sigma bonds. Multiple bonds have more strength than single bonds. Therefore, bond strength is inversely proportionate to bond length.

Electron Domain

The number of lone pairs of center atom + the number of sigma bonds.

• help determine the hybridized orientations

If the E.D. is 5, what is the hybridized orientations?

5 = sp³d

Hybridization

The process of combining orbitals to form new orbitals that are better suited for bonding and explaining molecular shapes. The energy level of the lower orbital gets promoted to an intermediate (balanced) level.

• when asked for what is the hybridization, give the orbital combination e.g. sp²

• Only occurs in compounds with a center atom

When hybridizing p-orbitals, why do you always show the left over orientations?

Because they form pi bonds.

V.S.P.E.R.

Valence Shell Pair Electron Repulsion

Who created VSPER and why?

• Nyholm and Gillespie created this theory to explain molecular geometry

What are the 2 principles of VSPER?

1) Lone electron pairs and bonding pairs (e.g. sigma bonds) in the valence orbital repel each other (least to most repulsion: b.p. b.p. < b.p. l.p. < l.p. l.p.)

2) Lone pairs repel each other more than bonding pairs

Linear

B.p. = 2

L.p. = 0

+ any molecule that doesn’t have a center atom

Triangular Planar

B.p. = 3

L.p. = 0

V-Shape

B.p. = 2

L.p. = 1

VSPER - Double/Triple Bonds

Ignore pi bonds when using VSPER - they don’t affect the shape.

VSPER - More than 1 Center Atom

Look at each center atom individually. Treat as two separate molecules that are attached together.

Types of Intramolecular Forces

1) Ionic Bond (anion + cation) △EN = 1.7 - 4

2) Polar Covalent Bonding (electrons are unequally shared, and are closer to the atom with the greater △EN) △EN = 0.4 - 1.7

3) Pure Covalent (electrons are equally shared between atoms) △EN = 0 - 0.4

Electronegativity

An element’s ability to attract a bonding pair of electrons.

% Ionic Character

How much more ionic it is. The closer the △EN is to 4, the greater the ionic character.

Polar Molecules

Have an overall partial charge

Dipole Vectors

Arrows that represent the direction of the partial (-) charge. Can only draw dipoles if the bond between atoms is polar covalent.

How to determine if a molecule is Non-polar?

A molecule is non-polar if:

1) Center atom doesn’t have lone pairs (linear, triangular planar, tetrahedral, trigonal bipyramid, octahedral, square planar)

2) All bonds are the same

3) All bonds are pure covalent

When do Dipole vectors cancel out?

If the vectors face each other or away from each other. Doesn’t count if they are in an angle. Also doesn’t count if the △EN are different.

Intermolecular forces

Bonds between molecules

Ion-dipole

Occurs between an ion and a molecule with a dipole. Is the strongest IMF because there is a full charge

Hydrogen Bonding

Occurs with molecules that have both an H and F, O, or N. Second strongest IMF because FON have greatest EN.

Dipole-Dipole

The attraction between 2 polar molecules. 3rd strongest because charges are partial.

London Dispersion Force

Forces between non-polar molecules, formed from temporary dipoles as electrons move closer to one atom. The larger the molecule (greater MM), the larger the LDF because more electrons mean more attraction. Is the weakest because the charge is temporary.

Van der Waals force

The sum of all IMF forces in a molecule.

Surface Tension

Strength of molecules at the surface of a liquid (elastic skin)

Cohesion

Intermolecular forces between “same molecules”

Adhesion

Intermolecular forces between “unlike molecules”

Network Covalent Crystal (Giant Covalent/Giant Molecular)

Intramolecular force

• substances made from several atoms that bond covalently in a regular repeated pattern

  • Usually formed with carbon and silicon

Allotropes (Different physical forms) of Carbon

1) Graphite - layers of hexagons of Carbon attracted by LDF, each carbon is only bonded 3 times, each carbon has a delocalized electron which moves from carbon to carbon, making graphite a conductor of electricity

2) Graphene - one layer of carbons in a hexagon, can conduct electricity because each carbon bonds 3 times, thus having a delocalized electron

3) Diamond - each carbon is bonded 4 times, can’t conduct electricity, tetrahedral

Metallic Bonding

Intramolecular force

• the electrostatic attraction between a lattice of cations and delocalized electrons

• electrons move from cation to cation, which is what allows metals to conduct electricity

• cation layers can slide past each other, making metals malleable

Ionic Compounds

Intramolecular Force

• the electrostatic attraction between cations and anions bonds the ion, forming an ionic lattice

• the lattice is not fixed, meaning it simply follows the ration of cations to anions

Molecular Crystals

Intermolecular Forcees

• solid molecular compounds which are attracted to each other by LDF in solid state

• E.g. I₂ (s)

Melting/Boiling Points

Determined by the strength of the intra and intermolecular forces, as they determine how much energy is needed to change states. Intra»Inter

Solubility

Substances dissolve in other substances when they are similar in polarity. “Like dissolves like”

1) Same polarity

2) Same IMF

3) Small molecule size

Conductivity

Something can conduct electricity if:

• they have delocalized electrons e.g. metallic bonding, graphite, graphene

  • mobile ions e.g. aqueous or liquid ionic compound

Volatility

How easily a substance evaporates

• stronger the IMF, the lower the volatility

Linear

Trigonal Planar

V-shape

Tetrahedral

Trigonal Pyramid

V-shape/angular

Trigonal Bipyramid

Seesaw

T-shape

Octahedral

Square Pyramid

Square Planar

NH₄⁺

Ammonium