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PChem 1110 > Unit 2 > Flashcards

Unit 2 Flashcards

1
Q

Bohr’s Model

A

Light is passed through a prism, separating it into wavelengths of visible light.

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2
Q

What does an emission spectrum tell us?

A

It represents the type of visible light emitted by an atom when its electrons are excited.

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3
Q

Quantization

A

An energy transfer wherein it can only be in integer values of the elementary unit of energy, the particle.

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4
Q

Planck’s Basic Assumptions

A

a. The energy transferred through light-matter interactions is proportional to the frequency of the wave, expressed as E = h (Planck’s constant) x v (frequency of the wave)
b. Energy transfer only occurs in integer multiples of the elementary unit of energy E, also expressed as the photon.

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5
Q

Planck’s Constant

A

6.626 x 10^-34 J

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6
Q

Einstein’s Theory

A

Each wave of EM radiation is composed of ‘quanta’ known as photons, whose energy is dependent on the frequency, as suggested by Planck.

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7
Q

Wavelength

A

Distance between two peaks or troughs, symbolized by lambda and measured in meters or nanometers
1 m = 10^9 nm or 1 nm = 10^-9 m

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8
Q

Frequency

A

Number of waves per cycle, symbolized by nu, and measured as cycles/second

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9
Q

Speed of waves

A

All travel at speed of light, and wavelength is inversely related to frequency

In a vacuum: 3 x 10^8 m/sec

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10
Q

True or false: matter can absorb or emit radiation in different regions of the electromagnetic spectrum.

A

True; it just does different things.
microwave: molecular rotation
infrared: molecular vibrations
visible light and uv light: electron transition

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11
Q

Quantum Mechanical Model

A

Atoms can exist in a variety of energy levels based on the distribution of the electrons in orbitals. Different electron states are also characterized by different mathematical functions that describe the corresponding electron waves.

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12
Q

Heisenberg Uncertainty Principle

A

We cannot know the exact location or movement of a particle (electron) at any given time

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13
Q

How is wavelength associated to kinetic energy?

A

Shorter wavelength, the higher the kinetic energy

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14
Q

How is an electron’s wavelength related to the space it resides in?

A

The smaller the volume, the shorter the wavelength, the higher the energy.

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15
Q

Can molecules be quantized?

A

Yes; radiation in the UV and visible regions of the EM spectrum induce electron excitation, while lower frequencies in the IR range produce different vibrational states, microwave radiation produces rotational states.

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16
Q

Why does IR absorption happen over a range of frequencies?

A

This is due to the multiple interactions between atoms in the same molecule and between different molecules in the system.

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17
Q

Delocalization

A

An electron ‘detaches’ from its original atom, producing bonds. In this case, the electrons move into the ‘bonding’ region’, allowing other electrons in each atom to get closer to their atomic nuclei, lowering the potential energy.

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18
Q

Wavenumber

A

The inverse of frequency

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19
Q

How does bond strength relate to wavenumber?

A

Stronger bond means a higher frequency. High frequency = high wavenumber

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20
Q

Bonding capacity

A

The amount of bonds a given atom is eligible to make considering the quantity of valence electrons in a given sample. Determined by occupancy limits and energy costs.

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21
Q

What do low dips in IR spectra represent in terms of bond movement?

A

Stretch

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22
Q

How was atomic radius initially measured?

A

Scientists measure the diameter of two bonded atoms and divided it by two using x-ray diffraction

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23
Q

How does atomic radius change throughout the periodic table and why?

A

It increases as you move down, but shrinks as you move right. This is because as you go across table, the number of electron shells doesn’t change, but the number of protons increases, causing the magnetic field to decrease the size of the atom. As you move down, you add the number of electron shells, but not the number of protons, so the atom continues to expand.

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24
Q

How does ionization energy change throughout the periodic table?

A

It increases going right and up. As the size of the atom gets smaller, the valence electrons are held tighter to the nucleus. Thus, you need more energy to remove them, or delocalize them.

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25
Q

Main Factors of Chemical Bonding

A
  1. Bonding is a stabilizing process for atoms with unfilled valence orbitals since some electrons gain access to lower energy states.
  2. The number of electrons that gain access to lower energy states is constrained by the number of unfilled valence orbitals.
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26
Q

Orbitals

A

Electron sub-subshells that consist of two electrons. In atomic subshells s, p, d, and f, there are 1, 3, 5, and 7 orbitals respectively. Orbitals define energy states where electrons can exist and the probability of finding them in those locations.

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27
Q

How to develop electron configurations

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

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28
Q

Pauli Exclusion Principle

A

No two electrons with the same properties can exist in the same orbital. Following this idea, the energy needed to pair the electrons should be smaller than the energy gained by delocalizing the electrons into the bonding orbital.

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29
Q

Covalent Bonds

A

Pairs of molecules become attracted to each other through electron delocalization, induced by a net force of electrostatic attraction between electrons and protons in different atoms.

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30
Q

How do you know which chemical bond is longer? (PE/Distance Graph)

A

On a Potential Energy/Distance graph, longer bonds are bonds where the lowest potential energy on the graph hits further right.

31
Q

How do you know which chemical bond is stronger?

A

Whichever graph has the lowest potential energy, since low potential energy indicates stability

32
Q

Electron density

A

The probability of having an electron in a given space, which increases during bond formation, due to electron delocalization.

33
Q

Why do IR spectra occur over a range of frequencies, not just one specific location?

A

Due to multiple atomic interactions between molecules and bonds between individual atoms, there are often times multiple bonds associated with one IR spectra. Thus, each dip represents a different bonding type.

34
Q

What happens to kinetic energy, repulsive PE, attractive PE, and the PE of other electrons during bonding? Total energy?

A

When electrons delocalize, they occupy a larger space. Thus, their kinetic energy goes down, repulsive PE goes down, attractive PE goes up, and the PE of other electrons goes down, causing total energy to decrease.

35
Q

What happens to potential energy as you add the amount of electrons that are delocalized?

A

Potential energy drops more drastically to demonstrate the stability maintained by more bonds.

36
Q

Are chemical bonds static entities?

A

No, but they require a specific quantity of energy/frequency in order to vibrate around their equilibrium positions as determined by atomic masses and bond strength.

37
Q

How does bond strength affect the energy of a vibrational state?

A

Absorption bands associated with a greater number of bonds or stronger bonds tend to appear at higher wavenumbers, indicating that wavenumber is directly proportional to frequency. At the same time, energy required also depends on mass, and therefore more energy is required to induce vibrational transitions in bonds between lighter atoms.

38
Q

Are chemical bonds quantized?

A

Absolutely! See IR spectra for information

39
Q

Photoelectron spectroscopy

A

A method by which atoms are bombarded with photons of a known energy quantity and measuring the kinetic energy of the electrons ejected. The more weakly the electron is held, the more kinetic energy it will have when ejected.

40
Q

How to analyze a PES

A

Using the shell model anatomy of an atom, determine the size of the peaks by how many electrons are in each shell. Remember that outer electrons have less energy required, while inner electrons require higher energy. Ionization energy decreases from left to right while electron number decreases going down.

41
Q

How does the limit in electron occupancy of orbitals determine molecular structure?

A

Electrons have an intrinsic property that constrains the amount of particles that can simultaneously occupy the same space. This property is electron spin, exemplified by the Pauli Exclusion Principle.

42
Q

Radicals

A

The number of valence electrons is odd, rather than even, and some bonding orbitals are occupied by one single electron. They are extremely reactive, and induce delocalization of electrons from other species. Think Nitrous Oxide

43
Q

Orbital Occupancy

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

44
Q

Octet Rule

A

Often, the most stable structures are those in which each atom has a full valence shell, generally corresponding to 8 electrons

45
Q

How to form Lewis Structures

A
  1. Choose the central atom (tends to have the highest bonding capacity)
  2. Count total valence electrons, and how many pairs of them will be formed.
  3. Use as many pairs as needed to form single bonds between the central atom and surrounding atoms.
  4. Use the remaining pairs to ensure each atom has a full valence shell. Start with outside atoms, but if not H, place any leftover electrons on the central atom.
  5. Ask: how can we maximize delocalization?
46
Q

Sigma bond

A

Head to head overlap

47
Q

Pi bond

A

side-to-side overlap

48
Q

Exceptions to the Octet Rule

A

Radicals, hypervalent, electron deficient

49
Q

Resonance

A

Electrons are delocalized over the entire system, not just one atom, resulting in several stable types of molecules.

50
Q

Line Structures

A
  1. Carbons not shown - assume at bends and ends
  2. Hydrogens bonded to carbons not shown - number determined by bonding capacity of Carbon
51
Q

Valence Shell Electron Pair Repulsion theory

A

Minimizing repulsions allows us to find the most stable shape. thus, regions of high electron density around any single atom will be as far as possible due to repulsions.

52
Q

How to determine number of electron regions?

A

Add bonds plus any lone pairs attached to the center.

53
Q

Linear EDG

A

Each electron domain forms an angle of 180 degrees

54
Q

Linear Molecular Geo

A

Three molecules with 2 electron regions form a straight line.

55
Q

Trigonal Planar EDG

A

3 Electron regions form a 120 degree angle from each other, but not exactly since high electron density creates a strong repulsion.

56
Q

Trigonal Planar Molecular Geo (of a Trigonal Planar EDG)

A

3 electron regions with no additional pairs creates a propeller shaped molecule with 118 degree angles

57
Q

Bent/Angular Molecular Geo

A

A molecular geometry containing 3-4 electron regions, where an ion pair(s) at the central atom of a molecule with three atoms causes the angle between the other two molecules to be smaller, depending on the amount of lone pairs. 116.8 degrees for 3, 104.5 for 4.

58
Q

Tetrahedral EDG

A

4 electron regions form a four-petaled flower of approximately 109.5 degrees and lower, dependent on the quantity of additional lone electron pairs present

59
Q

Tetrahedral Molecular Geo

A

4 electron pair structure shaped like a three-footed Eiffel Tower

60
Q

Trigonal Pyramidal Molecular Geo

A

4 electron pairs form a pyramid with three sides.

61
Q

How are the properties of chemical compounds determined?

A

It’s determined by the way in which valence electrons are distributed among the atoms in bond. It creates a ‘fingerprint’ for the molecule.

62
Q

How are partial charges formed?

A

Electrons are not shared equally in a covalent bond due to varying electronegativities. This creates higher electron density in more electronegative atoms.

63
Q

Why are some atoms more electronegative than others?

A

It has to do with the relative energy of the unfilled valence electron shells. Since a more full shell indicates lower energy and more stability, they are more likely pull electrons from weaker atoms.

64
Q

Electronegativity is

A

the measure of the decrease in total energy of an atomic system when a bonding electron is associated in the vicinity of the atom. The larger decrease in energy, the more electronegative the atom is.

65
Q

Electronegativity is ______ proportional to ionization energy.

A

directly

66
Q

What does electronegativity tell us about energy transfer?

A

It tells us which atoms are more likely to be more stable when getting a new bonding electron. It also tells us the energy cost associated with losing an electron.

67
Q

Calculation of partial electric charge

A

(x(A)-x(B))/[.5(x(A) +x(B))]
i.e. the difference of electronegativities divided by the average electronegativity

68
Q

Electric dipole

A

The presence of partial positive and partial negative charges separated by a bond. Its effect on its surroundings is dependent upon the magnitude and direction of the dipole moment

69
Q

Magnitude of the dipole moment

A

magnitude of the partial electric charge x distance of the bond length
-tells us the polarity of the bond

70
Q

Can there be polar bonds and still have a non-polar molecule?

A

Yes; since polar bonds are a vectorial quantity, consider the direction of the vectors using the head-to-tail method. If the vectors (same direction and magnitude) connect into a closed shape, the molecule is non-polar.

71
Q

What is a molecule’s interaction with photons dependent on?

A

Bond polarity and overall molecular polarity.

72
Q

Can a molecule have no electronegativity differences and still be polar with electron density differences?

A

Yes; think O3

73
Q

What is electrostatic potential?

A

The amount of energy required to put a unit positive charge in a particular space. The more positively charged, the more energy required.

PChem 1110 (4 decks)

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