Chapter 2

Question 1

1. State the major evidence for classifying light as a wave phenomenon. Use the de Broglie equation to show that matter has wavelike properties.

Answer 1

Lambda = h/mass*velocity

• Young’s double slit experiment

Question 2

2. Describe the electromagnetic spectrum and calculate wavelength from frequency or frequency from wavelength. Also be able to calculate the energy of a photon from its frequency or wavelength.

Answer 2

• Left to Right. Gamma rays, X - Rays, UV Light, Visible Light ( Blue - red) , Infrared, Microwave, Radio Waves
• Gamma Rays are smallest wavelength, Radio waves are the largest wavelength
• Gamma Rays are the highest frequency, Radio waves are the lowest frequency (Wavelength and frequency are inversely related)
• Wavelength*Frequency = Speed of light Change in Energy = h*frequency => h*c/wavelength

Question 3

3. Given the energies of a set of energy levels, predict the wavelengths or frequencies of the lines in a line spectrum. Alternatively, calculate the energy level separations from the wavelength or frequency of the lines in a line spectrum.

Answer 3

Practice Problems

Question 4

4. State the major evidence for classifying light as a particle.

Answer 4

• Photoelectric effect

Question 5

5. Explain the term “quantized” as it is applied to energy in this chapter.

Answer 5

• Energy can be gained or lost only in integer multiples of h*frequency.
• Each of these small “packets” of energy is called a quantum. A system can transfer energy only in whole quanta.
• = atoms and molecules can’t accept just any amount of energy, but instead only the amount of energy that exactly matches the size of these quanta.

Question 6

6. Describe the photoelectric effect and discuss what was learned from a study of this effect.

Answer 6

Refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it.

(a) Light with frequency less than the threshold frequency produces no electrons.
(b) Light with frequency higher than the threshold frequency causes electrons to be emitted from the metal.

Question 7

Summarize the series of electromagnetic radiation for the names

Answer 7

• Lyman: Emit to n = 1, Large transitions, Uv light
• Balmer: Emit to n= 2, Medium transitions, Visible light
• Paschen: Emit to n= 3, small transitions, Infrared “

Question 8

7. Explain what an atomic (line) spectrum is and describe the kind of information that it gives about the energy levels in an atom.

Answer 8

• A spectrum in which spectral line can be distinguished.
• Because each spectral line can be clearly distinguished, match each line to the transition between energy levels responsible for it.
• Changes in energy between discrete energy levels in hydrogen will produce only certain wavelengths of emitted light,

Question 9

8. Distinguish between a continuous spectrum and a line spectrum.

Answer 9

• Continuous spectrum results when white light is passed through a prism. This spectrum contains all the wavelengths of visible light.
• Line spectrum: emission spectrum, only shows wavelengths for which the photons were emitted

Question 10

9. Distinguish between an emission and an absorption spectrum.

Answer 10

• Absorption lines are where light has been absorbed by the atom thus you see a dip in the spectrum
• Emission spectra have spikes in the spectra due to atoms releasing photons at those wavelengths.
• EX: Hydrogen for example, indicates that only certain energies are allowed for the electron in the hydrogen atom.

Question 11

10. Explain how emission spectra and absorption spectra of the same element are related.

Answer 11

the lines in the emission spectrum of a substance are at the exact same wavelengths as the lines in the absorption spectrum of the same substance.

Question 12

11. State the Rydberg formula and explain its relationship to the spectrum of hydrogen atoms.

Answer 12

• Use Rydberg equation to calculate the wavelength of light emitted when energy is emitted from n final to n initial
• Look up formula - use exmple

Question 13

12. State the Bohr model of the hydrogen atom and explain its significance (work on this answer).

Answer 13

Key things to remember:

• Energy levels are stationary states
• Change in energy casues by electrons between stationary states
• When electron move, absorbing electron (increase in energy), or emitting energy (decreasing energy)

Question 14

13. Explain the relationship of a line in a spectrum to the atom’s energy levels and photons emitted by that atom.

Answer 14

The lines on the atomic spectrum relate to electron transitions between energy levels, if the electron drops an energy level a photon is released resulting in an emission line and if the electron absorbs a photon and rises an energy level an absorption line is observed on the spectrum.

Question 15

14. State the Heisenberg Uncertainty Principle and explain its meaning.

Answer 15

• The exact position of the electron is never known, which is consistent with the Heisenberg uncertainty principle:
• it is impossible to know accurately both the position and the momentum of a particle simultaneously.
• Think about how small the mass of an electron is - this helps to explain it’s uncertainty

Question 16

15. Describe wave‐particle duality.

Answer 16

Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well.

Question 17

16. Discuss the probability interpretation of an atomic orbital.

Answer 17

• The square of the wave function indicates the probability of finding an electron near a particular point in space.
• The maximum in the curve occurs because of two opposing effects.
  
  • The probability of finding an electron at a particular position is greatest near the nucleus, but the volume of the spherical shell increases with distance from the nucleus.
• Therefore, as we move away from the nucleus, the probability of finding the electron at a given position decreases, but we are summing more positions.
• Thus the total probability increases to a certain radius and then decreases as the electron probability at each position becomes very small. “

Question 18

17. Give the allowed values and names of the quantum numbers n, l, and ml and relate how each affects the energy, shape, orientation, and average radius of the orbital.

Answer 18

Make the table and make the filling in electron shell oribitals

Question 19

18. Sketch the shapes and orientations of s, p, and d orbitals.

Answer 19

Sketch the shapes

Question 20

19. Describe the Pauli Exclusion Principle.

Answer 20

no two electrons in a given atom can have the same set of quantum numbers

Question 21

20. Use the Aufbau principle and Hund’s rule to diagram an atom’s or ion’s electron configurations.

Answer 21

• Aufbau: As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogenlike orbitals.
• Hund: For an atom with unfilled subshells, the lowest energy is achieved by electrons occupying separate orbitals with parallel spins, as far as allowed by the Pauli exclusion principle. “

Question 22

Exceptions to Aufbau

Answer 22

Cr and Cu - both valence shell configurations move one from the 4s into the 4p - making them stable as half full

Question 23

Atomic radius (nuclear charge are the same thing)

Answer 23

• Down a group - radius increases, the size gets bigger, and the electrons further away = less effective nuclear charge
• Left to right on a period - radius decreases, adding electrons to the same shell, but also adding protons, the protons are better able to pull on the the eletrons

Question 24

Atomic radius (nuclear charge) - Ions

Answer 24

• An ion has the same number of protons
• Nuetral atom < Anion anions have more electrons - aka a fuller electron cloud
• Nuetral atom > Cation B/C cations got rid of their valence shell, made the entire atom smaller, made electron cloud smaller

Question 25

Ionization energy

Answer 25

• X => X- How tightly an atom is holding its electrons
• Left to right, inoization energy increases
• Going down a group ionization energy decreases

Question 26

Successive Ionization Energy

Answer 26

Overall trend - successive ionization energy increases, but BIG jump at core electrons

Question 27

Exceptions to Ionization Energy Trends

Answer 27

Phosphorus and sulfur are neighboring elements in Period 3 of the periodic table and have the following valence electron configurations: Phosphorus is 3s^2 3p^3 , and sulfur is 3s^2 3p^4 . Ordinarily, the first ionization energy increases as we go across a period, so we might expect sulfur to have a greater ionization energy than phosphorus. However, in this case the fourth p electron in sulfur must be placed in an already occupied orbital. The electron–electron repulsions that result cause this electron to be more easily removed than might be expected.

Question 28

Electron affinity

Answer 28

is the energy change associated with the addition of an electron to a gaseous atom

Question 29

What is group 1 on the periodic table?

Answer 29

“Alkali metals - *1s in valence shell *likely to form+1 cations “

Question 30

What is group 2 on the PT?

Answer 30

“Alkaline Earth Metals *s^2 valence shell electrons *likely to form +2 cation”

Question 31

What is group 17 on the PT?

Answer 31

“Halogens *7 valence electrons s^2,p^5 *likely to become -1 anion “

Question 32

What is group 18 on the PT?

Answer 32

“Nobel gases *Full outershell”

Question 33

What are groups 3-12 on the PT?

Answer 33

“Transition Metals *d block “

Question 34

Why are the values for emission negative?

Answer 34

Energy at infinity corresponds to the electron being infinitely far away from the nucleus. When it is bound to the nucleus at lower values of n, the atom is more stable, thus lower energy, which must then be negative.

Question 35

What is the periodic table mostly made of?

Answer 35

Metals!