Study Guide for Periodic Table Assessment – Chapter 5 Sections 1 & 2
Key Terms to know:
Valence Shell/Electrons / Core Electrons / Lewis Dot NotationSublevels (s, p, d & f) / Orbitals / Energy Levels
Hund’s Rule / Aufbau Principle / Pauli Exclusion Principle
Photoelectric Effect / Photon / Frequency
Wavelength / Quanta / Ground State
Excited State / Emission Spectra / Absorption Spectra
Electromagnetic Spectrum / Electromagnetic Radiation / Duality
Quantum Model / Planetary Model / Heisenberg Uncertainty Principle
Key Concepts:
· The reason that a new theory was needed (beyond Rutherford’s Model) was due to his inability to explain where the atom’s electrons were and why consecutive elements had such wildly different properties.
· Through the photoelectric effect we were able to conclude that light can behave as a wave and as a particle (Light particles are referred to as photons)
o The experiments that led to the observation of the photoelectric effect showed that each frequency of light produced photons with different energies
o The energy of a photon is directly related to the frequency of light transmitted by the photon
§ The higher the frequency of light the higher the energy
§ The lower the frequency of light the lower the energy
o The energy of the photon = Planck’s constant x frequency of the light (E = h n)
· De Broglie determined that there was a direct relationship between the energy of light and the energy of particles…the conclusion of his statement is known as wave particle duality.
o Wave particle duality tells us that electrons can exhibit the properties of both waves and particles but, we are only able to study one aspect at a time.
· Bohr combined De Broglie’s discovery of the dual nature electrons with Planck’s energy calculation to show that the energy of an electron can only change in certain increments.
o This led to his proposal of energy levels in the electron cloud.
o Bohr also proposed that each line in the absorption and emission spectra was representative of an electron’s energy change
§ Absorption lines show energy gained when moving from the ground state to the excited state
§ Emission lines show energy released when electron returned to the ground state from the excited state.
Review Questions:
What did Bohr mean when he said that electrons exist in energy levels?
According to the Modern Atomic Model (post Bohr), there is a higher probability of finding electrons in certain areas around the nucleus. What do we call these areas of higher probability? (hint there are four types of these commonly referred to as s, p, d & f)
What is the “valence” shell?
What are the “core” electrons?
What do atoms in the same column have in common? What about atoms in the same row?
How many valence electrons do each of the following elements have?
a) Boron / e) Lithiumb) Chlorine / f) Calcium
c) Nitrogen / g) Silicon
d) Sulfur / h) Xenon
Draw the correct Lewis Dot Structure and ion symbol for a, c, d & h from above.
Write the filling order through 7p for electrons as atoms fill their sublevels in the e- cloud
Write the electron configuration for Lead (Pb)
Write the Shorthand/Noble Gas Configuration for Gold (Au)
State Hund’s Rule and then determine which of the following orbital filling notations is correctly constructed then explain what is wrong with the other one.
What is electromagnetic radiation (EMR)?
At what speed do all forms of EMR travel?
Explain the relationship between wavelength (l), frequency (n) and Energy (E)
E = h v & C = l v
E / l Long / vE / l / v High
List the colors in the EMR in order of increasing energy
What do the lines in the emission/absorption spectrum for an element represent?
Calculate the energy of a photon whose light has a wavelength of 645 nm
(C = 3.00 x 1017nm/s and h = 6.626 x 10 -34 J*s)
Explain how an emission spectrum differs from a continuous spectrum.
What causes the lines of specific wavelengths that appear in the emission spectra for elements?