Download File - Roden`s AP Chemistry

Unit 4: Atomic Structure & Periodicity
Electromagnetic Radiation
Questions to consider:
 Why do we get colors?

Why do different chemicals give us different colors?
Characteristics of light
 One way that energy travels through space
 Equation c = λ ѵ
o Wavelength (λ) – distance between two consecutive peaks or troughs in a wave
o Frequency (ѵ) - number of waves (cycles) per second that pass a given point in space
o Speed (c) – speed of light (2.9979x108 m/s)

Planks constant (h) = 6.626x10-34 J•s
o E=hѵ
E = hc/ λ
1. Calculate the frequency of light that has a wavelength of 6.7x10-5 cm.
2. Calculate the energy of a photon of orange light with a frequency of 5.0x1014 sec-1.
3. Calculate the energy of a mol of photons of orange light with a frequency of 5.0x1014 sec-1.
4. Calculate the energy of a photon of light with a wavelength of 425nm.
5. The energy required to break the oxygen-oxygen bond in O2 is 496 kJ/mol. Calculate the minimum
wavelength of light that can break the bond in one molecule.
The Bohr Model
 Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.
 Bohr’s model gave hydrogen atom energy levels consistent with the hydrogen emission spectrum.
 Ground state – lowest possible energy state (n = 1)
Energy-Level Diagram for Electronic Transitions

Orbit-Transition Diagram, Which Accounts for the
Experimental Spectrum
For a single electron transition from one energy level to another:
 1
1 
E =  2.178  1018 J  2  2 
ninitial 
 nfinal




∆E = change in energy of the atom (energy of the emitted photon)
nfinal = final distance from nucleus
ninitial = initial distance from nucleus
The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain
allowed circular orbits for the electron.
As the electron becomes more tightly bound, its energy becomes more negative relative to the zeroenergy reference state (free electron). As the electron is brought closer to the nucleus, energy is released
from the system.
Bohr’s model is incorrect. This model only works for hydrogen.
Electrons move around the nucleus in circular orbits.
6. What color of light is emitted when an excited electron in the hydrogen atom falls from:
a. n = 5 to n = 2
b. n = 4 to n = 2
c. n = 3 to n = 2
d. Which transition results in the longest wavelength of light?
HW 4A
The Quantum Mechanical Model
 We do not know the detailed pathway of an electron.
 Heisenberg uncertainty principle:
o There is a fundamental limitation to just how precisely we can know both the position and
momentum of a particle at a given time.
h
x    m  

4
Δx = uncertainty in a particle’s position
Δ(mν) = uncertainty in a particle’s momentum
h = Planck’s constant
Physical Meaning of a Wave Function (Ψ)
o The square of the function indicates the probability of finding an electron near a particular point in
space.
 Probability distribution – intensity of color is used to indicate the probability value near a
given point in space.
Radial Probability Distribution

Relative Orbital Size
o Difficult to define precisely.
o Orbital is a wave function.
o Picture an orbital as a three-dimensional electron density map.
o Hydrogen 1s orbital:
 Radius of the sphere that encloses 90% of the total electron probability.
7. What is the maximum number of electrons in a:
a. 2s subshell
b. 3d subshell
c. n = 4
d. 2p orbital
Hund’s Rule
 The lowest energy configuration for an atom is the one having the maximum number of unpaired
electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals.

Aufbau Principle
 As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to
hydrogen-like orbitals.
 An oxygen atom has an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s
subshell, and four electrons in the 2p subshell.
Oxygen: 1s22s22p4
 Orbital Diagram
o A notation that shows how many electrons an atom has in each of its occupied electron orbitals.

Valence Electrons
o The electrons in the outermost principal quantum level of an atom.
1s22s22p6 (valence electrons = 8)
o The elements in the same group on the periodic table have the same valence electron
configuration.
o The Orbitals Being Filled for Elements in Various Parts of the Periodic Table
8. Determine the expected electron configurations for each of the following.
a. S
b. B
c. Eu
9. Write the electron configuration for Kr. Show both the full configuration and the core configuration.
10. Using the element silicon, show an orbital diagram. Is it paramagnetic or diamagnetic?
11. Write the core configurations for chromium and copper.
12. Write the core configurations for:
a. S2b. Al3+
c. Fe3+
HW 4B
Excited States
 Electrons in their lowest energy configuration in an atom are said to be in the ground state.
o Obey 3 rules governing electron configurations
 Aufbau Principle – electrons enter orbitals of lowest energy first
 Hund’s Rule – electrons enter orbitals of equal energy one at a time with parallel spins
 Pauli Exclusion Principle – no two electrons in the same atom can have the same four
quantum numbers
 When electrons in the ground state absorb energy they change their configuration to an excited state.
o Either Aufbau or Hund’s are violated
o Pauli Exclusion is NEVER violated
13. For each of the following, decide if the electrons are in a ground state, excited state or impossible state. If
any rule violation occurs, state the name.
_____State
Violation_________
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HW 4C
Periodic Trends in Atomic Properties
 Atomic Radius
 Ionic Radius
 Electronegativity
 Ionization Energy
 Electron Affinity
↓
↑
↑
↑

Atomic Radius
o In general as we go across a period from left to right, the atomic radius decreases.
 Effective nuclear charge increases, therefore the valence electrons are drawn closer to the
nucleus, decreasing the size of the atom.
o In general atomic radius increases in going down a group.
 Orbital sizes increase in successive principal quantum levels.
o
o
o
Valence electrons- electrons in the highest principle energy level
Core electrons – all other electrons
Effective nuclear charge
 the net charge experienced by a particular electron in an atom
 depends on the number of electrons that screen the electron of interest
o Screening – inner-core electrons act to offset the positive charge of the
nucleus as seen by an electron further away; they screen a portion of the
nuclear charge from valence electrons
 Zeff = Z-S
 Z – nuclear charge
 S – number of core electrons
14. Which should be the larger atom? Why?
a. Na or Cl
b. Li or Cs

Ionic Radius
o Ions compared to their own neutral element
o
Ions in the same family
o
Ions that are isoelectronic

Electronegativity
o Tendency of an atom to attract an electron in a bond

Ionization Energy
o Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e–
Mg → Mg+ + e–
I1 = 735 kJ/mol (1st IE)
Mg+ → Mg2+ + e–
I2 = 1445 kJ/mol (2nd IE)
2+
3+
–
Mg → Mg + e
I3 = 7730 kJ/mol *(3rd IE)
o
o
o
*Core electrons are bound much more tightly than valence electrons.
In general, as we go across a period from left to right, the first ionization energy increases.
 Why?
 Electrons added in the same principal quantum level do not completely shield the
increasing nuclear charge caused by the added protons.
 Electrons in the same principal quantum level are generally more strongly bound
from left to right on the periodic table.
In general, as we go down a group from top to bottom, the first ionization energy decreases.
 Why?
 The electrons being removed are, on average, farther from the nucleus.
The Values of First Ionization Energy for the Elements in the First Six Periods
15. Explain why the graph of ionization energy versus atomic number (across a row) is not linear.
16. Where are the exceptions?
17. Which atom, Na or Cl, would require more energy to remove an electron? Why?
18. Which atom, Li or Cs, would require more energy to remove an electron? Why?
19. Which atom, Li or Be, has the larger second ionization energy? Why?
o

Successive Ionization Energies (KJ per Mole) for the Elements in Period 3
Electron Affinity
o Energy change associated with the addition of an electron to a gaseous atom. X(g) + e– → X–(g)
o In general as we go across a period from left to right, the electron affinities become more
negative.
o In general electron affinity becomes more positive in going down a group.
20. Arrange the elements oxygen, fluorine, and sulfur according to increasing:
a. Ionization energy
b. Atomic size
21. Briefly explain each of the following statements:
a. The atomic radius of Mg is smaller than Ba.
b. The ionization energy of S is lower than P.
c. The energy required to remove the fourth electron in Al is significantly larger than the third
electron.
d. The energy change associated with gaining an electron, the electron affinity , is positive for some
atoms and negative for others.
22. The diagram shows the first ionization energies for the elements from Li to Ne. Briefly explain each of the
following in terms of atomic structure.
a. In general, there is an increase in the first ionization energy from Li to Ne.
b. The first ionization energy of B is lower than that of Be.
c. The first ionization energy of O is lower than that of N.
d. Predict how the first ionization energy of Na compares to those of Li and of Ne. Explain.
HW 4D