1. Chapter 2 Practice questions

2. Do Now 9/26: Fill in the chart
Particle (symbol)
Mass
Charge
Location
Proton p+
neutron n0
Electron e-

3. Understanding
Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons).
Guidance: Relative masses and charges of the sub-atomic particles should be known, actual values are given in section 4 of the data booklet. The mass of the electron can be considered negligible.
Negatively charged electrons occupy the space outside the nucleus.

4. Additional Practice Problems
How are ions formed?
Fill in the following table
Species
Protons
Neutrons
electrons 30 35 28
7Li+ 15 16
14N3- 3 7

5. Applications and skills
Use of the nuclear symbol notation AX to deduce the number of protons, neutrons, and electrons in atoms and ions. Z

6. Average Relative Atomic Mass
Rarely a whole number
Each element has several isotopes that make up the average mass
What are isotopes?
If you know the % abundance and the mass of the isotopes you can calculate the average relative atomic mass
Atoms of the same element with a different number of neutrons and thus a different relative mass.

7. Calculating average mass
Lithium has two isotopes. 7.59% of the atoms are lithium-6 while % of the atoms are lithium-7. What is the average mass?
(.0759 x 6) + (.9241 x 7)
6.92 amu or 6.92 g/mol

8. Information may be provided in a mass spectrum
We will learn more about how mass spec works later in the year.
Do need to know how to read the spectrum
X axis – shows the mass/charge ratio which can be considered equivalent to their mass
Y axis shows % abundance
(.6 x 69) + (.4 x 71)
69.80 amu or g/mol

9. You can also calculate % abundance
Must know the masses of the isotopes present
Worked Example: Boron exists in two isotopic forms, 10B and 11B. 10B is used as a control for nuclear reactors. Use your periodic table to find the abundances of the two isotopes.
(x * 10) + ((1-x) * 11) = 10.81
10x + 11 – 11x = 10.81
-x = -.19
x = .19
10B = 19% abundance
11B = 81% abundance

10. Practice
Use the average atomic mass of chlorine determine the relative abundance of 35Cl and 37Cl
(x * 35) + ((1-x) * 37) = 35.45
35x + 37 – 37x = 35.45
-2x = -1.55
x = .77.5
35Cl = 77.5% abundance
37Cl = 22.5% abundance

11. Understandings & Applications and skills
Mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition
Guidance: The operation of the mass spectrometer is not required
Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.
Guidance: Specific examples of isotopes need not be learned.

12. Section 2.2 Absorption vs emission spectrum
When white light is passed through H2 gas an absorption line spectrum is produced with some colors missing.
If a high voltage is applied to the gas a corresponding emission line spectrum is produced

13. Energy of the photon emitted is equal to the energy change in atom
DEelectron = Ephoton
Also related to the frequency of the radiation by the Plank eq (constant in data book)
Ephoton=hv
Leading to:
DEelectron = hv

14. Spectrum from hydrogen helped support discrete levels for electrons (bohr model).
High energy = transitions to 1st level (UV range)
Medium energy = transitions to 2nd level (visible range)
Low energy transitions to 3rd level (IR range)

15. c = nl speed of light = frequency x wavelength

16. Understandings & Applications and skills
Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.
The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.
Guidance: The names of the different series in the hydrogen line spectrum are not required. 
Description of the relationship between color, wavelength, frequency, and energy across the electromagnetic spectrum.
Guidance: Details of the electromagnetic spectrum are given in the IB data booklet in section 3.
Distinction between a continuous spectrum and a line spectrum.
Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second, and third energy levels.

17. Bohr model fails after hydrogen
Wave particle model
Suggests it is sometimes preferable to think of an electron as having wave properties.
Uncertainty Principle – cannot know where an e- is at any given moment in time best we can hope for is a probability picture of where the electron is likely to be.
Schrodinger model – wave equation (orbitals are shapes rather than rings)

18. Principal Quantum Number, n
Indicates main energy levels
n = 1, 2, 3, 4…
(these are the levels we have talked about so far)
Each main energy level has sub-levels that are not at the same exact energy level (kind of like a split level house)

19. Orbital Quantum Number, ℓ (Angular Momentum Quantum Number)
Indicates shape of orbital sublevels – each sublevel can hold 2 e-
ℓ sublevel s p d f

21. Quantum Number relationships in Atomic Structure
Principle Quantum #: main energy level (n)
Sublevels in main energy level (l)
# orbitals per sublevel
Total orbitals in energy level
Number of electrons per sublevel
Total # electrons per main energy level (2n2) 1 s
(1) 2
(2)
(4)
(8) p 3 6
(9)
(18) d 5 10 4
(16)
(32) f 7 14

22. #1 Aufbau Principle
Electrons occupy orbitals of lower energy first.

23. Aufbau Diagram

24. -Pauli Exclusion Principle -Electron Spin Quantum Number
An orbital can hold only two electrons and they must have opposite spin.
Electron Spin Quantum Number (ms):
+1/2, -1/2

25. Hund’s Rule
When electrons enter orbitals of equal energy, 1 electron enters each orbital before pairing occurs. Analogy: Students could fill each dorm room, one person at a time, before doubling up.

26. Orbital Diagram for Hydrogen
Electron Configuration
Standard:
1s1
Condensed:
none

27. Orbital Diagram for Helium
Electron Configuration
Standard:
1s2
Condensed:
none

28. Orbital Diagram for Lithium
Electron Configuration
Standard
1s22s1
Condensed:
[He]2s1

29. Orbital Diagram for Beryllium
Electron Configuration
Standard
1s22s2
Condensed
[He]2s2

30. Orbital Diagram for Boron
Electron Configuration
Standard
1s22s22p1
Condensed
[He]2s22p1

31. Orbital Diagram for Nitrogen
Electron Configuration
Standard
1s22s22p3
Condensed

32. Orbital Diagram for Fluorine
Electron Configuration
Standard
1s22s22p5
Condensed
[He]2s22p5

33. Orbital Diagram for vanadium
Electron Configuration
Standard
1s22s22p63s23p64s23d3
condensed
[Ar]4s23d3

34. Exceptions: Chromium 24 e- 1s22s22p63s23p64s13d5
Copper 29 e- 1s22s22p63s23p64s13d10

35. Understandings & Applications and skills
The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n2.
A more detailed model of the atoms describes the division of the main energy level into s, p, d, and f sub-levels of successively higher energies.
Sub-levels contain a fixed number of orbital, regions of space where there is a high probability of finding an electron.
Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.
Recognition of the shape of an s orbital and the px, py, and pz atomic orbitals.
Application of the Aufbau Principle, Hund’s Rule, and the Pauli Exclusion Principle to write electron configurations for the atoms and ions up to Z = 36.
Guidance:
Full electron configurations (e.g. 1s2 2s2 2p6 3s2 3p4) and condensed electron configurations (e.g. [Ne] 3s2 3p4) should be covered.
Orbital diagrams should be used to represent the character and relative energy of the orbitals. Orbital diagrams refer to arrow-in-box diagrams, such as the one given below:
The electron configurations of Cr and Cu as exceptions should be covered.