1. Chapter 2: Protecting the Ozone Layer
Why do we need to do to protect the ozone layer?
Isn’t ozone hazardous to human health?
Why is the ozone layer getting smaller?
What can we do (if anything) to help stop the depletion of our ozone layer?

2. Ozone Formation Energy + 3 O2 2 O3
Energy must be absorbed (endothermic) for this reaction to occur.
Ozone is an allotropic form of oxygen.
An allotrope is two or more forms of the same element that differ in their chemical
structure and therefore their properties.
Element Allotropes
oxygen O2, O3
carbon graphite, diamond, buckminister fullerenes
2.1

3. O Atomic number (Z) 8 16.00 Mass number (A) -The number of protons.
(nuclear charge) O 8
16.00
Mass number (A)
-The sum of the protons and neutrons.
2.2
2.2

4. The electrons in the outermost energy levels are called
The electrons in the outermost energy levels are called valence electrons.
The group number (of the representative elements) on the periodic table tells you the number of valence electrons.
Group 1A: 1 valence electron 1A
Group 3A: 3 valence electrons 8A 2A 3A 4A 5A 6A 7A
2.2

5. Isotopes of carbon: C-12, C-13, C-14 also written as: 12C 13C 14C
Isotopes are two or more forms of the same element (same number of protons) whose atoms differ in number of neutrons, and hence in mass.
Isotopes of carbon: C-12, C-13, C-14
also written as: 12C 13C C
2.2

6. Representing molecules with Lewis structures:
Consider water, H2O:
1. Find sum of valence electrons:
1 O atom x 6 valence electrons per atom = 6
+ 2 H atoms x 1 valence electron per atom = +2
8 valence electrons
2. Arrange the electrons in pairs; use whatever electron pairs needed to connect the atoms, then distribute the remaining electron pairs so that the octet rule is satisfied:
2.3

7. Representing molecules with Lewis structures:
Typical valence for selected atoms = the # of bonds an atom typically forms
Element
Typical valence
Classification H,
X (X= F, Cl, Br, I) 1
monovalent O 2
divalent N 3
trivalent C 4
tetravalent
2.3

8. Representing molecules with Lewis structures: Multiple bonds
Double bond
Triple bond
Occasionally a single Lewis structure does not adequately represent the true structure of a molecule, so we use resonance forms:
2.3

9. Try these; draw valid Lewis structures for:
HNO3
CO2
H2S
H2SO4
Can you draw other valid Lewis structures for HNO3?
Sulfur is under oxygen; think of H2O
Sulfur is in the 3rd period: it can have an expanded octet
2.3

10. The Nature of Light 2.4 Low E High E
Wavelength (l) = distance traveled between successive peaks (nm).
Frequency (n) = number of waves passing a fixed point in one second
(waves/s or 1/s or s-1 or Hz).
2.4

11. The Electromagnetic Spectrum
The various types of radiation seem different to our senses, yet they differ only in their respective l and n.
2.4

12. R O Y G B I V Visible: l = 700 - 400 nm Decreasing wavelength
Infrared (IR) : longest of the visible spectrum, heat ray absorptions cause molecules to bend and stretch.
Microwaves: cause molecules to rotate.
Short l range: includes UV (ultraviolet), X-rays, and gamma rays.
2.4

13. where c = 3 x 108 m/s (the speed of light)
The wavelength and frequency of electromagnetic radiation are related by: c = ln
where c = 3 x 108 m/s (the speed of light)
The energy of a photon of electromagnetic radiation is calculated by: E = h n
where h = 6.63 x J.s (Plank’s constant)
Energy and frequency are directly related-higher frequency means higher energy.
2.5

14. What is the energy associated with a photon of light with a wavelength of 240 nm?
C = ln
E = h n n
E = (6.63 x J.s) (1.3 x 1015 s-1)
= C l
E = 8.6 x J
n = 3 x 108 m/s
= 1.3 x 1015 s-1
10-9 m
240 nm x nm
UV radiation has sufficient energy to cause molecular bonds to break
2.5

15. 2.6

16. The Chapman Cycle
l ≤
A steady state
condition
l ≤
2.6

17. Biological Effects of Ultraviolet Radiation
The consequences depend primarily on:
The energy associated with the radiation.
The length of time of the exposure.
The sensitivity of the organism to that radiation.
The most deadly form of skin cancer, melanoma, is linked with the intensity of UV radiation and the latitude at which you live.
An Australian product uses “smart bottle” technology; bottle color changes from white to blue when exposed to UV light.
2.7

18. How CFCs Interact with Ozone
First, UV radiation breaks a carbon-halogen bond:
Photon (l < 220 nm) + CCl2F CClF2 + Cl. (free radicals)
2.9

19. The chlorine radical attacks an O3 molecule:
2Cl. + 2O ClO. + 2O2
Then two chlorine monoxide radicals combine:
2 ClO ClOOCl
The ClOOCl molecule then decomposes:
UV photon + ClOOCl ClOO. + Cl.
ClOO Cl. + O2
The Cl. radicals are free to attack more O3
Br radical is 45 times more effective at destroying ozone than Cl radical.
The Cl. radicals are both consumed and generated; they act as catalysts
The net reaction is: 2 O O2
2.9

20. Experimental analyses show that as ClO
Experimental analyses show that as ClO. concentrations increase, ozone concentration decreases.
2.9

21. HCFCs are alternatives to CFCs: they decompose more readily in the troposphere so they will not accumulate to the same extent in the stratosphere.
2.9