Chapter 2: Protecting the Ozone Layer

1. Chapter 2: Protecting the Ozone Layer Why do we need 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 = O3 • Ozone is an allotrope of oxygen. • Allotropes have the same elemental composition, but different formulas and structures. • Allotropes can have very different chemical and physical properties. • Graphite and diamond are both allotropes of carbon

3. Ozone can be produced by man-made activities • From high voltage discharges • As a secondary pollutant from car exhaust • Ozone is a very reactive chemical, and a hazardous pollutant at ground level • Bleaches fabric/paper • Kills bacteria • Can cause lung damage • Has a short lifetime due to it's high reactivity

4. “Ozone Layer” • Not actually a layer, but a region ~10-18 miles above the earth's surface, in the stratosphere, where ozone is created by a natural process.

5. The Regions of the Lower Atmosphere Atmospheric pressure changes with altitude 2.1

6. The ozone layer is a region in the stratosphere with maximum ozone concentration. 2.1

7. Stratospheric ozone is produced via interaction of oxygen and ultraviolet light from the sun: • step 1: O2 + energy → 2 O • step 2: O2 + O → O3 • overall reaction: 3 O2 + energy → 2 O3 (O3 can be lost by: O3 + O → 2 O2) • The stratospheric ozone concentration is very low. All of the ozone would form a pure layer over the earth less than 1 cm thick.

8. Stratospheric ozone blocks harmful ultraviolet (UV) light from the sun. • Most atmospheric gases are not able to block these rays. • To understand why some gases absorb UV, and how UV light damages the environment, we need to examine the atomic and electronic structures of matter.

9. amu = atomic mass unit

11. Most of the mass of an atom is in the small, dense nucleus. • The radius of an atom is about 100,000 times larger than the radius of the nucleus. • Electrons are located around the nucleus in orbitals. • Orbitals – not distinct like planetary orbits, but 3-D regions where electrons can probably be found (“electron clouds”). • The position of the outermost electrons determines the radius of the atom. • Most of an atom is empty space

12. An element is defined by the number of protons in its nucleus: This is called the atomic number (Z) The periodic table is arranged by atomic number, starting with #1 in the upper left and increasing across each period. • Most of the mass of an atom is in the nucleusmass number (A) = protons(p) + neutrons(n) • Atoms of elements are electrically neutral, so: # of electrons (e-)= # of protons (p)

14. All atoms of an element have the same atomic number. • All atoms of an element will therefore have the same number of electrons. • All atoms of an element may not have the same number of neutrons, so may also have different mass numbers. • Atoms of an element with different numbers of neutrons are called isotopes.

15. Isotope Notation • Z = atomic # = # of protons = p • A = mass # = p + n • Isotopes are denoted using the chemical symbol, X, or the element name: ZAX = AX = X-A = element name-A So a carbon isotope with 6 neutrons could be written as:612C = 12C = C-12 = Carbon-12

16. Atoms bond to attain more stable configurations of electrons. Electrons occupy “shells” or “energy levels” around an atom. Each shell can only hold a certain # of electrons. The inner shells are filled first. Only electrons in the outer shell can participate in bonding. Valence electrons (ve)

17. The outer electrons are called “valence electrons” or “bonding electrons” Most atoms want to get 8 ve (octet rule), which they do by gaining, losing, or sharing electrons. Main exception: Hydrogen wants 2 ve Noble gases already have 8 ve, which is why they almost never react. Valence electrons (ve)

20. Use chemical symbols and valence electrons to show bonding Used to predict chemical structures Used to predict molecular geometries This information is used to predict chemical and physical properties Lewis Structures

21. To draw Lewis Structures for molecules: • Add up valence electrons for all atoms • Draw backbone ('skeletal structure') for the molecule • Link atoms with single bonds (two shared electrons) • Usually elements down and to left on PT are central (carbon is usually central) • H and Group 17 elements almost always only form single bonds

22. Add electrons to satisfy the octet rule for all atoms: • Add electrons to outer atoms first • Electrons usually placed in pairs above, below, right, or left of symbol • Electrons between two atoms are bonding electrons, and count toward valence electrons for both atoms • 2 shared e- - one line – single bond • 4 shared e- - two lines – double bond • 6 shared e- - three line – triple bond

23. Use/move unshared pairs of electrons to form double and triple bonds, if needed, to satisfy octets • Double check: • Final structure must have the exact number of total valence electrons • Check that all atoms obey the octet rule • [(sum of desired v.e. over all atoms) – (total valence electrons)]/2 = total # of bonds

24. Representing molecules withLewis 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

25. In many cases, several valid Lewis Structures may be drawn for a molecule. • More advanced chemistry can eliminate some of these structures. • However, not all molecules can be accurately represented by Lewis Structures • The actual structure may be an average of several Lewis Structures • These are called Resonance Structures

26. Representing molecules withLewisstructures: Multiple bonds Doublebond Triplebond Occasionally, a single Lewis structure does not adequately represent the true structure of a molecule, so we useresonanceforms: 2.3

27. 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

28. Molecular Geometry • The geometric shape of a molecule may be determined from the Lewis Structure • Groups of valence electrons on an atom repel each other, and move to maximize their angles of separation • Groups of electrons are: • Unshared pairs of electrons • single bonds • double bonds • triple bonds

29. With two groups – 180° between groups • Linear • With three groups – 120° between groups • Trigonal • With four groups – 109.5° between groups • Tetrahedral • The actual shape of the molecule is dependent on the actual bonds.

30. Electronic Geometry • Use all groups of electrons to determine the angles between groups • Molecular Geometry • Look only at bonds to determine the shape of the molecule • Angles between bonds determined by electronic geometry • For central atoms with no unshared electrons, el. geom. = mol. geom.

31. Depending on it's geometry and elemental composition, a molecule may bend, twist, stretch, and compress. • The bonds act like springs, and these movements vibrate at specific frequencies, like a tuning fork. • Light also vibrates at specific frequencies, and if these frequencies match, the molecule can absorb that light.

32. Visible light is part of the electromagnetic spectrum. • Radio waves, microwaves, infrared, ultraviolet, x-rays, and gamma rays are also types of EM radiation, often generically called light. • Light acts like a particle and a wave: • The smallest particle of light is a photon • Photons have a wavelength and frequency, as do sound or ocean waves.

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

34. Wavelength (λ) – the distance at which a wave starts repeating itself, often measured from peak to peak. Frequency () – the number of waves that pass a fixed point per second. Velocity – how fast a wave moves For almost any wave, it's velocity is constant in a given medium. The speed of light (c) is constant in a vacuum and in air.

35. The Electromagnetic Spectrum The various types of radiation seem different to our senses, yet they differ only in their respective  and  2.4

36. Wavelength, frequency, and the speed of light are related by: c = λ*  So as the frequency increases, the wavelength drops, and vice versa. The energy (E) of a photon increases with it's frequency: E = h*  (h is a Plancks constant)

37. When a molecule absorbs a photon, this energy can cause: • stretching • bending • twisting • heating • ejecting electrons • breaking of bonds

38. Visible:  = 700–400 nm ROY GBIV Decreasingwavelength Infrared (IR): longest of the visible spectrum; heat ray absorptions cause molecules to bend and stretch. Microwaves: cause molecules to rotate. Short range: includes UV (ultraviolet), X-rays, and gamma rays. 2.4

39. What is the energy associated with a photon of light with a wavelength of 240 nm? E = h C =  E = (6.63 x 10–34 J.s) (1.3 x 1015 s–1)  = C  E = 8.6 x 10–19 J 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

40. When bonds are broken in living organisms, mutations or death may occur, depending on which bonds (DNA or RNA) and how many bonds are affected. • Gamma rays and x-rays can penetrate the body and break bonds internally. • Ultraviolet light is not able to penetrate the body, but can cause mutations on the surface of plants and animals. • Photons of visible light are not energetic enough to break most bonds

41. There are three classes of UV light: • UV-A (320-400 nm) • least energetic UV • Much UV-A reaches the earth's surface • UV-B (280-320 nm) • more energetic than UV-A • blocked by ozone in the stratosphere • UV-C (200-280 nm) • most energetic UV • blocked by O2 and ozone

42. 2.6

43. The Chapman Cycle ≤ + ≤ A steady state condition + 2.6

44. The ozone concentration in the stratosphere is reduced by some man made pollutants: Chlorofluorocarbons (CFCs) – contain Cl, F and C – produce chlorine free radicals in the stratosphere Free radicals are very reactive because they need another e- to get an octet of ve. Cl٠ + 2O3 → 3 O2 + Cl٠ Cl٠ is a catalyst – it speeds up this reaction, but is not consumed. One of these chlorine radicals can destroy many ozone molecules

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

46. How CFCs Interact with Ozone CFC-11 CFC-12 Freon 11 Freon 12 trichlorofluoromethane dichlorodifluoromethane CCl3F CCl2F2 First, UV radiation breaks a carbon-halogen bond: Photon  < 220 nm) + CCl2F2 .CClF2 + Cl. (free radicals) 2.9

47. The chlorine radical attacks an O3 molecule: 2 Cl. + 2O3 2 ClO. + 2 O2 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 The Cl. radicals are both consumed and generated; they act as catalysts The net reaction is: 2 O3 3 O2 2.9

48. The most common uses for CFCs have been: • air conditioners • aerosols • plastic foams The CFCs have been banned in most countries, and replaced with HCFCs, which do not pose the same threat to the ozone. But the CFCs are very inert, so have a long lifetime in the environment, and will still affect the ozone layer for many years to come.

49. HCFCs are alternatives to CFCs: they decompose more readily in the troposphere so they will not accumulate to the same extent in the stratosphere. HCFC-141b HCFC-22 dichlorofluoroethane chlorodifluoromethane C2H3Cl2F CHClF2 2.12

50. Some effects of increased UV light due to reduction of O3 concentration in the ozone layer: • Skin cancer • Skin damage • Eye damage • Harms land and aquatic animals • Harms plants – suppressed growth