You should review Sections 7.8, 7.9, and 7.10 before reviewing this slide show.

1. You should review Sections 7.8, 7.9, and 7.10 before reviewing this slide show.

2. Unit 27Applications of Acids and Bases (Chapter 7) Acid Rain (7.8) Antacids (7.9) Acids/Bases in the House (7.10) Acids/Bases in Health (7.10)

3. Acid Rain – The Natural Background There are three primary reactions that lead to the acidity of rain water, all occurring to some extent naturally: CO2 (g) + H2O (l) →H2CO3 (aq) SO3 (g) + H2O (l) →H2SO4 (aq) (The SO3 comes from oxygen reacting with SO2.) 3 NO2 (g) + H2O (l) → 2HNO3 (aq) + NO (g) (The NO2 comes from the reaction of N2 with O2 from processes such as lightning strikes.) So, if these are natural processes, where do the manmade contributions come into play?

4. Acid Rain – The Numbers Consider the following table that shows the difference in concentration of the three key components on the previous slide between natural sources and manmade sources. The term ppm is parts per million – just like a percent except instead of multiplying by 100 you multiply by one million (106). Reference: Table 18.2 in Chemistry: The Central Science, 12th Edition, T. L. Brown, H. E. LeMay, B. E. Bursten, C. J. Murphy, P. M. Woodward, Pearson, 2012.

5. Acid Rain – The Effects • Natural rainwater has a pH of about 5.6 due to the naturally occurring acidic oxides listed on the previous page. • Human activities nearly quadruple the concentration of SO2 in the air, in large part due to sulfur containing coal and fossil fuels. • The pH of acid rain can be as low as 4, and even lower in severe cases. That may not sound like much, but remember the pH scale is based on powers of 10. For a drop in pH of two units, the concentration of H+ increases by a factor of 102, or 100. • Freshwater lakes populated with living organisms typically have a pH between 6.5 and 8.5. The pH of freshwater does not match that of rainwater for a variety of reasons, but recall that freshwater is in contact with all sorts of rocks and minerals that can neutralize some of the acid nature. All biological life suffers severely if the pH of a lake drops below 4.0.

6. Acid Rain – A Schematic NOx refers to all of the nitrogen oxides. http://www.howstuffworks.com

7. Acid Rain – The Measurements This is a little dated, but notice how the Northeast has more acidic rain than other parts of the country. The coal in the east typically has higher sulfur content than that in the west. The difference would be even larger except the western coal has a lower energy content so you have to burn more of it to get the same amount of work done. http://pubs.usgs.gov/gip/acidrain/2.html

8. Freshwater pH – The Measurements Hydrogen ion concentration as pH from freshwater measurements made at the field laboratories, 2004. (Source: National Atmospheric Deposition Program/National Trends Network)

9. Improving Emissions Chemistry can help reduce some of the problems related to sulfur emissions. For example, the Clean Air Act was passed in 1970 with one of the goals to reduce emissions of acid rain gases as well as others. Making these changes required a combination of reducing high-sulfur content coal and fuel as well as modifying processes to reduce the emission of sulfur at the site – mainly power plants. If you want to see the effects over a period of time, look at the website: http://nadp.sws.uiuc.edu/data/animaps.aspx If you click, for example, the Concentration Animation for SO4, it gives an animation of the air concentration levels for about a twenty year period from 1985 to about 2005. It is kind of interesting to watch. You can also look at the Deposition Animation – the amount deposited in the water and soil and look at nitrate and ammonium animations as well.

10. Visible Effects http://weather-images.co.tv/images/mc/acid-rain-stone-erosion-of-statue-1.jpg Statue over entrance to castle in Westphalia, Germany built in 1702. The left picture was taken in 1908 and the right one in 1968.

11. More Visible Effects Acid rain can have a severe impact on forests. image from http://scienceclarified.com

12. Antacids • The pH of the stomach is about 1 – it is extremely acidic. • If the stomach becomes too acidic due to too much eating or perhaps stress, one will usually resort to an antacid – just a fancy name for a base. • Most antacids rely on one or more of a small list of bases to help neutralize the excess acid. This list includes: • - sodium bicarbonate • - calcium carbonate • - aluminum hydroxide • - magnesium carbonate • - magnesium hydroxide • You are already familiar with the hydroxides as bases. • The bicarbonate (HCO3- ) and carbonate (CO32-) ions also act as bases - they are the conjugate bases of a couple of weak acids. • A variety of other health effects can arise with overuse of any of these. For example, over indulgence of sodium can cause high blood pressure problems.

13. Other Acid-Base Uses and Functions • The most common acids you may find around the house include: • sulfuric acid – present in car batteries, some drain cleaners • hydrochloric acid (also known as muriatic acid) – used around swimming pools, cleaning agents • acetic acid – vinegar is about 5% acetic acid • Acids and bases are an integral part of biological processes: • Strong acids and bases can denature (break down) proteins making • them unable to carry out their function • The blood pH must be within a very narrow range centered on about 7.4 or else severe consequences can follow.

14. End of Unit 27