Final Stand: Which is worse? Water pollution or water shortage?

Although nobody desires to drink foul, polluted water, having water available is much better than not having any water at all; methods used to purify originally polluted water make water safe to consume.
As we have learned, there are many ways to purify water to some extent: during our foul water lab, we learned about oil-water separation, sand filtration, and charcoal adsorption. Recently we learned that sand filtration, charcoal adsorption, and the distillation processes are extremely similar to natural water purification executed by the hydrologic cycle; evaporation of water vapor and condensation of water droplets in clouds removes nearly all dissolved substances, and filtration through sand and gravel removes nearly all suspended matter. The water cycle, the cycle of evaporation, condensation, and precipitation, produces some of the cleanest water on earth due to its natural purification process. Through our other labs, we learned how to test pH levels, the presence and absence of ions, and solubility of water. This knowledge resulted in our expansion of knowledge as to how to test, clean, and balance water. Just recently, we learned about municipal water treatments, or the treatments of local water that comes through our taps. Although we often feel as though bottled water is cleaner than tap water due to tap water's slight odor of chlorine, it is often not the case. Both tap and bottled water can begin extremely polluted, but through extensive cleaning processes, it is made safe to drink. Municipal water treatments include the flow of water through screens to prevent large objects from entering the water-treatment plant, prechlorination, flocculation, post-chlorination, aeration, the addition of calcium oxide (or lime) to the water to neutralize acidic water, and fluoridation to reduce tooth decay. This series of steps makes water safe for consumption, and may even make tap water more pure than bottled water.
These methods of purification of water are not the only methods, however, they are some of the most popular methods of water treatment. As I previously mentioned, before purification, water is not completely pure. It is important to recognize that due to the presence of atmospheric gases, oxygen, nitrogen, and carbon dioxide, water is never completely purified, but through purification and filtration, can be made as pure as possible and safe to drink. This is why it is essential to have water, even if it is polluted, rather than not having any at all.

Water shortage:

Polluted water:

Water cycle: Natural water purification:

Municipal water treatment: in this picture, flocculation and sand filtration:

Extra Credit for Friday, July 1st: Drugs in the Netherlands: Closed shops: Why tourists in the Netherlands may have to stop smoking pot

Drugs in the Netherlands
Closed shops
Why tourists in the Netherlands may have to stop smoking pot
Jun 23rd 2011 | AMSTERDAM | from the print edition

Although the Dutch society is known to be liberal due to legalized prostitution, gay marriage, and the Dutch "coffee shops", where cannabis, or marijuana, is freely sold to the public for private use, there may be an end to these famous coffee shops. The Dutch government is beginning to institute rules against about 660 coffee shops, and may force them to become members-only clubs that involve a strict registration process. On top of this, these shops will only be accessible to Dutch residents, not drug tourists who greatly contribute to the business. The official policy regarding Dutch coffee shops will not be instituted immediately, but debated in parliament in September. Since the southern provinces have the most cannabis tourists, they will be the first to be effected by these new laws that go against the previously liberal governmental rule. The reason the government wishes to institue these laws is to stop the “nuisance” of drug tourists and to fight against organized crime, defined as criminal activity on the part of an organized and extensive group of people. However, the Dutch government is led by liberals and backed by others who feel as though this act compares to the "incompatible" Islamic values, which recently tried to regulate prostitution. Even though this failed, it is another example of the evolution of Dutch liberalism. Derrick Bergman, head of a lobby group that fights for the legalization of marijuana called VOC, feels as though approval of the new guidelines will bring an end to the coffee shops. Despite the intentions of protecting the people, with this end, tourism would decline drastically due to the virtual end of the multi-million euro cannabis market. Also, new and worse criminality, like illegal drug trade and street dealing, will increase exposure to hard drugs; this would be even more dangerous to the population, and even if people believe they are buying marijuana, it could be laced with hard, dangerous, and addictive substances. Even though people are anticipating these new guidelines of being passed, they may not be. Earlier on, a coffee-shop owner of a southern Dutch city filed a case against an earlier decision by the city’s mayor to initiate a membership program for only those of the Dutch population. This case is currently in the hands of the highest Dutch appeal body, Council of State. Whatever the Council decides will impact the fate of the regulation act and affect whether it will be passed or declined.

http://www.economist.com/node/18867682
http://dictionary.reference.com/browse/organized+crime

ISDS #1-9, 18-22

1) Make a diagram of the hydrologic cycle and label all processes.

2) List three major processes that occur in natural water purification and, for each, identify the contaminants that the process removes.

1) Evaporation of water vapor and condensation of water droplets in clouds: removes nearly all dissolved substances.
2) Bacterial action: converts dissolved organic contaminants into simple compounds.
3) Filtration through sand and gravel: removes nearly all suspended matter.

3) How are the properties of aluminum hydroxide related to the process of flocculation?

Aluminum hydroxide is formed by alum and slaked lime. It is a sticky, jellylike material that traps and removes the suspended particles in the water-- flocculation.

4) Why is calcium oxide (CaO) sometimes added in the final step of municipal water treatment?

Sometimes, water is acidic enough to dissolve metallic water pipes. This could result in the entrance of undesirable ions into the water supply. Lime, or calcium oxide, is added in the final step of municipal water treatment because it is basic and it neutralizes acidic water by raising its pH to a proper level.

5) Fluoride, an ingredient in many types of toothpaste, is sometimes added to municipal water supplies in the last stage of water treatment. How much fluoride is added and what is its purpose?

As much as 1 ppm of fluoride is sometimes added to municipal water supplies in the last stage of water treatment because it reduces tooth decay.

6) What are advantages of chlorinated drinking water compared to untreated water?

Chlorinated drinking water has saved countless lives because it controls waterborne diseases by killing disease-producing microorganisms. Chlorine kills dangerous bacteria and along with other purification processes, makes water safe to drink.

7) Is there a disadvantage to using chlorination in water treatment? Explain.

Yes. Chlorine in water can react with organic compounds made by decomposing animal and plant water to form substances that can be harmful to humans. One group of the substances, trihalomethane (THM), such as chloroform, is a cancer causing substance, or a carcinogen. These substances are very dangerous to human health.

8) Water from a clear mountain stream may require chlorination to make it safe for drinking. Explain.

Water from a clear mountain stream may require chlorination to make it safe for drinking due to various types of bacterias found on the ground, such as feces left by animals, or dirty substances added to the water through animals drinking and bathing in the water.

9) List two alternatives to the use of chlorination in municipal water treatment.

1) Charcoal filters: they remove most organic compounds from water, including THMs *(but they are expensive, hard to clean, and need to be replaced often).
2) Ozone or ultraviolet light: these disinfect water *(but do not protect the water once it has left the treatment center).

18) Explain what would happen to Earth's hydrologic cycle if water evaporation suddenly stopped.

Evaporation is the starting point of the hydrologic cycle, and if evaporation suddenly stopped, so would condensation and precipitation. Although the water underground may continue to flow for a period of time, it would never be replenished by rain water or snow and eventually dry out; the hydrologic cycle would stop.

19) One unique characteristic of water is that it is present in all three physical states (solid, liquid, and gas) in the range of temperatures found on Earth. How would the hydrologic cycle be different if this were not true.

The hydrologic cycle involves all three physical states (solid, liquid, and gas). Water evaporates as vapor, or a gas, condenses into tiny water droplets in clouds (liquid), and falls as either rain water (liquid), snow, or hail (solids). If this were not true, the hydrologic cycle would not function the way it does, and not be able to exist.

20) Why does the EPA limit the concentration of THMs to 80 ppb instead of requiring their total elimination from municipal water supplies.

The alternatives to the use of chlorination carry disadvantages, and since very little amounts (80 ppb) of THMs should not have much effect on the human body, the EPA recognizes the difficulties of using alternatives and accepts the substance in extremely low concentrations.

21) Compare how the various processes used in the foul-water investigation (page 10) are similar to steps in the natural purification of water.

The processes of evaporation and condensation mimic the process of distillation and what happens in a distillation apparatus on a grander scale. Also, when water seeps into the ground and it passes through gravel, sand, and rock, suspended mater are filtered out; this is like the sand filtration process used in the foul-water investigation.

22) Some physicians recommend consuming about 2 L of water daily. Municipal water supplies may contain up to 1 ppm fluoride. Assume that you drink 2 L of water per day. At 1 ppm fluoride, how many grams of fluoride ion would you consume in
a. one day?
b. one week?
c. one year?

a. 1 ppm fluoride.
b. 7 ppm fluoride.
c. 365 ppm fluoride.

Dissolved Oxygen Levels in Water: Riverwood Fish Kill?

Dissolved oxygen is essential for living organisms in water, however, too much or too little can be fatal. Since the solubility of gas within water decreases as temperature increases, the high temperatures of the hot summer months can create problems for dissolved oxygen levels in water. Fish are cold blooded creatures, meaning the temperature of their blood adjusts to the temperature of the surrounding environment. High temperatures cause the metabolisms of fish to increase. This makes the fish need more food, more exercise, and therefore, more oxygen. When the fish consume more oxygen, there is less oxygen in the water, and although it may sound ironic, fish suffocate when there are low dissolved oxygen levels in their waters.
There is also such thing as too much dissolved oxygen levels in water. Since oxygen and nitrogen are atmospheric gases, when there is dissolved oxygen in water, there is also dissolved nitrogen. During times of low temperatures in water or distribution of large quantities of air into water (see dam picture), water can become a supersaturated solution, with 110%-124% oxygen gas and nitrogen gas saturation. Although oxygen may be partially utilized by fish during metabolism, nitrogen gas clogs the capillaries of the fish, which leads to the development of bubble trauma syndrome. Not only does this kill fish, but it can only be diagnosed promptly after death through careful dissection of gills in search of gas bubbles.

Too little dissolved oxygen in water:

Too much dissolved oxygen gas *AND nitrogen gas in water: this creates supersaturated water with oxygen and nitrogen gas.

Solubility Lab Report: The Metals

ABSTRACT:

The objective of the solubility lab was to test the solubility of succinic acid (C4H604) at three different temperatures. Our main concern at the beginning of the lab was working with this slightly toxic acid. This placed much importance on being very careful and listening to and following instructions thoroughly. Although this lab may seem simple, one mistake may easily falsely change the results. In order to be safe, we wore latex gloves at all times and only handled the test tubes with tongs to protect us from the heat. Going into this experiment, we knew that the solubility of succinic acid would rise with temperature because succinic acid is a solid at room temperature; therefore, we knew if our results showed the solubility going down, that a mistake must have been made. In this experiment, we used a lot of laboratory equipment: a plastic weighing boat, a 400 mL beaker to heat the water, six test tubes, a stirring rod, a thermometer, a graduated cylinder (for measuring degrees in celsius), a scale, a beaker tong, a scale, a large beaker to cool the water in an ice bath, and a heater. First, we added 300 mL of water to our beaker, heated it to 45°C on the heater, added 4 grams of succinic acid and 15 mL of water to a test tube, and placed it in the heated water bath of the beaker. After the solution had been stirred for seven minutes in 30 second intervals, we put the test tube in a cold iced water bath for two minutes and then let it sit in room temperature in the beaker stand for five more minutes as the solute settled to the bottom of the test tube. We measured the solute after five minutes and repeated this process with 55°C and 65°C. At the end of the whole procedure, we recorded our data and found our average solute measurement (11.66 mm of succinic acid).

PROCEDURE:

Our first step to begin the process was to organize all of our laboratory equipment carefully; any mistakes to the seemingly simple process could lead to inaccurate data and disorganization. At first, our group, The Metals, had trouble with our heater and soon found it was broken. We began the lab experiment with what felt like chaos, but once we had a working heater and all of our equipment laid out in front of us, the procedure went smoothly! To begin the actual procedure, we filled our 400mL beaker with 300mL of tap water. We then placed it on the water heater that, and turned it on to its highest setting, 6. The reason why we set it to its highest setting was to get the water heated as quickly as possible, for due to our earlier mishap with the broken heater, we really needed to regain momentum. We placed our thermometer into the beaker, and once it reached our first temperature, 45°C, we lowered the setting on the heater, in order to maintain the heat of the water. We used the weighing boat and a very sensitive scale to carefully measure out 4 grams of succinic acid (C4H6O4) three times. We placed each 4 gram measurement of the succinic acid into 3 different test tubes and added 15 (not 20 due to the small test tubes) mL of water to each tube; by doing this, we would be prepared for each step of the process on time. Due to some confusion, we placed two tubes into the heated beaker, and although using the thermometer we were able to measure the heat of the tap water in the beaker to 45°C, the solution within the test tubes remained at 43°C; this was okay, however, because as long as the temperature of the succinic acid solution was within 2°C of the target temperature, it was acceptable. For 7 minutes, every 30 seconds we would use a stirring rod to mix the succinic acid solutions in hopes of increasing the solubility. After seven minutes had passed, we used the beaker tongs to remove the test tube from the water. We then extracted the liquid from the tube with a beral pipet and put it in a new, clean test tube (leaving the un-disolved solute in the original test tube). Then, we placed the new test tube into an iced water bath. Our test tube sat in the water bath for two minutes, and after two minutes, we took it out of the cold water and put it into a test tube rack to sit in room temperature for five minutes. At this time, our solution appeared to be clear. After waiting five minutes, the solution was still clear; therefore, due to our lack of solute, 0 mm, our group was able to conclude that succinic acid was practically insoluble at 45°C. Next, we increased the heat of the heater in order to raise the temperature of both the tap water within the beaker and the solution within the test tube to 55°C. Using the thermometer, we found that we were able to obtain a 55°C temperature in the water in the beaker, and a 54°C temperature in the succinic acid solution. Again, we used a stirring rod to mix the solution for for 7 minutes in thirty second intervals, in order to mix the solute into our distilled water solvent. After seven minutes had passed, we used our beaker tongs to remove the test tube from the heated water. We extracted the liquid from the tube with a beral pipet and put it in a new, clean test tube (leaving the un-disolved solute in the original test tube). Then, we placed the new test tube into our iced water bath for two minutes. This time, however, when we took the test tube out of the water bath, we noticed significant amounts of solute settling to the bottom of the tube. After five minutes of sitting in the test tube holder at room temperature, there was 15 mm of solute at the bottom of the tube. Finally, we were ready for our last portion of the procedure, finding the solubility of succinic acid in 15 mL water at 65°C. Since we had not placed our prepared test tube of distilled water and four grams of succinic acid in the beaker atop the heater in the beginning, we had to spend a bit more time heating our test tube. Once we brought the heat of both the water in the beaker and the water within the test tube to 65°C, (and repeatedly used the thermometer to check the temperatures of the water in the beaker and the succinic acid solution), we lowered the setting on the heater from 6 to about 1. We were then ready to begin the process. We repeated the process of using a stirring rod to mix the succinic acid solution for 7 minutes in thirty second intervals. After seven minutes had passed and the distilled water solvent was as mixed with succinic acid as it could be, we used our beaker tongs to remove the test tube from the heated water. We extracted the liquid from the tube with a beral pipet and put it in a new, clean test tube (leaving the un-dissolved solute in the original test tube). Then, we placed the new test tube into our iced water bath for two minutes. By the end of the two minutes, we noticed a more solute settling to the bottom of our test tube than the amount that settled in the solution created at 55°C. With this observation, we knew we were getting (semi) accurate data. After five minutes of sitting in the test tube rack, we measured 20 mm of succinic acid settled to the bottom of the tube. After carefully recording all of our data and observations, we were ready to clean up all of the equipment and our area, and prepare to find our average, and the class average of crystal height obtained at each temperature.

Water temperature being measured:

Solubility of succinic acid: From left: 45°C, 55°C, 65°C:

OVERVIEW OF RESULTS:

Us, The Metals:
-Solubility of succinic acid at 45°C: 0mm of settled succinic acid at the bottom of the test tube; 0mm solubility.
-Solubility of succinic acid at 55°C: 15mm of settled succinic acid at the bottom of the test tube; 15mm solubility.
-Solubility of succinic acid at 65°C: 20mm of settled succinic acid at the bottom of the test tube; 20mm solubility.
Our Average (of all three temperatures): 11.66mm solubility

Class Average: Since only two of the groups successfully obtained three sets of data for each temperature, 45°C, 55°C, and 65°C, we only included two groups, The Acids and The Metals, in our calculations for the class average.
-Solubility of succinic acid at 45°C: 1.5mm of settled succinic acid at the bottom of the test tube; 1.5mm solubility.
-Solubility of succinic acid at 55°C: 25mm of settled succinic acid at the bottom of the test tube; 25mm solubility.
-Solubility of succinic acid at 65°C: 28mm of settled succinic acid at the bottom of the test tube; 28mm solubility.

Class Data:

Class Average:

DATA ANALYSIS:

1. Find the mean crystal height obtained by your entire class for each temperature reported.

Because only two of the groups successfully obtained three sets of data for each temperature, 45°C, 55°C, and 65°C, we only included two groups, The Acids and The Metals, in our calculations for the class average. For the procedure executed to find solubility of succinic acid at 45°C, our group, The Metals, measured 0 mm of settled succinic acid at the bottom of the test tube; 0 mm solubility. The Acids measured 3 mm settled succinic acid at the bottom of the test tube; 3 mm solubility. The average of the two groups at 45°C was 1.5 mm settled succinic acid at the bottom of the test tube; 1.5 mm solubility. For the procedure executed to find solubility of succinic acid at 55°C, The Metals measured 15 mm of settled succinic acid at the bottom of the test tube, or 15 mm solubility, and The Acids measured 35 mm settled succinic acid at the bottom of the test tube, or 35 mm solubility. The average of the two groups at 55°C was 25 mm settled succinic acid at the bottom of the test tube; 25 mm solubility. The last test we conducted was to find solubility of succinic acid at 65°C. Once again, The Metals measured less than The Acids, measuring 20 mm of settled succinic acid at the bottom of the test tube, or 20 mm solubility, while the Acids measured 36 mm of settled succinic acid at the bottom of the test tube, or 36 mm solubility. The average of 65°C was 28 mm of settled succinic acid at the bottom of the test tube; 28 mm solubility.

2. Plot the mean crystal height in millimeters (y-axis) versus the water temperature in degrees Celsius (x-axis).

QUESTIONS:

1. Why is it useful to collect data from more than one trial at a particular temperature?

It is useful to do more than one trial at a particular temperature because results could vary every time due to various factors, such as small changes in temperature, the amount of time spent and intensity of mixing the solution, and the amount of time that the solution is left to sit in the hot water, the ice water, and in room temperature.

2. How did you make use of the properties of a saturated solution at different temperatures?

Since we knew that succinic acid is a solid solute and not a gas, we entered the procedure with the awareness that when the temperature of the water was risen, the solubility of the succinic acid should have risen as well. Although we did not run into the situation of solubility decreasing when the temperature was risen, if we had, this knowledge would have kept us from continuing the procedure with false data. Knowing that a supersaturated solution, if disturbed and cooled, rebalances itself and loses extra solute particles, it was interesting to see how many solutes sank to the bottom of the test tubes when the tubes were put in ice baths.

3. Did all the succinic acid that originally dissolved in the water crystallize out of the solution? Provide evidence to support your answer.

Yes, by cooling the clear solution in an ice bath after separating it from its original test tube and then allowing it to sit in room temperature for five minutes, we re-crystallized all of the succinic acid that originally dissolved in the heated water. By doing this, we were able to see the solubility of the succinic acid in each given temperature.

4. Given pooled class data, did you have enough data points to make a reliable solubility curve for succini acid? Would the curve be good enough to make useful predictions about succinic acid solubility at temperatures you have not yet investigated? Explain your answer.

No. Only two of the groups, The Metals and The Acids, actually completed the experiment with solubility data for each temperature. Because of this, we were only able to use these groups to average class data, even though there are many groups in the class. The curve we constructed only included the averaged data of these two groups; therefore, the curve we made would not be good enough to make useful predictions about succinic acid solubility at temperatures we have not yet investigated.

5. What procedures in this investigation could lead to errors? How would each error affect your data?

A huge error many groups made was putting the solute in the original test tube, and not the separated liquid, into the ice bath. This made these groups unable to complete their data tables. Also, wrong measurements of temperatures, too much or too little succinic acid, dirty tools, and too much or too little time being heated and cooled are all factors that would lead to errors and inaccurate data.

6. Using your knowledge of solubility, propose a different procedure for gathering data to construct a solubility curve.

A different procedure for gathering data to construct a solubility curve would be to take a substance and pour it into a beaker filled with water while mixing it at a given temperature. Once the water became a saturated solution, one would record the data, increase the temperature of the water, and repeat the procedure.

Explaining pH: pH in a nutshell!

The pH scale is used to measure and report whether a solution is acidic, basic (alkaline), or chemically neutral. The pH scale is numbered 0-14; at room temperature, 25°C, below 7 means the solution is acidic, 7 means the solution is chemically neutral, and above seven means the solution is basic. Each unit on the pH scale has a tenfold difference in acidity or alkalinity than the unity above or below. In order to measure pH, litmus (vegetable dye) paper is used. After applying a few drops of a solution or dipping the paper in a solution, if the paper turns blue, the solution is basic, and if the paper turns red, the solution is acidic. Concentrated acids and bases react chemically with many other substances, and they often corrode, or eat away at, materials. Most acids are made up of one or more hydrogen (H) atoms, and bases generally include hydroxide ions (OH-). Although ammonia and baking soda do not contain hydroxide ions, since they react with water to generate hydroxide ions, they produce basic solutions. Acidic and basic solutions affect living organisms. For example, low pH (acidity) in water can impair fish-egg development and cause increased concentrations of metal ions in natural waters. High pH (basic contamination) dissolves organic materials, like skin and scales.
Additionally, although 7 is the ideal pH for drinking water, the EPA allows drinking water to be in the range of 6.5-8.5 pH.

ISCS #20-27, 33, 35, page 83-84

20) Using Figure 1.44 on page 70, decide which is more acidic:
a. a soft drink or a tomato
b. black coffee or pure water
c. milk of magnesia or household ammonia

a. a tomato.
b. pure water.
c. milk of magnesia.

21) How many times more acidic is a solution at pH 2.0 than a solution at pH 4.0.

Since there is a tenfold difference between each number on the pH scale, a solution at pH 4.0 is 100 times more acidic than a solution at pH 2.0.

22) List three negative effects of inappropriate pH levels on aquatic organisms.

Too low pH (high acidity):
1) fish egg development impaired, reproduction harmed.
2) increased concentrations of metal ions in natural waters by leaching metal ions from surrounding soil.

High pH (basic contamination):
3) dissolves organic materials, such as scales on fish.

23) Distinguish between polar and nonpolar molecules.

A polar molecule is a molecule with an uneven distribution of electrical charge; each molecule has a partial positive region at one end and a partial negative region at the other end. H2O (water) is a polar molecule, since there 8 oxygen electrons and only 2 hydrogen electrons. Since there are more electrons in oxygen, there is a greater negative charge at the oxygen end, and a partially positive charge at the hydrogen end of the molecule. A nonpolar molecule's distribution of electrical charge is equal throughout, meaning it is balanced at both ends of the molecule.

24) Would you select ethanol, water, or lamp oil to dissolve a nonpolar substance? Explain.

Ethanol and water are both polar substances, so, to dissolve a nonpolar substance, I would select the nonpolar substance, lamp oil; polar substances dissolve in polar solvents and nonpolar substances dissolve in nonpolar solvents-- "like dissolves like."

25) Why does table salt (NaCl) dissolve in water but not cooking oil?

Since water is an extremely attractive polar molecule, the positive and negative charges of Na (cation) Cl (anion) are "pulled" by the charges of the water molecules. Cooking oil on the other hand, is nonpolar; therefore, cooking oil does not have the same effect on NaCl as water does, and NaCl does not dissolve in it.

26) Explain the phrase "like dissolves like."

As I stated in #24, polar substances dissolve in polar solvents and nonpolar substances dissolve in nonpolar solvents. This relationship between solvents that dissolve solutes with similar electrical charges is described as "like dissolves like."

27) Explain why you cannot satisfactorily clean greasy dishes with just plain water.

Since water is composed of polar molecules and grease, or oil, is composed of nonpolar molecules, water would not be able to dissolve the grease on the dishes. In order to dissolve the grease, the dishes would need to be washed with a solution composed of nonpolar molecules. "Like dissolves like!"

33) Many mechanics prefer to use waterless hand cleaners to clean their greasy hands. Explain
a. what kind of materials are likely to be found in these cleaners.
b. why using these cleaners is more effective than washing with water.

a. Nonpolar materials are likely to be found in these cleaners.
b. Using these cleaners would be more effective than washing with water because since water molecules are polar, they would not dissolve the grease off of the mechanics' hands. Instead, the water would sort of roll off of their hands without them being cleaned.

35) Fluorine has the highest electronegativity of any element. Fluorine and hydrogen form a polar bond. Which atom in HF would you expect to have a partial positive charge? Explain.

Since fluorine (F) has the highest electronegativity of any element, meaning the most electrons, it would make up the negative end of the polar bond, HF. Since hydrogen (H) has less negative charges, or less electrons, than fluorine, it would have the partial positive charge of the polar bond.

Riverwood Fish Kill!

http://www.flyfishingfrenzy.com/2010/04/15/community-conservation-clean-up-in-minnesota/

ISCS #9-19, p.82-83

9) Calculate the masses of water and sugar in a 55.0-g sugar solution that is labeled 20.0% sugar by mass.

.20 x 55 = 11
11g sugar.

55-11 = 44
44g water.

11g sugar, 44g water.

10) The EPA maximum standard for lead in drinking water is 0.015 mg/L. Express this value as parts per million (ppm).

15,000ppm.

11) What makes a water molecule polar?

Since water has an uneven distrobution of electrical charge, where the end of the oxygen atom is negatively charged and the ends of the hydrogen atoms are positively charged, a water molecule is polar.

12) Draw a model that shows how molecules in liquid water generally arange themselves relative to one another.

See drawing.

13) Which region of a polar water molecule will be attracted to a
a. K+ ion?
b. Br- ion?

a. The negatively charged end, the oxygen end.
b. The positively charged end, the hydrogen end.

14) Why are heavy metals called heavy?

Heavy metals are called heavy because thheir atoms have greater masses than those od essential metallic elements, and therefore, are harmful to humans and other organisms.

15) List three symptoms of heavy metal poisoning.

Numbness, staggered walk, tunnel vision, and brain damage are (four) symptoms of heavy metal poisoning. Heavy metal poisoning can lead to death; it causes damage the nervous system, brain, kidneys, and liver.

16) List two possible sources of human exposure to
a. lead
b. mercury.

a.
1) Soil surrounding heavily traveled roads, for it may have absorbed automobile exhaust from the gas from cars that contained tetraethyl lead before the 1970s.
2) Flaking leaded paint in houses built before 1978/.
b.
1) Thermometers and florescent light bulbs.
2) Fish exposed to mercury in the water.

17) What ion is found in many bases?

Hydroxide (OH-) is found in many bases.

18) What element is found in most acids?

Hydrogen (H) is found in most acids.

19) Classify each sample as acidic, basic, or chemically neutral:
a. seawater (pH = 8.6)
b. drain cleaner (pH = 13.0)
c. vinegar (pH = 2.7)
d. pure water (pH = 7.0)

a. basic.
b. basic.
c. acidic.
d. chemically neutral.

C.5 Questions #1-3, p. 62

1) Suppose you dissolved 40g potassium chloride (KCl) in 100g water at 50°C. You then let the solution cool to room temperature, about 25°C.
a. What changed would you see in the beaker as the solution cooled? See Figure 1.40.
b. Draw models of what the contents in the beaker would look like at the particulate level at 50°C, 40°C, and 25°C.

a. As the solution cooled, it would become a supersaturated solution, and therefore, if the beaker was undisturbed and no more KCl was added, nothing would happen to the solution. However, since the solution would be unstable, if one crystal of KCl was added or the beaker was knocked, the extra solutes in the solution would become precipitate, and the solution would rebalance itself into a saturated solution.
b. See drawing

2) An unsaturated solution will become more concentrated if you add more solute. Decreasing the total volume of water in the solution (such as by evaporation) also increases the solution's concentration. Consider a solution made by dissolving 20g KCl in 100g water at 40°C.
a. Draw a model of this solution.
b. Suppose that while the solution was kept at 40°C, one-fourth of the water evaporated.
i. Draw a model of this final solution and describe how it differs from your model of the original solution.
ii. How much water must evaporate at this temperature to create a saturated solution?

a and b i: See drawings.
b ii: 50 grams of water would have to evaporate at this temperature to create a saturated solution.

3) A solution may be diluted (made less concentrated) by adding water.
a. Draw a model of solution containing 10.0 KCl in 100g water at 25°C.
b. Suppose you diluted this solution by adding another 100g water with stirring. Draw a model of this 25°C new solution.
c. Compare your models in Questions 3a and 3b. What key feature is different in the two models? Why?

a and b: See drawings.
c. In 3a, the concentration of the solution expressed as a percent KCl by mass is 10%. Since the amount of water is double in 3b, 10g/200g X 100% = 5%. The percent concentration of 3a is 10%, where as the percent concentration of 3b is 5%.

Extra Credit for Friday, June 24: Obesity: Does light make you fat?: When—not just what—mice eat affects how much weight they put on.

Obesity
Does light make you fat?
When—not just what—mice eat affects how much weight they put on
Oct 14th 2010 | from the print edition

Although processed and fatty foods and lack of exercise are blamed for the obesity epidemic, recent studies have shown that another factor may be light nights. Since light regulates the body's biological clock which controls metabolism, brighter nights disturb the metabolism's perception of meal times and times for sleep. Laura Fonken of Ohio State University and her team of researchers examined how nocturnal light can affect weight, body fat, and glucose intolerance by conducting tests on mice; mice are like physiologically like humans but a large problem in the experiments is the fact that mice are nocturnal (awake at night) and humans are diurnal (asleep at night). Fonken arranged her mice into three cages for an eight week period: the first cage was lit constantly, simulating a never-ending overcast day, the second cage simulated a natural habitat, with 16 hours of overcast daylight followed by eight hours of darkness, and the third cage was also lit with 16 hours of overcast daylight, but the eight hours of darkness was replaced with a dim glow, simulating twilight at the first "flickers" of dawn. Over the eight-week period, even though they ate similar amounts of food and moved about just as much, the mice in the first and third cages gained 50% more weight than the mice in the second cage who were placed in a natural light and dark cycle; nocturnal light did affect weight, body fat, and glucose intolerance by increasing all three categories. The difference between each groups of mice and the reason why the mice in the first and third cages gained so much weight was not what the mice ate, but when they ate. Since mice are nocturnal, as expected, the mice in the natural condition only ate a 33% of their food in the overcast day part of their cycle, unlike the mice exposed to the twilight/dawn who ate over 55% of their food in the day part of their cycle. Even though the mice exposed to a day and twilight/dawn cycle did have eight hours of dim light, they did not have any time of complete darkness, and this really made a difference in their times of food consumption. Furthermore, Dr. Fonken observed that due to their disturbed biological clocks, the mice in the first cage that was lit constantly gained 10% more weight than the mice in the other two cages who faced light that simulated light and (semi)-dark cycles. Again, humans are mice are different because humans are diurnal and mice are nocturnal. Even so, like the mice that ate in the "day" due to the unnatural light in their environments, many humans eat their main meals late at night due to the spread of electric lighting. Although it has never actually been tested properly, many nutritionists believe eating late is a factor in putting on weight. Further tests are needed to prove this theory, but these experiments show that artificial lighting may be a part in weight gain.

http://www.economist.com/node/17248910

What I need to work on for Friday's test

Math problems:
-unit conversions (especially the complicated ones)
-solubility curve problems
-facts about the metric systems

Vocabulary
-beware of questions like "what particle of an atom is bigger?" *know your atoms and molecules!

Revisit ions

Look over element symbols

Study pH scale and the place of different substances

Refresh memory on drawings

Revisit labs and class demos

Look over notes

ISCS #1-8, p.82

1) Explain why three teaspoons of sugar will completely dissolve in a serving of hot tea, but will not dissolve in an equally sized serving of iced tea.

Due to the different temperatures, the solubility, or maximum quantity of a substance that will dissolve in a certain quantity of water at a specified temperature, for hot tea is higher than the solubility of iced tea. This is why three teaspoons of sugar will completely dissolve in a serving of hot tea, but not in an equally sized serving of iced tea.

2) What is the maximum mass of potassium chloride (KCl) that will dissolve in 100g water at 70°C?

The maximum mass of KCl that will dissolve in 100g water at 70°C is 48g.

3) If the solubility of sugar (sucrose) in water is 2.0g/mL at room temperature, what is the maximum mass of sugar that will dissolve in
a. 100mL water?
b. 355mL (12 oz) water?
c. 946mL (1 qt) water?

a. 200g.
b. 710g.
c. 1,892g.

4) Rank the substances in Figure 1.33 (page 54) from most soluble to least soluble at
a. 20°C.
b. 80°C.

a. NaCl (sodium chloride), KCl (potassium chloride), KNO3 (potassium nitrate).
b. KNO3 (potassium nitrate), KCl (potassium chloride), NaCl (sodium chloride).

5) Distinguish between the terms saturated and unsaturated.

Saturated means that too much of a substance has been added into a solvent for it to be able to dissolve into it, and therefore it settles at the bottom of the container; no amount stirring will make the crystals dissolve into the liquid. A saturated solution, however, is a solution in which the solvent contains as much dissolved solute as it normally can at that temperature. On the other hand, an unsaturated solution is a solution that contains less dissolved solute than the amount the solvent can normally hold at that temperature.

6) Using the graph on page 54, answer these questions about the solubility of potassium nitrate, KNO3:
a. What maximum mass of KNO3 can dissolve into 100g water if the water temperature is 20°C?
b. At 30°C, 55g KNO3 is dissolved in 100g water. Is this solution saturated, unsaturated, or supersaturated?
c. A saturated solution of KNO3 is formed in 100.0g water at 75°C. If some solute precipitates as the saturated solution cools to 40°C, what mass (in grams) of solid KNO3 should form?

a. 30g KNO3 can dissolve into 100g water if the water temperature is 20°C.
b. This solution is supersaturated.
c. 60g of solid KNO3 should form.

7) You are given a solution of KNO3 of unknown concentration. What will happen when you add a crystal of KNO3 if the solution is
a. unsaturated?
b. supersaturated?
c. saturated?

a. The solution would either become saturated or remain unsaturated (depending on how many crystals are needed for it to become saturated)
b. The one crystal would cause the other extra solute crystals to appear and settle. This would make the solution rebalanced.
c. The extra crystal would settle to the bottom of the container as a precipitate, since the solvent would have contained as many solutes as it could hold.

8) A 35-g sample of ethanol us dissolved in 115g water. What is the percent concentration of the ethanol, expressed as percent ethanol by mass?

115g water
35 x 1.15 = 40.25
1.15 x 40.25g = 46.2875g
46.2875/115 = 0.4025
The percent concentration of the ethanol, expressed as percent ethanol by mass, is 40.25%

C.2 Questions #1-3

1)
a. What mass (in grams) of potassium nitrate (KNO3) will dissolve in 100g water at 60°C?
b. What mass (in grams) of potassium chloride (KCl) wil dissolve in 100g water at the same temperature?

a. 105g of potassium nitrate will dissolve in 100g water at 60°C.
b. 45g of potassium chloride will dissolve in 100g water at 60°C.

2)
a. You dissolve 25g potassium nitrate in 100g water at 30°C, producing an unsaturated solution. How much more potassium nitrate (in grams) must be added to form a saturated solution at 30°C?
b. What is the minimum mass (in grams) of 30°C water needed to dissolve 25g potassium nitrate?

a. 20g of potassium nitrate must be added to form a saturated solution at 30°C.
b. 45g is the minimum mass of 30°C water needed to dissolve 25g potassium nitrate.

3)
a. A supersaturated solution of potassium nitrate is formed by adding 150g KNO3 to 100g water, heating until the solute completely dissolves, and then cooling the solution to 55°C. If the solution is agitated, how much potassium nitrate will precipitate?
b. How much 55°C water would have to be added (to the original 100g water) to just dissolve all of the KNO3?

a. If the solution is agitated, 55g potassium nitrate will precipitate.
b. 110.2°C water would have to be added to the original 100g water to just dissolve all of the KNO3.

What did you learn from this lab about water and about process?

I learned that "purified water" is not as pure as I thought it was. I never gave much thought to, or realized, that there are many other things in my water other than the components that create water. By doing this lab, I realized that there are various ions in each source of water on this earth, and unless the water is distilled, they are not completely removed from the water systems. I also learned that in order to test for the presence or absence of ions in different solutions, it is necessary to go through tedious processes that require much focus and observation. Testing water was definitely not as easy as I thought it would be!

Water Testing Lab Report

Abstract:

In this lab, the class was assigned to test water for different ions through confirming tests, tests confirming the presence or absence of the ion in question. Our group, the "Metals", conducted multiple tests in hopes of achieving a solution change in color, or by the appearance of a precipitate, an insoluble material. We worked with 5 different water samples: a reference solution, that we labeled as #1 in our dish, a control sample, #2, a tap water sample, #3, a distilled water sample, #4, and an ocean water sample, #5. The reference solution is a solution with a strong presence of the ion being tested; therefore, it was different each time we refilled our dish depending on the ion being tested for, and it acted as an example of what the precipitate should look like when the ion is present in the water sample. However, it is important to note that the precipitate of the other samples did not mimic the precipitate of the reference solution, for the other samples did not contain the same high quantities of the ions being tested for. In order to test for the calcium ion, (Ca^2+), our reference solution was calcium chloride, or (CaCl^2), and we added about three drops of sodium carbonate (Na2CO3) to each of the five liquid-holding wells. To test for the iron(III) ion, (Fe^3+), our reference solution was ferric nitrate, (Fe3^+), and we added three drops of potassium thiocyanate (KSCN) to each well. To test for the chloride ion, (Cl^-), our reference solution was, again, calcium chloride, or (CaCl^2), and we added three drops of (toxic) silver nitrate (AgNO3) to each well. Finally, in our last test for the sulfate ion, (SO4^2-), our reference solution was ferrous sulfate, (SO4^2-) and we added three drops of barium chloride (BaCl2) to each well. We analyzed how the substances in the wells reacted with the test reagents and what (if any) precipitates were formed. The qualitative and quantitative tests determined what substance and how much of the particular substance was in our samples, and therefore, reflected the tests necessary to conduct for the "Snake River fish kill" case in Riverwood in order to reach a conclusion as to what contaminated the water and killed the fish.

Procedure:

Although this procedure was repetitive, it was very important to remain focused and fully alert as to what was going on, and what was being done to prepare and execute each test properly. First, my group prepared for the procedure by making a list numbered 1-5, corresponding to the numbers 1-5 labeled on the wellplate. We decided that well #1 would always be the reference solution would be held, well #2 would always be the control solution would be held, well #3 would be where tap water would be held, well #4 would be where distilled water would be held, and well #5 would be where ocean water would be held (we would fill each well only 3/4 full each time we filled them). By doing this, we remained organized throughout the entire process. In our first test, for the calcium ion (Ca^2+), we filled our reference solution well with calcium chloride (CaCl2). As I said before, reference solutions are meant to be rich in the ion being tested for, and calcium chloride (CaCl2) has plenty calcium ions, (Ca^2+). Our test reagent was sodium carbonate (Na2CO3); we added 3 drops of this to each well (from the opening at the top of its bottle) to see what precipitates and color changes (if any) each solution produced. The reference solution (calcium chloride) formed cloud-like precipitates that were spaced out from each other and floating in the solution, the control solution became urine yellow and formed a round, watery and milky precipitate at the bottom of the well, the tap water and distilled water solutions did not change, and the ocean water became slightly milkier in color. With this data we concluded that the Ca^2+ cations were present in the reference solution, control solution, and ocean water, but not in the tap water or distilled water.

Calcium ion test wellplate before sodium carbonate:

Calcium ion test wellplate after sodium carbonate:

Before we proceeded to the next test, we thoroughly rinsed and dried our wellplate.

In our second test, for the iron(III) ion (Fe^3+), we filled our reference solution well with ferris nitrate (Fe^3+). Our test reagent was potassium thiocyanate (KSCN); we added three drops of this to each well (from the opening at the top of its bottle) to see what precipitates and color changes each solution produced. Our originally yellow reference solution (ferris nitrate) became dense black in color without any visible precipitate and our control solution became blood red with a yellowish-brown rim. Nothing happened to the tap water, distilled water, or the ocean water. With these results, we concluded that the Fe^3+ cations were present in the reference solution and in the control solution, but not the tap, distilled, or ocean waters.

Iron(III) ion test wellplate before potassium thiocyanate:

Iron(III) ion test wellplate after potassium thiocyanate:

Before we proceeded to the next test, we thoroughly rinsed and dried our wellplate.

In our third test, the test for the chloride ion (Cl-), we filled our reference solution well with calcium chloride (CaCl2), the same reference solution we used for the calcium ion test. Our test reagent was [toxic] silver nitrate (AgNO3), that we added three drops of to each well (from the opening at the top of its bottle) in order to see what precipitates were produced and what color changes occurred. Our originally clear reference solution (calcium chloride) became a milky and foggy color with some watery cracks. It almost looked like thin frozen ice over a lake in the winter. The control solution formed a thick milk-colored rounded precipitate that sat at the bottom of the well with some control solution surrounding it. The tap water became cloudy in color, nothing happened to the distilled water, and the ocean water became even foggier and whiter than the reference solution. With these results, we concluded that the Cl- anions were present in the reference solution, the control solution, the tap water, and the ocean water, but not the distilled water.

Chloride ion test wellplate before silver nitrate:

Chloride ion test wellplate after silver nitrate:

Before we proceeded to the next test, we thoroughly rinsed and dried our wellplate.

In our fourth and final test, the test for the sulfate ion (SO4^2-), we filled our reference solution well with ferrous sulfate (SO4^2-). Our test reagent was barium chloride (BaCl2), and we added three drops of it to each well (from the opening at the top of its bottle) in order to analyze what color changes occurred and what precipitates were produced. Our originally muddy-red reference solution (ferrous sulfate) formed a milky precipitate that colored much of the solution, but left the muddy-red color visible underneath. The control solution became slightly cloudier and the ocean water developed a fairly heavy white fog, yet the tap water and distilled water remained the same. With these results, we concluded that the SO4^2- anions were present in the reference solution, the control solution, and the ocean water, but not the tap or distilled water.

Sulfate ion test wellplate before ferrous sulfate:

Sulfate ion test wellplate after ferrous sulfate:

Overview: Final Results:

Analyzing my data:
-Calcium Ion (Ca^2+) Test: Ca^2+ cations were present in the reference solution, control solution, and ocean water, but not in the tap water or distilled water.
-Iron(III) Ion (Fe^3+) Test: Fe^3+ cations were present in the reference solution and in the control solution, but not the tap, distilled, or ocean waters.
-Chloride Ion (Cl-) Test: Cl- anions were present in the reference solution, the control solution, the tap water, and the ocean water, but not the distilled water.
-Sulfate Ion (SO4^2-) Test: SO4^2- anions were present in the reference solution, the control solution, and the ocean water, but not the tap or distilled water.

Questions on page 45:
1) Why were a reference solution and a blank used in each test?

A reference solution was used in order to see what reaction happens between a test reagent and a solution with a high value of the ions being tested. A blank was used to show that there is no reaction when the ion being tested for is not present in the solution.

2) What are some possible problems associated with the use of qualitative tests?

A qualitative test is a test that can be done when, for example, a tap is turned on. One immediately sees the color of the water and can smell the water, and determine that it is clear and odorless. Even if it is clear and odorless, the water can be contaminated. Also, qualitative tests only test for the presence or absence of a substance, not the amount of the substance.

3) These tests cannot absolutely confirm the absence of an ion. Why?

Sometimes these tests cannot absolutely confirm the absence of an ion because sometimes, the ion is present in such small amounts, it is not detected through the tests.

4) How might your observations have changed if you had not cleaned your wells or stirring rods thoroughly after each test?

If we had not cleaned our wells or stirring rods, our samples would have been contaminated by the previous substances. This could lead to unwanted reactions, inaccurate reactions, or no reactions at all; therefore, we would not be collecting correct data.

ISBS #25-34, p. 51-52

25) Describe differences between qualitative and quantitative tests.

A qualitative test is a test executed in order to identify the presence or absence of a particular substance in a sample and a qualitative test is a test given to determine how much of a specific substance is present in a sample.

26) What is a confirming test?

A confirming test is a positive or negative test that confirms the presence or absence of the ion in question. In a this test, one must look for a change in solution color or for the appearance of a precipitate (an insoluble material). If neither appears, indicating a negative test, it means that the ion is either not present, or present in very low quantities.

27) In the water-testing investigation (pages 42-45), what was the purpose of
a. the reference solution
b. the distilled-water bank?

a. The reference solution is a solution where the ion being tested for is present; therefore, it is an example of what the precipitate looks like when the ion is present in the water sample. Usually, however, when the ion is present in the water sample, the precipitate is weaker, for it is usually present in smaller amounts than in the reference solution.
b. Since the distilled-water lacks all ions, it shows that when an ion-stimulating solution is added to it, there is no reaction, but the water remains clear instead.

28) Using the procedure outlined in the water-testing investigation, a student tests a sample of groundwater for iron and observes no color change. Should the student conclude that no iron is present? Explain your answer.

For this class procedure, the student should conclude that no iron is present given no color change. However, if the student was working in a real science laboratory, he or she should conduct further tests, for iron could be present in very small quantities; quantities too small to visibly see through the water-testing investigation lab project.

29) Given an unknown mixture,
a. what steps would you follow to classify it as a solution, a suspension, or a colloid?
b. describe how each step would help you to distinguish among the three types of mixtures.

a. I would stir the mixture and then let it sit for a couple minutes in order to see if any solid particles settle on the bottom of its container. I would also use the Tyndall effect in order to classify the mixture as a solution, a suspension, or a colloid.
b. By stirring and allowing the mixture to sit for a few minutes, I would be able to find out if solid particles large enough to settle to the bottom of the mixture are present in the mixture. These large solid particles would make the mixture an example of suspension. Additionally, a positive Tyndall effect would prove the mixture is a colloid (or a suspension depending on the size of the particles). A negative Tyndall effect would show me that the particles within the water sample are too small, and therefore mixed in and a part of the liquid; a solution.

30) Explain the possible risks in failing to follow the direction "shake before using" on the label of a medicine bottle.

By indicating the necessity to shake the bottle before using, the label shows that the medicine bottle contains a mixture that is an example of suspension. Many of the larger particles of the medicine would settle to the bottom, and therefore require shaking before consuming, in order to mix the particles and liquid together. If one failed to follow the direction "shake before using" on the label of a medicine bottle, they may not get enough, or overdose on components of the medication. The medicine would not work to its potential.

31) Why is it useful for element symbols to have international acceptance?

Since scientists all over the world study the same elements (since they are all present on earth), it is very useful for element symbols to have international acceptance. It is much easier for scientists to deal with the same symbols representing the same things. This contrasts with the US metric system; unlike the SI system that is used internationally, it is only useful in the United States.

32) Draw a model of a solution in which water is the solvent and oxygen gas (O2) is the solute.

See drawing.

33) Is it possible for water to be 100% "chemical free?" Explain.

No. Even with the expensive process of distillation, it is impossible to have 100% "chemical free" water. The atmospheric gases, nitrogen, oxygen, and carbon dioxide all exist in water to some degree; they always dissolve into water.

34) Compare the physical properties of water (H2O) with the physical properties of the elements from which it is composed.

Water, made of 1 hydrogen atom and 2 oxygen atoms has very different physical properties than its components. Although water exists as a liquid at room temperature, oxygen and hydrogen exist as a gas at room temperature. Also, because water is a polar molecule and the attraction of its positive and negative charges create a very strong surface tension, it has a much higher surface tension than the gaseous (and therefore spaced apart) atoms hydrogen and oxygen. Also, the freezing point of water is 0°C, the boiling point is 100°C, and the density is 1g/mL or 1g/cm^3. The freezing, boiling, and densities of hydrogen and oxygen are different.

How does testing water help us?

Testing water is an extremely helpful tool to human beings. Testing water can help us find out what is polluting, and therefore damaging our water supply; this pertains to the contaminated water in Riverwood, causing the Riverwood population's water shortage. Testing water helps to provide data and understanding as to what is and what is not in a body or supply of water, and therefore helps us to prevent and control the spread of contaminants in water that can result in famine, death, and water shortages. By knowing what is in the water through water testing, people can oversee and change the condition of water and make it better and safer to use.

ISBS #19-24, p. 51

19) For each of the following elements, identify the number of protons or electrons needed for an electrically neutral atom.

a. carbon: 6 protons, 6 electrons.
b. aluminum: 13 protons, 13 electrons.
c. lead: 82 protons, 82 electrons.
d. chlorine: 17 protons, 17 electrons.

20) Decide whether each of the following atoms is electrically neutral.

a. sulfur: 16 protons, 18 electrons. No.
b. iron: 26 protons, 24 electrons. No.
c. silver: 47 protons, 47 electrons. Yes.
d. iodine: 53 protons, 54 electrons. No.

21) Classify each of the following as an electrically neutral atom, an anion, or a cation.

a. O^2-: anion
b. Li: electrically neutral atom.
c. Cl: electrically neutral atom.
d. Ag+: cation.
e. Hg^2+: cation

22) For each particle in Question 32, indicate whether the electrical charge or lack of electrical charge was from a neutral atom gaining electrons, losing electrons, or neither.

a. O^2-: gaining electrons.
b. Li: neither.
c. Cl: neither.
d. Ag+: losing an electron.
e. Hg^2+: losing electrons.

23) Write the symbol and show the electrical charge (if any) on the following atoms or ions:

a. hydrogen with 1 proton and 1 electron: H
b. sodium with 11 protons and 10 electrons: Na+
c. chlorine with 17 protons and 18 electrons: Cl-
d. aluminum with 13 protons and 10 electrons: Al^3+

24) Write the name and formula for the ionic compound that can be formed from these cations and anions:

a. Potassium iodine: KI
b. Calcium sulfide: CaS
c. Iron(iii) bromide: Fe^3+(Br-)3: (Fe)(Br)3
d. Barium hydroxide: Ba^2+ + (OH-)2: (Ba)(OH)2
e. Ammonium phosphate: (NH4+)3 + PO4^3-: (NH4)3(PO4)
f. Aluminum oxide: (Al^3+)2 + (O^2-)3: (Al)2(O)3

Water Diary

Day 1 total water use: 546.1 L
Day 2 total water use: 2827.75 L
Day 3 total water use: 761.5 L

3 day total water use: 4135.35 L

Average daily water use: 1378.45 L

Average water use per person per day: 459.48333 L

A.7 #1-7, p.20-21

1) Calculate the total water volume (in liters) used by your household during the three days.

The total water volume used by my household during the three days is 4135.35 liters.

2) How much water (in liters) did one member of your household use, on average, in one day?

On average, one member of my household used 459.48333 liters of water in one day.

3) Compile the answers to Question 2 for all members of your class by creating a histogram.

4) What is the range of the average daily personal water use within your class?

The range is 734 liters.

5) Calculate the mean and median values for the class data. Which do you think is more representative of the data set- the mean or median value? That is, which is a better expression of central tendency for these data?

The mean, or average, is 446 liters. The median is 459 liters. The better expression of central tendency for these data is the mean. The mean is lower than the median. Since our data only shows one extremely high number for water use per person, it is more appropropriate to express a lower central tendency; a closer number to the rest of the values.

6) Compare your answer from Question 2 with the estimated average volume of water, which is 370 L, used daily by each person in the United States. What reasons can you propose to explain any difference between your value and the national average value?

Our class has a higher value than the US average of 370 liters used per person per day. Since we have a large range of 734 liters due to one household with a very high water use per person per day (950 L) and two households with very low water uses per person per day (216 L), our higher average than usual makes sense. Desipte this higer mean, one household in our class had an average per person per day of 363 L, closer to the nation's 370 L average value

7) Which is closer to the nation average (mean) for daily water use per person, your answer to Question 2 or the class average in Question 5? What reasons can you give to explain why that value is closer?

The class average, 446 liters, is closer than my answer to Question 2, (459.48333 liters), to the nation average of 370 L for daily water use per person. I feel as though the average per person in my household per day is higher because of my large garden. My garden is much larger than most, and therefore, uses a lot more water than standard-sized gardens. My garden uses over 2,076 L of water three times a week, with additional, unaccounted for hosing in between. This increases the the average value of water use in my house significantly, and makes it higher than both the class average of 446 L, and the nation average of 370 L per person per day.

Friday, June 17th Homework Assignment for Monday, June 20th.

ISBS, p. 50-52

13) See drawing.

14) Look at models.
a. Which models represent elements?
b. Which models represent compounds?

a. models i, ii, iv, and vi.
b. models iii and v.

15) What two pieces of information does a chemical formula provide?

A chemical formula provides each chemical substance (each element), and a subscript below each element representing how many atoms of each element is in one unit of the substance.

16) Name the elements and list the number of each atom indicated in the following substances:
a. phosphoric acid, H3PO4 (used in soft drinks)
b. sodium hydroxide, NaOH (found in some drain cleaners)
c. Sulfur dioxide, SO2 (an air pollutant)

a. Elements: Hydrogen, Phosphorus, and Oxygen. 3 Hydrogen atoms, 1 Phosphorus atoms, and 4 Oxygen atoms.
b. Elements: Sodium, Oxygen, Hydrogen. 1 Sodium atom, 1 Oxygen atom, and 1 Hydrogen atom.
c. Sulfur and Oxygen. 1 Sulfur atom and 2 Oxygen atoms.

17) See drawings.

18) Write chemical equations that represent the following word equations:
a. Baking soda (NaHCO3) reacts with hydrochloric acid (HCl) to produce sodium chloride, water, and carbon dioxide.
b. During respiration, one molecule of glucose, C6H12O6, combines with six molecules of oxygen gas to produce six molecules of carbon dioxide and six molecules of water.

a. NaHCO3 + Hcl → NaCl + H2O + CO2
b. C6H12O6 + O6 → 6 CO2 + 6 H2O

Thursday, June 16th Homework Assignment for Friday, June 17th.

Sketches for #1 and #3 (will also turn in a clear hard copy in class):

2) What kind of matter does the following model represent? Explain your answer.

This model represents a mixture of two two-atom compounds, with one compound having two atoms of one element, and the other compound having three atoms of one element. Since the compound with two atoms of one element, (probably H2O) floats above the other compound, this is a heterogeneous mixture and an example of suspension.

Unit I B.1-B.4 Vocabulary List

matter: anything that occupies space and has mass.

physical properties: properties that can be observed and measured without changing the chemical makeup of the substance.

density: the mass of material within a given volume. Water: 1g/mL, 1g/cm^3.

freezing/melting point (of water): 0°C or 32°F.

boiling point (of water): 100°C or 212°F.

aqueous solution: water-based solution.

surface tension: shows the strong intermolecular force that holds water molecules together.

mixture: when two or more substances combine and the substances retain their individual properties. The components of a mixture can be separated by physical means, such as filtration and adsorption.

heterogeneous mixture: composition is not the same, or uniform, throughout. (not evenly distributed).

suspension: solid particles are large enough to settle out or can be separated by using filtration.

Tyndall Effect: the scattering of light that indicates that small, solid particles, are still in the water.

collid: the small, solid particles that remain still in the water. Make liquid cloudy.

homogeneous mixture: a mixture that is uniform throughout. (evenly distributed).

solutions: homogeneous mixtures. (such as salt solution)

solute: salt in a salt solution.

solvent: water in a salt solution; dissolving agent.

particulate level: the level of atoms and molecules.

atoms: building blocks of matter. all matter is made of atoms.

element: matter that is made up of only one kind of atom.

compound: a substance that is composed of the atoms of two or more elements linked together chemically in certain fixed proportions.

chemical formulas: represent compounds.

substance: all elements and compounds. has a uniform and definite composition, as well as distinct properties.

molecule: the smallest unit of a molecular compound that retains the properties of that substance (smallest representation of the substance) ex: H2O molecule represents water.

chemical bonds: hold atoms of molecules together.

molecular compound: such as H2O

Wednesday, June 15th Homework Assignment for Thursday, June 16th.

Page 50:

1) What is a physical property?

A physical property is a property that can be observed and measured without changing the chemical makeup of the substance.

2) Identify three physical properties of water.

-Boiling point: 100°C or 212°F.
-Freezing point: 0°C or 32°F.
-Density: 1g/mL or 1g/cm^3

3) How does the density of solid water compare to the density of liquid water?

Since ice floats on the suface of liquid water, it is less dense.

4) Describe a setting where you might observe water as a solid, a liquid, and gas all at the same time.

One might observe water as a solid, a liquid, and a gas all at the same time in winter, at a waterfall in the mountains. The mountains surrounding the waterfall would be frozen, and as the powerful waterfall would crash onto the body of water below, many particles of water would rise into the air as water vapor, or gas. The cold mountains tend to be foggy as well, and fog is defined as a thick cloud of water droplets suspended in the atmosphere.

5) How are heterogeneous and homogeneous mixtures different?

In a heterogeneous mixture, the composition of it is not evenly distributed, or uniform, throughout. On the other hand, a homogeneous mixture is a mixture that is uniform throughout.

6) When gasoline and water are poured into the same container, they form two distinct layers. What do you need to know to predict which liquid will be on top?

In order to predict which liquid will be on top, it is necessary to know the densities of gasoline and water. The substance with the lower density would be on top.

7) Identify each of the following materials as a solution, a suspension, or a colloid.
a. A medicine accompanied by instructions to "shake before using"
b. Italian salad dressing
c. Mayonnaise.
d. A cola soft drink.
e. An oil-based paint.
f. Milk

a. A suspension
b. A suspension
c. A colloid
d. A solution
e. A suspension
f. A colloid

8) You notice beams of light passing into a darkened room through the blinds on a window. Does this demonstrate that air in the room is a solution, a suspension, or a collid? Explain.

Just like the Tyndall Effect, the beam of light passing into a darkened room through the blinds on a window reflects that since the light shines all the way through the room, it shows that there are small particles present in the air. This means that the air in the room is a collid; these particles are too small to come out through a filter. Although air appears to be a homogeneous mixture, the Tyndall Effect done through natural light proves that it is a heterogeneous mixture.

9) Drawing: separate paper.

10) Suppose you have a clear, red liquid mixture. A beam of light is observed as it passes through the mixture. Over a period of time, no particles settle to the bottom of the container. Classify this mixture as a solution, a collid, or a suspension, and provide evidence to justify your choice.

This mixture is a collid. Since with the Tyndall Effect, the beam of light passes through the mixture, it is evident that the particles within the mixture are small, but not actually part of the mixture. Therefore, this mixture can not be a solution. Furthermore, since over a period of time, no particles settle to the bottom of the container, the mixture is not a suspension.

11) Define the term substance and give two examples.

A substance is an element or compound with a uniform and definite composition, as well as distinct properties. The smallest substance is a molecule, which is composed of atoms that are held together by chemical bonds. Two examples of substances are water, H2O, and ammonia, NH3.

12) Classify each of the following substances as an element or a compound.
a. CO
b. Co
c. HCl
d. Mg
e. NaHCO3
f. NO
g. I2

a. Compound
b. Element
c. Compound
d. Element
e. Compound
f. Compound
g. Compound