assignment 18 water essentials

Learning Objectives

By the end of this section, you will be able to do the following:

• Describe the properties of water that are critical to maintaining life
• Explain why water is an excellent solvent
• Provide examples of water’s cohesive and adhesive properties
• Discuss the role of acids, bases, and buffers in homeostasis

Why do scientists spend time looking for water on other planets? Why is water so important? It is because water is essential to life as we know it. Water is one of the more abundant molecules and the one most critical to life on Earth. Water comprises approximately 60–70 percent of the human body. Without it, life as we know it simply would not exist.

The polarity of the water molecule and its resulting hydrogen bonding make water a unique substance with special properties that are intimately tied to the processes of life. Life originally evolved in a watery environment, and most of an organism’s cellular chemistry and metabolism occur inside the watery contents of the cell’s cytoplasm. Special properties of water are its high heat capacity and heat of vaporization, its ability to dissolve polar molecules, its cohesive and adhesive properties, and its dissociation into ions that leads to generating pH. Understanding these characteristics of water helps to elucidate its importance in maintaining life.

Water’s Polarity

One of water’s important properties is that it is composed of polar molecules: the hydrogen and oxygen within water molecules (H 2 O) form polar covalent bonds. While there is no net charge to a water molecule, water's polarity creates a slightly positive charge on hydrogen and a slightly negative charge on oxygen, contributing to water’s properties of attraction. Water generates charges because oxygen is more electronegative than hydrogen, making it more likely that a shared electron would be near the oxygen nucleus than the hydrogen nucleus, thus generating the partial negative charge near the oxygen .

As a result of water’s polarity, each water molecule attracts other water molecules because of the opposite charges between water molecules, forming hydrogen bonds. Water also attracts or is attracted to other polar molecules and ions. We call a polar substance that interacts readily with or dissolves in water hydrophilic (hydro- = “water”; -philic = “loving”). In contrast, nonpolar molecules such as oils and fats do not interact well with water, as Figure 2.13 shows. A good example of this is vinegar and oil salad dressing (an acidic water solution). We call such nonpolar compounds hydrophobic (hydro- = “water”; -phobic = “fearing”).

Water’s States: Gas, Liquid, and Solid

The formation of hydrogen bonds is an important quality of the liquid water that is crucial to life as we know it. As water molecules make hydrogen bonds with each other, water takes on some unique chemical characteristics compared to other liquids and, since living things have a high water content, understanding these chemical features is key to understanding life. In liquid water, hydrogen bonds constantly form and break as the water molecules slide past each other. The water molecules' motion (kinetic energy) causes the bonds to break due to the heat contained in the system. When the heat rises as water boils, the water molecules' higher kinetic energy causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor). Alternatively, when water temperature reduces and water freezes, the water molecules form a crystalline structure maintained by hydrogen bonding (there is not enough energy to break the hydrogen bonds) that makes ice less dense than liquid water, a phenomenon that we do not see when other liquids solidify.

Water’s lower density in its solid form is due to the way hydrogen bonds orient as they freeze: the water molecules push farther apart compared to liquid water. With most other liquids, solidification when the temperature drops includes lowering kinetic energy between molecules, allowing them to pack even more tightly than in liquid form and giving the solid a greater density than the liquid.

The lower density of ice, as Figure 2.14 depicts, causes it to float at the surface of liquid water, such as in an iceberg or ice cubes in a glass of water. In lakes and ponds, ice will form on the water's surface creating an insulating barrier that protects the animals and plant life in the pond from freezing. Without this insulating ice layer, plants and animals living in the pond would freeze in the solid block of ice and could not survive. The expansion of ice relative to liquid water causes the detrimental effect of freezing on living organisms. The ice crystals that form upon freezing rupture the delicate membranes essential for living cells to function, irreversibly damaging them. Cells can only survive freezing if another liquid like glycerol temporarily replaces the water in them.

Link to Learning

Click here to see a 3-D animation of an ice lattice structure.

Water’s High Heat Capacity

Water’s high heat capacity is a property that hydrogen bonding among water molecules causes. Water has the highest specific heat capacity of any liquid. We define specific heat as the amount of heat one gram of a substance must absorb or lose to change its temperature by one degree Celsius. For water, this amount is one calorie . It therefore takes water a long time to heat and a long time to cool. In fact, water's specific heat capacity is about five times more than that of sand. This explains why the land cools faster than the sea. Due to its high heat capacity, warm blooded animals use water to more evenly disperse heat in their bodies: it acts in a similar manner to a car’s cooling system, transporting heat from warm places to cool places, causing the body to maintain a more even temperature.

Water’s Heat of Vaporization

Water also has a high heat of vaporization , the amount of energy required to change one gram of a liquid substance to a gas. A considerable amount of heat energy (586 cal) is required to accomplish this change in water. This process occurs on the water's surface. As liquid water heats up, hydrogen bonding makes it difficult to separate the liquid water molecules from each other, which is required for it to enter its gaseous phase (steam). As a result, water acts as a heat sink or heat reservoir and requires much more heat to boil than does a liquid such as ethanol (grain alcohol), whose hydrogen bonding with other ethanol molecules is weaker than water’s hydrogen bonding. Eventually, as water reaches its boiling point of 100° Celsius (212° Fahrenheit), the heat is able to break the hydrogen bonds between the water molecules, and the kinetic energy (motion) between the water molecules allows them to escape from the liquid as a gas. Even when below its boiling point, water’s individual molecules acquire enough energy from other water molecules such that some surface water molecules can escape and vaporize: we call this process evaporation .

The fact that hydrogen bonds need to be broken for water to evaporate means that bonds use a substantial amount of energy in the process. As the water evaporates, energy is taken up by the process, cooling the environment where the evaporation is taking place. In many living organisms, including in humans, the evaporation of sweat, which is 90 percent water, allows the organism to cool so that it can maintain homeostasis of body temperature.

Water’s Solvent Properties

Since water is a polar molecule with slightly positive and slightly negative charges, ions and polar molecules can readily dissolve in it. Therefore, we refer to water as a solvent , a substance capable of dissolving other polar molecules and ionic compounds. The charges associated with these molecules will form hydrogen bonds with water, surrounding the particle with water molecules. We refer to this as a sphere of hydration , or a hydration shell, as Figure 2.15 illustrates and serves to keep the particles separated or dispersed in the water.

When we add ionic compounds to water, the individual ions react with the water molecules' polar regions and their ionic bonds are disrupted in the process of dissociation . Dissociation occurs when atoms or groups of atoms break off from molecules and form ions. Consider table salt (NaCl, or sodium chloride): when we add NaCl crystals to water, the NaCl molecules dissociate into Na + and Cl – ions, and spheres of hydration form around the ions, as Figure 2.15 illustrates. The partially negative charge of the water molecule’s oxygen surrounds the positively charged sodium ion. The hydrogen's partially positive charge on the water molecule surrounds the negatively charged chloride ion.

Water’s Cohesive and Adhesive Properties

Have you ever filled a glass of water to the very top and then slowly added a few more drops? Before it overflows, the water forms a dome-like shape above the rim of the glass. This water can stay above the glass because of the property of cohesion . In cohesion, water molecules are attracted to each other (because of hydrogen bonding), keeping the molecules together at the liquid-gas (water-air) interface, although there is no more room in the glass.

Cohesion allows for surface tension , the capacity of a substance to withstand rupturing when placed under tension or stress. This is also why water forms droplets when on a dry surface rather than flattening by gravity. When we place a small scrap of paper onto a water droplet, the paper floats on top even though paper is denser (heavier) than the water. Cohesion and surface tension keep the water molecules' hydrogen bonds intact and support the item floating on the top. It’s even possible to “float” a needle on top of a glass of water if you place it gently without breaking the surface tension, as Figure 2.16 shows.

These cohesive forces are related to water’s property of adhesion , or the attraction between water molecules and other molecules. This attraction is sometimes stronger than water’s cohesive forces, especially when the water is exposed to charged surfaces such as those on the inside of thin glass tubes known as capillary tubes. We observe adhesion when water “climbs” up the tube placed in a glass of water: notice that the water appears to be higher on the tube's sides than in the middle. This is because the water molecules are attracted to the capillary's charged glass walls more than they are to each other and therefore adhere to it. We call this type of adhesion capillary action , as Figure 2.17 illustrates.

Why are cohesive and adhesive forces important for life? Cohesive and adhesive forces are important for transporting water from the roots to the leaves in plants. These forces create a “pull” on the water column. This pull results from the tendency of water molecules evaporating on the plant's surface to stay connected to water molecules below them, and so they are pulled along. Plants use this natural phenomenon to help transport water from their roots to their leaves. Without these properties of water, plants would be unable to receive the water and the dissolved minerals they require. In another example, insects such as the water strider, as Figure 2.18 shows, use the water's surface tension to stay afloat on the water's surface layer and even mate there.

pH, Buffers, Acids, and Bases

The pH of a solution indicates its acidity or basicity.

You may have used litmus or pH paper, filter paper treated with a natural water-soluble dye for use as a pH indicator. It tests how much acid (acidity) or base (basicity) exists in a solution. You might have even used some to test whether the water in a swimming pool is properly treated. In both cases, the pH test measures hydrogen ions' concentration in a given solution.

Hydrogen ions spontaneously generate in pure water by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H + ) ions and hydroxide (OH - ) ions. While the hydroxide ions are kept in solution by their hydrogen bonding with other water molecules, the hydrogen ions, consisting of naked protons, immediately attract to un-ionized water molecules, forming hydronium ions (H 3 O + ). Still, by convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.

The concentration of hydrogen ions dissociating from pure water is 1 × 10 -7 moles H + ions per liter of water. Moles (mol) are a way to express the amount of a substance (which can be atoms, molecules, ions, etc.). One mole represents the atomic weight of a substance, expressed in grams, which equals the amount of the substance containing as many units as there are atoms in 12 grams of 12 C. Mathematically, one mole is equal to 6.02 × 10 23 particles of the substance. Therefore, 1 mole of water is equal to 6.02 × 10 23 water molecules. We calculate the pH as the negative of the base 10 logarithm of this concentration. The log10 of 1 × 10 -7 is -7.0, and the negative of this number (indicated by the “p” of “pH”) yields a pH of 7.0, which is also a neutral pH. The pH inside of human cells and blood are examples of two body areas where near-neutral pH is maintained.

Non-neutral pH readings result from dissolving acids or bases in water. Using the negative logarithm to generate positive integers, high concentrations of hydrogen ions yield a low pH number; whereas, low levels of hydrogen ions result in a high pH. An acid is a substance that increases hydrogen ions' (H + ) concentration in a solution, usually by having one of its hydrogen atoms dissociate. A base provides either hydroxide ions (OH – ) or other negatively charged ions that combine with hydrogen ions, reducing their concentration in the solution and thereby raising the pH. In cases where the base releases hydroxide ions, these ions bind to free hydrogen ions, generating new water molecules.

The stronger the acid, the more readily it donates H + . For example, hydrochloric acid (HCl) completely dissociates into hydrogen and chloride ions and is highly acidic; whereas the acids in tomato juice or vinegar do not completely dissociate and are weak acids. Conversely, strong bases are those substances that readily donate OH – or take up hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly alkaline and give up OH – rapidly when we place them in water, thereby raising the pH. An example of a weak basic solution is seawater, which has a pH near 8.0 This is close enough to a neutral pH that marine organisms have adapted in order to live and thrive in a saline environment.

The pH scale is, as we previously mentioned, an inverse logarithm and ranges from 0 to 14 ( Figure 2.19 ). Anything below 7.0 (ranging from 0.0 to 6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is alkaline. Extremes in pH in either direction from 7.0 are usually inhospitable to life. The pH inside cells (6.8) and the pH in the blood (7.4) are both very close to neutral. However, the environment in the stomach is highly acidic, with a pH of 1 to 2. As a result, how do stomach cells survive in such an acidic environment? How do they homeostatically maintain the near neutral pH inside them? The answer is that they cannot do it and are constantly dying. The stomach constantly produces new cells to replace dead ones, which stomach acids digest. Scientists estimate that the human body completely replaces the stomach lining every seven to ten days.

Watch this video for a straightforward explanation of pH and its logarithmic scale.

How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers readily absorb excess H + or OH – , keeping the body's pH carefully maintained in the narrow range required for survival. Maintaining a constant blood pH is critical to a person’s well-being. The buffer maintaining the pH of human blood involves carbonic acid (H 2 CO 3 ), bicarbonate ion (HCO 3 – ), and carbon dioxide (CO 2 ). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, it removes hydrogen ions and moderates pH changes. Similarly, as Figure 2.20 shows, excess carbonic acid can convert to carbon dioxide gas which we exhale through the lungs. This prevents too many free hydrogen ions from building up in the blood and dangerously reducing the blood’s pH. Likewise, if too much OH – enters into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to put survival in jeopardy.

Other examples of buffers are antacids that some people use to combat excess stomach acid. Many of these over-the-counter medications work in the same way as blood buffers, usually with at least one ion capable of absorbing hydrogen and moderating pH, bringing relief to those who suffer “heartburn” after eating. Water's unique properties that contribute to this capacity to balance pH—as well as water’s other characteristics—are essential to sustaining life on Earth.

To learn more about water, visit the U.S. Geological Survey Water Science for Schools All About Water! website.

• 1 W. Humphrey W., A. Dalke, and K. Schulten, “VMD—Visual Molecular Dynamics,” Journal of Molecular Graphics 14 (1996): 33-38.

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EESA01 Laboratory Manual: Introduction to Environmental Science

Chapter 11 laboratory 3: water quality, 11.1 introduction.

In recent lectures, we have discussed hydrology (the physical processes governing where water is and how it moves through the environment), freshwater resources, and water pollution. The purpose of this lab is for you to become familiar with how water quality (and by extension, pollution) is assessed. In this lab, you will sample water from Highland Creek, which runs behind the UTSC campus. You will then analyze a number of physical, chemical, and biological properties of your water sample and arrive at a “water quality index” to assess the relative quality of the water you sampled. The main learning objective of this lab is for you to become comfortable working in a laboratory environment. Another objective is to expose you to practical, employable skills used every day in a number of environment-related jobs. Additionally, these techniques are relatively inexpensive and their application, for example, in the developing world is entirely plausible.

Water is an essential chemical for health and development, but is denied to a large portion of the Earth’s population. Poor water quality can lead to numerous adverse effects on human health, including frequent episodes of diarrhea, cognitive damage, and water- borne diseases from pathogens, which in the most severe cases can lead to death.

Most forms of water pollution are not visible to the human eye, so scientists and technicians measure certain physical, chemical, and biological properties of water to characterize water quality. Biological properties include dissolved oxygen concentrations, and the presence of fecal coliform bacteria and other disease-causing organisms, which result mostly from animal waste in runoff and sewer overflows. Dissolved oxygen is a powerful indicator of ecosystem health, because surface waters that are low in dissolved oxygen are less capable of supporting aquatic life and may be indicative of eutrophication problems.

Common chemical properties include nutrient concentrations, pH, electrical conductivity, and hardness. Nutrients are elements (e.g., P) or compounds (e.g., NO 3 - ) that are consumed by organisms and are required for growth and survival. The mass of a particular nutrient in a given volume of water is its concentration. “Hard” water contains naturally high concentrations of calcium and magnesium ions, which prevent soap from lathering and leave chalky deposits behind when the water is heated or boiled. Hardness is often related to alkalinity, or the ability of water to buffer pH. pH is a measure of the concentration of hydrogen ions in a solution. This is measured on a pH scale from 0 to 14, which is logarithmic—this means that for every step in the scale, the hydrogen ion concentration changes tenfold. 7 is considered neutral, solutions below 7 are acidic, and solutions above 7 are basic. Lastly, electrical conductivity is a measure of a material’s ability to allow the conduction of an electric charge. It is usually measured in Siemens per meter, or S/m. In water, it is the dissolved materials in water that provide the electrical conductivity, thus electrical conductivity is directly related to salinity.

Among physical characteristics, suspended sediment concentration is the dry mass of particles that are suspended above the bed sediment (i.e., flowing in the water) per volume of water. If scientists can measure only one variable, they will often choose suspended sediment concentration because it tends to correlate with many other, more specific measures of water quality (like metals concentrations), and is therefore a good indicator of overall water quality. Fast-moving rivers that cut through arid or easily eroded landscapes, like the Yellow River in China, carry a heavy load of sediment and are turbid (or muddy-looking) as a result. Yet another important physical characteristic of water quality is temperature. High temperatures can interfere with some biological processes. For example, warmer water holds less dissolved oxygen, and inputs of energy to streams can affect cold water fish habitat, e.g. for rainbow trout.

By measuring a number of different water quality variables, one can generate an overall water quality score or index. Values obtained in the lab or field can be converted to “scores” using the figures in Appendix ( E ). If you find that your values fall beyond the scale of these figures, simply extrapolate the curves and estimate to the best of your ability what the “score” should be . These “scores” are then tabulated and totaled into a final score/index with the use of weighing factors (Table 11.1 ), and an overall water quality score/index and “condition” are obtained (Table 11.2 ). The appendix can be found here: Appendix ( E ).

11.2 Procedure for Laboratory 3

You should have a written account of what you have done in a given experiment. Making detailed notes during the lab will help you in answering the questions in the assignment.

• You will be hiking to Highland Creek, via the pathway behind the H-wing patio to obtain water samples. To save time, meet your TA on the H-wing patio at your scheduled lab time (Figure 11.1 ). Please read chapter 7 for more information on health and safety in the field. Once there, listen to your TA about sampling techniques and be sure to obtain two good and representative water samples; one for microbiological analysis (using your autoclaved bottle) and one for suspended sediment concentration analysis.

Figure 11.1: Left: Map of UTSC . Right: Picture of H-Wing patio (Photo taken by Andrew Apostoli).

Please be on time! Your teaching assistant will not wait for you if you are late.

• Depending on the number of students present, you will need to get into groups of 2-4 people. Your TA will inform you on the group size once everyone arrives at the creek. One person from each group will need to wear a wader. Waders will be provided at the scene, but please make sure you find the correct size (the correct size is based on your shoe size, so look under the boot of the wader). The other group members will record the data. Due to there being two activities, you may switch the person in the wader if you wish.

Plate 11.1: A wader is a green waterproof garment for the legs and body. Photo taken by Tom Meulendyk and Chai Chen. Click here for source

• While at the creek, use the multiparameter sonde to measure temperature, pH, electrical conductivity, and dissolved oxygen concentration (in both mg/L and % saturation) directly in the creek and record your data.

Plate 11.2: A student using a multiparameter sonde. Photo taken by Tom Meulendyk and Chai Chen. Click here for source

• For your microbiological/suspended sediment analysis, you will collect your water sample using the autoclaved bottle that is provided (Figure 11.2 ). Your TA and one of the lab technicians will provide instructions on how to collect your water sample.

Figure 11.2: A diagram of a suspended sediment collector. Click here for source

Potenial pictures/videos to add here about the methodology??

• Next, analyze your sample in the autoclaved bottle for the presence of coliform bacteria. This is easily done through the use of peel plates (see procedure in Appendix ( D ). On the back of your peel plate, write your name, tutorial #, and email address. Take a peel plate, and press back and lift the tab on top. Further pull this tab to expose the culture disk. Onto this culture disk you will pipet 1mL of your sample about 1-2 cm from the surface within 2-3 seconds—ensure that your pipette is vertical when you are pipetting your sample (Figure 11.3 ). Next, reseal the adhesive cover to keep the culture disk from being exposed. At this point, you will place your peel plate into the incubator in the lab and leave it to incubate for 24 hours at 35±1 \(^\circ\) C. You will be sent (via email) a photo of your peel plate so that you can assess the number of colonies that appear.

Figure 11.3: Instructions on how to use a micropipette. Click here for source

Make sure you have finished the following before you exit the laboratory:

• You either have a picture or a written document showing the values and units for the following parameters measured by the multiparameter sonde: temperature, pH, electrical conductivity, and dissolved oxygen concentration (in both mg/L and % saturation)
• You have recorded the dry mass of your filter
• You have recorded the volume of your sample water
• You have written your name and practical number on both the petri dish and coliform peel plate
• You have cleaned up your work area (sample water can be poured down the drain)

11.3 Assignment #3

Assignment 3 is 32 marks total – worth 10% of your final grade. Please submit a pdf version of your assignment on Quercus under “Lab 3 submission”. Your lab 3 assignment is due before the beginning of lab 4.

List the values you obtained for each of temperature ( \(^\circ\) C), pH, electrical conductivity ( \(\mu\) S/cm), dissolved oxygen (% saturation), suspended sediment concentration (mg/L), and coliform concentrations (colonies per 1 mL). (12 marks)

Using the figures in Appendix ( E ), convert your data into quality values and complete Table 11.1 . Using Table 11.2 , what is the overall water quality score? (10 marks)

Refer to the Guidelines for Canadian Recreational Water Quality – Third Edition (section 4.1). Assuming your coliform measurement is equal to the number of E. coli bacteria in your water sample, is the water in Highland Creek at the time of sampling safe enough for recreational/primary contact activity (e.g., swimming)? Explain. (4 marks)

Consider the scores in your water quality index. What changes in the watershed could improve these values? What should be the most immediate course of action given your data? You’ll possibly have to do some research. Make your response less than 10 lines long, cite sources, and don’t just copy out other text. See Appendix ( H ) for more information on APA referencing. (6 marks)

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