Daily pH Cycle and Ammonia Toxicity

World Aquaculture, 34(2): 20-21.

William A. Wurts, Ph.D.

Senior State Specialist for Aquaculture

Kentucky State University CEP at the UK Research and Education Center

P.O. Box 469

Princeton, KY  42445-0469

www.ca.uky.edu/wkrec/Wurtspage.htm

 

Ammonia is a nitrogen waste released by aquatic animals into the production pond environment.  It is a primary byproduct of protein metabolism.   Ammonia is excreted directly from the fish gill into the water.  Ammonia concentrations are usually at their highest late in the production season when biomass of the cultured species and the amount of protein fed are greatest.  Ammonia is toxic to aquatic life and toxicity is affected by pond pH.  Ammonia-nitrogen (NH₃-N) has a more toxic form at high pH and a less toxic form at low pH, un-ionized ammonia (NH₃) and ionized ammonia (NH₄⁺), respectively.  In addition, ammonia toxicity increases as temperature rises.

 

The measure of whether water is acidic, basic (alkaline) or neutral is known as pH.  A scale of 1 to 14 is traditionally used, which represents the negative logarithm of the hydrogen ion concentration.  A pH of 7.0 is neutral; above 7.0 is basic and below 7.0 is acidic; close to 7.0 is weak and far from 7.0 is strong.  It is a common perception that the pH of water is neutral and constant at a value of 7.0.   In an environment free of carbon dioxide, aquatic life, and compounds other than H₂O; pond pH would remain 7.0 or neutral.  However, this combination of conditions is unlikely to occur on our planet.  The pH of water is naturally acidic because the atmosphere contains carbon dioxide (CO₂). Carbon dioxide readily dissolves into water, raindrops and other sources of water exposed to air, forming a weak acid (H₂CO₃, carbonic acid).  Therefore, events in the aquatic environment that affect CO₂ concentrations also affect pH.  There are minerals in soil that can dissolve in water to create acidity and alkalinity as well.

 

Photosynthesis and Respiration

Pond CO₂ concentrations and pH, are affected by respiration and photosynthesis.  Carbon dioxide is released during respiration and consumed for photosynthesis.  As a result, pond pH varies throughout the day (Fig. 1).

 

 

 

 

 

 

 

 

 

     The plant members of the pond plankton community, phytoplankton, absorb CO₂ for photosynthetic production of sugar.  As daylight progressively intensifies, the rate of photosynthesis increases and so does the uptake of CO₂.  The removal of CO₂ reduces the concentration of carbonic acid, and pond pH rises.  Late in the production season, high waste nutrient concentrations can promote dense phytoplankton   blooms   which,   in   turn,   can

remove all of the CO₂ from pond water during photosynthesis.  This can cause the water to become alkaline with pH levels greater than 9.0.  Pond pH is highest late in the afternoon -- a few hours before sunset.

 

After sunset, photosynthesis and CO₂ uptake stop. However, respiration continues day and night.  During respiration, plants and animals consume oxygen to free the energy stored in food.  The end product of respiration is CO₂, which is released directly into the water.  As photosynthesis is halted by the absence of light, CO₂ begins to accumulate and the carbonic acid concentration increases.  The rising concentration of carbonic acid causes the pH to fall. Toward the end of the production season, the biomass and respiration of cultured animals and phytoplankton is high. Nighttime concentrations of CO₂, and therefore carbonic acid, can become excessive, lowering pH below 7.0.  As such, pond pH would be lowest an hour or two before sunrise.

 

Effects of pH on Ammonia Toxicity

The daily interplay of photosynthesis and respiration creates a cyclical change in pond pH.  Pond water becomes most acidic just before the period of darkness ends and most alkaline after several hours of daylight.  The presence of un-ionized ammonia, the toxic form, increases as pH rises and decreases as pH falls which causes ammonia to become more ionized.  The concentration of un-ionized ammonia in production ponds is lowest just before dawn and highest late in the afternoon.

 

This has significant implications for water quality monitoring, especially several weeks prior to harvest when fish biomass is greatest.  For example (Table 1), a producer measures water quality at 0400 hr.  The total NH₃-N concentration is 2.7 mg/L, pH is 7.0, and water temperature is 28 ^oC.  The farmer then cross-references these values with a standard, pH-temperature table and calculates the concentration   of   “un-ionized”   NH₃-N  to  be 0.019 mg/L.  The  producer decides to check water quality again at 1600 hr and finds that total   NH₃-N  is  still  2.7  mg/L.  But,  pH   and water temperature have risen to 9.0 and 30 ^oC.  After checking the reference table, the farmer discovers that the un-ionized NH₃-N concentration is now 1.2 mg/L.  An un-ionized NH₃-N level of 0.019 mg/L would be considered    acceptable    for   channel   catfish production.  However, the un-ionized NH₃-N concentration of 1.2 mg/L recorded at 1600 hr could be lethal to channel catfish within several hours.  Over a 12-hr period, the un-ionized ammonia concentration increased approximately 63-fold.  The temperature change accounts for less than 10% of the increase in toxicity while the rise in pH from 7.0 to 9.0 is responsible for more than 90%.

 

 

Table 1.  Amount of total ammonia-nitrogen (Tot/NH₃-N) present as un-ionized ammonia-nitrogen (UI/NH₃-N), for early morning and late afternoon pH and temperature measurements in a hypothetical production pond.

 

Time	Tot/NH3-N (mg/L)	Temp °C	pH	UI/NH3-N (mg/L)
0400 hr 1600 hr	2.7 2.7	28 30	7.0 9.0	0.019 1.2

 

 

Measuring pH and Ammonia

Photosynthesis and respiration have significant effects on pond pH.  Because those processes affect pH, ammonia toxicity is influenced also.  When monitoring water quality, it is important for producers to understand the daily shifts in pH and their impacts on un-ionized ammonia concentrations.  First, for ammonia-nitrogen measurements to be useful, pH and NH₃-N must be measured at the same time.  Second, a morning pH determination is not meaningful for assessing whether daily ammonia concentrations have reached unsafe levels in ponds.  To have any practical value for pond management decisions, NH₃-N and pH should be tested late in the afternoon.  A solid grasp of the pH cycle and its interrelationship with NH₃ is critical for the successful culture of any aquatic species.

Impacts OF EUTROPHICATION

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Fig 1. Red tide in Xiamen, in China’s Fujian Province, April 21, 2007. Red tides are nutrient-fueled blooms of phytoplankton that discolor water with their pigments. Several species are known to have toxic effects on marine life and pose a risk to human health through the consumption of exposed fish. Image Credit: Gu Liuzhang | Asianewsphotos

Excess nutrients in coastal waters can cause excessive growth of phytoplankton, microalgae (i.e. epiphytes and microphytes), and macroalgae (i.e. seaweed).
In turn, the increase in phytoplankton and algae can lead to more severe secondary impacts such as:

• Loss of subaquatic vegetation as excessive phytoplankton, microalgae and macroalgae growth reduce light penetration.
• Change in species composition and biomass of the benthic (bottom-dwelling) aquatic community, eventually leading to reduced species diversity and the dominance of gelatinous organisms such as jellyfish.
• Coral reef damage as increased nutrient levels favor algae growth over coral larvae. Coral growth is inhibited because the algae outcompetes coral larvae for available surfaces to grow.
• A shift in phytoplankton species composition, creating favorable conditions for the development of nuisance, toxic, or otherwise harmful algal blooms.
• Low dissolved oxygen and formation of hypoxic or “dead” zones (oxygen-depleted waters), which in turn can lead to ecosystem collapse.

The scientific community is increasing its knowledge of how eutrophication affects coastal ecosystems, yet the long-term implications of increased nutrient fluxes in our coastal waters are currently not entirely known or understood. We do know that eutrophication diminishes the ability of coastal ecosystems to provide valuable ecosystem services such as tourism, recreation, the provision of fish and shellfish for local communities, sportfishing, and commercial fisheries. In addition, eutrophication can lead to reductions in local and regional biodiversity.

Fig 2. A series of phytoplankton blooms. A cyanobacterial (blue-green algae) in the Baltic Sea (upper left). Red tide bloom (dinoflagellate) in the Sea of Japan (upper right). Cyanobacterial bloom in the St John’s River Estuary, Florida (lower left). Cyanobacteria-chlorophyte bloom in New Zealand (lower right)

Today nearly half of the world’s population lives within 60 kilometers of the coast, with many communities relying directly on coastal ecosystems for their livelihoods. This means that a significant portion of the world’s population is vulnerable to the effects of eutrophication in their local coastal ecosystems.
Two of the most acute and commonly recognized symptoms of eutrophication are harmful algal blooms and hypoxia.

Harmful Algal Blooms

Harmful algal blooms (Figures 1 and 2) can cause fish kills, human illness through shellfish poisoning, and death of marine mammals and shore birds. Harmful algal blooms are often referred to as “red tides” or “brown tides” because of the appearance of the water when these blooms occur. One red tide event, which occurred near Hong Kong in 1998, wiped out 90 percent of the entire stock of Hong Kong’s fish farms and resulted in an estimated economic loss of $40 million USD.

Fig 3. A menhaden (Brevoortia sp.) fish kill in August 2003 was caused by severe hypoxic conditions in Greenwich Bay, part of Narragansett Bay, Rhode Island, USA. Image Credit: Chris Deacutis | IAN

Hypoxia

Hypoxia, considered to be the most severe symptom of eutrophication, has escalated dramatically over the past 50 years, increasing from about 10 documented cases in 1960 to at least 169 in 2007. Hypoxia occurs when algae and other organisms die, sink to the bottom, and are decomposed by bacteria, using the available dissolved oxygen. Salinity and temperature differences between surface and subsurface waters lead to stratification, limiting oxygen replenishment from surface waters and creating conditions that can lead to the formation of a hypoxic or “dead” zone. The formatioin of dead zones can lead to fish kills (Figure 3) and benthic mortality. Because benthic organisms are bottom dwelling and cannot easily flee low-oxygen zones, they are often the most severely impacted.

Acid-Base Titrations

Modified 23:44, 21 Feb 2013 by Delmar

Acid-Base titrations are usually used to find the the amount of a known acidic or basic substance through acid base reactions. The analyte (titrand) is the solution with an unknown molarity. The reagent (titrant) is the solution with a known molarity that will react with the analyte. 

1. 1. Procedure
   1. 1.1. Choosing an Indicator
2. 2. Calculations
3. 3. References
4. 4. Outside Links

Procedure

The analyte is prepared by dissolving the substance being studied into a solution. The solution is usually placed in a flask for titration. A small amount of indicator is then added into the flask along with the analyte. The reagent is usually placed in a burette and slowly added to the analyte and indicator mixture. The amount of reagent used is recorded when the indicator causes a change in the color of the solution.

Some titrations requires the solution to be boiled due to the CO2 created from the acid-base reaction. The CO₂ forms carboinic acid (H₂CO₃) when dissolved in water. The carbonic acid then acts as a buffer, reducing the accuracy of data. After boiling most of the CO₂ will be removed from the solution allowing the solution to be titrated to a more accurate endpoint. The endpoint is the point where all of the analyte has be reacted with the reagent.

Choosing an Indicator

 A useful indicator has a strong color that changes quickly near its pKa. These traits are desirable so only a small amount of an indicator is needed. If a large amount of indicator is used, the indicator will effect the final pH, lowering the accuracy of the experiment. The indicator should also have a pKa value near the pH of the titration's endpoint. For example a analyte that is a weak base would require an indicator with a pKa less than 7. Choosing an indicator with a pKa near the endpoint's pH will also reduce error because the color change occurs sharply during the endpoint where the pH spikes, giving a more precise endpoint.

Figure 1: A Basic Titration Curve

Notice that this reaction is between a weak acid and a strong base so phenolphthalein with a pKa of 9.1 would be a better choice than methyl orange with a pKa of 3.8. If in this reaction we were to use methyl orange as the indicator color changes would occur all throughout the region highlighted in pink. The data obtained would be hard to determine due to the large range of color change, and inaccurate as the color change does not even lie with the endpoint region. Phenolphthalein on the other hand changes color rapidly near the endpoint allowing for more accurate data to be gathered.

Calculations

Multiply the volume of reagent added to get to the endpoint, with the molarity of the reagent to find the moles of reagent used. With the balanced equation of the acid-base reaction in question to find the moles of unknown substance. Then the original molarity can be calculated by dividing through with the initial volume.

For example an unknown molarity of HCl acts as the analyte. 50mL of it is placed into a flask and a 0.1M solution of NaOH will be the reagent. The endpoint's pH is 7 so litmus, with a pKa of 6.5 is chosen. The color of the solution changes when 10mL of 0.1M NaOH is added.

The balanced chemical equation related to is

        HCl(aq) + NaOH(aq) → H₂O(l) + Na⁺ + Cl^-

Or just the net ionic equation

        H⁺ + OH^- → H₂O_(l)

The following equation can then be derived

        

        X = 0.0010 mol of HCl

The molarity is now easily solved for

        0.0010 mol HCl / 0.050 L = 0.020M HCl

Uncertainties in Measurements

Modified 23:40, 21 Feb 2013 by Delmar

All measurements have a degree of uncertainty regardless of precision and accuracy. This is caused by two factors, the limitation of the measuring instrument (systematic error) and the skill of the experimenter making the measurements (random error). 

 1. Introduction
    2. Systematic and Random Error
    3. Calculating Error
    4. Methods of Reducing Error
    5. Outside links
    6. References
    7. Problems
    8. Solutions

Introduction

As the illustration below demonstrates:

(If you look at the slight curve the meniscus makes, the lowest part is where you actually record)

The graduated cylinder in the picture contains a certain amount of water to be measured. The amount of water is somewhere between 40ml and 50ml according to the marked lines. By checking to see where the botton of the meniscus lies, referencing the ten smaller lines, the amount of water lies between 44ml and 45ml. The next step is to estimate the uncertainty between 44ml and 45ml. Making an approximate guess, the level is less than 44.5 ml but greater than 44.0 ml. We then report that the measured amount is approximately 44.1ml. The graduated cylinder itself may be distorted such that the graduation marks contain inaccuracies providing readings slightly different from the actual volume of liquid present. 

Systematic and Random Error

When we use tools meant for measurement, we assume that they are correct and accurate, however measuring tools are not always right. In fact, they have errors that naturally occur called systematic errors. An example of a systematic error is a weighing scale. There are two specific types of systematic errors:

• Offset or Zero setting errors: When the measuring tools cannot read zero while the quantity measure is zero.
• Multiplier or Scale Factor Error: When the measuring tools consistently give readings to changes that are greater or smaller than the actual change.

Random error: Sometimes called human error, random error is determined by the experimenter's skill or ability to perform the experiment and read scientific measurements. These errors are random since the results yielded may be too high or low. Often random error determines the precision of the experiment or limits the precision. For example, if we were to time a revolution of a steadily rotating turnable, the random error would be the reaction time. Our reaction time would vary due to a delay in starting (an underestimate of the actual result) or a delay in stopping (an overestimate of the actual result).

Random vs Systematic Errors

In this experiment a series of shots is fired at a target. Random  errors are caused by anything that makes the shots inconsistent and arrive at the target at random different points. For example, the shooter has an unsteady hand or a change in the environment may distort the shooter's view. These errors would result in the scattering of shots shown by the right target in the figures to the left. A systematic error, on the other hand, would include consistent errors that always arise. For example, the gun may be misaligned or there may be some other type of technical problem with the gun. This type of error would yield a pattern similar to the left target with shots deviating roughly the same amount from the center area. 

 

When measuring a defined length with a ruler, there is a source of uncertainty and the measurement may need estimation or rounding between two points. When doing this estimation, it is possible to over estimate and under estimate the measured value, meaning there is a possibility for random error. Also, the ruler itself may be too short or too long causing a systematic error. For example, the illustration to the right shows a pencil whose length lies between 25cm and 26cm. With an intermediate mark, the ruler shows in greater detail that the pencil length lies somewhere between 25.5cm and 26cm. Therefore, one may reasonably approximate that the length of the pencil is 25.7cm. The presence of a systematic error, however, would likely be more subtle than a random error because the environment may affect the ruler in a difficult to notice way or the ruler itself may have slightly inaccurate markers.

Precision vs. Accuracy

Precision is often referred to as reproducibility or repeatability. For example, consider the precision with which the golf balls are shot in the figures below. A set of shots that are only precise would mean you are able to cluster your shots near each other on the green but you are not reaching your goal, which would be to get the golf balls into the hole. This concept is illustrated in the left picture of the two figures below. Accuracy, on the other hand, is how close a value is to the true or accepted value. The picture to the right demonstrates accuracy showing that the balls all get into the hypothetically large hole but are all at different corners of the hole. Therefore, the shots are not precise since they are relatively spread out but they are accurate because they all reached the hole. To sum up this concept, accuracy is the ability to hit the desired target area or measured value while precision is the agreement of shots or measured values with each other but not with the intended target or value. 

Calculating Error

Since equipment used in an experiment can only report a measured value with a certain degree of accuracy, calculating the extent to which a measurement deviates from the value accepted by the scientific community is often helpful in gauging the accuracy of equipment. Such a calculation is referred to as the percent error of a measurement and is represented by the following formula: 

Percent Error = [(Experimental Result - Accepted value) / Accepted Value] X 100%

Consider this real world example to understand the role of percent error calculations in determining the accuracy of measuring equipment: A toy company that ships its products around the world must calculate fuel costs associated with transporting the weight of their standard 2 by 3 foot box. To predict shipping costs and create a reasonable budget, the company must obtain accurate mass measurements of their boxes. The accepted mass of a standard box is 0.525 kg. The company measures a sample of three dozen boxes with a sophisticated electronic scale and an analog scale each yielding an average mass of 0.531 kg and 0.49 kg, respectively. A calculation of percent error for each device yields the following results: 

Percent Error of Electronic Scale = [(0.531kg - 0.525kg) / 0.525kg] X 100% = 1.14 %

Percent Error of Analog Scale = [(0.49kg - 0.525kg) / 0.525kg] X100% = -6.67%

Immediately, one notices that the electronic scale yields a far more accurate measurement with a percent error almost six times lower than the measurement obtained from the analog scale. Also note that percent error may take on a negative value as illustrated by the calculation for the analog scale. This simply indicates that the measured average lies 6.67% below the accepted value. Conversely, a positive percent error indicates that the measured average is higher than the accepted value. 

Methods of Reducing Error

While inaccuracies in measurement may arise from the systematic error of equipment or random error of the experimenter, there are methods that can be employed to reduce error:

Weighing by difference: Mass is an important measurement in many experiments and it is critical for labs to reduce error in mass measurements whenever possible. A simple way of reducing the systematic error of electronic balances commonly found in labs is to weigh masses by difference. This procedure entails the following:

1) finding the mass of both the desired material and the container holding the material,

2) transferring an approximate amount of the material to another container,

3) remeasuring the mass of the original container, and

4) calculating the mass of the removed sample by taking the difference between the initial and final weights of the original container. The following formula illustrates the procedure used for weighing by difference: 

(mass of container + mass of material) - (mass of container + mass of material after removing material) = mass of removed material

While most electronic balances have a "tare" or "zero" function that allows one to automatically calculate a mass by difference, equipment can be faulty so it is important to remember the fundamental logic behind weighing by difference.

Averaging Results: Since the accuracy of measurements are limited in part to the capacity of an experimenter to interpret their equipment, it makes sense that the average of several trials would be taken rather than a single trial. The reasoning behind averaging results is that an error of a measured value that falls below the actual value may be accounted for by averaging with an error that is above the actual value. By performing a series of trials (the more trials the more accurate the averaged result), an experimenter can account for some of their random error and yield a measurement with higher accuracy. 

Calibrating Equipment: Just as random error can be reduced by averaging several trials, systematic error of equipment can be reduced by calibrating a measuring device. This usually entails comparing a standard device of well known accuracy to the second device requiring calibration. Additionally, procedures exist for different kinds of equipment that can reduce the systematic error of the device. For example, a typical buret in a lab may be used to carry out a titration involving neutralization of an acid and base. If the buret formerly held acid but must now hold a base, then it would benefit the experimenter to condition the buret with the base before carrying out the titration so that the buret may acclimate to the new substance and provide a more accurate reading. Such procedures, together with calibration, can reduce the systematic error of a device. 

Classification of Matter

Modified 17:13, 22 May 2013 by Delmar

Matter can be identified by its characteristic inertial and gravitational mass and the space that it occupies. On earth matter is commonly found in three different states: solid, liquid, and gas.

    1. Introduction
    2. Classifying Matter
        2.1. Examples of Homogeneous Mixtures, also known as Solutions
        2.2. Examples of Heterogeneous Mixtures
        2.3. Example of a Colloidal Mixture Whose Components Tend Not to Settle Out
    3. Separation of Mixtures
    4. States of Matter
        4.1. Solids
        4.2. Liquids
        4.3. Gases
    5. Other States of Matter
    6. Outside links
    7. References
    8. Problems
    9. Contributors

Introduction

A substance is a sample of matter whose physical and chemical properties are the same throughout the sample because the matter has a constant composition. It is common to see substances changing from one state of matter to another.  To differentiate the states of matter at least at a particle level, we look at the behavior of the particles within the substance. When substances change state, it is because the spacing between the particles of the substances is changing due to a gain or loss of energy.  For example, we all have probably observed that water can exist in three forms with different characteristic ways of behaving: the solid state (ice), liquid state (water), and gaseous state (water vapor and steam). Due to water's prevalence, we use it to exemplify and describe the three different states of matter. As ice is heated and the particles of matter that make up water gain energy, eventually the ice melts in to water that eventually boils and turns into steam.  

Before we examine the states of matter,  we will consider some ways samples of matter have been classified by those who have studied how matter behaves.

Classifying Matter

Evidence suggests that substances are made up of smaller particles that are ordinarily moving around. Some of those particles of matter can be split into smaller units using fairly strong heat or electricity into smaller rather uniform bits of matter called atoms. Atoms are the building blocks of elements.  Elements are all those substances that have not ever been decomposed or separated into any other substances through chemical reactions, by the application of heat, or by attempting to force an direct electric current through the sample.

Atoms in turn have been found to be made up of yet smaller units of matter called electrons, protons, and neutrons. 

Elements can be arranged into what is called the periodic table of elements based on observed similarities in chemical and physical properties among the different elements. When atoms of two or more elements come together and bond, a compound is formed. The compound formed can later be broken down into the pure substances that originally reacted to form it.

Compounds such as water are composed of smaller units of bonded atoms called molecules. Molecules of a compound are composed of the same proportion of elements as the compound as a whole since they are the smallest units of that compound. For example, every portion of a sample of water is composed of water molecules.  Each water molecule contains two hydrogen atoms and one oxygen atom, and so water as a whole has, in a combined state, twice as many hydrogen atoms as oxygen atoms..

Water can still consist of the same molecules, but its physical properties may change. For instance, water at a temperature below 0 degrees Celsius (32 degrees Fahrenheit) is ice, whereas water above the temperature of 100. degrees Celsius (212 degrees Fahrenheit) is a gas, water vapor. When matter changes from one state to another, temperature and pressure may be involved in the process and the density and other physical properties change. The temperature and pressure exerted on a sample of matter determines the resulting form of that the matter takes, whether solid, liquid, or gas.

Since the properties of compounds and elements are uniform,  they are classified as substances.  When two or more substances are mixed together, the result is called a mixture. Mixtures can be classified into two main categories: homogeneous and heterogeneous. A homogeneous mixture is one in which the composition of its constituents are uniformly mixed throughout. A homogeneous mixture in which on substance, the solute, dissolves completely in another substance, the solvent, may also be called a solution. Usually the solvent is a liquid, however the solute can be either a liquid, solid, or a gas. In a homogeneous solution, the particles of solute are spread evenly among the solvent particles and the extremely small particles of solute cannot be separated from the solvent by filtration through filter paper because the spaces between paper fibers are much greater than the size of the solute and solvent particles. Other examples of homogeneous mixtures include sugar water, which is the mixture of sucrose and water, and gasoline, which is a mixture of dozens of compounds.

A heterogeneous mixture is a nonuniform mixture in which the components separate and the composition varies. Unlike the homogeneous mixture, heterogeneous mixtures can be separated through physical processes. An example of a physical process used is filtration, which can easilty separate the sand from the water in a sand-water mixture by using a filter paper. Some more examples of heterogeneous mixtures include salad dressing, rocks, and oil and water mixtures. Heterogeneous mixtures involving at least one fluid are also called suspension mixtures and separate if they are left standing long enough. Consider the idea of mixing oil and water together. Regardless of the amount of time spent shaking the two together, eventually oil and water mixtures will separate with the oil rising to the top of the mixture due to its lower density.

Mixtures that fall between a solution and a heterogeneous mixture are called colloidal suspensions (or just colloids). A mixture is considered colloidal if it typically does not spontaneously separate or settle out as time passes and cannot be completely separated by filtering through a typical filter paper.  It turns out that a mixture is colloidal in its behavior if  one or more of its dimensions of length, width, or thickness is in the range of 1-1000 nm. A colloidal mixture can also be recognized by shining a beam of light through the mixture. If the mixture is colloidal, the beam of light will be partially scattered by the suspended nanometer sized particles and can be observed by the viewer. This is known as the Tyndall effect. In the case of the Tyndall effect, some of the light is scattered since the wavelengths of light in the visible range, about 400nm to 700 nm, are encountering suspended colloidal sized particles of about the same size.  In contrast, if the beam of light were passed through a solution, the observer standing at right angles to the direction of the beam would see no light being reflected from either the solute or solvent formula units that make up the solution because the particles of solute and solvent are so much smaller than the wavelength of the visible light being directed through the solution.

• Solutions: molecules ~0.1-2 nm in size
• Colloids: molecules ~ 2-1000 nm in size
• Suspensions: molecules greater than ~ 1000 nm in size

Examples of Homogeneous Mixtures, also known as Solutions

Filtered seawater is solution of the compounds of water, salt (sodium chloride), and other compounds.

Examples of Heterogeneous Mixtures

   

separation of sand and water             separation of salad dressing                 various mixtures within a rock

Example of a Colloidal Mixture Whose Components Tend Not to Settle Out

Figure 4: milk is a colloid of liquid butterfat globules suspended in water

 

Separation of Mixtures

Most substances are naturally found as mixtures, therefore it is up to the chemist to separate them into their natural components. One way to remove a substance is through the physical property of magnetism. For example, separating a mixture of iron and sulfur could be achieved because pieces of iron would be attracted to a magnet placed into the mixture, removing the iron from the remaining sulfur. Filtration is another way to separate mixtures. Through this process, a solid is separated from a liquid by passing through a fine pored barrier such as filter paper. Sand and water can be separated through this process, in which the sand would be trapped behind the filter paper and the water would strain through. Another example of filtration would be separating coffee grounds from the liquid coffee through filter paper.  Distillation is another technique to separate mixtures. By boiling a solution of a non-volatile solid disolved in a liquid in a flask, vapor from the lower boiling point solvent can be driven off from the solution by heat, be condensed back into the liquid phase as it comes in contact with cooler surfaces, and be collected in another container.  Thus a solution such as this may be separated into its original components, with the solvent collected in a separate flask and the solute left behind in the original distillation flask. An example of a solution being separated through distillation would be the distillation of a solution of copper(II) sulfate in water, in which the water would be boiled away and collected and the copper(II) sulfate would remain behind in the disllation flask.

The picture above depicts the equipment needed for a distillation process. The homogeneous mixture starts out in the left flask and is boiled. The vapor then travels down chilled tube on the right and condenses back into a liquid and drips into the flask.

States of Matter

Everything that is familiar to us in our daily lives - from the land we walk on, to the water we drink and the air we breathe - is based upon the states of matter called gases, liquids, and solids.  

Solids

When the temperature of a liquid is lowered to the freezing point of the substance (for water the freezing point is 0^oC), the movement of the particles slows with the spacing between the particles changing until the attractions between the particles lock the particles into a solid form. At the freezing point, the particles are closely packed together and tend to block the motions of each other. The attractions between the particles hold the particles tightly together so that the entire ensemble of particles takes on a fixed shape. The volume of the solid is constant and the shape of a solid is constant unless deformed by a sufficiently strong external force.  (Solids are thus unlike liquids whose particles are slightly less attracted to one another because the particles of a liquid are a bit further apart than those in the corresponding solid form of the same substance.)  In a solid the particles remain in a relatively fixed positions but continue to vibrate.  The vibrating particles in a solid do not completely stop moving and can slowly move into any voids that exist within the solid.

           

Figure 2: The diagram on the left represents a solid whose constituent particles are arranged in an orderly array, a crystal lattice. The image on the right is a ice cube. It has changed from liquid into a solid as a result of absorbing energy from its warmer environment. 

Liquids

When the temperature of a sample increases above the melting point of a solid, that sample can be found in the liquid state of matter. The particles in the liquid state are much closer together than those in the gaseous state, and still have a quite an attraction for each other as is apparent when droplets of liquid form. In this state, the weak attractive forces within the liquid are unable to hold the particles into a mass with a definite shape. Thus a liquid's shape takes on the shape of any particular container that holds it.  A liquid has a definite volume but not a definite shape. Compared to to the gaseous state there is less freedom of particle movement in the liquid state since the moving particles frequently are colliding with one another, and slip and slide over one another as a result of the attractive forces that still exist between the particles, and hold the particles of the liquid loosely together.  At a given temperature the volume of the liquid is constant and its volume typically only varies slightly with changes in temperature.

    

Figure 3: The diagram on the left represents a container patially filled with a liquid. The image on the right is of water being poured out of a glass. This shows that liquid water has no particular shape of its own.

Gases

In the gas phase, matter does not have a fixed volume or shape. This occurs because the molecules are widely separated with the spaces between the particles typically around ten times further apart in all three spatial directions, making the gas around 1000 times less dense than the corresponding liquid phase at the same temperature. (A phase is a uniform portion of mater.) As the temperature of a gas is increased, the particles to separate further from each other and move at faster speeds.  The particles in a gas move in a rather random and independent fashion, bouncing off each other and the walls of the container.  Being so far apart from one another, the particles of a real gas only weakly attract each other such that the gas has no ability to have a shape of its own. The extremely weak forces acting between the particles in a gas and the greater amount of space for the particles to move in results in almost independent motion of the moving, colliding particles. The particles freely range within any container in which they are put, filling its entire volume with the net result that the sides of the container determine the shape and volume of gas. If the container has an opening, the particles heading in the direction of the opening will escape with the result that the gas as a whole slowly flows out of the container.

      

Figure 4: The image on the left represents an enclosed container filled with gas. The images are meant to suggest that the gas particles in the container are moving freely and randomly in myriad directions.The image on the right shows condensing water forming from the water vapor that escaped from the container. 

Other States of Matter

Besides of the three classical states of matter, there are many other states of matter that share characteristics of one more of the classical states of matter. Most of these states of matter can be put into three categories according to the degrees in varying temperature. At room temperature, the states of matters include liquid crystal, amorphous solid, and magnetically ordered states. At low temperatures the states of matter include superconductors, superfluids, and Bose-Einstein condensate state of matter. At high temperatures the states of matter include, plasma and Quark-gluon plasma. These other states of matter are not typically studied in general chemistry.

Chemical Change vs. Physical Change

Modified 01:30, 14 Apr 2013 by Delmar

The difference between a physical reaction and a chemical reaction is composition. In a chemical reaction, there is a change in the composition of the substances in question; in a physical change there is a difference in the appearance, smell, or simple display of a sample of matter without a change in composition. Although we call them physical "reactions," no reaction is actually occurring. In order for a reaction to take place, there must be a change in the elemental composition of the substance in question. Thus, we shall simply refer to physical "reactions" as physical changes from now on. 

1. 1. Introduction
2. 2. Common Physical Changes    2.1. Texture
       2.2. Color
       2.3. Temperature
       2.4. Shape
       2.5. Change of State
   
   3. Physical Properties
   
       3.1. Luster
       3.2. Malleability
       3.3. Ability to be drawn into a thin wire
       3.4. Density
       3.5. Viscosity
   
   4. Common Chemical Changes
   
       4.1. Change in Temperature
       4.2. Change in Color
       4.3. Noticeable Odor
       4.4. Formation of a Precipitate
       4.5. Formation of Bubbles
   
   5. References
   6. Outside Links
   7. Problems
   8. Answers
   9. Contributors

Introduction

Physical changes are limited to changes that result in a difference in display without changing the composition. Some common changes (but not limited to) are: 

• Texture
• Color
• Temperature
• Shape
• Change of State (Boiling Point and Melting Point are significant factors in determining this change.)

Physical properties include many other aspects of a substance. The following are (but not limited to) physical properties.

• Luster
• Malleability
• Ability to be drawn into a thin wire
• Density
• Viscosity
• Solubility
• Mass
• Volume 

Any change in these physical properties is referred to as a physical change. For further information, please refer to Properties of Matter.
Chemical changes, on the other hand, are quite different. A chemical change occurs when the substance's composition is changed. When bonds are broken and new ones are formed a chemical change occurs. The following are indicators of chemical changes:

• Change in Temperature
• Change in Color
• Noticeable Odor (after reaction has begun)
• Formation of a Precipitate
• Formation of Bubbles

Note: When two or more reactants are mixed and a change in temperature, color, etc. is noticed, a chemical reaction is probably occurring. These are not definite indicators; a chemical reaction may not be occurring. A change  in color is not always a chemical change. If one were to change the color of a substance in a non-chemical reaction scenario, such as painting a car, the change is physical and not chemical. This is because the composition of the car has not changed. Proceed with caution.

Common Physical Changes

Texture

The texture of a substance can differ with a physical change. For example, if a piece of wood was sanded, waxed, and polished, it would have a very different texture than it initially had as a rough piece of wood.

Rough Wood                                                          Finished Wood

   

As you can see, the texture of the finished wood is much smoother than the initial grainy wood.

Color

The changing of color of a substance is not necessarily an indicator of a chemical change. For example, changing the color of a metal does not change its physical properties. However, in a chemical reaction, a color change is usually an indicator that a reaction is occurring. The guy painting the metal car is not changing the composition of the metallic substance.

Painting a Car

Temperature

Although we cannot see temperature change, unless if a change of state is occurring, it is a physical change.

Hot Pan

One cannot see the pan physically changing shape, color, texture, or any of the other physical properties. However, if one were to touch the pan, it would be incredibly hot and could cause a burn. Sitting idle in a cupboard, this pan would be cold. One cannot assess this change only through visual exposure; the use of a thermometer or other instrument is necessary.

Shape

The shape of an object can be changed and the object will still remain true to its chemical composition. For example, if one were to fold money, as shown by the figure below, the money is still chemically the same.

Origami Money

Change of State

The change of state is likewise a physical change. In this scenario, one can observe a number of physical properties changing, such as viscosity and shape. As ice turns into water, it doesn't retain a solid shape and now becomes a viscous fluid.
The physical "reaction" for the change of ice into liquid water is:

H₂O (s) → H₂O (l)

Ice Melting

The following are the changes of state:

Solid→Liquid	Melting
Liquid→Gas	Vaporization
Liquid→Solid	Freezing
Gas→Liquid	Condensation
Solid→Gas	Sublimation

If heat is added to a substance, such as in melting, vaporization, and sublimation, the process is endothermic. In this instance, heat is increasing the speed of the molecules causing them move faster.

If heat is removed from a substance, such as in freezing and condensation, then process is exothermic. In this instance, heat is decreasing the speed of the molecules causing them move slower.

Physical Properties

Luster

The luster of an element is defined as the way it reacts to light. Luster is a quality of a metal. Almost all of the metals, transition metals, and metalloids are lusterous. The non-metals and gases are not lusterous. For example, oxygen and bromine are not lusterous. Shown below is are lusterous paper clips.

Lusterous Paperclips

Malleability

Malleability is also a quality of metals. Metals are said to be malleable. This means that the metals can deform under an amount of stress. For example, if you can hit a metal with a mallet and it deforms, it is malleable. Also, a paperclip can be shaped with bare hands.

Bent Paperclip

The image shows the malleability of a certain metal as stress is applied to it.

Ability to be drawn into a thin wire

In materials science, this property is called ductility. For example, raw copper can be obtained and it can be purified and wrapped into a cord. Once again, this property is characteristic of mainly metals, nonmetals do not possess this quality.

Copper Wire

Density

The density of an object is its mass divided by its volume (d=m/v). A substance will have a higher density if it has more mass in a fixed amount of volume. For example, take a ball of metal, roughly the size of a baseball, compressed from raw metal. Compare this to a baseball made of paper. The baseball made of metal has a much greater weight to it in the same amount of volume. Therefore the baseball made out of metal has a much higher density. The density of an object will also determine whether it will sink or float in a particular chemical. Water for example has a density of 1g/cm³. Any substance with a density lower than that will float, while any substance with a density above that will sink.

Oil Sinking in a Glass of Water

Viscosity

Viscosity is defined to be the resistance to deformation of a particular chemical substance when a force is applied to it. In the example below, one can see two cubes falling into two different test tubes. The upper substance shows a violent reaction to the dropping of the cube. The lower substance simply engulfs it slowly without much reaction. The upper substance has a lower viscosity relative to the lower substance, which has very high viscosity. One may even think of viscosity in terms of thickness. The substance with more thickness has higher viscosity than a substance that is deemed "thin." Water has a lower viscosity than honey or magma, which have relatively high viscosities. 

Viscosity of Fluids

Common Chemical Changes

The follow are all indicators of chemical reactions. For further information on chemical reactions, please refer to Chemical Reactions.

Change in Temperature

A change in temperature is characteristic of a chemical change. During an experiment, one could dip a thermometer into a beaker or Erlenmeyer Flask to verify a temperature change. If temperature increases, as it does in most reactions, a chemical change is likely to be occurring. This is different from the physical temperature change. During a physical temperature change, one substance, such as water is being heated. However, in this case, one compound is mixed in with another, and these reactants produce a product. When the reactants are mixed, the temperature change caused by the reaction is an indicator of a chemical change.

Violent Reaction (Fireworks) that has Heat as a Product

In the reaction above, a hot chemical is shooting out of an Erlenmeyer flask. The temperature after the reactants were mixed increased incredibly in this reaction.

The following reaction has heat as a product. It is exothermic. However, do not be fooled; every exothermic change is not a chemical change. The melting of an ice cube, which is also exothermic, is a change in a physical property and not composition. Thus, it is a physical change.

Reaction: Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + Heat

Change in Color

A change in color is also another characteristic of a chemical reaction taking place. For example, if one were to observe the rusting of metal over time, one would realized that the metal has changed color and turned orange. This change in color is evidence of a chemical reaction. However, one must be careful; sometimes a change in color is simply the mixing of two colors, but no real change in the composition of the substances in question. 

Metal Rusting

The reaction above is that of the rusting of iron.
Reaction: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃

Noticeable Odor

When two or more compounds or elements are mixed and a scent or odor is present, a chemical reaction has taken place. For example, when an egg begins to smell, (a rotten egg) a chemical reaction has taken place. This is the result of a chemical decomposition. 

Spoiled Egg

Formation of a Precipitate

The formation of a precipitate may be one of the most common signs of a chemical reaction taking place. A precipitate is defined to be a solid that forms inside of a solution or another solid. Precipitates should not be confused with suspensions, which are solutions that are homogeneous fluids with particles floating about in them. For instance, when a soluble carbonate reacts with Barium, a Barium Carbonate precipitate can be observed.

Test Tube

Reaction: Ba²⁺(aq) + CO₃^2-(aq) → BaCO₃(s)
For further information, please refer to Classification of Matter.

Formation of Bubbles

The formation of bubbles, or rather a gas, is another indicator of a chemical reaction taking place. When bubbles form, a temperature change could also be taking place. Temperature change and formation of bubbles often occur together. For example, in the following image, one can see a gas spewing. This is the formation of a gas. 

Gas Formation

However, most reactions are much more subtle. For instance, if the following reaction occurs, one may notice Carbon Dioxide bubbles forming. If there is enough Hydrochloric Acid, bubbles are visible. If there isn't, one can't readily notice the change.

Reaction: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

References

1. Chang, Raymond. General Chemistry: the Essential Concepts. Boston, MA: McGraw-Hill Higher Education, 2006. Print.    
2. Chemistry for Dummies. For Dummies, 2008. Print.
3. Petrucci, Ralph H. General Chemistry Principles and Modern Applications. Upper Saddle River, NJ: Pearson/Prentice Hall, 2007. Print.

Outside Links

• http://en.wikipedia.org/wiki/Physical_property
• http://en.wikipedia.org/wiki/Viscosity
• http://www.chemtutor.com/ms.htm#ionize
• http://en.wikipedia.org/wiki/Density
• http://en.wikipedia.org/wiki/Precipitate
• http://en.wikipedia.org/wiki/Chemical_reaction
• http://www.sxc.hu

All images are courtesy of http://www.sxc.hu, which provides royalty free images that are free to be copied without restrictions.
The viscosity image is also free to be duplicated as per permission of author on wikipedia.com.

Problems

1. Which of the following is a chemical reaction?

a) Freezing liquid Mercury

b) Adding yellow to blue to make green

c) Cutting a piece of paper into two pieces

d) Dropping a sliced orange into a vat of Sodium Hydroxide 

e) Filling a balloon with natural air

2. Which of the following is a physical reaction?

a) Shattering Glass with a baseball

b) Corroding Metal

c) Fireworks Exploding

d) Lighting a match

e) Baking a cake

3. Which of the following is a chemical reaction?

a) Painting a wall blue

b) A bicycle rusting

c) Ice cream melting

d) Scratching a key across a desk

e) Making a sand castle

4. Which of the following is a physical reaction?

a) Frying an egg

b) Digesting carrots

c) A Macbook falling out of a window

d) Creating ATP in the human body

e) Dropping a fizzy tablet into a glass of water

5. Write C for Chemical Reaction or P for Physical Reaction.

a) Burning Leaves

b) Cutting Diamonds

c) Crushing a pencil

d) Salivary Amylase breaking down food in the mouth

e) Salt mixing in with water

Answers

1. D
2. A
3. B
4. C
5.     a) C

b) P

c) P

d) C

e) Neither. This is one of the gray areas of chemical change and physical change. Although the salt has dissociated into Sodium and Chloride ions, it is still salt in water. Salt, initially is actually just a conglomerate of sodium and chloride ions and by dissociating them, just the arrangement of the ions has changed. Please click here for more information.

Contributors

• Mandeep Sohal (mdsohal@ucdavis.edu)

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About Eutrophication

Within the past 50 years, eutrophication—the over-enrichment of water by nutrients such as nitrogen phosphorus—has emerged as one of the leading causes of water quality impairment. The two most acute symptoms of eutrophication are hypoxia (or oxygen depletion) and harmful algal blooms, which among other things can destroy aquatic life in affected areas.

Fig 1. The eutrophication process and subsequent formation of sea-bottom hypoxia in coastal waters 

Fig 2. This divided body of water shows the remarkable difference between mesotrophic (moderately enriched) (upper basin) and eutrophic water (lower basin). Image Credit: Fisheries and Oceans Canada

The rise in eutrophic and hypoxic events has been attributed to the rapid increase in intensive agricultural practices, industrial activities, and population growth which together have increased nitrogen and phosphorus flows in the environment. The Millenium Ecosystem Assessment (MA) found that human activities have resulted in the near doubling of nitrogen and tripling of phosphorus flows to the environment when compared to natural values. Before nutrients—nitrogen in particular—are delivered to coastal ecosystems, they pass through a variety of terrestrial and freshwater ecosystems, causing other environmental problems such as freshwater quality impairments, acid rain, the formation of greenhouse gases, shifts in community food webs, and a loss of biodiversity.
Once nutrients reach coastal systems, they can trigger a number of responses within the ecosystem. The initial impacts of nutrient increases are the excessive growth of phytoplankton, microalgae (e.g., epiphytes and microphytes), and macroalgae (i.e., seaweed). These, in turn, can lead to other impacts such as: loss of subaquatic vegetation, change in species composition, coral reef damage, low dissolved oxygen, and the formation of dead zones (oxygen-depleted waters) that can lead to ecosystem collapse.