This lab tries to model the rate of radioactive decay using candy that goes by the M & M brand name. It explores the quantities of radioactive isotopes before and after decay. These candies have two-sided flat shapes with one side printed with the letter ‘M’ and the other side left blank.
In this lab, a hundred pieces of candy were put on a flat plate and made to lie flat at random. Another flat plate was used to cover the plate after which the plate was shaken twice to mix the candies. After that, the numbers of candy that had blank sides facing up were counted and placed separately. These were considered the decay products of radioactive isotopes. The pieces of candy that remained on the plate (with the printed faces facing upwards) were also counted and recorded. These were considered the radioactive isotopes. The first round of tossing was regarded as the first half-life, which is the amount of time taken for a radioactive substance to decay to half the original amount. The plate was covered and tossed a second time, which was considered the second half-life. Pieces of candy with the number side facing upwards were counted and added to the first lot to give the total of the decay products following the second half-life. The same procedure was repeated for the third and fourth half-lives.
It was observed that with each radioactive decay, the number of isotopes reduced while the number of decay products increased. For example, there were 100 isotopes and 0 decay products at zero half-life. This number changed to 49 and 51 after the first decay and 29 and 71 after the second decay for the isotopes and decay products. The third and fourth half-lives yielded 11 isotopes and 89 decay products and 5 isotopes and 95 decay products respectively.
These results reinforce the supposition that the rate of radioactive decay is dependent on the probability of disintegration (Snyder, 2003). It was observed that the quantity of decay product was a fraction of the original isotope. Therefore, the rate of radioactive decay is also proportional to the quantity of the radioactive isotope that is present at the beginning. Additionally, the findings show that the half-life is not influenced by the physical state, temperature, or any outer influence. However, the experimental finding were not accurate representatives of the theoretical rate of radioactive decay due to experimental errors that cannot be avoided. The numbers of candy that end up with their plain sides or numbered sides facing up are determined by chance (probability). It is almost impossible for two similar experiments to yield exact outcomes. However, the trend is the same if the experiments are repeated.
The Glass Where Pollution Begins
This lab examines the concept of pollution and tries to determine the presence of small quantities of matter in all substances. It shows that it is difficult to find pure substances because everything is contaminated with another substance to a certain extent. The purpose of this lab was to find out the level at which a contaminant can be detected in water.
Eight glasses of identical sizes were cleaned. The first glass, which was used as the standard measuring glass, was marked near the top. The other glasses were then labeled from 1 to 7. A teaspoon of salt was added to the standard glass followed by clean, warm tap water up to the mark. The salt was stirred to ensure uniform distribution and poured to glass 1. A mark was made approximately one-tenth from the surface of the fluid using a pen and a ruler. Thereafter, a tenth of the salt solution was then transferred back to the standard glass, filled with water up to the mark and stirred for uniform distribution of the salt water. A tenth of the water was then transferred to glass 2 (Snyder, 2003). The same steps were followed for the remaining five glasses. A similar procedure was repeated using a spoonful of salt instead of salt. In the end, water from each of the glasses was tasted starting from the most dilute to the highly concentrated to determine the concentration where the taste of salt could be tasted.
The salt was first tasted at 1/10 teaspoonful of salt per glass. The water in glass 1 could be termed as polluted with salt because it was not pure water. It was possible to drink the water in glass 7, which contained 1/1000000 teaspoonful of salt. However, the water in glass 7 was still polluted with salt because it contained salt albeit in a much smaller quantity than glass 1. Similar observations were made with the sugar solutions. The presence of sugar, however, was first tasted in glass 3, which contained 1/100 teaspoonful of sugar.
This experiment demonstrated that pure substances were rare and that pollution could still occur even when the presence of the contaminant is not obvious. This observation was corroborated by the fact that water in glass 7 tasted normally even though it still contained traces of salt. In addition, the presence of salt was first detected in glass 2 that contained 1/10 spoonful of salt. The findings of this lab were subject to the variability due to size of the glass, how much the teaspoon was filled with salt and the sensitivity of individual taste buds. An individual with more sensitive taste buds could taste the presence of salt at a much lower salt concentration than a person with less sensitive taste buds.
Lab – Squeezing the Air Out of a Bottle
This lab illustrates the effect of temperature and pressure on the volume of a substance using air as an example. Therefore, the main aim of the experiment was to squeeze air out of a bottle.
A clean empty glass bottle was placed in a freezer for thirty minutes. After that, the bottle was inverted with its mouth below the surface of a glass of water. The bottle was then wrapped with both hands without actual squeezing. A similar procedure was repeated by placing a coin on a slightly wet rim of a previously chilled bottle (Snyder, 2003). Both hands were then wrapped around the glass bottle.
It was observed that air bubbles formed in the glass containing water after holding the glass for some time. In the second scenario, the coin made a pinging noise after holding on to the glass bottle. Placing both hands around the glass transferred warmth from the hands to the bottle. The heat, in turn, caused the air inside the bottle to expand and escape through the mouth of the bottle hence the formation of bubbles in the water and the production of the pinging noise by the coin. The pinging noise was a result of pressure exerted on the coin by the escaping gas, which caused the coin to move slightly.
These findings demonstrate that at a constant pressure, the volume of air increases with an increase in temperature. Therefore, air is sensitive to changes in temperature. This observation is useful in the operations of a volume-gas thermometer. In the first demonstration, the air bubbles only appeared or the first few seconds then stopped. Certain experimental errors could have influenced the outcomes of this experiment. For example, the diameter of the mouth of the bottle could have played a substantial effect on the force of the bubbles. A bottle with a smaller opening could have produced better results since air would leave the bottle with greater force. In the second demonstration, the size and weight of the coin could have influenced the magnitude of the pinging sound. A larger and heavier coin was likely to produce a lower sound than a smaller and lighter coin. In addition, holding the chilled bottle for a prolonged period led to slight discomfort due to the freezing of hands hence bringing the experiment to an end. Another source of error was the initial temperature at which the bottle was frozen. Freezing the bottle at different temperatures for the same amount of time would yield different results in terms of the time taken to warm the air in the bottle before the formation of bubbles. Air in the bottle kept at lower temperatures was likely to take longer to warm and expand hence increasing the level of discomfort during the experiment.
Hydrogen is the first element of the periodic table with an atomic number of 1 and an atomic mass of approximately 1.00794 atomic mass units (Snyder, 2003). The name of this element is usually denoted by the letter ‘H’ and is obtained from two Greek names ‘hydro’ and ‘genes’ that mean ‘water-forming.’ Hydrogen belongs to group 1 because it has only one electron in its outermost (and only) energy shell. Similarly, hydrogen is a period one element because it has only one energy level. Hydrogen rarely exists as a free element in nature. Instead, it exists as a compound with elements such as oxygen, carbon and nitrogen. Consequently, it can be found almost everywhere because it occupies approximately nine tenths of the weight in the entire universe. In the earth’s crust and oceans, it is estimated that the hydrogen occupies approximately 1.40×103 milligrams per kilogram and 1.08×105 milligrams per liter respectively (Aldersay-Williams, 2011).
The History of Hydrogen
Many scientists have contributed to the elucidation of the physical and chemical properties of hydrogen. In 1671, Robert Boyle realized that hydrogen was liberated in the reaction between iron filings and dilute acids. In the year 1766, Henry Cavendish recognized hydrogen as an element and named it “the inflammable air from metals”. Subsequently, a different scientist named Antoine Lavoisier named the element Hydrogen. In 1839, Sir William Robert Grove realized that delivering an electric current through water, which contained hydrogen and water, destroyed the chemical bonds between hydrogen and oxygen. Conversely, inverting the process led to the production of electricity, which made hydrogen an important component of fuel cells.
Physical Attributes of Hydrogen
Hydrogen is a colorless and odorless gas, which causes the element to fall under non-metals. It has a density of 0.00008988 g/cm3, which makes it far much lighter than air (Snyder, 2003). Hydrogen melts at temperatures of -259.14 °C and boils at temperatures of -252.87 °C. Furthermore, hydrogen can develop hexagonal crystals under certain conditions. An alteration in the atomic weight of hydrogen confers it the ability to exist in more than one form as Deuterium or Tritium, which are known as isotopes of hydrogen. It reacts differently to various environmental conditions such as temperature and pressure. For example, the compression of hydrogen at elevated pressures causes the gas to liquefy for ease during storage.
Chemical Properties of Hydrogen
Hydrogen is an extremely reactive element that reacts chemically with several elements to form stable compounds. For example, its reaction with oxygen leads to the formation of water (H2O) while its reactions with carbon gives rise to a unique group of substances known as organic compounds. As a result, energy is stored in the chemical bond that form between hydrogen and other elements. Hydrogen gas combines with oxygen and chlorine to form volatile mixtures that can burst into flames in the presence of glints, heat or sunlight. Hydrogen can ignite extemporaneously in air at temperatures of 500 °C. Unadulterated hydrogen-oxygen blazes produce ultraviolet light. However, when the proportion of oxygen surpasses the amount of hydrogen in the mixture, the flame is almost invisible to the ordinary eye. Therefore, hydrogen leaks, which are extremely dangerous can only be identified by flame detectors. Hydrogen also reacts readily with all oxidizing elements. For example, at temperatures of approximately 25 oC, hydrogen can react without warning with chlorine and fluorine to yield the related hydrogen halides.
Commercial Use and Applications of Hydrogen
The physical and chemical properties of hydrogen facilitate it usage in various applications. In liquid form, hydrogen is used as a fuel especially in the propulsion of rockets. The reaction of hydrogen with carbon yields products such as methane and butane, which undergo combustion in engines to produce energy. Hydrogen is also used to generate electricity in fuel cells following its reaction with oxygen (Sorensen, 2012). Aside from fuel, hydrogen is a vital component of hydrogen peroxide that is often used as a bleach or wound sanitizer (Direct & Thacker, 2010). The weight and density of hydrogen makes it the most suitable gas in filling of balloons.
The manufacturing industry also use hydrogen as a raw material in the manufacture of other chemicals such as ammonia in the Haber process, methanol and cyclohexane (Scott, Denton, & Nicholls, 2013). Hydrogen has a tendency of filling up any vacant spaces within the structure of unsaturated organic compounds and for this reason, it is used in the hydrogenation of fats and oils. Hydrogen can produce hot flames with temperatures as high as 2000 oC hence is useful in welding under water.
In its purest form, hydrogen is considered one of the most environmental friendly energy sources. Nevertheless, environmental hazards may arise after using other fuel sources that contain hydrogen. For instance, partial combustion of gases such as methane and butane leads to the production of carbon monoxide, which is a toxic gas that leads to suffocation. However, this problem can be circumvented by ensuring that total combustion of such fuels occurs in a well-ventilated environment.
Aldersay-Williams, H. (2011). Periodic tales: The curious lives of elements. USA: HarperCollins.
Direct, J. & Thacker, E. (2010). The magic of hydrogen peroxide. USA: James Direct Inc.
Scott, R. B., Denton, W. H., Nicholls, C. M., (2013). Technology and uses of liquid nitrogen. Burlington, MA: Elsevier.
Snyder, C. H. (2003). The extraordinary chemistry of ordinary things (4th ed.). Hoboken, NJ: John Wiley & Sons.
Sorensen, B. (2012). Hydrogen and fuel cells: Emerging technologies and applications. Burlington, MA: Elsevier.