The benefit of heat absorption and retention in water cannot be overemphasized as it is majorly responsible for temperature regulation and stability in both the human body and the planet as a whole. The effect of impurities on the heat retention and absorption of freshwater was investigated with a view on identifying what impurity gave freshwater the best heat retention and absorption ability. The effect of concentrations of impurities on the heat absorption and retention capacity of water was also investigated. Two methods were used each for the determination of both heat retention and absorption of freshwater. Measured values (5g, 10g and 15g) of sugar samples was dissolved in a 100ml polypropylene beaker of water and kept in a freezer simultaneously till the solutions attained freezing point, with the temperature drop recorded at intervals of ten minutes with a digital thermometer. The beakers were removed simultaneously from the freezer with the temperature rise recorded till room temperature was attained. Measured values (5g, 10g and 15g) of sugar were each added to 100ml of water, and the solution heated to boiling point. Time taken for each sample to reach boiling point was also recorded. A cooling system was setup with the aid of a copper calorimeter and stirrer, to enable uniform temperature during the cooling process. A digital thermometer was used to record temperature drop at each ten minute interval till room temperature was attained. This was done for samples of Salt, Milk, Alum, Baking powder and CaC0₃. Graph of temperature against time was plotted using Microsoft Excel Spreadsheet, in which the rate of heat retention and absorption of freshwater was determined. The result shows that generally, impurities reduce both the heat retention and absorption capacity of water. Also, for all concentrations of impurities subjected to the freezing process (heat retention), Baking powder gave water the best heat retention ability in a time of 2hrs 50mins and Sugar was the least in 1hr 50mins. It was also, deduced that Baking powder and Salt gave freshwater a high heat absorption rate both in a time of 1hr 10mins, while Milk took the least time in absorbing heat in 2hrs 40mins.
TABLE OF CONTENTS
Title page i
Table of Content v-viii
List of figures ix-x
List of tables xi-xii
CHAPTER ONE – INTRODUCTION
1.0 INTRODUCTION 1
1.1. Physical and Chemical properties of water 2
1.1.1 Electrical conductivity of water 2
1.1.2 Dissolving capacity of water 3
1.1.3 PH of water 3
1.2 Distinct states at which water can exist 4
1.3 Density of water 5
1.4 Water at Excited State 5
1.5 Solubility of Different Materials in Freshwater 6
1.6 Boiling of freshwater 7
1.6.1 Boiling point of Freshwater 7
1.6.2 Factors that affect boiling point of a substance 7
1.7 Freezing of freshwater 8
1.7.1 Freezing point of freshwater 9
1.8 Anomalous nature of water 10
1.9 Water’s high heat capacity 11
1.10 Impurities 12
1.10.1 Effect of impurities on boiling point of a liquid 12
1.10.2 Effect of impurities on freezing point of a liquid 12
1.10.3 Sugar 13
1.10.4 Salt 14
1.10.5 Powdered Milk 14
1.10.6 Alum 15
1.10.7 Baking powder 15
1.10.8 CaC0₃ 16
1.11 Thermodynamics 16
1.12 The concept of heat 16
1.12.1 Heat Transfer 18
1.12.2 Methods of heat transfer 19
184.108.40.206 Conduction 19
220.127.116.11 Convection 20
18.104.22.168 Radiation 21
1.13 Importance and significance of the work 22
1.14 Aims and Objectives of the project 23
CHAPTER TWO – LITERATURE REVIEW
2.0 INTRODUCTION 24
CHAPTER THREE – MATERIALS AND METHODS
3.0 INTRODUCTION 27
3.1 MATERIALS USED 27
3.1.1 Mextech Digital Thermometer 28
3.2 METHODOLOGY 29
3.2.1 EFFECTS ON IMPURITIES ON FREEZING POINT OF WATER
(HEAT RETENTION) 29
3.2.2 EFFECTS OF IMPURITIES ON THE MELTING RATE OF
FRESHWATER (HEAT ABSORPTION) 30
3.2.3 EFFECT OF IMPURITIES ON THE BOILING POINT OF
WATER (HEAT ABSORPTION) 30
3.2.4 EFFECT OF IMPURITIES ON THE COOLING RATE OF WATER
(HEAT RETENTION) 31
CHAPTER FOUR – RESULTS AND DISCUSSIONS
4.1 TEMPERATURE CHANGES DURING FREEZING 32
4.2 TEMPERATURE CHANGES DURING MELTING 50
4.3 TEMPERATURE CHANGES DURING COOLING 68
CHAPTER FIVE – CONCLUSION AND RECOMMENDATION
5.0 CONCLUSION 86
EFFECT OF IMPURITIES ON THE HEAT ABSORPTION AND RENTENTION CAPACITY OF FRESHWATER
Water is a transparent fluid which forms the world’s streams, rivers, lakes, rain, and oceans. As a chemical compound, a single water molecule contains one atom of oxygen, and two atoms of hydrogen connected by covalent bonds.
Water covers 71% of the total Earth’s surface.96.5% of earth’s water is found in seas and oceans, 1.7% in groundwater, 1.7% in glaciers and ice caps of Antarctica and Greenland, a small fraction of water can be found in other large water bodies, and 0.001% in the air as vapor, clouds (formed of ice and liquid water suspended in air), and precipitation .Only 2.5% of the Earth’s water is freshwater, and 98.8% of that water is ice (excepting ice in clouds) and groundwater. Only less than 0.3% of all freshwater is in rivers, lakes, and the atmosphere (Gleick, P.H 1993).
Many substances dissolve in water and it is commonly referred to as the universal solvent. For this cause, water in nature and in use is rarely pure and some properties may vary from those of the pure substance. However, there are also many compounds that are essentially, if not completely insoluble in water. Pure Water has a boiling and melting point of 100⁰C and 0⁰C respectively. Water is the only common substance found naturally in all three common states of matter and it is essential for all life on Earth. Water makes up 55-78% of the human body, so its importance cannot be overemphasized.
A water molecule has no net charge because the number of positively charged protons equals the number of negatively charged electrons. However, because the hydrogen ends of the molecule have a slight positive charge and the oxygen end has a slight negative charge, it is called a polar molecule. The negative and positive ends of different water molecules slightly attract each other, forming hydrogen bonds. These hydrogen bonds are about twenty times weaker than the covalent bonds between hydrogen and oxygen.
1.1 Physical and Chemical Properties of Water
The polar nature of the water molecule and the hydrogen bonds are responsible for many of water’s unique physical and chemical properties.
1.1.1 Electrical Conductivity of Water
Most natural waters contain dissolved ions (atoms or molecules possessing a charge) derived from the water’s interaction with soil, bedrock, atmosphere, and biosphere. As a result of these ions, water is able to conduct electricity much better than it otherwise can: for example, sea water with its dissolved salts can conduct electricity about 100 times more readily than distilled water. In any case, the ability of water to conduct electricity is the reason for the warning labels that appear on most electrical appliances warning consumers not to operate them near water.
The electrical conductivity (or specific conductance) of water depends on the concentration and charge of the dissolved ions (Weast, Robert C. 2000). Because of this relationship, conductivity often is used as an indicator of the total dissolved solids (TDS) in the water. The TDS is an important chemical property of water that provides information regarding the water’s history of “evolution” (for instance, its movement through underground aquifers). Conductivity is only an estimate of TDS, however, because a given value of conductivity can be produced by several different combinations of ion concentration and charge.
1.1.2 Dissolving Capacity of Water
Water molecules also can be attracted to surfaces of minerals in soils and rocks. This attraction allows dissolving (dissolution) and other chemical reactions to occur so readily that water often is called the “universal solvent.” Because water can dissolve and carry a wide range of chemicals, minerals, and nutrients, it plays an essential role in almost all biological and geochemical processes.
1.1.3 PH of Water
One important chemical property of natural water that affects its ability to dissolve minerals and influence chemical reactions is its pH. The pH, which indicates the acidity of water, measures the abundance of positively charged hydrogen ions (H +), and is defined numerically as the negative logarithm of the concentration of H + ions. Because pH is measured on a logarithmic scale, the concentration of H+ ions is ten times greater in water with a pH of 5 than water with a pH of 6, for example.
Water with a pH greater than 0 and less than 7 is termed acidic; a pH equaling 7 is neutral at temperatures at the Earth’s surface; and a pH between 7 and 14 is termed alkaline (or basic). Distilled water is considered neutral, and has a pH of 7. Natural waters also can be neutral, but more often are either slightly acidic or slightly basic.
Water in some volcanic caldera lakes can be very acidic, with pH values sometimes less than 1. If a wristwatch were dropped into water this acidic, the watch would be damaged beyond recognition within minutes. Rain in unpolluted areas has a pH of about 5.6 due to the dissolution of carbon dioxide in the atmosphere. This slightly acidic nature enhances rain-water’s dissolving power. In some locations, rainwater’s acidity is greatly increased (the pH is lowered) by the absorption of certain atmospheric pollutants, causing what is called acid rain. Rainwater is neutralized by chemical reactions with minerals in soils and rocks so that the pH of most streams and lakes is between about 6.5 and 8.5. Aquatic organisms often are very sensitive to the pH of water. Below a pH of about 5, most fish will die.
1.2 Distinct States at Which Water Can Exist
Water can exist in three distinct states: water (liquid), steam (gas), and ice (solid). At sea level, ice melts (solid changes to a liquid) at 0°C (32°F) and boils (liquid changes to a gas) at 100°C (212°F). The temperature at which water changes from one state to another depends on atmospheric pressure. At the elevation of Denver, Colorado, where air pressure is about 17 percent lower than at sea level, water boils at about 94°C (201°F).
Boiling hot water thrown from a cup into very cold air will almost instantly freeze in midair and create a shower of ice crystals. At sufficiently high pressures and temperatures, water and steam are no longer distinct phases, and instead comprise a supercritical fluid. The critical temperature is 374°C (705°F) and the critical pressure is 22.06 mega-pascals (3,198.70 pounds-force per square inch), or the pressure reached at a depth of 2.2 kilometers (1.4 miles) in the ocean. At those depths, of course, sea water is well below critical temperature. The hot water circulating in vents above the active volcanic system at mid-ocean ridges at the bottom of the ocean is thus a supercritical fluid.
1.3 Density of Water
Water is unique among common substances in that its density decreases when it freezes; that is, water goes from 1.000 gram per cubic centimeter in liquid form to 0.915 grams per cubic centimeter in solid form. As a result, water expands about 9 percent in volume when it freezes. Consequently, ice floats on lakes and rivers, and icebergs float in the ocean. This also explains the common phrase “only the tip of the iceberg,” implying that there is considerably more present than can be seen. Because of the relative densities of ice and liquid water, approximately 90 percent of the mass of ice remains hidden below water level.
In water, hydrogen bonds produce clusters of water molecules with a more open (less dense) structure than water itself. Cluster formation reaches a minimum at about 4°C (39°F). Because of this, water at 4°C is denser than water at any other temperature and will sink to the bottom in a pool or lake. In lakes, as winter air temperatures fall below freezing, this phenomenon helps to keep the lake from freezing entirely, because when the water near the surface cools to 4°C, it sinks below the crust of surface ice, which is at 0°C. As a result, water remains unfrozen at the bottom of the lake. This is the reason people with fish ponds in the northern latitudes are amazed in the spring to find their goldfish still alive even though the pond surface froze completely during the winter.
1.4 Water at Exited State
When heat is added to water, its temperature increases. Specific heat is the measure of the amount of heat or energy required to raise the temperature of 1 gram of a material 1°C. The specific heat of water is 1 calorie per gram, or 4.18 joules per gram. The specific heat of water is larger than that of most other substances (about four times greater than rocks). Even more heat must be added in order to melt ice and vaporize water. The amount of heat required to melt ice is called the latent heat of fusion; its value is 80 calories per gram. The latent heat of vaporization, which is the amount of energy required to convert one gram of water to one gram of vapor at a constant temperature, is even larger—about 540 calories per gram. Thus, the amount of energy required to boil or evaporate water without changing its temperature is about five times the energy it takes to warm water from its freezing to melting point.
The amount of energy required to vaporize water and to melt ice is exactly matched by the energy released during the reverse of these processes. When 1 gram of water vapor condenses, it releases 540 calories; when 1 gram of water freezes, it releases 80 calories. As a result, large bodies of water such as oceans and lakes have a moderating effect on climate by storing (through evaporation) and releasing (through condensation) large amounts of energy as the temperature increases or decreases, respectively.
1.5 Solubility of Different Materials in Freshwater
The chemical and physical properties of water depend on the amount and composition of dissolved materials. For example, the melting point of water decreases if salt is added. For this reason, sea water can remain liquid at temperatures below 0°C (32°F), and salt is sometimes used on roads and sidewalks during the winter in order to prevent water on their surfaces from freezing. Fresh water is generally characterized by having low concentrations of dissolved salts and other totally dissolved solids. It excludes specifically seawater and brackish water although it does not include mineral-rich waters like the chalybeate springs. The term “sweet water” has been commonly used to describe fresh water compared to salt water.
1.6 Boiling of Freshwater
A liquid boils at a temperature at which its vapor pressure is equal to the pressure of the gas above it (Goldberg David.E 1988). The lower the pressure of a gas above a liquid, the lower the temperature at which the liquid will boil. Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point, the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding environmental pressure
1.6.1 Boiling point of Freshwater
Boiling point is the temperature at which the pressure exerted by the surroundings upon a liquid is equaled by the pressure exerted by the vapor of the liquid; under this condition, addition of heat results in the transformation of the liquid into its vapor without raising the temperature. At any temperature a liquid partly vaporizes into the space above it until the pressure exerted by the vapor reaches a characteristic value called the vapor pressure of the liquid at that temperature. As the temperature is increased, the vapor pressure increases; at the boiling point, bubbles of vapor form within the liquid and rise to the surface. The boiling point of a liquid varies according to the applied pressure; the normal boiling point is the temperature at which the vapor pressure is equal to the standard sea-level atmospheric pressure (760 mm [29.92 inches] of mercury). At sea level, water boils at 100° C (212° F). At higher altitudes the temperature of the boiling point is lower.
1.6.2 Factors that affect the boiling point of a substance
The factors that affect the boiling point of a substance include:
- Pressure: when the external pressure is:
- Less than one atmosphere, the boiling point of the liquid is lower than its normal boiling point.
- Equal to one atmosphere, the boiling point of a liquid is called the normal boiling point.
- Greater than one atmosphere, the boiling point of the liquid is greater than its normal boiling point.
Less than one atmosphere, the boiling point of the liquid is lower than its normal boiling point.
Equal to one atmosphere, the boiling point of a liquid is called the normal boiling point.
Greater than one atmosphere, the boiling point of the liquid is greater than its normal boiling point.
- Types of Molecules – the types of molecules that make up a liquid determine its boiling point. If the intermolecular forces between molecules are:
- Relatively strong, the boiling point will be relatively high.
- Relatively weak, the boiling point will be relatively low.
1.7 Freezing of Freshwater
When a substance is changed from its liquid state to a solid state when its temperature is lowered below its freezing point. Example is when a cup of water (liquid state) is placed in a refrigerator and then changes to ice block (solid state).
1.7.1 Freezing point of Freshwater
Freezing point is the temperature at which a liquid becomes a solid at normal atmospheric pressure. Alternatively, a melting point is the temperature at which a solid becomes a liquid at normal atmospheric pressure.
A more specific definition of freezing point is the temperature at which solid and liquid phases coexist in equilibrium. The normal freezing point is the temperature at a substance melts (or freezes) at one atmosphere (760 torr = 760 mm Hg = 14.7 psi = 101.3 kPa) of pressure.
As with the melting point, increased pressure usually raises the freezing point. The freezing point is lower than the melting point in the case of mixtures and for certain organic compounds such as fats. As a mixture freezes, the solid that forms first usually has a composition different from that of the liquid, and formation of the solid changes the composition of the remaining liquid, usually in a way that steadily lowers the freezing point. This principle is used in purifying mixtures, successive melting and freezing gradually separating the components. The heat of fusion, the heat that must be applied to melt a solid, must be removed from the liquid to freeze it. Some liquids can be super cooled i.e., cooled below the freezing point without solid crystals forming. Putting a seed crystal into a super cooled liquid triggers freezing, whereupon the release of the heat of fusion raises the temperature rapidly to the freezing point.
Super cooling: The process of lowering the temperature of a liquid below its freezing point without it becoming a solid.
Super heating: In physics, this is a phenomenon in which a liquid is heated to a temperature higher than its boiling point.
Heat of Fusion: This is the amount of heat (energy) required to change a gram of a substance from its solid to its liquid state, without a change in its temperature.
Heat of Vaporization: This is the amount of heat (energy) required to change a gram of liquid into its gaseous state at its boiling point.
1.8 Anomalous nature of water
The anomalous properties of water are those where the behavior of liquid water is quite different from what is found with other liquids .No other material is commonly found as solid, liquid and gas. Frozen water (ice) also shows anomalies when compared with other solids. (P. Ball 2008). Although it is an apparently simple molecule (H2O), it has a highly complex and anomalous character due to its intra-molecular hydrogen bonding. As a gas, water is one of lightest known, as a liquid it is much denser than expected and as a solid it is much lighter than expected when compared with its liquid form. It can be extremely slippery and extremely sticky at the same time (Connor Court, Ballarat, 2014). Many other anomalies of water may remain to be discovered, such as the possible link of water to room temperature superconductivity.
As liquid water is so common-place in our everyday lives, it is often regarded as a ‘typical’ liquid. In reality, water is most atypical as a liquid, behaving as a quite different material at low temperatures to that when it is hot. It has often been stated that life depends on these anomalous properties of water (Plenum: New York, 1985). In particular, the high cohesion between molecules gives it a high freezing and melting point, such that we and our planet are bathed in liquid water. The large heat capacity, high thermal conductivity and high water content in organisms contribute to thermal regulation and prevent local temperature fluctuations, thus allowing us to more easily control our body temperature. The high latent heat of evaporation gives resistance to dehydration and considerable evaporative cooling. Water ionizes and allows easy proton exchange between molecules, so contributing to the richness of the ionic interactions in biology. Also, it is an excellent solvent due to its polarity, high relative permittivity (dielectric constant) and small size, particularly for polar and ionic compounds and salts.
At 4 °C water expands on heating or cooling. This density maximum together with the low ice density results in
(i) The necessity that all of a body of fresh water (not just its surface) is close to 4 °C before any freezing can occur
(ii) The freezing of rivers, lakes and oceans is from the top down, so permitting survival of the bottom ecology, insulating the water from further freezing, reflecting back sunlight into space
1.9 Water’s High Heat Capacity
The capability for a molecule to absorb heat energy is called heat capacity. Water’s high heat capacity is a property caused by hydrogen bonding among water molecules. When heat is absorbed, hydrogen bonds are broken and water molecules can move freely. When the temperature of water decreases, the hydrogen bonds are formed and release a considerable amount of energy. Water has the highest specific heat capacity of any liquid. Specific heat is defined 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, or 4.184 Joules. Heat retention refers to the amount of heat an object or material can store overtime (Camilloni and Barros, 1997). As a result, it takes water a long time to heat and a long time to cool. In fact, the specific heat capacity of water is about five times more than that of sand. This explains why the land cools faster than the sea.
Heat capacity is an extensive property of matter, connoting that it is proportional to the size of the system (Camilloni and Barros, 1997).
Impurities are substances inside a confined amount of liquid, gas, or solid, which differ from the chemical composition of the material or compound. Impurities are either naturally occurring or added during synthesis of a chemical or commercial product. (Cheng, E et al 2004)
1.10.1 Effect of impurities on the boiling point of a liquid
Considering the problem in terms of entropy. The boiling point of a liquid is defined as the temperature at which the liquid and gas phases are at equilibrium. Mathematically, this means that the change in free energy from going to liquid to gas is zero . Adding a (non-volatile) impurity to the liquid phase increases the entropy of the liquid without affecting the entropy of the gas; the end result is that smaller for an impure liquid than a pure liquid. Adding an impurity to the liquid phase causes to decrease without changing . Therefore, adding an impurity to water will cause it to boil at a higher temperature. In essence, vaporization is a process that is driven by the increase in entropy associated with going from the liquid phase to the gas phase. By making the liquid phase more disordered, this gain in entropy becomes smaller, and vaporization becomes slightly less favorable.
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