{"id":170,"date":"2017-01-23T16:35:42","date_gmt":"2017-01-23T16:35:42","guid":{"rendered":"https:\/\/pressbooks.ccconline.org\/introduction-to-oceanography\/chapter\/5-1-properties-of-water\/"},"modified":"2021-10-25T21:39:38","modified_gmt":"2021-10-25T21:39:38","slug":"5-1-properties-of-water","status":"publish","type":"chapter","link":"https:\/\/pressbooks.ccconline.org\/introduction-to-oceanography\/chapter\/5-1-properties-of-water\/","title":{"raw":"5.1 Properties of Water","rendered":"5.1 Properties of Water"},"content":{"raw":"The most obvious feature of the oceans is that they contain water. Water is so ubiquitous that it may not seem like a very interesting substance, but it has many unique properties that impact global oceanographic and climatological processes. Many of these processes are due to [pb_glossary id=\"840\"]hydrogen bonds[\/pb_glossary]<\/strong> forming between water molecules.\r\n\r\n \r\n\r\n[caption id=\"attachment_164\" align=\"aligncenter\" width=\"300\"]\"Hydrogen<\/a> Figure 5.1.1<\/strong> Hydrogen bonds (dashed lines) between water molecules. Oxygen atoms are shown in red, hydrogen atoms in white (Public domain, via Wikimedia Commons).[\/caption]\r\n\r\nThe water molecule consists of two hydrogen atoms and one oxygen atom. The electrons responsible for the bonds between the atoms are not distributed equally throughout the molecule, so that the hydrogen ends of water molecules have a slight positive charge, and the oxygen end has a slight negative charge, making water a [pb_glossary id=\"1044\"]polar molecule[\/pb_glossary]<\/strong>. The negative oxygen side of the molecule forms an attraction to the positive hydrogen end of a neighboring molecule. This rather weak force of attraction is called a hydrogen bond (Figure 5.1.1). If not for hydrogen bonds, water would vaporize at -68o<\/sup> C, meaning liquid water (and thus life) could not exist on Earth. These hydrogen bonds are responsible for some of water\u2019s unique properties:\r\n\r\n1. Water is the only substance to naturally exist in a solid, liquid, and gaseous form under the normal range of temperatures and pressures found on Earth. This is due to water\u2019s relatively high freezing and vaporizing points (see below).\r\n\r\n2. Water has a high [pb_glossary id=\"828\"]heat capacity[\/pb_glossary]<\/strong>, which is the amount of heat that must be added to raise its temperature. [pb_glossary id=\"1164\"]Specific heat[\/pb_glossary] is the heat required to raise the temperature of 1 g of a substance by 1o<\/sup> C. Water has the highest specific heat of any liquid except ammonia (Table 5.1.1).\r\n\r\n[table id=3 \/]\r\n\r\nWater is therefore one of the most difficult liquids to heat or cool; it can absorb large amounts of heat without increasing its temperature. Remember that temperature reflects the average [pb_glossary id=\"1609\"]kinetic energy[\/pb_glossary] of the molecules within a substance; the more vigorous the motion, the higher the temperature. In water, the molecules are held together by hydrogen bonds, and these bonds must be overcome to allow the molecules to move freely. When heat is added to water the energy must first go to breaking the hydrogen bonds before the temperature can begin to rise. Therefore, much of the added heat is absorbed by breaking H bonds, not by increasing the temperature, giving water a high heat capacity.\r\n\r\nHydrogen bonds also give water a high latent heat; the heat required to undergo a [pb_glossary id=\"1022\"]phase change[\/pb_glossary] from solid to liquid, or liquid to gas. The [pb_glossary id=\"890\"]latent heat of fusion[\/pb_glossary]<\/strong> is the heat required to go from solid to liquid; 80 cal\/g in the case of ice melting to water. Ice is a solid because hydrogen bonds hold the water molecules into a solid crystal lattice (see below). As ice is heated, the temperature rises up to 0o<\/sup> C. At that point, any additional heat goes to melting the ice by breaking the hydrogen bonds, not to increasing the temperature. So as long as ice is present, the water temperature will not increase. This is why your drink will remain cold as long as it contains ice; any heat absorbed goes to melting the ice, not to warming the drink.\r\n\r\nWhen all of the ice is melted, additional heat will increase the temperature of the water 1o<\/sup> C for each calorie of heat added, until it reaches 100o<\/sup> C. At that point, any additional heat goes to overcoming the hydrogen bonds and turning the liquid water into water vapor, rather than increasing the water temperature. The heat required to evaporate liquid water into water vapor is the\u00a0[pb_glossary id=\"892\"]latent heat of vaporization[\/pb_glossary]<\/strong> which has a value of 540 cal\/g (Figure 5.1.2).\r\n\r\n[caption id=\"attachment_165\" align=\"aligncenter\" width=\"1000\"]\"Latent<\/a> Figure 5.1.2<\/strong> Latent heat required for phase changes in water. Latent heat of fusion is the heat required to melt ice (80 cal\/g), and latent heat of vaporization is the heat required to turn liquid water into water vapor (540 cal\/g) (PW).[\/caption]\r\n\r\nThe high [pb_glossary id=\"828\"]heat capacity[\/pb_glossary] of water helps regulate global climate, as the oceans slowly absorb and release heat, preventing rapid swings in temperature (see section 8.1<\/a>). It also means that aquatic organisms aren't as subjected to the same rapid temperature changes as terrestrial organisms. A deep ocean organism may not experience more than a 0.5o<\/sup> C change in temperature over its entire life, while a terrestrial species may encounter changes of more than 20o<\/sup> C in a single day!\r\n\r\n3. Water dissolves more substances than any other liquid; it is a \"universal solvent\"<\/strong>, which is why so many substances are dissolved in the ocean. Water is especially good at dissolving ionic salts; molecules made from oppositely charged [pb_glossary id=\"872\"]ions [\/pb_glossary] such as NaCl (Na+<\/sup> and Cl-<\/sup>). In water, the charged ions attract the [pb_glossary id=\"1044\"]polar [\/pb_glossary] water molecules. The ions get surrounded by a layer of water molecules, weakening the bond between the ions by up to 80 times. With the bonds weakened between ions, the substance dissolves (Figure 5.1.3).\r\n\r\n[caption id=\"attachment_166\" align=\"aligncenter\" width=\"600\"]\"Attraction<\/a> Figure 5.1.3<\/strong> Attraction between polar water molecules and charged ions (such as in NaCl) is greater than the attraction between the charged ions, causing the ions to dissociate and the salt to dissolve (PW).[\/caption]\r\n\r\n4. The solid phase is less dense than the liquid phase<\/strong>. In other words, ice floats. Most substances are denser in the solid form than in the liquid form, as their molecules are more closely packed together as a solid. Water is an exception: the density of fresh water is 1.0 g\/cm3<\/sup>, while the density of ice is 0.92 g\/cm3<\/sup>, and once again, this is due to the action of hydrogen bonds.\r\n\r\nAs water temperature cools the molecules slow down, eventually slowing enough that hydrogen bonds can form and hold the water molecules in a crystal lattice. The molecules in the lattice are spaced farther apart than the molecules in liquid water, which makes ice less dense than liquid water (Figure 5.1.4). This is familiar to anyone who has ever left a full water bottle in the freezer, only to have it burst as the water freezes and expands.\r\n\r\n \r\n\r\n[caption id=\"attachment_167\" align=\"aligncenter\" width=\"500\"]\"Crystal<\/a> Figure 5.1.4<\/strong> Crystal lattice structure of ice, showing water molecules held together by hydrogen bonds (By Adam001d (Own work) [CC BY-SA 3.0], via Wikimedia Commons).[\/caption]But the relationship between temperature and water density is not a simple linear one. As water cools, its density increases as expected, as the water molecules slow down and get closer together. However, fresh water reaches its maximum density at a temperature of 4o<\/sup> C, and as it cools beyond that point its density declines as the hydrogen bonds begin to form and the intermolecular spacing increases (Figure 5.1.5 inset). The density continues to decline until the temperature reaches 0o<\/sup> C and ice crystals form, reducing the density dramatically (Figure 5.1.5).\r\n\r\n \r\n\r\n[caption id=\"attachment_168\" align=\"aligncenter\" width=\"573\"]\"Graph<\/a> Figure 5.1.5<\/strong> As temperature declines, the density of water increases until it reaches maximum density at 4o<\/sup> C (inset). Density then declines slightly down to 0o<\/sup> C, where it declines dramatically as ice crystals form (Klaus-Dieter Keller, [CC BY-SA 3.0], via Wikimedia Commons).[\/caption]There are a number of important implications to ice being less dense than water. Ice floating on the surface of the ocean helps regulate ocean temperatures, and therefore global climate, by influencing the amount of sunlight that is reflected rather than absorbed (see section 8.1<\/a>). On a smaller scale, surface ice can prevent lakes and ponds from freezing solid during the winter. As fresh surface water cools, the water gets denser, and sinks to the bottom. The new surface water then cools and sinks, and the process is repeated in what is referred to as [pb_glossary id=\"1002\"]overturning[\/pb_glossary]<\/strong>, with denser water sinking and less dense water moving to the surface only to be cooled and sink itself. In this way, the entire body of water is cooled somewhat evenly. This process continues until the surface water cools below 4o<\/sup> C. Below 4o<\/sup> C, the water becomes less<\/em> dense as it cools, so it no longer sinks. Instead, it remains as the surface, getting colder and less dense, until it freezes at 0o<\/sup> C. Once fresh water freezes, the ice floats and insulates the rest of the water beneath it, reducing further cooling. The densest bottom water is still at 4o<\/sup> C, so it does not freeze, allowing the bottom of a lake or pond to remain unfrozen (which is good news for the animals living there) no matter how cold it gets outside.\r\n\r\nThe dissolved salts in seawater inhibit the formation of the crystal lattice, and therefore make it harder for ice to form. So seawater has a freezing point of about -2o<\/sup> C (depending on salinity), and freezes before a temperature of maximum density is reached. Thus seawater will continue to sink as it gets colder, until it finally freezes.\r\n\r\n5. Water has a very high [pb_glossary id=\"1192\"]surface tension[\/pb_glossary]<\/strong>, the highest of any liquid except mercury (Table 5.1.2). Water molecules are attracted to each other by hydrogen bonds. For molecules not at the water surface, they are surrounded by other water molecules in all directions, so the attractive forces are evenly distributed in all directions. But for molecules at the surface there are few adjacent molecules above them, only below, so all of the attractive forces are directed inwards, away from the surface (Figure 5.1.6). This inwards force is what causes water droplets to take on a spherical shape, and water to bead up on a surface, as the spherical shape provides the minimum possible surface area. These attractive forces also cause the surface of the water to act like an elastic \"skin\" which allows things like insects to sit on the water's surface without sinking.\r\n\r\n[caption id=\"attachment_169\" align=\"aligncenter\" width=\"600\"]\"The<\/a> Figure 5.1.6<\/strong> The net attractive force between molecules at the surface is inwards, leading to surface tension. For molecules in the center, the force is equal in all directions (PW).[\/caption]\r\n\r\n[table id=8 \/]","rendered":"

The most obvious feature of the oceans is that they contain water. Water is so ubiquitous that it may not seem like a very interesting substance, but it has many unique properties that impact global oceanographic and climatological processes. Many of these processes are due to hydrogen bonds<\/a><\/strong> forming between water molecules.<\/p>\n

 <\/p>\n

\"Hydrogen<\/a>
Figure 5.1.1<\/strong> Hydrogen bonds (dashed lines) between water molecules. Oxygen atoms are shown in red, hydrogen atoms in white (Public domain, via Wikimedia Commons).<\/figcaption><\/figure>\n

The water molecule consists of two hydrogen atoms and one oxygen atom. The electrons responsible for the bonds between the atoms are not distributed equally throughout the molecule, so that the hydrogen ends of water molecules have a slight positive charge, and the oxygen end has a slight negative charge, making water a polar molecule<\/a><\/strong>. The negative oxygen side of the molecule forms an attraction to the positive hydrogen end of a neighboring molecule. This rather weak force of attraction is called a hydrogen bond (Figure 5.1.1). If not for hydrogen bonds, water would vaporize at -68o<\/sup> C, meaning liquid water (and thus life) could not exist on Earth. These hydrogen bonds are responsible for some of water\u2019s unique properties:<\/p>\n

1. Water is the only substance to naturally exist in a solid, liquid, and gaseous form under the normal range of temperatures and pressures found on Earth. This is due to water\u2019s relatively high freezing and vaporizing points (see below).<\/p>\n

2. Water has a high heat capacity<\/a><\/strong>, which is the amount of heat that must be added to raise its temperature. Specific heat<\/a> is the heat required to raise the temperature of 1 g of a substance by 1o<\/sup> C. Water has the highest specific heat of any liquid except ammonia (Table 5.1.1).<\/p>\n

[table id=3 \/]<\/p>\n

Water is therefore one of the most difficult liquids to heat or cool; it can absorb large amounts of heat without increasing its temperature. Remember that temperature reflects the average kinetic energy of the molecules within a substance; the more vigorous the motion, the higher the temperature. In water, the molecules are held together by hydrogen bonds, and these bonds must be overcome to allow the molecules to move freely. When heat is added to water the energy must first go to breaking the hydrogen bonds before the temperature can begin to rise. Therefore, much of the added heat is absorbed by breaking H bonds, not by increasing the temperature, giving water a high heat capacity.<\/p>\n

Hydrogen bonds also give water a high latent heat; the heat required to undergo a phase change<\/a> from solid to liquid, or liquid to gas. The latent heat of fusion<\/a><\/strong> is the heat required to go from solid to liquid; 80 cal\/g in the case of ice melting to water. Ice is a solid because hydrogen bonds hold the water molecules into a solid crystal lattice (see below). As ice is heated, the temperature rises up to 0o<\/sup> C. At that point, any additional heat goes to melting the ice by breaking the hydrogen bonds, not to increasing the temperature. So as long as ice is present, the water temperature will not increase. This is why your drink will remain cold as long as it contains ice; any heat absorbed goes to melting the ice, not to warming the drink.<\/p>\n

When all of the ice is melted, additional heat will increase the temperature of the water 1o<\/sup> C for each calorie of heat added, until it reaches 100o<\/sup> C. At that point, any additional heat goes to overcoming the hydrogen bonds and turning the liquid water into water vapor, rather than increasing the water temperature. The heat required to evaporate liquid water into water vapor is the\u00a0latent heat of vaporization<\/a><\/strong> which has a value of 540 cal\/g (Figure 5.1.2).<\/p>\n

\"Latent<\/a>
Figure 5.1.2<\/strong> Latent heat required for phase changes in water. Latent heat of fusion is the heat required to melt ice (80 cal\/g), and latent heat of vaporization is the heat required to turn liquid water into water vapor (540 cal\/g) (PW).<\/figcaption><\/figure>\n

The high heat capacity<\/a> of water helps regulate global climate, as the oceans slowly absorb and release heat, preventing rapid swings in temperature (see section 8.1<\/a>). It also means that aquatic organisms aren’t as subjected to the same rapid temperature changes as terrestrial organisms. A deep ocean organism may not experience more than a 0.5o<\/sup> C change in temperature over its entire life, while a terrestrial species may encounter changes of more than 20o<\/sup> C in a single day!<\/p>\n

3. Water dissolves more substances than any other liquid; it is a “universal solvent”<\/strong>, which is why so many substances are dissolved in the ocean. Water is especially good at dissolving ionic salts; molecules made from oppositely charged ions <\/a> such as NaCl (Na+<\/sup> and Cl–<\/sup>). In water, the charged ions attract the polar <\/a> water molecules. The ions get surrounded by a layer of water molecules, weakening the bond between the ions by up to 80 times. With the bonds weakened between ions, the substance dissolves (Figure 5.1.3).<\/p>\n

\"Attraction<\/a>
Figure 5.1.3<\/strong> Attraction between polar water molecules and charged ions (such as in NaCl) is greater than the attraction between the charged ions, causing the ions to dissociate and the salt to dissolve (PW).<\/figcaption><\/figure>\n

4. The solid phase is less dense than the liquid phase<\/strong>. In other words, ice floats. Most substances are denser in the solid form than in the liquid form, as their molecules are more closely packed together as a solid. Water is an exception: the density of fresh water is 1.0 g\/cm3<\/sup>, while the density of ice is 0.92 g\/cm3<\/sup>, and once again, this is due to the action of hydrogen bonds.<\/p>\n

As water temperature cools the molecules slow down, eventually slowing enough that hydrogen bonds can form and hold the water molecules in a crystal lattice. The molecules in the lattice are spaced farther apart than the molecules in liquid water, which makes ice less dense than liquid water (Figure 5.1.4). This is familiar to anyone who has ever left a full water bottle in the freezer, only to have it burst as the water freezes and expands.<\/p>\n

 <\/p>\n

\"Crystal<\/a>
Figure 5.1.4<\/strong> Crystal lattice structure of ice, showing water molecules held together by hydrogen bonds (By Adam001d (Own work) [CC BY-SA 3.0], via Wikimedia Commons).<\/figcaption><\/figure>\n

But the relationship between temperature and water density is not a simple linear one. As water cools, its density increases as expected, as the water molecules slow down and get closer together. However, fresh water reaches its maximum density at a temperature of 4o<\/sup> C, and as it cools beyond that point its density declines as the hydrogen bonds begin to form and the intermolecular spacing increases (Figure 5.1.5 inset). The density continues to decline until the temperature reaches 0o<\/sup> C and ice crystals form, reducing the density dramatically (Figure 5.1.5).<\/p>\n

 <\/p>\n

\"Graph<\/a>
Figure 5.1.5<\/strong> As temperature declines, the density of water increases until it reaches maximum density at 4o<\/sup> C (inset). Density then declines slightly down to 0o<\/sup> C, where it declines dramatically as ice crystals form (Klaus-Dieter Keller, [CC BY-SA 3.0], via Wikimedia Commons).<\/figcaption><\/figure>\n

There are a number of important implications to ice being less dense than water. Ice floating on the surface of the ocean helps regulate ocean temperatures, and therefore global climate, by influencing the amount of sunlight that is reflected rather than absorbed (see section 8.1<\/a>). On a smaller scale, surface ice can prevent lakes and ponds from freezing solid during the winter. As fresh surface water cools, the water gets denser, and sinks to the bottom. The new surface water then cools and sinks, and the process is repeated in what is referred to as overturning<\/a><\/strong>, with denser water sinking and less dense water moving to the surface only to be cooled and sink itself. In this way, the entire body of water is cooled somewhat evenly. This process continues until the surface water cools below 4o<\/sup> C. Below 4o<\/sup> C, the water becomes less<\/em> dense as it cools, so it no longer sinks. Instead, it remains as the surface, getting colder and less dense, until it freezes at 0o<\/sup> C. Once fresh water freezes, the ice floats and insulates the rest of the water beneath it, reducing further cooling. The densest bottom water is still at 4o<\/sup> C, so it does not freeze, allowing the bottom of a lake or pond to remain unfrozen (which is good news for the animals living there) no matter how cold it gets outside.<\/p>\n

The dissolved salts in seawater inhibit the formation of the crystal lattice, and therefore make it harder for ice to form. So seawater has a freezing point of about -2o<\/sup> C (depending on salinity), and freezes before a temperature of maximum density is reached. Thus seawater will continue to sink as it gets colder, until it finally freezes.<\/p>\n

5. Water has a very high surface tension<\/a><\/strong>, the highest of any liquid except mercury (Table 5.1.2). Water molecules are attracted to each other by hydrogen bonds. For molecules not at the water surface, they are surrounded by other water molecules in all directions, so the attractive forces are evenly distributed in all directions. But for molecules at the surface there are few adjacent molecules above them, only below, so all of the attractive forces are directed inwards, away from the surface (Figure 5.1.6). This inwards force is what causes water droplets to take on a spherical shape, and water to bead up on a surface, as the spherical shape provides the minimum possible surface area. These attractive forces also cause the surface of the water to act like an elastic “skin” which allows things like insects to sit on the water’s surface without sinking.<\/p>\n

\"The<\/a>
Figure 5.1.6<\/strong> The net attractive force between molecules at the surface is inwards, leading to surface tension. For molecules in the center, the force is equal in all directions (PW).<\/figcaption><\/figure>\n

[table id=8 \/]<\/p>\n

definition<\/span>