Module 1: Intermolecular Forces PDF

MODULE FOR GENERAL CHEMISTRY 2 ------------------------------------------------------- - MODULE 1: INTERMOLECULAR FORCES Most Essential Learning Competencies: 1. Use the kinetic molecular model to explain the properties of liquids and solids...

MODULE FOR GENERAL CHEMISTRY 2 ------------------------------------------------------- - MODULE 1: INTERMOLECULAR FORCES Most Essential Learning Competencies: 1. Use the kinetic molecular model to explain the properties of liquids and solids (STEM_GC11IMFIIIa-c-99) 2. Describe and differentiate the types of intermolecular forces (STEM_GC11IMFIIIa-c-100) 3. Describe the following properties of liquids, and explain the effect of intermolecular forces on these properties: surface tension, viscosity, vapor pressure, boiling point, and molar heat of vaporization (STEM_GC11IMFIIIa-c-102) 4. Explain the properties of water with its molecular structure and intermolecular forces (STEM_GC11IMFIIIa-c-103) LESSON 1: The Kinetic Molecular Model and the Properties of Liquids and Solids Matter can exist in three main different states: namely, solid, liquid, and gas. The most common example of which is water. You only have to think about water to appreciate how different the three states of matter are. Steam bathing, drinking, and ice skating are all done in contact with water in its various forms. But how do these states of matter differ from each other? Understanding the kinetic molecular model of the three states will answer this question. Activity 1: What's the Matter? Directions: Based on the picture below, identify what state of matter is being represented. Write a brief description based on the arrangement of their particles and give three examples for each state. (1) (2) (3) State (1) Arrangement of Particles: __________________________________________________________________________ Examples: ______________________________________________________________________________________ State (2) Arrangement of Particles: __________________________________________________________________________ Examples: ______________________________________________________________________________________ State (3) Arrangement of Particles: __________________________________________________________________________ Examples: ______________________________________________________________________________________ What is the Kinetic Molecular Theory? The kinetic-molecular theory is based on the idea that matter is composed of tiny particles that are always in motion. The theory helps explain the observable properties and behaviors of solids, liquids, and gases. It helps to explain why matter exists in different phases (solid, liquid, and gas) and how matter can change from one phase to another. 1  The state of a substance depends on the balance between the kinetic energy of the individual particles (molecules or atoms) and the intermolecular forces. The kinetic energy keeps the molecules apart and moving around, and is a function of the temperature of the substance. The intermolecular forces are attractive forces that try to draw the particles together. Postulates of Kinetic Molecular Theory The Kinetic Molecular Theory (KMT) is based on a series of postulates. Some of the postulates of KMT are as follows: 1. Matter is made of particles that are constantly in motion. This energy in motion is called kinetic energy. 2. The amount of kinetic energy in a substance is related to its temperature. Increased temperature means greater speed. 3. There is space between particles. The amount of space between particles is related to the substance's state of matter. 4. Phase changes happen when the temperature of the substance changes sufficiently. 5. There are attractive forces in between particles called intermolecular forces. The strength of these forces increases as particles get closer together. KMT in Relation to Liquids and Solids The principal difference between the condensed states (liquids and solids) and the gaseous state is the distance between molecules. In a liquid, the molecules are so close together that there is very little empty space between particles. Thus, liquids are much more difficult to compress than gases, and they are also much denser under normal conditions. Molecules in a liquid are held together by one or more types of attractive forces. A liquid also has a definite volume, because molecules in a liquid do not break away from the attractive forces. The molecules can, however, move past one another freely. So, a liquid can flow, can be poured, and assumes the shape of its container. In a solid, molecules are held rigidly in position with virtually no freedom of movement, so they only vibrate only about fixed positions. There is even lesser empty space between particles in a solid than in a liquid because their particles are tightly packed. Thus, solids are almost incompressible and possess definite shape and volume. This is due to the stronger intermolecular force of attraction compared to liquids. Figure 1 shows the orientation of particles in each state of matter and the direction of phase change due to the addition and removal of kinetic energy. Figure 1 Activity 2: Be a Science Detective! Directions: Investigate and analyze the given situation. Explain the phenomenon. Based on the kinetic molecular model, solids usually have higher densities compared to liquids. However, ice, a solid form of H 2O floats on liquid water which means ice is less dense than water. What do you think is the reason for this observation? What is the biological significance of this concept? Activity 3: Describe Me Directions: Compare the properties of solids and liquids by completing the table based on the kinetic molecular model. Provide a short description of each characteristic for the given state of matter. CHARACTERISTICS SOLIDS LIQUIDS Intermolecular force Shape Volume Density Compressibility Arrangement of particles Motion of molecules Fluidity --------------------------------------------------------- End of Module 1 – Lesson 1 ------------------------------------------------------------ 2  LESSON 2: Types of Intermolecular Forces In the preceding lesson, we have noted the differences in the properties of matter in the gas phase from those in the liquid and solid phases. Such difference can be attributed to the strong attractive forces in solid and liquid molecules. Gas molecules have negligible or no attractions at all. The condensation of gaseous substance to form liquids which in turn form solids could be explained by the attractive forces called intermolecular forces. Intermolecular forces vs. Intramolecular forces It is important to note the difference between intermolecular forces and intramolecular forces. As discussed in General Chemistry 1, atoms can form stable units called molecules by sharing or transfer of electrons. This is called intramolecular bonding. Intramolecular (within molecules) forces hold atoms together in a molecule. Intramolecular forces stabilize individual molecules. Generally, these forces are simply chemical bonds such as ionic and covalent bonding. On the other hand, intermolecular forces are attractive forces between molecules. Intermolecular forces are responsible for the non-ideal behavior of gases, but they exert more influence in the condensed phases of matter - liquids and solids. Types of Intermolecular Forces The intermolecular forces of attraction in substances include dipole-dipole, London dispersion forces, hydrogen bonding, and ion-dipole forces. 1. London Dispersion Forces London dispersion forces, or simply dispersion forces, are intermolecular forces of attraction between all atoms and molecules. In addition, dispersion forces are the only kind of intermolecular forces present among symmetrical nonpolar substances such as O2 and CO2 and monoatomic species such as noble gases. Without dispersion forces, such substances could not condense to form liquids or solidify to form solids. Dispersion forces are weak attractive forces that result from the continuous movement of electrons in particles. Nonpolar molecules have zero dipole moment because their electron density is uniform and symmetrical. Nevertheless, the electrons have some freedom to move around the molecule. This induces temporary dipoles (instantaneous dipoles) in neighboring atoms or molecules. As electron clouds become larger and more diffuse, they are attracted less strongly by their own positive nuclei. Thus, they are more easily distorted or polarized by the adjacent/nearby nuclei. Polarization increases with increasing numbers of electrons and therefore with increasing size of molecules. Therefore, dispersion forces are generally stronger for molecules that are larger or have more electrons. For example, between Helium and Argon, two Argon atoms will have greater dispersion force because they are bigger than Helium atoms. 2. Dipole-dipole Forces Dipole-dipole forces are attractive forces between polar molecules, that is, between molecules that possess dipole moments. Their origin is electrostatic, and they can be understood in terms of Coulomb's law. The larger the dipole moment, the greater the force. Dipole-dipole forces are the attraction between the positive end of one molecule and the negative end of another. Dipoles form when there is a large difference in electronegativity between two atoms joined by a covalent bond. 3. Hydrogen Bonding The hydrogen bond is a special case of very strong dipole-dipole interaction. It is not a chemical bond in a formal sense. Strong Hydrogen bonding occurs among polar covalent molecules containing H and one of the three small, highly electronegative elements – F, O, or N. Like ordinary dipole-dipole interactions, Hydrogen bonds result from the attraction between + (partial positive) atoms of one molecule, in this case H atoms and the − (partial negative) atoms of another molecule. The + H 3 is attracted to a lone pair of electrons on an F, O, or N atom. Typically, a Hydrogen bond is about five to ten times stronger than other dipole-dipole interactions. 4. Ion-Dipole Forces Ion-dipole force acts between an ion (either cation or anion) and a polar molecule. When an ionic compound is placed in an aqueous solution, the positive end of the ionic compound becomes surrounded by the partial negative end of the ionic compound. In turn, it becomes surrounded by the partial positive Hydrogen ion in water. In short, the positive pole is attracted to the negative ion (anion), while the negative pole is attracted to a positive ion (cation). Ion-dipole interactions are involved in the dissolution process, like in the case of sodium chloride (table salt) dissolving in water. The Na+ and Cl- ions are dispersed among water molecules. The Na+ ions will be surrounded by the partial negative Oxygen of the water molecule, while the Cl- ions will be surrounded by the partial positive H of the water molecule. The strength of this interaction depends on the charge and size of the ion and on the magnitude of the dipole moment and size of the molecule. The charges on cations are generally more concentrated because cations are usually smaller than anions. Therefore, a cation interacts more strongly with dipoles than does an anion having a charge of the same magnitude. These four intermolecular forces vary in strength. Ion dipole forces are the strongest of the four, followed by Hydrogen bonding being a special type of dipole-dipole. Dipole-dipole is weaker than the ion-dipole and Hydrogen bonding, while London dispersion forces are the weakest. Activity 4: Operation Crossword Puzzle Directions: Complete the crossword by filling in the boxes to form a word that fits each clue. Disregard space between two-word phrases or hyphens. Across: 1. This is a special case of a very strong dipole-dipole interaction. 4. The force that holds atoms together in a molecule. 6. Intermolecular forces present among symmetrical nonpolar substances. 8. The attractive force between molecules. Down: 2. This is an intermolecular force that acts between an ion and a polar molecule. 3. A collective term used to define the attraction of intermolecular forces. 5. These are attractive forces between polar molecules. 7. The atom of this element forms hydrogen bonding. Activity 5: What If? Directions: Investigate and analyze the given situation. Provide a detailed explanation for each case. Water is present in almost every living thing, including human beings. It was discussed that water molecules are held by Hydrogen bonds. What do you think will happen if Hydrogen bonding in water does not exist at all? -------------------------------------------------------- End of Module 1 – Lesson 2 ------------------------------------------------------------- 4  LESSON 3: Properties of Liquids Liquids are made up of particles that are close to each other and have kinetic energy. The particles are not confined to a rigid position, and they move, but they can only travel at a short distance before they collide with each other and change the direction of motion. They roll and slide on top of one another and flow. Since the molecules flow, they take the shape of their container and diffuse moderately to a fixed volume. Liquids have moderately high density since they occupy a fixed volume, and the particles are attracted to each other. They also have low compressibility and thermal expansion. The kinetic energy of the molecules break away from their neighbor, and thus, the particles are joined by intermolecular forces. Most liquids exist as molecules at room temperature. The presence of the intermolecular forces results in special properties. The physical properties of liquids depend on the type of the different intermolecular forces. Surface Tension Surface tension is the force that causes the surface of a liquid to contract. It is the property of the surface of a liquid that allows it to resist an external force due to the cohesive nature of its molecules. Phenomena such as insects walking on the surface of the water, droplets of liquid being spherical in shape, and needles remaining suspended on the surface of the water can all be explained in terms of surface tension. The strength of surface tension depends on the intermolecular force of attraction. If the intermolecular force of attraction of a liquid is strong, then there is a greater force needed to break through the surface and the greater the surface tension is. Since the intermolecular forces vary in nature and strength, surface tension is different for various forms of liquids. Water has a high surface tension because of its ability to form a Hydrogen bond. Temperature affects surface tension. An increase in the liquid's temperature causes water molecules at the surface to evaporate, resulting in the weakening of the force of attraction. Therefore, there is less force needed to break through the surface of the molecules, and this decreases surface tension. Molecules within a liquid are pulled in all directions by intermolecular forces. Molecules at the surface are pulled downward and sideways by other molecules, not upward away from the surface. These intermolecular forces tend to pull the molecules into the liquid and cause the surface to tighten like an elastic film or "skin". Capillary action is the tendency of a liquid to rise in narrow tubes or be drawn into small openings such as those between grains of a rock. Capillary action, also known as capillarity, is a result of the intermolecular attraction between the liquid and solid materials. Capillary action is shown by water rising spontaneously in capillary tubes. A thin film of water adheres to the wall of the glass tube as water molecules are attracted to atoms making up the glass (SiO 2). Surface tension causes the film of water to contract and pulls the water up the tube. 2 Types of forces are involved in capillary action: 1. Cohesion is the intermolecular attraction between like molecules (the liquid molecules). 2. Adhesion is an attraction between unlike molecules (such as those in water and in the particles that make up the glass tube). These forces also define the shape of the surface of a liquid in a cylindrical container (the meniscus!) Viscosity Viscosity is the resistance of fluids to flow. A liquid's resistance (friction) to flow exists between the molecules of liquid when they move past each other. The greater the resistance in flowing, the more viscous the liquid is. Maple syrup in pancakes is usually made from the xylem sap of sugar maple, red maple, or black maple trees. It is boiled down, so it becomes a more concentrated and viscous liquid. Maple syrup is more viscous than water. The difference in viscosity between the two liquids is a measure of their intermolecular force of attraction. In order to flow, molecules must move, roll and slide over one another. A liquid with low intermolecular force allows its molecules to move freely and has a lower viscosity. An increase in temperature causes kinetic energy to increase. Heat breaks the intermolecular forces causing the liquid molecules to move faster. This makes the molecules flow more readily. Therefore, an increase in temperature decreases viscosity. 5  Since the structure of maple syrup contains a lot of O-H bond compared to water, more H-bonds are formed in maple syrup. The greater the number of H-bonds, the stronger the intermolecular force of attraction is, and the higher the viscosity of the liquid. Viscosity is expressed in units of centipoise. The table gives the viscosities of liquids of some pure substances. Water has 1 centipoise or 0.001 Pa/s at 20 °C. Substances with lower viscosities include carbon tetrachloride and benzene. Glycerol has a resistance to the flow of more than a thousand times greater than water. Liquids that have strong intermolecular forces have higher viscosities than those that have weak intermolecular forces. Viscosity decreases as temperature increases: hot molasses flows much faster than cold molasses. Vapor Pressure Vaporization is a phase change from liquid to gas, while the opposite process (gas to liquid) is condensation. When liquid molecules break free from their neighbors and escape into the gas phase, the process is called evaporation. Vaporization is a broader term that includes evaporation and boiling. Gas and vapor are similar but not the same. Vapor is used to refer to the gaseous phase of a substance, which is normally a liquid or solid at room temperature. The average kinetic energy of the liquid molecules of a substance depends on temperature. Most liquid particles have higher kinetic energy, and some others move at a slower pace. Substances that evaporate readily are volatile. They have weak intermolecular forces of attraction. Some examples of volatile liquids are alcohol, gasoline, paint thinner, and dry-cleaning solvents. Volatile substances burn more readily since they easily combine with Oxygen. Since the kinetic energy of a molecule is proportional to its temperature, evaporation proceeds more quickly at higher temperatures. As the faster-moving molecules escape, the remaining molecules have lower average kinetic energy, and the temperature of the liquid decreases. Therefore, evaporation is accompanied by cooling. In a closed container half-filled with liquid, the fast-moving molecules also escape into the gas phase forming vapor at the space above the liquid. Gas molecules move in random directions, collide with other gas particles and the walls of the container. Some will strike the liquid surface and condense back into it. In the closed flask, none of the gas particles are able to get out of the container. Eventually, the number of molecules that goes into the gaseous state would equal the number of molecules that condenses back. When the rate of condensation of the gas becomes equal to the rate of evaporation of the liquid, the gas in the container is said to be in equilibrium with the liquid. Like any gas sample, the molecules in the gaseous state over its liquid create a pressure. The greater the number of gaseous particles, the greater the pressure exerted by the gas. The pressure exerted by the gas in equilibrium with a liquid in a closed container at a given temperature is called the equilibrium vapor pressure or simply vapor pressure of the liquid. The equilibrium vapor pressure is the maximum vapor pressure of a liquid at a given temperature and that it is constant at a constant temperature. It increases with temperature. Vapor pressure is independent of the amount of liquid as well as the surface area of the liquid in contact with the gas. When the temperature is high, more molecules have enough energy to escape from the liquid. At a lower temperature, fewer molecules have sufficient energy to escape from the liquid. When liquids evaporate, the molecules have to have sufficient energy to break the attractive forces that hold them in the liquid state. The stronger these intermolecular forces are, the greater the amount of energy needed to break them. For some substances with weak intermolecular forces, the energy requirement is easily obtained from collisions with other molecules and absorption of energy from the surroundings. Many molecules can vaporize, resulting in high vapor pressure. For molecules with strong intermolecular forces, gathering enough energy may not be as easy and register low vapor pressures. The stronger the intermolecular forces of attraction, the lower the vapor pressure of a liquid. 6 Molar Heat of Vaporization The relationship between vapor pressure and strength of intermolecular forces is consistent with the trends in two other properties of liquids, the enthalpy or molar heat of vaporization and the boiling point of the liquid. The molar heat of vaporization (ΔHvap) is the energy required to vaporize 1 mole of a liquid at a given temperature. H is the symbol for enthalpy, which means heat content at a given standard condition. The heat of vaporization may be considered a measure of the strength of intermolecular forces in a liquid. If the intermolecular attraction is strong, it takes a lot of energy to free the molecules from the liquid phase, and the heat of vaporization will be high. It is easier to vaporize acetone (lower Hvap) than water (higher Hvap) at a given temperature, and more acetone escapes into the vapor phase at a given temperature. Acetone is a polar substance but has no H-bonding. It has weaker intermolecular forces than water, and therefore acetone molecules are held less tightly to one another in the liquid phase. A practical way to demonstrate differences in the molar heat of vaporization is by rubbing acetone on your hands. Compare what you feel when water is used. Acetone has a lower ΔHvap than water, so that heat from our hands is enough to increase the kinetic energy of these molecules and provide additional heat to vaporize them. As a result of the loss of heat from the skin, our hands feel cool. Boiling Point The boiling point of a liquid is the temperature at which the liquid changes into a gas. A liquid boils when its vapor pressure equals the pressure acting on the surface of the liquid. The boiling point is the temperature at which the vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which the liquid converts to a gas when the external pressure is 1 atm. The normal boiling point of water is 100 °C. The boiling point of a liquid depends on the external pressure. For example, at 1 atm, water boils at 100 °C, but if the pressure is reduced to 0.5 atm, water boils at only 82 °C. The boiling point is related to the molar heat of vaporization; the higher ΔHvap, the higher the boiling point. Activity 6: I Can Do It! Directions: Identify the concept that is described in each statement below. Choose the correct answer from the list below. Surface tension Boiling point Molar heat of vaporization Viscosity Vapor Vaporization Vapor pressure Liquid flow Capillary action Fluid 1. The measure of the elastic force on the surface of a liquid. 2. A gas or a liquid; a substance that can flow. 3. The tendency of a liquid to rise in narrow tubes or to be drawn into small openings. 4. The measure of a fluid’s resistance to flow. 5. A gaseous substance that exists naturally as a liquid or solid at normal temperature. 6. The change of phase from liquid to vapor (gaseous phase). 7. The equilibrium pressure of a vapor above its liquid; that is, the pressure exerted by the vapor above the surface of the liquid in a closed container. 8. The temperature at which a liquid boils. 9. The energy (usually in kilojoules) required to vaporize 1 mole of a liquid at a given temperature. 10. The movement of liquids and gases; describes how fluids behave and how they interact with their surrounding environment. Activity 7: Picture Shows What I Know Directions: Describe what is happening to the water molecules in the 2 flasks shown in the picture. Questions: 1. What happens to the molecules of water in the container when the temperature increases? 2. Container B shows equilibrium. How does it happen in such a balance? ------------------------------------------------------------- End of Module 1 – Lesson 3 -------------------------------------------------------- 7 LESSON 4: Properties of Water Water makes up a large proportion of the entire biosphere, where 95% is saltwater, and the remaining 5% is freshwater. Water is locked up in ice and glaciers, deep and shallow underground lakes, soil, atmosphere, and rivers. The human body consists of 50-75% water. Water serves important purposes for life on earth. Water's unique properties result from the strong intermolecular force of attraction characterized by the hydrogen bond. Some substances, like common table salt, NaCl, dissolve in water very easily. When placed in water, sodium chloride molecules fall apart. The positively charged sodium ion (Na+) binds to Oxygen, while the negatively charged chloride ion (Cl-) attaches to hydrogen. This property of water allows for the transport of nutrients vital to life in animals and plants. A drop of rainwater falling through the air dissolves atmospheric gases. When rain reaches the earth, it affects the quality of the land, lakes, and rivers. The following are properties of water: Boiling point and freezing point. The high boiling point of water is a consequence of its strong intermolecular forces of attraction caused by the formation of the H-bond. It also explains why water is liquid at room temperature. Due to Hydrogen bonding, water molecules cling to each other (cohesion) and remain in the liquid state under temperatures favorable to plants and other living organisms. Pure water at sea level boils at 100 0C and freezes at 0 0C, but extra energy is needed to push water molecules into the air. This is called latent heat—the heat required to change water from one phase to another. At higher elevations (lower atmospheric pressure), water's boiling temperature decreases. This is why it takes longer to boil an egg at higher altitudes. The temperature does not get high enough to cook the egg properly. If a substance is dissolved in water, the freezing point is lowered. Energy is lost when water freezes. A great deal of heat is released into the environment when liquid water changes to ice. It is lost when the high energy phase of liquid water moves to the low energy phase of ice. Thus, nights when ice freezes often feel warmer than nights when the ice melts. Specific heat. Specific heat refers to the amount of heat needed to change the temperature of 1 gram of a substance by 1 oC. For water, its specific heat is 1cal/g oC. It means that water can absorb and release large quantities of heat without a change in temperature. This is the reason why body temperature remains at 37 oC even when there's a change in the surrounding. This also explains why oceans and lakes exert an influence on the climate. If there were no large bodies of water, the earth would experience significant temperature variations. Water has high specific heat. The amount of energy required to raise the temperature of water by one degree Celsius is quite large. Because so much heat loss or heat input is required to lower or raise the temperature of the water, the oceans and other large bodies of water have relatively constant temperatures. Thus, many organisms living in the oceans are provided with a relatively constant environmental temperature. The high-water content of plants and animals living on land helps them to maintain a relatively constant internal temperature. The specific heat of water is five times greater than that of sand. Density in its liquid form. Water is the only substance that contracts when cooled. For most substances, their solid form is denser than their liquid form. This is because the H-bond is more extensive in its solid state than in its liquid state. Ice has an open structure because the hydrogen bonds could not get inside the hexagonal ring structure. This more open structure of the solid form of water causes the ice to have a smaller number of molecules packed in a given volume. This causes the mass to be lower. Hence, the density of ice is lesser than the liquid water, and, as a result, ice floats on water. This also causes the water in ponds or lakes to freeze from the top down. Water is most dense at 4 0C and then begins to expand again (becoming less dense) as the temperature decreases further. This expansion occurs because its Hydrogen bonds become more rigid and ordered. The expansion of water takes place even before it actually freezes. As water temperature drops, the colder water (0-4 0C), where it is less dense— rises to the pond or lake surface. It freezes to form a lid of ice. This ice insulates the water below from the wintry chill so that it is less likely to freeze. Organisms that inhabit the pond are able to survive the frigid winter below the icy surface. For most substances, solids are denser than liquids. But the special properties of water make it less dense as a solid. Ice floats on water! Strong hydrogen bonds formed at freezing 0 0C lock water molecules away from each other. When ice melts, the structure collapses, and molecules move closer together. Liquid water at 4 0C is about 9% denser 8 than ice. This property plays an important role in water ecosystems. Floating ice often insulates and protects animals and plants living in the water below. Surface tension. The hydrogen bond formation among water molecules causes water to have high surface tension, as described earlier. This high surface tension causes water to move from the roots of a tree to the top of very tall trees and explains why water moves into the fibers of a towel. This phenomenon is called capillarity. Water molecules at the surface (next to air) hold closely together, forming an invisible film. Water's surface tension can hold a weight that would normally sink. You can carefully float a paper clip on top of the water. Some aquatic insects, such as the water strider or pond skater, rely on surface tension to walk on water. Surface tension is essential for the transfer of energy from wind to water to create waves. Waves are necessary for rapid oxygen diffusion in lakes and seas. Next to mercury, water has the highest surface tension of all commonly occurring liquids. Cohesion—Water molecules stick to each other. This is due to the hydrogen bonds among the molecules. Water molecules at the surface have a much greater attraction for each other than for molecules in the air. This cohesiveness creates a high surface tension at the surface of the water. The water molecules at the surface crowd together, producing a strong layer as they are pulled downward by the attraction of other water molecules beneath them. Adhesion—Water molecules stick to other substances. You can see this property when water creeps up the inside of a drinking glass. Think of a sponge or a paper towel used to "soak up" spilled water. This is how water makes things wet. Water also clings to living things. Most plants have adapted to take advantage of water's adhesion that helps move water from the roots to the leaves. This is called capillary action. This can also be seen as blood moves through our capillaries, carrying nutrients to each cell within our body. One of the tallest plants is the redwood tree. Water moves from its roots to its leaves, more than 90 m above the ground. As a plant loses water through pores in the leaves, more water moves up from roots and stems to replace the lost water. The process of water loss by leaves is known as transpiration. Thermal properties - Water absorbs or releases more heat than many substances for each degree of temperature increase or decrease. Because of this, it is widely used for cooling and for transferring heat in thermal and chemical processes. Differences in temperature between lakes and rivers and the surrounding air may have a variety of effects. For example, local fog or mist is likely to occur if a lake cools in the surrounding air enough to cause saturation—small water droplets are suspended in the air. Large bodies of water, such as the oceans or the Great Lakes, have a profound influence on climate. They are the world's great heat reservoirs and heat exchangers and the source of much of the moisture that falls as rain and snow over adjacent landmasses. When water is colder than the air, precipitation is curbed, winds are reduced, and fog banks are formed. These properties of water are crucial in stabilizing temperatures on earth. Heat of vaporization. A large amount of heat is needed to vaporize a given amount of water. This causes a significant drop in temperature during evaporation. When molecules of water absorb heat energy, they move fast in the water. Eventually, the speed of movement of some molecules becomes so fast, allowing them to overcome the intermolecular attraction, detach from the multi-molecular water, form bubbles, and leave the water surface in the gas state. This property of water helps to cool down the body of living organisms. This is called evaporative cooling. In humans, body heat is used to vaporize sweat; in plants, heat is likewise used in converting liquid water to water vapor, which then escapes into the atmosphere. This natural process of vaporizing plant water is called transpiration. pH. Water molecules have a tendency to ionize. They dissociate into ions (charged particles), Hydrogen ions (H +), and hydroxide ions (OH-). In pure water, a very small number of water molecules form ions in this way. The tendency of water to dissociate is balanced by the tendency of Hydrogen ions and Hydroxide ions to reunite to form water. A neutral solution contains an equal number of Hydroxide ions and Hydrogen ions. A solution with a greater concentration of Hydrogen ions (H+) is said to be acidic. A solution with a greater concentration of Hydroxide (OH -) ions is said to be alkaline or basic. 9 Activity 8: Mind Power Direction: Identify what property of water each picture shows below. 1. _________ 2. _________ 3. __________ 4. ________ 5. _________ Activity 9: Word Hunt Part A. Directions: Fill in the blanks with words that correspond to each statement below. Choose the word from the word bank. covalent solvent deposition dissolve cohesion adhesion polar viscosity negatively positively 1. In a water molecule, the hydrogen and oxygen atoms are held together by ___________ bonds. 2. The electrons are not shared equally between covalently bonded atoms in a __________ molecule. 3. The polarity of water allows it to __________ most substances. 4. _________ refers to the attraction of molecules for other molecules of the same kind. 5. For adjacent water molecules, hydrogen bonds form between a hydrogen with a partial _________ charge and the __________ charged end of oxygen. Part B. Direction: Answer the following questions briefly and concisely. You may use a separate sheet of paper for your answer. Questions: 1. When you warm up oil and water, which temperature will rise faster? Support your answer. 2. What items can you gently "float" on the water surface? (e.g., paperclips, needles, etc.). Explain. 3. What happens to the bonds (Hydrogen bonds) when water boils? ---------------------------------------------------------- End of Module 1 – Lesson 4 ----------------------------------------------------------- ASSESSMENT (MODULE 1) Direction: Read each item carefully. Write only the letter that corresponds to your answer. 1. Which statement below is NOT consistent with the Kinetic Molecular Theory (KMT) A. Matter is made of particles that are constantly in motion. B. The amount of kinetic energy in a substance is related to its temperature. C. There is space between particles. D. Phase changes do not happen when the temperature of the substance changes sufficiently. E. There is an attractive force between particles called intermolecular forces. 2. Many substances, for example, salt (NaCl) and sucrose, dissolve quickly in water. Which property of water is related to this phenomenon? A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form hydrogen bonds with polar surfaces. C. Water is liquid at normal physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 3. Water drops that fall on a surface tend to form rounded drops or beads. A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form hydrogen bonds with polar surfaces. C. Water is a liquid at normal physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 4. If you put the end of a paper towel to colored water, the water will move up into the towel. Which property of water is related to this phenomenon? A. Water molecules are cohesive; they form hydrogen bonds with each other. 10  B. Water molecules are adhesive; they form H-bonds with polar surfaces. C. Water is a liquid at average physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 5. A paper clip can float on water. Which property of water explains this? A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form H-bonds with polar surfaces. C. Water is a liquid at average physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 6. When you place a straw into a glass of water, the water seems to climb up the straw before you even place your mouth on the straw. Which property of water is related to this phenomenon? A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form H-bonds with polar surfaces. C. Water is a liquid at normal physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 7. Water is most dense at about 4°C. As a result, the water at the bottom of a lake or the ocean usually has a temperature of about 4°C. Which property of water is related to this phenomenon? A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form H-bonds with polar surfaces. C. Water is a liquid at normal physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 8. If you drop a tiny amount of water onto a very smooth surface, the water molecules will stick together and form a droplet rather than spread out over the surface. Which property of water is related to this phenomenon? A. Water molecules are cohesive; they form hydrogen bonds with each other. B. Water molecules are adhesive; they form H-bonds with polar surfaces. C. Water is a liquid at normal physiological temperature. D. Water has high specific heat. E. Water has a high heat of vaporization. 9. Rank the matter based on decreasing the relative strength of attractive forces between particles. A. liquid, solid, gas D. liquid, gas, solid B. solid, liquid, gas E. liquid, solid, gas, plasma C. gas, liquid, solid 10. What happens to water molecules when cooled? A. The water molecules become excited. B. The water molecules slide past each other. C. The water molecules become fixed in position. D. The forces between molecules break. 11. The energy (usually in kilojoules) required to vaporize 1 mole of a liquid at a given temperature. A. Density C. Capillary action B. Viscosity D. Molar heat of vaporization 12. The tendency of a liquid to rise in narrow tubes or to be drawn into small openings. A. Density C. Capillary action B. Viscosity D. Molar heat of vaporization 13. A measure of a fluid's resistance to flow. A. Density C. Capillary action B. Viscosity D. Molar heat of vaporization 14. The hydrogen and oxygen atoms are held together by __________ bonds. A. Polar B. Viscosity C. Covalent D. Molecular 15. Electrons that are not shared equally between or among covalently bonded atoms creating a _________ molecule. A. Polar B. Viscosity C. Covalent D. Molecular 11  MODULE 2: TYPES OF SOLIDS, PHASE CHANGES, AND DIAGRAMS Most Essential Learning Competencies 1. Describe the difference in structure of crystalline and amorphous solids. (STEM_GC11IMFIIIa-c-104) 2. Interpret the phase diagram of water and carbon dioxide. (STEM_GC11IMFIIIa-c-107) 3. Determine and explain the heating and cooling curve of a substance. (STEM_GC11IMFIIIa-c-109) LESSON 1: Crystalline and Amorphous Solids Solid from what you have learned in Junior High School is one of the phases of matter. You mastered that solids have a more organized pattern arrangement of their particles than liquids and gases. Did you know that there are two main categories of solids? Yes, there are two types of solids! They can be classified as Crystalline solids and Amorphous solids. Amorphous solids are any non-crystalline solid in which the atoms and molecules are not organized in a definite pattern. In contrast, crystalline solids, or crystals, have distinctive internal structures that lead to distinctive flat surfaces or faces. The faces intersect at angles that are characteristic of the substance. Refer to the flow chart below for the summary on the classification of solids. Activity 1: Amorphous Solids vs. Crystalline Solids A. Directions: Categorize the following physical properties of solids by writing "A" for Amorphous Solids and "C" for Crystalline Solids on the space provided. _____ 1. Pseudo-solids or super-cooled liquids. _____ 2. Particles are arranged in a repeating pattern. They have a regular and ordered arrangement resulting in a definite shape. _____ 3. They have a sharp melting point. _____ 4. They do not have definite heat of fusion. _____ 5. Isotropic in nature. This means that magnitude of the physical properties is the same along with all directions of the solid. _____ 6. Anisotropic in nature. This means that the magnitude of physical properties (such as refractive index, electrical conductivity, thermal conductivity, etc) is different along with the crystal's different directions. _____ 7. When cutting with a sharp edge, the two new halves will have smooth surfaces. _____ 8. When cutting with a sharp edge, the two resulting halves will have irregular surfaces. _____ 9. They are rigid solids, and applying mild forces will not distort their shape. _____ 10. They are not rigid, so that that mild effects may change the shape. 12 B. Directions: Complete the table by describing the differences between the two kinds of solids. Types of Solids Physical Properties Amorphous Crystalline Nature Geometry Melting Heat of Fusion (The change in enthalpy when a substance is heated to change its state from solid to liquid.) Isotropism Cleavage Rigidity --------------------------------------------------------------- End of Module 2 – Lesson 1 ------------------------------------------------------ LESSON 2: Phase Changes and Phase Diagram Activity 2: Phase Diagrams Directions: You may choose to color the diagram to your liking as long as it will lead you to understand it better. Also, unleash your creativity in giving the diagram a name. Make it concise yet catchy, and it should wrap everything that is seen in the diagram. Moreover, answer the following Guide Questions provided right below each diagram. Name of Diagram A. ______________________________ Questions: 1. What do you think happens at Point B? 2. Based on the diagram, explain how sublimation takes place? 3. Based on the diagram, explain how deposition takes place? 4. Based on the diagram, what are the factors that affect the phase changes of matter? B. ______________________________ C. ______________________________ 13 Questions: 1. Indicate the triple points of diagrams B and C. 2. Indicate the boiling points of diagrams B and C. 3. Indicate the freezing points of diagrams B and C. In contrast to the phase diagram of water, the phase diagram of CO2 (Figure C of Question 4 in Activity 3 below) has a more typical melting curve, sloping up and to the right. The triple point is −56.6°C and 5.11 atm, which means that liquid CO2 cannot exist at pressures lower than 5.11 atm. At 1 atm, therefore, solid CO 2 sublimes directly to the vapor while maintaining a temperature of −78.5°C, the normal sublimation temperature. Solid CO 2 is generally known as dry ice because it is a cold solid with no liquid phase observed when it is warmed. Activity 3: Heating and Cooling Curves of a Substance D. ______________________________ E. ______________________________ Questions: 1. Which of the two graphs shows a heating curve? Graph D or Graph E? _________ 2. Which of the two graphs shows a cooling curve? Graph D or Graph E? _________ 3. Which process releases heat? ___________________ 4. Which process absorbs heat? ___________________ TYPES AND PROPERTIES OF SOLIDS Amorphous Solids have fixed shape and volume; however their particles are not arranged in regular geometric patterns. The term 'amorphous,' in Greek roots, translates to "without form". They are referred to as "super-cooled liquids" since this type of solids appears to have been cooled at very low temperatures, and their viscosity are very high, preventing the flow of the The atomic arrangements in (A) a crystalline liquid. solid, (B) an amorphous solid During the cooling process, the particles are trapped in the disarranged characteristic manner of liquids. Many polymers are amorphous solids. Examples of amorphous solids are glass, rubber, gels, some plastics, and nanostructured materials. Crystalline Solids referred to as "true solids," are solids with highly ordered arrangements of particles (atoms, ions, and molecules) in microscopic structures. The latter make up a crystal lattice that accounts for the structure of the solid. Examples of crystalline solids include salt (sodium chloride), diamond, and sodium nitrate. There are 4 types of crystalline solids, namely: Ionic Solids, Covalent Solids, Molecular Solids, and Metallic Solids. Ionic solids have positive and negative ions held together by electrostatic attractions, which can be quite strong. It accounts for the high melting points many ionic crystals have. Although ionic solids are hard, they also tend to be 14 brittle, and they shatter rather than bend. Ionic solids do not conduct electricity, but once it is molten or dissolved, it may be a good conductor since their ions are free to move. Examples of this type of solids are sodium chloride and nickel oxide. A covalent solid contains a three-dimensional network of covalently bonded atoms. Some solids form covalent bonds resulting in the formation of molecules. However, in some instances, molecules may not form but rather covalent networks that extend throughout the solid Structure of (A) diamond and (B) graphite crystals. These covalent solids are very hard, have high melting points, and often low thermal conductivity. The structures of diamond and graphite explain these observed properties. Graphite is an exceptional example, composed of planar sheets of covalent crystals held together in layers by noncovalent forces. Unlike typical covalent solids, graphite is very soft and electrically conductive. Molecular solids are soft, have low to moderately high melting points, and have poor thermal and electrical conductivity. The particles of this type of solids are either composed of atoms or molecules held together by intermolecular forces. The strengths of the attractive forces between the units present in different crystals in molecular solids vary widely. Small symmetrical molecules (nonpolar molecules), such as H2, N2, O2, and F2, have weak attractive forces and form molecular solids with very low melting points (below −200 °C). On the other hand, molecular solids composed of molecules with permanent dipole moments (polar molecules) melt at higher temperatures. Examples of molecular solids include ice that melts at 0 °C and table sugar that melts at 185 °C. Metallic Solids/Crystals are often described as a uniform distribution of atomic nuclei within a "sea" of delocalized electrons. The atoms within such a metallic solid are held together by a unique force known as metallic bonding that gives rise to many useful and varied metallic properties. This type of solids has high thermal and electrical conductivity, metallic luster, and malleability. Many are very hard and quite strong. A solid substance may change into another phase under certain conditions. These phase changes are better shown in a graphical way known as diagrams, which show the effects of pressure and temperature on a specific phase of matter. PHASE CHANGES AND PHASE DIAGRAMS The transformation of matter from one phase to another is what we call phase change. It always involves absorption or release of heat. The illustration shows the different phase changes of matter. You even have already encountered some of these processes in the lower grade levels. Evaporation and melting processes absorb heat in order to change material from one phase to another. On the other hand, condensation and freezing involve the release of heat to change the phase of the material. Sublimation is the process of changing the phase of a substance from solid to gas without passing the liquid phase. It involves the absorption of heat. Deposition is the process of changing the phase of a substance from gas to solid, and the change involves the release of heat. Shown is the general shape for a phase diagram exhibiting three states or phases of matter (solid, liquid, and gas). Factors that affect phase change are pressure and temperature. Point B is the Triple Point. It is the point at which all three distinct phases of matter (solid, liquid, gas) coexist. Any other point on the three curves represents an equilibrium between two phases. The segment between Point B and Point C is where evaporation or condensation takes place. The line is the vapor-pressure curve of the liquid. At any point in this segment (considering pressure and temperature), the substance 15 is at equilibrium between the liquid and gas phases. This curve ends at Point C known as the Critical Point. It is the point in temperature and pressure on a phase diagram where the liquid and gaseous phases of a substance merge together into a single phase. Beyond this point, the liquid and gas phase become indistinguishable. The diagram above shows the phase change of (a) Water (H2O) and (b) Carbon dioxide (CO2) The gas phase is the most stable at low pressure and temperature. Solid phase is stable upon extending to low temperature and high pressures, while the stability of liquid phase lies between the range of the other two regions of the solid and gas phases. The triple point of water is represented by the letter A (0.0098 0C, 4.58 torr). The normal melting or freezing point of water is represented by the letter B (0 0C, 1.00 atm while the normal boiling point is represented by the letter C (1000C, 1.00 atm). The critical point is represented by the letter D (3740C, 218 atm). On the other hand, the triple point of CO2 is represented by the letter X (-56.40C, 5.11 atm). Its normal sublimation point is represented by the letter Y (-78.50C, 1.00 atm), and the critical point is represented by the letter C (-31.10C, 73.0 atm). HEATING AND COOLING CURVES Graph (a) Graph (b) The Graphs show (a) Heating Curve and (b) Cooling Curve. A typical heating curve for a substance shows changes in temperature that result as the substance absorbs increasing amounts of heat. Plateaus in the curve (regions of constant temperature) are exhibited when the substance undergoes phase transitions. 16  Graph (a) generally shows an increase in temperature as the substance changes from solid to liquid and then to gas. Melting and evaporation require the absorption of heat for the process to take place. Graph (b) on the other hand, shows a decrease in temperature indicating the release of heat as the processes involved take place. Condensation and freezing involve the release of heat as the substance changes from one phase to another phase. Segment B shows a constant temperature despite the continued input of heat (from an external heat source). As the solid melts, its temperature does not rise. All of the energy that is being put into the solid substance goes into the melting process and not into any increase in temperature. During the melting process, the two phases – solid and liquid coexist and are in equilibrium with one another. Continued heating of the substance after it has completely melted will now increase the kinetic energy of the liquid molecules and the temperature will start to rise (segment C). Assuming that the atmospheric pressure is standard, the temperature will rise steadily until it reaches 100°C. At this point, the added energy from the heat will cause the liquid to begin to vaporize. As with the previous phase change, the temperature will remain at 100°C while the water molecules are going from the liquid to the gas or vapor state (Segmet D). Once all the liquid has completely boiled away, continued heating of the steam (remember the container is closed) will increase its temperature above 100°C (Segment E). Activity 4. Crystalline Solids Direction: Identify the type of crystalline solid (metallic, network covalent, ionic, or molecular) formed by each of the following substances. __________1. CaCl2 __________6. CH3CH2CH2CH3 __________2.SiC __________7. HCl __________3.N2 __________8. NH4NO3 __________4.Fe __________9. K3PO4 __________5.C (graphite) __________10. SiO2 Question: Explain why ice, which is a crystalline solid, has a melting temperature of 0 °C, whereas butter, which is an amorphous solid, softens over a range of temperatures. Activity 5. Phase Diagrams Directions: At Standard Temperature and Pressure (STP), Bromine (Br2) is in the liquid phase. It undergoes sublimation once it reaches -250C and a pressure of 101.3 kPa. Please refer to the phase diagram of Bromine below and use the diagram to answer the questions that follow. 1. Label each region in the diagram as solid, liquid, or gas. 2. Label the triple point on the diagram. What are the temperature and pressure of the triple point? ________________ 3. At what point will Bromine turns into solid? ______________________ 4. At what point will Bromine turns into liquid? ______________________ 5. At what point will Bromine turns into a gas? ______________________ Activity 6: Perfect Match Directions: Read each statement below and decide whether it is TRUE or FALSE. 1. Phase change neither involves absorption nor release of heat. 2. The triple point of a substance is the temperature and pressure at which the three phases of that substance co- exist. 3. Ionic, covalent, molecular, and metallic solids are classified as crystalline solids. 4. Crystalline solids and amorphous solids are the major categories of solids. 5. A cooling curve is a line graph that represents the change of phase of matter, typically from a gas to a solid or a liquid to a solid. ------------------------------------------------------------- End of Module 2 – Lesson 2 -------------------------------------------------------- ASSESSMENT (MODULE 2) Direction: Write only the letter of your answer. 1. What type of solids is made up of unlike charge of particles and later results in the formation of electrostatic force? A. Ionic Solids B. Covalent Solids C. Molecular Solids D. Metallic Solids 2. These are solids which particles are arranged in regular geometric patterns. 17  A. Crystalline Solids B. Amorphous Solids C. Melted Solids D. Frozen Solids 3. Which of the following statements is TRUE about the triple point? A. A substance can exist in all of the different states depending on temperature and pressure. B. A substance can exist in all of the different states depending on temperature. C. A substance can exist in all of the different states depending on pressure. D. A substance cannot exist in all of the different states regardless of temperature and pressure. 4. Some solids have fixed shapes and volume but their particles are not arranged in regular geometric patterns. What type of solid is described? A. Crystalline Solids B. Amorphous Solids C. Melted Solids D. Frozen Solids 5. The following are crystalline solids EXCEPT? A. Ionic Solids B. Covalent solids C. Metallic Solids D. Amorphous Solids 6. Which of the following is a/are characteristic/s of an amorphous solid? I. Pseudo-solids or super-cooled liquids. II. Particles are arranged randomly. They do not have an ordered arrangement resulting in irregular shapes. III. They do not have definite heat of fusion. IV. They are rigid solids and applying mild forces will not distort their shape. A. I & II only C. All of the situations given B. All choices EXCEPT IV D. None of the situations given For questions Nos. 7 to 9 refer to the Phase Diagram and the following choices: A. At any point on that specific line, the substance is both solid and liquid B. At any point on that specific line, the substance is both solid and gas C. At any point on that specific line, the substance is both liquid and gas D. At any point on that specific line, the substance is solid, liquid, and gas 7. What does the line from Point B to Point D indicate? 8. What does the line from Point B to Point C indicate? 9. What does the line from Point A to Point B indicate? 10. What happens at the critical point? I. The vapor pressure curve ends at this point. II. The temperature above where the gas cannot be liquefied no matter how much pressure is applied III. The substance at this point cannot be distinguished as a liquid or a gas IV. The condition of temperature and pressure where all phases of matter exist in equilibrium A. I & II only C. All of the situations given B. All choices EXCEPT IV D. None of the situations given For numbers 11 - 15 refer to the Phase Diagram for an unknown compound X 11. What is the critical temperature of compound X? A. ~400 0C B. ~500 0C C. ~770 0C D. ~800 0C 12. If you were to have compound X in your room for quite some time, what do you think the phase of the compound could be? A. It is most likely a solid B. It is most likely a liquid C. It is most likely a gas D. All of the choices since it is at equilibrium. 13. At what temperature and pressure will all phases (solid, liquid, and gas) coexist? A. ~51 atm, 3500C C. ~51 atm, 4000C 0 B. 350 C, ~61 atm D. 3500C, ~51 atm 18 14. If I have compound X at 100 0C and 45 atm, what do you think will happen if I raise the temperature to 400 0C? It will undergo ___________. A. condensation B. deposition C. evaporation D. sublimation 15. Why can’t compound X be boiled at 2000C? It’s because it only forms liquid when it ______________________. A. reaches above 3500C temperature. C. reaches above 2500C temperature. B. reaches above 3000C temperature. D. reaches above 3100C temperature. MODULE 3: SOLUTION CONCENTRATION AND STOICHIOMETRY Most Essential Learning Competencies 1. Use different ways of expressing the concentration of solutions: percent by mass, mole fraction, molarity, molality, percent by volume, percent by mass, ppm. (STEM_GC11PPIIId-f-111); and 2. Perform stoichiometric calculations for reactions in solution. (STEM_GC11PPIIId-f-112) 3. Describe laboratory procedures in determining the concentration of solutions. (STEM_GC11PPIIId-f-119) LESSON 1: Concentration of Solution The concentration of solution refers to the amount of solute present in a given amount of solvent or solution. A solution can be qualitatively described as dilute: a solution that contains a small proportion of solute relative to solvent, or concentrated: a solution that contains a large proportion of solute relative to the solvent. Quantitatively, one type of solution may be prepared and expressed in different concentrations. One teaspoon of sugar in a cup of water is a different solution from a cup of water with five teaspoons of sugar. The amount of solute in a solution may be expressed in several ways. These include percent concentration - by mass, by volume or by mass- volume, mole fraction, molality, molarity, and parts per million. A. Percent Concentration (by Mass, by Volume and by Mass-Volume) I. Percent by Mass. This expresses the mass of solute per 100g of solution. In most applications, “percent concentration” means weight/weight percent (% weight/weight) which is equal to the number of grams of solute per 100 grams of solution. A solution that contains 30% by mass sugar means that the solution contains 30g of sugar dissolved in 70g of water. It also means that there are 30g of sugar per 100g of solution. The formula for percent by mass is: 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐮𝐥𝐭𝐞 % (𝐰𝐭⁄𝐰𝐭) = 𝐱 𝟏𝟎𝟎 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐥𝐮𝐭𝐞 + 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐥𝐯𝐞𝐧𝐭 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐥𝐮𝐭𝐢𝐨𝐧 = 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐥𝐮𝐭𝐞 + 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐬𝐨𝐥𝐯𝐞𝐧𝐭 Sample Problem 1. If 7.5 g of sodium nitrate is dissolved in 85 ml of water, calculate the concentration of sodium nitrate in the solution. Solution: To find the total mass of solution, we must add the mass of solute, 7.5 grams, and the mass of water. Since the density of water is 1.0 g/ml, 85 ml of water is equivalent to 85 grams of water. Therefore, the total mass of the solution is 92.5 grams. 19 2. A common laboratory reagent is 10% (wt/wt) NaOH solution. How would you prepare 750 grams of the solution? Solution: You can prepare 750 g of 10% (wt/wt) NaOH solutions as follows: Activity 1: Knowledge Check on Percent by Mass Concentration! Direction: Solve the given problem. You may use a separate sheet of paper. If 7.5 g of sodium nitrate, NaNO3 is dissolved in 85ml of H2O (density of H2O = 1.0g/ml), calculate the %(wt/wt) concentration of H2O in the solution. II. Percent by Volume, Volume/Volume Percent or % (vol/vol) When both solute and solvent are liquids, it is sometimes convenient for you to describe the concentration as percent by Volume (%vol/vol) which is the number of Volume of the solute in 100 volume of solution. “Volume” may be any volume unit provided you use the same unit for both solute and solution. Sample Problem 1. A 40% (vol/vol) solution of ethylene glycol in water is used to give protection to a car’s cooling system. What Volume of ethylene glycol would you use to make five liters of this solution? Solution: a) You can translate the label 40% (vol/vol) ethylene glycol as: In these factors, “vol” may be any volume unit you want – mL, liter, or whatever is required. The problem specified liters, therefore you can use the second factor and compare the required volume as: b) You can calculate the needed Volume of ethylene glycol using the first conversion factor 20 or you may solve it this way: Activity 2: Knowledge Check on Percent by Volume Concentration! Direction: Solve the given problem. You may use a separate sheet of paper. A brand of rubbing alcohol says, it contains 70% (vol/vol) isopropyl alcohol. How many mL of isopropyl alcohol are there in 600 ml of the solution in the bottle? III. Weight/Volume Percent or % (wt/vol) When it is impractical to express both the solute and solvent in mass or volume units, a hybrid expression for percent is to be used. Hybrid because their units do not cancel as they ought to. Sample Problem 1. A solution is prepared by dissolving 5.0 grams of glucose in enough water to make 250 mL of solution. Calculate % (wt/vol) glucose. Solution: Fill in the given to the equation, 2. A 50 mL of 12% by mass-volume solution was used in an experiment. How many grams of solute does the solution contain? Activity 3: Knowledge Check on Percent by Mass/Volume Concentration! Direction: Solve the given problem. You may use a separate sheet of paper. The label of betadine skin cleanser says 7.5 % solution. Taking it to be % (wt/vol), how many grams of betadine (Providone-Iodine) are present in 50 mL bottle? 21 B. Mole Fraction, X The variation in some physical and/or chemical properties of a solution especially those containing only two components is sometimes described over the entire range of concentration. The concentration of the solution is best described by mole fraction or mole percent. Mole fraction is usually designated as X that relates to the number of moles of a particular solute to the total number of moles in the solution. Sample Problem 1. Compute the mole fraction of acetone X (acetone) and of chloroform (X chloroform) in a solution prepared by mixing 50.0 g each of acetone (Molar Mass = 58.0) and chloroform (Molar Mass = 119.5) Solution: a. The first step is for you to compute the number of moles (n) of each substance as well as the total number of moles. (solute: acetone; solvent: chloroform) b. Compute the X of acetone and X of chloroform 22 Activity 4: Knowledge Check on Mole Fraction! Direction: Solve the given problem. You may use a separate sheet of paper. Calculate the mole fraction of sulfuric acid (H2SO4) in 8% (wt/wt) aqueous H2SO4 solution (molar masses: H2SO4 = 98 g/mol, H2O = 18 g/mol C. Molar Concentration or Molarity, M Molarity is the ratio of the moles of solute to the volume of solution in liters. Sample Problem 1. Calculate the molar concentration of the solution that contains 15 grams of potassium hydroxide (KOH) in 225 ml of solution. (Molar mass of KOH = 56 g/mol) Solution: 1. Convert 15 grams of KOH to moles using conversion factor, 1 mol KOH = 56 g (the molar mass of KOH). 2. Convert 225 ml of solution to Liter of solution using the conversion factor: 1L = 1000 ml. Use the formula in computing molarity and substitute the values obtained above. Activity 5: Knowledge Check on Molarity! Direction: Solve the given problem. You may use a separate sheet of paper. Calculate the molar concentration of a solution that contains 23 g of potassium hydroxide. KOH in 250 ml of solution. Molar mass of KOH is 56 g/mol. D. Molal Concentration or Molality, m Solutions may also be expressed in molality. The molality of a solution is the ratio of moles of solute to the mass of solvent in kilograms. Sample Problem 1. Determine the molal concentration, m of a solution that contains 18 grams of NaOH in 100 ml of water. The molar mass of NaOH is 40 g/mole. Solution: a. Convert 18 grams of NaOH into moles using the molar mass of NaOH 23  b. Convert 100 ml of water into grams using the density of water, 1.0 g/ml. Then convert the grams to kilograms using the conversion factor, 1Kg = 1000 c. Use the formula in computing molality and substitute the values obtained above. Activity 6: Knowledge Check on Molality! Direction: Solve the given problem. You may use a separate sheet of paper. How many grams of solute is present in 0.4 moles Magnesium hydroxide, Mg(OH) 2 in 550 g water, H2O? The molar mass of Mg(OH)2 is 58 g/mol. E. Parts per Million, ppm For a very dilute solution where the concentration of the solute is very low like pesticide residue in water or heavy metal like Hg2+ concentration in effluents, or even hardness of water, it is convenient to express concentration in terms of parts per million, ppm. Parts per million (ppm) expresses the number of parts of solute per one million parts of the solution. The last expression is approximately true for water as a solvent because the density of water is 1.0 g/ml. Also, for a very dilute solution, the amount of the solution could be equated to the amount of the solvent, water. While these concentrations are very small, but we should not neglect their importance. Some of the industrial pollutants that are being released daily into the water we drink and the air we breathe can be extremely harmful in concentrations as small as 1 ppm. Sample Problem 1. A water sample was reported to contain 250 ppm CaCO3. How many grams of CaCO3 is present in 4 liters of water. Solution: 250 ppm CaCO3 can be translated as 250 mg/liter of solution. Since this is a very dilute solution and the solvent is water, the liter of solution could be equated to the Volume of water. So, we can use the expression below, The problem asks for the mass in grams of CaCO3 present in 4 liters of water: 24 Activity 7: Knowledge Check on Molality! Direction: Solve the given problem. You may use a separate sheet of paper. A commercial pesticide formulation contains 1.0 g deltametrin in 1L solution. What is its concentration in ppm? ------------------------------------------------------End of Module 3 – Lesson 1 ---------------------------------------------------------------- LESSON 2: Solution Stoichiometry Activity 8: Remember Me! Direction: Identify the concept related to solution stoichiometry that is described in each statement. Get your answer from the word bank. Law of Multiple Proportion Superscript Law of Constant Composition Balanced Hess law Coefficient Law of Definite Proportion Subscript Law of Conservation of Mass 1. It states that the mass of the products is equal to the mass of the reactants. 2. It states that the mass of one element combines with a fixed mass of another element in a ratio of whole numbers 3. It states that all samples of a given chemical compound have the same elemental composition 4. For a chemical equation to be correct, it must be _________. 5. A number written before the symbol of an element or formula of a compound. Activity 8 enabled you to recall important concepts and laws related to chemical equations. Mass relations are based on the three important laws. If you keep these laws in mind, you'll be able to make valid predictions and calculations for chemical reactions including those that involve solutions. Stoichiometry Involving Solution In General Chemistry 1, you have done several calculations using a balanced chemical equation. Stoichiometry deals with solving quantitative problems using a balanced chemical equation. Recall that in these types of calculations, we used the following steps: 1. Convert from the given units to moles, if not given in moles; 2. Convert from moles of the given quantity to moles of the desired quantity, using the balanced equation; and 3. Convert from moles to any other desired units. The number of moles of a substance can be related to its molar mass and number of molecules. It can also be related to the volume at Standard Temperature and Pressure (STP). Sample Problem 1. What volume of 0.556 M HCl has enough HCl to combine exactly with 25.4 mL of 0.458 M NaOH? The equation for the reaction is, HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) Solution: 1. Find the moles of NaOH in 25.4 mL of 0.458 M NaOH 2. From the expression of molarity, M = n of solute / L of solution 3. The molarity of the solution gives you two conversion factors: 4. When you use the first conversion factor, you can get the moles of NaOH as: 25  5. Use the coefficients of NaOH and HCl from the balanced equation to calculate how much 0.0116 mole NaOH is equivalent to in mole of HCl. From the balanced equation, the ratio of coefficients of NaOH and HCl is 1:1 6. Find the Volume of HCl using 0.556 M, the given molarity of HCl aqueous solution. 2. Calculate the mass (in grams) of calcium nitrate, Ca(NO3)2 that can be produced by reacting 136 ml of 4.00 M nitric acid, HNO3 with excess calcium hydroxide, Ca(OH)2. The molar mass of Ca(NO3)2 = 164 g/mol Solution: 1. Write the balanced equation for the reaction 2HNO3 + Ca(OH)2 → 2H2O + Ca(NO3)2 2. Use the molarity and Volume of the solution to get the number of moles of HNO 3. 3. Find the number of moles of Ca(NO3)2 using the stoichiometric factor: 4. Find the mass of Ca(NO3)2 using its molar mass: Activity 9: Knowledge Check on Solution Stoichiometry! Direction: Solve the given problem. You may use a separate sheet of paper. How many mL of 0.250 M HCl would react exactly with 30.0 mL of the 0.150 M solution of Ca(OH)2 solution? The chemical reaction involved is: HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) ------------------------------------------------------------- End of Module 3 – Lesson 2 ------------------------------------------------------- LESSON 3: Effects of Solution on the Colligative Properties of Solution Colligative properties of solutions are properties that depend upon the concentration of solute molecules or ions, but not upon the identity of the solute. Colligative properties include vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure. 26 Lowering the Vapor Pressure Vapor pressure is the pressure of a vapor in thermodynamic equilibrium with its condensed phase in a closed container. When a non- volatile solute is dissolved in a solvent, the vapor pressure of the solvent is lowered. The presence of solute decreases the rate of escape of solvent molecules resulting in lower vapor pressure. Boiling Point Elevation The boiling point of a liquid is defined as the temperature at which the vapor pressure of that liquid equals the atmospheric pressure (760mm Hg). The addition of the solute increases the boiling point of the solution. The atmospheric pressure remains the same while the vapor pressure of the solution is lowered resulting in the increase of the difference in atmospheric pressure and vapor pressure of the solution. Therefore, a higher temperature is required to boil the solution. Freezing Point Depression Normal freezing or melting point is the temperature at which solid and liquid are in equilibrium under 1 atm. The addition of solute will decrease the vapor pressure and so will decrease the freezing point. For a liquid to freeze it must achieve a very ordered state that results in the formation of a crystal. If there are impurities in the liquid, i.e. solutes, the liquid is inherently less ordered. The presence of impurities in a liquid or a substance makes a variation in the freezing point by making them low or high. Therefore, a solution is more difficult to freeze than a pure solvent so a lower temperature is required to freeze the liquid. Osmotic Pressure This is the external pressure that must be applied to the solution to prevent it from being diluted by the entry of solvent via osmosis. Osmosis is the movement of solvent particles across a semipermeable membrane from a dilute solution (low concentration) into a concentrated solution. The solvent moves to dilute the concentrated solution and equalize the concentration on both sides of the membrane. Osmotic pressure is directly proportional to the concentration of the solution. Therefore, doubling the concentration will also double the osmotic pressure. The osmotic pressure of two solutions having the same molal concentration is identical. Activity 10: True or False Direction: Read each statement and evaluate if it is correct (True) or not (False). Write your answer in the given space. _______1. Colligative properties arise from the fact that solute affects the concentration of solvent. _______2. Vapor pressure is a colligative property. _______3. Lowering of vapor pressure is not dependent on the number of species present in the solution. _______4. The colligative properties of solutions depend on the nature of the solute and the solvent. _______5. Colligative molality is the molality times the number of solute particles per formula unit. _______6. Osmotic pressure is directly proportional to the concentration of the solution. ______ 7. Relative lowering of vapor pressure is a colligative property. ______ 8. The boiling point of a solution decreases in direct proportion to the molality of the solute. ______ 9. When a non-volatile solute is dissolved in a solvent, the vapor pressure of the solvent is lowered. ______10.The depression of the freezing point is directly proportional to the molality of the solvent. ------------------------------------------------------------- End of Module 3 – Lesson 3 -------------------------------------------------------- 27 ASSESSMENT (MODULE 3) Direction: Write only the letter of your answer. 1. A solution contains 28% phosphoric acid by mass. This means that: A. 100 ml of this solution contains 28 g of phosphoric acid B. 1 ml of this solution contains 28 g of phosphoric acid C. 1 L of this solution contains 28 mL of phosphoric acid D. 1 L of this solution has a mass of 28 g 2. Calculate the concentration in % (wt/wt) of a solution containing 20.0 g of NaCl dissolved in 250.0 g of H2O. A. 6.76% (m/m) B. 7.41% (m/m) C. 8.00% (m/m) D. 8.25% (m/m) 3. What is the concentration in % (m/v) of a NaCl solution prepared by dissolving 9.3 g of NaCl in a sufficient amount of water to give 350 mL of solution? A. 3.26% (m/v) B. 0.455% (m/v) C. 37.6% (m/v) D. 2.66 (m/v) 4. Calculate the grams of NaOH present in 5.0 mL of a 1.0% (m/v) NaOH solution. A. 0.050 g B. 0.10 g C. 0.50 g D. 1.0 g 5. How many grams of NaOH are there in 500.0 mL of a 0.175 M NaOH solution? A. 14 g B. 3.50 g C. 114 g D. 0.00219 g 6. What is the molarity of an aqueous solution containing 22.5 g of sucrose (C 12H22O11) in 35.5 mL of solution? A. 1.85 M B. 1.85 m C. 1.85 M D. 1.85 m 7. How many grams of H3PO4 are in 175 mL of a 3.5 M solution of H3PO4? A. 4.9 B. 20 C. 60 D. 612 8. What is the molality of 6 grams of table salt, NaCl in 10 grams of a solution? (MM of NaCl = 58.45 g/mol) A. 1.027 m B. 10.27 M C. 1.027 m D. 10.27 M 9. What is the molality of an aqueous NaOH solution made with dissolving 5.0 Kg of water and 3.6 moles of NaOH? A. 3.6 m C. 1.4 m B. 0.72 m D. 0.090 m 10. After mixing 10.00 g of compound A with 20.00 g of compound B, it is found that the mole fraction of compound A is 0.400. The mole fraction of compound B must be: A. 0.200 B. 0.400 C. 0.600 D. 0.800 11. What is the mole fraction of CaCl2 (molar mass=111 g/mole) when 3.75 g of it is placed in 10.1 g of water (molar mass=18.0 g/mole)? A. 0.8752 B. 0.5280 C. 0.1043 D. 0.0568 12. How many grams of Calcium phosphate, Ca3(PO4)2 can be produced from the reaction of 2.50 L of 0.250 M Calcium chloride, CaCl2 with an excess Phosphoric acid, H3PO4. The balanced equation is, 3CaCl2 + 2H3PO4 → Ca3(PO4)2 + 6HCl A. 34.4 g B. 46.6 g C. 76.4 g D. 64.6 g 13. How many liters of 0.53 M HCl is required to neutralize 0.78 g sodium carbonate, Na 2CO3? The balanced equation is 2HCl + NaCO3 → 2NaCl + H2CO3 A. 0.028 L B. 0.082 L C. 1.128 L D. 1.182 L 14. Adding salt to water will make the freezing point of the resulting solution ______0oC. A. equal to B. higher than C. lower than D. greater than or equal to 15. As the temperature increases, the solubility of a gas in a carbonated drink _________. A. decreases B. increases C. stays the same D. decreases then increases 28