General Chemistry 2 Reviewer - Basic Properties of Water (PDF)

**Properties of water** Properties of water arise from its structure. Figure 1 below shows the molecular structure of water. Each oxygen (O) and hydrogen (H) atom attains a filled outer level by sharing electrons in single bonds. With two bonding pairs and two lone pairs around O and a large ele...

**Properties of water** Properties of water arise from its structure. Figure 1 below shows the molecular structure of water. Each oxygen (O) and hydrogen (H) atom attains a filled outer level by sharing electrons in single bonds. With two bonding pairs and two lone pairs around O and a large electronegativity difference in each O¬H bond, the H2O molecule is bent and highly polar. Open photo -This arrangement allows each molecule of water to engage in four hydrogen - bonds with its neighboring water molecules as seen Figure 2 below. From these basic atomic and molecular facts emerges some unique and remarkable macroscopic behavior of water.  **The Unique Properties of Water** -**Water is a good solvent** One of the unique properties of water is its ability to dissolve a large variety of chemical substances. Water dissolves salts and other ionic compounds, including polar covalent compounds like alcohols and organic substances that has the ability to form hydrogen and bond with water. For example, oxygen dissolved in water are needed by animals like fish that are provided by plants while plants need carbon dioxide. This is the reason why water is sometimes called the universal solvent because it can dissolve so many things. **B. Water has a high specific heat** -The specific heat of water is 1 calorie/g-oC (4.18 J/g-oC). Among the many liquids, water has the highest specific heat. Specific heat is the amount of heat or energy needed to raise the temperature of one gram of a substance by 1ᵒ C. If the temperature of water rises slightly, water can absorb a large amount of heat. But in order to raise the temperature of water, the intermolecular hydrogen bonds should break. On the contrary, it is also true that water can give off much heat with only a slight decrease in its temperature. This allows large bodies of water to help moderate the temperature on earth. The boiling point of water is unusually high Many compounds similar in mass to water have much lower boiling points. The strong intermolecular forces in water permit it to be a liquid at a large range of temperatures. Figure 3 below shows the plot of broken lines that directed one to the estimated boiling points of HF, H2O and NH3 if H bonding was not present in these three substances. Open photo **Surface Properties of water** -Hydrogen bonding is also responsible for water's high surface tension and high capillarity. Except for some molten metals and salts, water has the highest surface tension among any liquids. It is the reason for a spherical shape of water droplets when placed on a waxy surface or plastic sheets.Surface tension is a measure of the tendency of a liquid to maintain a minimum surface area. The surface tension of a liquid also increases as intermolecular forces increase in strength. Water has a strong surface tension because of its strong hydrogen bonds. -It keeps plant debris resting on a pond surface, providing shelter and nutrients for fish and insects. -High capillarity means water rises through the tiny spaces between soil particles, so plant roots can absorb deep groundwater during dry periods. **The Unusual Density of Solid water** -Unlike all other liquids, the molecules in solid water are actually farther apart than they are in liquid water. As shown in Figure 4 below, when solid water forms, the hydrogen bonds result in a very open structure with unoccupied spaces causing the solid to occupy a larger volume than the liquid. This makes ice less dense than liquidwater,causing ice to float on water.  **Solutions** **Mixtures** Mixture: a combination of two or more substances that do not combine chemically, but remain the same individual substances; can be separated by physical means. Two types: -- Heterogeneous -- Homogeneous **Heterogeneous Mixture** "Hetero" means "different" Consists of visibly different substances or phases (solid, liquid, gas) Can be separated by filtering **Homogeneous Mixture** "Homo" means the same has the same uniform appearance and composition throughout; maintain one phase (solid, liquid, gas) Commonly referred to as solutions **Solution** Solution: a mixture of two or more substances that is identical throughout (homogeneous) can be physically separated composed of : **[solutes-]**the substance being dissolved. **[solvents]**-the substance that dissolves the solute. A solution is a homogeneous mixture made up of atoms, ions, or molecules. It has two mediums , namely, the solute ( the dissolved medium) and the solvent (the dissolving medium). In a mixture of instant coffee powder and hot water, the coffee powder is the solute and the solvent is water. In a mixture of two liquids where water is one of the substances, the water is always the solvent, and the solution is called an aqueous solution. Open photo **AIR**  **LIQUID SOLUTION** Open photo **Concentration** the amount of solute dissolved in a solvent at a given temperature described as dilute if it has a low concentration of solute dissolved described as concentrated if it has a high concentration of solute dissolved [Unsaturated-] has a less than the maximum concentration of solute dissolved [Saturated -] has the maximum concentration of solute dissolved (can see solid in bottom of solution) [Supersaturated-]contains more dissolved solute than normally possible (usually requires an increase in temperature followed by cooling) **Solubility** refers to the amount of solute that can dissolve in a given amount of solvent at room temperature under given conditions. Ex. if 1.0 g of sugar is placed in 100 g of water at 35 ̊C, all sugar dissolves. **Factors Affecting Solubility** 1\. Nature of solute and solvent 2\. Effect of temperature Miscible liquids can easily dissolve in one another. Immiscible liquids are not soluble in each other. **Polarity and Dissolving** Chemists use the saying "like dissolves like": **Polar solutes** tend to dissolve in polar solvents. **Nonpolar solutes** tend to dissolve in nonpolar solvents. -Oil is nonpolar while water is polar. They are immiscible. **Mathematical methods of expressing the concentration of solution.** The concentration of solutions refers to the amount or volume of solute dissolved in a specific amount of solvent. There are many ways of computing accurately the concentration of solutions. Some of these methods are : percent by mass, percent by volume, mole fraction, molarity , and molality. **\ ** **Phase Diagram of Water** Water is undergoing phase changes. Based on the illustration when ice cube melts and becomes water it evaporates and becomes water vapor. Phase changes require either the addition of heat energy (melting, evaporation, and sublimation) or subtraction of heat energy (condensation and freezing). Changing the amount of heat energy usually causes a temperature change. When a substance changes from one state, or phase, of matter to another it has undergone a change of state, or a change of phase. These changes of phase always occur with a change of heat. A change of state occurs when matter is converted from one physical state to another. For example, when water is heated, it changes from a liquid to a gas---when cooled water will eventually freeze into a solid which is commonly called ice. A change of state is usually accompanied by a change in temperature and/or pressure. **Phase change** There are six ways a substance can change between these three phases; melting, freezing, evaporating, condensing, sublimination, and deposition. These processes are reversible and each transfers between phases differently: **Melting:** The transition from the solid to the liquid phase **Freezing**: The transition from the liquid phase to the solid phase Evaporating: The transition from the liquid phase to the gas phase **Condensing:**The transition from the gas phase to the liquid phase **Sublimination**: The transition from the solid phase to the gas and phase **Deposition**: The transition from the gas phase to the solid phase **Phase diagram** Phase diagram is a graphical representation of the physical states of a substance under different conditions of temperature and pressure. A typical phase diagram has pressure on the y- axis and temperature on the x- axis. As the line crosses or curves on the phase diagram , a phase change occurs. The phase diagrams indicate the physical states that exist under specific conditions of pressure and temperature and also provide the pressure dependence of the phase-transition temperatures (melting points, sublimation points, boiling points).  Open photo  Open photo Point A is the point of intersection of all lines, where the solid/ liquid, liquid/gas, and solid/ gas lines intersect, which is called the triple point. The only combination of temperature and pressure at which all three phases (solid, liquid , and gas ) are in equilibrium. The pressure lower than the triple point cannot exist as a liquid, regardless of the temperature. The phase diagram shown is for a single pure substance in a closed system , not for a liquid in an open container in contact with air at 1 atm pressure.  The phase diagram of water is divided into three regions each of which represents a pure phase. The line separating two regions indicates conditions under which two phases can exist in equilibrium. For example, the curve between the liquid and the water vapor phase shows the variation of vapor pressure with temperature. The other two curves similarly indicate conditions for equilibrium between ice and liquid water water vapor. The point at which all three curves meet is called a triple point, the only condition under which all three phases can be in equilibrium with one another. The phase diagram of water shows that the triple point occurs at 0.0098 ᵒC and 4.58 mmHg ( 0.01 ᵒC and 0.0006 atm). The normal melting (**freezing) point is 0ᵒ C**. The normal **boiling point is 100 ᵒC**. A **critical point is 374 ᵒC and 218 atm.** **Kinetic Molecular Model of Liquids and Solids Solids** Solid has an ordered and dense molecule but has fixed shape and volume. It is slightly compressible because of the presence of strong intermolecular forces. The diffusion is extremely low because its particles are held together by a strong force of attraction called bonds. The particles vibrate to and fro near their fixed locations and never come in contact with any particles other than their immediate neighbor. Solids have definite shape, definite volume, definite melting point, high density, incompressible, and has low rate of diffusion. Examples are iron, concrete, marble, pebbles, wood, glass and others **Liquids** Liquids have disordered, slightly dense but has definite volume. The attractive forces in liquids are not strong enough to keep the molecules from moving past from one another this is why liquids can be poured, and it can fill any shape. They are also called fluid. Fluids have loosely packed molecules. They vibrate vigorously that the attraction cannot hold them in a fixed position. Liquids are almost incompressible, have fixed volume but no fixed shape, it readily flows, and have boiling point above room temperature. Examples: Water, alcohol, oil **STRUCTURE OF CRYSTALLINE AND AMORPHOUS SOLIDS** \* Solids can be categorized into two groups: the crystalline solids and the amorphous solids. The differences in properties of these two groups of solids arise from the presence or absence of long range order of arrangement of the particles in the solid. **1. Arrangement of particles** **Crystalline Solid**-they can form a regular repeating three-dimensional structure called a crystal lattice, thus producing a crystalline solid. **Amorphous Solid**- they can aggregate with no particular long range order, and form an amorphous solid (from the Greek ámorphos, meaning \"shapeless\"). Open photo \- **[Crystalline solids]** are arranged in fixed geometric patterns or lattices. Examples of crystalline solids are ice and sodium chloride (NaCl), copper sulfate (CuSO4), diamond, graphite, and sugar (C12H22011).
The ordered arrangement of their units maximizes the space they occupy and are essentially incompressible. **Examples of crystalline solids:** -sodium chloride -copper sulfate -diamonds -graphite -sugar -blue sapphire and calcite \- **[Amorphous solids]** have a random orientation of particles. Examples of amorphous solids are glass, plastic, coal, and rubber. They are considered supercooled liquids where molecules are arranged in a random manner similar to the liquid state. **Examples of amorphous solids:** -glass -plastics -coal -rubber More than 90% of naturally occurring and artificially prepared solids are crystalline. Minerals, sand, clay, limestone, metals, alloys, carbon (diamond and graphite), salts (e.g. NaCl and MgSO4), all have crystalline structures. They have structures formed by repeating three dimensional patterns of atoms, ions, or molecules. The repetition of structural units of the substance over long atomic distances is referred to as long-range order. Amorphous solids (e.g. glass), like liquids, do not have long range order, but may have a limited, localized order in their structures.   **Behavior of solids when heated** The presence or absence of long-range order in the structure of solids results in a difference in the behavior of the solid when heated. The structures of crystalline solids are built from repeating units called crystal lattice. The surroundings of particles in the structure are and the attractive forces experienced by the particles are of similar types and strength. These attractive forces are broken by thuds a a spect emperature thus, emailing solids be thine temperature, physical properties of the crystalline solids change sharply. \- Amorphous solids soften gradually when they are heated. They tend to melt over a **wide range** **of temperature**. This behavior is a result of the variation in the arrangement of particles in their structures, causing some parts of the solid to melt ahead of other parts.