Chemistry Chapters 5, 11, and 12
Chemistry Chapters 5, 11, and 12 CHM 160 001
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This 6 page Bundle was uploaded by Devin Mart on Tuesday March 29, 2016. The Bundle belongs to CHM 160 001 at Missouri State University taught by Dr. Richter in Spring 2016. Since its upload, it has received 8 views. For similar materials see General Chemistry 1 in Chemistry at Missouri State University.
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Date Created: 03/29/16
Chapter 5: Gases ● Pressure (the force per unit area) is a characteristic property of gases. ○ Pressure is the force per unit area exerted by gas molecules colliding with surfaces. ● Gaseous molecules flow from the region of higher pressure to the region of lower pressure. ○ In shallow wells, atmospheric pressure is used to push water up the pipes. ● All mater is composed of atoms. ○ Atoms move about in perpetual motion and are never at rest. ● Gaseous atoms move at several hundred miles per hour. ○ Pressure is the result of molecular collisions. ● Temperature is the average kinetic energy of atoms, molecules or ions that compose a substance. ○ As heat is added to the system atoms, molecules and ions begin to move (vibrate, rotate and translate). ■ The warmer the substance, the more the atoms move. ○ Movement of water molecules becomes so rapid that the water can’t stay as ice. ■ As the average kinetic energy continues to increase the water molecules will eventually gain enough energy to break free from the liquid. ● Pressure increases when a larger number of gas particles are present (if temperature is kept constant). ○ A pressure imbalance occurs when the external pressure does not equal the internal pressure. ● As the number of moles of CO₂ (g) increases, the force exerted on the walls of the container increases while the area remains “constant”. ● Evangelista Torricelli (1646) proposed that the atmosphere has a pressure. ○ At higher altitudes the pressure decreases. ○ Confirmed Torricelli’s hypothesis that there is an ‘atmospheric pressure’ or ‘blanket of air’. ● The units of pressure are (usually) atm, mmHg, Torr or Pascal. ● With no internal pressure to compensate, the external pressure crushes the can. ● Measuring pressure in a gas tank before and after can be related directly to the amount of gas used. ● Gases have four basic physical properties that are interrelated. ○ Pressure (P), Temperature (T), Amount in moles (n), and Volume (V). ● Pressure is inversely related to volume, keeping temperature and number of moles constant. ○ Pressure times volume equals a constant. ● Pressure is proportional to temperature, keeping n and V constant. ● As the average kinetic energy increases the atoms hit the walls with more force. ○ The pressure of a gas approaches zero at 270℃. ○ Kelvin must be used in all calculations. ● Volume is proportional to temperature, keeping n and P constant. ○ As the temperature increases the pressure exerted on the walls increases, and the volume must increase until equilibrium is established. ● Ratio of volumes of gases consumed or produced is equal to the ration of simple whole numbers (Temperature and Pressure remain constant). ● Avogadro assumed that equal volumes of different gases collected under similar conditions contain the same number of particles. ○ Equal volumes of different gases collected under similar conditions contain the same number of particles. ○ As the amount of gas increases (moles), the volume increases. ● The ideal gas equal is a general relationship to explain and predict the properties of gases. ○ Boyle’s law → P ∝ 1/V ○ Amonton’s law → P ∝ T ○ Charles's law → V ∝ T ○ Avogadro’s law → V ∝ n ○ Avogadro’s law → P ∝ n ● R is the ideal gas “constant” ○ R = 0.08206 at STP ● Postulates of Kinetic Theory. ○ 1. Gases are composed of molecules whose size can be ignored when compared to the average distance between them. ■ Molecules move randomly in straight lines in all directions at various speeds. ○ 2. The average kinetic energy of a molecule is proportional to the absolute temperature (in K). ○ 3. The forces of attraction or repulsion between 2 molecules (intermolecular forces) in a gas are very weak or negligible, except when they collide. ■ When molecules collide with one another the collisions are elastic. ● Kinetic molecular theory is used to explain the gas laws. ● The kinetic theory of gases assumes there are no attractions between gas molecules. ○ At low temperatures and high pressures the ideal gas laws fail. ● Real pressure is always lower than ideal pressure. Chapter 11: Liquids, Solids, and Intermolecular Forces ● Intermolecular forces are the attractions between all chemical molecules and atoms. ○ DNA is held together by intermolecular forces. ● Liquids have properties between the extremes of solids and gases. ● One phase of matter can be transformed to another by changing the temperature or pressure or both. ○ Most substances that are liquids at room temperature and pressure are composed of individual molecules. ■ Generally liquids at room temperature and pressure are covalent systems. ● Movement of water molecules becomes so rapid that the water can’t stay as ice. ● The attractive forces between molecules are called intermolecular forces. ○ Intermolecular forces are often called Van Der Waals forces. ● Intramolecular bonds are much stronger than intermolecular bonds. ○ Increasing the average kinetic energy of molecules disrupts intermolecular forces. ● The strength of the intermolecular forces between molecules or atoms determines the phase solid, liquid or gas of the substance at a given temperature. ● Dispersion forces (London forces) exist in all molecules ○ The larger the molar mass, the stronger the Dispersion forces and the higher the boiling point. ● Dipoledipole interactions occur between polar molecules. ● Hydrogen bonding involves a hydrogen ‘bridging’ two very electronegative atoms (Nitrogen, Oxygen or Fluorine). ● The iondipole force occurs when an ionic compound is mixed with a polar compound. ○ 1. Dispersions → all molecules and atoms ○ 2. Dipoledipole → Polar molecules ○ 3. Hydrogen bonding → Molecules containing H bonded to F, O, or N ○ 4. Iondipole → Mixtures of ionic compounds and polar molecules ● The stronger the intermolecular forces, the higher the boiling point. ● The force that controls the shape of a liquid is called the surface tension. ○ The stronger the force of attraction, the higher the surface tension. ○ Water drops are spherical due to surface tension. ● The stronger the intermolecular force, the higher the viscosity of a liquid will be. ○ Viscosity the resistance of a liquid to flow. ● Vaporization occurs when thermal energy overcomes intermolecular forces and produce a phase change from liquid to gas. ○ A liquid can form a gas by either boiling or evaporation ○ At a given temperature, some of the particles in a liquid will have enough energy to form a gas. ● Vapor pressure is the pressure of a gas in dynamic equilibrium with its liquid in a closed container ○ Vapor pressure is only a constant at a specific temperature. ● At equilibrium, the number of moles in each state remains constant. ● The strength of the intermolecular forces can be estimated by measuring the amount of heat required to transform liquid to gas. ○ The enthalpy of vaporization can be calculated from tabulated data. ● Sublimation is the transformation of a solid directly to a gas. ● The strength of the intermolecular forces can be also estimated by measuring the amount of heat required to transform solids to liquids. ● Heat can sometimes enter a system without changing its temperature. ● At 0℃ (melting point), all the heat is used to melt the ice (i.e., break some intermolecular forces). ○ At 100℃ (boiling point), all the heat is used to vaporize the liquid water (i.e., break some intermolecular forces). ● The normal boiling point is the temperature at which the liquid boils at 1 atm. ○ A liquid boils when the pressure of the vapor escaping from the liquid is equal to the pressure exerted on the liquid by its surroundings. ○ Atmospheric pressure effects boiling point (the lower the atmospheric pressure the lower the boiling point and pressure needed). ● Phase diagrams describe the state or phase of a substance at different temperatures and pressures. ○ At the critical point, the meniscus separating the liquid and gas disappears and it becomes supercritical neither a liquid nor a gas. ● Water is the most common and important liquid on earth, and among liquids, water is unique. ○ Water has a higher heat capacity than any other common liquid. ■ Heat capacity: the ability to absorb heat without undergoing a change in temperature. ○ The amount of heat required to raise the temperature depends on the strength and amount of the intermolecular forces. ● Specific heat is the heat needed to raise the temperature of 1 gram of a substance by 1℃. ● Molar heat capacity is the heat needed to raise the temperature of 1 mol of a substance by 1℃. ● Absorption of heat by equal moles of substances will raise the temperature more in the substance with the lower heat capacity. ○ The ability of water to absorb lots of heat energy limits the range of temperatures in lakes and rivers (summer vs. winter). ○ The water that surround San Francisco absorbs a majority of the heat without a significant rise in temperature. ● Oceans act as thermal reservoirs. ○ This moderates swings in temperature from winter to summer on a global scale. ● Water expands when it freezes. ○ Solid water is less dense than liquid water because of the air pockets between each molecule that form when it is frozen. ● The melting point of water is 100℃ higher than expected while the boiling point of water is 200℃ higher than expected. ● Water is excellent at dissolving ionic solids. Chapter 12: Solutions ● In osmosis, water flows from a less concentrated solution to a more concentrated solution. ○ As it flows through the stomach and intestine, seawater draws water out of bodily tissues, promoting dehydration. ○ These are sometimes called “thirsty solutions” since the solution with the high solute concentration draws water away from the solution with a lower solute concentration. ● Solutions are homogeneous mixtures of two or more substances. ○ Solutions contain two components, a solute and a solvent. ● Unless it is highly unfavorable energetically, substances tend to combine to form uniform mixtures. ○ An aqueous solution contains water as the solvent. ● Intermolecular forces may promote the formation of a solution or prevent it, depending on the nature of the forces in the solute and solvent. ○ When a solute and solvent are mixed, there is a competition between the solventsolute interactions and the solutesolute interactions. ■ If the solventsolute interactions are stronger, the material will dissolve. ■ If the solutesolute or solventsolvent interactions are stronger, the material will not dissolve. ● In general, like dissolves like polar solvents dissolve polar solutes, nonpolar solvents dissolve nonpolar solutes. ○ Polar solutes interact with the partial charges in water and dissolve. ○ Nonpolar solutes do not interact with the partial charges in water and don’t dissolve. ■ Nonpolar solutes will dissolve in nonpolar solvents. ● Energy changes occur when a solution forms. ○ Exothermic reactions will produce energy (negative) and give off heat, while endothermic reactions will take energy (positive) and absorb head. ● To understand the energy changes that occur when a solution forms envision the process occurring in three steps: ○ 1. Separate the solute into its constituent parts. ■ This step is always endothermic because energy is required to overcome the bonding forces that hold particles together. ○ 2. Separate the solvent particles from one another. ■ Again, this step is always endothermic for the same reason as the first step. ○ 3. Mix the solute particles with the solvent particles. ■ This step is always exothermic because energy is released as the solute particles interact (through intermolecular forces) with the solvent particles. ● The overall enthalpy change upon solution formation, the enthalpy of solution is the sum of each step in the reaction. ○ In general, the more negative the enthalpy of solution is the greater the likelihood that a solution will form. ● The enthalpy of hydration is the enthalpy change that occurs when 1 mol of gaseous solute ions dissolve in water. ○ Since the iondipole interactions that form between the ions and water are much greater than the hydrogen bonds in water the enthalpy of hydration is large and negative. ● The dissolution of a solute in a solvent is an equilibrium process. ○ A saturated solution occurs when the dissolved solute is in dynamic equilibrium with the undissolved solid solute. ■ If you add more solute to a saturated solution it will not dissolve, while if you add more solute to an unsaturated solution it will dissolve. ● The solubility of most solids in water increases with increasing temperature. ● The solubility of most gases in water decreases with increasing temperature. ○ The solubility of gases in water increases with increasing pressure. ● The amount of solute in a solution is variable. ○ Solution concentration is usually measured in units of molarity.
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