Chapter 11: Gases

Matter is paticulate and ts behavior can be understood in terms of particles

7.1 Describe measuring gas pressures using barometers and manometers. Convert between pressure units. (11.1 and 11.2) Explain and measure gas pressure. Relate pressure units. 

• The behavior of gases can be explained and predicted by a model called the kinetic molecular theory. The core of this model is that gases are composed of particles in constant motion.
• 1.) Observations on the properties of gases
• 2.) Observations led to scientific laws (that summarized the observed behavior)
• 3.) The observations and laws led to the development of the kinetic molecular theory-the model the nature of a gas. 
• Observations---> Scientific Laws ---> Kinetic Molecular Theory
• Gases:
• Have no definite shape and no definite volume for gases will expand to fill any container. 
• Characteristics of Gases:
• Compressibility: Change in volume of a sample resulting from a pressure change acting on the sample. 
• Solids and gases are not significantly compressible but gases can be compressed under sufficient pressure. 
• Gas molecules are widely spread and therefore have low densities. 
• Pressure: Particles in a gas collide with each and with the surfaces around them. Each collisions exerts only a small force, but when the forces of many particles are summed, they quickly add up. 
• The result of the constant collision between the atoms and molecules in a gas and the surfaces around them is pressure. 
• The pressure that a gas sample exerts is the force that results from the collisions of gas particles divides by the area of the surface with which they collide.
• Pressure= (Force/ Area) 
• P= (F/A)
• Pressure is the force (F) that acts on a given area (A)
• Atmospheric Pressure: The pressure that is exerted upon the earth's surface by the atmosphere. 
• Variation in pressure in Earth's atmosphere create wind, and changes in pressure help us to predict weather. 
• The number of gas particles in a given volume decreases with increasing altitude (See Below)*
• Pressure decreases with increasing altitude
• Pressure Observations:
• The total pressure exerted by a gas depends on several factors such as:
• Concentration: Number of gas particles in a given volume. The lower the concentration (fewer gas particles), the lower the pressure. Likewise, a high density of gas partiles results in high pressure. 
• Altitude *: Because the number of gas particles in a given volume decreases with increasing altitude, pressure decreases with increasing altitude. 
• Ex. In the air, airplane cabins have to be artificially pressurized due to lack of oxygen at high altitudes. 
• Ex. When you ascend in the plane or hike up a mountain, the external pressure (the pressure that surrounds you) drops, while the pressure within the ear cavity (internal pressure) stays the same. This imbalance, creates the ear popping effect. 
• Volume of the Container: A low density of gas particles results in low pressure.
• Average Speed of the Gas Particles: A high density of gas particles results in high pressure. 
• SI Units of  Pressure:
• Force: 
• (N) newton
• 1 N= 1 kg x m/s2
• Pressure:
• (Pa) Pascal
• 1 Pa=1 N/m2
• Non SI Units of Pressure that are Commonly Used:
• atmosphere (atm)
• millimeter of mercury (mm Hg)
• torr
• Note that 1 mmHg=1 torr

Manometer: The instrument used to measure the pressure of a gas sample in the laboratory. They are U shaped tubes partially filled with a liquid that are connected to the gas sample on one side and open to the air on the other.

7.2 Apply the ideal gas law to relate and calculate values for pressure, volume, temperature, and amount of a gas (11.3-11.5) Apply the ideal gas law to relate and calculate values for pressure, temperature, and amount of gas. 

Pressure/Gas Laws:

• Measurable properties of gases include: 
• Pressure (P) in atm
• Volume (V) in L
• Temperature (T) in Kelvin (C+273.15)
• Amount in moles (n)
• The properties are interrelated-when one changes, it affects the others. 
• The Gas Laws: mathematical relationships that describe the quantitative relationships and behavior of gases as they are mixed, or subjected to pressure or temperature changes. 

These simple gas laws describe the relationship between pairs of these properties.

• Boyle's Law: describes how volume (V) varies with pressure (P) at constant temperature and amount of gas. 
• Boyle and Hook used a J shaped tube to measure the volume of a gas sample of gas at different pressures. They trapped a sample of air in the J tube and added mercury to increase the pressure on the gas. 
• Boyle's law states that volume and pressure have an inverse relationships where an increase in one causes a decrease in the other. 
• If the volume is increased, the pressure will decrease
• We can use Boyle's ;aw to calculate the volume of a gas following a pressure change or the pressure of a gas following a volume change. 
• As long as the temperature change and the amount of gas remain constant. 

p1v1=v1v2; Where p1 and v1 are the initial pressure and volume of the gas and p2 and v1 are the final volume and pressure. 

Charles's Law: The volume of a fixed quantity of gas maintained at constant pressure is directly proportional to its absolute temperature * (T in Kelvin) assuming that the pressure and the number of moles are constant. 

• AJ Charles was interested in gases and was among the first people to ascend in a hydrogen filled balloon. 
• When the temperature of a gas sample increases, the gas particles move faster; collisions with the walls are more frequent 
• The volume of a gas increases with increasing temperature. 
• Volume and temperature are linearly related. If two variables are linearly related, plotting one against the other produces a straight line. 
• Absolute zero * refers to the temperature at 0 K-colder temperatures do no exist. Anything less would have a negative volume which is physically impossible. 
• If a balloon is moved from an ice water bath to a boiling water bath, its volume expands as the gas particles within the balloon move faster (due to the increased temperature) and collectively occupy more space. 
• Ice water (low kinetic energy)----> Boiling water (high kinetic energy)

(v1/t1)=v2/t2 at constant pressure (P) and number of moles (n)

The pressure exerted on a sample of a fixed amount of gas is doubled at constant temperature, and then the temperature of the gas in kelvins is doubled at constant pressure. What is the final volume of the gas?

(a) The final volume of the gas is twice the initial volume.

(b) The final volume of the gas is four times the initial volume.

(c) The final volume of the gas is one-half the initial volume.

(d) The final volume of the gas is one-fourth the initial volume.

(e) The final volume of the gas is the same as the initial volume.

Avogradro's Law: The volume (v) of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas (n)

• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fills more space. 
• More gas molecules=larger volume
• We can use Avogrado's law to calculate the volume of a gas following a change in the amount of the gas as long as the pressure and temperature are constant. 
• The volume of a gas sample increases linearly with the number of moles in the gas sample. 
• When the amount of gas in a sample increases at constant temperature and pressure, its volume increases in direct proportion because the greater number of gas particles fill more space. 
• Equal volumes of gases contain equal numbers of molecules (gas doesn't matter)

• R (gas constant) is the proportionality constant in the ideal gas equation
• R relates pressure, volume, temperature, and the number of moles of gas in the ideal gas equation.
• Molar Volume: Volume of one mole of a gas determines at STP (standard temperature and pressure)
• An ideal gas is a gas whose volume, pressure, and temperature may be described by the ideal gas equation

7.3 Apply Dalton's Law of Partial Pressure to calculate the pressure of combined gases and to calculate the partial pressures of gases in mixtures. (11.6)

7.4 Describe gases in terms of the Kinetic Molecular Theory (11.7, 11.8)

• Kinetic Model Theory: The simplest model for the behavior of gases. In this theory, a gas is modeled as a collection of particles (either molecules or atoms, depending on the gas) in constant motion. The basic postulates (or assumption) of kinetic molecular theory are listed below:
• 1.) The size of a particle is negligibly small. KMT assumes that the particles themselves occupy no volume, even though they have mass. 
• Smaller molecules move faster then larger molecules. 
• 2.) The average kinetic energy of a particle is proportional to the temperature in kelvins. The motions of atoms ad molecules in a gas is due to thermal energy, which distributes itself among the particles in the gas.  
• The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature. Nothing else. 
• The temperature of the gas increases, the average speed of the particle increases. 
• However, not all the gas particles are moving at the same speed. 

• 3.) The collision of one particle with another (or with the walls of its container) is completely elastic. This means that when two particles collide, they may exchange energy, but there is no overall loss of energy. 

7.5 Differentiate Effusion and Diffusion (11.9)

Diffusion: The process by which gas molecules spread out in response to a concentration gradient, and even though the particles undergo many collisions, the root square velocity still influence the rate of diffusion.

Effusion: The process by which a gas escapes from a container into a vacuum through a small hole.

7.6 Distinguish between Ideal and Real Gases (11.11)

• Ideal gases are hypothetical-there is no gas that will exactly follow the ideal gas behavior. 
• However, may gases will behave very closely to ideal gases under certain conditions. 
• According to KMT, ideal gas laws assume:
• No attractions between gas molecules
• gas molecules do no take up space
• At low temperatures and high pressures these assumptions are not valid. 
• Real gases will behave more like ideal gases under conditions of:
• Low pressure (at higher pressures, the volume of the gas molecules become significant
• High temperatures (at lower temperatures, the inter molecular attractions can lower the expected pressures)

Chapter 9: Introduction to Solutions and Aqueous Reactions

9.3 Solution Stoichiometry (307)

• In aqueous reactions, quantities of reactants and products are often specified in terms of volumes and concentration. We can use the volume and concentration of a reactant of a reactant or product to calculate its amount in moles. 
• Volume A==> Amount A (in moles) ==> Amount B (in moles) ==> Volume B

9.4 Types of Aqueous Solutions and Solubility

• When a solid is put into a liquid solvent, the attractive forces that hold the solid together (the solute-solute interactions) compete with the attractive forces between the solvent molecules and the particles that compose the solid (solvent-solute interactions)
• Electrolytes and Non electrolyte Solutions
• A salt solution (Electrolyte) conducts electricity while a sugar solution (Non electrolyte) does not. 
• Salt (ionic) 
• Electrolyte: Substances that dissolve in water to form solutions that conduct electricity.
• Strong Electrolyte: Substances that completely dissociate into ions when they dissolve in water. Strong acids are strong electrolytes. 
• Weak Electrolyte: Does not completely ionize in water. Weak acids are weak electrolytes that conduct electricity weakly. 
• Sugar (molecular)
• Most molecular compounds-with the exception of acids- dissolve in water as intact molecules. Sugar dissolves because the attraction  between sugar molecules and water molecules, overcome the attraction of sugar molecules to each other. 
• Compounds such as sugar that do not dissociate into ions when dissolved in water are nonelectrolytes that do not conduct electricity. 
• Acids (Molecular)
• Ionize to form H+ ions when they dissolve in water. 
• Ex. HCI ionizes into H+ and Cl- when it dissolves in water.  HCI (aq) ==> H+(aq) + Cl- (aq)
• Strong acids completely ionizes in solution. 
• In general, a compound is soluble if it dissolves in water and insoluble if it does not.

9.5 Precipitation Reactions

• Precipitation Reaction: a reaction in which a solid forms upon the mixing of two solutions. 
• However, they do not always occur when two aq solutions are mixed. 
• The key to predicting reactions is understanding that only insoluble compounds form precipitate. Two solutions containing soluble compounds combine and an insoluble compound precipitate. 
• Practice Problems (317)

9.6 Representing Aqueous Reactions: Molecular, Ionic, and Complete Ionic Equations

• Molecular Equation: A chemical equation showing the complete, neutral formulas for every compound in a reaction. 
• Complete Ionic Equation: Chemical equation showing all of the species as they are actually present in solution
• Net Ionic Equation: Equation showing only the species that actually change during the reaction.

9.7 Acid-Base Reactions

9.9 Oxidation-Reduction Reactions

Chapter 8: Chemical Reactions and Chemical Quantities

8.3 Writing and Balancing Chemical Equations

• Chemical changes occur via chemical reactions. 
• A combustion reaction is a particular type of chemical reaction in which a substance combines with oxygen to form one or more oxygen containing compounds. Combustion also emit heat. 
• We represent a chemical reaction with a chemical equation. The substances on the left are reactants and the ones on the right are products. 
• When a chemical equation is not balanced, it violates the law of conservation of mass because an atom is not formed out of nothing. 
• To correct an unbalanced equation-or write an equation that closely represents what actually happens-we must balance it by changing the coefficients to ensure that the number of each type of atom on the left is equal to the right side. New atoms do not form during a reaction, nor do atoms vanish-matter must be conserved. 
• What quantity or quantities must always be the same on both sides of a chemical equation?
• the number of atoms of each kind. 

8.4 Mole to mole and mass to mass conversions

• The coefficients in a chemical equation specify the relative amounts in moles of each substance involved in the reactions. 
• The numerical relationships between chemical amounts in a balanced chemical equation are called stoichiometry, which allows the prediction of the amount of products that will form in a chemical reaction based on the amount of reactants that react. It also allows is to determine the amount of reactants necessary to form a given amount of product. 
• This is essential in knowing the quantity of chemical reactants to obtain products in the desired quantities. 
• Think of it as a recipe. The amount of ingredients (cheese, crust, sauce) make a certain amount of pizzas (products). 

2 C8H18 (l) + 25 O2 (g) ====> 16 CO2 (g) + 18 H2O (g)

• Making Molecules: Mole to Mole Conversions
• In a balance chemical equation, we have a "recipe" for how reactants combine to form products. From our balanced equation for the combustion of octane for example is the following stoichiometric ratio:
• 2 mol C8H18: 16 mol CO2
• We can use this ratio to determine how many moles of CO2 from when a given number of moles of C8H18 burns. 
• Suppose we burn 22.0 mol C8H18, how many moles of CO2 form?
• 22.0 mol C8H18 X (16 mol CO2/2 mol C8H18)= 176 mol CO2
• The combustion of 22 moles of C8H18 adds 176 moles of CO2 to the atmosphere. 
• Making Molecules: Mass to Mass Conversions
• This calculation is similar to the Mole to Mole conversions exept that we are now given mass of octane instead of the amount of octane in moles, Consequently we must first convert mass (in grams) to the amount (in moles). (pg. 280)
• Mass A ==> Amount A (in moles) ==> Amount B (in moles) ==> Mass B
• Practice Problems 8.4 and 8.5 (pg. 281 and 282)

Video Overview

Mole to Mole Conversions

Mass to Mass Conversions

Mass to Mass Conversions Part II

8.5 Limiting Reactant, Theoretical Yield, and Percent Yield

• The three most important concepts in reaction stoichiometry: limiting reactant, theoretical yield, and percent yield.
• Limiting Reactant (Reagent):The substance that limits the amount of products that can be made
• Excess Reactant: is any reactant that occurs in a quatity greater than is required to completely react with the limiting reactant.
• Theoretical Yield: The amount of product that can (theoretically) be made based on limiting reactants
• Actual Yield: The amount of product that is (actually) made. Always equal or less than theoretical yield. 
• Percent Yield: Ratio percentage. (Actual Yield/Theoretical Yield) X 100