Chem weeks 4-6
Popular in Chemistry: Principles and Applications
Popular in Chemistry
This 7 page Class Notes was uploaded by Catherine Carter on Friday October 14, 2016. The Class Notes belongs to CHE 105 at University at Buffalo taught by Melvyn Churchill in Summer 2015. Since its upload, it has received 2 views. For similar materials see Chemistry: Principles and Applications in Chemistry at University at Buffalo.
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Date Created: 10/14/16
Thermochemistry – study of heat absorbed or emitted from the system during a chemical reaction o Thermodynamics – the study of energy and heat transfer o Energy – the ability to do work or transfer heat work – energy used to cause an object that has mass to move w=Fd (w is work, F is force, d is distance the force is exerted) heat – energy used to cause the temperature of an object to rise represented by q heat always and only flows from warmer to cooler o Potential energy – the energy an object possesses because of its position or chemical composition o Kinetic energy – energy an object possesses because of its motion o electrostatic potential energy – the interaction between two charged particles k is a constant, Q1 is the charge of particle 1, Q2 is the chage of particle 2, and d is the distance in meters o Radiant energy – electromagnetic radiation (E=hv) o chemical energy – potential energy stored in the bonds between atoms o thermal energy – “heat” caused by the motion of molecules o the SI unit of energy is the joule (J) 1 calorie = 4.184 J 1 “Calorie”(like in food science) = 1kcal (kilocalorie) The System and Surroundings o the system is what we want to study o the surroundings are everything else o system + surroundings = universe First Law of Thermodynamics Energy is neither created nor destroyed o the total energy of the universe is constant (if the system loses energy then the surroundings are gaining the energy lost) o Internal Energy (E) – sum of all the kinetic and potential energies of all components of the system you can’t measure E , but you can measure any change in E (ΔE) if ΔE > 0 then the final energy is greater than the initial and the system has absorbed energy from the surroundings which is called endergonic if ΔE < 0 then the final energy is less than the initial energy and the system released energy to the surroundings which is called exergonic The changes always refer to the SYSTEM internal energy, E, is a “State Function” meaning it depends on the present state of the system and the path that it arrived at that state does not matter q and w are NOT state functions o when a process happens in an open container the only work done is between the volume of a gas pushing on the surroundings or the surroundings pushing on it Enthalpy (H) – measuring enthalpy lets us account for the heat flow during the process H = E + PV o when P is a constant pressure o at a constant pressure, the change in enthalpy is the heat gained (+) or lost () by the system because: o Enthalpy is a “state function” o the enthalpy of a system can be found by: o the enthalpy for a reaction can be found by: o endothermic reaction is when ΔH is positive (heat is taken in by the system) o exothermic is when ΔH is negative (heat is given out by the system) o ΔH is also called the enthalpy of reaction or the heat of reaction o enthalpy is an extensive property and it depends on the number of moles present which we get from the coefficients of the balanced equation o for the reverse it is equal in magnitude but with the opposite sign o it also depends on the states of the products and the reactants Calorimetry is the measurement of heat flow, we use this since we cannot know the exact enthalpy of the reactants and the products the heat capacity of an object is the amount of energy needed to raise the temperature of the object by exactly 1K (1C) the specific heat capacity is the amount of energy needed to raise the temperature of exactly 1g of a substance by exactly 1K molar heat capacity is the amount of energy needed to raise the temperature of exactly 1mol of a substance by exactly 1K o s is specific heat constant pressure calorimetry o reactions carried out in aqueous solutions in a simple calorimeter can indirectly measure the heat change for the system through the temperature change of the water o if the water gained heat, then the system lost heat so q is negative and vice versa o the experiment is insulated from the outside world so that the heat lost by the system is gained by the water (the surroundings) bomb calorimetry o the volume in bomb calorimetry is constant so you are measuring the change in internal energy ΔE not ΔH, for most reactions the difference is very small o in order to calculate ΔE you need to determine the heat capacity of the C cal qreaction calorimeter ( ) burn a substance of known and measure the change in the temperature use this equation to find the heat capacity of a calorimeter o for the unknown substance, measure the temperature difference and then use this equation (basically the same one using the heat capacity of the calorimeter that you found before): Hess’s Law o o this works because enthalpy is a state function so the pathway doesn’t matter just the initial and final states (reactants and products) o Coefficients matter so if the reactant/the products have it INCLUDE IT o is basically o the enthalpy of formation ( ) is the enthalpy change for the reaction where a compound is made from its elements in their elemental forms for an element in its standard state o n and m are the stoichiometric coefficients o standard enthalpies of formation are measured under standard conditions (25C and 1.00 atm pressure) Gases Substances that are gaseous under ambient conditions (1atm, 25 Celsius) include 11 elements and some smaller molecular compounds o monatomic gases (all 6 noble gases), 5 diatomic gases (NClHOF), and one triatomic gas (O3) o lots of things containing hydrogen, oxygen, and fluoride gases expand to fill their containers, are highly compressible, have extremely low densities (at STP (1atm, 0 Celsius)), don’t have fixed volumes or shapes the volume a gas occupies is equal to the volume of the container it is in less commonly Pressure is the amount of force applied to an area, atmospheric pressure is the weight of the air per unit of area o measure in mm Hg or torr o o The Pascal is a very small unit of pressure (SI Unit) o The bar is useful in meteorology o the manometer is used to measure the difference between the atmospheric pressure and the pressure of a gas in a vessel o standard pressure is the normal atmospheric pressure at sea level Boyle’s Law o the volume of a fixed quantity of a gas at a constant temperature is inversely proportional to the pressure o Charle’s Law o the volume of a fixed amount of gas at a constant pressure is directly proportional to its temperature o Avogadro’s Law o the volume of a gas at a constant temperature and pressure is directly proportional to the number of moles of the gas o GayLusac’s Law (Law of Combining Volumes) – gases combine in simple ratios by volume Avogadro’s hypothesis – equal volumes of gas at the same temperature and pressure contain the same number of molecules 1 mol of an ideal gas occupies 22.414 L at STP when you combine all of the equations you get the equation R is a gas constant and has a different value depending on the units used (MAKE SURE THEY MATCH) o o Temperature is ALWAYS kelvin The density of gas is the moles over the volume (moles/L) to get grams you’d have to multiply the other side by the molar mass M to find the molar mass from the density: Dalton’s law of partial pressure says that the total pressure of a mixture of gases equals the sum of the pressures that each would exert if it was present alone so: o o When you collect gas over water, you have to subtract the vapor pressure of the water from the total pressure Kinetic Molecular Theory (KMT) o Gases consist of large numbers of molecules that are in continuous, random motion o The pressure of a gas is due to the molecules colliding with the walls of the container o The combined volume of the molecules of the gas is negligible relative to the volume of the container of the gas o All attractive and repulsive forces between gas molecules are negligible o When molecules collide, the collisions are elastic so that there is no net loss of momentum of energy o The average kinetic energy of the molecules does not change as long as the temperature of the gas remains constant The average kinetic energy gas molecules is proportional to the absolute temperature o o There’s a bell curve for how many molecules move at what speed at a certain temperature o o Effusion is gas particles going through a small hole, diffusion is gas particles spreading out within their container o Through the KMT you can relate the root mean square velocity to the mass of a particle n is the number of velocities this is how you calculate rms value, u is the rms velocity o Average and rms values are not the same o You can compare the rates of effusion of two different gases through a pinhole by their rms velocities: this is called graham’s law Real gases o Gases behave ideally only under high temperature and low pressure o Assumptions made under KMT do not apply for real gases o At high pressures, the volume of the gas molecules is no longer negligible since they have infinite volume o The molecules do attract and repel each other so the partial pressures are not independent of each other o Fixed by the van der Waals’ equation. B is a constant and must be determined experimentally for each gas, it is the effective molar volume. You then have to correct for the pressure from the attraction of the molecules which is the constant a and is determined experimentally for each gas. The equation becomes: o