Energy Chem Honors Chapter 10
7 Forms of Energy • Sound- from vibration of sound waves • Chemical- fuel, gas, wood, battery • Radiant (light)- electromagnetic energy • Electrical energy- electrons moving among atoms- as in the conductive wire of an electrical cord • Atomic- nuclear (from nucleus of an atom) • Mechanical- walk, run • Thermal- heat
What is energy? • Energy is the ability to do work or produce heat • Two types: • Potential energy • Kinetic energy
Potential energy- energy due to position or composition • Ex: • water behind a dam • Attractive and repulsive forces
Kinetic energy – is energy due to the motion of the object and depends on the mass of the object and depends on the mass (m) and velocity (v) • KE = ½ mv2
Law of Conservation of Energy • States that energy can be converted from one form to another but can never be created or destroyed
Chemical energy stored in bonds of gasoline moleculesconverted to mechanical and heat energy through combustion
Hoover Dam- hydropower plant converts mechanical energy (flowing water) to electromagnetic, which is transported to homes, and then converted back into mechanical energy in a blender
State Function • Property of a system that does not depend on the pathway to its present state. • Ex: Displacement is a state function: I send 2 students to the cafeteria…the cafeteria is a specific distance from here….even though the 2 students take different routes to get to the cafeteria, they bith end up the same distance away from this room.
Energy is a state function : • Work and heat are not state functions:
Temperature and Heat • What is the difference between warm water and cold water?
Thermal energy: random motions of components of an object • Temperature: measures the random motions of the components of a substance (thermal energy) • Heat: flow of energy due to a temperature difference
System and Surroundings • System- part of the universe that we are focusing our attention on • Surroundings- everything else in the universe
Exothermic process – a process that results in the evolution of heat- energy flows out of the system • Endothermic process- a process that absorbs energy from the surroundings- energy flows into the system
Exothermic or endothermic? • 1. Your hand gets cold when you touch ice • 2. ice melts when you touch it • 3. Ice cream melts • 4. Propane is burning in a propane torch. • 5. Water drops on your skin evaporate after swimming • 6. Two chemicals mixing in a beaker give off heat
Exothermic • Endothermic • Endothermic • Exothermic • Endthermic • exothermic
Energy from an exothermic reaction comes from the difference in potential energy between the products and reactants • Which has lower energy, the reactants or products?
-Total energy is conserved -exothermic so energy flows from system to surroundings - So the energy gained by the surroundings must be equal to the energy lost by the system Burning match
Burning match cont. • The potential energy lost from the burned match came from the stored energy in the bonds of the reactants • In any exothermic reaction some of the potential energy stored in the chemical bonds is converted to thermal energy- which is random kinetic energy- as heat
Thermodynamics • Thermodynamics- study of energy • Law of Conservation of energy is also known as • THE FIRST LAW OF THERMODYNAMICS!! • The energy of the universe is constant.
1st Law- A Quantitative application • We can use the first law to analyze energy changes in chemical systems • Internal energy (E) of a system is the sum of all of the PE (potential energy) and KE (kinetic energy) of a system E = PE + KE
Internal Energy, E H2(g) , O2(g) H2O (l)
The internal energy of a system can be changed by a flow of work or heat or both so mathematically we have ΔE = q + w Δ = change in q = heat added or liberated by the system w = work done on the system
Thermodynamic quantities have • Number – indicates the magnitude of the change • Sign (+ or -) – indicates the direction of the flow from the SYSTEM’s point of view
q and w • When heat is added to a system, or work is done on a system, the internal energy increases so • When heat is transferred from the surroundings to the system q is positive • When work is done on a system by the surroundings w is positive
q and w • Work and heat created by the system transferred to the surroundings lowers the internal energy of the system • When work is done by the system on the surroundings, w is negative • When heat flows from the system to the surroundings, q is negative
In an exothermic process, what is the sign of q? • In an endothermic process, what is the sign of q?
Measuring Energy Changes • Common units of energy change • Calorie – amount of energy (heat) required to raise the temperature of one gram of water one degree Celsius • Joule – in terms of a calorie • 1 calorie = 4.184 J
Measuring Energy Changes cont. • Express 60.1 cal of energy as Joules • 60.1 cal x4.184 J = 251 J • 1 cal
Measuring Energy Changes Cont. • The energy (heat) required to change the temperature of a substance depends on • The amount of substance being heated (grams) • The temperature change • Identity of the substance
Measuring Energy changes cont • Identity of the substance- definition of a calorie is based on water, what about other substances?
Measuring Energy Changes cont. • Specific heat – the amount of energy required to change the temperature of one gram of any substance by one degree Celsius • Water 4.184 J/g °C • Iron (s) 0.45 J/g °C • Silver (s) 0.24 J/g °C
Measuring Energy Changes cont. • q = mcΔT q = energy in the form of heat m = mass c = specific heat ΔT = change in temperature in °C
Measuring Energy Changes cont. • Determine the amount of energy (heat) in Joules required to raise the temperature of 7.40 g of water from 29°C to 46°C
Measuring Energy Changes cont. • Ex: What quantity of energy (in Joules) is required to heat a piece of iron weighing 1.3 g from 25°C to 46°C ?
Calorimetry • A calorimeter is used to determine the heat associated with a chemical reaction
Calorimeter Constant • All parts of a calorimeter heat up or cool down as heat is released or absorbed in the chemical process • Heat capacity of the calorimeter, C, also known as the calorimeter constant • Defined as the sum of the products of the specific heat and the mass of all components of the calorimeter
Calorimeter Constant cont. • Formula for the heat energy produced in the calorimeter is q = Ccalorimeter ΔT
Calorimeter Constant Ex • A calorimeter has a heat capacity of 1265 J/°C • A reaction causes the temperature of the calorimeter to change from 22.34 C to 25.12 C. How many joules of heat were released in this process?
Calorimetry • Can determine the temperature change of the reaction, and use the heat capacity of the calorimeter to determine the ΔH of the reaction • What is H?
Enthalpy (H) • For most chemical reactions chemists are interested in the heat generated at constant pressure, which is denoted as qp= H • H is the heat content of a compound • ΔH= which is the difference in heat between the reactants and products • so ΔH= Hproducts - Hreactants
Thermochemistry • Because chemical reactions occur at constant pressure, ΔH= qp for chemical reaction • ΔH is a state function- independent of the path