Chemical Principles Ch. 9

Chemical Principles (Zumdahl/DeCoste)

Chapter 9

9.1 – The Nature of Energy

  • Energy: The capacity to do work or produce heat
    • Law of conservation of energy: energy can be converted form one form to another but can neither be created nor destroyed
    • Potential energy: energy due to position or composition
    • Kinetic energy: due to the motion of an object, depends on mass and velocity
    • KE = ½ mv2
  • Energy can be converted from one form to another easily
  • Ways to transfer energy
    • Heat: the transfer of energy between two objects due to a temperature difference (not a substance contained in an object)
    • Work: a force acting over a distance
    • Frictional heating: the transfer of energy to a surface as heat
  • State function: (also known as state property) a property of the system that depends only on its present state; does not depend in any way on the system’s past or future
    • The value of a state function does not depend on how the system arrived at the present state; only the characteristics of the present state

Chemical Energy

  • System: the part of the universe where we focus attention
  • Surroundings: everything else in the universe
  • Exothermic: reaction that results in the evolution of heat; energy flows out of the system
  • Endothermic: reaction that absorbs energy form the surroundings; energy flows into a system

Thermodynamics

  • Thermodynamics: the study of energy and its inter conversions
    • First law of thermodynamics: The energy of the universe is constant
    • Internal energy (E): internal energy of a system is the sum of kinetic and potential energies of all the “particles” in the system
  • Thermodynamic quantities:
    • Number: magnitude of the change
    • Sign: direction of the flow
    • Reflects the system’s point of view
    • If energy flows into the system, q is +x and if energy flows out of the system, q is -x
    • If the system does work on the surroundings, w is negative; if the surroundings do work on the system, w is positive

Compression and expansion of gases

∆V = final volume – initial volume = A * ∆h

w = –P∆V

9.2 Enthalpy

H = E + PV

  • H: enthalpy of the system
  • E: internal energy of the system
  • P: pressure of the system
  • V: volume of the system
  • Enthalpy is also a state function
    • At a constant pressure, the change in enthalpy of the system is equal to the energy flow as heat
    • For a chemical reaction, the enthalpy change is given by

9.3 Thermodynamics of Ideal Gases

  • It is often useful to refer to the properties of matter in the simplest possible context
    • Ideal gases – hypothetical condition approached by real gases at high temperatures and low pressures
    • For an ideal gas:
    • The only way to change the KE of a gas is to change the temperature
    • Energy required to change 1 mole of ideal gas by ∆T:
    • Molar heat capacity: Energy required to raise the temperature of 1 mole of that substance by 1 K (

Constant Volume

  • No PV work done (∆V = 0)
  • Cv = (3/2)R = “heat” required to change the temperature of 1 mole of gas by 1 K at constant volume

Constant Pressure

  • Volume increases, PV work occurs
    • Energy must be supplied to translational energy of the gas and to provide work of the gas as it expands
    • Heat required to raise T of 1 mole by 1 K =
    • Cp, the molar heat capacity of an ideal gas at constant pressure, is

Heating a Polyatomic Gas

  • We assume that ideal gases consist of “particles” with no structure
    • Monoatomic real gases measure values of Cv very close to
  • Real polyatomic molecules have much higher observed values for Cv
    • These molecules absorb energy to increase rotational and vibrational movement in addition to translational movement
    • Energy that is absorbed for vibrational and rotational energies do not contribute to translational kinetic energy, and therefore does not increase the temperature
    • Elevated Cv value is not caused by no ideal behavior; does not depend on whether the gas obeys the ideal gas law

Energy and Enthalpy

  • Enthalpy: H = E + PV
    • Change in enthalpy: ∆H = ∆E ∆(PV) => ∆H = ∆E + nR∆T
    • Substituting in energy, we get that ∆H = nCp ∆T

9.4 Calorimetry

  • Calorimetry: the science of measuring heat
    • Calorimeter: used to determine the heat associated with a chemical reaction experimentally
    • Heat capacity:
    • Specific heat capacity: the heat capacity given per gram of a substance

9.5 Hess’s Law

  • Hess’s Law: Enthalpy is a state function, so the change in enthalpy going from initial to final state is independent of the pathway

Characteristics of Enthalpy Changes

  1. If a reaction is reversed, the sign for ∆H is also reversed

    1. The sign of ∆H indicates the direction of heat flow at constant pressure

  2. The magnitude of ∆H is directly proportional to the amount of reactants and products

    1. If the coefficients of a balanced reaction are multiplied by some integer, then ∆H is also multiplied by the same integer

9.6 Standard Enthalpies of Formation

  • Standard enthalpy of formation (∆H°f): the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states
    • Superscript 0 (°): process has been carried out under standard conditions
    • Standard state: precisely defined reference state
    • Always given per mole of product, with the product in the standard state

Definition of Standard States

  1. For a gas, the standard state is a pressure of 1 atm
  2. For a substance in solution, the standard state is a concentration of 1 M at an applied pressure of 1 atm
  3. For a pure substance in condensed state (liquid or solid), the standard state is pure liquid or solid
  4. For an element the standard state is the form in which it exists (is most stable) under conditions of 1 atm and temperature of interest (usually 25°C)

∆H°reaction = ∑∆H°f (products) - ∑∆H°f (reactants)

Enthalpy Calculations

  1. When a reaction is reversed, the sign of ∆H is reversed, but the magnitude stays the same
  2. When the balanced equation for a reaction is multiplied by an integer, ∆H is multiples by the same integer
  3. The change in enthalpy for a reaction can be calculated from the enthalpies of formation of the reactants and products:

∆H°reaction = ∑∆H°f (products) - ∑∆H°f (reactants)

  1. Elements in their standard states are not included in the ∆H reaction calculations (∆H°f for an element in its standard state is zero)

9.7 Present Sources of Energy

Petroleum and Natural Gas

  • Petroleum: thick, dark liquid composed mainly of hydrocarbons, which contain hydrogen and carbon
  • Natural gas: associated with petroleum deposits, consists mostly of methane but also significant amounts of ethane, propane, and butane
  • Pyrolytic (high-temperature) cracking: process by which the heavier molecules of kerosene fraction are heated until they break into smaller molecules of hydrocarbons in the gasoline fraction
    • Kerosene (fraction C10 – C18) – excellent lamp oil
    • Gasoline (fraction C5 – C10 ) – used to power gasoline-powered engines

Coal

  • Coal: formed form the remains of plants that were buried and subjected to pressure and heat over long periods of time
    • Over time, percentage of carbon content increases
    • Energy available from combustion increases as the carbon content increases
  • Pollution
    • Burning coal can pose pollution problems
    • High-sulfur coal yields sulfur dioxide, which leads to acid rain
    • The carbon dioxide produced when burned has significant effects on the climate

Effects of carbon dioxide on climate

  • The earth receives a tremendous amount of energy from the sun, some of which is reflected by the atmosphere, and the rest of which passes through to the surface
    • Molecules in the atmosphere, principally H2O and CO2, strongly absorb infrared radiation and trap it in the atmosphere
    • Greenhouse effect: raising the temperature of the earth due to greenhouse gases that absorb and trap infrared radiation

9.8 New Energy Sources

  • Coal conversion possibilities
    • Produce a gaseous fuel by breaking many of the carbon-carbon bonds, which hare replaced by carbon-hydrogen and carbon-oxygen bonds (syngas – synthetic gas)
    • Formation of coal slurries – suspension of fine particles in a liquid, night replace solid coal and residual oil as fuel for electricity-generating power plants
  • Hydrogen as a fuel
    • Heat combustion of hydrogen gas is highly exothermic, but virtually none of the hydrogen on earth exists as free gas
    • Potential methods of obtaining hydrogen gas: electrolysis of water, thermal decomposition of water, and biological decomposition of water
    • Potential problems:
    • Storage and transportation of hydrogen; H2 decomposes to atoms on metal surfaces, which can cause structural changes that would lead to pipeline failure
    • Energy available per unit of volume: although the energy per gram of hydrogen is greater, the energy per unit volume is much less
  • Other energy alternatives
    • Oil shale: complex carbon-based material called kerogen contained in porous rock formations; most of the fuel is not fluid and cannot be pumped
    • Ethanol: readily available from many sources, but does not vaporize easily when temperatures are low
    • Seed oil: renewable fuel source and energy efficient; net energy gain