Energy

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Energy

• Chapter 10

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Energy

• Objective to understand the general properties
  of energy.
• Energy- is the ability to cause change. Energy
  comes in many forms. Some of these forms include
  radiant, electrical, chemical, thermal and
  nuclear energy. The basic unit of energy is the
  joule (JEWL)
• Kinetic Energy (KE)- is energy in the form of
  motion, is determined by mass and speed.
• Potential Energy (PE)- is stored energy. The
  amount of potential energy a sample of matter has
  depends on its position or condition. The greater
  the height of an object, the greater its
  potential energy.

KE 1/2mv2
P.E mgh
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10.2 Temperature and Heat

• Objective to understand the concepts of
  temperature and heat.
• Temperature is a measure of the random motions
  of the components of a substance.
• Heat is a flow of energy due to a temperature
  difference.
• Always the final temperature is the average of
  the original temperatures

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10.3 Exothermic and Endothermic Processes

• Objective To consider the direction of energy
  flow as heat.
• Exothermic is a process that results in the
  release of heat. For example, in the combustion
  of a match, energy flows out of the system as
  heat.
• Endothermic is a process that absorb energy from
  the surrounding. For example boiling water to
  form steam.

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10.5 Measuring Energy Changes

• Objective to understand how heat is measured.
• One calorie is the amount of energy(heat) needed
  to raise the temperature of one gram of water by
  1C
• kcal energy needed to raise the temperature of
  1000 g of water 1C
• joule (an SI unit)
• 1cal 4.184 J
• Practice Express 60.1 cal of energy in units of
  joules.

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• The specific heat capacity (s) of a material is
  the amount of energy transferred as heat that
  will raise the temperature of 1 g of a substance
  by 1 degree Celsius.
• Practice A 2.8-g sample of pure metal requires
  10.1 J of energy to change its temperature from
  21C to 36 C. What is this metal?

• Specific heat (C) is measured in J/gC

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• Using Specific Heat
• Specific heat can be used to measure changes in
  energy.
• Energy (heat) Q mass(m) Change (?T)
  specific heat(s)
• 
  in temperature
• Q m ?T s
• J g
  C J/gC
• The symbol ? (delta) means change in
• ?T is the change in temperature
• ?T Tfinal Tinitial

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• How does the specific heat of water compare with
  the specific heats of the other materials?

Specific Heat of Some Common Materials (J/gC)
Water (l) liquid 4.186 Water (s)
ice 2.03 Water(g) steam 2.0 Alcohol 2.450 Alu
minum 0.89 Carbon (Graphite) 0.71 Sand 664 I
ron 0.45 Copper 385 Silver 0.24 Gold 0.
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10.6 Thermochemistry (Enthalpy)

• Objective to consider the heat (enthalpy) of
  chemical reactions.
• Enthalpy (H) at constant pressure (p), the
  change in enthalpy equals to the energy that
  flows as heat.
• ? Hp heat
• Practice When 1 mol of methane (CH4) is burned
  at constant pressure, 890 kJ of energy is
  released as heat. Calculate ? H for a process in
  which a 5.8g sample of methane is burned at
  constant pressure.

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10.7 Hesss Law

• Objective to understand Hesss law
• Calorimeter is a device used to determine the
  heat associated with a chemical reaction.
• Hesss Law States that the overall enthalpy
  change in a reaction is equal to the sum of the
  enthalpy changes for the individual steps in the
  process.
• The overall reaction for oxidation of nitrogen to
  produce nitrogen dioxide is
• N2(g) 2O2(g) ? 2NO2(g) ?H1 68 kJ
  or
• N2(g) O2(g) ? 2NO(g) ?H2 180
  kJ
• 2NO(g) O2(g) ? 2NO2(g) ?H3 -112
  kJ
• N2(g) 2O2(g) ? 2NO2(g) ?H2 ?H3 68
  kJ

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Bomb Calorimeter
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Practice

• Two forms of carbon are graphite, the soft,
  black, slippery material used in lead pencils
  and as a lubricant for locks, and diamond, the
  brilliant, hard gemstone. Using the enthalpies
  of combustion for graphite (-394 kJ/mol) and
  diamond (-396 kJ/mol), calculate ?H for the
  conversion of graphite to diamond
• Cgraphite(s)? Cdiamond(s)

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Gibbs Free Energy

• The energy free to do work is the change in Gibbs
  free energy.
• DGº DHº - TDSº
• (T must be in Kelvin)
• All spontaneous reactions (reactions that will
  happen) release free energy.
• So DG lt0 for a spontaneous reaction.
• Entropy (S) Is the degree of randomness or
  disorder.

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• Entropy lt Entropy ltlt Entropy Of
• of a solid of a liquid of a
  gas
• Standard entropy is the entropy at 25ºC and 1
  atm pressure. Abbreviated Sº, measured in J/K.
• The change in entropy for a reaction is
• DSº Sº(Products)-Sº(Reactants).
• A change with positive DS and negative DH is
  always spontaneous.
• A change with negative DS and positive DH is
  never spontaneous.

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DGDH-TDS
Spontaneous?
DH
DS
DG
At all Temperatures
At high temperatures, entropy driven
At low temperatures, enthalpy driven
Not at any temperature, Reverse is spontaneous
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Practice

• Determine if the following changes are
  spontaneous at 25ºC.
• 2H2S(g) O2(g) 2H2O(l) S(rhombic)
• At what temperature does it become spontaneous?
• ?Hf for each component is
• H2S -20.1 kJ O2 0 kJ H2O -285.8 kJ
  S 0 kJ
• Then ?H (Products Reactants)
• ?H 2 (-285.8) - 0 - 2 (-20.1) (0)
  -531.4 kJ

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• ?S for each component is
• H2S 205.6 J/K O2 205.0 J/K
• H2O 69.94 J/K S 31.9 J/K
• Then ?S (Products Reactants)
• ?S2 (69.94) - 2(31.9) - 2 (205.6)
  205 -412.5 J/K
• ?G ?H - T ?S
• ?G -531.4 kJ - 298K (-412.5 J/K)
• ?G -531.4 kJ - -123000 J
• ?G -531.4 kJ 123 kJ
• ?G -408.4 kJ lt 0, So it is Spontaneous