CHEM 105N - Chapter 7: Energy and Chemical Processes

Energy and Chemical Processes

Energy

• Energy is the capacity to do work or transfer heat.
  • Work: Energy transferred when a force exerted on an object causes a displacement of that object.
  • Heat: Energy used to cause the temperature of an object to increase.

Two Fundamental Forms of Energy

• Kinetic Energy
  • Energy of motion or movement.
• Potential Energy
  • Energy that is waiting to happen.
  • Depends on the relative position of an object compared to other objects.
  • Stored in the chemical bonds that make up substances.

Kinetic and Potential Energy

• Kinetic Energy (KE) = $\frac{1}{2} m v^2$, where m is mass and v is velocity.
• Potential energy is waiting to happen.
• Kinetic energy is energy that is happening.

Law of Conservation of Energy

• Energy cannot be created or destroyed.
• The total energy of a system does not change in any reaction or transformation.
• Potential Energy
  • Stored energy, or energy which has the “potential” to do work.
  • Static objects have more potential energy.
  • Moving objects have less potential energy.
• Kinetic Energy
  • Slower moving objects have less kinetic energy.
  • Faster-moving objects have more kinetic energy and can do more work.
  • Heat loss (to the surroundings) lowers the kinetic energy of the object, and heat gain (from the surroundings) raises its kinetic energy.

Units of Energy

• A calorie (cal) is the amount of energy needed to raise the temperature of 1 g of water by 1 °C.
• A joule (J) is another unit of energy.
• $1 \text{ cal} = 4.184 \text{ J}$
• Both joules and calories can be reported in the larger units kilojoules (kJ) and kilocalories (kcal).
  • $1,000 \text{ J} = 1 \text{ kJ}$
  • $1,000 \text{ cal} = 1 \text{ kcal}$
  • $1 \text{ kcal} = 4.184 \text{ kJ}$

Questions

• How many kJ of energy are eaten in a normal 2,000 Calorie diet?
• How many kcal are used riding a bicycle for 2 hours if 200 kJ per hour are needed?
• How many kJ of energy are consumed when eating a glazed donut from Krispy Kreme, being worth 190 Calories (kcal)?

First Law of Thermodynamics

• Energy can be converted from one form to another, but it is neither created nor destroyed.
• To heat your home, chemical energy needs to be converted to heat.
• Sunlight is converted to chemical energy in green plants.

Definitions: System and Surroundings

• The portion of the universe that we single out to study is called the system.
• The surroundings are everything else.

Internal Energy

• By definition, the change in internal energy, $\,\Delta E$, is the final energy of the system minus the initial energy of the system:
  • $\Delta E = E<em>{\text{final}} - E</em>{\text{initial}}$

Changes in Internal Energy

• If \,\Delta E > 0, $E<em>{\text{final}} > E</em>{\text{initial}}$, the system absorbed energy from the surroundings.
• If \,\Delta E < 0, $E<em>{\text{final}} < E</em>{\text{initial}}$, the system released energy from the surroundings.

Thermodynamic Quantities

• Have Three Parts
  • A number
  • A unit
  • A sign
• Note about the sign:
  • A positive $\,\Delta E$ results when the system gains energy from the surroundings.
  • A negative $\,\Delta E$ results when the system loses energy from the surroundings.

Exchange of Heat Between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.
• When heat is released by the system into the surroundings, the process is exothermic.

Energy Changes

• $\Delta E = E<em>{\text{products}} - E</em>{\text{reactants}}$
• Energy is released in exothermic reactions, and energy is gained in endothermic reactions.
• $\,\Delta E$ is negative for an exothermic reaction and positive for an endothermic reaction.

Questions

• Endothermic or Exothermic?
  • Initial energy is +50.7 kJ and final energy is +82.1 kJ.
  • Initial energy is -72.4 kJ and final energy is -98.2 kJ.
  • Initial energy is +66.3 kJ and final energy is -82.1 kJ.

Phase Changes

• Conversion from one state of matter to another is called a phase change.
• Energy is either added or released in a phase change.
• Phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.

Energy Change and Change of State

• The heat of fusion is the energy required to change a solid at its melting point to a liquid.
• The heat of vaporization is the energy required to change a liquid at its boiling point to a gas.
• The heat of sublimation is the energy required to change a solid directly to a gas.

Heating Curves

• A graph of temperature vs. heat added is called a heating curve.
• The temperature of the substance does not rise during a phase change.
• For phase changes, the product of mass or moles and heat of fusion or vaporization is heat.

Question

• How much energy is needed to melt 43.8 g of Au at its melting point of 1064 °C? ($\Delta H_{\text{fus}} = 15.3 \text{ cal/g}$

Changes in Internal Energy

• $\Delta E = q + w$
• When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w).

Energy Changes in Reactions

• When molecules come together and react, bonds are broken in the reactants, and new bonds are formed in the products.
• Bond breaking always requires an input of energy.
• Bond formation always releases energy.
  • To cleave a bond, energy must be added.
  • To form a bond, energy is released.

Energy Changes in Reactions

• $\,\Delta H$ is the heat (q) absorbed or released in a reaction; it is called the heat of reaction or the enthalpy change.
• When energy is absorbed, the reaction is said to be endothermic, and $\,\Delta H$ is positive (+).
• When energy is released, the reaction is said to be exothermic, and $\,\Delta H$ is negative (−).

Bond Dissociation Energy

• The bond dissociation energy is the $\,\Delta H$ for breaking a covalent bond by equally dividing the $e^-$ between the two atoms.

• Bond dissociation energies are positive values because bond breaking is endothermic (\Delta H > 0).

• Bond formation always has negative values because bond formation is exothermic (\Delta H < 0).

• $H - H \rightarrow H \cdot + H \cdot$ $\,\Delta H = +104 \text{ kcal/mol}$

• $H \cdot + H \cdot \rightarrow H - H$ $\,\Delta H = -104 \text{ kcal/mol}$

Bond Dissociation Energy

• The stronger the bond, the higher its bond dissociation energy.
• In comparing bonds formed from elements in the same group, bond dissociation energies generally decrease going down the column.

Bond Enthalpy

• The enthalpy associated with breaking one mole of a particular bond in a gaseous substance.
• The bond enthalpy is always positive because energy is required to break chemical bonds.
• Energy is always released when a bond forms.
• The greater the bond enthalpy, the stronger the bond.

Average Bond Enthalpies (kJ/mol)

• C-H: 413
• N-H: 391
• O-H: 463
• F-F: 155
• C-C: 348
• N-N: 163
• O-O: 146
• C=C: 614
• N-O: 201
• O=O: 495
• Cl-F: 253
• C-N: 293
• N-F: 272
• O-F: 190
• Cl-Cl: 242
• C-O: 358
• N-Cl: 200
• O-Cl: 203
• C=O: 799
• N-Br: 243
• O-I: 234
• Br-F: 237
• C-F: 485
• Br-Cl: 218
• C-Cl: 328
• H-H: 436
• Br-Br: 193
• C-Br: 276
• H-F: 567
• C-I: 240
• H-Cl: 431
• I-Cl: 208
• H-Br: 366
• I-Br: 175
• H-I: 299
• I-I: 151

Questions

• Which bond is the strongest using the table of bond dissociation energy?

Endothermic Reaction

• Heat is absorbed.
• $\,\Delta H$ is positive.
• The bonds broken in the reactants are stronger than the bonds formed in the products.
• The products are higher in energy than the reactants.

Exothermic Reaction

• Heat is released.
• $\,\Delta H$ is negative.
• The bonds formed in the products are stronger than the bonds broken in the reactants.
• The products are lower in energy than the reactants.

Enthalpy of Reaction

• The change in enthalpy, $\,\Delta H$, is the enthalpy of the products minus the enthalpy of the reactants:
  • $\Delta H<em>{\text{rxn}} = H</em>{\text{products}} - H_{\text{reactants}}$
• This quantity, $\,\Delta H_{\text{rxn}}$, is called the enthalpy of reaction or the heat of reaction.
• Exothermic reactions have \,\Delta H < 0.

Questions

• Identify each of the following reactions as exothermic or endothermic.
  • N2(g) + 3H2(g) → 2NH3(g) + 22 kJ
  • CaCO3(s) + 133 kJ → CaO(s) + CO2(g)
  • 2SO2(g) + O2(g) → 2SO3(g) + heat
• Identify the following reaction as exothermic or endothermic: 88 kJ + 2 H2O (l) → 2 H2O (g)
• Identify the following reaction as exothermic or endothermic: 2 H2O2 (l) → 2 H2O (l) + O2 (g) + 196 kJ
• If 4.50 g of CH4 are combusted according to the following reaction, how many kJ were released? CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (l) $\,\Delta H = -890 \text{ kJ}$
• If 15.0 g of NO are produced according to the following reaction, how many kJ were absorbed?
• How many grams of O2 reacted if 306 kJ were released in the following reaction?

How Do Reactions Occur?

• The collision model is based on the kinetic molecular theory.
• Molecules must collide to react.
• If there are more collisions, more reactions can occur.
• So, if there are more molecules, the reaction rate is faster.
• In a chemical reaction, bonds are broken, and new bonds are formed.
• Molecules can only react if they collide with each other.

Orientation of Molecules

• Molecules can often collide without forming products.
• Aligning molecules properly can lead to chemical reactions.
• Bonds must be broken and made, and atoms need to be in proper positions.

Energy Needed for a Reaction to Take Place (Activation Energy)

• The minimum energy needed for a reaction to take place is called activation energy.
• An energy barrier must be overcome for a reaction to take place.

Transition State (Activated Complex)

• Reactants gain energy as the reaction proceeds until the particles reach the maximum energy state.
• The organization of the atoms at this highest energy state is called the transition state (or activated complex).
• The energy needed to form this state is called the activation energy.

Reaction Progress

• Plots are made to show the energy possessed by the particles as the reaction proceeds.
• At the highest energy state, the transition state is formed.
• Reactions can be endothermic or exothermic after this.
• Rate constant depends on the magnitude of Ea.

Reaction Conditions

There are three conditions required for a chemical reaction to occur:

1. Collision: The reactants must collide.
2. Orientation: The reactants must align properly to break and form bonds.
3. Energy: The collision must provide the energy of activation.

Activation Energy

• Activation energy is the amount of energy required to break the bonds between atoms of the reactants.
• If the energy of a collision is less than the activation energy, the molecules bounce apart without reacting.
• Many collisions occur, but only a few actually lead to the formation of product.
• The activation energy is the energy needed to convert reacting molecules into products.

Enthalpy of Reaction

• The heat of reaction:
  • Amount of heat absorbed or released during a reaction.
  • Difference between the energy of breaking bonds in the reactants and forming bonds in the products.
  • $\,\Delta H = \Delta H<em>{\text{products}} - \Delta H</em>{\text{reactants}}$

Exothermic Reactions

• In an exothermic reaction,
  • the energy of the reactants is greater than that of the products, and heat is released along with the products.
  • Enthalpy of reaction ($\Delta H$) value is negative (–), indicating heat is released.

Endothermic Reactions

• In an endothermic reaction,
  • the energy of the reactants is lower than that of the energy of the products.
  • Enthalpy of reaction ($\Delta H$) value is positive (+), indicating heat is absorbed.