# Energy & change

## Learning objectives

• Describe, distinguish, and relate the following properties. Predict whether these properties increase, decrease, or stay the same over the course of a given chemical or physical change.
• temperature
• thermal energy
• Understand heat on both theoretical and experimental levels.
• Relate heat transferred to changes in thermal energy when no work is done.
• Relate heat to an object's mass and initial and final temperatures. Clearly distinguish heat and temperature.
• Explain how heat can be measured experimentally (calorimetry).
• Estimate the final temperature when hot and cold objects are brought into contact.
• Define heat capacity and specific heat. Describe how these quantities can be measured experimentally.
• Define enthalpy. Distinguish enthalpy from thermal energy.
• Describe how changes in enthalpy and thermal energy accompanying a chemical reaction can be measured calorimetrically.
• Define bond energy. Use tables of bond energies to estimate the enthalpy of a reaction.
• Write and manipulate thermochemical equations.
• Combine a set of step thermochemical equations to obtain a net thermochemical equation (Hess's Law)
• Write thermochemical equations for combustion and formation reactions.

## Before you start...

• energy, energy units, and the difference between kinetic and potential energy
• review the relationship between temperature and average molecular velocity

## Lecture outline

In all chemical change, chemical bonds are broken or formed. Energy is required to break a chemical bond (just as energy is required to stretch a spring until it breaks). Conversely, forming a chemical bond releases energy. Virtually all chemical reactions absorb or release energy because bond making seldom exactly balances bond breaking in the reaction. In this unit, we will learn to measure and predict the amount of heat absorbed or released by a chemical reaction.

### The concept of energy

• the usual definition of energy: the ability to do work
• work is moving an object against an opposing force
• work = distance × opposing force
• SI unit of work or energy: the joule (J)
• two basic forms of energy
• potential energy: energy of position
• examples
• boulder on a ledge
• cations and anions
• chemical bonds
• kinetic energy: energy of motion
• examples
• pool balls
• molecules
• why is the concept of energy useful?
• if something is isolated from everything else, its total energy never changes
• this allows seemingly unrelated behaviors of the system to be connected
• example: the pendulum
• Two things energy is NOT
• some sort of invisible fluid
• something which can be measured directly

### Thermal energy

• definition: energy due to chaotic molecular motions
• three factors affecting thermal energy
• temperature
• higher temperature leads to higher thermal energy
• sample size
• a cup of hot coffee has more energy than a teaspoon of coffee, all other things being equal.
• composition
• E(solid) < E(liquid) < E(gas), all other things being equal
• anything that changes temperature, sample size and/or composition of an object can change its thermal energy

### Heat

• definition: transfer of thermal energy due to a temperature difference
• thermal energy isn't measurable, but heat is
• Three factors affect how much heat an object absorbs or loses
• mass of object
• temperature change of object
• final temperature - initial temperature
• if there is no change in temperature, no heat flows
• composition of object
• specific heat: heat required to raise the temperature of 1 g of material by 1 K
• different materials have different specific heats  material at298 K and 1 atm specific heat(J/g K) ice 2.09 water 4.18 steam 1.86 sodium 1.23 aluminum 0.9 iron 0.45
• heat capacity: heat required to raise the temperature of an object by 1 K
• computing heat
• heat = mass x specific heat x temperature change = heat capacity x temperature change
• examples
• 100.0 g of water cools from 30.10°C to 25.05 °C. How much heat is released?
• 100.0 g of water at 25.00 °C absorbs 100 J of heat. What is its final temperature?
• A stone weiging 2.0 g absorbs 5.0 J of heat and warms by 3.0 °C. What is the specific heat of the stone? What is the heat capacity of the stone?

### Enthalpy

• enthalpy change: heat absorbed or released by a process running at constant pressure
• symbol: H = final enthalpy - initial enthalpy
• note: enthalpy changes depend only on initial and final states, not on the route between them!
• state function: a quantity that depends only on the present state (properties) of the system, not on the process used to arrive at that state.
• enthalpy changes are slightly different from thermal energy changes
• constant pressure processes must use a little energy to push back the atmosphere
• enthalpy change is thermal energy change, minus work against atmosphere, for a constant pressure process

### Comparing Thermochemical Quantities

 definition SI units type temperature hotness/coldness property that controls direction of heat flows K intensive property thermal energy energy due to molecular motions J extensive property heat transfer of thermal energy due to a temperature difference J process enthalpy adjusted thermal energy J extensive property

### Calorimetry

• calorimetry is the experimental measurement of heat flows
• bomb calorimetry
• constant pressure calorimetry: heat generated by a constant pressure process
• strategy for solving calorimetry problems
1. identify all q's by deciding which parts of the system absorb or release significant amounds of heat
2. set up an energy conservation equation. set the sum of all heat flows to zero.
3. introduce T's. replace experimental q's with temperature changes, using q = mcT or q = CT.
4. solve the equation for the desired quantity.

### Enthalpy of Reaction

• chemical reactions usually absorb or release heat
• energy must be absorbed to break a chemical bond
• energy is released when a chemical bond forms
• exothermic vs. endothermic reactions  Reaction type: exothermic endothermic heat is: released absorbed reaction vessel temperature: rises falls enthalpy change is negative positive net bond: formation breaking

### Thermochemical equations

• example: spacecraft reentry
• shockwave processes involve bond breaking  N2(g) 2N(g) H = +941 kJ O2(g) 2O(g) H = +502 kJ N2(g) + O2(g) 2NO(g) H = + 168 kJ
• heat shield processes involve bond making  2N(g) N2(g) H = -941 kJ 2O(g)O2(g) H = -502 kJ N(g) + O(g) NO(g) H = -638 kJ
• these are thermochemical equations: stoichiometric equations with reaction enthalpy
• whatever you do to the stoichiometric equation, do also to the reaction enthalpy!
• reversing the reaction reverses the sign on the reaction enthalpy
• scaling the reaction scales the reaction enthalpy
• How to combine 'step' thermochemical equations to get a 'target' equation:
1. write the step reactions.
2. write the target reaction.
3. reverse step reactions so products/reactants match the target reaction.
4. scale step reactions so products/reactants that don't appear in the target reaction will cancel out.
5. add the step reactions.
6. scale the resulting reaction so it matches the target reaction.
• H depends on pressures, concentrations, and temperatures of reactants and products!
• to keep things simple, define standard conditions:
• all solution concentrations are 1 M
• all gases have a partial pressure of 1 atm
• all liquids and solids are under an external pressure of 1 atm
• reaction occurs at 25°C
• write when the reaction is run under standard conditions
• special reaction enthalpies
• the following are often tabulated for use as 'step' reactions:  definition symbol sign enthalpy of formation enthalpy of formation of one mole of compound from its elements in their most stable forms Hf + or - enthalpy of combustion enthalpy of complete combustion of one mole of compound Hc always -
• use the same procedure we outlined earlier to combine formation or combustion reactions to get a target reaction

### Enthalpies of phase changes

 definition symbol sign enthalpy offusion Heat to melt 1 mole of solid to liquid Hfus always + enthalpy ofvaporization Heat to evaporate 1 mol of liquid Hvap always + enthalpy ofsublimation Heat to vaporize 1 mol of solid Hsub always +
• heating & cooling curves
• obtain heat capacities from slopes of curve where temperature changes
• plateaus are regions where melting or boiling is occuring
• temperatures at plateaus indicate melting and boiling points
• length of plateau is enthalpy of phase change
• mixtures give curves without flat plateaus

### Molecular view of enthalpy changes

• bond enthalpy: enthalpy change per mole when a bond is broken in the gas phase for a particular substance.
• average bond enthalpy: average enthalpy change per mole when the same type of bond is broken in the gas phase for many similar substances.

Average Bond Enthalpies in kJ/mol.
= denotes a double bond;
denotes a triple bond.  Cl S F O N C H H 432 368 563 463 391 415 436 C 328 259 477 = 441 351 728 = 292 615 = 890 348 615 = N 200 270 175 638 161 418 = 941 O 203 185 139 502 = F 251 310 158 S 277 266 Cl 243

• bond enthalpies are always positive: bond breaking is endothermic
• estimating H from bond enthalpies
• strategy: imagine reaction as a) dissociation of reactants into atoms, b) recombination of atoms into products.
1. Add enthalpies for all product bonds
2. Add enthalpies for all reactant bonds
3. H is approximately the difference between the product and reactant bond enthalpies
• limitations
• procedure doesn't account for molecular attractions/repulsions, so doesn't work well for liquid/solid phase reactions
• bonds interact with each other within molecules, so bond enthalpies really aren't additive

General Chemistry Online! Energy & change