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
- boulder on a ledge
- cations and anions
- chemical bonds
- kinetic energy: energy of motion
- 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
- definition: energy due to chaotic molecular motions
- three factors affecting thermal energy
- 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.
- 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
- 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
298 K and 1 atm
- 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
- 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 change: heat absorbed or released by a process running at constant pressure
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
||hotness/coldness property that controls direction of heat flows
||energy due to molecular motions
||transfer of thermal energy due to a temperature difference
||adjusted thermal energy
- 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
- identify all q's by deciding which parts of the system absorb or release significant amounds of heat
- set up an energy conservation equation. set the sum of all heat flows to zero.
- introduce T's. replace experimental q's with temperature changes, using q = mcT or q = CT.
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 vessel temperature:||rises
|enthalpy change is||negative
when the reaction is run under standard conditions
- 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
- adding reactions adds reaction enthalpies
- How to combine 'step' thermochemical equations to get a 'target' equation:
- write the step reactions.
- write the target reaction.
- reverse step reactions so products/reactants match the target reaction.
- scale step reactions so products/reactants that don't appear in the target reaction will cancel out.
- add the step reactions.
- 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 H°