Energy and Chemical Change

Energy is the ability to do work and supply heat
Work is motion against an opposing force
Kinetic energy is the energy an object has because of its motion
For an object of mass m with velocity v
The law of conservation of energy states that energy cannot be created or destroyed
This can be applied to the collision of two particles with only kinetic energy
Potential energy (PE) is the energy of position or internal arrangement
KE can be converted into PE and vice versa
Work is required to pull the negatively charged electron away from the positively charged nucleus
The gain and loss of PE can be summarized
Pushing or pulling an object against an opposing force requires energy. The objects PE will rise.
When the opposing force is not resisted, the object’s PE falls
The SI unit of energy is the joule (J)
A 2 kg object moving at 1 meter per second has 1 J of kinetic energy
You may also encounter the calorie (cal)
The dietary Calorie (note capital), Cal, is actually 1 kilocalorie
When a cold and hot object come into contact, they eventually reach thermal equilibrium (the same temperature)
The energy that is transferred as heat comes from the object’s internal energy
The energy associated with the motion of the object’s molecules is referred to as its molecular kinetic energy
The internal energy is often given the symbol E or U
We are interested in the change in E:
The internal energy change is positive if the system absorbs energy from the surroundings and negative if it releases energy to its surroundings
The temperature of an object is related to the average kinetic energy of its atoms and molecules
Heat is a transfer of energy due to a temperature difference
Isolated warm (left) and cold (right) objects
Thermal contact is made: thermal energy is transferred from left to right
Thermal equilibrium: the same average KE for molecules in both objects
The energy of an object depends only on its current condition
The current condition is called the state
Internal energy is a state function because it is a measure of energy
An important property of state functions is that they are independent from the mechanism or method by which a change occurred
The object we are interested in is called the system
Everything outside the system is called the surroundings
A boundary separates the system from the surroundings
The system and surroundings together are called the universe
Systems are classified according to what can cross its boundary
Open systems can gain or lose mass and energy across their boundaries
Closed systems can absorb or release energy, but not mass, across their boundaries
Isolated systems cannot exchange energy or matter with their surroundings
Consider heat flowing between the system and surroundings
The sign of the heat change is used to say whether it was gained or lost
When heat is gained by an object, it is written as a positive number
When heat is lost by an object, it is written as a negative number
A spontaneous change is one that continues on its own
Heat flows spontaneously from a warmer to colder object
The heat directly gained or lost by an object is directly proportional to the temperature change it undergoes
The object’s specific heat (C) relates the heat (q) to the objects temperature change
The heat capacity is the amount of heat needed to raise the object’s temperature by one degree Celsius and has the units J/°C
C is an extensive property that can be determined from experiment, and is proportional to the sample mass
The specific heat capacity (s) is an intensive property, and is unique for each substance
For example
Example: The temperature of 251 g of water is changed from 25.0 to 30.0 °C. How much heat was transferred to the water?
ANALYSIS: Connect heat to the temperature change.
SOLUTION:
Note: Heat was absorbed because q is positive
Chemical bonds are the net attractive force between nuclei and electrons in compounds
Breaking a chemical bond requires energy
Making a chemical bond releases energy
The potential energy that resides in chemical bonds is called chemical energy
Chemical reactions generally involve both making and breaking chemical bonds
The net gain or loss of energy is often in the form of heat
Any reaction where heat is a product is called exothermic
Reactions that consume energy are called endothermic
Reactions can release heat by replacing “weak” bonds with “strong” ones
The amount of heat absorbed or released by a chemical reaction is called the heat of reaction
A calorimeter can be used to measure the heat of reaction
Calorimeters are usually designed to measure heats of reaction under conditions of constant volume or constant pressure
Pressure is the amount of force acting on a unit area:
Atmospheric pressure is the pressure exerted by the mixture of gases in the atmosphere
At sea level the atmospheric pressure is about 14.7 lb/in²
Other common pressure units are the atmosphere (atm) and bar:
14.696 lb/in² = 1.0000 atm = 1.0133 bar
qv and qp are used to show heats measured at constant volume or pressure, respectfully
In reactions where gases are produce or consumed qv and qp can be very different
If the volume change is , work (w) is
Note that the work of expansion is negative
Work and heat are alternate ways to transfer energy
Their sum is the change in internal energy the system undergoes
This is a statement of the first law of thermodynamics, which says that energy cannot be created or destroyed
Heat and work are not state functions because they depend on the path between the final and initial state
The heat produced by a combustion reaction is called the heat of combustion
The heats are measured in closed containers because the reactions consume and produce gases
The instrument used to measure these heats is called a bomb calorimeter
The reaction is run at constant volume so that
Heats of reactions in solution are usually run in open containers at constant pressure
They may transfer heat and expansion work
The heat change measured at constant pressure is the enthalpy, H
Enthalpy is also a state function
is negative for an exothermic process
is positive for an endothermic process
The the difference in the values of the internal energy and enthalpy change can be large for reactions that consume or release gases
The amount of heat that a reaction produces or absorbs depends on the number of moles of reactant that react
A set of standard states have been defined for reporting heats of reactions
Standard thermodynamic states are: 1 bar pressure for all gases and 1 M concentration for aqueous solutions
A temperature of 25 °C (298 K) is often specified as well
The standard heat of reaction is the value of the enthalpy change occurring under standard conditions involving the actual number of moles specified the the equation coefficients
An enthalpy change for standard conditions is denoted
For example, the thermochemical equation for the production of ammonia from it elements at standard conditions is:
The physical states are important
The law of conservation of energy requires
Enthalpy is a state function
An enthalpy diagram is a graphical representation of alternate paths between initial and final states
Remember to include the physical states of reactants and products in thermochemical equations.
Enthalpy changes for reactions can be calculated by algebraic summation
This is called Hess’s Law: The value of the enthalpy change for any reaction that can be written in steps equals the sum of the values of the enthalpy change of each of the individual steps.
Enthalpy changes for a huge number of reactions may be calculate using only a few simple rules
Rules for Manipulating Thermochemical Equations:
When an equation is reversed the sign of the enthalpy change must also be reversed.
Formulas canceled from both sides of an equation must be for substances in identical physical states.
If all the coefficients of an equation are multiplied or divided by the same factor, the value of the enthalpy change must likewise be multiplied or divided by that factor.
An enormous database of thermochemical equations have been compiled:
The standard heat of combustion is the amount of heat released when 1 mol of a fuel completely burns in pure oxygen gas with all products brought to 25 °C and 1 bar
Standard heats of combustion are always negative and produce water in liquid form
The standard enthalpy of formation of a substance is the amount of heat absorbed when 1 mole of the substance if formed at 25 °C and 1 bar from its elements in their standard states
The standard enthalpy of formation for elements in their standard states are zero
These are the values most commonly used to calculated standard enthalpy changes for reactions
Standard enthalpies of formation are given in Table 7.2 and Appendix C
Hess’s law can be restated in terms of standard enthalpies of formation:
Example: Calculate the enthalpy of reaction for 2NO(g)+O2(g)2NO2(g)
ANALYSIS: Use Hess’s law and Table 7.2
SOLUTION: