Chapter 16: Reaction Energy Study Guide

Section 1: Thermochemistry

Define Thermochemistry

Know the differences between temperature and heat

Define Calorimeter

Be able to use the temperature change and known mass of a rxn to determine the amount of energy released or absorbed

Define temperature

Relationship between the amount of kinetic energy and temperature in a reaction

• what does this mean for what the measurement of temperature is based on?

Define heat of reaction

Phase change = change in potential energy, whereas heat change= change in kinetic energy

Define Heat

How energy transferred as heat is moved

Know what the quantity of energy transferred as heat during aa temperature change depends on

Know what differences in energy absorption in metals depends on

Define specific heat

Know what pressure conditions specific heat is measured under

know how and what types of problems to use cp=q/m x delta T

know how and what types of problems to use q=cp x m x delta T

what happens if a mixture of hydrogen and oxygen is ignited

• understand the energy changes involved in the reaction, including the release of heat and the formation of water vapor.

• understand what type of reaction this is (endothermic or exothermic)

define thermochemical equations

• be able to identify thermochemical equations

• know the rules to writing thermochemical equations

define enthalpy change

define the enthalpy of reaction

know application of the formula:  delta H = Hproducts- Hreactants

what happens to the enthalpy of products in an endothermic reaction? what about exothermic reactions?

Is the decomposition of water vapor endothermic or exothermic? why?

Why is it important for the physical states of reactants and products to be included in thermochemical equations?

When is delta H negative? When is it positive?

What are the 4 things to consider when looking at all thermochemical equations?

Understand potential energy diagrams:

• be able to determine if its endothermic or exothermic

• what represents the potential energy of the reactants?

• what represents the potential energy of the products?

• What represents the heat of reaction delta H?

• What represents the activation energy of the forward reaction?

• What represents the activation energy of the reverse reaction?

• What represents the potential energy of the activated complex?

• Determine if the reverse reaction is endothermic or exothermic.

• What points in the diagram would change if a catalyst were added?

Be able to use the temperature change formula: Q= mass x delta T x specific heat capacity

• Test questions will ask how many joules are

Be able to use the phase change formula: Q= mass x heat of fusion or Q= mass x heat of vaporization depending on the type of phase change occurring.

516-518: Vonese

Driving Force of Reactions

• The change in energy of a reaction system is one of the two factors that allow chemists to predict whether a reaction will occur spontaneously and to explain how it occurs

• The randomness of the particles in a system is the second factor affecting whether a reaction will occur spontaneously

Reactions generally move to a lower-energy state.

• The majority of chemical reactions in nature are exothermic, resulting in products that are more stable and have lower energy than the reactants.

• Reactions naturally tend towards a lower energy state

• Some endothermic reactions can occur spontaneously, indicating that energy (such as continued heating) is not the only factor determining spontaneity

Entropy measures randomness in a system.

•  A naturally occurring endothermic process is melting.

  • Such as an ice cube melting spontaneously at room temperature as energy is transferred from the warm air to the ice

• During melting, the well-ordered arrangement of water molecules in the ice crystal is lost, and a less-ordered liquid phase of higher energy content is formed

• A system that can go from one state to another without a decrease in enthalpy does so with an increase in entropy

• There is a general tendency in nature to proceed in a direction that increases the randomness of a system

• A random system lacks a regular arrangement of its parts. Tendency toward randomness is called entropy.

• Entropy, S, can be defined in a simple qualitative way as a measure of the degree of randomness of the particles, such as molecules, in a system.

• Entropy in states of matter

  • Solids: particles in fixed position, vibrating; randomness is LOW, so entropy is LOW

  • Liquids: particles are moving rapidly and are much farther apart; MORE random, entropy is HIGHER in liquids (compared to solids)

  • Gases: particles are moving rapidly and are much farther apart; MUCH MORE random, entropy is HIGHER in gases

  • General rule: Entropy in gases>>liquids>solids 

    • (exception of liquid mercury, which is less than some solids)

• The entropy of a pure crystalline solid at absolute zero is zero

• As energy is added, randomness of molecular motion increases

• Measures of energy absorbed and calculations that are used to determine absolute entropy (standard molar energy) are kJ/(mol • K)

• Entropy change, which can also be measured, is defined as the difference between the entropy of the products and the reactants.

  • An increase in entropy is + ΔS.

  • A decrease in entropy is - ΔS.

• The process of forming a solution almost always involves an increase in entropy because there is an increase in randomness.

• This applies to mixing gases, dissolving a liquid in another liquid, and dissolving a solid in a liquid

• Example: Sugar in tea

  • Initial: solid sugar has low entropy as molecules are in one region, separate from water molecules

  • Then: after dissolving, sugar molecules are thoroughly mixed throughout tea; both sugar and water can be found anywhere, increasing system’s randomness and + ΔS

Free energy changes determine if a reaction is endothermic or exothermic.

• Processes in nature are driven towards two directions: towards least enthalpy and greatest entropy

• The direction that LOWERS free energy of a system will be the direction natural processes proceed in 

• Enthalpy and Entropy fight for dominance. The dominant factor determines if forward or reverse reaction is favored

• Free energy: the combined enthalpy-entropy function, also known as “Gibbs free energy”

• Assess the tendency of enthalpy and entropy to change 

• Free-energy change: denoted by ΔG; defined as the difference between the change in ΔH and TΔS

Typical Measurements for  this equation

• TΔS: kj or J

• ΔH: kj or J

• ΔG: kj or J

• ΔS: kj/K

There are 4 possible combinations of terms with this formula

Ex1. Exothermic Reaction

• Within this rxn, entropy decreases as we go from 2 moles to 1 mol of gas

• Since ΔH is negative, this rxn is exothermic

• The enthalpy term predominates in this rxn

Ex2. Endothermic Reaction

• Even though the entropy increase would normally push the forward reaction to occur spontaneously at room temperature, the positive ΔG tells us otherwise

• Since ΔH is positive, this rxn is endothermic

   Thermochemistry

501-508: Sage

• Thermochemistry- is the study of the transfers of energy as heat that accompany chemical reactions and physical changes

Temperature and heat are related but not identical 

• Calorimeter- the energy absorbed or released as heat in a chemical or physical change is measured in a calorimeter

• Energy given off (or absorbed) during the reaction is equal to the energy absorbed (or given off) by the known quantity of water.

  • Amount of energy is determined from the temperature change of the known mass of surrounding water

• Data collected from calorimetry experiments are temperature changes because energy cannot be measured directly; but temperature, is directly measurable

• Temperature- is a measure of the average kinetic energy of the particles in a sample of matter. 

• The greater the kinetic energy of the particles in a sample, the higher the temperature is and the hotter it feels

• To assign numerical values to temperature, it is necessary to define a temperature scale. 

• For calculations in thermochemistry, we use the Celsius and Kelvin scales. 

  • Celsius and Kelvin temperatures are related by the following equation

K= 273.15 + ^oC

• The ability to measure temperature is thus based on energy transfer; measured in joules

• A joule is the SI unit of heat as well as all other forms of energy

N x m = kg x m2s2

• Heat can be thought of as the energy transferred between samples of matter because of a difference in their temperatures

• Energy transferred as heat always moves spontaneously from matter at a higher temperature to matter at a lower temperature

Energy transfer varies from reaction to reaction

• The quantity of energy transferred as heat during a temperature change depends on the nature of the material changing temperature, the mass of the material changing temperature, and the size of the temperature change

• EX: one gram of iron heated to 100.0^oC and cooled to 50.0^oC in a calorimeter transfers 22.5 J of energy to the surrounding water.

  • But one gram of silver transfers 11.8 J of energy under the same conditions.

  • Difference depends on the metals’ differing capacities for absorbing this energy.

• Specific heat is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree (1^oC) or one kelvin (1 K) (because the sizes of the degree divisions on both scales are equal)

  • Values of specific heat can be given in units of joules per gram per Celsius degree, J/(g x ^oC), joules per gram per kelvin, J/(g x K), or calories per gram per Celsius degree, cal/(g x ^oC)

• Specific heat is measured under constant pressure conditions, so its symbol, c_p, has a subscripted p as a reminder. In the equation, c_p is the specific heat at a given pressure, q is the energy lost or gained, m is the mass of the sample, and T represents the change in temperature

cp=qm x T

• This equation can be rearranged to give an equation that can be used to find the quantity of energy gained or lost with a change in temperature

Energy Lost or Gained     q=cp m T

Heat energy is transferred during a reaction

• Energy absorbed as heat during a chemical reaction at constant pressure is represented by H. The H is the symbol  for quantity called enthalpy. It is not practical to talk about enthalpy as a quantity, because we have no way to directly measure the enthalpy of a system

• Only changes in enthalpy can be measured. The Greek letter (“delta”) stands for “change in.” 

  • Therefore, H is read as “change in enthalpy”. 

• An enthalpy change is the amount of energy absorbed by a system as heat during a process at constant pressure H = Hproducts- Hreactants

• The enthalpy of reaction is the quantity of energy transferred as heat during a chemical reaction. (HEAT OF REACTION)

Enthalpy of Reaction in Exothermic Reactions

• If a mixture of hydrogen and oxygen is ignited, water will form and energy will be released explosively

• Energy that is released comes from the reactants as they form products

  • Because energy is released, the reaction is exothermic, and the energy of the product, water, must be less than the energy of the reactants. 

• EX: 2 mol of hydrogen gas at room temp are burned, 1 mol of oxygen gas is consumed and 2 mol of water vapor are formed

2H₂(g) + O₂(g) —> 2H₂O(g)

• Experiments have shown that 483.6 kJ of energy are evolved when 2 mol of gaseous water are formed from its elements at 298.15 K. 

• Modifying the chemical equation to show the amount of energy as heat released during the reaction gives the following expression.

2H₂(g) + O₂(g) —> 2H₂O(g) + 483.6 kJ

• This expression is an example of a thermochemical equation, an equation that includes the quantity of energy released or absorbed as heat during the reaction as written

• In  any thermochemical equation, we must always interpret the coefficients as numbers of moles and never as numbers of molecules. 

• The quantity of energy released as heat in this or any other reaction depends on the amount of reactants and products

• The quantity of energy as heat released during the formation of water from H₂ and O₂ is proportional to the quantity of water formed; require twice as many moles of reactants and would release 2x more energy as heat

4H₂(g) + 2O₂(g) —> 4H₂O(g) + 967.2 kJ

• Producing one-half as much water requires one-half as many moles of reactants and releases only one-half as much energy, or ½ x 483.6 kJ. 

H₂(g) + ½ O₂(g) —> H₂O(g) + 241.8 kJ

Enthalpy of Reaction in Endothermic Reactions

• The situation is reversed in an endothermic reaction– products have a larger enthalpy than reactants.

• Decomposition of water vapor is endothermic; it is the reverse of the reaction that forms water vapor

• In endothermic reactions, enthalpy now appears on the reactant side of the thermochemical equation but no changed value

   2H_2,O(g) + 483.6 kJ —> 2H₂(g) + O₂(g)

• The physical states of reactants and products must always be included in thermochemical equations because they influence the overall amount of energy as heat gained or lost. 

  • EX: energy need to decompose water would be greater than 483.6 kJ if we started with ice, becauze extra energy would be needed to go from ice to liquid and then to vapor

Thermochemical Equations

• Thermochemical equations are usually written by designating the value of H

• For exothermic reaction, H is always negative because the system loses energy.

2H₂(g) + O₂(g) —> 2H₂O(g)   H= -483.6kJ

• For endothermic reaction H is always positive because the system gains energy.

2H₂O(g) —-> 2H₂(g) + O₂(g)  H= +483.6kJ

• Since energy as heat is absorbed, the enthalpy of the reactants is lower than the final enthalpy of the products, and H is positive.

• When looking at all the thermochemical equations, consider the following. 

1. The coefficients in a balanced thermochemical equation represent the number of moles of reactants and products and never the number of molecules. They can be fractions when necessary

2. The Physical state of the product or reactant involved in a reaction is an important factor, and, therefore, must be included in the thermochemical equation

3. The change in Enthalpy represented by a thermochemical equation is directly proportional to the number of moles of substances undergoing a change. If 2 mol of water are decomposed, twice as much enthalpy

4. The value of the enthalpy change, H, is usually not significantly influenced by changing temperature

Enthalpy of formation is the energy change when elements form one mole of a compound

• Thermochemical data are often recorded as the enthalpies of such composition reactions. 

• The molar enthalpy of formation is the enthalpy change that occurs when one moles of a compound is formed from its elements in their standard state at 25⁰C and 1 atm. 

• Enthalpies of formation are given for the standard states of reactants and products; usually atmospheric pressure and room temp 298.15K

• To signify that a value represents measurements on substances in their standard states, a ^o sign is added to the enthalpy symbol, giving delta H for the standard enthalpy of a reaction.

• Adding a subscript f, as in Hfo further indicates a standard enthalpy of formation. 

508-514 (start from “exothermic compounds tend to be very stable”): Pham

Exothermic compounds tend to be very stable

• A compound with a large negative enthalpy of formation releases a large amount of energy as heat when it’s formed → stable

• Elements in standard states have ∆H⁰_f = 0

• The majority of enthalpies of formation are negative

• ∆H⁰_f of CO2 is -393.5 kJ/mol therefore, it is more stable than the elements from which it was formed

• Compounds with relatively positive or slightly negative values are unstable

• Ex: HI is a colorless gas that decomposes at room temperature 

  • Has an enthalpy of formation of +26.5 kJ/mol

  • As it decomposes, the violet iodine vapor, I₂, becomes visible through the container 

• Compounds with a high positive enthalpy of formation are very unstable and may react or decompose violently 

• Ex#1: C₂H₂ reacts violently with oxygen and must be stored in cylinders as a solution in acetone

• Ex#2: HgC₂N₂O₂ has a very large enthalpy of formation of +270 kJ/mol which makes it useful as a detonator for explosives

Enthalpy changes in combustion

• Combustion reactions produce energy in the form of light and heat when a substance is combined with oxygen 

• Enthalpy of combustion: the enthalpy change that occurs during the complete combustion of one mole of a substance

• Enthalpy of combustion is defined in terms of one mole of reactant, and enthalpy of formation is defined in terms of one mole of product

• All substances are in their standard states

• General enthalpy notation, ∆H, applies to enthalpies of reaction

• ∆H_c refers to the enthalpy of combustion

• A combustion calorimeter is a common instrument used to determine enthalpies of combustion

• A similar apparatus under constant pressure is used to obtain enthalpy measurements

Change in enthalpy is calculated using Hess’s Law

• Thermochemical equations can be rearranged and added to give enthalpy changes for reactions not included in the data tables

• Hess’s law: the overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process

• Energy difference between reactants and products is independent of the route 

• Measured enthalpies of reaction can be combined to calculate enthalpies of reaction that are difficult or impossible to actually measure

• Calculate the enthalpy of formation of methane gas from its elements, hydrogen gas and solid carbon (graphite), at 298.15 K: 

• In order to calculate the change in enthalpy, we use the combustion reactions of each element 

• Principles for combining thermochemical equations:

1. If a reaction is reversed, the sign of ∆H is also reversed

2. Multiply the coefficients of the known equations so that, when added together, they give the desired thermochemical equation. Multiply ∆H by the same factor as the corresponding equation.

• In this case, since methane is on the right side of the thermochemical equation, we must reverse the combustion equation of methane and change the sign of ∆H. This will turn the reaction to an endothermic one.

• Since we now have 2 moles of water as a reactant, we will need 2 moles of water as a product

• For the combustion of hydrogen, it only produces one mole of water so we would need to multiply everything by 2 

• Now add the three equations together 

• Hess’s law states that the enthalpy difference between reactants and products is independent of the pathway

• Any enthalpy of reaction may be calculated using enthalpies of formation for all the substances in the reaction of interest, without knowing anything else about how the reaction occurs

Enthalpy of formation is the sum of its sub-reaction enthalpies

• When carbon is burned in a limited supply of oxygen, CO is produced

• Carbon is first oxidized to CO₂ then part of it is reduced with carbon to give some CO

• These two reactions occur simultaneously so we get a mixture of CO and CO₂ 

• It’s not possible to directly measure the enthalpy of formation of CO(g_ from C(s) and O₂(g)

• However, we do know the enthalpy of formation of CO₂ and enthalpy of combustion of CO

• Reverse second equation because we need CO as a product

• This diagram is a model for a reaction that takes place in two steps

• If we plot the reactions based on their relative energy, you can see the relationship among the values for the enthalpy of formation of CCO

• The formation of CO₂ is at a level of -393.5 kJ/mol

• It shows the reverse of the combustion reaction (+293.0 kJ/mol) is added to that level

• The value of the formation of CO is -110.5 kJ/mol.