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Energetics. 5.1 Endothermic and Exothermic Reactions. Endothermic rxn  heat is taken in from the surroundings ( rxn vessel gets cooler) Exothermic rxn  rxn that result in the release of heat ( rxn vessel gets warmer)

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5 1 endothermic and exothermic reactions
5.1 Endothermic and Exothermic Reactions

  • Endothermicrxn heat is taken in from the surroundings (rxn vessel gets cooler)

  • Exothermic rxn rxn that result in the release of heat (rxn vessel gets warmer)

  • Enthalpy change amount of heat energy taken in/ given out in a chemical rxn (ΔH)

    • Positive for endo. Rxn

    • Negative for exo. rxn


  • Exo. produces a more stable product

    • Ex. Cdiamond CgraphiteΔH= -1.9 kJ mol-1

  • Kinetic vs. Thermodynamic stability

    • Graphite is more thermodynamically more stable than diamond

    • Diamond is kinetically stable

    • What does this mean?

Activation energy
Activation Energy

Define Activation Energy

Activation energy1
Activation energy

  • High AE= faster or slower rxn?

  • Can the conditions be altered?

  • If yes, in what ways?

  • Does how endo- or exothermic a rxn is tell us how quickly the rxn will run?

    • Ex. Which rxn will run faster?

      • -52 kJ mol-1

      • -16 kJ mol-1

  • Do endo. or exo. rxns usually occur spontaneously? (under normal conditions)

5 2 calculations of enthalpy changes from experimental data
5.2 Calculations of enthalpy changes from experimental data

  • Specific heat capacity (c) energy required to raise the temp. of 1 g of substance by 1 K (1°C) or, the energy to raise 1 kg of substance by 1 K.

    • Units-> J g-1 K-1

      J g-1 °C-1

      kJ kg-1 K-1

      J kg-1 K-1

Calculating c
Calculating c

  • q= mcΔT

    • q= heat energy

    • m= mass

    • ΔT= change in temperature

  • How does heat capacity affect how easily a substance can be heated?

  • Can c be calculated for a substance undergoing cooling?

Measuring enthalpy change of combustion rxns
Measuring enthalpy change of combustion rxns

  • Worked ex. Page 185

  • Literature value for ΔH ethanol is -1371 kJ mol-1

  • What are some reasons the calculated value is different?

  • A bomb calorimeter could have been used so that the system was heavily insulated and provide a plentiful supply of oxygen

Enthalpy changes in solution
Enthalpy changes in solution

  • General method for measuring

    • Measure known amounts of reagents

    • Record initial temps

    • Mix in a polystyrene cup

    • Record max/ min temperatures observed

  • Assume that c for the final solution is the same as water

H of solutions definitions
ΔH of Solutions:Definitions

  • Enthalpy change of neutralisation (ΔHn) enthalpy change when 1 mol of water molecules are formed when acid reacts with alkali under standard conditions

    • H+(aq) + OH-(aq)  H2O(l)

  • Enthalpy change of solution (ΔHsol) the enthalpy change when 1 mol of solute is dissolved in excess solvent to form a solution of ‘infinite dilution’ under standard conditions

    • NH4NO3(s) NH4+(aq) + NO3-(aq)

5 3 hess s law
5.3 Hess’s Law

  • The enthalpy change accompanying a chemical reaction is independent of the pathway between the initial and final states

  • What does this mean?

  • Ex. Find ΔHr for the reaction of AB

    • Knowns:

      • ΔHr= ΔH1 +ΔH2

      • AC =ΔH1

      • BC= ΔH2

      • What is C?

      • What needs to change about the BC step?

Hess s law definitions
Hess’s law:Definitions

  • State function pathway does not matter

  • Standard conditions pressure= 1 atm (or, 1.01E5 Pa), 298K (or, 25°C)

  • Standard enthalpy change (ΔHrΘ) the enthalpy change when molar amounts of reactants as shown in the stoichiometric equation react together under standard conditions to give products (Θ= under standard conditions)

Working out enthalpy changes
Working out enthalpy changes

  • Hess’s law can be used to determine enthalpy changes of unknows from knowns

  • P. 194 Worked example

5 4 bond enthalpies
5.4 Bond enthalpies

  • The enthalpy change when 1 mole of covalent bonds, in a gaseous molecule, are broken under standard conditions (aka bond energy)

  • Ex. Then enthalpy of H-H bond is 436 kJ mol-1

    H2(g)  2H(g) ΔHΘ= +436 kJ mol-1

    How many H-H bonds were broken?

Bond enthalpy
Bond Enthalpy

  • What state must a substance be in to calculate bond enthalpy?

  • Consider this process:

    Br2(l)  2Br(g) Br-Br= 193 kJ mol-1

    What is the ΔHΘ? +224 kJ mol-1

    Why is this higher than the bond enthalpy?

    The reactants are not in a gaseous state

    We must also account for the energy required for vaporisation of the reactants

    This process is called atomisation

Bond breaking
Bond breaking

  • Bond breaking is…

    endothermic or exothermic?

    Endothermic! What does that mean about ΔH?

    Positive! What will bond making be?

    Exothermic with a negative ΔH

Average bond enthalpy
Average bond enthalpy

  • The average amount of energy required to break 1 mole of covalent bonds, in a gaseous molecule under standard conditions

  • These are the values used to calculate bond enthalpies

Using bond enthalpies to work out enthalpy changes in a rxn
Using bond enthalpies to work out enthalpy changes in a rxn

  • Must draw out the structural formulas for rxn

  • Imagine the rxn happening and ALL bonds being broken

    • Add up the total energy for all broken bonds

  • Draw in all the bonds formed in products

    • Add up the total energy of all bonds made

  • Determine signs for the total enthalpy changes

    • Broken positive

    • Made negative

  • Add the changes to get the overall enthalpy change of the rxn


  • Consider the rxnbetween ethene and bromine, to produce 1,2-dibromoethane,

    C2H4(g) + Br2(g) C2H4Br2(g)

    What bonds are broken?

    What bonds are made?

    Follow your steps!

Using a cycle
Using a cycle

  • Same concept as previous calculations, but a process is drawn out to see all the steps

  • P. 203

5 5 calculating enthalpy changes definitions
5.5 Calculating enthalpy changes:Definitions

  • Standard enthalpy change of combustion (ΔHcΘ) the enthalpy change when 1 mole of a substance is completely burnt in oxygen under standard conditions.

    • If ΔHcΘis always negative, what does this mean?

  • Standard enthalpy change of formation (ΔHfΘ) the enthalpy change when 1 mole of the substance is formed from its elements in their standard states under standard conditions

    • Endo and exorxns are dependent on the type of substance

    • ΔHfΘ for any element in its standard state is zero

Using h c to calculate enthalpy change
Using ΔHcΘto calculate enthalpy change

  • Method 1: Construct an enthalpy cycle

    • P. 208

  • Method 2: rearrange the equations to give the overall equations related to the enthalpy change

    • P. 209

  • Method 3: use an enthalpy level diagram for calculations

    • P. 210

  • Method 4: use the equation,

    ΔHr = ΣΔHc (reactants)- ΣΔHc(products)

Using h f to calculate other enthalpy chages
Using ΔHfΘto calculate other enthalpy chages

  • Method 5: similar to method 1, but used for formation rather than combustion

    • P. 214

  • Method 6: refer method 2 (be sure equations are running in the correct direction)

  • Method 7: draw enthalpy level diagram for formation (method 3)

  • Method 8: use the equation,

    ΔHr = ΣΔHf (products)- ΣΔHf (reactants)

Choosing your method
Choosing your method

  • Choose a method based on the data you are given, NOT on what needs to be found

  • If needing the enthalpy of combustion and given the enthalpy of formation, use one of the methods 5-8

  • Once the basic principle of the methods are understood, there is no need to have any distinctions between them

5 6 enthalpy changes for ionic compounds
5.6 Enthalpy changes for ionic compounds

  • First ionisation energy

  • Second ionisation energy

  • First electron affinity enthalpy change when one electron is added to each atom in 1 mol of gaseous atoms under standard conditions (always EXOTHERMIC)

    X (g) +e-  X- (g)

  • Second electron affinity (always ENDOTHERMIC) why?

  • Lattice enthalpy(ΔHΘlatt) the enthalpy change when 1 mol of an ionic compound is broken apart into iest constituent gaseous ions under standard conditions

Born haber cycles
Born-Haber cycles

  • Enthalpy level diagram breaking down the formation of an ionic compound into a series of simpler steps

  • 1. put the equation for the enthalpy of formation

  • 2. add lattice enthalpy

  • 3. convert to gaseous form (why?)

    • Two steps

  • Must convert ALL reactants to gaseous form

  • Connect the cycle by adding the electrons removed from one reactant to the more electronegative reactant

Draw a born haber cycle
Draw a Born-Haber cycle

  • P. 219-221

  • Na and Cl example

Comparisons of lattice enthalpy
Comparisons of lattice enthalpy

  • P. 224

  • What is lattice enthalpy the result of?

    Electrostatic attractions of + and – ions

  • If the attractions of great, will more or less energy need to be supplied to break the bonds?


Effect of charge and size
Effect of charge and size

  • How does the charge of the ions effect lattice enthalpy?

    The higher the ion charge, the greater the lattice enthalpy

  • Does NaCl or MgCl2 have great lattice enthalpy?


  • How does size effect lattice enthalpy?

    The larger the ions the weaker the forces, the smaller the lattice enthalpy

  • Which has the larger lattice enthalpy, CsCl or NaCl?


Theoretical vs experimental
Theoretical vs. experimental

  • Theoretical assumes a totally ionic model

    • What is this?

      Bonding is solely due to attractive forces between oppositely charged ions

  • Experimental use the Born-Haber cycle to find

  • These are compared to determine how ionic a particular compound is

How to use theoretical and experimental values
How to use theoretical and experimental values

  • If values are exactly the same, complete ionic bonding is suggested

  • If values are significantly different, it is suggested that the bonding has a significant degree of covalency

  • Ex. Silver iodide

    Theoretical value/ kJ mol-1 736

    Experimental value/ kJ mol-1 876

    What do the values suggest?

Covalent character
Covalent character

  • What is covalent character the result of?

    Polarisation of the negative ion by the positive one

  • How does size of the anion effect this?

    The polarisation effect is greater

5 7 en t r o p y
5.7 ENtropY

  • A measure of randomness or disorder of a system

  • Especially significant in the case of endothermic processes occurring at standard conditions (ice melting at room temp, water evaporating, NaCl dissolving in water etc)

  • Endo rxns can only occur if there is an increase in entropy

  • Represented by S

  • Units: J K-1 mol-1

Standard entropy
Standard entropy

  • Represented by SΘ

  • Positive ΔSΘ indicates increased entropy

    • Less order

  • Negative ΔSΘ indicates decreased entropy

    • More order

Predicting sign of entropy change
Predicting sign of entropy change

  • Which state of matter has the higher entropy?


  • Which state has the least entropy?


Predicting sign
Predicting sign

  • Must consider whether the system’s disorder increases or decreases

  • Good to consider whether moles of gas have increased or decreased

  • What would an increase in moles of gas mean?

  • If moles of gas remain constant, our prediction of a change in entropy would be approximately zero

Calculating entropy change for a rxn
Calculating entropy change for a rxn

ΔSΘ= ΣSΘproducts -- ΣSΘreactants

5 8 spontaneity
5.8 Spontaneity

  • Spontaneous reaction one that occurs without any outside influence

Predicting spontaneity
Predicting spontaneity

  • A reaction being spontaneous does not mean it will run quickly!

  • Whether a rxn is spontaneous or not under a certain set of conditions can be deduced by looking at the change in the “entropy of the Universe”

    ΔSUniverse = ΔSsuroundings + Δssystem

    If ΔSUniverseis positive, the entropy of the universe increases and the rxn occurs spontaneously

    When heat is given out in a rxn, the entropy of the surroundings get hotter

Gibbs free energy
Gibbs free energy

  • Represented by ΔG

    • Also called free energy change

  • ΔG = ΔH – TΔS

  • ΔH and ΔS are referring to the system

  • Under standard conditions this symbol is used: ΔGΘ

  • For a reaction to be spontaneous, ΔG for the rxn must be NEGATIVE

  • Units kJ mol-1

Calculating g
Calculating ΔG

  • Method 1 use ΔGΘ = ΔHΘ – TΔSΘ

  • Method 2 use ΔGΘ = ΣΔGfΘ(products) – ΣΔGfΘ(reactant)

  • Standard free energy of formation the free energy change for the formation of 1 mol of substance from its elements in their standard state under standard conditions

  • Temp must be in K

  • If no temp is given and it is standard conditions, assume 298K

Non spontaneous rxns
Non-spontaneous rxns

  • If a rxn is non-spontaneous, does that mean it will never happen?


  • What will make it run?

    Outside influence such as; temp, catalysts, etc.

Effects of temp on spontaneity
Effects of temp on spontaneity

  • Refer to the equation:

  • ΔG = ΔH – TΔS

  • If ΔS is positive, temp must be high or low to be spontaneous?

  • High

  • ΔG must be negative to be spontaneous, thus TΔS must be higher than ΔH

Temp and spontaneity
Temp and spontaneity

  • If ΔS is negative --TΔS will be positive and the rxn cannot be spontaneous

  • Endo rxns will only occur spontaneously in entropy is increased and temp is significantly high

  • Exorxns will always be spontaneous at some temp

  • If the rxn involves a decrease in entropy the rxn will be spontaneous at a lower temp; becomes less spontaneous as temp increases