Unit C

ΔH_rxn= ? - ?

ΔH_rxn= ? - ?

H(products) - H(reactants)

If ΔH is large and negative….

the reaction is THERMODYNAMICALLY FAVORABLE

if ΔH is large and positive….

the reaction is NOT thermodynamically favorable

If the energy of the products is GREATER than the energy of the reactants….

the reaction is ENDOthermic

If the energy of the products is LESS than the energy of the reactants….

the reaction is EXOthermic

Exothermic reactions feel….

Hot

Endothermic reactions feel…

Cold

1 pair of electrons shared; longest and weakest bond

Single bond

2 pairs of electrons shared; shorter and stronger bond

Double bond

3 pairs of electrons shared; shortest and strongest bond

Triple bond

Breaking of bonds is…

Endothermic (+ΔH)

Forming of bonds is…

Exothermic (-ΔH)

In an exothermic reaction: BDE of reactants is _____ than BDE of products

GREATER (products are lower than reactants on the graph)

In an endothermic reaction: BDE of reactants is _____ than BDE of products

LESS (products are higher than reactants on the graph)

BDE to estimate ΔH_rxn:

ΔH_rxn= ???

ΔH_rxn = (BDE reactants) + (BDE products)

REMEMBER: BDE is ADDING

Multiply ____ by the same number the ____ are multiplied by when balancing equations for Hess’s Law

the ΔH value, coefficients

If something is mixing/dissolving, it is…

Thermodynamically favorable (even without added stirring)

What are IMFs?

Intermolecular Forces; attractions between molecules

• tells if the substance is mixable or not

• weaker than bonds

Dispersion Forces

• present in all molecules and atoms

• weakest IMF

• randomly caused by electron movement in orbital

• if electrons gather on one side of the molecule, they can create an instantaneous dipole

Dipole-Dipole

• negative end of 1 polar molecule is attracted to the positive end of another polar molecule

• stronger IMF than Dispersion Forces because electrons always gathered on that side (permanent Dipole rather than instantaneous)

Hydrogen Bonding

• found in polar molecules with Hydrogen attached to Florine, Oxygen, or Nitrogen

• stronger than Dispersion Forces and Dipole-Dipole forces because F, O, and N are more electronegative (chemlab coming in clutch)

Ion-Dipole

• cations attracted to the negative end of dipole

• anions attracted to the positive end of dipole

• stronger than Dispersion, Dipole-Dipole, and Hydrogen Bonds

• strength increases as charge increases

Lattice Energy

Bonds/IMFs between solute particles of an ionic crystal are broken

• ENDOTHERMIC REACTION (breaking bonds=endothermic)

If IMFs between solvent molecules broken, then…

the reaction is ENDOthermic (breaking bonds=endothermic)

IMFs between solute and solvent particles FORMED

reaction = EXOthermic (forming bonds=exothermic)

Hydration

• Energy IN to break apart crystal structure

• Energy IN to separate Hydrogen Bonds between H2O molecules

• Energy OUT as cations and anions attract to the water molecules

ΔH_solute=

Bonds/IMFs breaking

ΔH_solvent=

IMFs/Bonds breaking

ΔH_mix=

IMFs forming

ΔH_soln=

ΔH_solute + ΔH_solvent + ΔH_mix

If ΔH_mix > ΔH_solute + ΔH_solvent then

yes dissolve

If ΔH_mix < ΔH_solute + ΔH_solvent then

no dissolve

Experimental ΔH_soln calculation:

q = mcΔT = ΔH * moles

Experimental ΔH_comb calculation:

q = mcΔT = -ΔH * moles

Definition of ΔH

Enthalpy change; heat transferred between system and surrounding under constant pressure

0th Law of Thermodynamics

2 systems in equilibrium with the 3rd system are in thermal equilibrium with each other

1st Law of Thermodynamics

Energy is neither created nor destroyed (conservation of energy)

2nd Law of Thermodynamics

Entropy of an isolated system always increases (hot —> cold)

• ΔE = q + w

  • ΔE: change in internal energy

  • q: heat absorbed

    • if q>0 then sys gaining heat from surr

  • w: work

    • if w>0 then sys gaining work from surr

3rd Law of Thermodynamics

Entropy of a system approaches a constant as temperature approaches 0K (absolute zero)

ΔH units

kJ/mol

Entropy (S) and ΔS units

J/molK

If ΔS is positive, then…

Entropy is INCREASING

• products are more randomly arranged than reactants

If ΔS is negative, then….

Entropy is DECREASING

• reactants are more randomly arranged than products

When determining entropy, look at…

1. Phase (gas has highest entropy)

2. Number of molecules (more molecules means more entropy)

3. Molecule size/shape (larger, more “floppy” molecules have greater entropy)

4. Larger atoms (larger atoms have more entropy)

5. Temperature (higher temperature means greater entropy)