Class2-notes.pdf - Chem 202 u2013 class 2 Introduction to...

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Chem 202 – class 2 Introduction to Free Energy Definition of enthalpy, H Bond dissociation enthalpy – Use in estimating ∆H of a reaction What does it mean for a reaction to be spontaneous? Concept of entropy, S Qualitatively determine the sign of an entropy change ∆S Free energy change ∆G = ∆H – T∆S determines spontaneity (must be negative) ∆H negative AND/OR ∆S positive Spontaneous does not equal fast!
A reaction is a change in the way atoms or molecules are bonded together Bonds hold a molecule together – bonded state is one of lowered potential energy (PE) (If the PE of a pair of isolated atoms is defined as zero, the energy of the bonded molecule is negative ) Atoms exert forces on each other, pulling or pushing to reach an arrangement with lower P.E.
Raising the P.E. to break a bond requires an input of energy from somewhere else: kinetic energy . Temperature is a measure of K.E. at the molecular level. * Heat is a transfer of energy in the form of molecular-level (randomized) K.E. If heat energy is added to a system (a collection of molecules), where might it go? *This is a good working definition for this class. there are cases where K.E. and temperature are not so clearly correlated due to quantum mechanical effects. Increases K.E. of system (raising temperature) Converts to P.E., breaking bonds and/or deforming them away from minimum P.E. lengths and angles Converted to work against an external pressure, making system expand Changes to KE and PE add up to change in internal energy ”U” Energy lost as work at constant pressure equals (pressure × change in volume)
Heat entering a system (at constant pressure, with no exchange of energy other than work of expansion/compression) equals the change in enthalpy of the system Breaking bonds increases enthalpy,

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