Physical

• Energy= capacity to do work
• work= process of moving against an opposing force
• work= force x distance
• Chemical reactions- including those taking place in a cell release or absorb energy

  Energy can only be transferred to another place or converted to another form- it cannot be created or annihilated
• Thermodynamics follow the heat, work and energy changes of the system

  This explains the reactivity and equilibrium- allowing predictions
• The First Law (energy)

  ISOLATED SYSTEM: the internal energy, U of an isolated system is constant

  CLOSED SYSTEM: the change in internal energy of a closed system is the heat that flows into the system, plus the work done on the system

  ∆U=q + w (integrated form)
• *q is positive (+) if heat flows from the surroundings to the system, negative (-) if heat flows from the system to its surroundings

  *w is positive (+) if work is done on the system, negative if work is done by it
• Heat and Heat capacity

  The internal energy depends on the temperature through its heat capacity (for a small change in temp of a system of fixed volume via heat dq (w=0))

  AT CONSTANT VOL: dU= Dqᵥ = cᵥdT
• Mechanical work at pressure, P

  work= force x distance = P.Adx = P.dV

  (P= pressure, dx= distance moved)

  Work is done by the system, against atmospheric pressure, so change in U due to work is negative (q=0):

  dU= dw = -PdV
• Enthalpy

  In space, external pressure P=0 (therefore -PdV=0 ~ no work can be done by expanding gas in free expansion)

  For an increase of volume of the system, work is done against the opposing pressure (i.e. on the surroundings)

  At constant pressure (dP=0): dH = dqₚ
• Standard enthalpy of formation= the enthalpy change on forming 1 mol of the substance in its standard state from its constituent elements in their reference states

  (The standard enthalpy of formation of a pure element in its reference state is 0)
• standard reaction enthalpy= the enthalpy change on reaction of 1 mol substance with all reactants and products in their standard states

  EXOTHERMIC: ∆H < 0

  ENDOTHERMIC: ∆H > 0
• Standard enthalpy of combustion= the standard reaction enthalpy for the complete combustion of the substance with oxygen

  (Typically a hydrocarbon reacting with oxygen to form CO₂, H₂O and heat)
• Hess' Law (∆ᵣH) = the net enthalpy change for an overall reaction is the sum of the enthalpy changes for the individual reactions

  ∆ᵣH = ∑∆բH (products) - ∑∆բH (reactants)
• Entropy = matter and energy to become more dispersed/ disordered
• Change in entropy, dS, is defined as the heat transferred to a system reversibly, divided by the absolute temperature:

  dS = dqᵣₑᵥ/ T
• Third Law of Thermodynamics= the entropy of a perfect crystal approaches zero as the absolute temperature approaches zero

  (entropy when determined from this point is called the absolute energy)
• The change in a state function (energy, U, enthalpy, H, entropy, S) depends only on the initial and final states, and not on the path between them
• Path Functions depend on the particular path that was followed in getting from the initial and final states (e.g. work (w) and heat (q))
• The entropy, S, of a system is directly related to the no of equivalent ways, W, in which a given overall state of a system could come around:

  S = Kᴮ lnW
• Second Law of Thermodynamics= entropy of the universe tends to increase
• A small change in a system of interest will lead to a total entropy change in the universe

  dSₜₒₜ = dS + dsₛᵤᵣᵣ

  (where dS is the entropy change of the system itself, and dₛᵤᵣᵣ is the entropy change of the surroundings)
• dSₛᵤᵣᵣ= - dH/ T
• dSₜₒₜ = dS - dH/T
• Gibbs Energy

  Gibbs Free Energy is the maximum non-expansion work that can be extracted from a process at constant temperature and pressure

  ∆G=∆H - T∆S

  Spontaneous change at constant temperature and pressure: ∆G >0

  Non-expansion work is work arising from any work other than the from expansion of the system

  ∆G = Wₘₐₓ, ₙₒₙ₋ₑₓₚ
• Chemical Kinetics is the study of:

  *measures consumption of reactants/ formation of products

  *effects of pressure, temp + catalysts

  *mechanisms of reactions

  *identify the sequence of elementary steps
• Stoichiometry = numbers of reactant and product molecules

  (studies how the concentrations change over time, where T is held constant)
• Experimental Techniques to measure concentrations:

  1. pressure measurement (for gaseous species)
  2. spectrophotometry (very fast)
  3. conductivity measurement (ionic solutions)
  4. pH measurements (for reactions involving H+/OH-)
  5. polarimetry (for chiral species)
• Rate of reaction = number of moles of reaction per unit time
• Instaneous Rate

  rate of reaction = k[A][B]

  Units= dm^3 mol^-1 s^-1
• Order of reaction = power to which the concentration is raised

  e.g. 1st order: rate = k[A][B]

  2nd order: rate = k[A]^2[B]
• Zero Order: rate independent of concentration

  ROR = [A]^0 = 1
• Fractional Order

  ROR = k[A]^1/2[B]
• Negative Order

  ROR =K[A][B]/[C]
• Isolation Method for finding the Order

  *reactants are kept constant except 1

  Although the true law might be: rate=[A][B]^2; we can approximate [B] by its initial value [B]₀, so that:

  rate=(k[A]²₀)[B] = k'[A], where k'=dₑf k[B]²₀

  *This is now a pseudo first order reaction
• Initial Rate method for finding order

  *measurements made before concentrations have changed

  *order found by plotting logarithm of initial rate, r₀, against the logarithm of concentration
• Arrhenius Equation. $k=$$Ae^-Ea/RT$
• Arrhenius Equation with the Collision Theory

  Pre-exponential factor: how frequently molecules collide in the right orientation to react

  Activation energy Ea: minimum collision energy needed for this rearrangement

  e^-Ea/RT is the Arrhenius Law: fraction of collisions with the necessary energy
• $lnK=$$lnA-Ea/RT$

  link against 1/T

  Intercept= lnA

  Gradient= -Ea/RT
• Use:

  ln K'/K = Ea/R(1/T - 1/T')

  to find rate k' at a different temperature, T