CHEM1020 Module 2 - Chemical Balance

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Gwen Lawrie

CHEMICAL BALANCE HETEROGENEOUS EQUILIBRIA Heterogenous equilibria: reactions where reactants and products exist as more than one phase  Equilibrium constant does not include pure solids or liquids as these are constants  Pressure is not affected by adding more liquids or solids STATES OF MATTER Phase change: transition of substance from one phase to another (no chemical reaction)  Depend on temperature, pressure and magnitudes of intermolecular forces  Phase changes require that energy is supplied or removed from the substance undergoing the change Molar enthalpy of fusion: heat required to melt 1 mol of substance at its normal melting point Solid  liquid Molar enthalpy of vaporisation: heat required to vaporise 1 mol of substance at its normal boiling point Liquid  gas  I.e. sweat (an endothermic process) has a cooling effect Molar enthalpy of sublimation: heat required to vaporise 1 mol of substance from solid phase Solid  gas REVERSE CHANGES (ΔH < 0) Solidification: liquid  solid Condensation: Gas  liquid Deposition: Gas  solid PHASE DIAGRAMS High temperature + low pressure = gas High temperature + high pressure = liquid Low temperature + high pressure = solid Boundary lines: neighbouring phases are in equilibrium Triple point: where all three phases are in equilibrium Supercritical fluid: form upon compression of gases at high temperature or heating a liquid to high pressure  At critical point the densities in the gas phase and liquid phase become equal VAPOUR PRESSURE The number of molecules of a liquid that have enough energy to escape into the vapour phase depends on:  The strength of intermolecular forces (stronger bonds = less volatile)  The temperature (higher temperatures = more molecules with sufficient energy to escape) CLOSED CONTAINER Vapour pressure: the pressure at which dynamic equilibrium is achieved in a closed container No. Molecules leaving the surface (evaporation) = no. Molecules re-entering the liquid (condensing) P = Pvap OPEN CONTAINER When the vapour pressure of the liquid in an open container equals the external pressure, the liquid starts to boil (normal boiling point) Occurs when P = P vap TEMPERATURE DEPENDENCE OF P vap Substitute K for vapP (using equilibrium constant for vaporisation of liquid) Can be used for liquid  gas or solid  gas At boiling point T = ΔH/ ΔS GIBBS FREE ENERGY Two gases will always mix spontaneously as it leads to an increase in entropy relative to the unmixed gases  Has negative Δmix Total Gibbs energy change for system = Gibbs energy change for chemical reaction + Gibbs energy change for mixing Magnitude and sign of ΔG indicate whether an observable spontaneous reaction occurs  Lowering Gibbs energy = spontaneous reaction ΔrG tells us about the composition at equilibrium (not the spontaneity)  Positive r G means no observable products – reaction will not go  Negative ΔrG means no observable reactants – reaction will proceed in forward direction THE RELATIONSHIP BETWEEN Δ G AND Q r ΔrG = ΔrG + RTlnQ With p for gasec, Q for solutions And Δ G is the standard Gibbs energy for reaction r ΔrG = ΔrG+ RT lnQ and ΔrG = -RT lnK: ΔrG = RT ln(Q/K) ΔrG = 0 Q/K = 1 (Q = K) ΔrG < 0 Q/K < 1 (Q < K) ΔrG > 0 Q/K > 1 (Q > K) SOLUBILITY Solution: a homogenous mixture of two or more pure substances  Liquid solutions contain solvent + solute (gas, liquid or solid) Solvation: when a solute molecule is surrounded by solvent molecules Hydration: when solutes become surrounded by water molecules Saturated: when a dynamic equilibrium is set up between ions in solution and the solid Solubility: molar concentration of a saturated solution at a particular temperature GAS-LIQUID SOLUTIONS For gas to dissolve in a liquid, gas molecules must be able to be dispersed evenly throughout solvent  Intermolecular forces between solvent molecules are strong and thereforsol H (enthalpy) is large  Δsolis enthalpy change when dissolving 1 mole of solute in 1L solvent Gas dissolving in organic solvent = endothermic Gas dissolving in water = exothermic Temperature increases  solubility decreases Henry’s Law* Pressure increases  more soluble C g K Ph g Where C is concentration of gas P gass partial pressure of gas above solution (Pa) K is Henry’s law constant (mol Pa ) h -1 -1 Partial pressure of N in air is 7.8x10 Pa 2 *Henry’s Law does not hold for reactions! SOLID-LIQUID SOLUTIONS A salt can dissolve endothermically or exothermically in water  Overcoming lattice enthalpy Temperature increases  solubility increases (generally) Supersaturation: prepared by dissolving a solid at high temperature and then slowly cooling this solution SOLUBILITY OF SALTS All ionic solids are soluble to some extent  Some are very sparingly soluble i.e. AgBr Ionic compounds and polar molecules can dissolve in water when: Solute-solvent attractions > water dipoles and solute attractions Dissolution of NaCl in water: NaCl  Na + + Cl- (s) (aq) (aq) When excess salt is present a heterogenous equilibrium is established NaCl (s)<==> Na +(aq)+ Cl-(aq) K spsolubility product (same as equilibrium constant with anions and cations)  Concentration of one ion multiplied by concentration of another  Larger K sp= more soluble Molar solubility (s) = molar concentration of a salt in its saturated solution If molar solubility of AgBr is 6.7x10 mol L -1 + - -7 -1 Then s = [Ag ] = [Br ] = 6.9x10 mol L And K =spAg ][Br ] = s = 4.8x10 -13 USE ICE TABLE! SOLUBILITY RULES  All Na , K and NH 4+ compounds are soluble -  All nitrates (NO 3 are soluble  Most chlorides (Cl ) are soluble (except Ag, Hg and Pb)  Most sulphates (SO 42-) are soluble (except Sr, Ba, Pb)  Most carbonates (CO 2-)are insoluble (except group 1A & NH ) + 3 4  Most hydroxides (OH ), oxides (O ) & sulphides (S ) are insoluble (except group 1A & NH ) 4+ PREDICTING PRECIPITATION Same as reaction quotient (Q) Once Q passes K spprecipitation will occur If Q > K spsolid will precipitate until Q = K sp If Q =K sp,he solution is saturated and is at equilibrium; not precipitate will form If Q
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