CHEM1020 Module 2 - Chemical Balance

6 Pages
52 Views
Unlock Document

Department
Chemistry
Course
CHEM1100
Professor
Gwen Lawrie
Semester
Spring

Description
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
More Less

Related notes for CHEM1100

Log In


OR

Join OneClass

Access over 10 million pages of study
documents for 1.3 million courses.

Sign up

Join to view


OR

By registering, I agree to the Terms and Privacy Policies
Already have an account?
Just a few more details

So we can recommend you notes for your school.

Reset Password

Please enter below the email address you registered with and we will send you a link to reset your password.

Add your courses

Get notes from the top students in your class.


Submit