IBCH4 definitions

thermodynamics

• law of cons. of energy → E can be converted from 1 form to another but not destroyed/created (total E remains constant)

• heat (Q) in (J) → form of thermal E, transferred between objects because of temp. diff. 

  • conduction, convection, radiation

• temp. (T) in (K) → measure of avg. kinetic energy of particles

  • how much heat has been transferred → measured by temp. change

• heat capacity (c) in (J/kg K) → amt. of E needed to raise temp. of substance by 1 degree

  • specific heat capacity → amt. of heat E needed for 1 unit of mass of substance to cause inc. of one unit in temp.

open/closed/isolated systems

• open → allows transfer of heat/matter in/out of system

• closed → allows transfer of only heat 

• isolated → doesn’t allow transfer of heat/matter

endo/exo

• exothermic reaction gives out heat, warming mixture (heat lost to surroundings) (ΔH<0)

• endothermic reaction takes in heat, cooling mixture (heat gained from surroundings) (ΔH>0)

standard enthalpy change of combustion/formation

• standard enthalpy change of combustion (∆H_c) is heat E released when 1 mol of pure substance is completely burnt in excess O₂ under standard conditions

• standard enthalpy change of formation (∆H_f) is heat change at constant production of 1 mol of pure substance from its elements in standard states under standard conditions

• can see these values in data booklet + when writing reaction equation can use fractions

measurement uncertainty in calorimetry

• heat transfer due to calorimeter not being 100% insulated

• reactions rarely immediate → heat lost while reaction takes place

• heat absorbed by calorimeter (e.g. can)

• standard heat capacity not exactly same for sol. and water

• experiment not under standard conditions

fossil fuels

• crude oil → HC + sulfur, nitrogen etc, refined to petroleum etc

• coal → combustible sedimentary rock with high amt. C, HC

• natural gas → naturally occurring mixture of gaseous HC (mainly methane, little bit of other higher alkanes)

• cheap/efficient, found in large amounts, reliable, infrastructure exists alr

• non-renewable, produces GG → global warming, pollution (NO_x, SO_x), isn’t found everywhere

biofuel

• = fuel produced over short timespan from biomass (renewable E source)

• biological carbon fixation → production of organic compounds from CO₂ (photosynthesis)

• bioethanol → ethanol produced by fermentation of glucose/biomass, used as fuel/fuel additive

• biodiesel → produced in transesterification reaction

• renewable, reduced GG, sustainable (many plant mats+waste can be used), economic security (reduced dependency on imported oil)

• high cost, possible deforestation as demand inc., monoculture may result in reduced biodiversity, food products diverted to produce biofuel, use of agricultural land/water/fertiliser/pesticides for growing crops

specific energy

• = amt of heat E released per mass of fuel (kJ/kg) (enthalpy/molar mass)

• larger HC → lower volatility bc stronger dispersion forces

  • affects how HC interact with oxygen + type of combustion 

  • longer HC chain → greater tendency of fuel to undergo incomplete combustion, which releases poisonous CO and produces less specific E when compared to same HC’s complete combustion

collision theory

chem. reaction only occurs when → collisions between molecules have enough E to break bonds in reagents, molecules collide w/ proper orientation, new bonds form

• change in concentration (c) or amount (n) of reactant/product with time (t)

  • r/v = ∆/c∆t (absolute value) in  mol/dm3 s

  • if half mass → half volume produced

  • if less V but same c → slower but reaches same volume

• avg rate measures over particular time interval (usually slower)

• instantaneous rate measures at particular time from gradient (e.g. initial, 10hr…)

factors affecting RoR

• concentration / mixing

  • mixing or inc. c inc. freq. of collisions between particles → greater RoR 

• pressure

  • inc. pressure forces particles closer tog. → inc. collision rate → inc. RoR (only gases)

• temperature

  • inc. in temp. → inc. velocity of particles → greater E_k → more collisions → greater RoR

• catalyst (= inc. RoR but remains chemically unused at end of reaction)

  • provide alt. pathway for reaction w/ lower E_a

    • temp. has no effect on E_a. slower tend to have greater E_a

• light

  • VIS/UV radiation breaks bonds in reactant molecules (some rates can be inc.)

  • silver halides…are photosensitive + undergo partial decomp. in light

• particle size

  • dec. size inc. SA of solid substance → more freq. collisions → greater RoR

dynamic equilibrium + equilibrium constant

• only in closed systems!

• macroscopic properties don’t change → amt./concentrations/densities of reactants/products same

• at microscopic lvl → particles continue to take part in forward/reverse reaction (rates of these equal)

• aA + bB ⇌ cC + dD

• K_c (equilibrium constant) expression  = [C]^c[D]^d/[A]^a[B]^b is constant at given temp. 

  • [] = concentration

  • only for homogeneous equilibrium!

  • if K_c<1 → equilibrium at reactant’s side (more reactants)

  • if K_c=1 → equal amt. of reactants/products

  • if K_c>1 → equilibrium at product’s side (more products)

• deduce K_c equilibrium, use it to find expression for reactions w/ diff. coefficients

le chatelier’s principle

arrhenius + BL acids/bases

Arrhenius acids and bases:

• acid → any species that inc. concentration of hydrogen ions/protons in (aq) sol.

• base → any species that inc. concentration of hydroxide ions in (aq) sol.

• doesn’t take into account reactions w/o water

Bronsted-Lowry acids and bases

• acid → any species capable of donating a proton

• base → any species capable of accepting a proton

• conjugate acid of BL base → species formed after base accepts a proton

• conjugate base of BL acid → species formed after acids donates a proton

  • e.g. NH₃ + H₂O → NH₄⁺ + OH⁻ 

strong/weak acids/bases

• strong acids/bases dissociate/ionise completely (irreversible)

  • result in weak conjugate bases/acids

  • strong acids → HCl, HNO₃, H₂SO₄ (diprotic acid, donates 2H), HI (hydroiodic), HBr (hydrobromic)

  • strong bases → group 1 hydroxides e.g. NaOH, KOH

• weak acids/bases dissociate/ionise partially (equilibrium)

  • result in stronger conjugate bases/acids

  • weak acids → organic acids, salicylic acid, ethanoic acid

  • weak bases → NH₃ (ammonia), amines e.g. aminoethane, CO₂2- (carbonates), HCO₃- (hydrogen carbonates)

• if dissociation not complete → less OH ions produced

  • check pH, conductivity, amt. of H₃O ions produced to distinguish strong/weak