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R1 What drives chemical reactions?, R1.1 - MEASURING ENTHALPY CHANGE (ΔH)…
R1 What drives chemical reactions?
R1.2 - ENERGY CYCLES IN REACTIONS
Born Haber Cycle
an enthalpy cycle that relates the lattice enthalpy and enthalpy formation of an ionic solid using a series of one step
the standard enthalpy of formation of the ionic compound,∆ Hf
the ionization energy of the metal, ∆HIE
the electron affinity of the non-metal, ∆HEA
the standard enthalpy of atomization of a constituent element, ∆Hat
the standard lattice enthalpy of the ionic compound, ∆Hlat
HESS LAW
ALGEBRAIC
METHOD
DIAGRAM METHOD
Write the equation
Write the other species from equation 1,2 and 3 that are not included in the original equation. (Write at bottom of equation)
Draw two arrows from reactant in equation to the other species at bottom to represent equation 2 and 3.
Draw an arrow from products to species at bottom to represent equation 3.
Label the arrows with enthalpy change values.
Calculate the sum of enthalpy change value.
If the arrow are in opposite direction from the pathway, reverse the sigh of enthalpy change value.
Bond-breaking and bond-forming
Bond-break : endothermic process (energy needed to break chemical bond)
Bond-forming: exothermic process (energy released when bond are formed)
Bond enthalpy: Energy required to break one mole of bonds in one mole of gaseous molecules under standard conditions.
Bond enthalpies are average values and therefore it is only an approximate. The values of bond enthalpies are provided in data booklet.
enthalpy change of reaction formula: reactant- product
STANDARD ENTHALPY CHANGE OF COMBUSTION AND FORMATION
When describing, standard enthalpy change, there are number of steps that must be completed.
Determine the type of reaction.
Write an equation to describe the reaction.
Include state symbols in the equation.
The standard enthalpy change of combustion, ΔHc⦵, is the enthalpy change that occurs when one mole of a substance in its standard state is burned completely in oxygen.
The standard enthalpy change of formation, ΔHf⦵, of a substance is the energy
change that occurs when one mole of a substance is formed from its constituent
elements in their standard states.
The standard state of a pure substance is the form that it
takes under standard conditions: 25.00°C (298.15K), which is taken as being room temperature, and a pressure of 1.00×10^5
Pa.
Value of standard enthalpy of combustion: Section 14 of data booklet
Value of standard enthalpy of formation: Section 13 of data booklet
Enthalpy of formation for element: 0kJ/mol
Lattice Enthalpy
IB DEFINITION: Lattice enthalpy is the energy required to convert one mole of a solid ionic compound into gaseous ions (ENDO)
2nd DEFINITION:
energy required to form one mole of a solid ionic compound from gaseous ions (EXO)
FACTORS AFFECTING LATTICE ENTHALPY
IONIC CHARGE (higher ionic charge, higher electrostatic attraction, higher lattice enthalpy)
IONIC RADIUS (smaller ionic radius, higher charge density, higher electrostatic attraction, higher lattice enthalpy)
BOND ENTHALPY
ENTHALPY: average value of a particular type of bond which has been measured over a range of molecule. always has positive value as bonds are being broken and so energy is absorbed
Calculation of bond enthalpy
ΔHrxn = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
HOW TO CALCULATE??
use data in table 11
draw structure of reactant and product
apply equation
DEFINITION: the enthalpy change when one mole of specific bonds are broken (must be in gaseous phase)X-Y(g) → X(g) + Y(g)
R1.4 - ENTROPY
About Entropy
Definition
Entropy is a measure of the dispersal or distribution of matter and/or energy in a system.
The more ways the energy can be distributed, the higher the entropy
2nd Law Of Thermodynamic = As time moves foward, matter and energy become more dispersed and the total entropy pf the universe increase
Predicting Entropy Changes
)
Factors Affecting Entropy
1) Number Of Gas Molecule
a) Increase = +ve
b) Decrease = -ve
2) Change Of State Of Matter
a) Solid ----> Liquid ( +ve )
b) Liquid ----> Solid (-ve)
3) Temperature Change
a) Temp Increase = +ve
b) Temp decrease = -ve
Absolute Entropy
Zero Entropy = A perfectly ordered solid at absolute zero
Calculating Entropy Changes
ΔS° = ΣS°(product) - ΣS°(reactant)
=The value of entropy can be obtained from the Data Booklet (Section 13)=
Example :
Calculate entropy change of this equation
C2H4(g) + H2(g) ----> C2H6(g)
Solution = 230 - (220+131) = -121JK-1mol-1
ΔG (Gibbs Free Energy)
Gibs Free Energy & Spontanity
1) G<0 (Spontaneous)
2) G=0 (Equlibrium)
3) G>0 (Non-Spontaneous)
ΔG, ΔH & ΔS
Relationship Between ΔG, ΔH, ΔS & Spontanity
Reference : ΔH,ΔS = ΔG/Spontaneity
+ve, +ve = Depends On Temp / Yes (Temp High Enough)
-ve, +ve = Always -ve / Always
-ve, -ve = Depends On Temp / Yes (Temp Low Enough)
+ve, -ve = Always +ve / Never
Calculation
Example:
Calculate the value of for the reaction of ethanol and oxygen to produce ethanoic acid and waterat 298 K using the and S values in section 13 of the Data Booklet
Predict whether the reaction will be spontaneous or non-spontaneous
Answer:
= i- Calculate (-286 + (-484)) - (0 + (-278) ) = -492
ii- Calculate (70 + 160) - (205 + 161 ) = -136
iii- Convert to kJ -136/1000 = -0.136
iv- Formula -492 - 298(-0.136) = -451.472 kJmol-1
ΔG & Equlibrium
Formula
a) Standard Condition (ΔG = -RT ln K)
b) Non Standard (ΔG = -ΔG - RT ln Q)
R1.1 - MEASURING ENTHALPY CHANGE (ΔH)
energy transfer in chemical reactions :star:
isolated system:
matter and energy can neither enter nor exit the system.
closed system:
no transfer of matter, though energy may be transferred across boundary
open system:
transfer of matter and energy is possible across its boundary
energy is commonly transferred in the form of heat, but may also be in electricity, sound, light.
chemical potential energy is stored in chemical bonds of
REACTANTS
and
PRODUCTS
type of chemical reactions :explode:
endothermic, ΔH (+ve)
absorb heat from surrounding
feel cold
reactant
(
more
stable, bond enthalpy
high
)
product
(
less
stable, bond enthalpy
low
)
thus energy absorbed
exothermic, ΔH (-ve)
release heat to the surrounding
feel hot
reactant
(
less
stable, bond enthalpy
low
)
product
(
more
stable, bond enthalpy
high
)
thus energy released
standard condition:
temperature (298K), pressure (101.3 kPa), concentration (mol dm–3)
colorimetry:
measurement of heat flow
colorimeter:
apparatus to measure heat flow
heat capacity:
amount of energy required to raise the temperature of an object by one degree (unit JK–1)
specific heat capacity:
heat capacity of a substance (unit Jg-1 K-1)
ΔH neutralisation :!!:
definition:
enthalpy change when 1 mole of H+ ion from an acid combine wit 1 mole of OH- form base to form 1 mole of water under standard condition.
example :
a student mix 50 cm3 of 1.0 mol dm-3 HCl and 50 cm3 of 1.0 mol dm-3 NaOH in a coffee-cup colorimeter, the temperature of solution increases from 21.0°C to 27.5°C. calculate enthalpy change. (shc of H2O = 4.18 Jg-1K-1)
solution
1) write equation
2) find Q
3) find mole of water
HCl + NaOH --> NaCl + H2O
Q = mcΔT
= (100g)(4.18Jg-1K-1)(27.5°C - 21.0°C) = -2717 J
= -2.7 kJ
n HCl = n NaOH
(no limiting reagent)
n =
1.0 mol dm-3 x 0.05 dm3 = 0.05 mol
thus, n H2O = 0.05 mol
ΔH per mole of H2O = Q ÷ n
-2.7kJ ÷ 0.05 mol = -54kJ/mol
must find ΔH for 1 mole of water
ΔH combustion :fire:
definition:
enthalpy change when 1 mole of substance is completely burnt in oxygen under standard condition.
example 1
: 0.25g ethanol, C2H5OH was burnt and the heat given out raised the temperature of 500cm3 of water from 20.1°C to 23.4°C. calculate heat of combustion of 1 mole of ethanol
solution1
1) write equation
2) find Q
3) find mole of substance
C2H5OH + 3O2--> 2CO2 + 3H2O
Q = mcΔT
= (500g)(4.18Jg-1K-1)(23.4°C - 20.1°C)
= -6897 J
= -6.897 kJ
n, ethanol =
mass ÷ molar mass
n =
0.25g ÷ 46g/mol
n =
0.00543 mol
ΔH per mole of ethanol = Q ÷ n
-6.897kJ ÷ 0.00543 mol = -1269kJ/mol
example 2:
what is the heat evolved,Q from the combustion of 15.5g of C3H8? given that ΔH combustion of C3H8 is -2219kJ/mol.
solution 2
1) write equation
2) find ΔH
3) find mole of substance
4) find Q
C3H8 + 5O2--> 3CO2 + 4H2O
ΔH = -2219kJ/mol
n = 15.5 g ÷ 44.0 g/mol
= 0.352 mol
Q = ΔH x n
= -2219 x 0.352
= -779 kJ
example 3
calculate enthalpy of reaction using values in Table 14 in Data Booklet.
C2H4 + H2 --> C2H6
solution 3
use formula ΔHrxn
= ΣΔH reactant - ΣΔH product
ΔHrxn
= [-1411+ (-286)] - (-1561)
= -136 kJ/mol
can measure using 2 ways:
1) ΔH=Q/n
2) formula ΔHrxn
= ΣΔH reactant - ΣΔH product
ΔH formation :check:
definition:
enthalpy change when 1 mole of substance is formed from it's constituent elements under standard conditions
example 1 :
calculate ΔHrxn using table 14 from data booklet.
C(graphite) + O2(g) --> CO2(g)
solution
C (graphite) = 0kJ/mol
O2 (g) = 0kJ/mol
CO2 (g) = -3944kJ/mol
ΔHf= product - reactant
= (0+0) - (-394)
= +394kJ/mol
example 2 :
c Using following data, calculate ΔHf for CS2(l)
C(graphite)(s) + O2(g) --> CO2(g) ΔHrxn= -393.5kJ/mol
S(rhombic)(s) + O2(g) --> SO2(g) ΔHrxn= -296.4kJ/mol
CS2(l) + 3O2(g) --> CO2(g) + 2SO2(g) ΔHrxn= -1073.6kJ/mol
solution
equation: C(graphite) + 2S --> CS2
method: algebraic method
= -393.5 + (0296.4)(2) + (+1073.6)
= +87.3 kJ/mol
standard enthalpy of formation for an
ELEMENT
is
ZERO
example:
C (graphite) = 0
O2 (g) = 0
CO2 (g) = -3944kJ/mol
methods to calculate
1) formula,
ΔHrxn = ΔHf(products) - ΔHf(reactants)
2) algebraic method
R1.3 - FOSSIL FUEL
Biofuel
definition:
Biofuels are renewable source which is produced from the biological fixation of carbon over short period of time
ethanol as biofuel
equation : C6H12O6 --> 2C2H5OH + 2CO2
it can be used in internal combustion engine
produced from fermentation of glucose from plants that are high in starches
the process is carried out at 37 °C, without oxygen
advantages and disadvantages of using biofuel
carbon fixation diagram
non-renewable energy
Nonrenewable energy comes from sources that will run out or will not be replenished in our lifetimes
fossil fuels such as coal and natural gas
nuclear energy
renewable energy
geothermal
wind
tidal power
solar (sun)
biomass
Renewable energy is energy derived from natural sources that are replenished at a higher rate than they are consumed.
gasohol
mixture of 10% ethanol and 90% gasoline
gasohol can be substitute for gasoline
advantage of using 10% ethanol:
renewable
low CO and nitrogen oxides emission
decrease dependency on oil
disadvantage of using ethanol:
ethanol absorbs water as it can form hydrogen bonds with H2O in the atmosphere
leads to ethanol separating from hydrocarbon components in fuel
can cause corrosion
(methanol can also be used)
Fuel Cells
what is fuel cell??
used to convert chemical energy of fuel ----> electrical energy
Hydrogen fuel cell
acidic electrolyte
Worked example for half equation at anode and cathode
diagram
what is the process? :smiley:
1) PEM only allow H+ ions to diffuse between anode and cathode
2) H2 gas is
oxidised at anode
3) electrons cannot pass through PEM, so it has to leave through the external circuit, producing electric current
diagram of hydrogen cell
alkaline electrolyte
diagram
hydrogen can be used as fuel also because it doesnt produce CO2
the redox reaction in it (transfer of e from hydrogen to oxygen) can produce electric current, heat, and water
the reactants are continuously supplied to different electrodes
diagram
direct methanol fuel cell (DMFC)
worked example for two half equations
cleaner than hydrogen because it produce less greenhouse gas
has greater energy density than hydrogen
can be used to supply hydrogen has for hydrogen fuel cell
Fossil Fuels
Fossil fuels is a fuel formed from the remains of living organisms over millions of years. Fossil fuels have a high carbon content and release energy when combusted. Fossil fuels include crude oil, coal and natural gas.
COAL
Originated as prehistoric forest which were flooded, buried and then gradually compressed by layer upon layer of soil
Forest-Peat-Lignite-Bituminous (soft coal)-Anthracite (hard coal)
Advantages:
cheap and plentiful
reserves last much longer than oil reserves
large infrastructure already exist for transporting and burning of coal
Disadvantages:
Lowest specific energy
release the most carbon dioxide per unit of energy produced
can lead to acid rain
CRUDE OIL (petroleum)
A complex mixture of hydrocarbon that supply us with fuels for transportation and electric generation
Advantages:
Very versatile
High volatility
Sulfur impurities can be removed easily
Disadvantages:
Limited reserves
note evenly distributed around the world
used in vehicles can cause pollution and risk health
NATURAL GAS
Advantages:
Produce fewer pollutant per unit energy
easily transported in pipeline and pressurized containers
Flammable gas that formed naturally and composed of methane, hydrogen sulphide and nitrogen that trapped in geological formation capped by impermeable rocks
Disadvantages :
limited supplies
contributed to global warming
risk of explosion dues to leaks
Specific Energy
Gasoline and kerosene have a high specific energy, small mass of fuels, offer very long ranges
A higher carbon content, more carbon dioxide being produce
natural gas produces the leas amount of CO2 per unit energy release and coal produces the most
Specific energy = energy released from the fuel / mass of fuel consumed
CO2 Levels and Greenhouse Effect
CO2 is considered as a greenhouse gas due to its ability to absorbed infrared radiation.
CO2 will emit IR back into the atmosphere
increase in global temperature and lead to global warming
CO2 trap heat energy inside atmosphere
Impact on Climate:
Changes in agriculture-crop yield
Rising sea levels dues to melting glaciers
Changes in average global temperature - CO2 contributed 50% towards global warming
Combustion
incomplete combustion
produces CO, or C (soot) with water
equation:
1) 2C3H8 +7O2 —> 6CO2 + 8H2O
2) C3H8 +2O2 —> 3C + 4H2O
releases less heat than complete combustion
larger hydrocarbon tends to undergo incomplete combustion more and produces more soot
diagram
complete combustion
its called complete combustion because it has completely oxidised products
produces C02 and H20 only as the products
metals
metal+O2 --> metal oxide
equation: 4Li + O2 --> 2LiO2
half equation: Li-->Li+e (loss of e)
results in oxidation of metal and reduction of oxygen
oxidation meaning
Oxidation is the loss of electrons/hydrogen during a reaction by a molecule, atom or ion.
Oxidation occurs when the oxidation state of a molecule, atom or ion is increased.
oxygen is added
redox reaction:
reduction meaning
gain of electrons/hydrogen
the oxidation state of an atom, molecule, or ion decreases
oxygen is removed
non metals
non metal+02--> non metal oxide
equation: S+O2-->SO2
organic compounds
alkanes
alcohol
both produces C02+H20 as byproduct and both is exothermic
