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Average KE same
Heat (q) transfer thermal energy from hot to cool due to temp diff
2..2
1vmKE
Average translational energy/KE per particle
Heat Temperature
Heat vs Temperature
Symbol Q Unit - Joule
Form of Energy Symbol T Unit – C/K
Not Energy
At 80C
Distribution of molecular speed, Xe, Ar, Ne, He at same temp
2.2
1vmKE
80oC
Diff gases have same average KE per particle.
Click here Heat vs Temperature Click here specific heat capacity
He – mass low ↓ - speed v high ↑ Xe – mass high ↑ - speed v low ↓
Temp ᾳ Average KE Higher temp - Higher average KE
2..2
1vmKE
Movement of particle, KE.
Heat energy (energy in transfer)
80oC 50oC
degree of hotness/coldness
Total KE/PE energy of particles in motion
1 liter water 80C 2 liter water 80C
Same Temp Same average kinetic energy per particle Same average speed
Same temp
Diff amt heat
Specific heat capacity Amount of heat needed to increase
temp of 1g of substance by 1C
Q = Heat transfer
Click here specific heat capacity Click here specific heat capacity
80oC 50oC
Warmer body higher amt average KE
Energy transfer as heat
Gold
0.126J/g/K
Silver
0.90J/g/K
Water
4.18J/g/K
Cold body lower amt average KE
Q = mcθ
Heat
Total KE/PE energy of particles in motion
Symbol Q Unit - Joule
Form of Energy
Amt heat energy Q, need to change temp depend
m = mass c = specific heat capacity
θ = Temp diff
Lowest Highest specific heat capacity 0.126J 4.18J ↓ ↓ to raise 1g to 1 K to raise 1g to 1K
Click here themochemistry notes
Coffee-cup calorimeter
constant pressure – no gas
Calorimetry - techniques used to measure enthalpy changes during chemical processes.
Bomb calorimeter
Constant vol – gas released
80C
50C
Heat capacity bomb Heat released
∆Hc is calculated.
Combustion-exo - temp water increase.
Specific heat capacity Amount of heat needed to increase temp of 1 g of substance by 1C
Q = Heat transfer
Q = mcθ m = mass c = specific
heat capacity θ = Temp diff
Coffee-cup calorimeter
constant pressure – no gas
Calorimetry - techniques used to measure enthalpy changes during chemical processes.
Bomb calorimeter
Constant vol – gas released
Cup calorimeter Determine specific heat capacity of X
m = 1000g
Heated 200 C
5000 ml water
m = 5000g
c = 4.18
Ti = 20 C
Tf = 21.8 C
Heat lost by X = Heat gain by water mc∆T = mc∆T
X
1000 x c (200 – 21.8) = 5000 x 4.18 x (21.8 – 20) c = 37620/ 178200 c = 0.211J/g/K
Benzoic acid – used std – combustion 1g release 26.38 kJ Combustion 0.579 g benzoic acid cause a 2.08°C increase in temp. 1.732g glucose combusted, temp increase of 3.64°C. Cal ΔHcomb X.
Bomb calorimeter (combustion) Find heat capacity of bomb and ∆Hc X
Bomb sealed, fill with O2.
1g – 26.38kJ 0.579g – 26.38 x 0.579 Q = - 15.3kJ
∆Hc X = Qbomb Find heat capacity bomb
Q bomb = c∆T
KkJc
T
Qc
TcQ
/34.7
08.2
3.15
Qbomb = c∆T = 7.34 x 3.64 = 26.7 kJ
Insulated with water.
Combustion X
Find Q using
benzoic acid
1.732g – 26.7 kJ 180g – 2.78 x 103 kJmol-1
Click here bomb calorimeter
X
X
1. 2. 3.
System – rxn vessel (rxn take place)
open system closed system isolated system
Enthalpy – Heat content/Amt heat energy in substance - Energy stored as chemical bond + intermolecular force as potential energy
Exchange
energy
Exchange
matter
Exchange
energy
NO Exchange
energy
NO Exchange
matter
Heat(q) transfer from system to surrounding ↓
Exothermic. ∆H < 0 ↓
HOT
Surrounding – rest of universe
Heat(q) transfer to system from surrounding ↓
Endothermic. ∆H > 0 ↓
COLD
H
Time
H
Time
Heat
energy
Heat
energy ∆H = + ve
∆H = - ve
∆H system = O
reaction system
surrounding
No heat loss from system (isolated system)
∆Hrxn = Heat absorb water (mc∆θ)
∆Hrxn = mc∆θ
water
Enthalpy Change = Heat of reaction = -∆H
2Mg(s) + O2(g) →2MgO(s) ∆H = -1200kJ mol-1 Enthalpy/H (heat content)
2Mg + O2
2MgO
∆H= -1200kJ mol-1
- Energy neither created nor destroyed - Converted from one form into another. - Amt heat lost by system equal to amt heat gain by its surrounding. - Total energy system plus its surrounding is constant, if energy is conserved.
change
Energy Flow to/from System
System – KE and PE energy – Internal Energy (E)
Change Internal energy, ∆E = E final – E initial
Energy transfer as HEAT or WORK Lose energy to surrounding as heat or work
E = sum kinetic energy/motion of molecule, and potential energy represented by chemical bond bet atom
∆E = q + w
∆E = Change internal energy
q = heat transfer
w = work done by/on system
Thermodynamics Study of work, heat and energy on a system
Change Internal energy, ∆E = E final – E initial
Energy transfer as HEAT or WORK Gain energy from surrounding as heat or work
Heat add , q = + 100 J Work done by gas, w = - 20 J ∆E = + 100 – 20 = + 80 J
Q = Heat gain
+ 100J
w = work done by system = -20 J
w = work done on system = +20 J
Q = Heat lost
- 100J Heat lost , q = - 100 J Work done on gas, w = + 20 J ∆E = - 100 + 20 = - 80 J
∆E universe = ∆E sys + ∆E surrounding = 0
System – KE and PE energy – Internal Energy (E)
Heat and work Heat only
Q = Heat gain
+ 100 J
No work – no gas produced
Heat add , q = + 100 J No work done = 0 ∆E = q + w ∆E = + 100 = + 100 J
Q = Heat lost
- 100J
Heat lost, q = - 100 J No work done = 0 ∆E = q + w ∆E = - 100 = - 100 J
No work – no gas produced
H = E + PV ∆H = ∆E + P∆V
Enthalpy change w = work done by/on gas
1st Law Thermodynamics
∆E = q + w ∆E = q
∆E = q + w + q = sys gain heat - q = sys lose heat + w = work done on sys - w = work done by sys
∆E = q + w
Work done by gas
No gas – No work
change
Energy Flow to/from System
System – KE and PE energy – Internal Energy (E)
Change Internal energy, ∆E = E final – E initial
Energy transfer as HEAT or WORK Lose energy to surrounding as heat or work
E = sum kinetic energy/motion of molecule, and potential energy represented by chemical bond bet atom
∆E = q + w
∆E = Change internal energy
q = heat transfer
w = work done by/on system
Thermodynamics Study of work, heat and energy on a system
Change Internal energy, ∆E = E final – E initial
Energy transfer as HEAT or WORK Gain energy from surrounding as heat or work
No work done by/on system ∆E = q + w w = 0 ∆E = + q (heat flow into/out system) ∆H = ∆E = q (heat gain/lost)
∆E universe = ∆E sys + ∆E surrounding = 0
System – KE and PE energy – Internal Energy (E)
Heat only – Exothermic and Endothermic reaction
Q = Heat gain
+ 100 J
No work – no gas produced
Heat add , q = + 100 J No work done = 0 ∆E = q + w ∆E = + 100 J ∆E = ∆H = + 100 J
Q = Heat lost
- 100J Heat lost, q = - 100 J No work done = 0 ∆E = q + w ∆E = - 100 J ∆E = ∆H = - 100J
No work – no gas produced
H = E + PV ∆H = ∆E + P∆V
Constant pressure
Enthalpy change w = work done by/on gas
1st Law Thermodynamics
P∆V = 0
∆E = q + w ∆H = ∆E + P∆V
∆E = q + 0 ↓
∆E = q
No gas produced V = 0
∆H = ∆E + 0 ↓
∆H = ∆E
At constant pressure/no gas produced
∆H = q ∆Enthalpy change = Heat gain or lost
No work done w = 0
H
E
E
∆H = + 100J
H ∆H = - 100J
Enthalpy Change
Heat(q) transfer from system to surrounding ↓
Exothermic ∆H < 0 ↓
HOT
Heat
energy ∆H = - ve
Enthalpy Change = Heat of rxn = -∆H
Mg(s) + ½ O2(g) → MgO(s) ∆H = -1200kJ mol-1
Mg + ½ O2
MgO
∆H= -1200
- Energy neither created nor destroyed - Converted from one form into another. - Amt heat lost by system equal to amt heat gain by its surrounding. - Total energy system plus its surrounding is constant, if energy is conserved.
Reactant (Higher energy - Less stable/weaker bond)
Product (Lower energy - More stable/strong bond)
Temp surrounding ↑
Exothermic rxn Combustion C + O2 → CO2
Neutralization H+ + OH- → H2O Displacement Zn + CuSO4 → ZnSO4 + Cu Condensation H2O(g) → H2O(l)
Freezing H2O(l) → H2O(s)
Precipitation Ag+ + CI- → AgCI(s)
Endothermic rxn Dissolve NH4 salt NH4CI (s) → NH4
+ + CI -
Dissolve salt MgSO4. 7H2O(s) → MgSO4(aq
CuSO4. 5H2O(s) → CuSO4(aq) Na2CO3.10H2O(s) → Na2CO3(aq Evaporation/Boiling H2O(l) → H2O(g)
Melting H2O(s) → H2O(l)
Heat(q) transfer to system from surrounding ↓
Endothermic. ∆H > 0 ↓
COLD
Heat
energy
Reactant (Lower energy - More stable/strong bond)
Product (Higher energy - Less stable/weak bond)
∆H = + ve Temp surrounding ↓
Click here thermodynamics
∆H= + 16
NH4CI (s)
NH4CI (aq)
Enthalpy Change = Heat of rxn = -∆H
NH4CI (s) → NH4CI (aq) ∆H = + 16 kJ mol-1
Click here enthalpy
3000
1800
116
100
E
X
O
E
N
D
O
NaCI (s)
Na(s) + ½CI2 (g))
LiCI (s)
Li+(g) + CI– (g)
AgCI
Ag+ + CI-
NaCI + H2O
HCI + NaOH
ZnSO4 + Cu
Zn + CuSO4
Li+(aq)
Li+(g) + H2O
LiCI(aq)
LiCI + H2O
2CO2 + 3H2O
C2H5OH + 3O2
- Energy neither created nor destroyed - Converted from one form into another. - Amt heat lost by system equal to amt heat gain by its surrounding. - Total energy system plus its surrounding is constant, if energy is conserved.
Std Enthalpy Changes ∆Hθ
Std condition
Pressure 100kPa
Conc 1M All substance at std states
Temp 298K
Bond Breaking Heat energy absorbed – break bond0
Bond Making Heat energy released – make bond
Std ∆Hcθ combustion
Std Enthalpy Changes ∆Hθ
∆H for complete combustion 1 mol sub in std state in excess O2
∆H when 1 mol solute dissolved form infinitely dilute sol
Std ∆Hsolθ solution
∆H when 1 mol gaseous ion is hydrated
Std ∆Hhydθ hydration
∆H when 1 mol metal is displaced from its sol
Std ∆Hdθ displacement
∆H when 1 mol H+ react OH- to form 1 mol H2O
C2H5OH + 3O2 → 2CO2 + 3H2O LiCI(s) + H2O → LiCI(aq)
Ag+ + CI - → AgCI (s)
Zn + CuSO4 → ZnSO4 + Cu
Std ∆Hlat θ lattice
∆H when 1 mol precipitate form from its ion
Std ∆Hpptθ precipitation Std ∆Hn
θ neutralization
∆H when 1 mol crystalline sub form from its gaseous ion
HCI + NaOH → NaCI + H2O Li+(g) + CI–(g) → LiCI (s)
Li+(g) + H2O→ Li+
(aq)
∆H = - ve ∆H = - ve ∆H = - ve ∆H = - ve
∆H = - ve ∆H = - ve ∆H = - ve
∆H when 1 mol form from its element under std condition
∆H = - ve
Std ∆Hf θ formation
Na(s) + ½CI2 (g)→ NaCI (s)