PDP Chemistry
Energetics
Thermochemistry: energy changes in chemical reactions
Topic 5: Energetics
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energy
exothermic and endothermic reactions
calorimetry
enthalpy change
bond enthalpy
temperature and heat
absorption and emission of radiation
burning fuels
PDP moles tests comments
 time was short…
 show working for quantitative questions
 13.b) graph scales- we need to do more work with these…
 Chemical equations- balancing, subscripts and coefficients, charges
 — 1. work out reactants and products (a word equation may help
here)
 — 2. write the formula for each compound (use the charge on ions
to find the formula, but the charges are not written in the formula)
 — 3. balance. change only the coefficients, NOT the formulas
 13.d) it is hard to pick out the relevant information and ignore what
you don’t need
Energy
Energy: the ‘capacity to do work’
 the unit of energy is the Joule (J)
 forms of energy:
 kinetic
 gravitational potential
 elastic potential
 chemical potential
 electric potential
 nuclear potential
 thermal (internal energy)
 radiant
Law of conservation of energy
The total energy of an isolated system remains constant- it is
said to be conserved over time.
Energy can be neither created nor be destroyed, but it can
change form.
Exothermic and
endothermic reactions
Why are reactions endothermic or exothermic?
In chemical reactions bonds break and new bonds form.
Different bonds have different amounts of chemical energy.
 Exothermic reaction: less energy in the bonds after the reaction 
energy released as heat
 Endothermic reaction: energy is needed as there is more energy in
the bonds after the reaction  energy taken in
Graphs
 sharp pencil and ruler
 graph paper- use the whole sheet
 choose a linear scale for each axis
 plot each point with an x or a + (or error bars)
 titles and units on axes
 title for graph
 trendline (or maybe join points with straight lines)- no snakes
Description and explanation of results
Conclusion
 Description and explanation of results from one exothermic
and reaction and upload to MB
Temperature
Temperature and energy
 Which gas is at the higher temperature?
 Which gas contains more energy?
Temperature- average KE per particle
higher average speed  higher
temperature
more particles at same average
speed  same temperature
Measuring temperature
Absolute scale of temperature: zero Kelvin = zero kinetic energy
zero K = -273°C
Convert to Celcius
1.
melting point of ice 273K
2.
room temperature 300K
3.
temperature of deep space 4K
4.
boiling point of oxygen at standard pressure 90K
5.
surface of the sun 6000K
6.
temperature of an exploding nuclear bomb 106K
Distribution of speeds of particles
number of
particles
speed
Distribution of speeds
Calorimetry
How much energy does it take to heat
water?
4.18Joules of energy are needed to heat up 1 gram of water
(approx 1mL) by 1°C.
(4.18 is the specific heat capacity of water, and aqueous
solutions are almost the same)
Calculating how much energy is released
1.
Calculate how many Joules are released in each of your
experiments from last week
2.
Calculate how many kJ are released by each mole of
reactant in the experiments.
4.18Joules of energy are needed to heat up 1 gram of water
(approx 1mL) by 1°C.
Exothermic and endothermic reactions
Ammonium nitrate and water
Iron and oxygen
Hot pack (more than one is true)
A. The temperature of the hot pack decreases during the reaction
B. The temperature of the hot pack increases during the reaction
C. The hot pack transfers heat to the person touching it
D. Heat is taken in from the person to the hot pack
E.
The hot pack has less energy stored in chemical bonds after it is used
F.
The hot pack has more energy stored in chemical bonds after it is
used
Energy changes
Exothermic reactions
Stored
chemical
energy, J
reactants
∆H
products
time
Endothermic reactions
Stored
chemical
energy, J
products
∆H
reactants
time
Enthalpy change
Enthalpy, H
 Energy stored in chemical bonds of reactants (in Joules)
 PE and KE of particles + energy to make space for substance
Image:
https://en.wikipedia.org/wiki/File
:Ammonium_Nitrate.jpg
Burning fuels
Fuels combine with oxygen in combustion
most chemical fuels are organic compounds, often
hydrocarbons (contain C, H) or carbohydrates (contain C, H, O)
 carbon oxidized to CO2
 hydrogen oxidized to H2O
Calculate a theoretical ∆H in kJ/mol for
the reactions from yesterday’s lab
C2H6O + 3O2  2CO2 + 3H2O
Energy density
Energy density used to compare fuels
Fuel
hydrogen
natural gas (methane)
ethanol
coal
wood
Energy density, kJ/g
142
56
26
24
16
energy released by when 1g of fuel is burned
Standard enthalpy
Standard enthalpy change of a reaction
∆H⊖ to compare energy changes in reactions
in kJ/mol
(measured at STP 298K and 1atm)
Calculating standard enthalpy
0.2g of magnesium reacts with 10mL excess hydrochloric acid to
increase the temperature by 50°C. Calculate the standard
enthalpy of reaction.
Bond enthalpy
Bond enthalpy: the enthalpy change when one
mole of bonds is formed in the gaseous state
X (g) + Y(g)  X-Y(g)
Forming bonds is exothermic (negative ∆H)
Breaking bonds is endothermic (positive ∆H)
Exothermic reaction
combustion of methane
CH4 + 2 O2  CO2 + 2H2O + heat
Reactants
Products
DH = (Products) – (Reactants)  negative value
Exothermic reactions: DH < O
Endothermic reactions DH > O
Enthalpy of combustion
Bond
Average bond
enthalpy kJ/mol
C-C
347
C=O
746
C-H
413
O=O
498
O-H
464
C-O
358
Image:
https://en.wikipedia.org/wiki/Burning
Calculating standard enthalpy change
What is the enthalpy change in kJ per mole if 45kJ are given out
when 0.8g of methane is burned?
What is the enthalpy change in kJ per mole if 1.6g of methanol is
used to heat 200mL water from 20C to 38C?
Using temperature to calculate ∆Hᶱ
Heat energy = mass x specific heat capacity x temperature change
Q = mc∆T
Pollution from combustion
Acid rain
Greenhouse gases
Carbon monoxide
Particulate pollution
Image: http://en.wikipedia.org/wiki/File:Periodic_table.svg