Transcript Chapter Two Atoms & The Periodic Table

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Exam over Chapters 8 & 9
Chapter ten notes
 Section 10.1
 Section 10.2
 Section 10.3
 Section 10.4
 Section 10.5
 Section 10.6
 Section 10.7
 Section 10.8
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System: Part we care about
 Reactants & Products
Surroundings: Everything
else in the universe
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Aopen system (mass and heat pass through)
Bclosed system (heat only pass through)
Cisolated system (no heat or mass transfer)
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For chemical reactions to happen
spontaneously, the final products must be
more stable than the starting reactants
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Higher energetic substances
 Typically less stable, more reactive
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Lower energetic substances
 Typically more stable, less reactive
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Thermal energy flows from warmer to cooler
H2O(s)  H2O(l)
2H2(g) + O2(g)  2H2O(l)
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Study of heat and its transformations into
other energies
 Thermochemistry is a part of this
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Thermodynamics studies changes in the
state of a system
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State functions are properties that are
determined by the state of the system,
regardless of how it was achieved
 Final – Initial
 Ex:
▪ Energy
▪ Pressure
▪ Volume
▪ Temperature
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Has 2 components:
 Kinetic energy: various types of molecular and
electron motion
 Potential energy: attractive and repulsive
interactions between atoms and molecules
 ΔU = U(products) – U(reactants)
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ΔU = q + w
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q = heat (absorbed or released by the system)
w = work (done on or by the system)
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Calculate the overall change in internal
energy (ΔU) for a system that absorbs 188 J
of heat and does 141 J of work on its
surroundings.
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Convert 723.01 J into calories
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SKETCH and LABEL what an exothermic and
endothermic energy vs. time graph would
look like.
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Calculate the overall change in internal
energy for a system that releases 43 J in heat
and has 37 J of work done on it by its
surroundings
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Reactions can be carried out in two ways:
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In a closed container (constant volume):
 qv = ΔU
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In an open container (constant pressure):
 qp = Δ H
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Combustion of propane gas:
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ΔH = H(products) – H(reactants)
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“+” = endothermic
“—” = exothermic
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H2O(s)  H2O(l)
ΔH = +6.01 kJ/mol
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH = -890.4 kJ/mol
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH = -890.4 kJ/mol
How much energy is release from 18.4 g of methane
being burned?
If 924.3 kJ of energy was released, how many grams
of water was produced?
If you change the AMOUNTS in a balanced
equation, you change the enthalpy the same
way
1)
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Ex: if coefficients are doubled, so is the enthalpy
If you reverse the equation, you reverse the
sign of the ΔH
2)
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Ex: H2O(s)  H2O(l) ΔH = +6.01 kJ/mol
H2O(l)  H2O(s) ΔH = -6.01 kJ/mol
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Measurement or heat changes within a
system
 Using a calorimeter
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Specific Heat (s): amount of heat required to
raise the temperature of 1 g of a substance by
1°C (ex: liquid water is 4.184 J/(g*°C)
q = (s)(m)(ΔT)
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Heat Capacity (C): amount of heat required
to raise the temperature of an object by 1°C
q = (C)(ΔT)
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What is the amount of heat (in kJ) required to
heat 255 g of water from 25.2 °C to 90.5 °C?
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Can calculate changes in heat using
styrofoam cups and known mass of water
 Assuming constant pressure
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Therefore…
qp = msΔT = ΔH
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System: reactants and products (the reaction)
Surroundings: water in calorimeter
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For an exothermic reaction:
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 The system loses heat
 The surroundings gain (absorb) heat
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A 30.4-g piece of unknown metal is heated up
in a hot bath to a temperature of 92.4°C. The
metal is then placed in a calorimeter
containing 100. g of water at 25.0°C. After
the calorimeter is capped, the temperature of
the calorimeter raises to 27.2°C. What was
the specific heat of the unknown metal?
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Ex: 50.0 mL of 1.00 M HCl and 50.0 mL of
1.00 M NaOH are mixed in a calorimeter with
100 g of water and capped at room temp
(25°C). The reaction reaches a max of 31.7°C.
What is the ΔH°rxn?
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125.0-g of a metal is heated to 100.0°C. It is
then placed into a calorimeter containing
100.0 mL (100.0 g) of water at 25.0°C and
capped. The energy is transferred and the
max temperature of 34.1°C is reached. What
is the specific heat of the metal?
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Given the following, determine the ΔH for
3H2(g) + O3(g)  3H2O(g)
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Standard Enthalpy of Formation (ΔH°f): heat
change that results when 1 mole of a
compound is formed from its constituent
elements in their standard states
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“Standard State” means “stable form”
 1 atm and 25°C typically
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Example: O(g) (249.4), O2(g) (0), O3(g) (142.2)
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ΔH°rxn: enthalpy of a reaction under standard
conditions
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When we know reactions go to completion or
can be done in one step, we can use a direct
method
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Ex: Calculate ΔH°rxn for
2SO(g) + 2/3O3(g)  2SO2(g)
From Appendix 2: SO(g): (5.01), O3(g): (142.2),
SO2(g): (-296.4)
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When a reaction is too slow or side reactions
occur, enthalpy of reaction can be calculated
using Hess’s Law
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Recall: when bonds are made, energy is
given off (exo); when bonds break, energy is
needed (endo)
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Bond Enthalpy: the measure of stability of a
molecule
 Enthalpy change associated with breaking a
particular bond in 1 mole of gaseous molecules
▪ H2(g)  H(g) + H(g) ΔH = 436.4 kJ/mol
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The higher the bond enthalpy, the stronger
the bond
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The bonds in different compounds have
different bond enthalpies
 Ex: O—H bond in water vs. O—H bond in
methanol are different
 Therefore, we speak of AVERAGE bond enthalpy
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Recall: amount of energy required to convert
1 mole of ionic solid to its constituent ions in
the gas phase
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Ex: NaCl(s)  Na+(g) + Cl-(g)