Chapter Two Atoms & The Periodic Table
Download
Report
Transcript Chapter Two Atoms & The Periodic Table
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
System: Part we care about
Reactants & Products
Surroundings: Everything
else in the universe
Aopen system (mass and heat pass through)
Bclosed system (heat only pass through)
Cisolated system (no heat or mass transfer)
For chemical reactions to happen
spontaneously, the final products must be
more stable than the starting reactants
Higher energetic substances
Typically less stable, more reactive
Lower energetic substances
Typically more stable, less reactive
Thermal energy flows from warmer to cooler
H2O(s) H2O(l)
2H2(g) + O2(g) 2H2O(l)
Study of heat and its transformations into
other energies
Thermochemistry is a part of this
Thermodynamics studies changes in the
state of a system
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
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)
ΔU = q + w
q = heat (absorbed or released by the system)
w = work (done on or by the system)
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.
Convert 723.01 J into calories
SKETCH and LABEL what an exothermic and
endothermic energy vs. time graph would
look like.
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
Reactions can be carried out in two ways:
In a closed container (constant volume):
qv = ΔU
In an open container (constant pressure):
qp = Δ H
Combustion of propane gas:
ΔH = H(products) – H(reactants)
“+” = endothermic
“—” = exothermic
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)
Ex: if coefficients are doubled, so is the enthalpy
If you reverse the equation, you reverse the
sign of the ΔH
2)
Ex: H2O(s) H2O(l) ΔH = +6.01 kJ/mol
H2O(l) H2O(s) ΔH = -6.01 kJ/mol
Measurement or heat changes within a
system
Using a calorimeter
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)
Heat Capacity (C): amount of heat required
to raise the temperature of an object by 1°C
q = (C)(ΔT)
What is the amount of heat (in kJ) required to
heat 255 g of water from 25.2 °C to 90.5 °C?
Can calculate changes in heat using
styrofoam cups and known mass of water
Assuming constant pressure
Therefore…
qp = msΔT = ΔH
System: reactants and products (the reaction)
Surroundings: water in calorimeter
For an exothermic reaction:
The system loses heat
The surroundings gain (absorb) heat
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?
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?
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?
Given the following, determine the ΔH for
3H2(g) + O3(g) 3H2O(g)
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
“Standard State” means “stable form”
1 atm and 25°C typically
Example: O(g) (249.4), O2(g) (0), O3(g) (142.2)
ΔH°rxn: enthalpy of a reaction under standard
conditions
When we know reactions go to completion or
can be done in one step, we can use a direct
method
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)
When a reaction is too slow or side reactions
occur, enthalpy of reaction can be calculated
using Hess’s Law
Recall: when bonds are made, energy is
given off (exo); when bonds break, energy is
needed (endo)
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
The higher the bond enthalpy, the stronger
the bond
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
Recall: amount of energy required to convert
1 mole of ionic solid to its constituent ions in
the gas phase
Ex: NaCl(s) Na+(g) + Cl-(g)