# Chapter 10 ```Energy
10.1 The Nature of Energy
 Energy: the ability to do work or
produce heat
 Potential energy (store energy):
energy due to position or
composition
 Kinetic energy (motion energy):
energy due to motion of the object
and depends on the mass of the
object and its velocity

KE = &frac12; (mv2)
 Law of conservation of energy:
that energy can be converted from
one from to another but can be
neither created or destroyed.

Energy of the universe is constant
The nature of energy
 Work: force acting over a
distance
 W = F/d
 State function: property of
the system that changes
independently of its pathway
 Which of the following are state
functions, and which are not
 The temperature of an ice cube
 The volume of an aerosol can
 The amount of time required
for a 10 mi bike ride
10.2 Temperature and Heat
 Temperature: is a measure of the random motions of the
components of a substance
 E.g H2O molecules move rapidly in hot water than in cold
water
 Heat: a flow of energy due to a temperature difference
 Tfinal = average temp from mixing (hot &amp; cold temp)
10.3 Exothermic and Endothermic
Process

System – everything
we focus on in
experiment

Surroundings –
everything other the
system


exothermic (energy
flows out of system to
surrounding (via heat)
endothermic ( energy
flows into system from
surrounding (via heat)
Examples
 Identify whether these process are exothermic or
endothermic
 Your hand gets cold when you touch ice
 The ice melts when you touch it
 Propane is burning in a propane torch
 Two chemicals mixing in a beaker give off heat
10.4 Thermodynamics
 Is the study of energy.
 First law of thermodynamics: the energy of the universe is
constant
 Internal energy – energy of the system
 ∆E = q x w
 ∆ =&gt; change in the function
 q =&gt; represents heat
 W =&gt; represents work
10.5 Measuring Energy changes
 calorie: the amount of energy (heat) required to raise the
temperature of one gram of water by 1oC
 1Calorie = 1000 calories
 Joule (J) – SI unit
 1 calories = 4.184 joules
Converting Calories to Joules
 Express 60.1 cal of energy in units of joules
 How many calories of energy corresponds to 28.4 J?
Specific heat
 The amount of energy
required to change the
temperature of one gram
of a substance by 1oC
 Denoted as s
 Heat required = specific
heat x mass x change in
temp
 q = s m ∆T
Calculating Energy Requirements
 Determine the amount of energy (heat) in joules required to raise
the temperature of 7.40 g water from 29.0oC to 46.0oC
 A 5.63 g sample of solid gold is heated from 21oC to 32oC. How
much energy in Joules and calories is required?
 A sample of gold requires 3.1 J of energy to change its
temperature from 19oC to 27oC. What is the mass of this sample
of gold
 A 55.0 g aluminum block initially at 27.5oC absorbs 725 J of
heat. What is the final temperature?
10.6 Enthalpy
 Heat of reaction (∆H) = the amount of heat that being
absorbed or release by a given reaction
 ∆Hp = heat
Enthalpy
 When 1 mol of methane (CH4) is burned at constant
pressure, 890 kJ of energy is released at heat. Calculate
∆H for a process in which a 5.8 g sample of methane
is burned at constant pressure
 The reaction that occurs in the heat packs used to treat
sports injuries is
4Fe(s) + 3 O2(g)  2Fe2O3(s) ∆H= -1652 kJ
How much heat is released when 1.00 g of Fe(s) is
reacted with excess O2(g)
Entropy
 Entropy (S) – measure of
disorder or randomness
 Randomness increase -&gt; S increases
 Sgas&gt; Sion &gt;Smolecule&gt;Ssolid
 Ssolid = 0
 E.gH2O(s)  H2O (l)
H2O (l)  H2O (g)
H2O (g)  H2O (l)
S=+
S=+
S=-
 2nd law of thermodynamic: the
entropy of the universe is always
increasing
```