Energy & States of Matter – Part 2

Energy & States of Matter – Part 2
The Nature of Energy
 Energy - The ability to do work or produce heat.
Two types:
1. Potential energy – energy due to
position or composition (i.e. energy
stored in chemicals). PE = mgh

Examples include water behind a dam or
energy stored in chemical bonds.
2. Kinetic energy – energy due to motion.
KE = ½mv2

Examples include water flowing downhill
over a turbine or gasoline exploding.
10-2
The Nature of Energy
Law of Conservation of Energy – Energy can be
converted from one form to another but can neither
be created nor destroyed.
10-3
Temperature and Heat
Temperature - Measure of the random motions
(average kinetic energy) of the components of a
substance.
Heat – Flow of energy due to a temperature
difference.
Remember:
All particles are in motion. As the
temperature increases, the thermal energy or vigor of
their motion increases.
Heat is the transfer of this thermal energy from one
object to another.
10-4
Figure 10.2: Equal masses of hot and cold water.
10-5
Figure 10.3: H2O molecules in hot and cold water.
10-6
Figure 10.4: H2O molecules in same temperature water.
10-7
T
T
final
hot
T
initial
2
cold
initial
90 C 10 C
2
50 C
Change in temperature (hot) = Tf – Ti =
Thot = 50. C – 90. C = -40. C
Change in temperature (cold) = Tf – Ti =
Tcold = 50. C – 10. C = 40. C
 T =Tf – Ti
In science we use the Greek letter Delta
( ) to mean “change in.” It is ALWAYS
the final value minus the initial value.
Exothermic and Endothermic Processes
The universe is divided into two halves: the system
and the surroundings.
The system is the part we are concerned with. We
define its boundaries.
The surroundings are the rest.
Every reaction has an energy change associated with
it.
Energy is stored in bonds between atoms.
10-9
Exothermic reactions release energy to the
surroundings
Energy exits the
system.
Endothermic reactions absorb energy from the
surroundings.
Energy
Examples:
enters the system.
 A match feels hot because energy is exiting the
system (the match) and entering the surroundings (your
skin, for example).
 An ice cube feels cold for the same reasons, but
energy flows in the opposite directions.
10-10
Thermodynamics
Thermodynamic quantities always consist of two
parts:
A number giving the magnitude of the
change.
A sign indicating the direction of the
flow (from the system’s point of view).
For an endothermic process,
gaining energy).
For an exothermic process,
losing energy).
10-11
E >0 (the system is
E <0 (the system is
Measuring Energy Changes
Different materials respond differently to being
heated.
Units are needed to explore differences:
 calorie (cal) – The amount of energy (heat) required to
raise the temperature of one gram of water by one
Celsius degree.
 joule
(J) = SI unit;
1 calorie = 4.184 joules
FYI:In chemistry calories are written with a small
“c”; calories on food labels are 1000 of these (or
one kilocalorie). We distinguish these with a capital
“C”
10-12
Measuring Energy Changes
 Practice Conversions:
1) Express 34.8 calories of energy in
units of joules.
2) Express 47.3 J of energy in units of
cal.
10-13
 Specific heat capacity (s) – The amount of
energy required to change the
temperature of one gram of a substance
by one Celsius degree.
 Each substance has its own unique
specific heat capacity
 (units are J/g C).
Measuring
Energy
Changes
The Specific Heat Capacities of Some Common Substances
Substance
Specific Heat (J/g C)
Water (l)
4.184
Water (s)
2.03
Water (g)
2.0
Aluminum (s)
0.89
Iron (s)
0.45
Mercury (l)
0.14
Gold
0.13
10-15
Measuring Energy Changes
We need the following equation to calculate energy:
 Q = s x m x
T
Where:
 Q = energy (heat) required
 s = specific heat capacity
 m = mass of the sample (in grams)
 T = Change in temperature (in C)
This equation always applies when a substance is being
heated (or cooled) and no change of physical
state occurs.
10-16
Measuring Energy Changes: Practice Problems
REMEMBER
T =
Change in temperature (in C) =
Tfinal - Tinitial
1) A 2.6 –g sample of a metal requires 15.6 J of
Energy to change its temperature from 21 C to
34 C. Use specific heat values from your notes to
identify this metal.
2) A sample of gold requires 3.1 J of energy to
change its temperature from 19 C to 27 C. What
is the mass of this sample of gold?
3) Calculate the amount of energy required (in
calories) to heat 145-g of water from 22.3 C to
75.0 C.
10-17
Practice Problems
1. If it takes 526J of energy to warm 7.40g of water by 17oC,
how much energy would be needed to warm 7.40 g of
water by 55oC?
2. If 72.4 kJ of heat is applied to a 952 g block of metal, the
temperature increases by 10.7oC. Calculate the specific
heat capacity of the metal in J/goC.
Energy Calculations
 Use Q=sm∆T ONLY when there is a temperature change.
 If no temperature change occurs – only a phase change
(i.e. solid to liquid, etc. . .) then use the formula
 Q = m x Hf
OR Q = m x Hv
 m= mass
 Hf = Heat of fusion (melting or freezing) = 334 J/g
 Hv = Heat of vaporization (evaporating or condensing) = 2260 J/g
Practice
How much energy is required to melt 357g of ice to
liquid water?
2. How much energy will be released when 780g of
water vapor is condensed to liquid water?
3. How much energy will be required to melt 255g of
-5.0˚C, then vaporize it and heat the water vapor to
125˚C?
1.