What is Energy?
1. Energy – The ability to do work.
2. Potential energy – stored energy in chemical
bonds.
3. Average Kinetic energy – energy of motion.
Forms of energy
1) Thermal
2) Chemical
3) Light
4) Nuclear
5) Electromagnetic
Activation Energy
 Activation energy – energy needed to start a
chemical reaction.
Examples:
 Heat – flame to start barbecue
 Mechanical – squeezing ice pack to mix chemicals
Physical Behavior of Matter
1) Matter – anything that has mass and
occupies some volume.
2) Phases of matter:
a) Solid(s) – Definite volume, definite shape,
rigid, fixed, regular, crystalline structure.
b) Liquid(l) – Definite volume, no definite
shape. Takes the shape of its container.
c) Gas(g) – No definite shape and no definite
volume. Takes the shape of the container
uniformly. Gas molecules move in constant,
random, straight-line motion.
Heat and Temperature
Heat
Temperature
 MEASUREMENT of Average
KE energy of a substance.
Units for measuring temp.
1. Kelvin (K)
2. Celsius (0C)
3. Absolute zero = zero k
1. Kinetic energy of
Temp. at which particles stop
molecules.
moving
2. Heat comes from the
Conversion scale (table T)
random movement of
molecules in a substance. K = 0C + 273
3. Heat flows from source to a 0C = K - 273
sink. (high to low)
Convert the following Temperatures
Table T
1. 110oC  K
383 K
2. -35oC  K
238 K
3. 300K  oC
27 oC
4. 75K 
oC
5. 500K 0C
-198
oC
227 oC
K = oC + 273
oC
= K - 273
Convert the following:
use table T
500 C to K
K = 50 +273 =
323K
238 K to 0 C
0C
= 238 – 273 =
-350C
Which cup of coffee has the greatest average Kinetic Energy?
Explain
80 oC
90 oC
70 oC
85 oC
How can we describe matter changing phase when
energy is added or removed?
CONSERVATION of ENERGY - In any chemical reaction or
physical change, energy is conserved.
1) Energy can be changed from one form to another.
(PE  KE)
(KE  PE)
2) Energy can be transferred from one body or system to
another. Heat is transferred from a source to a sink.
3) BUT……THE TOTAL AMOUNT OF ENERGY REMAINS
CONSTANT!!!!!!!
Units of Energy
Joules
1. The joule is the SI unit for energy symbolized (J).
2. The amount of heat or energy needed to increase the
temperature of a substance.
Example:
1. The amount of heat required to raise 1 gram of water by 1
Celsius degree is 4.18 J
2. One KJ = 1000 Joules
Heating and Cooling Curves
System + Surroundings = Universe
1) System – what we want to study or examine.
The reaction in a flask: NH4Cl(s) + H2O(l)
2) Surroundings – everything outside the
system.
3) Universe – System + Surroundings.
Surroundings
Surroundings
System
System
Surroundings
Heat Transfer
Exothermic
1. Temperature of the
surroundings is measured
2. Increase in Temperature
3. Heat is released from the
system to the surroundings
1. Reactions that release
energy.
2. Heat is a product.
Endothermic
Sulfuric acid and sugar
Phases of Matter (water)
1. Temperature of the
surroundings is measured
2. Decrease in Temperature
3. Heat is absorbed from the
surroundings to the system
1. Reactions that absorb
energy.
2. Heat is a reactant.
Exothermic or Endothermic?
Why?
Endothermic or exothermic?
How do you know?
1) KNO3(s) + 34.89 kJ
H2O
K+ (aq) + NO3- (aq)
•Endothermic
•Heat is a reactant (left side of arrow)
•Bonds Broken
•Temperature of surroundings decreases
2) 4Al(s) + 3O2(g)
2Al2O3(s) + 3351 kJ
•Exothermic
•Heat is a product (right side of arrow)
•New Bonds Formed
•Temperature of surroundings increases
Exothermic or Endothermic?
Why?
Exothermic or Endothermic?
Why?
Increase in energy
Solid  Liquid  Gas
Summary: PHASE CHANGES
Energy is absorbed
Particles spread out and
move faster
Melting
SOLID
Vaporizing
SUBLIMATION: SOLID TO GAS
LIQUID
Freezing
Condensing
Energy is released
Particle speed slows down
and the particles move closer
Gas
Exothermic or Endothermic?
Why?
Aim: Does the Temperature of Water change
during a Phase Change?
BBC - GCSE Bitesize: Exothermic and endothermic reactions
1.
Exothermic reactions release energy from the system
(reaction occurs here) to the surroundings.
INCREASE in TEMPERATURE
2. Endothermic reactions have the system absorbing
energy from the surroundings.
DECREASE in TEMPERATURE.
Concept Check:
1) Which sample of water contains particles having
the highest average kinetic energy?
(1) 25 mL of water at 95°C
(2) 45 mL of water at 75°C
(3) 75 mL of water at 75°C
(4) 95 mL of water at 25°C
2) Which change is exothermic?
a)Freezing of water
b)Melting of iron
c)Vaporization of ethanol
d)Sublimation of iodine
3) As the temperature of a substance decreases, the
average KE of its particles
a)Increases
b)Decreases
c)Remains the same
4) In a laboratory where the air temperature is
22°C, a steel cylinder at 100.°C is submerged in
a sample of water at 40.°C. In this system, heat
flows from
(1) both the air and the water to the cylinder
(2) both the cylinder and the air to the water
(3) the air to the water and from the water to
the cylinder
(4) the cylinder to the water and from the water
to the air
Units of Energy
Joules
1. The joule is the unit for energy symbolized (J).
2. The amount of heat or energy needed to increase the
temperature of a substance.
Example:
1. The amount of heat required to raise 1 gram of water by 1
Celsius degree is 4.18 J
2. One KJ = 1000 Joules
Heating and Cooling Curves
Phase Changes at Constant Temperature
Heat of Fusion
Q=mHf
1. Solid  Liquid (endothermic)
2. Solid  Liquid (exothermic)
Heat of Vaporization
Q=mHv
1. Liquid  Gas (endothermic)
2. Liquid  gas (exothermic)
Specific Heat
Heat capacity of a substance
1. Amount of heat needed to increase the temperature of
one gram of a substance by 1 degree Celsius.
2. Because substances have different compositions, each
substance has its own specific heat.
Specific Heat of Water
4.18 J/g 0C or 4.18 J/gK
Which has a higher specific heat, pure water or concrete?
Water = 4.18 joules /g 0C
Concrete = 0.84 joules /g 0C
Specific heat formula
Table T
Q = mcΔT
1) q is the heat gained or lost (J)
2)m is the mass of the heated substance (grams)
3)c is the specific heat capacity (J/g 0C)
4)ΔT (pronounced delta T) is the change in
temperature.
5)ΔT = (Tfinal - Tinitial) in 0C or K
How many joules are absorbed by 30.0g of
water when its temperature is raised from
20.0 0C to 40.0 0C?
Formula for measuring specific
heat
Table T
1. Unit for q =
2. Unit for mass =
3. Unit for specific heat =
4. Unit for change in temperature =
1) Heat is added to a 200. gram sample of H2O(s) to melt the
sample at 0°C. Then the resulting H2O(ℓ) is heated to a
final temperature of 65°C.
a) Determine the total amount of heat required to
completely melt the sample.
b) Calculate the total amount of heat required to raise the
temperature of the H2O(ℓ) from 0°C to its final
temperature.
c) Compare the amount of heat required to vaporize a 200.
g sample of H2O(ℓ) at its boiling point to the amount of
heat required to melt a 200. g sample of H2O(s) at its
melting point.
1) The temperature of a sample of water changes
from 10°C to 20°C when the sample absorbs
418 joules of heat. What is the mass of the sample?
(1) 1 g
(3) 100 g
(2) 10 g
(4) 1000 g
2) A 36-gram sample of water has an initial
temperature of 22°C. After the sample absorbs
1200 joules of heat energy, the final temperature
of the sample is
(1) 8.0°C
(3) 30.°C
(2) 14°C
(4) 55°C
Change in State - YouTube
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