Transcript Soap Bubbles

1
 Joseph Black – explained heat in terms of a
fluid (Lavoisier had called this fluid “caloric”
from Latin word for heat.
 Count Rumford – friction could convert
mechanical energy into heat (motion as cause)
 John Dalton – idea of atoms
2
 James Prescott Joule – tried to find the
mechanical equivalent of heat (where a given
amount of energy produces the same amount
of heat)
 James Clerk Maxwell – developed a solid
explanation showing relationship between
motion of atoms and heat.
3
 Heat flows from hot to colder areas due to a
temperature difference only (till thermal
equilibrium is established).
 Heat is a form of internal energy which is
transferred from one object to another due to
a difference in temperature between the
objects.
4
 The heat content of a substance is the total
energy of all the particles of that substance.
 The total energy combines both kinetic and
potential energies.
5
 The temperature of a body of matter is a
measure of the average kinetic energy of the
random motion of its particles.
 Temperature is that property of a substance
which determines whether it is in thermal
equilibrium with another object.
6
 Thermal equilibrium is the situation in which
no heat moves from one object to another
(they have the same temperature).
7
 Thermometers work on idea of thermal
expansion = the amount of expansion or
contraction is always the same for the same
increase or decrease in temperature.
 3 types: gas (air), liquid (Hg & alcohol), solid
(bimetallic)
 Know creation and calibration ideas
8
K
= °C + 273.15
 °C
= K – 273.15
 °F
= 9/5 °C + 32
 °C = 5/9 (°F - 32)
9
Fahrenheit Celsius
Kelvin
Boiling Pt. H2O 212
100
373
Body temp
98.6
37
310
Freezing
32
0
273
Coincidence -40
-40
233
Absolute zero -460
-273
0
10
 15° calories
= the amount of heat needed to
raise 1 gram of water from 14.5° to 15.5° C at 1
atmosphere of pressure
 kilocalorie = kcal or Calorie = 1000 cal
 1 calorie = 4.185 Joule
 1 kcal = 4185 Joule
11
 Specific heat is the amount of heat needed to
raise 1 gram of water 1° C at 1 atmosphere of
pressure
 What is the degree change if 1 calorie of heat
is added to 1 gram samples of:
water
helium ice gold
12
Q
= m x ∆t x cp
where
Q = heat flow
m = mass
∆t = change in temperature
cp = specific heat
13
 The amount of heat lost by a substance is
equal to the amount of heat gained by the
substance to which it is transferred.
 m x ∆t x cp = m x ∆t x cp
heat lost
heat gained
14
 Specific heat – how well a substance resist
changing its temperature when it absorbs or
releases heat
 Water has high cp – results in coastal areas
having milder climate than inland areas
(coastal water temp. is quite stable which is
favorable for marine life).
15
 Organisms are primarily water – thus are able
to resist more changes in their own
temperature than if they were made of a liquid
with a lower cp
16
When calories of heat are added to water there is a
small change in temperature because most of the
heat energy is used to disrupt hydrogen bonds before
water molecules can begin to move faster.
 Temp. of water drops – many additional hydrogen
bonds form, releasing a considerable amount of heat
energy.

17
18
Metal
Specific Heat
Thermal
Conductivity
Density
cp (cal/g° C)
k (watt/cm K) (g/cm3)
Electrical
Conductivity
1 E 6/ Ώm
Brass
0.09
1.09
8.5
Iron
0.11
0.803
7.87
11.2
Nickel
0.106
0.905
8.9
14.6
Copper
0.093
3.98
8.95
60.7
Aluminum
0.217
2.37
2.7
37.7
0.0305
0.352
11.2
Lead
19
 Conduction – faster vibrating particles collide
with less energetic neighbor and transfer
energy to it
 Convection – motion of hot fluid, displacing
cold fluid in path setting up convection current
 Radiation – energy transmitted by
electromagnetic waves
20
 From 0° to 4° the volume of water in a sample
decreases (the greatest density is at 4° c)
 Ice floats: body of water freezes from top
down allowing life underneath to continue
21
 Water mlcl can participate in 4 bonds with
other water mlcl (solid mlcl can have as many
as a dozen bonds with surrounding mlcl
resulting in a more compact substance).
 The spaces between mlcl in ice are greater
than the same spaces in liquids.
22
 Density of ice increases from 0 to 4° as large
clusters of mlcl break into smaller clusters that
takes up less space in the aggregate. Above 4°
normal thermal expansion is seen with a
decrease in density.
23
 Heat of Fusion – amount of heat needed to
change solid to liquid at its melting point
 Heat of Vaporization – heat needed to
change liquid to gas at boiling point
 Heat of Sublimation – heat to change a solid
to gas
 Heat of Condensation – heat released when
gas condenses to a liquid
24
 Matter is defined as any material that has
mass, occupies volume, and exhibits inertia
(resistance to movement).
25
 Solids – definite shape and volume,
resist deformation
 Very close spacing of particles
 Particles appear to vibrate around fixed points
 Particles vibrate faster at higher temp.
26
Crystalline – particles arranged in regular, repeated
patterns (long-range order) example: NaCl (s)
 Amorphous – solids that lack the definite
arrangement of crystalline solids (have short-range
order)

 Examples: pitch, glass, plastics
27
 Definite volume, resist compression, take
shape of container
 Greater spacing between particles, particles
appear to travel in straight line paths between
collisions but appear to rotate and/or vibrate
about moving points
28
 Have no definite shape or volume, take shape
and volume of container
 Can be compressed or dispersed, particles
vibrate very rapidly, relatively far apart
 There are no intermolecular forces holding
particles together
29





Very high temperature ionized gas
No fixed volume or shape
Most are mixtures that are not easily containable
Particles are electrically charged and of low density
Example: the Milky Way
30
 Energy – having the ability to do work
 Work – a push or pull over a distance
 Force – a push or pull
 Momentum – mass x velocity
 Linear momentum of a moving body is a
measure of its tendency to continue in motion
at a constant velocity
31
 Potential Energy – the energy a body
possesses by virtue of its position,
composition, and/or condition
 P.E. is the stored energy
 P.E. = mass x gravity x height
32
 K.E. = the energy of motion
 K.E. is conserved in all elastic collisions
 K.E. = ½ m v2
(m = mass, v = velocity)
 Heat energy flows from hot objects to cooler
ones by transfer of K.E. when particles collide
(conduction).
33
 P.E. forces that hold mlcl together and in
correct position in solids.
 P.E. forces that hold mlcl together in liquids.
 These forces are between mlcl.
 Gases have enough K.E. to prevent formation
of these forces.
34
 Gases are mlcl in continuous motion.
 An increase in temp. increases speed thus
increasing K.E.
 All gases are compressible
 Gases display diffusion
 Gases can be liquified (called liquifaction)
35
 Nothing escapes or enters system
 All mlcl in motion (have K.E.)
 Mlcl exert uniform pressure against walls of
container
 Mlcl exert pressure on other mlcl as they
collide, push, bounce off other mlcl
36
 Pressure = Force / Area
 Atmospheric Pressure = cumulative net force
per area generated by weight of our
atmosphere
 Values = 14.7 lb/in2, 101.3 kPa,
760 mm of Hg, 1 atm, 1033 g/cm2
37
 The pressure a gas exerts on the walls of its
container is the sum of the forces acting (= the
frequency of collisions plus the force of each
collision) due to the random collisions of near
limitless numbers of moving molecules.
38
Inelastic collisions – the normal type in which
objects lose energy and slow down
 Elastic collisions – particles bounce off, exchange
energies but there is no loss of energy (energy is
conserved but may be redistributed)

39
Energy is conserved only in elastic collisions
Momentum is conserved in every collision in which
there is no friction.


40
Gay-Lussac
P≈T
Holding volume constant, the pressure is
proportional to the absolute temp.
 P1 / T 1 = P 2 / T 2


41
 Boyle’s Law
V ≈ 1/P
 If the temp. is held constant, the volume
of a gas varies inversely with the pressure
 P1V1 = P2V2
42
 Charles’ Law
V ≈ T
 If the pressure is held constant, the
volume of a gas is proportional to its
absolute temp.
 V1 / T1 = V2 / T2
 For every degree increase in temp. the
volume increases by 1/273 of its original
volume
43
 Combined Gas Law:
P1V1/T1 = P2V2/T2
 Ideal Gas Law:
PV = n R T (where n = # moles, and
R = gas constant)
44
Chemical properties are those properties
of a substance that can be determined by
a chemical test. They are seen by the
material’s tendency to change, either
alone or by interaction, and in doing so
form different materials.
45
 Does the substance support
combustion? Burn itself?
 How does it react with acids? With
oxygen? With electricity?
 Examples: alcohol burns, wood decays,
sodium explodes and burns in water
46
 Physical properties are those properties
used in identifying substances when we
use our senses. These do not require
chemical analysis.
47
 Color, hardness, density, texture, magnetic
attraction, solubility, taste, light transmission,
viscosity, refractive index, specific heat, boiling
point, melting-freezing point, odor, expansioncontraction coefficients
48
 This is a change in the physical properties of a
substance without a change in the chemical
composition.
 The arrangement of molecules may be
changed but the molecular makeup remains
the same.
 These changes involve intermolecular forces
which increase or decrease during the change.
49
+ heat  steam (100° C)
36 g 25 920 cal
36 g
 Ice (0° C)
 Steam (100° C)  ice (0° C) + heat
36 g
36 g
25 920 cal
50
 The molecular makeup (specific
arrangement of atoms) is changed,
resulting in new substances being formed
and energy changes occurring.
 Two types: exothermic and endothermic
51
 Any chemical change that releases
energy
 The amount released must be greater
than the amount used to start reaction
 Bond making is exothermic (energy is
released into surroundings
52
 Oxidation – wooden splint burning
(giving off light, heat, CO2, H2O
 Burning H2 in air, body reactions,
dissolving metals in acid, mixing acid and
water, sugar dehydration, plaster of Paris
in water
53
 Any chemical change that absorbs
energy
 Energy continues to be absorbed as long
as reaction continues
 Bond breaking is endothermic (energy is
absorbed from surroundings
54
 Electrolysis (breaking water down into
H2 and O2 by running electricity in it)
 Photosynthesis, pasteurization, canning
vegetables
 2H2O + energy
4 g + 32 g 36 g 136 600 cal
 2H2O + energy  2H2 + O2
36 g 136 600 cal 4g + 32 g
 2 H2 + O2
55
 Physical change – strength of intermolecular
forces increased or decreased
 Chemical change – bonds formed or broken
 Energy absorbed – bonds broken or
intermolecular forces overcome
 Energy released – bonds formed or
intermolecular forces strengthened
56
 Dry ice sublimates
 CO2 + H2O + sunlight  glucose
 Air in heated tire expands
 Burning coal
 Water frozen into ice
 Acid dissolves metal
57
 Sublimation is the direct change of a solid to
a gas
 Deposition is the change of a gas to a solid
 Examples: moth balls
(naphthalein),paradichlorobenzene, camphor,
iodine crystals, CO2 fire extinguishers
58
 Melting-freezing point – this is the same
temperature at which a pure substance can
change into solid from liquid or solid into liquid.
 The solid-liquid phases are in equilibrium.
59
 When heat is added to a solid, the temp. will
increase till it reaches the melting-freezing
point. It will remain at that temp. until all the
solid has melted and then the temp. can rise
again according to its specific heat.
60
 Boiling point is defined as the temperature at
which the liquid’s vapor pressure is equal to
outside (atmospheric pressure usually).
 When the vapor pressure equal atmospheric
pressure as many mlcl are leaving the surface
as are re-entering the surface of the liquid.
61
 Boiling point varies with elevation.
 Cooking times must adjust due to elevation.
 Pressure cookers can cook food more rapidly
due to increased pressure, resulting in high
boiling points (which cooks food faster).
62
63
64



Study of heat changes that take place in a
change of state or chemical reaction
If heat is released, process is called
exothermic
If heat is absorbed, process is called
endothermic
65
Heat is energy transferred from one object to
another due to a difference in temperature.

We measure the temperature change that
accompanies heat transfer.

We have to measure the temperature change of
the surroundings (the solvent, container,
atmosphere).

The system is the reactants and products of the
reaction.

66


When a system releases heat to
surroundings, the temperature of the
surroundings increases (exothermic). An
example would be combustion of propane in
a barbecue grill.
When a system absorbs heat, the
temperature of the surroundings decreases
(endothermic). An example would be
melting ice.
67
The amount of heat transferred depends on the
energy stored in each substance. This stored
energy is called heat content or enthalpy and is
represented by H.

∆H = qp
Enthalpy = heat transferred


qp = m ∙ ∆t ∙ cp
68
cp reflects that ability of a substance to absorb
heat (defined as the amount of heat needed to
raise the temperature of 1 gram by 1 degree Celsius)

cp of water = 1.00 cal/g° C or 4.185 J/g ° C


In most situations it is the temperature change of
the surroundings that is measured (which equals
the heat releases/absorbed from the reaction itself)
69



For the increase in temperature of the
surroundings, heat must be released by the
system.
The surroundings increase is positive while
the heat release by the system must be
negative.
Exothermic reactions always have negative
values.
70



Heat absorbed by system results in
temperature decrease for surroundings
(negative quantity).
Heat absorbed by system must have
positive value.
Enthalpy change for endothermic is always
a positive value.
71


Heats of solution deal with the process of a
solute dissolving in a solvent.
In the case of an ionic solute, there are two
processes:
 Energy to break apart the ionic bonds in the crystal lattice
(called crystal lattice energy)
 Energy released when the free ions form attractive forces
with water molecules (called heat of hydration)
 The heat of solution is the sum of these two effects
72





Crystal lattice energy of KCl:
KCl (s)  K1+ (g) + Cl1- (g)
∆H = + 167.6 kcal
Heat of hydration of KCl:
K1+ (g) + Cl1- (g)  K1+ (aq) + Cl1- (aq)
∆H = - 163.5 kcal
Overall: KCl (s)  K1+ (aq) + Cl1- (aq)
∆H = + 4.1 kcal - Endothermic Reaction
73







NH4NO3
NaOH
KNO3
KClO3
KOH
NaCl
NaC2H3O2
+ 6.1
- 10.6
+ 8.0
+ 9.89
- 13.77
+ 0.93
+4.085
(endothermic)
(exothermic)
(endothermic)
(endothermic)
(exothermic)
(endothermic)
(endothermic)
74
Usually these reactions are exothermic but adding
vinegar to baking soda is slightly endothermic.

The neutralization reaction is slightly exothermic.
HC2H3O2 (aq) + NaHCO3 (aq)  CO2 (g) + NaC2H3O2 (aq) + H2O (l)

net bond formation

Evaporation of the liquid occurs as the CO2
escapes from solution. Evaporation absorbs heat,
cooling the liquid (along with expansion of bubbles
also helps to cool the surroundings) = net result is
endothermic reaction.
75






Mixing a strong acid with water is exothermic.
Breaking chemical bonds requires energy.
Forming chemical bonds releases energy.
HCl (g)  H1+ (aq) + Cl1- (aq)
It looks like heat would be absorbed because the
bond between the H and Cl is broken. The
hydrogen reacts with water to form a complex:
H3O·(H2O)+ n (where n is between 1 and 9).
This hydration makes the overall reaction strongly
exothermic.
76