Gases, Liquids, , q
,
and Solids Chapter 4
1
States of Matter
Solids are packed close together
Gases are far apart Gases
are far apart
from each other, lots of empty space.
Liquids are Li
id
farther apart but relatively close
close.
2
STATES OF MATTER
• One major factor responsible for the physical state of substances is intermolecular forces:
• ion‐ion ion‐dipole
• hydrogen bonding London forces
• dipole‐dipole
•
• Physical properties and the behavior of solids, liquids, and gases are the nature of the interaction that g
attracts one particle (atom, ion, or molecule) to another.
3
Changes Between the States of Matter
Changes Between the States of Matter
• Energy is involved.
Energy is involved
• Energy is the capacity to do work.
•
Work is defined in terms of motion of an object.
is defined in terms of motion of an object
•
EEnergy is associated with any change in physical or i
i t d ith
h
i h i l
chemical properties of matter.
4
Types of Energy
Types of Energy
• Kinetic energy is the energy associated with motion
associated with motion.
• Potential energy is energy associated with position or
associated with position or composition. • Faster and larger objects have more kinetic enrgy.
5
Heat and Temperature
Heat and Temperature
• Heat is a form of kinetic energy. It is the energy that spontaneously flows from a hotter body to the colder one.
• Temperature is a measure of the average kinetic energy of the molecules in an object.
j
•
The greater the kinetic energy of the molecules of g,
g
p
something, the higher its temperature.
6
Changes Between the States of Matter
• Involve energy
Gas
Condensation
Evaporation
Sublimation
Liquid
Deposition
Melting
Freezing
Solid
7
Adding Heat
g
8
Phase Change
• Heat
Heat involved in the phase change, q.
involved in the phase change q
• Phase change happens at constant temperature.
•
q = m
q m x ∆H
x ∆Hphase
h
•
•
m = mass of substance
m mass of substance
•
∆Hphase p
g
phase = heat of phase change
• ∆Hvap.
vap ; ∆Hfusion ; ∆Hcond.
cond
• ∆Hvap. = ‐ ∆Hcond.
= 540 cal/g
∆Hvap g
vap
∆Hfusion = 79.8 cal/g
9
Phase Changes
• Heat of fusion is the heat required to melt a solid.
• Heat of vaporization
Heat of vaporization is the heat required to evaporate a is the heat required to evaporate a
liquid.
• Amount of heat involved in the process for 1.00 gram of substance.
• Units are cal/gram 10
Phase Change
• Heat released is negative. Exothermic
• Heat absorbed is positive. Endothermic
Exothermic
•
q= (-)
Solid Liquid
ld
d
Endothermic
q= (+)
11
Temperature Changes
• When the temperature changes energy is involved.
• The amount of heat involved in changing the temperature of a substance depends on:
temperature of a substance depends on:
• a. The type of substance because not all substances absorb heat the same way. (Specific heat)
absorb heat the same way. (Specific heat)
• b. The quantity of matter.
• c. The range of temperature
c The range of temperature
12
Specific Heat
p
• The amount of heat required to raise the temperature of 1 gram of the substance by 1°C
temperature of 1 gram of the substance by 1
C.
• It is a physical property of a substance.
• Units are:
U i
• sp.ht. = cal or Joule
g °C
°
g °C
°
13
Specific Heat of Diff
Different t
Substances
14
Heat
•
q= m x sp.ht. x ∆T
•
∆T = Tfinal – Tinitial
• C
Could be used for a single substance or for chemical ld b
df
i l
bt
f
h i l
reactions.
•
qreleased = ‐(q
(qabsorbed)
Law of Energy Conservation
Law of Energy Conservation
• Used to determined caloric content on food but remember that food calories is 1000 times the calories in this process. Called big C.
15
Heat
• What’s the amount of heat involved in:
• a. raising the temperature of 10.0 grams of water from 19 °C to 100 °C.
• b. decrease the temperature of 10.0 grams of water from 100 °C to 20 °C. •
A 800 cal
A. 800 cal b. b ‐800
800 cal c. 8 cal d. cal
c 8 cal
d ‐8
8 cal e. 0.125 cal
cal
e 0 125 cal
16
Gases
• No definite shape or volume. • Easily compressible. Molecules are free to move.
•
•
•
•
•
Their properties or changes are followed by studying:
a. Pressure
P
b. Volume
c Temperature
c. Temperature
d. Moles • Pressure – force applied per unit of area
17
Pressure
Gas pressure is the force of collisions that take place between the particles and an object (the walls of a container that holds the gas).
18
Units of Pressure
• Atmosphere –
p
pressure p
needed to hold a column of mercury 76 cm tall at sea level.
• 1 atm = 760 mm Hg
• 1 atm 1 atm = 760 torr
760 torr
• 1 atm = 14.696 psi
19
The higher the altitude, lower the atmospheric pressure.
20
Standard Temperature Standard
Temperature
and Pressure (STP)
0 °C and 1 atm
21
The Gas Laws
22
Whether we are dealing with inhaled anesthetic Whether
we are dealing with inhaled anesthetic
gases, the helium in a balloon, or the air in a tire, a gases be a e a s a a e .
all gases behave in a similar manner.
For that reason we called them
Ideal gases
Ideal gases
23
Pressure and Volume
•What happens to volume of gas as pressure increases?
Boyle’s law says that pressure and volume are inversely proportional.
24
Boyle’ss Law
Boyle
Law
25
B l ’ L
Boyle’s Law
P1V1 = P2V2
initial
final
26
K = ºC + 273.15
CHARLES LAW
Absolute Absolute
zero
V1 = V2 T1
T2
27
Pressure and Temperature
p
•How does the temperature affects the pressure of a gas?
Gay‐Lussac’s law says that, for a gas with a constant volume, pressure and temperature are directly related
directly related.
28
Moles and Volume
Moles and Volume
• Moles are quantity of matter in terms of number of molecules.
• Avogadro’s law says that, at a given temperature and pressure, volume and the number of moles are directly related.
29
Ideal Gas Law
• Pressure is: T
•
n (moles)
•
1/V
/
• P nT/V
•
P = nRT
V
• R = 0.08206 L∙atm∙K–1∙mol–1
Ideal gas constant
• PV = nRT
30
Combined Gas Law
• When some of the properties of a sample of gas (P, T, and V) are changed the remaining properties will changed. h
• Based on the ideal gas law equation:
• PV = nRT or n = PV/RT
/
• n1 = n2
P1 ⋅ V1
P2 ⋅ V2
=
T1
T2
initial
final
31
The Combined Gas Law P1 ⋅ V1
P2 ⋅ V2
=
T1
T2
If pressure is constant:
If pressure is constant: V1
V2
=
T1
T2
Charles’ Law
If temperature is constant: P1 ⋅ V1
=
P2 ⋅ V2
Boyle’s Law
32
At Standard Temperature At
Standard Temperature
and Pressure (STP)
0 °C and 1 atm
The volume of 1 mol of any gas is 22.4L
at STP (Molar volume)
at STP (Molar volume)
Molar volume units: 22.4 L/mol
Example: How many moles are 44.8 L of Carbon dioxide?
33
Partial
a t a Pressure
essu e
34
• Dalton’s Law of Partial Pressure states that the total pressure of a mixture of gases is the sum of the partial pressures of its components.
– The partial pressure of a gas in a mixture is the pressure that the gas would exert if alone.
total
total pressure
Pt = Pa + Pb + Pc
gas a
gas b
gas c
35
Gases dissolved in a Liquid
q
• Oxygen and Carbon dioxide are normally dissolved in blood.
• Henry’s Law – The amount of gas dissolved in a liquid is directly related to the pressure of the gas above the liquid.
• P = k x concentration
• k=
k Henry’s constant characteristic of the substance.
H
’
h
i i f h
b
• c= units of conc. will be discussed in the next chapter.
• Important application for anesthesiologists. I
t t
li ti f
th i l i t
• What is better to use as anesthesia, a substance with a large or small k?
a large or small k?
36
Gases dissolved in a Liquid
Gases dissolved in a Liquid
37
Liquids
q
38
Vapor
p Pressure
Due to collisions that take place between particles (
(atoms or molecules), particles at the surface are l l )
i l
h
f
continually evaporating ‐ being “bounced” off into the gas phase. At the same time gas phase molecules are being trapped and converted to liquid.
How is vapor pressure related to temperature? As H
i
l t dt t
t ?A
temperature increases, vapor pressure?
How is vapor pressure related to intermolecular forces? As strength of Inter Forces increases, vapor pressure? 39