CHAPTER 9: LIQUIDS AND SOLIDS Section 9.1 Liquid/Vapor Equilibrium Vaporization – process in which a liquid vapor open container - evaporation continues until all liquid evaporates closed container 1) Liquid evaporate. 2) Vapor particles collect and Condense. 3) Eventually, Rate of Evaporation Rate of Condensation DYNAMIC EQUILIBRIUM
22
Embed
Rate of Evaporation Condensation Chemistry/AP Chem Lectures/… · CHAPTER 9: LIQUIDS AND SOLIDS Section 9.1 Liquid/Vapor Equilibrium Vaporization – process in which a liquid vapor
This document is posted to help you gain knowledge. Please leave a comment to let me know what you think about it! Share it to your friends and learn new things together.
Transcript
CHAPTER 9: LIQUIDS AND SOLIDS
Section 9.1 Liquid/Vapor Equilibrium Vaporization – process in which a
liquid vapor open container
- evaporation continues until all liquid evaporates
closed container
1) Liquid evaporate. 2) Vapor particles collect and
Condense. 3) Eventually,
Rate of Evaporation
Rate of Condensation
DYNAMIC EQUILIBRIUM
Vapor Pressure ● At equilibrium, # molecules/volume is constant. ● Pressure of gas over liquid is constant. ● As long as both liquid and vapor are present,
the pressure exerted bythe vapor is independent of the volume of the container.
Vapor Pressure is dependent on:
a) Characteristics of liquid b) Temperature
Important As long as BOTH liquid and solid are present, the vapor pressure will be constant.
If volume
More liquid will evaporate
Equilibrium Re-establish
Vapor Pressure vs. Temperature
In general, Vapor pressure as Temp. For example, H2O Vapor Pressure of H2O Temp 24 mmHg 25°C 92 mmHg 50°C 760 mmHg 100°C
Higher the Temperature
More Molecules Vaporize
A plot of pressure vs. temperature does not produce a straight line.
This is not a direct relationship!
● Graphs of curves are often difficult to
interpret.
The solution to this problem: Graph manipulated variables of pressure and temperature that will represent a
straight line.
Instead of P vs. T, graph ln P vs. 1/T
● Recall that the general equation of a straight line is y = mx + b (m = slope & b = y-intercept) Here, y = 1n P x = 1/T m = -∆Hvap/R Therefore,
ln P = - TR
H vap 1 + b
If 2 different temps are evaluated:
at T2 : ln P2 = -2
1
TR
Hvap
+ b
at T1 : ln P1 = -1
1
TR
Hvap
+ b
Clausius- Clapeyron Equation
121
2
12
11lnlnln
TTR
H
P
PPP
vap
where R = 8.314 J/K●mole
Boiling Point
Vapor Pressure of Liquid
EQUALS
Pressure Above the Surface of the Liquid
Normal Boiling Point = the temperature a liquid boiling at 1 atm of pressure above the liquid.
● The boiling point of any liquid can be
lowered by reducing pressure above liquid. Varies with altitude.
Critical Temperature – the
temperature above which the liquid state of a pure substance cannot exists regardless of the pressure.
Critical Pressure – the pressure that be applied to cause the condensation of a pure liquid at the critical temperature.
Section 9.2 Phase Diagrams ● Phase Diagram – a graphical way to summarize the conditions under which the different states of a substance are stable.
Water’s Phase Diagram
Liquid
Vapor
Solid
Pre
ssu
re (
atm
)
Temperature (oC)
Critical Point
Normal Boiling Pt.
1.00
.0060
217.75
NormalFreezing Point
Triple Point
0.00 0.01 100.00 373.99
*Not to Scale
D C
B
A
● Melting-Point Curve - Observe the solid/liquid states at different
pressures. - Along the curve both phases are in
equilibrium.
- Special Note: When conditions indicate that a substance is in the liquid or solid state, the vapor of that substance is also present (in equilibrium).
- The question one should ask is “How much
vapor is present?”
- Answer: It depends!!! Sublimation – Transformation of a solid
directly into a vapor. Melting Point – The opposite process
freezing.
Triple Point – the point on a phase diagram representing the temperature and pressure at which three phases of a substance coexist in equilibrium.
● Most water solutions are also nonconductors ● Some polar molecules form ions when they react with H2O
conduct electricity
For example: HF(g) H+1 (aq) + F-1(aq)
2. Generally, molecular compounds are insoluble in water. 3. They have low melting & boiling points. ● Many are gases (N2, O2, …) ● Some are liquids with melting points <25oC (like H2O, mp = 0oC).
● Some are solids with melting points <300oC (like I2, mp = 114oC).
The boiling point and melting point of molecular substances is
directly related to the strength of their INTERMOLECULAR ATTRACTIVE FORCES
among molecules.
Intermolecular Forces 1. (London) Dispersion Forces ● Found in all molecular substances.
involves a temporary or induced dipole.
Consider the H2 molecule
● Nonpolar bond equal sharing of electrons
For an instant, the electrons within the molecule can concentrate closer to one atom in the molecule.
● Produces a +/ - (dipole) within the
molecule. ● This temporary dipole induces a similar dipole
within another molecule. ● These temporary dipoles result in the two molecules attracting each other.
This attraction is the Dispersion Force! The attraction is dependant on:
1. The # of electrons in the molecules involved.
2. The ease of the electrons in the molecules to be dispersed within the individual molecules.
Larger atoms/molecules
Produce Greater Temporary Dipoles.
In general,
As Molar Mass ,
The dispersion forces ,
The bp & mp of nonpolar
molecules .
2) Dipole Forces ● These interactions occur in polar
molecules. ● The +/ - (dipole) of one polar molecule
lines up with +/ - (dipole) of another polar
molecule (opposites attract). The greater the dipole moments (the
measure of the polarity of a molecule) of the molecules, the stronger the attractive force. ● This interaction (attraction) really only works when the molecules are close together. ● When the molecules are in the gas phase,
the dipole forces of attraction are negligible (as is the case for dispersion forces).
3) Hydrogen Bonding
This attraction occurs in polar molecules, HOWEVER, only in molecules that have X – H bonds where X = N, O, or F.
This attractive force is an unusually strong dipole force of attraction.
Why is the hydrogen bonding such a powerful attractive force?
2 – Reasons: 1. There is a large difference in the
electronegativity of the X and H H(2.2) F(4.0)
H(2.2) O(3.5) H(2.2) N(3.0)
The H – atom almost behave as a “naked” proton.
2. The H – atom is very small.
The lone pair of electrons on F, O, and N can get really close to H.
Let’s look at the pattern of boiling points.
bp(oC) bp(oC) bp(oC)
NH3 -33 H2O 100 HF 19
PH3 -88 H2S -60 HCl -85
AsH3 -63 H2Se -42 HBr -67
SbH3 -18 H2Te -2 HI -35
Note the effect of hydrogen bonding in the first row of boiling points.
IMPORTANT Although these intermolecular force are very important, they are very weak compared to a covalent bond.
Section 9.4: Network Covalent, Ionic, and Metallic Solids
● Most Molecular substances are gases or
liquids at room temperature. ● Most NON-molecular substances (network
covalent, ionic, and metallic) are solids at room temperature.
1. Network Covalent Solids
These solids are made of atoms joined by a continuous network of covalent bonds.
In general, these solids are: a. High melting (over 1000oC)
- In order for this type of solid to melt, bonds need to be broken
- Remember: When molecular solids to
melt, only interactions need to be broken!
b. This type of solid is typically insoluble in all common solvents.
- Why? Bonds need to be broken!
NOT EASY TO DO!!!
c. These solids are poor conductors of electricity.
- Why? No mobile electrons are available
Example: Carbon 2 types of solids exist: Graphite & Diamond Both have very high melting points >3500oC.
Ionic Solids These solids are held together by very strong electrostatic attractive forces (ionic bonds).
1. They are composed of cations/anions. 2. They are non-volatile (do not become gases
very easily). 3. They are high melting (600 – 2000oC). 4. They do not conduct electricity (they only do
when they form aqueous solutions or they are molten).
5. Many (but not all) are soluble in water.
Metallic Solids A structural unit of electrons and metal cations.
Positive metal ions anchored in position with electrons moving around from one metal ion to another.
Metals are highly conductive - Very mobile electrons – metals have very low
electronegativities.
● Metals have a high thermal conductivity - Very mobile electrons vibrate
● Metals are ductile and mobile.
● Metals have high luster.
- Electrons within a metal can absorb and emit light energy very easily.