liquids and solids
TRANSCRIPT
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Liquids and Solids Polarity of Molecules In the previous chapter, learned how
atoms bond to each other and how these bonds affect the shape of the molecules they form.
We also learned how to determine
the type of bond that exists between two atoms.
The shape of molecules, and the type of bonds they possesses, affect the distribution of electron density within them.
When a molecule has an
unsymmetrical distribution of e- within the molecule, we say that the molecule is POLAR.
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When a molecule has a symmetrical
distribution of e- within the molecule, we say that the molecule is NONPOLAR.
The covalent bonds that make up molecules come in two types:
1) Covalent Bond - a bond resulting in an equal sharing of e-. - the two atoms in the bond have identical electronegativities.
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2) Polar Covalent Bond -a bond resulting in an unequal sharing of e-. - the two atoms in the bond have different electronegativities.
- the more electronegative atom in the bond will have a partial negative charge ( -).
- the less electronegative atom in the bond will have a partial positive charge ( +).
- the existence of +and - in a
bond is called a DIPOLE. When a molecule has an unsymmetrical distribution of e- (it is POLAR), an overall +and - exists
within the molecule. When this occurs, we say that the molecule has a DIPOLE.
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The measure of how much of a dipole a molecule possesses is called the Dipole Moment. The greater the dipole moment, the more polar the molecule.
However, merely having polar bonds does not necessarily mean that the molecule is polar.
The geometry of a molecule must be
considered when determining a molecules’ polar character. Linear Molecules:
H F F Be F
H F
Polar
BeF F
Nonpolar
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Trigonal Planar:
B
Cl
Cl
Cl
B
Cl
F
Cl
PolarNonpolar Tetrahedral Molecules:
CCl ClCl
F
PolarNonpolar
Cl
CCl ClCl
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INTERMOLECULAR FORCES 1) Dipole Forces These interactions occur in polar
molecules. This intermolecular force is only 1%
as strong as a covalent bond.
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.
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When the molecules are in the gas phase, the dipole forces of attraction are negligible.
2) 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)
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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 bp bp
NH3 -33oC H2O 100 oC HF 19oC
PH3 -88oC H2S -60oC HCl -85oC
AsH3 -63oC H2Se -42oC HBr -67oC
SbH3 -18oC H2Te -2oC HI -35oC
*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.
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3) (London) Dispersion Forces Found in all molecular-nonpolar
substances.
involves a temporary or induced dipole.
Consider the H2 molecule Nonpolar covalent 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.
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● 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.
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MOLECULAR INTERACTIONS To begin our discussion, let’s talk about the simplest phase of matter – the gas phase.
When these particles pass each other, 3 possible scenarios can occur.
1) They are indifferent to each
other. 2) They are attracted to each
other. 3) They are repelled by each
other. Question: In order for an ice cube to form, which of these options must exist? Answer: ATTRACTION!!!
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3 Types of attractive forces exist. They are (in order of weakest to strongest): 1. Dispersion Forces 2. Dipole/Dipole Forces 3. Hydrogen Bonding Force (a stronger version of Dipole/Dipole Forces) All 3 of these attractive forces can exist in all three phase of matter. Let’s force our discussion on the Liquid and Solid.
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Liquids & Solids Liquid State
Much more ordered than a gas. Greater attractive force than a gas.
Keeps molecules together. Attractive Forces:
- Hydrogen-Bonding - Dipole/Dipole - Dispersion
Fluidity – can flow and assume the shape of the container. Density – most liquids are 1000X more dense than a gas. -They are 10% as dense as a solid.
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Compressibility If 1000 atm of pressure is applied to a gas – the volume can be decreased by 99.9%.
For liquids (1000 atm) will decrease volume by 4%.
Diffusion Liquid can diffuse Much more
slowly than a gas. Affect by Temp. Surface Tension
A force that tends to pull adjacent parts of a liquid’s surface together.
Result: Decrease surface area.
H2O
Air
Surface Why? Intermolecular Forces
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Capillary Action (Related to Surface Tension)
Attraction of a liquid to a surface Water in a tube Upward pull
Evaporation vs. Boiling (Vaporization) Evaporation = Process in which a liquid is converted into a gas by particles leaving the surface of the liquid. (Not boiling) Boiling = A process as above, however, the vapor pressure of the liquid is equal to atmospheric pressure.
*This varies with altitude.
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Solidification (Freezing) Occurs when heat is lost by a liquid. Solids
The particles in a solid are much more ordered than in a liquid
Packed together.
Vibrate in place.
Very dense and is NOT compressible.
Very low rate of diffusion
- If a zinc plate and copper plate are clamped together for a very long time, a few atoms will diffuse into each other.
Diffusion occurs millions of
times slower in solid that liquids.
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Types of Solids 1) Crystalline Solids
Exist as single or group of atoms 3D arrangement of particles in a crystal is called a Crystal Structure
7-types of crystal structures exist:
Triclinic
Hexagonal
Cubic Tetragonal Orthorhombic
Monoclinic Rhombohedral
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4 Types of Crystalline Solids exist: Ionic Molecular Covalent & & Metallic Network Covalent A) Ionic Crystalline Solids - Made of Metal/Nonmetals - Strong bonding forces between + & - ions - Hard & Brittle - High Melting Points B) Metallic Crystalline Solids - Made only of metals - Metals atoms surrounded by a “Sea of Electrons” - Good conductors of electricity - Melting Points vary
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C) Network Covalent Crystalline Solids - Made of single atoms (nonmetals) - Atoms covalently bonded to each other. - Bonding makes up a network -Very hard & brittle - Non-conductor - High melting point D) Molecular Covalent Crystalline Solids - Covalently bonded molecules held to each by intermolecular forces. - Tend to have low M.P. - Vaporize easily - Tend to be soft - Good insulators
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2) Amorphous Solids Not crystalline–no regular shape Very slow flowing Example: foam, glass, plastic
Equilibrium & Phase Changes
Equilibrium A dynamic condition in which 2 opposing changes occur at equal rates in a closed system. Example: A 2-lane highway with vehicles traveling at 55 mph in either direction.
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WaterConsider a dry flask in which room temperature water is added.
One phase exist in the flask.
Phase = any part of a system that has
uniform composition and properties.
When the flask is capped, and the some surface water molecules have enough energy, the liquid water will overcome the intermolecular forces and become a gas.
Liquid + heat energy vapor
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If some of the water vapor falls down to the surface of the water, it will condense.
vapor Liquid + heat energy
After some time, these two process will occur at the same rate – equilibrium.
These two processes can be
summarized with the following equation: Liquid + heat energy vapor
The ideas of equilibrium were first Understood and explained in 1888
by the French chemist, Henri Louis Le Chatelier.
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Le Chatelier develop of the following principle:
Le Chatelier Principle - When a
system at equilibrium is disturbed by a stress, the system will achieve a new equilibrium position in order to counter act the stress.
Consider this principle for this equilibrium process:
Liquid + heat energy vapor
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When all three phases of matter are considered together, the following equilibrium situation must be considered:
Solid Liquid GasVaporization
Freezing
Melting
Sublimation
Deposition
Condensation
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Equilibrium Vapor Pressure The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. Greater equilibrium vapor pressure when the intermolecular forces decrease. Volatile/Nonvolatile Liquids Volatile Liquid Nonvolatile Liquid
Weak Intermolecular
Forces
StrongIntermolecular
Forces
EvaporateEasily
Do NotEvaporate
Easily
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Heat, Energy & Phase Changes
Energy - the potential or capacity to do work or produce heat.
Kinetic Energy – energy associated
with motion. KE = ½ mv2
Potential Energy – energy of an
object due to its position.
PE = mgh
SI Unit: joule (J)
Non-SI Unit: calorie (cal) 1 cal = 4.184 J
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Heat Flow Heat – energy that is transferred from one object to another due to a difference in their temperatures. Heat flow accompanies all physical
and chemical processes.
All physical and chemical processes occur within the universe.
The universe can be divided into 2 parts: System & Surroundings
System Surroundings
THE UNIVERSE
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Direction & Sign of Heat Flow Heat flow is given the symbol “q”.
If q is “+”, heat flows from the surroundings to the system.
endothermic
Temperature of the system If q is “-”, heat flows to the
surroundings from the system exothermic
Temperature of the system
When a chemical or physical process occurs, we must always be consistent with the Law of Conservation of Mass and Energy.
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Heat flow can be calculated for a chemical or physical process, however, let us first discuss how heat flows to and from a system.
The heat flow, q, can be calculated.
For a pure substance,
q = m x c x T
c = specific heat of the system = amount of heat needed to raise the
temperature of 1 gram of pure substance 1oC Unit = J/g oC
m = mass of the sample
(different for all substances) Unit = g
T = Tfinal - Tinitial
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Unit = oC Specific Heat of Several Pure
Substances at 20oC
Substance Specific Heat (J/g oC)
Air(g) 1.00
Al(s) 0.895
C(graphite) 0.711
C(diamond) 0.502
CO2(g) 0.832
Cu(s) 0.387
Au(s) 0.129
Fe(s) 0.448
Pb(s) 0.128
Ag(s) 0.233
H2O(l) 4.18
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Calorimetry
To measure the heat flow for a chemical or physical process, a device called a calorimeter is used to measure the heat transferred to or from a chemical reaction or physical process by making the following assumption:
qprocess = -qcalorimeter
So,
qprocess = s x m x T
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Example #1: Calculate the quantity of heat (in J) needed to warm water from 25°C to 95°C for making tea. The cup will hold 250 g of water. Answer: Specific Heat (H2O): s = 4.18 J/g°C Mass: m = 250 g ∆T = 95°C - 25°C = 70°C SO, q = (s)(m)(∆T) q = (4.18 J/g°C)(250 g)(70°C) q = 73000 J
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Example #2: 100.00 g of water at 22.4°C is placed in a calorimeter. A 75.25 g sample of aluminum is removed from hot water at a temperature of 99.3°C and quickly placed in the calorimeter. The substances reach a final temperature of 32.9°C. Determine the specific heat of the aluminum. Answer: Step #1: Determine the heat transferred to the water from Al. Step #2: Determine the specific heat of the aluminum.
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Step #1 q = (s)(m)(∆T) s = 4.18 J/g°C m = 100.0 g ∆T = 32.9°C - 22.4°C = 10.5°C q = (4.18 J/g°C)(100.00 g)(10.5°C) q = 4390 J Step #2 q = (s)(m)(∆T) s = ??? m = 75.25 g ∆T = 99.3°C - 32.9°C = 66.4°C q = 4390 J 4390J = (s)(75.25 g)(66.4°C) s = 0.879 J/g°C
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Energy Changes Accompanying a Change in State
During a change in phase, the energy of a substance changes.
Solid Liquid Gas
Freezing Condensation
Melting Evaporation
Heat isabsorbed
Heat isabsorbed
Heat isLost
Heat isLost
What temperature does ice (H2O) melt?
What temperature does H2O freeze?
What temperature does H2O evaporate?
What temperature does H2O condense (liquefy)?
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Regardless of the substance, the melting & the freezing points are the same.
Also, regardless of the substance, the evaporation & the condensing points are the same.
The heat needed to vaporize a liquid into a gas is called the
Heat of Vaporization.
The heat lost when a gas turns into a liquid is called the
Heat of Condensation.
Heatvaporization - Heatcondensation
Similarly, the heat needed to melt a solid into a liquid is called the
Heat of Fusion.
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The heat lost when a liquid turns into a solid is called the
Heat of Solidification.
Heatfusion - Heatsolidification
These relationships can be represented graphically:
Tem
per
atu
re
Boiling
Point
MeltingPoint
Heat Added
Heat of Fusion
Heat of Vaporization
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The heats of fusion & vaporization are expressed in the units of J/g.
Heats of Fusion & Vaporization Substance Heat of
Fusion (J/g) Heat of
Vaporization(J/g)
Cu 205 4726
Ethanol 109 876
Au 64.5 1578
Pb 24.7 858
Ag 88 2300
H2O 334 2260
What is the heat transferred when 20.0 g of ice melts (s = 4.18 J/g oC)?
Let’s use what we have learned:
q = (s)(m)(ΔT), s = 4.18 J/g oC m = 20.0 g ΔT = 0 oC (at melting, no ΔT occurs!) q = (4.18 J/g oC)(20.0 g)(0) = 0 J *** Does this make sense?! ***
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Melting requires no heating??!! *** THAT MAKES NO SENSE! ***
Therefore, calculating the heat transferred during a phase change must be calculated differently.
HeatTransferred
(heat of phase )(mass)
Example #1: Calculate the heat to change a 250.0 g sample of liquid water into steam a 100.0oC. Answer: (Hint: What phase change is happening?) q=(heat of phase change)(m) q=(heat of vaporization)(m) =(2260 J/g)(250.0 g) =565000 J
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Example #2: Calculate the heat needed to melt a 15 g ice cube at 0.0oC to give water, also at 0.0oC. Answer: q=(heat of phase change)(m) q=(heat of fusion)(m) =(334 J/g)(15 g) =5.0 x 103 J = 5.0 kJ