Transcript Introductory Chemistry, 2nd Edition Nivaldo Tro
Chapter 12 Liquids and Solids
2006, Prentice Hall
Interactions Between Molecules
• many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction your taste and smell organs work because molecules in the thing you are sensing interact with the receptor molecule sites in your tongue and nose • in this chapter we will examine the physical interactions between molecules and the factors that effect and influence them 2
The Physical States of Matter
• matter can be classified as solid, liquid or gas based on what properties it exhibits
State
Solid
Shape
Fixed
Volume
Fixed
Compress
No
Flow
No
Liquid
Indef. Fixed No Yes
Gas
Indef. Indef. Yes Yes •Fixed = keeps shape when placed in a container, •Indefinite = takes the shape of the container 3
Structure Determines Properties
• the atoms or molecules have different structures in solids, liquid and gases, leading to different properties 4
Properties of the States of Matter Gases
• low densities compared to solids and liquids • fluid the material exhibits a smooth, continuous flow as it moves • take the shape of their container • expand to fill their container • can be compressed into a smaller volume 5
Properties of the States of Matter Liquids
• high densities compared to gases • fluid the material exhibits a smooth, continuous flow as it moves • take the shape of their container • keep their volume, do not expand to fill their container • can not be compressed into a smaller volume 6
Properties of the States of Matter Solids
• high densities compared to gases • nonfluid they move as entire “block” rather than a smooth, continuous flow • keep their own shape, do not take the shape of their container • keep their own volume, do not expand to fill their container • can not be compressed into a smaller volume 7
The Structure of Solids, Liquid and Gases
8
Gases
• in the gas state, the particles have complete freedom from each other • the particles are constantly flying around, bumping into each other and the container • in the gas state, there is a lot of empty space between the particles on average 9
Gases
• because there is a lot of empty space, the particles can be squeezed closer together – therefore gases are compressible • because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow 10
Liquids
• the particles in a liquid are closely packed, but they have some ability to move around • the close packing results in liquids being incompressible • but the ability of the particles to move allows liquids to take the shape of their container and to flow – however they don’t have enough freedom to escape and expand to fill the container 11
Solids
• the particles in a solid are packed close together and are fixed in position though they are vibrating • the close packing of the particles results in solids being incompressible • the inability of the particles to move around results in solids retaining their shape and volume when placed in a new container; and prevents the particles from flowing 12
Solids
• some solids have their particles arranged in an orderly geometric pattern – we call these
crystalline solids
salt and diamonds • other solids have particles that do not show a regular geometric pattern over a long range – we call these
amorphous solids
plastic and glass 13
Why is Sugar a Solid But Water is a Liquid?
• the state a material exists in depends on the attraction between molecules and their ability to overcome the attraction • the attractive forces between ions or molecules depends on their structure the attractions are electrostatic depend on shape, polarity, etc.
• the ability of the molecules to overcome the attraction depends on the amount of kinetic energy they possess 14
Properties of Liquids Viscosity
• some liquids flow more easily than others • the resistance of a liquid to flow we call
viscosity
• larger the attractive forces between the molecules = larger the viscosity • also, molecules whose shape is not round will have a larger viscosity 15
Properties of Liquids Surface Tension
• liquids tend to minimize their surface – a phenomenon we call
surface tension
• this tendency causes liquids to have a surface that resists penetration 16
Surface Tension
• molecules in the interior of a liquid experience attractions to surrounding molecules in all directions • but molecules on the surface experience an imbalance in attractions, effectively pulling them in • to minimize this imbalance and maximize attraction, liquids try to minimize the number of molecules on the exposed surface by minimizing their surface area • stronger attractive forces between the molecules = larger surface tension 17
Forces of Attraction within a Liquid
• •
Cohesive Forces
= forces that try to hold the liquid molecules to each other surface tension
Adhesive Forces
= forces that bind a substance to a surface capillary action meniscus 18
Escaping from the Surface
• the process of molecules of a liquid breaking free from the surface is called
evaporation
also known as vaporization • evaporation is a physical change in which a substance is converted from its liquid form to its gaseous form the gaseous form is called a
vapor
19
Evaporation
• over time, liquids evaporate – the molecules of the liquid mix with and dissolve in the air • the evaporation happens at the surface • molecules on the surface experience a smaller net attractive force than molecules in the interior • but all the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape 20
Factors Effecting the Rate of Evaporation
• increasing the surface area increases the rate of evaporation • increasing the temperature increases the rate of evaporation • weaker attractive forces between the molecules = faster rate of evaporation • liquids that evaporate quickly are called
volatile
liquids, while those that do not are called
nonvolatile
21
Escaping the Surface
• the
average
kinetic energy is directly proportional to the kelvin temperature • but not all molecules in the sample have the same kinetic energy • those molecules on the surface that have enough kinetic energy will escape raising the temperature increases the number of molecules with sufficient energy to escape 22
Escaping the Surface
• since the higher energy molecules from the liquid are leaving, the total kinetic energy of the liquid decreases, and the liquid cools • the remaining molecules redistribute their energies, generating more high energy molecules • the result is the liquid continues to evaporate 23
Reconnecting with the Surface
• when a liquid evaporates in a closed container, the vapor molecules are trapped • the vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid – this process is called
condensation
a physical change in which a gaseous form is converted to a liquid form 24
Dynamic Equilibrium
• evaporation and condensation are opposite processes • eventually, the rate of evaporation and condensation in the container will be the same • opposite processes that occur at the same rate in the same system are said to be in
dynamic equilibrium
25
Evaporation and Condensation
26
Vapor Pressure
• once equilibrium is reached, from that time forward, the amount of vapor in the container will remain the same as long as you don’t change the conditions • the partial pressure exerted by the vapor is called the
vapor pressure
• the vapor pressure of a liquid depends on the temperature and strength of intermolecular attractions 27
Boiling
• in an open container, as you heat a liquid the average kinetic energy of the molecules increases, giving more molecules enough energy to escape the surface so the rate of evaporation increases • eventually the temperature is high enough for molecules in the interior of the liquid to escape – a phenomenon we call
boiling
28
Boiling Point
• the temperature at which the vapor pressure of the liquid is the same as the atmospheric pressure is called the
boiling point
the normal boiling point is the temperature required for the vapor pressure of the liquid to be equal to 1 atm • the boiling point depends on what the atmospheric pressure is the temperature of boiling water on the top of a mountain will be cooler than boiling water at sea level 29
Temperature and Boiling
• as you heat a liquid, its temperature increases until it reaches the boiling point • once the liquid starts to boil, the temperature remains the same until it all turns to a gas • all the energy from the heat source is being used to overcome the attractive forces in the liquid 30
Energetics of Evaporation
• as it loses the high energy molecules through evaporation, the liquid cools • then the liquid absorbs heat from its surroundings to raise its temperature back to the same as the surroundings • processes in which heat flows into a system from the surroundings are said to be
endothermic
• as heat flows out of the surroundings, it causes the surroundings to cool as alcohol evaporates off your skin, it causes your skin to cool 31
Energetics of Condensation
• as it gains the high energy molecules through condensation, the liquid warms • then the liquid releases heat to its surroundings to reduce its temperature back to the same as the surroundings • processes in which heat flows out of a system into the surroundings are said to be
exothermic
• as heat flows into the surroundings, it causes the surroundings to warm 32
Heat of Vaporization
• the amount of heat needed to vaporize one mole of a liquid is called the
heat of vaporization
D H vap it requires 40.7 kJ of heat to vaporize one mole of water at 100°C endothermic D H vap depends on the initial temperature • since condensation is the opposite process to evaporation, the same amount of energy is transferred but in the opposite direction D H cond = D H vap 33
Liquid
Heats of Vaporization of Liquids at their Boiling Points and at 25°C
Chemical Formula Normal Boiling Point, °C
D
H vap at Boiling Point, (kJ/mol)
D
H vap at 25°C, (kJ/mol)
water isopropyl alcohol H 2 O C 3 H 7 OH 100 82.3
40.7
39.9
44.0
45.4
acetone diethyl ether C 3 H 6 O C 4 H 10 O 56.1
34.5
29.1
26.5
31.0
27.1
34
Example: Calculate the amount of water in grams that can be vaporized at its boiling point with 155 kJ of heat.
Information Given: 155 kJ Find: g H 2 O CF: 40.7 kJ = 1 mol; 18.02 g = 1 mol SM: kJ → mol → g • Apply the Solution Map: 155 kJ 1 mol H 2 O 40.7
kJ 18
.
02 g H 2 O 1 mol H 2 O = 68.626 g H 2 O • Sig. Figs. & Round: = 68.6 g H 2 O 35
Temperature and Melting
• as you heat a solid, its temperature increases until it reaches the
melting point
• once the solid starts to melt, the temperature remains the same until it all turns to a liquid • all the energy from the heat source is being used to overcome the attractive forces in the solid that hold them in place 36
Energetics of Melting and Freezing
• when a solid melts, it absorbs heat from its surroundings, it is
endothermic
• as heat flows out of the surroundings, it causes the surroundings to cool as ice in your drink melts, it cause the liquid to cool • when a liquid freezes, it releases heat into its surroundings, it is
exothermic
• as heat flows into the surroundings, it causes the surroundings to warm 37
Heat of Fusion
• the amount of heat needed to melt one mole of a solid is called the
heat of fusion
D H fus fusion is an old term for heating a substance until it melts, it is not the same as nuclear fusion • since freezing is the opposite process to melting, the same amount of energy transferred is the same, but in the opposite direction D H crystal = D H fus • in general, D H vap > D H fus because vaporization requires breaking all attractive forces 38
Heats of Fusion of Several Substances
Liquid
water isopropyl alcohol acetone diethyl ether
Chemical Formula
H 2 O C 3 H 7 OH C 3 H 6 O C 4 H 10 O
Melting Point, °C
0.00
-89.5
-94.8
-116.3
D
H fusion , (kJ/mol)
6.02
5.37
5.69
7.27
39
Sublimation
• sublimation is a physical change in which the solid form changes directly to the gaseous form without going through the liquid form • like melting, sublimation is endothermic 40
Intermolecular Attractive Forces
Why are molecules attracted to each other?
• intermolecular attractions are due to attractive forces between opposite charges • + ion to - ion • • + end of polar molecule to - end of polar molecule H-bonding especially strong
larger charge = stronger attraction
• even nonpolar molecules will have a temporary induced dipoles 42
Dispersion Forces
• also known as London Forces or Induced Dipoles • caused by electrons on one molecule distorting the electron cloud on another • all molecules have dispersion forces + + + + + + + + + + 43
Instantaneous Dipoles
44
Strength of the Dispersion Force
• • depends on how easily the electrons can move, or be
polarized
• the more electrons and the farther they are from the nuclei, the larger the dipole that can be induced
strength of the dispersion force gets larger with larger molecules
45
Attractive Forces and Properties
• stronger attractive forces between molecules = higher boiling point in pure substance • stronger attractive forces between molecules = higher melting point in pure substance though also depends on crystal packing 46
Dispersion Force and Molar Mass
Noble Gas
He
Molar Mass, (g/mol)
4.00
Boiling Point, (K)
4.2 Ne Ar Kr Xe 20.18 39.95 83.80 131.29 27 87 120 165
47
250 200 150 100 -150 -200 -250 -300 50 0 1 -50 -100
Relationship between Dispersion Force and Molecular Size
2 3 4 5
Period
BP, Noble Gas BP, Halogens BP, XH4 6 48
Permanent Dipoles
• because of the kinds of atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole • all polar molecules have a permanent dipole 49
Dipole-to-Dipole Attraction
• polar molecules have a permanent dipole a + end and a – end • the + end of one molecule will be attracted to the – end of another 50
Polarity and Dipole-to-Dipole Attraction
CH 3 CH 2 CH 3 CH 3 -O-CH 3 CH 3 - CH=O CH 3 -C N MolarMass (g/mol) 44 46 44 41 Boiling Point, °C -42 -24 20.2 81.6 Dipole Size, D 0 1.3 2.7 3.9 51
Attractive Forces
Dispersion Forces – all molecules + + + + _ _ _ _ Dipole-to-Dipole Forces – polar molecules + + + - + 52
Attractive Forces and Properties
• Like dissolves Like miscible = liquids that do not separate, no matter what the proportions • polar molecules dissolve in polar solvents water, alcohol, CH 2 Cl 2 molecules with O or N higher solubility in H 2 O due to H-bonding with H 2 O • nonpolar molecules dissolve in nonpolar solvents ligroin (hexane), toluene, CCl 4 • if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition 53
Immiscible Liquids
When liquid pentane, a nonpolar substance, is mixed with water, a polar substance, the two liquids separate because they are more attracted to their own kind of molecule than to the other.
54
Hydrogen Bonding
• Molecules that have
HF
,
OH
or
NH
groups have particularly strong intermolecular attractions unusually high melting and boiling points unusually high solubility in water • this kind of attraction is called a
Hydrogen Bond
55
Properties and H-Bonding
Name
Ethane
Form ula Molar Mass (g/mol) Structure
C 2 Ethanol CH H 4 6 O 30.0
32.0
H H H C H H C H O H C H H H
Boiling Point, °C Melting Solubil Point, ity in °C Water
-88 64.7
-172 immisc -97.8
misc ble 56
Intermolecular H-Bonding
57
Hydrogen Bonding
• When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it.
• Since hydrogen has no other electrons, when it loses the electrons, the nucleus becomes deshielded exposing the proton • The exposed proton acts as a very strong center of positive charge, attracting all the electron clouds from neighboring molecules 58
H-Bonds vs. Chemical Bonds
• hydrogen bonds are
not
chemical bonds • hydrogen bonds are attractive forces
between
molecules • chemical bonds are attractive forces that
make
molecules 59
150 100 50 0 1 -50 -100 -150 -200
Relationship between H-bonding and Intermolecular Attraction
H 2 O BP, HX BP, H2X BP, H3X BP, XH4 HF NH 3 2 3 H 2 S H 2 Se 4 H 2 Te 5 SnH 4 GeH 4 SiH 4 CH 4
Period
60
Attractive Forces & Properties
CH 3 CH 2 OCH 2 CH 3 CH 3 CH 2 CH 2 CH 2 CH 3 CH 3 CH 2 CH 2 CH 2 OH
Molar Mass, (g/mol)
74
Boiling Point, °C
34.6
Solubility in water, (g/100 g H 2 O)
7.5 72 74 36 117 INSOL. 9 61
Types of Intermolecular Forces
Type of Force
Dispersion Force
Relative Strength Present in
weak, but increases with molar mass all atoms and molecules
Example
H 2 Dipole – Dipole Force moderate only polar molecules HCl Hydrogen Bond strong molecules having H bonded to F, O or N HF 62
Type of Force Type of Interaction
Ionic Bond Covalent Bond Hydrogen Bond Ion to Dipole Dipole to Dipole Ion to Induced Dipole to Induced Induced to Induced cation + anion shared electrons H (bonded to O, N, or F) attracted to an electronegative atom Ion + polar molec.
polar + polar ion + nonpolar polar + nonpolar nonpolar + nonpolar
Energy (kJ/mol)
300 - 600 200-400 20-40 10-20 1-5 1-3 0.005-2 0.005-2 63
Crystalline Solids
Types of Crystalline Solids
65
Molecular Crystalline Solids
• Molecular solids are solids whose composite units are molecules • Solid held together by intermolecular attractive forces dispersion, dipole-dipole, or H-bonding • generally low melting points and D H fusion 66
Ionic Crystalline Solids
• Ionic solids are solids whose composite units are formula units • Solid held together by electrostatic attractive forces between cations and anions cations and anions arranged in a geometric pattern called a
crystal lattice
to maximize attractions • generally higher melting points and D H fusion than molecular solids because ionic bonds are stronger than intermolecular forces 67
Atomic Crystalline Solids
• Atomic solids are solids whose composite units are individual atoms • Solid held together by either covalent bonds, dispersion forces or metallic bonds • melting points and D H fusion vary depending on the attractive forces between the atoms 68
Types of Atomic Solids
69
Types of Atomic Solids Covalent
• Covalent Atomic Solids have their atoms attached by covalent bonds effectively, the entire solid is one, giant molecule • because covalent bonds are strong, these solids have very high melting points and D H fusion • because covalent bonds are directional, these substances tend to be very hard 70
Types of Atomic Solids Nonbonding
• Nonbonding Atomic Solids are held together by dispersion forces • because dispersion forces are relatively weak, these solids have very low melting points and D H fusion 71
Types of Atomic Solids Metallic
• Metallic solids are held together by metallic bonds • metal atoms release some of their electrons to be shared by all the other atoms in the crystal • the metallic bond is the attraction of the metal cations for the mobile electrons often described as islands of cations in a sea of electrons 72
Metallic Bonding
• the model of metallic bonding can be used to explain the properties of metals • the luster, malleability, ductility, electrical and thermal conductivity are all related to the mobility of the electrons in the solid • the strength of the metallic bond varies, depending on the charge and size of the cations – so the melting points and D H fusion of metals vary as well 73
Water A Unique and Important Substance
• water is found in all 3 states on the earth • as a liquid, it is the most common solvent found in nature • without water, life as we know it could not exist the search for extraterrestrial life starts with the search for water 74
Water
• liquid at room temperature most molecular substances that have a molar mass (18.02 g/mol) similar to water’s are gaseous • relatively high boiling point • expands as it freezes most substances contract as they freeze causes ice to be less dense than liquid water 75