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Copyright © Houghton Mifflin Company. All rights reserved. 6a–1 Gases, Liquids, and Solids The Phases, or States, of Matter

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Copyright © Houghton Mifflin Company. All rights reserved. 6a–1

Gases, Liquids, and Solids

The Phases, or States,of Matter

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Physical States of Matter A solid has a definite shape and a definite

volume. A liquid has an indefinite shape - it takes the

shape of its container – a definite volume. A gas has an indefinite shape and an indef-

inite volume – it expands to fill its container.

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Table 6.1 Distinguishing Properties of Solids, Liquids, and Gases.

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6.1 The Kinetic Molecular Theory of Matter

1. Matter is composed of tiny particles.

2. The particles are in constant random motion and possess kinetic energy.

3. The particles interact with each other through attractions and repulsions and possess potential energy.

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6.1 The Kinetic Molecular Theory of Matter

4. The velocity of particles increases with temperature, as does their kinetic energy.

5. The particles transfer energy to each other through elastic collisions.

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Figure 6.2 Upon release, the steel ball on the left transmits its kinetic energy through elas-tic collisions to the ball on the right.

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6.2 Kinetic Molecular Theory and Physical States

In solids, cohesive forces (potential energy) dominate over disruptive forces (kinetic energy).

In liquids, cohesive forces (potential energy) and disruptive forces (kinetic energy) are similar in magnitude.

In gases, disruptive forces (kinetic energy) dominate over cohesive forces (potential energy).

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(a) Particles in a solid (atoms, molecules, or ions) are close together and vibrate about fixed sites. (b) Particles in a liquid, though still close together, freely slide over one another. (c) Particles in a gas are in constant random motion, each particle being independent of the others present.

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When a gas is compressed, the amount of empty space in the container is decreased. The size of the molecules does not change; they simply move closer

together.

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6.3 Gas Law Variables

VolumeExpressed in milliliters or liters

TemperatureExpressed in K (°C + 273 = K)

AmountExpressed in moles (n)

Pressure (Force per unit area, F/A)Expressed in atmospheres, mm Hg, or torr

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Figure 6.5 The essential components of a mercury barometer are a graduated glass tube, a glass dish, and liquid mercury.

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6.4 Boyle’s Law: A Pressure-Volume Relationship

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6.4 Boyle’s Law: A Pressure-Volume Relationship

P1 x V1 = P2 x V2

A sample of O2 gas occupies 1.50 L at a pressure of 735 mm Hg and a temperature of 25C. What volume will it occupy if the pressure is increased to 770 mm Hg, and the temperature does not change?

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Figure 6.8 Filling a syringe with a liquid is an application of Boyle's law.

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6.5 Charles’ Law:A Temperature-Volume Relationship

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6.5 Charle’s Law:A Temperature-Volume Relationship

_V1_ = _V2_

T1 T2

A sample of gaseous anaesthetic has a volume of 425 mL at a temperature of 37C. What is its volume if it is cooled to 20 C at constant pressure?

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6.6 The Combined Gas Law

_P1V1_ = _P2V2_

T1 T2

A 1.50 L sample of N2O gas at a pressure of 755 mm Hg has a temperature of OC. What volume will the gas occupy at 50C and 725 mm Hg?

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6.7 The Ideal Gas Law

The ideal gas law includes the quantity of gas, in moles.

PV = nRT

R = 0.0821 L • atm

mol • K

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6.7 The Ideal Gas Law

Carbon monoxide, CO, is a colorless, odorless, tasteless gas that forms from incomplete combustion of carbon compounds.

What volume is occupied by 1.52 moles of this gas at 0.992 atm pressure, and 65 C? (IV-1)

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6.7 The Ideal Gas Law

How many moles of gas are present in a flask that holds 100.0 mL, at the boiling point of water (100C), and a pressure of 760 mm Hg?

If the gas has a mass of 0.750 g, what is its molar mass?

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6.8 Dalton’s Law of Partial Pressures

In a mixture of gases, each gas behaves as if the others were not there, and exerts its own pressure.

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6.8 Dalton’s Law of Partial Pressures

PTotal = P1 + P2 + P3 + ….

Air is a mixture of N2, O2, and small amounts of other gases, mostly water vapor. If the atmospheric pressure is 758 torr, the partial pressure of oxygen is 146 torr, and the par-tial pressure of nitrogen is 594 torr, what is the partial pressure of water vapor?

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6.9 Changes of State

A process in which a substance is trans-formed from one physical state to another.

Heating or Cooling

Changing Pressure

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Figure 6.13 There are six changes of state possible for substances.

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Figure 6.14 Sublimation and deposition of iodine. (a) The beaker contains iodine crystals.(b) Iodine has an appreciable vapor pressure be-

low its melting point. When heated, the solid sublimes. The vapor deposits crystals on the cool surface.

Source: James Scherer

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6.10 Evaporation of Liquids

Evaporation is the process by which molecules escape from the liquid phase to the gas phase.

Rate of evaporation is proportional to temperature.

Evaporation isn't necessarily boiling.

Vapor describes gaseous molecules of a substance that is mostly present as a liquid or a solid.

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6.11 The Vapor Pressures of Liquids

In a closed container, a liquid evaporates until the vapor and liquid reach equilibrium.Rate of evaporation equals rate of condensation

Vapor pressure of the liquid can be measured

Vapor pressure varies with temperature

Volatile substances have high vapor pressure

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Evaporation of a liquid in a closed container (a) Liquid level drops for a time(b) Liquid level becomes constant (ceases to drop). (c) Equilibrium has been reached; Rate of evaporation equals rate of condensation. Liquid exerts characteristic vapor pressure

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Table 6.2 Vapor Pressure of Water at Various Temperatures.

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6.12 Boiling and Boiling Point

Boiling is a form of evaporation that takes place throughout a liquid, and involves bubble formation

Boiling occurs when a liquid's vapor pressure equals that of the external pressure

Normal boiling point is the temperature at which a liquid boils at 760 mm Hg.

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Figure 6.16 Bubbles of vapor form within a liquid when the temperature of the liquid reaches the liquid's boiling point.

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Table 6.3 Boiling Point of Water at Various Locations That Differ in Elevation.

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6.13 Intermolecular Forces

Why is CH4 a gas, H2O a liquid, and C6H12O6 (glucose) a solid?

Different intermolecular forcesLondon dispersion forces

Dipole-dipole forces

Hydrogen bonds

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6.13 Intermolecular Forces

Dipole-dipole forces

Occur between polar molecules

Molecules align so dipoles of opposite charge interact

5 - 25 kJ/mol

(covalent bond 100 - 1000 kJ/mol)

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Figure 6.18 Dipole-dipole interactions between randomly arranged ClF molecules.

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6.13 Intermolecular Forces

Hydrogen bonds

A special case of dipole-dipole forces

Occurs in compounds with NH, OH, or FH bonds

N, O, H are very electronegative

H is very small

Molecules are very close

10 - 50 kJ/mol

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Figure 6.19 Depiction of hydrogen bonding among water molecules. The dotted lines are the hydrogen bonds.

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Figure 6.20 Diagrams of hydrogen bonding between selected simple molecules. Solid lines represent covalent bonds; dotted lines represent hydrogen bonds.

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Figure 6.21 If there were no hydrogen bonding be-tween water molecules, the boiling point of water would be approximately -80C.

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6.13 Intermolecular Forces

London dispersion forces

Temporary or induced dipoles

Are found in all molecules

Increase with "polarizability"Large atoms and/or molecules

pi bonds (double or triple bonds)

1 - 50 kJ/mol

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Figure 6.22

Nonpolar molecules such as H2 can develop instant-aneous dipoles and in-duced dipoles.

The attractions between such dipoles, through they are transitory, create London forces.

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6.13 Intermolecular Forces

Intermolecular forces are heirarchical and additive

All molecules have London dispersion forces; they are the only intermolec-ular forces in nonpolar molecules

All polar molecules have dipole-dipole forces

Hydrogen bonds are special

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6.13 Intermolecular Forces

Gases:

Intermolecular forces between molecules are not large enough to overcome kinetic energy (thermal energy) of individual molecules

Molecules that are gases at room tem-perature usually have only London dis-persion forces and/or are very small

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6.13 Intermolecular Forces

Liquids:

Intermolecular forces between molecules are about large enough to overcome most of the thermal energy of the molecules

Molecules that are liquids at room tem-perature usually have some combination of intermolecular forces, or are of moder-ate size

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6.13 Intermolecular Forces

Solids:

Intermolecular forces between molecules are large enough to overcome almost all of the thermal energy of the molecules

Molecules that are solids at room tem-perature have some combination of inter-molecular forces, or are large

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6.13 Intermolecular Forces

Phase changes and intermolecular forces:

A substance melts at a temperature where the thermal energy of the molecules is large enough to overcome some intermol-ecular forces

A substance boils at a temperature where the thermal energy of molecules is large enough to overcome most intermolecular forces.