States of Matter

Vocabulary:Surface tensionTriple pointUnit cellCrystalline solidAllotropeAmorphous solidViscosityGas pressureVapor pressure

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States of Matter

Gases expand, diffuse, exert pressure, and can be compressed because they are in a low density state consisting of tiny, constantly-moving particles.

Section 12-1

The Kinetic-Molecular Theory

• Kinetic-molecular theory explains the different properties of solids, liquids, and gases.

• Atomic composition affects chemical properties.

• Atomic composition also affects physical properties.

• The kinetic-molecular theory describes the behavior of matter in terms of particles in motion.

Describes the behavior of matter in terms of particles in motionMakes several assumptions about the size, motion, and energy of gas particles

Assumptions of the Kinetic Molecular Theory

• Gasses consist of small particles that take up little volume relative to the volume of empty space around them– Gas molecules are very far apart and therefore don’t

experience attractive or repulsive forces

• Gas particles move in constant, random straight lines until they collide with other particles or with the walls of the container.

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• Gas particles are in constant random motion.

• An elastic collision is one in which no kinetic energy is lost.

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• Kinetic energy of a particle depends on mass and velocity.

• Temperature is a measure of the average kinetic energy of the particles in a sample of matter.

• A gas particle has a mass of 5.31 x 10-21 kg and a velocity of 1.00 x 102 m/s. What is the kinetic energy of the particle?

KE = .5 x (5.31 x 10-21 kg) x (1.00 x 102 m/s)2

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Explaining the Behavior of Gases

• Great amounts of space exist between gas particles.

• Compression reduces the empty spaces between particles.

Explaining the Behavior of Gasses

• Kinetic molecular theory helps explain the behavior of gasses– Blowing up a balloon

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• Gases easily flow past each other because there are no significant forces of attraction.

• Diffusion is the movement of one material through another.

Example: perfume, popcorn• Effusion is a gas escaping through a tiny

opening.

Example: flat tire

• Gasses have low densities– Chlorine gas = 0.00295 g/mL– Germanium = 5.323 g/mL– Why?

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• Graham’s law of effusion states:

• Graham’s law also applies to diffusion.

• Rates of diffusion

• Ammonia has a molar mass of 17.0 g/mol; hydrogen chloride has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

√ 36.5 g/mol

17.0 g/mol

• What is the molar mass of a gas that takes four times as long to effuse as hydrogen?

√ X g/mol = 4 2.0 g/mol

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Gas Pressure

• Pressure is defined as force per unit area.

• Gas particles exert pressure when they collide with the walls of their container.

• The particles in the earth’s atmosphere exert pressure in all directions called air pressure.

• There is less air pressure at high altitudes because there are fewer particles present, since the force of gravity is less.

• Pressure = force per unit area– Smaller area = more pressure

• Air pressure – pressure of the atmosphere– 1 atm– 760 mm Hg– 760 Torr– 101.3 kPa – 101,300 Pa– 14.7 psi

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• Torricelli invented the barometer.

• Barometers are instruments used to measure atmospheric air pressure.

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• Manometers measure gas pressure in a closed container.

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• The SI unit of force is the newton (N).

• One pascal(Pa) is equal to a force of one Newton per square meter or N/m2.

• One atmosphere is equal to 760 mm Hg or 101.3 kilopascals.

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• Dalton’s law of partial pressures states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases of the mixture.

• The partial pressure of a gas depends on the number of moles, size of the container, and temperature and is independent of the type of gas.

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Ptotal = P1 + P2 + P3 +...Pn

• Partial pressure can be used to calculate the amount of gas produced in a chemical reaction.

• A mixture of oxygen, carbon dioxide, and nitrogen has a total pressure of 0.97 atm. What is the partial pressure of O2 if the partial pressure of CO2 is 0.70 atm and the partial pressure of N2 is 0.12 atm?

Forces of attraction

•Intermolecular forces (dispersion forces, dipole-dipole forces, and hydrogen bonds) determine a substance’s state at a given temperature

Intermolecular forces

• Inter- means between or among

• Intermolecular forces can hold together identical particles or two different types of particles

• Weaker than intramolecular forces (bonds)

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Intermolecular Forces

• Attractive forces between molecules cause some materials to be solids, some to be liquids, and some to be gases at the same temperature.

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• Dispersion forces are weak forces that result from temporary shifts in density of electrons in electron clouds.

• Exist between all particles– Weak for small particles– Get stronger as the number of electrons

involved increases

– F2

– Cl2– Br2

– I2

Dipole-dipole forces

• Attraction between oppositely charged regions of polar molecules– Polar molecule =

• Neighboring polar molecules orient themselves so that oppositely charged regions align

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• Dipole-dipole forces are attractions between oppositely charged regions of polar molecules.

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• Hydrogen bonds are special dipole-dipole attractions that occur between molecules that contain a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons, typically fluorine, oxygen, or nitrogen.

• Explain why water is liquid at room temperature while compounds of similar masses are gasses

12.3 Liquids and Solids

• The particles in solids and liquids have a limited range of motion and are not easily compressed.

Liquids

• Kinetic molecular theory also applies to liquids and solids– Must take intermolecular forces into account

to apply it

• Density and compression– Much denser than gasses

• Due to intermolecular forces holding particles together

– Incompressible • Why can you compress a gas but not a liquid?

• Fluidity – both gases and liquids are classified as fluids because they can flow and diffuse– Liquids diffuse more slowly because

intermolecular attractions interfere with the flow

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Liquids

• Forces of attraction keep molecules closely packed in a fixed volume, but not in a fixed position.

• Liquids are much denser than gases because of the stronger intermolecular forces holding the particles together.

• Large amounts of pressure must be applied to compress liquids to very small amounts.

• Fluidity is the ability to flow and diffuse; liquids and gases are fluids.

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• The stronger the intermolecular attractive forces, the higher the viscosity.

• Particle size -larger molecules create higher viscosity.

• Long chains of molecules result in a higher viscosity.

• Temperature – lower temperature = higher viscosity

Increasing the temperature decreases viscosity because the added energy allows the molecules to overcome intermolecular forces and flow more freely.

• Viscosity is a measure of the resistance of a liquid to flow and is determined by the type of intermolecular forces, size and shape of particles, and temperature.

• Surface tension – the energy required to increase the surface area of a liquid by a given amount– Caused by intermolecular forces pulling down

on the particles on the surface of a liquid which stretches it tight like a drum

• Stronger the attraction between particles in a liquid = greater surface tension

• Surfactant – lowers the surface tension of water by disrupting hydrogen bonds between water molecules

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• Cohesion is the force of attraction between identical molecules.

• Adhesion is the force of attraction between molecules that are different.

• Capillary action is the upward movement of liquid into a narrow cylinder, or capillary tube.

Solids

• Solid particles have as much kinetic energy as liquids or gasses but much stronger attractive forces between particles– Limit the motion of particles to vibrations

• Most solids are more dense than liquids.• Ice is not more dense than water.

• Density of solids – almost always greater than density of liquids– Exception = water

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• Crystalline solids are solids with atoms, ions, or molecules arranged in an orderly, geometric shape.

– Unit cell = smallest arrangement of atoms in a crystalline solid that has the same shape as the whole crystal

• Categories of Crystalline Solids– Classified based on the types of particles they

contain and how they are bonded together

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• Categories of Crystalline Solids– Classified based on the types of particles they

contain and how they are bonded together

• Molecular solids– Molecules are held together by dispersion

forces, dipole-dipole forces or hydrogen bonds

– Most are not solid at room temperature – Poor conductors

• Covalent network solids– C or Si, can form multiple covalent bonds

which allow it to take many forms – Allotrope – element that can exist in different

forms at the same state

• Ionic solids– Made of cation + anion– Each ion is surrounded by ions of the

opposite charge– High melting point– Brittle

• Metallic solids– Positive metal ions surrounded by a sea of

mobile electrons

Amorphous solids

• Particles are not arranged in a regular, repeating pattern

• Does not contain crystals• Forms when molten material cools too

quickly for crystals to form– Glass– Rubber– Some plastics

Phase changes

• Matter changes phases when energy is added or removed

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Phase Changes That Require Energy

• Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature.Melting

• Heat flows from an object at a higher temperature to an object at a lower temperature

• Ice absorbs heat which does not raise temperature but is used to break hydrogen bonds

• When hydrogen bonds are broken molecules can move further apart into the liquid phase

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• When ice is heated, the ice eventually absorbs enough energy to break the hydrogen bonds that hold the water molecules together.

• When the bonds break, the particles move apart and ice melts into water.

• Melting point – temperature in which forces holding a solid together are broken and it becomes a liquid

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• Particles with enough energy escape from the liquid and enter the gas phase.

• Vaporization – process by which liquid changes to vapor– Vapor – gaseous state of a substance that is

normally liquid at room temperature– Evaporation – when vaporization occurs only

at the surface of a liquid– Vapor pressure – the pressure exerted by a

vapor over a liquid

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• In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure.

– Boiling – temperature at which the vapor pressure of a liquid equals the atmospheric pressure

– Energy being input causes molecules to move around more and vaporize

• Sublimation – changing from solid to gas without becoming a liquid– Dry ice– Moth balls– Solid air fresheners

Phase changes that release energy

• Freezing – Heat flows out of warmer object into cooler

object– Molecules slow down & become less likely to

flow past one another– Intermolecular forces cause the molecules to

become fixed into set positions– Freezing point – temperature in which a liquid

becomes a solid

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• As energy flows from water vapor, the velocity decreases.

• Condensation – process by which a gas or vapor becomes a liquid- shower mirror or glass door

• Deposition – substance changes from gas or vapor to solid without first becoming a liquid– frost

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• Temperature and pressure both effect the phase of a substanceHave opposite effectsPhase diagram – graph of pressure vs temperature that shows which phase a substance will be in under different conditions.

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• The triple point is the point on a phase diagram that represents the temperature and pressure at which all three phases of a substance can coexist.

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• The phase diagram for different substances are different from water.