Liquids and Solids Chapter 12. 2 Liquid Has a definite volume and indefinite shape Particles are in...

Liquids and Solids

Chapter 12

Chemistry chapter 12 2

Liquid

Has a definite volume and indefinite shape

Particles are in constant motion Closer together than gases Less kinetic energy than gases Greater attractive forces than gases


Fluid

Substance that can flow and therefore take the shape of its container


Fluid density

At normal pressure, most liquids are thousands of times denser than their gases. Particles are closer together

Different liquids can vary greatly in density


Density


Incompressibility

Liquids are much less compressible than gases Particles are closer together Liquids can transmit pressure

throughout themselves


Diffusion Liquids diffuse in other liquids in

which they can dissolve Much slower than gases

Particles closer together Attraction between particles slows

them down Faster at higher temperatures

More kinetic energy

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Surface Tension

Force that pulls parts of a liquid’s surface together, causing it to have the smallest possible size

From attractive forces between molecules

Liquid droplets take a spherical shape


Capillary action

The attraction of the surface of a liquid to the surface of a solid Causes meniscus Water can travel up paper Water traveling up a plant Drawing blood in capillary tube


Vaporization

Changing from a liquid to a gas Evaporation

Higher energy particles escape from the surface of a nonboiling liquid

perfume

Boiling Bubbles of vapor that appear

throughout the liquid and travel to the surface


Freezing

Changing a liquid to a solid by removing heat

Energy of particles decreases until they are pulled into a more orderly arrangement


Discussion

Describe the liquid state using kinetic molecular theory.

Explain why liquids in a test tube form a meniscus.

Compare and contrast vaporization and evaporation.


Solid

Has definite volume and definite shape

Particles are in constant motion Much closer together than liquid or

gas Much stronger intermolecular forces Held in relatively fixed position – only

vibrate Most ordered state of matter


High density and incompressibility

Substances are generally the most dense in the solid state Slightly denser than liquids, much

denser than gases Virtually incompressible

Sometimes we compress air pockets in the solids

Wood, cork, etc.


Diffusion

Very, very slow A few atoms may diffuse if

clamped together for a long time


Melting

Change of a solid to a liquid by addition of heat

Melting point – temperature at which something melts


Crystalline solids

Consist of crystals Particles are arranged in an orderly,

geometric, repeating pattern Fragments have geometric shapes Have definite melting points

When the crystal structure breaks apart


Crystal structure

Total three-dimensional arrangement of particles in a crystal


Lattice

Coordinate system that represents the arrangement of particles in a crystal.


Unit cell

Smallest portion of a crystal lattice that shows the 3D pattern of the entire lattice

Each crystal lattice contains many unit cells packed together

Has one of seven types of symmetry – see page 369


Ionic crystals

Positive and negative ions in a regular pattern

Hard and brittle High melting points Good insulators


Covalent network crystals Individual atoms connected by

covalent bonds Giant molecules

Diamond Quartz

Very hard and brittle Rather high melting points Nonconductors or semiconductors


Metallic crystals

Metal atoms surrounded by a sea of electrons

High electrical conductivity Varying melting points


Covalent molecular crystals

Covalently bonded molecules held together by intermolecular forces

Low melting points Easily vaporized Relatively soft Good insulators


Amorphous solids

Noncrystalline solids The particles are arranged

randomly Glass Plastics

Can be molded Fragments have irregular shapes


Amorphous solids

Made by cooling molten substances in a way that prevents crystallization

Also called supercooled fluids Retain certain fluid characteristics

even at temperatures at which they appear to be solid

Can flow over a wide range of temperatures


Discuss

Account for each of the following properties of solids: Definite volume Relatively high density Extremely low rate of diffusion

What is the difference between an amorphous solid and a crystalline solid?


Possible changes of state

Melting: solid to liquid Sublimation: solid to gas Freezing: liquid to solid Vaporization: liquid to gas Condensation: gas to liquid Deposition: gas to solid


Closed system

Matter cannot enter or leave, but energy can


Equilibrium

A dynamic condition in which two opposing changes occur at equal rates in a closed system.

The same number of particles are entering and leaving.

The total number stays the same.


Equilibrium


Phase

Any part of a system that has uniform composition and properties. Liquid or gas


An Equilibrium equation

When a substance changes state, it either absorbs or gives off energy, usually as heat.

vaporenergy heat liquid


Le Châtelier’s Principle

A system remains in equilibrium until a stress occurs on the system. Stress: change in concentration,

pressure, or temperature When a system is disturbed by a

stress, it attains a new equilibrium position that minimizes the stress.


Shifting equilibrium

Shifts to the right or left, depending on which part of the equation gains concentration.

See table 12-3 on page 375


Equilibrium vapor pressure

The pressure exerted by a vapor in equilibrium with its liquid at a given temperature

Increases as temperature increases But not directly


Kinetic-molecular theory

Increasing the temperature increases the energy and speed of the liquid particles

This means more particles evaporate, leading to higher vapor pressure


Caution

Equilibrium vapor pressure depends only on temperature.

If the system is not in equilibrium, gas laws must be used.


Volatile liquids

Evaporate easily Weak forces of attraction between

particles Ether, acetone


Boiling

Conversion of a liquid to a vapor within the liquid as well as at the surface.

Occurs when the equilibrium vapor pressure equals atmospheric pressure.

All the heat absorbed goes to evaporate the liquid, so the temperature remains constant.


Cooking If atmospheric pressure is lower (high

altitudes), liquids boil at lower temperatures and food takes longer to cook.

If the pressure is increased (pressure cooker), liquids boil at higher temperatures and food cooks faster.

If the pressure is decreased, it boils at low enough temperatures to avoid scorching milk and sugar. (evaporated and sweetened condensed milk)


Molar heat of vaporization

The amount of energy needed to vaporize one mole of liquid at its boiling point. (or the amount of energy released

when one mole of vapor condenses) A measure of the attraction

between particles.


Normal freezing point

Temperature at which the solid and liquid are in equilibrium at 1 atm pressure.

When a liquid freezes, energy is lost and order is gained.


Clarification

Boiling point is the same as condensation point.

Freezing point is the same as melting point.


Molar heat of fusion

The amount of heat energy required to melt one mole of solid at its melting point (or the amount of energy released

when one mole of a liquid freezes) Depends on the attraction between

particles.


Phase diagram

A graph of pressure versus temperature that shows the conditions under which the phases of a substance exist

Reveals how the states of a system change with changes in temperature or pressure


Water’s phase diagram


Curves on diagram

AB shows where solid and vapor can exist at equilibrium

AC shows liquid and vapor at equilibrium

AD shows solid and liquid at equilibrium Usually a positive slope, but water is

different


Triple point

Point A Shows the temperature and

pressure at which solid, liquid, and vapor can exist in equilibrium


Critical temperature

Temperature above which the substance cannot exist in the liquid state.


Critical pressure

The lowest pressure at which the substance can exist as a liquid at the critical temperature


Critical point

Point C Where critical pressure meets

critical temperature


Water’s phase diagram


Carbon Dioxide’s phase diagram


Discuss

Section review on page 382


Water Bent molecule

Bond angle = 105° Hydrogen bonding in liquids and

solids Usually 4 – 8 molecules per group in

liquid water Without them, water would be a gas at

room temperature Ice has hexagonal arrangement

Empty spaces lead to low density


Water

Has highest density at 3.98 °C When it melts, molecules can crowd

together When it gets hotter, increased kinetic

energy makes them spread apart Needs a lot of energy to

completely break hydrogen bonds and vaporize


Water Pure liquid water is transparent,

odorless, tasteless, and almost colorless Odors or tastes are caused by impurities

Molar heat of fusion: 6.009 kJ/mol Relatively large Density of ice: 0.917 g/cm3

Molar heat of vaporization: 40.79 kJ/mol Quite high


Calculating heat energy

energyheat ation or vaporizfusion ofheat amount

kJmol

kJmol


Example

Find the mass of liquid water required to absorb 5.23 x 104 kJ of heat energy on boiling.

2.31 x 104 g


You try

How much heat energy is absorbed when 16.3 g of ice melts?

5.44 kJ


You try

Calculate the quantity of heat energy released when 783 g of steam condenses.

1.77 x 103 kJ


Specific Heat The amount of energy required to

raise the temperature of one gram of substance by one one kelvin. J/(g∙K)

Used to calculate the energy absorbed or released by a substance during a temperature change No change of state

Example: warm water cool waterpQ c m T

Liquids and Solids Chapter 12. 2 Liquid Has a definite volume and indefinite shape Particles are in...

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