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Liquids and Solids. H 2 O (g). H 2 O (s). H 2 O (  ). Three States of Matter. The state (or phase) of matter is determined by the arrangement and motion of particles. The motion of particles is governed by the kinetic energy (KE) of the particles (Remember that KE = 1/2mv 2 )

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slide1

Liquids and Solids

H2O (g)

H2O (s)

H2O ()

slide2

Three States of Matter

The state (or phase) of matter is determined by the arrangement and motion of particles.

The motion of particles is governed by the kinetic energy (KE) of the particles (Remember that KE = 1/2mv2)

Solids Liquids Gases

Increase KE

Increase KE

slide3

Changes of State

sublimation

sublimation

boiling

melting

vaporization

condensation

freezing

deposition

slide5

Attractive Forces In Molecules

Intermolecular Forces: Attractive forces between molecules

Types of inter-molecular forces

  • dipole-dipole (1% as strong as covalent bonds)

A special type of dipole-dipole force is the hydrogen bond. These bonds form between molecules that contain a hydrogen atoms bonded to a very electronegative element like N, O or F. Hydrogen bonds are very strong compared to an ordinary dipole-dipole bond.

E.g HF, NH3, H2O all form hydrogen bonds

Hydrogen bonding 10% as strong as covalent bonds

slide6

Water molecules are polar molecules. The  - oxygen forms intermolecular bonds with the  + hydrogen of another water molecules. This is an example of a special type of intermolecular bond called a hydrogen bond.

Inter-molecular forces

slide8

Non-polar molecule

This instantaneous dipole will effect any nearby molecules

Movement of electrons causes an instantaneous dipole

This induces a dipole in a nearby molecule

slide9

Properties of Liquids

As we consider the properties of liquids (and solids) that KE and intermolecular attractions are governing the behavior of the substance

slide10

Properties of Liquids

Vaporization: Change from liquid to gas via boiling process

Evaporation: Change from liquid to gas at the surface of a liquid, not caused by boiling

--This happens because the molecules at the top a of the liquid don’t have as strong of an attraction to the other molecules. (If they have high enough KE, they can escape)

slide11
Open Containers: Evaporation causes liquid molecules to leave as gases and escape (amount of liquid decreases)

Closed Containers:

Evaporation causes liquid molecules to vaporize, but they get caught in the container, creating : VAPOR PRESSURE

Evaporation: A Closer Look

slide12

Vapor Pressure

This is known as Dynamic Equilibrium because the rates of evaporation and condensation are EQUAL .

In a sealed container, molecules will start to evaporate and the liquid’s volume will decrease.

Evaporation and Condensation DO NOT stop happening once the flask has reached equilibrium

But, after the air above the liquid becomes “saturated”, some of these molecules will then condense. After a short time, the volume of the liquid will not change.

The rate of evaporation = the rate of condensation

slide13

Properties of Liquids, cont.

Boiling: When all the molecules of a liquid have enough kinetic energy to vaporize, the liquid is said to be boiling.

Boiling Point (bp): The temperature at which the vapor pressure of a liquid is just equal to the external pressure on the liquid.

Normal Boiling Point: The temperature at which a substance boils at atmospheric pressure (101.3kPa)

slide14

Boiling point and Vapor pressure

When water is heated, the kinetic energy of the molecules increases and eventually bubbles of vapor form within it. The vapor pressure in the bubble is the same as the vapor pressure of the water at that temperature.

When the temperature of the water reaches a point that the vapor pressure of the bubbleequals atmospheric pressure, the bubbles get larger, rise to the surface, and escape as steam. The water begins to boil.

slide15

Boiling point and Vapor pressure

REMEMBER:

vapor pressure of the bubbleequals atmospheric pressure

The water begins to boil

atmospheric pressure

450mm Hg

58o C

At lower atmospheric pressures, the kinetic energy does not have to be as high to make the vapor pressure in the bubble equal to atmospheric pressure.

450 mm Hg

By reducing the atmospheric pressure, The water begins to boil at a lower temperature.

slide16

Other Properties of Liquids

1. Why, when you pour a liquid onto a surface does it form droplets?

2. Why do some liquids exhibit capillaryaction?

Hg H2O

3. Why are some liquids more viscous than others?

slide17

Viscosity:is the resistance to motion of a liquid.

Maple syrup is more viscous than water. But water is much more viscous than gasoline or alcohol.

The stronger the attraction between molecules of a liquid, the greater itsresistance to flow and so the more viscous it is.

Consider the following substances

a) molasses b) water c) ethyl alcohol

  • Which is the least viscous?
  • Which substance has the strongest intermolecular attractions?

Ethyl alcohol

Molasses

slide18

Surface tension

This water strider uses surface tension to his advantage

The inward force or pull which tends to minimize the surface area of any liquid is surface tension.

slide19

Surface tension occurs because the molecules on the surface of the liquid cannot bond to the outside molecules. As a result, they look for something else to bond to (in order to increase stability). They get “pulled” in towards each other until their surface area is minimized, thus minimizing the contact with the outside.

Surface tension in water is caused by hydrogen bonding between polar molecules. The more polar a liquid the stronger its surface tension.

Hg pure H2O H2O with detergent

**Surfactants are compounds that reduce the surface tension of a liquid (soaps and detergents are examples)**

The smallest surface area a liquid can form is a sphere.

slide20

Capillary action is the spontaneous rising of a liquid in a narrow tube.

Two forces are responsible for this action:

Cohesive forces:the intermolecular forces between molecules of the liquid

Adhesive forces: the attractive forces between the liquid molecules and their container

If the container is made of a substance that has polar bonds then a polar liquid will be attracted to the container.

This is why water forms a concave meniscus while mercury forms convex meniscus

Hg H2O

slide21

Properties of Solids

Solids generally have an orderly arrangement of atoms

Melting: When the kinetic energy of all the atoms in a solid is increased to a point where the atoms are able to freely flow around one another, the solid is said to have melted

Melting Point (mp): The temperature at which a solid turns into a liquid.

slide22
Crystalline

-Most solids are crystalline

-Contain particles arranged in an orderly, repeating, 3-D pattern called a crystal lattice

Non-Crystalline

- Amorphous solids have no set crystal structure

Examples:

1. Glass

2. Asphalt

3. Rubber

4. Plastic

5. Candles (Wax)

Types of Solids

slide23

Allotropes

Some crystalline solids (pure substances) occur in a variety of different forms, known as Allotropes. Each allotrope has a different crystalline pattern that connects the atoms of the solid

Carbon, Sulfur, Phosphorus, Oxygen, Boron and Antimony all have allotropes

The most common examples of allotropes are found in elemental Carbon:

Diamond Graphite Buckminsterfullerene

(Bucky Balls)

slide24

ICE

Solids are almost always more dense than their liquid forms, however, there is one exception: ICE

Ice molecules are locked in fixed positions, held by intermolecular-bonds.

Ice is less dense than liquid water because the molecules are further apart than in liquid water.

slide25

Tracking Changes in State

sublimation

sublimation

boiling

melting

vaporization

condensation

freezing

deposition

slide26

Tracking Changes in State

Melting/Vaporization Curves: tracks temperature changes as a function of time and shows all state changes

boiling

slide27

Tracking Changes in State

Phase Diagrams: Shows the various conditions at which each state of a substance can occur

boiling