Prof.P. Ravindran, - Universitetet i oslofolk.uio.no/ravi/cutn/cmp/5.States_of_Matter_old.pdf ·...

P.Ravindran, PHY075- Condensed Matter Physics, Spring 2015: States of Matter

http://folk.uio.no/ravi/CMP2015

Prof.P. Ravindran, Department of Physics, Central University of Tamil

Nadu, India

States of Matter

1


• Matter is anything that takes up space and has mass.

• Matter doesn’t have to be visible—even air is matter.

What is matter?

Matter

Everything in this photo is matter.


• All matter is made up of tiny particles, such as atoms, molecules, or ions.

• Each particle attracts other particles.

States of Matter

• These particles also are constantly moving.

• The motion of the particles and the strength of attraction between the particles determine a material’s state of matter.


• There are three familiar states of matter— solid, liquid, and gas.

• A fourth state of matter known as plasma occurs at extremely high temperatures. Plasma is found in stars, lightning, and neon lights.

States of Matter


States of Matter: Solids, Liquids, & Gases

• Intermolecular forces can be used to explain properties of solids, liquids and gases


Solids, Liquids, and Gases

deposition

sublimation

freezing

melting

boiling

condensing


• The particles that make up all types of matter are in constant motion.

• Although you can’t see them, a solid’s particles are vibrating in place.

Particles in Motion

• These particles do not have enough energy to move out of their fixed positions.


liquid

Three familiar states of matter:

Solid, Liquid, and Gas.

gassolid


• Gas is matter that does not have a definite shape or volume.

Gases

• Gas particles move at high speeds in all directions.

• The particles in gas are much farther apart

than those in a liquid or solid.


• Matter that exists in the gas state but is generally a liquid or solid at room temperature is called vapor.

Vapor

• Water, for example, is a liquid at room temperature. Thus, water vapor is the term for the gas state of water.


GASES

• Vapor Pressure

• The pressure of the gas “supported” over a liquid

• Indirectly related to the intermolecular forces within a liquid

• Higher intermolecular forces means more energy needed to escape.

• Lower vapor pressures give higher boiling points.

•Volatility

•Liquids with high vapor pressures (evaporate easier) are volatile


Gases

Low density

Indefinite shape and volume

Particles move around

Diffuse & Effuse

Gas particles spread out evenly as far

apart as possible.


You can change the volume of gas too.


Gases

Particles are spread apart

Move quickly

No definite shape

– Takes shape of

container

No definite volume

– Fills container


• A liquid is matter that has a definite volume but no definite shape.

Liquids

• Liquid takes the shape of the container.

• The volume of a liquid, however, is the same no matter what the shape of the container.

Particles are free to move around

one another


States of Matter: Basic Concepts

• Particles are in constant motion

• Fluidity is the ability to flow.

• Gases and liquids are classified as fluids because they can flow.

• A liquid diffuses more slowly than a gas at the same temperature (Why?)

• Because intermolecular attractions interfere with the flow.

Liquids


• The particles in a liquid move more freely than the particles in a solid.

Free to Move

• The particles in a liquid have enough energy to move out of their fixed positions but not enough energy to move far apart.


• Some liquids flow more easily than others.

Viscosity

• A liquid’s resistance to flow is known as the liquid’s viscosity.

• The slower a liquid flows, the higher its viscosity is.

• For many liquids, viscosity increases as the liquid becomes colder.


States of Matter: Liquids

• Viscosity: the resistance of a liquid to flow.

• ↑ Temperature = ↓Viscosity

• Why?

• With the increase in temperature, there is an increase in the average kinetic energy (velocity) of the molecules.


Viscosity

The resistance of a

fluid to flow

– Ketchup = high

– Water = low

Fluid

– A substance that flows


Fun fact: The Pitch Drop Experiment

Pitch, before and after being

hit with a hammer

If one heat a sample of

pitch and poured it into

glass funnel with a sealed

stem.

It is so viscous that it takes

10 years for a drop to drip.


• The uneven forces acting on the particles on the surface of a liquid are called surface tension.

Surface Tension

• Surface tension causes the liquid to act as if a thin film were stretched across its surface.


States of Matter: Liquids

• Particles in the middle of the liquid can be attracted to particles above them, below them, and to either side.

• The overall attractive force is pulling down on particles at the surface.

• The energy required to increase the surface area of a liquid

• Surface tension: the inward pull by particles under the surface.


Surface Tension

A result of an inward pull among the molecules of a liquid

– Brings the molecules at the surface closer together


• Adhesive Forces

• Substances bind to surfaces

• Causes a meniscus

• Cohesive Forces

• Binds molecules to each other

Practical Applications - LIQUIDS


• A solid is matter with a definite shape and volume.

• A solid does not take the shape of a container in which it is placed. This is because the particles of a solid are packed closely together.

Solids


States of Matter: Solids

• Particles are in constant motion.

• For a substance to be a solid rather than a liquid at a given temperature, there must be stronger attractive forces acting between particles in the solid.


Solids Liquids

High density

Definite shape and volume

Amorphous or Crystalline

depending on structure

Medium Density

Indefinite shape, definite

volume

Particles move around

Evaporate to create vapor

pressure


PROPERTIES OF SOLIDS

Density: In general, the particles of a solid are more closely

packed than those of a liquid. So most solids are more dense

than most liquids, but there are exceptions (wax, cork).

When solid and liquid states of a substance coexist, the solid

is more dense than the liquid, so the solid will sink in the

liquid.

Water is an important exception (ice floats in water). This is

because the H2O molecules are less closely packed in ice than

in liquid water.


Solids

Particles are packed together

and stay in a fixed position

– Vibrate

Definite shape

Definite volume


Types of Solids

Crystalline Solids

– Made up of crystals

Sand

Salt

Snow

sugar

Amorphous Solids

– Not arranged in a

regular pattern

Plastics

Rubber

glass


Solids

CrystallineAmorphous

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33

Fundamental Concept: Energy and Packing

• Non dense, random packing

• Dense, ordered packing

Dense, ordered packed structures tend to have

lower energies.

Energy

r

typical neighbor

bond length

typical neighbor

bond energy

Energy

r

typical neighbor

bond length

typical neighbor

bond energy


Solids

Amorphous solids

• show no definite structure, supercooled liquids.

SiO2 (glass), plastic, rubber

Crystalline solids• orderly, repeating, 3 dimensional pattern• atomic, ionic, and molecular


Phase ChangesT

emp

eratu

re E. Q=mCΔT

D. Q= mHvTb

B. Q =mHf

C. Q=mCΔT

Tm

A. Q= mCΔT

Thermal Energy (Heat)


Thermal Energy (heat)

• Phase Changes:

– B and D represent phase changes

– Occur at constant temperature

• Temperature Changes

– A,C, and E

– Temp is changing

– Sloping portion of the graph


Thermal Energy (heat )

• What is happening at each part of the graph:

• A. Substance is a solid. Can heat it up or cool it down along this line.

• B. Phase change: solid-liquid. The temperature at B (Tm) is the melting (freezing) point.

• C. Substance is a liquid. Can heat it up or cool it down along this line.

• D. Phase change: liquid-gas. The temperature at D (Tb) is the boiling (condensation) point.

• E. Substance is a gas. Can heat it up or cool it down along this line.


Heat Calculations

• Q = m c ∆ T

– Q = heat or thermal energy in Joules (J) or calories (cal)

– m= mass in grams (g)

– C = specific heat in J/(g oC) or cal/(g oC)

– ∆T= change in temperature in oC

• Q =mHf



– Hf = Heat of fusion (J/g or cal/g); use for liquid-solid phase

change


Heat Calculations

• Q =mHv



– Hv = Heat of vaporization (J/g or cal/g); use for gas-liquid

phase change

• If you move left to right:

– all the processes are endothermic (heat must be supplied).

• If you move right to left:

– all the processes are exothermic (heat is removed/released).


The six possible phase changes

Phase Changes


The boiling point of a substance is defined as the temperature

at which its vapor pressure equals the external atmospheric

pressure.

The molar heat of vaporization (ΔHvap) is the amount of heat

required to vaporize a mole of substance at its boiling point.

Phase Changes


The transformation of a liquid to a solid is called freezing.

The reverse process is called melting.

The melting point (freezing point) of a solid (or liquid) is the

temperature at which the solid and liquid phases coexist in

equilibrium.

ice ⇌ water

H2O(s) ⇌ H2O(l)

In dynamic equilibrium, the forward and reverse process are

occurring at the same rate.

Phase Changes


The molar heat of fusion (ΔHfus) is the energy required to

melt 1 mol of a solid.

Phase Changes


Heating curves:

Phase Changes

Solid

Boiling point Vapor

Liquid

Solid and liquid in

equilibrium

Liquid and vapor in

equilibrium

Time

Te

mp

era

ture

Melting point


Sublimation is the process by which molecules go directly from the solid

phase to the vapor phase.

Deposition is reverse process of sublimation.

The molar enthalpy of sublimation (ΔHsub) of a substance is the energy

required to sublime 1 mole of a solid.

Phase Changes

ΔHsub = ΔHfus + ΔHvap

Solid I2 in equilibrium with its vapor


Solids

Solids have “resistance” to changes in both shape and volume

Solids can be Crystalline or Amorphous

Crystals are solids that consist of a periodic array of atoms, ions, or molecules

– If this periodicity is preserved over “large” (macroscopic) distances the solid has “Long-range Order”

Amorphous solids do not have Long-Range Order

– Short Range Order


Atomic Packing: Metals and Ceramics

Cu2Mg

Al, Cu, Au, …

CdTe

Si

CrystallineRegular

arrangement of

atoms with

long range order


Types of Solids

Covalent

Molecular

(H2O)

Covalent

Network

(SiO2 - quartz)

Amorphous

(SiO2 - glass)


Types of Crystalline Solids

Type

Ionic Molecular Network Metal

StructureCrystal lattice of

ions

Crystal lattice of

molecules

Single 3-D crystal

of covalently

bonded atoms

Crystal of

individual metal

atoms

Interparticle

ForcesElectrostatic Intermolecular

Covalent

bondingMetal bond

Melting pointVery high

NaCl: 801 C

Typically < 300 C,

increases with

molar mass

Extremely high

SiO2: 1200 CVaries

Solubility in

WaterMany are soluble

Depends on

polarityInsoluble Insoluble

Conducts

ElectricityNo No No Yes

Other Dull, brittle Varies Very hardShiny, malleable,

ductile


Crystalline Solids

Crystalline state: defined by a 3-D periodic ordering of atoms

- perfect (no defect)

- infinite

Crystal: limited part of crystalline state

- smooth faces

- regular geometric shapes: set of equal faces (cube, octahedron,

prismatic…)

- possibility of cleavage (not possible for amorphous solids)

Geometric model of crystalline state

- infinite point lattices

- 3 non-coplanar vectors: a, b, c (in bold) or (with arrows)c,b,a


Crystalline solids

Diamond

A network solid

Sodium Chloride

An ionic solid

Water Ice

A molecular solid


Solids

Crystals Solids:– Short-range Order

– Long-range Order

Amorphous solids:– ~Short-range Order

– No Long-range Order


Types of Solids

• Crystalline - repeating geometric pattern

– covalent network

– metallic

– ionic

– covalent molecular

• Amorphous –

– no geometric pattern

decreasing

m.p.


What happens when many atoms come together to form a solid?

Regular structures (crystalline) or Irregular structures (amorphous)

Crystalline:• “A solid characterized by

• “long-range order” , “repetitive 3D pattern”

• Often opaque

• Most metals

• Portions of some polymers (“semi-crystalline”)

Amorphous (Non-crystalline)• “A solid which

• “without form”

• Often transparent

• Ceramic glasses

• Amorphous metals

• Some polymers are 100% amorphous


Quartz: crystalline SiO2 Glass: amorphous SiO2

What is the difference between quartz and glass?

Long-range order Only short-range order

http://en.wikipedia.org/wiki/Image:Window2006-05-21.JPG

http://en.wikipedia.org/wiki/Image:Window2006-05-21.JPG


States of Matter: Crystalline

• Crystalline solids: solids which have a well defined arrangement

• Flat surfaces - Definite angles

• A lattice: the 3D structure of a crystalline solid.


Particles in crystalline solids are arranged in

ordered manner.

NaCl (table salt) diamond

carbon

Chlorine

sodium

http://www.chem.neu.edu/Courses/1105Tom/05Lecture21.html





States of Matter: Crystals


• Crystals: the individual pieces of a crystalline solid.

• ex) quartz, diamond


Crystalline Solids

The molecules of crystalline solids are arranged in repeating

symmetrical patterns.

Metals

Minerals such as diamond

Salts

Ice


Crystalline solids

Type Unit particles

atomic Atoms (noble gases)

Covalent molecular Molecules (nonmetals)

covalent network

(strongest)

atoms connected by covalent bonds

(Cdia, Si, SiO2, SiC, Cgra)

ionic Ions (metal + nonmetal)

metallic atoms surrounded by mobile valence

electrons (metals)

• Crystalline solids can be classified into 5 categories based on the types of particles they contain:


States of Matter: Crystalline


• A unit cell: the smallest arrangement of connected points that can be repeated to form the lattice.


• Particles in amorphous

solids have a disordered

arrangement of

atoms/molecule.


States of Matter: Amorphous

• Amorphous solids: No orderly structure

• Lack of well defined faces or angles

• ex) rubber, glass

• Remember:• Ionic solids dissolve in water, conduct electricity and heat

• Covalent/Amorphous solids do not dissolve in water, will not conduct electricity and heat


• Some solids come together without forming crystal structures. Instead, the particles are found in a random arrangement.

Amorphous Solids

• These solids are called amorphous (uh MOR fuhs) solids.

• Rubber, plastic, and glass are examples of amorphous solids.


Amorphous solid doesn’t have discontinuity in the heating curve.

T(℃)

m.p.

Time0

T(℃)

Time0

Heating curve of crystalline solid and amorphous solid.


Amorphous Solids

Materials which don’t have the long range repetitive

pattern of crystals are called amorphous materials.

Amorphous means “without form”.

2SiOCeramic Compound

Crystalline Amorphous


Amorphous Solids

During the rapid cooling of a liquid, if atoms or

molecules do not find sufficient time to arrange

themselves in a long-range repetitive pattern amorphous

solids will form unlike crystalline solids obtained by

gradual cooling.

Glasses

Polymeric materials

Some ceramics


Amorphous Solids

Amorphous solids have molecules arranged in no

particular order.

Rubber

Wax

Some plastics


Amorphous Solids

Amorphous solids lack a regular three-dimensional arrangement

of atoms.

Glass is an amorphous solid.

Glass is a fusion product.

SiO2 is the chief component.

Na2O and B2O3 are typically fused with molten SiO2 and allowed

to cool without crystallizing.


Amorphous Solids


Amorphous Solids

Crystalline quartz Noncrystalline (amorphous)

quartz glass


Crystalline vs. Amorphous:

A crystalline solid has particles which are arranged in an orderly, geometric, 3-D structure.

– Examples: sodium chloride, ice, gems and minerals

In an amorphous solid, the particles are not arranged in any particular pattern.

– Examples: rubber, plastics, glass.


Amorphous glasses:

No long range order

Properties are:

Isotropic(same in all

directions)


Polymorphic Transformation

Materials having the same chemical composition can have

more than one crystal structure. These are called allotropic

or polymorphic materials.

– Allotropy for pure elements.

– Polymorphism for compounds.

These transformations result in changes in the properties of

materials and form the basis for the heat treatment fo steels

and alloys.


Carbon may exist in two forms:

Polymorphism

Graphite

( 2D layers)

Diamond

(3D structure)


POLYMORPHISM

Iron (Fe) may also exist in several forms:

BCC at room temperature → α iron

FCC at 910°C → γ iron

BCC at above 1400°C → β iron

Above 1539°C → liquid

Upon heating an iron from room temperature to above 910°C, its

crystal structure changes from BCC to FCC accompanied by a

contraction (reduction in volume).


Amorphous vs. Liquid

Amorphous solids have a structure that keeps them

together in a specific shape while a liquid will spread out to

fit whatever might be containing it.


Amorphous solids

Not all solids are crystalline in structure

Some are amorphous – they lack an ordered internal structure.

Rubber, plastic, and asphalt are amorphous solids

Glasses are a subgroup of amorphous solids

Transparent fusion of inorganic substances that have cooled to a rigid state without crystallizing

Sometimes called supercooled liquids

They have no fixed melting point


Phase Diagram

The lines represent the temperature / pressure range for equilibrium


Difference Between Amorphous Solids and Crystalline Solids

S.No. Amorphous solids Crystalline solids

1. Don't have definite geometrical shape. Characteristic geometrical shape.

2. melt over a wide range of temperature. They have sharp melting point

3. Physical properties are same in different

direction, i.e. isotropic.

Physical properties are different in different

directions. This phenomenon is known as

Anisotropy.

4. Amorphous solids are unsymmetrical. They are symmetrical

5. Don't break at fixed cleavage planes. Cleavage along particular direction at fixed

cleavage planes.


Phase Diagrams

A phase diagram is a graphical way to summarize the conditions under which the different states of a substance are stable.

• The diagram is divided into three areas

representing each state of the substance.

• The curves separating each area represent the

boundaries of phase changes.


Phase Diagrams

Below is a typical phase diagram. It consists of three curves that divide the diagram into regions labeled “solid, liquid, and gas”.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams

Curve AB, dividing the solid region from the liquid region, represents the conditions under which the solid and liquid are in equilibrium.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams

Usually, the melting point is only slightly affected by pressure. For this reason, the melting point curve, AB, is nearly vertical.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams

Curve AC, which divides the liquid region from the gaseous region, represents the boiling points of the liquid for various pressures.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams

Curve AD, which divides the solid region from the gaseous region, represents the vapor pressures of the solid at various temperatures.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams The curves intersect at A, the triple point,

which is the temperature and pressure where three phases of a substance exist in equilibrium.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.


Phase Diagrams

The temperature above which the liquid state of a substance no longer exists regardless of pressure is called the critical temperature.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Tcrit


Phase Diagrams The vapor pressure at the critical

temperature is called the critical pressure. Note that curve AC ends at the critical point, C.

B

temperature

pre

ssu

re

A

C

D

solid liquid

gas

.

.

Tcrit

Pcrit

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