Unit 2: Classification of Matter.

www.unit5.org/chemistry

Unit 2: Classification of Matter

http://www.unit5.org/chemistry

www.unit5.org/chemistry

Unit 2

“Chemical change has always been part of the Universe, even before human beings evolved. Indeed, scientists believe that life

began on Earth as a result of complex chemicals reproducing themselves over billions of years. Chemistry is a physical

science; it lies between the biological sciences helping to explain many of life’s processes, and the laws of physics, which include

matter and energy. Chemical processes are constantly occurring within us – when our bodies move, a series of chemical reactions takes place to give the muscles the energy that is taken in from

food. Many species of the animal world make use of chemistry to defend themselves, to kill their prey, and to build fragile structures

that have incredible strength. Modern methods of chemical analysis have led to greater understanding of the chemistry of

nature, so that it is possible to identify those chemical compounds that produce the color, taste, and smell of a flower or a fruit.”

Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 8

http://www.chem1.com/acad/webtext/pre/index.html

http://www.unit5.org/chemistry

Classification of Matter

MATTER(gas. Liquid,

solid, plasma)

PURESUBSTANCES MIXTURES

HETEROGENEOUSMIXTURE

HOMOGENEOUSMIXTURESELEMENTSCOMPOUNDS

Separated by

physical means into

Separated by

chemical means into

Kotz & Treichel, Chemistry & Chemical Reactivity, 3rd Edition , 1996, page 31

MatterMatter

SubstanceDefinite composition

(homogeneous)

SubstanceDefinite composition

(homogeneous)

Element(Examples: iron, sulfur,

carbon, hydrogen,oxygen, silver)



Mixture ofSubstances

Variable composition



Compound(Examples: water.

iron (II) sulfide, methane,Aluminum silicate)



Homogeneous mixtureUniform throughout,also called a solution

(Examples: air, tap water,gold alloy)



Heterogeneous mixtureNonuniform

distinct phases(Examples: soup, concrete, granite)



Chemicallyseparable

Physicallyseparable

The Organization of Matter

MATTER

PURESUBSTANCES


HOMOGENEOUSMIXTURES

ELEMENTS COMPOUNDS

Physical methods

Chemical methods

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 41

1) Chemistry:The study of matter, the changes matter goes through, and the associated energy changes.2) Matter:

Anything that has mass and takes up space (volume)

Pure SubstanceA sample of matter in which all parts have the same properties

Mixture2 or more substances that are physically combined – individual properties are

retained

ElementA substance that can’t be broken down into any other substance by ordinary chemical change

CompoundA substance made of 2 or more elements chemically

combined

HomogeneousConsistent properties

throughout

HeterogeneousUneven, inconsistent distribution

of particles

8) Mixtures vs Elements & Cmpds

a)Mixtures retain properties of constituents

b)Composition of a mixture can vary

c)Mixtures can be homogeneous or heterogeneous

10) Ways to Make Mixtures

a)Element + 1 or more elements

b)Cmpd + 1 or more elements

c)Cmpd + 1 or more cmpds

Gold

24 karat gold 18 karat gold 14 karat gold

Gold

Copper

Silver

18/24 atoms Au24/24 atoms Au 14/24 atoms Au

http://en.wikipedia.org/wiki/Gold

Elements, Compounds, and Mixtures12) If you have H2 and O2 in a container,

do you have water?

(a)an element(hydrogen)

(b)a compound(water)

(c)a mixture(hydrogen and oxygen)

(d)a mixture(hydrogenand oxygen)

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 68

hydrogenatoms hydrogen

atoms

oxygen atoms

= Oxygen atom = hydrogen atom

MatterMatter

Pure SubstanceDefinite composition

(homogeneous)

Pure SubstanceDefinite composition

(homogeneous)





















Chemicallyseparable

Physicallyseparable

How would you categorize elements and compounds?

During a“physical change”

a substance changes some physical property…

During a“physical change”

a substance changes some physical property…

H2O

…but it is still the same material with the same chemical composition.

…but it is still the same material with the same chemical composition.

H2Ogas

solid

liquid

Physical and Chemical PropertiesExamples of Physical Properties –

properties that can be observed without changing the substance

Boiling point Color Slipperiness Electrical conductivity

Melting point Taste Odor Dissolves in water

Shininess (luster) Softness Ductility Viscosity (resistance to flow)

Volatility Hardness Malleability Density (mass / volume ratio)

Examples of Chemical Properties – the ability of a substance to undergo a chemical reaction

Burns in air Reacts with certain acids Decomposes when heated

Explodes Reacts with certain metals Reacts with certain nonmetals

Tarnishes Reacts with water Is toxic

Ralph A. Burns, Fundamentals of Chemistry 1999, page 23Chemical properties can ONLY be observed during a chemical reaction!

The formation of a mixture

The formation of a mixture

The formation of a compound

The formation of a compoundChemical Change

Chemical Change

Physical Change

Physical Change

Physical & Chemical Changes

Limestone,CaCO3

crushing

PHYSICALCHANGE

Crushed limestone,CaCO3

heating

CHEMICALCHANGE

PyrexPyrex

CO2

CaO

Lime andcarbon dioxide,

CaO + CO2

PyrexPyrex

O2

H2O

PyrexPyrex

H2O2

Light hastens the decomposition of hydrogen peroxide, H2O2. The dark bottle in which hydrogen peroxide is usually storedkeeps out the light, thus protecting the H2O2 from decomposition.

Sunlight energy

Properties of Matter

http://antoine.frostburg.edu/chem/senese/101/matter/slides/sld001.htm

PyrexPyrex PyrexPyrex

ExtensiveProperties

IntensiveProperties

volume:mass:

density:temperature:

100 mL99.9347 g

0.999 g/mL20oC

15 mL14.9902 g

0.999 g/mL20oC

Solubility – a measure of the amount of solute that can be dissolved in a solvent at a given temperature.

Dissolving of Salt in Water

NaCl(s) + H2O Na+(aq) + Cl-(aq)

Cl-

ions

Na+

ions Water molecules

Dissolving of NaCl

Timberlake, Chemistry 7th Edition, page 287

HH

O

Na+

+

-- + -+

+

-

Cl-

+ -

+

hydrated ions

Density

• Density is an INTENSIVEINTENSIVE property of matter.

- does NOT depend on quantity of matter. - color, melting point, boiling point, odor, density

• Contrast with EXTENSIVEEXTENSIVE

- depends on quantity of matter.- mass, volume, heat content (calories)

Styrofoam Brick

http://en.wikipedia.org/wiki/Intensive_and_extensive_properties

http://en.wikipedia.org/wiki/Density

Styrofoam Brick

?It appears that the brick is ~40x more dense than the Styrofoam.

MMMM

VV= =DD

VVDD

BrickBrickStyrofoamStyrofoam

Styrofoam Brick

Determining the volume of an irregular solid

Vfinal = 98.5 cm3

- Vinitial = 44.5 cm3

Vfishing sinker = 54.0 cm3

Before immersion

Water

44.5 cm3

After immersion

Fishing sinker

98.5 cm3

Thread

Density

D

M

Vensity

ass

olume

D = M V

M = D x V

V =M D

Cube Representations

1 m3 = 1 000 000 cm3


Consider Equal Volumes

The more massive object(the gold cube) has the_________ density.

Equal volumes…

…but unequal masses

aluminum gold

GREATER

Density = Mass

Volume


Consider Equal MassesEqual masses……but unequal volumes.

The object with the larger volume (aluminum cube) has the density.

aluminum

gold


smaller

Christopherson Scales

Made in Normal, Illinois USA

Two ways of viewing density


Equal volumes…

…but unequal masses

The more massive object(the gold cube) has thegreater density.

aluminum gold

(A)

Equal masses……but unequal volumes.

(B)

gold

aluminumThe object with the larger volume (aluminum cube) has the smaller density.

Density of Some Common Substances


Substance Density (g / cm3)

Air 0.0013* Lithium 0.53 Ice 0.917 Water 1.00 Aluminum 2.70 Iron 7.86 Lead 11.4 Gold 19.3


Substance Density (g / cm3)

Air 0.0013* Lithium 0.53 Ice 0.917 Water 1.00 Aluminum 2.70 Iron 7.86 Lead 11.4 Gold 19.3

*at 0oC and 1 atm pressure

Which liquid has the highest density?

52

3

1

4

Coussement, DeSchepper, et al. , Brain Strains Power Puzzles 2002, page 16

least dense 1 < 3 < 5 < 2 < 4 most dense

Specific Gravity

Jaffe, New World of Chemistry, 1955, page 66

0.90.25

water 1.0

ice

cork

aluminum

2.7

reference

substance Density

Density

It is a unitless quantity that expresses a ratio of the substance’s density compared to the reference’s density

Chapter 13: States of Matter Solid, Liquid, and Gas (13.3) (13.2) (13.1)


Gas Liquid Solid

Average Kinetic Energy and Temperature – they’re the same thing

Kinetic energy

Fra

ctio

ns o

f pa

rtic

les

Average KE1 = lower temperature

Average KE2 = higher temperature

minimum energyfor reaction

Hot vs. Cold Tea

Kinetic energy

Many molecules have anintermediate kinetic energy

Few molecules have avery high kinetic energy

Low temperature(iced tea)

High temperature(hot tea)

Perc

ent o

f mol

ecul

es

~~~

SOLIDS – Chapter 13, Section 3 (p. 396-399)

True solids, or crystaline solids, have a crystal lattice structure – a characteristic, geometric arrangement of particles in a solid

Amorphous(Glass)Crystalline

Crystalline Amorphous

Particle Arrangement

Regular, geometric lattice structure

Irregular, no specific internal order

Shape Characteristic crystal structure

Irregular(broken glass)

Attractive Forces

Strong Variable / Weak

Melting Point

Definite, specific Indefinite – softens gradually

Examples Quartz Glass

Diamond Butter

Allotropes of Carbon

Graphite


Diamond


Diamond

http://en.wikipedia.org/wiki/Diamond

Sodium Chloride Crystal


= Cl-

= Na+

crystal lattice structure

a characteristic,

geometric

arrangement of

particles in a solid

Macromolecules and Allotropes of Carbon

Graphite BuckminsterfullereneDiamond



C60 & C70

“Buckytubes”

Buckminsterfullerene“Buckyballs”

Fullerenes

http://en.wikipedia.org/wiki/Fullerenes

Credit: Baughman et al., Science 297, 787 (2002)

Trojan Horse

Can use ‘camouflage’ to hide things. Be careful what’s in the Trojan!

Buckyballs can hide medicine to treat the human body.

Solid

H2O(s) Ice


Liquid

H2O(l) Water


In a liquid• molecules are in constant motion

• there are appreciable intermolecular forces

• molecules are close together

• Liquids are almost incompressible

• Liquids do not fill the container

Gas

H2O(g) Steam


Characteristics of GasesGases expand to fill any container.

– random motion, no attraction

Gases are fluids (like liquids).– no attraction

Gases have very low densities.– no volume = lots of empty space

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Phase Changes

Kinetic Theory

Kinetic Molecular Theory

• The main assumptions of KMT are…– Gases are made of tiny, individual particles.– The particles move in rapid, random straight-line motion.– Collisions between particles are elastic


Kinetic Molecular Theory• Particles in an ideal gas…

– have no attractive forces between them. – have no individual volume.


Kinetic Molecular Theory• Particles in an ideal gas…

– have no attractive forces between them. – have no individual volume.

Why do we use a gas that doesn’t really exist?– It simplifies the model and makes it easier to use– Under most normal conditions real gases behave like ideal

gases


Real Gases

• Particles in a REAL gas…– have their own volume– attract each other

• Gas behavior is most ideal…– at low pressures– at high temperatures– in nonpolar atoms/molecules


8

Elastic vs. Inelastic Collisions

8v1

elastic collision

inelastic collision

v2

v3 v4

Model Gas Behavior

• All collisions must be elastic • Take one step per beat of the

metronome • Container

– Class stands outside tape box

• Higher temperature – Faster beats of metronome

• Decreased volume– Divide box in half

• More Moles – More students are inside box

Mark area of container with tape on ground.

Add only a few molecules of inert gas

Increase temperature Decrease volume Add more gas Effect of diffusion Effect of effusion

(opening size)

Average Kinetic Energy and Temperature – they’re the same thing

Kinetic energy

Fra

ctio

ns o

f pa

rtic

les

Average KE1 = lower temperature


minimum energyfor reaction

Microscopic view of a liquid near its surfaceThe high energymolecules escapethe surface.

Kinetic energy

Fra

ctio

ns o

f pa

rtic

les Average KE1 = lower temperature


minimum energyto change phase

Evaporation

H2O(g)molecules

(water vapor)

H2O(l)molecules

A dynamic equilibrium can only be achieved in a closed container

Liquid/Vapor Dynamic EquilibriumThe two key properties we need to describe are

EVAPORATIONEVAPORATION and its opposite CONDENSATIONCONDENSATIONHH22O (O (l l ) ) →→ H H22O (O (gg) H) H22O (O (gg) ) → H→ H22O (O (l l ) )

At At equilibriumequilibrium the the raterate of evaporation of evaporation is equal to the is equal to the raterate of condensation of condensation

add energy and break intermolecular attractions

EVAPORATION

release energy and form intermolecular attractions

CONDENSATION

Water Molecules in Liquid and

Steam

Microscopic view of a liquid near its surface

The high energymolecules escapethe surface.

•Evaporation can only take place at the surface of the liquid

•At higher temperatures, more particles can escape intermolecular attractions

Behavior of a liquid in a closed container

A dynamic equilibriumcan only be achieved in a closed container

• To evaporate, molecules must have sufficient energy to break IM forces.

• Molecules at the surface break away and become gas.

• Only those with enough KE escape.• Breaking IM forces requires energy. The

process of evaporation is endothermicendothermic.• Evaporation is a cooling process.• It requires heat.

Evaporation

Change from gas to liquid

Achieves a dynamic equilibrium with vaporization in a closed system.

What is a closed system?

A closed system means matter can’t go in or out. (put a cork in it)

What the heck is a “dynamic equilibrium?”

Condensation

When first sealed, the molecules gradually escape the surface of the liquid.

As the molecules build up above the liquid - some condense back to a liquid.

The rate at which the molecules evaporate and condense are equal.

Dynamic Equilibrium

As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense.

Equilibrium is reached when:Rate of Vaporization = Rate of Condensation

Molecules are constantly changing phase “dynamic”

The total amount of liquid and vapor remains constant “equilibrium”

Dynamic Equilibrium

Vapor Pressure

more“sticky”

less likely tovaporize

In general:LOW v.p.

not very“sticky”

more likely tovaporize

In general:HIGH v.p.

measure of the tendency for liquid particles to enter gas phase at a given temp.

a measure of “stickiness” of liquid particles to each other

NOT all liquids have same v.p. at same temp.

Vapor Pressure of:

Propanone @ 40°C =

Ethanol @ 40°C =

Water @ 40°C =

58 kPa

17 kPa

8 kPa

101.3 kPa is Standard Atmospheric Pressure

Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Component Percent composition

Nitrogen, N2 78%Oxygen, O2 21%Argon, Ar 0.9%Water, H2O 0 – 4% (variable)Carbon dioxide, CO2 0.034% (variable)

The Earth’s AtmosphereFrom Space

Pressure

KEY UNITS AT SEA LEVELKEY UNITS AT SEA LEVEL

101.3 kPa (kilopascal)

1 atm

760 mm Hg

760 torr

14.7 psi

1013 mbar

14.7 psi


2m

NkPa

Sea level

Formation of a bubble is opposed by the pressure of the atmosphere

When the vapor pressure is equal to atmospheric pressure the bubble can expand and the liquid boils


0 20 40 60 80 1000

20

40

60

80

100

TEMPERATURE (oC)

PRESSURE (kPa)

CHLOROFORM

ETHANOL

WATER

Volatile substances evaporate easily (have high v.p.’s).

BOILING when vapor pressure = confining pressure (usually from atmosphere)

b.p. = 78oC

b.p. = 100oC

atmospheric pressure is 101.3 kPa

Vapor Pressure

93.3

80.0

66.6

53.3

40.0

26.7

13.3

0 10 20 30 40 50 60 70 80 90 100

61.3oC 78.4oC 100oC

chlo

rofo

rm

ethy

l alc

ohol

water

Pre

ssur

e (K

Pa)

Temperature (oC)

101.3

Boiling vs. EvaporationBoiling point: temperature at which

vapor pressure = atmospheric pressure

Revolutionary process - fast

AIRPRESSURE

90 kPa

VAPORPRESSURE

90 kPa

Normal Boiling point: temperature at which vapor pressure = standard atmospheric pressure

AIRPRESSURE

101.3 kPa

VAPORPRESSURE

101.3 kPa

Evaporation vs. Boiling

Boiling: Change from liquid to gas at the boiling point temperature

Evaporation: molecules change from liquid to gas phase below boiling point temperature.

If water is boiling at 89°C, what is the pressure?

Think: if a liquid is boiling, then it MUST have a VAPOR PRESSURE = TO ATMOSPHERIC PRESSURE


The pressure is approximately 65 kPa

Boiling Point of:

Propanone @ 80 kPa =

Ethanol @ 80 kPa =

Water @ 80 kPa =

48°C


73°C

94°C

Boiling Point on Mt. Everest

Water exerts a vapor pressure of 101.3 kPa at a temperature of 100 oC. This is defined as its normal boiling point: ‘vapor pressure = atmospheric pressure’

101.3

93.3

80.0

66.6

53.3

40.0

26.7

13.3

0 10 20 30 40 50 60 70 80 90 100

61.3oC 78.4oC 100oC

chlo

rofo

rm

ethy

l alc

ohol

water

Temperature (oC)

Pre

ssur

e (K

Pa)

On top of Mt. Everest

760 mm Hgx kPa = 253 mm Hg

101.3 kPa= 33.7 kPa

Why is boiling a cooling process?Which particles would change phase andhow would that effect the average KE?

Kinetic energy

Fra

ctio

ns o

f pa

rtic

les

minimum energyTo change phase

Average KE = temperature

LiquefactionA gas will change from a gas to a liquid under

conditions of:• High pressure = particles are closer together

causing ↑ attractive forces• Low temperature = particles move slower

allowing more attractive forces to develop

Heating CurvesEnergy is added at a constant rate

over time

Tem

per

atu

re (

oC

)

40

20

0

-20

-40

-60

-80

-100

120

100

80

60

140

Time

Melting - PE

Solid - KE

Liquid - KE

Boiling - PE

Gas - KE

Heating Curves (Chapter 17)

• Temperature Change– change in KE (molecular motion) – depends on the specific heat capacity

• Specific Heat (Cstate)– energy required to raise the temp of 1 gram of a

substance by 1°C– water has a very high specific heat capacity


Specific Heats of Some Substances

Specific Heat

Substance (cal/ g oC) (J/g oC)

Water 1.00 4.18Alcohol 0.58 2.4Wood 0.42 1.8Aluminum 0.22 0.90Sand 0.19 0.79Iron 0.11 0.46Copper 0.093 0.39Silver 0.057 0.24Gold 0.031 0.13

Heating Curves• Phase Change

– change in PE (molecular arrangement)– temp remains constant

• Heat of Fusion (Hf)– energy required to melt 1 gram of a substance at its

m.p.

Heating Curves• Heat of Vaporization (Hv)

– energy required to change 1 gram of a substance from liquid to gas

– usually larger than Hf…why?


Joule (J ) –

The SI unit used to measure the amount of heat absorbed or released during a reaction.

Calculating Energy Changes - Heating Curve for Water

q = heat m = mass Cstate = specific heat ΔT = change in temp

Tem

per

atu

re (

oC

)

40

20

0

-20

-40

-60

-80

-100

120

100

80

60

140

Time

q = m x Hf

q = m x Hv

q = m x Cwater x

q = m x Csteam x

q = m x Cice x

Csteam = 1.86 J / g • °C

Hf = 334 J / gHv = 2260 J / g

Cwater = 4.18 J / g • °C

Cice = 2.14 J / g •°C

112: Calculating Energy Changes - Heating Curve for Water

How many Joules of heat are needed to completely

melt 100g of ice at 0°C to water at 0°C ?

Tem

per

atu

re (

oC

)

40

20

0

-20

-40

-60

-80

-100

120

100

80

60

140

Time

q = m x HfHf = 334 J / g

q = m Hf

q = (100g)(334 J/g)

q = 33,400 J


How many Joules are needed to completely boil 200g of water at 100°C to steam at 100°C ?

Tem

per

atu

re (

oC

)

40

20

0

-20

-40

-60

-80

-100

120

100

80

60

140

Time

q = m x HvHv = 2260 J / g

q = m Hv

q = (200g)(2260 J/g)

q = 452,000 J


How many Joules are needed to change the temperature of 50g of water from 25°C to 95°C ?

Tem

per

atu

re (

oC

)

40

20

0

-20

-40

-60

-80

-100

120

100

80

60

140

Time

q = m x Cwater x Cwater = 4.18 J / g • °C

q = m Cwater ΔT

q = (50g)(4.18J/g • °C)(70°C)

q = 14,630 J

ΔT = 95°C - 25°C = 70°C

Cooling Curve for Water – energy is being removed at a constant rate

solid

liquid

gas

Heat added

Te

mp

era

ture

(oC

)

A

B

C

DE

Heating Curve for Water

0

100

LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

Energy Changes Accompanying Phase Changes Exothermic – Energy is released Endothermic – Energy is absorbed

Solid

Liquid

Gas

Melting Freezing

Deposition

CondensationVaporization

Sublimation

Ene

rgy

of s

yste

m

End

othe

rmic

cha

nges

Brown, LeMay, Bursten, Chemistry 2000, page 405

Energy of system

Exotherm

ic changes

solid

liquid

gas

vaporization

condensation

melting

freezing

Heat added

Te

mp

era

ture

(oC

)

A

B

C

DE

Heating Curve for Water

0

100

LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

Endotherm

ic Changes

Exotherm

ic Changes

Reaction RatesAmount of reactant used or product

produced per unit of time

A catalyst is a substance that increases the rate of a chemical reaction.

Ex., Enzymes, catalytic convertors

An inhibitor is a substance that decreases the rate of a chemical reaction.

Ex., Preservatives

Elements, Compounds and MixturesClassification of Matter

MATTER(gas. Liquid,

solid, plasma)

PURESUBSTANCES MIXTURES


HOMOGENEOUSMIXTURESELEMENTSCOMPOUNDS

Separated by

physical means into

Separated by

chemical means into

Kotz & Treichel, Chemistry & Chemical Reactivity, 3rd Edition , 1996, page 31

Classification of Matter

uniformproperties?

fixedcomposition?

chemicallydecomposable?

no

no

no

yes

hetero-geneousmixture

solution

element

compound

http://antoine.frostburg.edu/chem/senese/101/matter/slides/sld003.htm

The Chemical Parts List

Diatomic Elements, 1 and 7H2

N2 O2 F2

Cl2

Br2

F2

The 7 Diatomic Elements

Chemical Formulas – shorthand notation for a chemical compound

5 H2SO4Subscript – tells the number of atoms of the preceding element in one formula unit

Coefficient – tells the number of formula units in the entire expression

Chemical Formula – tells the number of atoms of each element in one formula unit of the substance

Number of atoms of an element in an each formula unit is given by the subscript.

Coefficient x Subscript = Number of atoms in an expression

What Do Chemical Formulas Represent?

Chemical Formulas – shorthand notation for a chemical compound

Expression # of formula units

# of atoms of each type in a formula unit

# of atoms of each type in entire expression

Total atoms in the expression

2 NaOH

NH4OH

5 Ca(OH)2

Subscripts outside of parentheses apply to everything inside the parentheses

2

1

5

7

25

8Na = 1 O = 1

H = 2

Na = 2 O = 2

H = 4

Ca = 1

O = 2 x 1 = 2

H = 2 x 1 = 2

N = 1 O = 1

H = 4 +1 = 5

N = 1 O = 1

H = 4 +1 = 5

Ca = 5

O = 10

H = 10

Methods of Separating Mixtures

• Filtration

• Evaporation

• Distillation

• Fractional Distillation

• Crystallization

• Decanting

Filtration separates

a liquid from an

insoluble solid


Mixture ofsolid andliquid Stirring

rod

Filtrate (liquidcomponentof the mixture)

Filter papertraps solid

Funnel

A Distillation Apparatusused to separate a soluble solid from a

liquid when either or both are to be retained

liquid with a soliddissolved in it

thermometer

condenser

tube

distillingflask

pure liquid

receiving flaskhose connected to

cold water faucetDorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 282

The solution is boiled and steam is driven off.


Salt remains after all water is boiled off.


No chemical change occurs when salt water is distilled.


Saltwater solution(homogeneous mixture)

Distillation(physical method)

Salt

Pure water

long spout helpsvapors to condense

mixture for distillation placed in here

Furnace

Glass retortA Hero’s Fountain

Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 13

Fractional Distillation: separation of two or more miscible liquids based on deiiference

in boiling point

Used in petroleum industry to separate crude oil into its varoius components

Decanting is used to separate a liquid from an insoluble solid when either the liquid,

solid or both are to be retained

Separation of a sand-saltwater mixture.


Separation of Sand from Salt1. Gently break up your salt-crusted sand with a plastic spoon.

Follow this flowchart to make a complete separation.

Salt-crusted

sand.

Dry

sand.

Wetsand.

Weigh themixture.

Decant clearliquid.

Evaporateto

dryness.

Pour intoheat-resistant

container.

Fill with water.

Stir and letsettle 1 minute.

Weighsand.

Calculateweight of

salt.

Repeat3 times?

Yes

No

2. How does this flowchart insure a completeseparation?

Unit 2: Classification of Matter.

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Transcript of Unit 2: Classification of Matter.