Liquids and Solids

Liquids and Solids

…if it’s not a gas…

Well, duh.

Ingredients: Water

Physical Parameters of Water

• Formula

• FM

• Shape

• Polar?

• Density

Physical Parameters of Water

• Formula H2O

• FM 18.02 g/mol

• Shape bent, 104.5o

• Polar? Yes

• Density 1.00 g/ml

Thermal characteristics of Water

• MP

• BP

• C

• Hfus

• Hvap


• MP 0.0oC

• BP 100.0oC

• C 4.18 J/goC

• Hfus 6.0 kJ/mol

• Hvap 41 kJ/mol

If you add heat to matter, it may…

a)

b)

c)

d)

If you add heat to matter, it may…

a) warm up.

b) melt

c) boil

d) expand (tough to calculate, don’t bother)

Let’s try to warm up a cup of cold coffee.

Step 1: Add heat.


Step 1: Add heat.

Well, that was easy.


What if you add half as much heat?



a)

b)

c)



a) Raise the temperature only half as much.

b)

c)




b) Use half as much coffee (and cup)

c)




b) Use half as much coffee (and cup)

c) Use a different substance

The effect of heat, q!

• When something warms up:

The heat, q, depends on:

• The mass of the sample (m)

• The change in temperature (T)

• The nature of the sample (C)

The effect of heat (q)

• When something warms up:

The heat, q, depends on:

• The mass of the sample (m)

• The change in temperature (T)

• The nature of the sample (C)

C is the specific heat capacity for a given substance. Its units are (J/goC)

If you add heat to a sample, it may…

a) warm up. q=mCT

b) melt

c) boil

d) expand (tough to calculate, don’t bother)

q=mCT

• q – heat, in Joules

• m –mass, in grams

• C –specific heat capacity, in J/goC

• T—change in temperature (Tfinal-Tinitial)

Cwater=4.184 J/goC

• Cethanol =2.4 J/goC

• Cice =2.1 J/goC

• CAl =.90 J/goC

• CFe =.46 J/goC

• Cglass =.50 J/goC

• CAg =.24 J/goC

How much heat?

• How much heat does it take to raise 50.g water from 15oC to 80.oC?

• q=mCT

How much heat?


• q=mCT = 50.g x 4.18 J/goC x (80.oC-15oC)

How much heat?


• q=mCT = 50.g x 4.18 J/goC x (80.oC-15oC) = 50.g x 4.18 J/goC x (65oC)

How much heat?


• q=mCT = 50.g x 4.18 J/goC x (80.oC-15oC) = 50.g x 4.18 J/goC x (65oC)

=14000 J (14 kJ)

What is the change in temperature?

• If you add 1550 J to 12 g water, how much will it heat up?

• T =q/mC



• T =q/mC1550 J / (12 g x 4.18 J/goC )



• T =q/mC1550 J / (12 g x 4.18 J/goC )

= 31oC



• T =q/mC1550 J / (12 g x 4.18 J/goC )

= 31oC

If the temperature starts at 25oC, it will heat up to …



• T =q/mC1550 J / (12 g x 4.18 J/goC )

= 31oC

If the temperature starts at 25oC, it will heat up to 56oC

Calorimetry

• --the measurement of heat.

Calorimetry


• If one thing gains heat…

Calorimetry


• If one thing gains heat…

…something else lost it.

• If 75 g of a metal at 96oC is placed in 58 g of water at 21oC and the final temperature reaches 35oC, what is the specific heat capacity of the metal?

Step 1

• How much heat did the water gain?

Step 1

• How much heat did the water gain?

q=mCT

Mass of water, in grams

Specific heat of water, 4.18 J/goC

Change in the temperature of water, in oC

Step 2

• How much heat did the metal lose?

Step 2

• How much heat did the metal lose?

• Heat lost = - heat gained

• qlost=-qgained

Step 3

• What is the specific heat capacity of the metal?

Step 3

• What is the specific heat capacity of the metal?

C=q/mT

Mass of metal, in grams

Specific heat of metal, in J/goC

Change in the temperature of metal, in oC

Heat lost by metal



.74 J/goC

Heats of fusion and vaporization

• How much heat is required to melt 150 g of water (at its melting point)?



• (Step 1: Convert to moles)



• 150 g x 1mol/18.02 g =



• 150 g x 1mol/18.02 g = 8.3 mol



• 150 g x 1mol/18.02 g = 8.3 mol

• (Step 2: Apply the heat of fusion)



• 150 g x 1mol/18.02 g = 8.3 mol

• Q=nHf=8.3 mol x 6.01 kJ/mol =



• 150 g x 1mol/18.02 g = 8.3 mol

• Q=nHf=8.3 mol x 6.01 kJ/mol = 50.kJ


• How much heat is required to boil 250 g of water (at its boiling point)?



• (Step 1: Convert to moles)



• 250 g x 1mol/18.02 g =



• 250 g x 1mol/18.02 g = 13.9 mol



• 250 g x 1mol/18.02 g = 13.9 mol

• (Step 2: Apply the heat of vaporization)



• 250 g x 1mol/18.02 g = 13.9 mol

• Q=nHf=13.9 mol x 41 kJ/mol =



• 250 g x 1mol/18.02 g = 13.9 mol

• Q=nHf=13.9 mol x 41 kJ/mol = 570 kJ

Heat problems.

Remember:

Every substance has its own specific heat capacity,

heat of fusion, and

heat of vaporization!


• MP 0.0oC

• BP 100.0oC

• C 4.18 J/goC

• Hfus 6.0 kJ/mol

• Hvap 41 kJ/mol,

Thermal characteristics of acetone

• MP -95.4oC

• BP 56.3oC

• C .126 J/goC

• Hfus 5.7 kJ/mol

• Hvap 29 kJ/mol

Thermal characteristics of copper

• MP 1085oC

• BP 2567oC

• C .385 J/goC

• Hfus 13 kJ/mol

• Hvap 231 kJ/mol

Is this special?

• MP and BP of molecules of similar size• Formula FM(g/mol) MP (oC) BP (oC)

• CH4

• NH3

• H2O• HF• Ne

Is this special?


• CH4 16• NH3 17• H2O 18 = 18 g/mol ± 11%• HF 20• Ne 20

Is this special?


• CH4 16 -183• NH3 17 -78• H2O 18 0 • HF 20 -83• Ne 20 -249

For the covalent hydrogen compounds of the second period:

Melting points

-300

-250

-200

-150

-100

-50

0

0 2 4 6

Substance #

Mel

tin

g p

oin

t

Series1

Water!

Is this special?


• CH4 16 -183 -164• NH3 17 -78 -33• H2O 18 0 100• HF 20 -83 20• Ne 20 -249 -246

Boiling points

-300-250-200-150-100-50

050

100150

0 2 4 6

Substance #

Bo

ilin

g p

oin

t

Series1


Water!

Melting

• Melting occurs when particles have enough motion to escape their solid structure

Melting

• Melting occurs when particles have enough motion to escape their solid structure

• A substance whose particles stick together better has a higher melting point

Boiling

• Boiling occurs when particles have enough motion to escape their liquid neighbors

Boiling

• Boiling occurs when particles have enough motion to escape their liquid neighbors

• A substance whose particles stick together better has a higher boiling point

Liquids

• The liquid range is all of those temperatures where the particles move around each other, but are unlikely to escape

Liquids

• The liquid range is all of those temperatures where the particles move around each other, but are unlikely to escape

• A substance whose particles stick together better, even while moving, has a larger liquid range.

Therefore…

• Water molecules stick together very well—in a solid, and as a liquid.

Why do molecules stick together?

Why do molecules stick together?

• Attractions between molecules are called intermolecular forces (IM forces)

• Some forces are stronger than others.

Non-polar molecules…

Show dispersion forces

• very weak

• very brief, small charge imbalances due to the motion of electrons.

• They unbalance and attract their neighbors.

Polar molecules…

Show dipole interactions

• fairly weak.

• permanent, small charge imbalances due to the polarity of their bonds.

• They attract their polar neighbors.

(But, not all polar bonds are created equal)

• When hydrogen is the less electronegative end of a polar bond:

+ -

H Cl--the hydrogen is more positive

--it is losing custody of its last electron

Polar molecules with hydrogen…

Show hydrogen bonding

• strongest of the weak bonds.

• permanent, larger charge imbalances than other polar bonds.

• They attract their polar neighbors better.

…and if it’s not weak…

Strong IM forces include…

Ionic bonds (in ionic compounds)

Metallic bonds (in pure metals and alloys)

Covalent bonds (in covalent network solids)

(None of these particles are molecules, but they are still called intermolecular forces.)

Why do particles stick together?


• In order, from weakest to strongest:

Dispersion Forces

Dipole Interactions

Hydrogen Bonding

Ionic Bonds

Metallic Bonds

Covalent Bonds


• In order, from weakest to strongest:

Dispersion Forces —between non-polar molecules

Dipole Interactions —between polar molecules

Hydrogen Bonding -between polar molecules w/H

Ionic Bonds —between ions

Metallic Bonds —between metal atoms

Covalent Bonds —in a network solid


If you are given a substance:

• --describe the type of substance

• --describe the strongest IM force between the particles

• --you may be asked to compare it to another substance

What kind of substance?

• barium

• chlorine

• tin (II) chloride

• sulfur dioxide

• solid sulfur

• helium

• dinitrogen monoxide• iron• sodium oxide• iodine• barium sulfide• sulfuric acid


• Watch out for a trick question.


• Watch out for a trick question.

--Ready?

Quiz

Q: What holds water together?

Quiz

Q: What holds water together?

There are TWO answers.

HA! It’s a trick question!

Answer 1:

• Polar covalent bonds between the hydrogen and oxygen atoms hold the atoms together as water molecules

• Answer 2:

• Hydrogen bonding attracts water molecules to each other as a liquid or a solid.

Is this special?

• Formula Type of substance• CH4

• NH3

• H2O• HF• Ne

Is this special?

• Formula Type of substance• CH4 non-polar covalent molecule• NH3 polar covalent molecule• H2O polar covalent molecule• HF polar covalent molecule• Ne non-polar individual atoms

Is this special?

• Formula Type of IM Forces• CH4 dispersion forces• NH3 hydrogen bonding• H2O hydrogen bonding• HF hydrogen bonding• Ne dispersion forces


Melting points

-300

-250

-200

-150

-100

-50

0

0 2 4 6

Substance #

Mel

tin

g p

oin

t

Series1

Water!


If you are given a substance:

• --describe the type of substance

• --describe the strongest IM force between the particles

• --you may be asked to compare it to another substance

Which has stronger intermolecular forces, NaCl or HCl?


• NaCl: ionic compound

HCl: polar covalent molecule




• NaCl: held together by ionic bonds

HCl molecules: attracted to each other by hydrogen bonds.




• NaCl: held together by ionic bonds

HCl molecules: attracted to each other by hydrogen bonds..

• The ionic bonds in NaCl are stronger than hydrogen bonds between HCl molecules

What kind of IM forces?

• barium

• chlorine


• sulfur dioxide

• solid sulfur

• helium


Rank in order of strength of IM forces.

• barium

• chlorine


• sulfur dioxide

• solid sulfur

• helium


Compare N2 and CO

• What type of substance?

• What type of IM forces

• Which is stronger?

• What will this do to the MP and BP?

There is an overlap.

• The strongest dispersion forces are stronger than average dipole interactions

• In general, a larger molecule has stronger dispersion forces.

• There is a big overlap between ionic and metallic bonds.

“’The time has come’, the walrus said…”

• The stronger the IM forces, the higher the:

MP, BP, Hfus, Hvap, C, surface tension, cohesion, viscosity,

strength and hardness of the solid…

…etc. Usually.

List in order of MP (low to high)

• barium

• chlorine


• sulfur dioxide

• solid sulfur

• helium


Properties of Substances

• In a pure substance, particles are identical.

• The way each particle holds its electrons defines the physical properties

• Melting and boiling separate the particles, it won’t destroy them

Particles can be…

• Molecules—for molecular compounds and non-metallic elements

• Atoms—for metals and noble gasses

• Formula units —for ionic compounds. The formula unit can be dissociated (separated) by physical means, but the ions cannot be isolated.

What kind of particles?

• barium• chlorine gas• tin (II) chloride• sulfur dioxide• water• solid sulfur• nitrogen gas• helium gas• nitrogen dioxide

• nitrous oxide• iron• sodium oxide• iodine• gold• barium sulfide• ammonia• sulfuric acid

Identical particles make regular crystals

• Teacher Imitation Day Spring 2010

Ionic compounds

• --are composed of identical formula units of (+) and (-) ions

• --have valence e- transferred to (-) ions

• --make a regular crystal

Simple cubic

Body centered cubic

Face centered cubic

Hexagonal

Rhombohedral

Tetragonal

Orthorhombic

Triclinic

Monoclinic

Ionic compounds

• --are hard but brittle. If you deform the crystal, positive ions meet and repel. The crystal shatters

• --have high melting and boiling points

Ionic compounds

• --might dissolve in water

• When ionic compounds are molten (melted) or aqueous (dissolved in water), the ions dissociate (separate).

Free ions can carry a current.

Metals, generally…

• --are solid at room temperature

• --are malleable and ductile. If you deform the metal, the sea of valence electrons still attract the new arrangement of nuclei.

• --dissolve in each other.

Electrons don’t care what nuclei are inside.

Covalent compounds

• Can be solid, liquid or gas at room temperature.

• Are generally soft, and easily melted or boiled.

Each molecule stands alone with its electrons

Define:

• Melting point Boiling point

• Heat of vaporizationHeat of fusion

• Specific heat capacity Adhesion

• Cohesion Surface tension

• Density Solubility

• Solution Solute

• Solvent Dissociation

Solutions: A solution is…

• --a homogeneous mixture


• --a homogeneous mixture

Components are mixed at a molecular level.

Any two samples of the same solution will have identical proportions of the components


• --a solute dissolved in a solvent


• --a solute dissolved in a solvent

Usually there is more solvent in a solution

The solvent is usually a liquid or a gas


• --a physical combination of indefinite proportions


• --a physical combination of indefinite proportions

Dissolving is a physical (not chemical) process

Two solutions can have different proportions

The components retain their own chemical and physical properties

Oh, yeah…

• With two gasses or two liquids that dissolve in each other (miscible liquids), either one can be the solvent

• Aqueous (aq) = “dissolved in water”

Oil and water don’t mix.

Why not?

“Like dissolves like”

• Water is a polar solvent, oil is a non-polar substance.

• They are not alike

Non-polar solvents dissolve non-polar solutes

Polar solvents dissolve polar and ionic solutes

Metallic solvents dissolve metallic solutes.

“Like dissolves like”

Does it dissolve?

1) CH3OH/H20

2) KBr/H2O

3) CCl4/H2O

4) S8/H2O

5) Hg/H2O6) Ag/Hg

7) S8/CCl48) NaCl/KBr

Does it dissolve?

1) CH3OH/H20—polar/polar =YES

2) KBr/H2O

3) CCl4/H2O

4) S8/H2O

5) Hg/H2O6) Ag/Hg

7) S8/CCl48) NaCl/KBr

Does it dissolve?

1) CH3OH/H20 —polar/polar =YES

2) KBr/H2O —ionic/polar =YES

3) CCl4/H2O —nonpolar/polar =NO

4) S8/H2O —nonpolar/polar =NO

5) Hg/H2O —metallic/polar =NO6) Ag/Hg —metallic/metallic =YES

7) S8/CCl4 —nonpolar/nonpolar=YES8) NaCl/KBr —ionic/ionic(if melted)=YES

How?

• Particles on the surface of a solid get surrounded by solvent particles (solvation) and lifted out of the solute.

The solvation of sodium and chloride ions

Water is special.

• The attraction of the polar water molecules lifts polar molecules and individual ions out of solids.

• Ions dissociate, making an electrolyte solution

Colloids and emulsions

• Colloids and emulsions are (barely) heterogeneous mixtures

• The particles are just barely too large to be called “molecular sized”

• Colloids and emulsions do not separate themselves, but appear cloudy or opaque

Suspensions

• If the particles are too large to dissolve or form a colloid, they can still be suspended in a fluid.

• Suspensions settle out eventually.

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Liquids and Solids