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# Liquids and Solids - PowerPoint PPT Presentation

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 H 2 O FM 18.02 g/mol Shape bent, 104.5 o Polar? Yes Density 1.00 g/ml.

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### Liquids and Solids

…if it’s not a gas…

Ingredients: Water

• Formula

• FM

• Shape

• Polar?

• Density

• Formula H2O

• FM 18.02 g/mol

• Shape bent, 104.5o

• Polar? Yes

• Density 1.00 g/ml

• MP

• BP

• C

• Hfus

• Hvap

• MP 0.0oC

• BP 100.0oC

• C 4.18 J/goC

• Hfus6.0 kJ/mol

• Hvap41 kJ/mol

a) warm up.

b) melt

c) boil

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

Well, that was easy.

What if you add half as much heat?

What if you add half as much heat?

a)

b)

c)

What if you add half as much heat?

a) Raise the temperature only half as much.

b)

c)

What if you add half as much heat?

a) Raise the temperature only half as much.

b) Use half as much coffee (and cup)

c)

What if you add half as much heat?

a) Raise the temperature only half as much.

b) Use half as much coffee (and cup)

c) Use a different substance

• When something warms up:

The heat, q, depends on:

• The mass of the sample (m)

• The change in temperature (DT)

• The nature of the sample (C)

• When something warms up:

The heat, q, depends on:

• The mass of the sample (m)

• The change in temperature (DT)

• The nature of the sample (C)

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

a) warm up. q=mCDT

b) melt

c) boil

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

q=mCDT

• q – heat, in Joules

• m –mass, in grams

• C –specific heat capacity, in J/goC

• DT—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 does it take to raise 50.g water from 15oC to 80.oC?

• q=mCDT

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

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

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

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

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

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

=14000 J (14 kJ)

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

• DT =q/mC

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

• DT =q/mC= 1550 J / (12 g x 4.18 J/goC )

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

• DT =q/mC= 1550 J / (12 g x 4.18 J/goC )

= 31oC

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

• DT =q/mC= 1550 J / (12 g x 4.18 J/goC )

= 31oC

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

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

• DT =q/mC= 1550 J / (12 g x 4.18 J/goC )

= 31oC

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

• --the measurement of heat.

• --the measurement of heat.

• If one thing gains heat…

• --the measurement of heat.

• 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?

• How much heat did the water gain?

• How much heat did the water gain?

q=mCDT

Mass of water, in grams

Specific heat of water, 4.18 J/goC

Change in the temperature of water, in oC

• How much heat did the metal lose?

• How much heat did the metal lose?

• Heat lost = - heat gained

• qlost=-qgained

• What is the specific heat capacity of the metal?

• What is the specific heat capacity of the metal?

C=q/mDT

Heat lost by metal

Mass of metal, in grams

Change in the temperature of metal, in oC

Specific heat of metal, in J/goC

• 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?

• 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?

.74 J/goC

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

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

• (Step 1: Convert to moles)

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

• 150 g x 1mol/18.02 g =

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

• 150 g x 1mol/18.02 g = 8.3 mol

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

• 150 g x 1mol/18.02 g = 8.3 mol

• (Step 2: Apply the heat of fusion)

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

• 150 g x 1mol/18.02 g = 8.3 mol

• Q=nHf=8.3 mol x 6.01 kJ/mol =

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

• 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)?

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

• (Step 1: Convert to moles)

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

• 250 g x 1mol/18.02 g =

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

• 250 g x 1mol/18.02 g = 13.9 mol

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

• 250 g x 1mol/18.02 g = 13.9 mol

• (Step 2: Apply the heat of vaporization)

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

• 250 g x 1mol/18.02 g = 13.9 mol

• Q=nHf=13.9 mol x 41 kJ/mol =

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

• 250 g x 1mol/18.02 g = 13.9 mol

• Q=nHf=13.9 mol x 41 kJ/mol = 570 kJ

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

• Hfus6.0 kJ/mol

• Hvap41 kJ/mol,

• MP -95.4oC

• BP 56.3oC

• C .126 J/goC

• Hfus5.7 kJ/mol

• Hvap29 kJ/mol

• MP 1085oC

• BP 2567oC

• C .385 J/goC

• Hfus13 kJ/mol

• Hvap231 kJ/mol

• MP and BP of molecules of similar size

• Formula FM(g/mol) MP (oC) BP (oC)

• CH4

• NH3

• H2O

• HF

• Ne

• MP and BP of molecules of similar size

• Formula FM(g/mol) MP (oC) BP (oC)

• CH4 16

• NH3 17

• H2O 18  = 18 g/mol ±11%

• HF 20

• Ne 20

• MP and BP of molecules of similar size

• Formula FM(g/mol) MP (oC) BP (oC)

• CH4 16 -183

• NH3 17 -78

• H2O 18 0

• HF 20 -83

• Ne 20 -249

• MP and BP of molecules of similar size

• Formula FM(g/mol) MP (oC) BP (oC)

• CH4 16 -183 -164

• NH3 17 -78 -33

• H2O 18 0 100

• HF 20 -83 20

• Ne 20 -249 -246

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

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

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

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

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

• 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.

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

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

• Some forces are stronger than others.

Showdispersion forces

• very weak

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

• They unbalance and attract their neighbors.

Show dipole interactions

• fairly weak.

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

• They attract their polar neighbors.

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

d+ d-

H Cl

--the hydrogen is more positive

--it is losing custody of its last electron

Show hydrogen bonding

• strongest of the weak bonds.

• permanent, larger charge imbalances than other polar bonds.

• They attract their polar neighbors better.

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.)

• 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

• 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.

Q: What holds water together?

Q: What holds water together?

HA! It’s a trick question!

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

!

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

• Formula Type of substance

• CH4

• NH3

• H2O

• HF

• Ne

• 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

• Formula Type of IM Forces

• CH4 dispersion forces

• NH3 hydrogen bonding

• H2O hydrogen bonding

• HF hydrogen bonding

• Ne dispersion forces

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

• NaCl: ionic compound

HCl: polar covalent molecule

• NaCl: ionic compound

HCl: polar covalent molecule

• NaCl: held together by ionic bonds

HCl molecules: attracted to each other by hydrogen bonds.

• NaCl: ionic compound

HCl: polar covalent molecule

• 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

• barium

• chlorine

• tin (II) chloride

• sulfur dioxide

• solid sulfur

• helium

• dinitrogen monoxide

• iron

• sodium oxide

• iodine

• barium sulfide

• sulfuric acid

• barium

• chlorine

• tin (II) chloride

• sulfur dioxide

• solid sulfur

• helium

• dinitrogen monoxide

• iron

• sodium oxide

• iodine

• barium sulfide

• sulfuric acid

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?

• 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 stronger the IM forces, the higher the:

MP, BP, Hfus, Hvap, C, surface tension, cohesion, viscosity, strength and hardness of the solid…

…etc. Usually.

• barium

• chlorine

• tin (II) chloride

• sulfur dioxide

• solid sulfur

• helium

• dinitrogen monoxide

• iron

• sodium oxide

• iodine

• barium sulfide

• sulfuric acid

• 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

• 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.

• 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

• Teacher Imitation Day Spring 2010

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

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

• --make a regular crystal

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

• --have high melting and boiling points

• --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.

• --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.

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

• Are generally soft, and easily melted or boiled.

Each molecule stands alone with its electrons

• Melting point Boiling point

• Heat of vaporization Heat of fusion

• Cohesion Surface tension

• Density Solubility

• Solution Solute

• Solvent Dissociation

• --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

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

• Aqueous (aq) = “dissolved in water”

Why not?

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

• They are notalike

Likedissolveslike

Non-polar solvents dissolve non-polar solutes

Polar solvents dissolve polar and ionic solutes

Metallic solvents dissolve metallic solutes.

• CH3OH/H20

• KBr/H2O

• CCl4/H2O

• S8/H2O

• Hg/H2O

• Ag/Hg

• S8/CCl4

• NaCl/KBr

• CH3OH/H20—polar/polar =YES

• KBr/H2O

• CCl4/H2O

• S8/H2O

• Hg/H2O

• Ag/Hg

• S8/CCl4

• NaCl/KBr

Like dissolves like

• CH3OH/H20 —polar/polar =YES

• KBr/H2O —ionic/polar =YES

• CCl4/H2O —nonpolar/polar =NO

• S8/H2O —nonpolar/polar =NO

• Hg/H2O —metallic/polar =NO

• Ag/Hg —metallic/metallic =YES

• S8/CCl4 —nonpolar/nonpolar=YES

• NaCl/KBr —ionic/ionic(ifmelted)=YES

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

• 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 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

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

• Suspensions settle out eventually.