Chapter 10: States of Matter

Concept Base: Chapter 1: Properties of Matter Chapter 2: Density Chapter 6: Covalent and Ionic Bonding

Pressure

FP =

A

standard pressure – the pressure exerted at sea level in dry air

760. mmHg 29.9 inHg 760. torr 101.325 kPa 1.01325 x 105 N/m2 (Pa) 1.00 atm 14.7 psi

Pressure

A column of air 1.00 m2 in cross-sectional area extending from the earth’s surface through the upper atmosphere has a mass of about 10,300 kg (22700 lb), producing an atmospheric pressure of approximately 101,000 Pa.

What do you think?

Which city has higher atmospheric pressure, Pittsburgh or Denver?

atoms combine to form

ionic bonds covalent bonds

(M + NM) (NM + NM)

chemical bond – a mutual electrical attraction between the nuclei and valence electrons of two atoms that binds the atoms together

Covalent Vs. Ionic Bonding

ionic bond – when electrons are taken by one atom from another atom

metal and a nonmetal

NaCl

cation and anion

Ionic Bonding

Covalent Bonding

covalent bond – when electrons are shared between two atoms

two nonmetals

No ions formed! (no electrons are taken)

H-H O

H H

.. ..

Pure

Covalent

Ionic

There is another type of bond, not purely covalent and not purely ionic.

Polar

Covalent

Nonpolar

Covalent

The Kinetic-Molecular Theory of Matter

In reality, all atoms are moving: vibrating, rotating, translating.

The only time something would NOT be moving is at ABSOLUTE ZERO,

(0 Kelvin or -273 oC).

The Kinetic-Molecular Theory of Gases

ideal gas – hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory.

real gas – a gas that does not fit all the assumptions of the kinetic-molecular theory.

Although an ideal gas does not exist, many gases behave nearly ideally if pressure is not very high and temperature is not very low.

The Kinetic-Molecular Theory of Gases

1. Gases consist of large numbers of tiny particles that are far apart relative to their size

• Gases take up much more space than solids or liquids

• Gases are easily compressed

The Kinetic-Molecular Theory of Gases

2. Collisions between gas particles and between particles and container walls are elastic collisions.

elastic collision –

a collision with no

loss of energy

inelastic collision –

a collision with some

loss of energy

The Kinetic-Molecular Theory of Gases

3. Gas particles are in continuous, rapid, straight-line, random motion. They therefore possess kinetic energy, which is energy of motion.

• The KE that the molecules have as a gas overcomes any attractive forces that they might have.

The Kinetic-Molecular Theory of Gases

4. There are no forces of attraction between gas particles.

The Kinetic-Molecular Theory of Gases

5. The temperature of a gas depends on the average kinetic energy (energy of motion) of the particles of the gas

KelvinKE T

The Nature of Gases

Gases do not have a definite shape or volume

They completely fill any container because they have no attraction for the other molecules in the container (unlike liquids and solids).

The Nature of Gases

fluid – a nonsolid state of matter in which the atoms or molecules are free to move past each other, as in a gas or liquid (flows)

Attractive forces are insignificant, thus gas particles glide easily past one another and are called fluids.

The Nature of Gases

Gases have low densities.

MD =

V

The Nature of Gases

Gases can be compressed

Thus, many more molecules can be inside a container under great amounts of pressure.

The Nature of Gases

Gases diffuse readily into one another.

Flasks are NOT connected. Valve is opened to connect the flasks.

The Nature of Gases

diffusion – the spontaneous movement of particles of gas caused by random motion

effusion – process by which gas particles pass through a tiny opening due to pressure being exerted upon them

KE = ½mv2

Gas molecules move at different speeds or velocities, depending on

•the temperature at which the molecules are

•the molar mass of the molecules

The higher the MM, the slower they move.

Check for Understanding

What kind of states of matter can be poured?

Which molecules are moving faster, water at 50oC or water at 20oC?

Which state(s) of matter can be compressed to great extents?

Intermolecular Forces pages 189-193 – Chapter 6

intermolecular forces - attractive forces between molecules

These forces vary in strength, however are generally much weaker than covalent bonds (not intermolecular).

The stronger the intermolecular force the closer the molecules get to one another, thus perhaps creating a solid versus a liquid, etc.

What do you think?

Now that we learned about intermolecular forces, what do you think an intramolecular force is? Give two examples.

Intramolecular Forces > Intermolecular Forces (stronger than)

Dipole-Dipole Forces

dipole – a molecule or a part of a molecule that contains both partial positive and partial negative regions

A dipole is created when there is a large difference in electronegativity.

Dipole-Dipole Forces

The partial positive end is attracted to the partial negative end.

Dipole-Dipole Forces

Compare ICl to Br2 because they have approximately the same molar mass

ICl BP = 97oC

Br2 BP = 59oC

What makes ICl have a higher BP?

Dipole-Induced Dipole Attraction

Ion-Induced Dipole Attraction

What do you think?

Looking at the previous intermolecular forces, do you think it would be possible to have an ion-dipole attraction? Explain.

Hydrogen Bonds

F-H (HF) O-H (H2O) N-H (NH3)

All of these bonds have a large difference in electronegativity, thus creating a large dipole, or a highly polar bond.

These highly polar bonds have a very strong attraction.

These very strong attractions are called hydrogen bonds.

Hydrogen Bonds

Hydrogen Bonding in Water

Hydrogen Bonding in Water

Because of the hydrogen bonding in water, an open, rigid structure is formed when freezing.

As a solid, there is more hydrogen bonding than as a liquid.

Dice at 0oC = 0.917 g/mL

Dwater at 0oC = 1.000 g/mL

Hydrogen Bonds

Hydrogen Bonding in Acetic Acid

HC2H3O2

Hydrogen Bonds

Compare H2O to H2S

H2O BP = 100oC

H2S BP = -61oC

Compare NH3 to PH3

NH3 BP = -33oC

PH3 BP = -88oC

What makes H2O and NH3 have higher

boiling points?

Hydrogen Bonds

Snowflakes are large ice crystals that have a unique shape. The shape reflects the rigid position of the hydrogen bonding of the solid.

London Dispersion Forces

Nonpolar molecules will also exhibit a weak attraction for one another.

The constant motion of electrons within a molecule can create a temporary dipole, or London dispersion force, that attracts to another temporary dipole.

London Dispersion Forces

London forces are dependent upon the motion of electrons. Therefore, the more electrons, the greater the London forces.

molar mass London forces

Check for Understanding

Compare Cl2, Br2, and I2 and arrange them according to the strength of their London forces.

What are their states of matter at room temperature? WHY?

Intermolecular Forces

Strength of Intermolecular Forces

from Highest to Lowest

hydrogen bonding

dipole-dipole attraction

dipole-induced dipole attraction

London dispersion forces

van der Waals forces - any dipole forces and London forces are groups as these vdW forces

Check for Understanding

What physical property directly correlates with the strength of the London dispersion forces?

What types of molecules have the strongest intermolecular forces?

What do you think accounts for NH3, ammonia, having a boiling point 130oC higher than CH4, methane?

Comparison of Boiling Points

Liquids

How do you know that this

contains a liquid without opening it?

Liquids

Liquid nitrogen changing to

gaseous nitrogen

Liquids

LIQUIDS…

have a definite volume and take on the shape of their container (unlike gases)

have a high density

are not compressed well (brake fluid)

diffuse (like food coloring in water)

have surface tension

are fluids (but fluid liquid)

Liquids

surface tension – a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size.

Surface tension results from attractive forces between the particles in the liquid. The stronger the attractive force, the higher the surface tension.

Liquids

Notice that the “pull” on the mercury atoms at the top is not symmetrical. That is what gives the characteristic spherical shape to drops of liquid.

Liquids

viscosity – the resistance of a liquid to flow

high viscosity = “thick” liquid

low viscosity = “thin” liquid

Liquids with stronger intermolecular forces have higher viscosity.

An increase in temperature will decrease the viscosity.

KelvinKE T

Liquids

volatile liquid - a liquid that evaporates readily at low temperatures

The higher the volatility of a liquid, the weaker the intermolecular forces of attraction between their particles.

An increase in T will increase evaporation.

KelvinKE T

Liquids

cohesive forces

forces of attraction between like molecules (H2O to H2O)

adhesive forces

forces of attraction between unlike molecules (H2O to glass)

In each cylinder, can you

describe which forces are

greater: cohesive or

adhesive?

Liquids

capillary action - the attraction of the surface of a liquid to the surface of a solid (adhesive forces)

Many liquids will “creep” along a solid, like water does to paper or cloth fibers until the pull of gravity is too much for it to overcome.

Check for Understanding

Order the following liquids from highest surface tension to lowest? Explain.

Br2, H2O, H2S

Which has a higher viscosity, water or carbon tetrachloride, CCl4?

How is this water- strider able to walk on water?

Evaporative Cooling

Why do you feel cold when you get out of the shower?

Why does fanning yourself cool you down?

Why do you sweat when you have a fever?

Solids

Solids are in a relatively fixed position.

Solids have only vibrational movements around fixed points.

Solids have definite shape and volume.

Solids are almost incompressible.

Solids do not diffuse (practically).

Solids

Solids are either crystalline or amorphous.

crystalline – consist of crystals, particles arranged in an orderly, geometric repeating pattern

amorphous – Greek for “without shape”; consist of particles, randomly arranged

Crystalline Solids

Crystalline solids break into orderly

pieces. After

breaking salt, the

cubic structure is still visible.

NaCl is cubic.

Crystalline Solids

crystal structure – three-dimensional arrangement of particles of a crystal, represented by a lattice

unit cell – The smallest portion of a crystal lattice that shows the 3-D pattern Be familiar with the

seven basic crystalline systems: Figure 11 - p.339

Crystalline Systems

Amorphous Solids

Most plastics are

amorphous.

Amorphous solids break into random

pieces.

They usually shatter into

irregular shapes.

Amorphous Solids

The freezing point of amorphous solids can vary according to how slowly the material cools. (Ex: butter)

http://math.ucr.edu/home/baez/physics/General/Glass/glass.html

Crystal Types

ionic

covalent

covalent network – each atom is covalently bonding to its “neighbors”

covalent molecular – each molecule is held together by intermolecular forces

metallic – metallic bonding

Ionic Crystals

ionic bonding

high melting points

positive and negative ions

Ex: NaCl, MgF2

Covalent Network Crystals

Diamond (C) and SiO2 are examples of solids that form

these “giant

molecules”.

Held together by covalent bonds.

Covalent Network Crystals

Covalent Molecular Crystals

Water is an example of covalent molecular

crystals due to its

hydrogen bonding.

Held together by intermolecular forces.

Metallic Crystals

metallic bonding – the chemical bonding that results from the attraction between metal atoms and the surrounding “sea of electrons”

Metallic bonding allows metals to be good conductors of electricity, malleable, and ductile.

Ex: Hg, Cu, Fe, W

Checking for Understanding

Name at least two common examples of amorphous substances.

What type of crystal would ammonia, NH3, as a solid be likely classified as?

Why do you think ionic crystals have such high melting points, thus are usually found as solids at room temperature?

Checking for Understanding

What is glass: crystalline or amorphous? Support your answer.

The Kinetic Theory of Heat and Temperature

When a phase change is occurring, the temperature does not change, only the position of the particles. (PE)

When something is being heated and it is not changing phase, the temperature will rise. (KE)

Potential Energy Differences PE – “energy of position”

PE vs. KE

Only differences in kinetic energy are reflected by temperature differences.

Difference in potential energy are NOT reflected in temperature differences.

THERE IS NO TEMPERATURE CHANGE DURING A PHASE

CHANGE.

Heating Curve for Water

Phase Changes During solidification or melting

H = Kfm

Kf = heat of fusion, the amount of heat needed to melt/freeze 1g of a substance

• During boiling or condensing

H = Kvm

Kv = heat of vaporization, the amount of heat needed to boil/condense 1g of a substance

Phase Changes

Kf for water = 333 J/g

Kv for water = 2260 J/g

Heating Curve for Water

Frost represents Deposition

Vapor Pressure

vapor pressure – the pressure due to a vapor above a liquid

How do you increase the vapor pressure?

How does the vapor get above the liquid?

Open System:

Closed System

Pressure

How can you increase the pressure, P, of a gas inside a container?

increase the temperature, T (KE) of the gas

decrease the size (volume, V) of the container

add more molecules to the container (increase the number of moles, n)

Dynamic equilibrium

occurs when the rate of

evaporation = rate of

condensation.

Equilibrium

dynamic equilibrium – Although there are two opposing processes going on, they are occurring at the same rate.

liquid

vapor

solid

vapor

Vapor Pressure

Boiling

boiling – the conversion of a liquid to a vapor not only at its surface, but within the liquid as well

boiling point – The temperature at which the vapor pressure is equal to the external pressure (usually atmospheric pressure)

normal boiling point – The temperature at which the vapor pressure is equal to standard pressure (1.00 atm)

Pvapor = Pexternal

Boiling

When Pvapor Patmosphere, the liquid will boil.

What do you think?

How is it possible that this beaker of

water is boiling at room

temperature?

vacuum pump

Phase Diagrams

TRIPLE POINT

CRITICAL POINT

CRITICAL TEMPERATURE (no more liquids

above this)

CRITICAL PRESSURE (the lowest pressure at which the substance can still be a liquid at the critical T)

Phase Diagram for Water

Carbon Phase

Diagram

What do you think?

Why is it that a closed bottle of water, of which part has been consumed, has condensation all over the inside of the container after time?

What do you think?

If you can boil water in a

vacuum pump, can you hard boil

an egg in the boiling water in

the vacuum pump?

What do you think?

Why is it that the snow in your driveway melts when you back over it with the car?