States of Matter. I. Review: Phases of Matter A.Solid –Fixed volume and shape –Molecules are...

States of Matter

I. Review: Phases of Matter

A. Solid– Fixed volume and shape– Molecules are tightly packed and in a set position

B. Liquid– Fixed volume, indefinite shape– Molecules are tightly packed, but not in a set

position.

C. Gas– No fixed volume or shape– Molecules are widely dispersed.

II. Forces of Attraction

A. Intramolecular Forces– Forces that act within a molecule.

• Ex: – Covalent and ionic bonds


B. Intermolecular Forces

Short range forces between molecules


• 1. Hydrogen Bonding• A particularly Strong force• Occurs in compounds in which a hydrogen

atom is attached to Fluorine, Oxygen, or Nitrogen atom.

• The hydrogen atom will act like a positive ion

• The hydrogen atom will then be attracted to negative charge on a nearby molecule.


1. Hydrogen Bonding

• Hydrogen bonding increases melting and boiling points because more energy is required to break the forces between molecules.


2. Dipole-Dipole ForcesAn intermolecular force• Occurs in polar molecules (dipoles) because they have

uneven charge distribution (a positive end and a negative end)

• A positive end of one molecule is attracted to a negative end of a nearby molecule.

• Dipole – dipole forces will also increase melting and boiling points.

• A dipole can also temporarily attract electrons from another molecule causing an temporary dipole.


3. London Dispersion Forces• A weaker intermolecular force (a type of induced

dipole).• Occurs in nonpolar molecules and noble gases.• Electrons are constantly moving, therefore

electron density changes.• This effect increases with increasing number of

electrons. F2 Gas

Br2 LiquidI2 Solid

III. Liquids and Solids

A. Liquids1. Density and Compression

a. density similar to that of solidsb. Not very compressible

2. Fluidity a. Weak intermolecular forces allow molecules to flow past each other.

3. Viscosity a. The resistance a material has to flowing.b. Increases with increasing intermolecular forces.

III. Liquids and Solids

B. Solids1. Density of Solids

a. Usually higher than liquids and solids for the same substance.2. Crystalline Solids

a. Solids consisting of ionically bound ions. 3. Network Covalent Solids

a. Solids consisting of covalently bound atoms. 4. Metallic Solids

a. Metal atoms bound with metallic bonds.5. Amorphous Solid

a. A solid with no molecular pattern.

• Crystalline solid

• Network Covalent solids

Metallic Solid

Amorphous solid

IV. Gases and the Kinetic Molecular Theory

A. Kinetic Molecular Theory– Many of the properties of gases can be explained

by the KMT

B. Assumptions:1. Particle Size

Atoms of gases are so much smaller than the spaces between atoms, that their size is negligible.

2. Particle Motiona. Collisions

elastic – No energy is lost during a collision.

inelastic – some energy is lost during a collision.


b. Constant, rapid, random motion

c. No forces of attraction or repulsion between gas particles


3. Particle Energy– Kinetic Energy = (1/2)(Mass) (velocity)2

– Depends on upon two factors• Particle velocity • Particle mass


Temperature1. Definition: The average kinetic energy of the particle

of a substance.

2. Three scales:

a. Celsius• Based on the boiling and freezing points of water.

b. Kelvin• Sets to coldest theoretical temperature to “zero.”

• The difference between two temperatures is the same in both Celsius and Kelvin

c. Fahrenheit• Based on the human body temperature and the coldest

temperature that the creator could get water to freeze at.


3. Conversions • Tf = ((9/5)Tc) + 32

• Tc = (5/9)(Tf – 32)

• Tk = Tc + 273


4. Practice:a. 28o C = _____________K

b. 200 K = _____________ oC

c. – 15o C = ____________oF

d. 10 K = _______________oF


Physical Properties1. Density – Very Low2. Fluidity – Flows very easy3. Compression – Highly compressible4. Expansion – Expands to fill container5. Diffusion – The movement from an area of

higher concentration to lower concentration.6. Effusion – The movement from an area of

higher concentration to lower concentration through an opening.


E. Graham’s Law of Diffusion/Effusion

Essentially, small molecules move faster

than larger ones.


• Sample: Which gas will diffuse the fastest: ammonia, NH3, or hydrogen chloride, HCl?


F. Gas Pressure1. Definition of Pressure: The force of

molecules on their container

2. Mathematical Formula: P = (Force/Area)

3. Air Pressure Conversions-Conversion Factors

760 mmHg (torr) = 1 atm = 101.3 kPa = 30.0 in Hg = 1 torr = .133 kPa


Sample: Convert 0.70 atm to mmHg.

Practice:

Convert 2.0 atm to torr (mmHg).

Convert 725 torr to atm.


4. Atmospheric (air) Pressure– The pressure that the air around us places on

everything

5. Measuring Atmospheric Pressure

a. Hg Barometer = glass tube sealed at one end with the other end immersed in a container of Hg

A typical Mercury Barometer


5. Measuring Pressure of Gasesa. Closed Manometer

b. Open Manometer


• 5. Standard Pressure

• 6. Air Pressure Varies with Altitude

• Dalton’s Law of Partial Pressures

• Ptotal = P1 + P2 + P3 …….


• Practice:

• 1. Determine the total pressure, in mm Hg, for a mixture that contains four gases with partial pressures of 5.0 atm, 4.56 atm, 3.02 atm, and 1.2 atm.


• 2. What is the partial pressure of hydrogen gas in a mixture of hydrogen and helium if the total pressure is 600 mmHg and the partial pressure of helium is 439 mmHg?


• V. Phase Changes

• Phase changes that require energy (endothermic)– Melting– Vaporization– Sublimation


• Phase changes that release energy (exothermic)

– Condensation– Deposition– Freezing


• Heating Curve

States of Matter. I. Review: Phases of Matter A.Solid –Fixed volume and shape –Molecules are...

Documents

Transcript of States of Matter. I. Review: Phases of Matter A.Solid –Fixed volume and shape –Molecules are...