Chapter 1 PHY 351

transcript

CHAPTER 1: GAS AND CONDENSED MATTER

PREPARED BY: NOR HALIZA YAAKOB,

Faculty of Applied Science Universiti Teknologi MARA Campus of Negeri Sembilan 72000 Kuala Pilah, Negeri

Sembilan, MALAYSIA. PHY 351 (Materials Science)

013-9495135 norhaliza@ns.uitm.edu.my

1.1: Force between particle

1.2: Ideal Law, Real Gas, Van der Waals

of State Equation

1.3: Condensed Matter-solid and liquid

Triple Point

SCOPE OF STUDY

INTRODUCTION

Materials may be defined as substance of which something is composed or made.

Materials Science is a scientific discipline that is primarily concerned with the search for basic knowledge about the internal structure, properties and processing of materials.

Material engineering is an engineering disciple that is primarily concerned with the use of fundamental and applied knowledge of materials, so that they can be converted into products needed or desired by society.

Gas and Condensed matter physics: The branch of physics dealing with the physical properties and to understand the behaviour of condensed phases of matter such as gas, liquids and solids.

The gaseous state of matter is found between the liquid and plasma states .

Liquids and solids are the most well known forms of condensed matter.

The atoms in condensed matter are closer together and more closely bounded together than in a gas, as a result condensed matter tends to be some form of liquid or solid.

All matter is held together by force.

Particle physicists think of forces as interactions between particles that produce structure.

The force between atoms within a molecule is a chemical or intramolecular force.

The force between molecules are a physical or intermolecular force.

These physical forces are what we overcome when a chemical changes its state (e.g. gas liquid).

FORCE BETWEEN PARTICLE

Bonding model for covalent molecular substances

Bonding for covalent molecular substances falls into two categories i. The strong forces of attraction which holds atoms together

within molecules (Intramolecular).

ii. The weak forces of attraction between molecules (Intermolecular).

What would intramolecular forces be?

Forces within molecules e.g covalent, metallic or ionic.

intra means within.

Intramolecular bonds are stronger than intermolecular forces.

Intramolecular Forces

Forces that occur between molecules.

Intermolecular Forces

What causes intermolecular forces?

Molecules are made up of charged particles: positive nuclei and negative electrons.

When one molecule approaches another there is a multitude of forces between the particles in the two molecules.

Each electron in one molecule is attracted to the nuclei in the other molecule but also repelled by the electrons in the other molecule.

The same applies for nuclei.

Types of Intermolecular forces

The three main types of intermolecular forces are:

i. Dipole-dipole attraction - occur only btw polar molecules.

ii. H bonding – only with Hydrogen and Oxygen, Fluorine and Nitrogen).

iii. Dispersion forces (London Dispersion Forces).

Dipole moment – molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.

Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.

Only about 1% as strong as covalent or ionic bonds.

Dipole–Dipole Attraction

Hydrogen Bonding

The electromagnetic attractive interaction between polar molecules in which Hydrogen is bound to a highly electronegative atom – nitrogen, oxygen, or fluorine.

Strong dipole-dipole forces.

Hydrogen Bonding in Water

Blue dotted lines are the intermolecular forces between the water molecules.

Hydrogen Bonding

Affects physical properties:

i. Boiling point

London Dispersion Forces

The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary (Instantaneous) dipoles.

Significant in large atoms/molecules.

Occurs in all molecules, including nonpolar ones.

Nonpolar Molecules

Become stronger as the sizes of atoms or molecules increase.

Strength of Intermolecular Interactions

Hydrogen Bonding

Dipole – Dipole

Melting and Boiling Points

In general, the stronger the intermolecular forces, the higher the melting and boiling points.

The kinetic theory of matter is based on the following postulates:

1. The particles are in constant random motion- They

possess kinetic energy due to their motion

2. There is no interaction between molecules, molecules obey laws of classical mechanics and interact only when colliding.(There are repulsive and attractive forces (PE) between particles but a weak force).

3. No energy is lost when the particles collide, called

elastic collision.

4. Average particle speed increases with temperature.

5. The molecules are separated by great distances relative to their size.

Kinetic Molecular Theory

The kinetic energy of a particle is given by

the equation:

Where:

m = particle mass in kg

v = particle velocity in m/s

KE = kgm2/s2 = J (joule)

According to postulate 4 of our kinetic theory particle velocity increases with temperature. This means as temperature increases then kinetic energy increases.

Kinetic Energy

Potential energy is the sum of the attractive and repulsive forces between particles.

Alternatively we can say forces between particles may be either cohesive or disruptive.

Cohesive forces include dipole-dipole interactions, dispersion forces, attraction between oppositely charged ions. Cohesive forces are largely temperature independent.

Potential Energy

Disruptive forces are those forces that make particles move away from each other.

These forces result predominately from the particle motion.

Disruptive forces increase with temperature in agreement with postulate 4.

We can conclude that as we increase the temperature particles will become further apart from each other.

The Gas Laws and Absolute Temperature

The relationship between the volume, pressure, temperature, and mass of a gas is called an equation

of state.

We will deal here with gases that are not too dense.

Boyle’s Law: the volume of a given amount of gas is inversely proportional

to the pressure as long as the temperature is constant.

The volume is linearly proportional to the temperature, as long as the temperature is somewhat above the condensation point and the pressure is constant:

Extrapolating, the volume becomes zero at −273.15°C; this temperature is called absolute zero.

The concept of absolute zero allows us to define a third temperature scale – the absolute, or Kelvin, scale.

This scale starts with 0 K at absolute zero, but otherwise is the same as the Celsius scale.

Therefore, the freezing point of water is 273.15 K, and the boiling point is 373.15 K.

Finally, when the volume is constant, the pressure is directly proportional to the temperature:

The Ideal Gas Law

We can combine the three relations just derived into a single relation:

What about the amount of gas present? If the temperature and pressure are constant, the volume is proportional to the amount of gas:

The Ideal Gas Law

A mole (mol) is defined as the number of grams of a substance that is numerically equal to the molecular mass of the substance:

1 mol H2 has a mass of 2 g

1 mol CO2 has a mass of 44 g

The number of moles in a certain mass of material:

• We can combine these into a general gas law:

The Ideal Gas Equation

), (constant 1

), (constant PnTV

),(constant TPnV

• Boyle’s Law:

• Charles’s Law:

• Avogadro’s Law:

R = gas constant, then

The ideal gas equation is:

Real Gases behave ideally at low P and high T.

The Ideal Gas Law

We can now write the ideal gas law:

P, pressure = Pa V, volume = m3

n, number of moles = moles R, universal gas constant = J mol-1 K-1

T, temperature = K

At sea level the atmospheric pressure is about ;

This is called one atmosphere (atm).

𝟏 𝒂𝒕𝒎 = 𝟏. 𝟎𝟏𝟑 𝒙 𝟏𝟎𝟓 𝑵/𝒎𝟐

1 Pascal (Pa) = 1 N/m2

Ideal Gas Law in Terms of Molecules: Avogadro’s Number

Since the gas constant is universal, the number of molecules in one mole is the same for all gases. That number is called Avogadro’s number:

The number of molecules in a gas is the number of moles times Avogadro’s number:

Therefore we can write:

where k is called Boltzmann’s constant.

𝑷𝑽 = 𝑵𝒌𝑻

Calculate the pressure exerted by 84.0 g of ammonia, NH3, in a 5.00 L container at 200. oC using the ideal gas law.

PV = nRT

P = nRT/V

n = 84.0g * 1mol/17 g

T = 200 + 273

P = (4.94mol)(0.08206 L atm mol-1 K-1)(473K)

P = 38.3 atm

Real Gases: Deviations from Ideality

Real gases behave ideally at ordinary temperatures and pressures.

At low temperatures and high pressures real gases do not behave ideally.

The reasons for the deviations from ideality are:

The molecules are very close to one another, thus their volume is important.

The molecular interactions also become important.

J. van der Waals, 1837-1923,

Professor of Physics,

Amsterdam. Nobel Prize 1910.

The dashed curve A’ and B’ represents the behavior of a gas as predicted by the ideal gas law (Boyle’s Law) for several different values of the temperature.

van der Waals’ equation accounts for the behavior of real gases at low temperatures and high pressures.

P + n a

VV nb nRT

The van der Waals constants a and b take into account two things (correction van der Waals made): 1) a accounts for intermolecular attraction

i. For nonpolar gases the attractive forces are London Forces

ii.For polar gases the attractive forces are dipole-dipole attractions or hydrogen bonds.

2) b accounts for volume of gas molecules

At large volumes a and b are relatively small and van der Waal’s equation reduces to ideal gas law at high temperatures and low pressures.

The van der Waals Equation

• General form of the van der Waals equation:

Real Gases: Deviations from Ideal Behavior

nRTnbVV

Corrects for

molecular

volume

Corrects for

molecular

attraction

The three states of matter.

Condensed matter

Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior

gas liquid solid

assumes the shape and volume of its container. particles can move past

one another

assumes the shape of the part of the

container which it occupies

particles can move/slide past one

another

retains a fixed volume and shape

rigid - particles locked into place

Compressible. lots of free space between particles

not easily compressible little free space

between particles

not easily compressible little free space

between particles

flows easily. particles can move past

one another

flows easily particles can

move/slide past one another

does not flow easily rigid - particles cannot

move/slide past one another

Clearly, a theory used to describe the condensed states of matter must include an attraction between the particles in the substance

• Condensed States of Matter:

– Liquids – Solids

Kinetic Theory Description of the Liquid State.

Like gases, the condensed states of matter can consist of atoms, ions, or molecules.

What separates the three states of matter is the proximity of the particles in the substance.

For the condensed states of matter the particles are close enough to interact.

Phase Changes

Triple Point Diagram of Water

Regions: Each region corresponds to one phase which is stable for any combination of P and T within its region

Lines Between Region: Lines separating the regions representing phase-transition curves

Triple Point: The triple point represents the P and T at which all 3 phases coexist in equilibrium

Critical Point: At the critical point the vapor pressure cannot be condensed to liquid no matter what pressure is applied.

Chapter 1 PHY 351

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