WHAT CHANGES OCCUR DURING A REACTION?

Kinetics of Liquids Molecules of a cold sample of liquid

have lower kinetic energy than those in a warmer sample

If a particle near the surface has enough kinetic energy, it can overcome the attractive forces in a liquid and escape into the gaseous state

Known as a phase change

What properties do liquids have?

Viscosity:The friction or resistance to motion that

exists between the molecules of a liquid when they move past one another

The stronger the attraction between the molecules in a liquid, the greater the resistance to flow

Liquids with large intermolecular forces tend to be highly viscous

Surface Tension: The resistance of a liquid to an increase in its

surface area Which liquids will have high surface tensions and

why?

Because of decreased volume and increased molecular interaction, liquids expand and contract only very slightly with temperature change

Boiling Point: The point at which the liquid’s vapor pressure is

equal to the atmospheric pressure Rapidly converting from liquid to the vapor

phase within the liquid as well as at the surface

Capillary Action The attraction of the surface of a liquid to the surface of

a solid Liquids will rise very high in a narrow tube if a strong

attraction exists between the liquid molecules and the molecules that make up the tubing

Pulls liquid up against force of gravity Concave meniscus Polar liquids exhibit capillary action The spontaneous rising of a liquid in a narrow tube, due

to: Cohesive forces – the intermolecular forces among the

molecules of the liquid Adhesive forces – the forces between the liquid and its

container Which of these are stronger for water? Adhesive

Vapor PressureEvaporation (vaporization) – a process by

which the molecules of a liquid can escape the liquid’s surface and form a gas

Endothermic processHeat of vaporization (enthalpy of

vaporization) – energy required to vaporize one mole of a liquid at a pressure of 1 atm

Symbol: Δhvap

Vapor Pressure Condensation – process by which vapor

molecules re-form a liquid

Phase Equilibrium Eventually, enough

vapor molecules are present so that the rate of condensation equals the rate of evaporation

The system is said to be at equilibrium

The pressure of the vapor present at equilibrium is called vapor pressure

Phase Equilibrium What will happen if the temperature is

increased? The number of liquid molecules will be

reduced The number of gaseous molecules will be

increased The rates of evaporation and

condensation will become equal again This illustrates what is known as ;

Le Châtelier’s Principle

Le Châtelier’s Principle Reversible reactions – conversion of reactants to products

and vice versa occur simultaneously Change in conditions is imposed on a system at equilibrium,

the equilibrium will shift in the direction that tends to reduce that change in conditions.

CHANGES IN CONCENTRATION: Substance is added reaction consumes added substance Substance is removed reaction shifts to produce more

 2NO2 (g) N2O4 (g)

  Which direction will the above reaction shift if we add

NO2? Which direction will the above reaction shift if we

remove N2O4?

CHANGES IN PRESSURE: Increase: shift in direction that produces fewer molecules (moles) of

gas. Decrease: shifts in direction that produces more molecules of gas.  In the reaction below, if we increase the pressure which

direction will the reaction shift? NH4Cl (s) NH3 (g) + HCl (g)

CHANGES IN TEMPERATURE:  EXOTHERMIC: Reaction gives off heat (product) ENDOTHERMIC: Reaction absorbs heat (reactant) Consider heat as a component of the reaction. 

H2 (g) + I2 (g) 2 HI (g) + Heat   If we raise the temperature (add heat) which way will the

reaction shift? If we want the reaction to go to the right do we add or remove

heat? 

What is the Haber Process?

During WWII, Allied forces blocked the Germans from acquiring sodium nitrates used for explosives from mines in Chile in hopes of shortening the war.

Fritz Haber used Le Chatelier’s Principle to come up with a new process of making ammonia.

N2(g) + 3H2(g) ―› 2NH3(g) + Heat

What do you think he did? Concentration: Remove the products

(NH3) Pressure: Increased the pressure Temperature: Kind of complicated;

reducing heat meant less pressure, but increasing heat would shift toward reactants therefore kept temperature moderate (500C)

The Solid State Can be classified into very broad

categories:1. Crystalline solids – highly regular

arrangement of components2. Amorphous solids – have considerable

disorder in their structure3. Polycrystalline solid – an aggregate of

a large number of small crystals in which the structure is regular but the crystals are arranged in random fashion

Crystalline Solid

Amorphous Solid

Polycrystalline Solid

Crystalline Solids Lattice structure – a 3D system of

points designating the positions of the components

Unit cell – the smallest portion of a crystal lattice that is repeated throughout the crystal

Crystalline Solids

Network Solids Atomic solid containing strong directional covalent

bonds Allotropes – forms of the same element that differ in

crystalline structure• Differ in properties because of differences in

structure• Example: Diamond is one allotrope of carbon in

which each carbon is covalently bonded to four other carbon atoms in a tetrahedral direction.

• Graphite is another allotrope of carbon, covalently bonded to form hexagonal sheets

• What is a buckyball?

Allotropes of Carbon

Changes of State Melting point – the temperature at which

atomic or molecular vibrations of a solid become so great that the particles break free from their fixed positions and start to slide past each other in a liquid state

Heating curve – a plot of temperature versus time for a substance where energy is added at a constant rate

Sublimation – when a solid goes directly to a gaseous state without passing through the liquid phase

Heating Curve for Water

Terms Heat of fusion(ΔHfus) – the amount of

energy required at the melting point temperature to cause the change of phase to occur

Heat of vaporization (ΔHvap) – the amount of heat needed to vaporize 1 gram of a liquid at constant temperature and pressure

Phase Diagrams Way to represent the phases of a substance as

a function of temperature and pressure Triple point – the point at which all three

states of a substance are present Critical temperature – the temperature

above which the vapor cannot be liquefied no matter what pressure is applied

Critical pressure – pressure required to produce liquefication at the critical temperature

Together, the critical temperature and critical pressure define the critical point

Phase Diagram for Water

http://www.teamonslaught.fsnet.co.uk/co2%20phase%20diagram.GIF

Phase Diagram for carbon dioxide

Intermolecular Forces Both solids and liquids are condensed states of

matter Relatively weak forces which occur between

molecules Both dipole-dipole and London dispersion forces

are known as Van der Waals forces*It is important to recognize that when a

substance such as water changes from solid to liquid to gas, the molecules remain intact. The changes in state are due to changes in the forces among the molecules rather than within the molecules*

Dipole-dipole Forces•The attractive force resulting when polar molecules line up so that the positive and negative ends are close to each other

•Try to maximize the (+)----(-) interactions

•In the gas phase, these forces are unimportant•Weaker than ionic or covalent bonds

London Dispersion Forces•Forces which exist among all covalent molecules but is the only force for nonpolar molecules.

•Weak attractive forces between molecules resulting from the small, instantaneous dipoles that occur because of the varying positions of the electron during their motion about nuclei.

Hydrogen Bonding•Unusually strong dipole-dipole attractions that occur among polar molecules in which hydrogen is bonded to a highly electronegative atom such as O-H, N-H, F-H

Physical Properties Nonpolar tetrahedral

hydrides show a steady increase in boiling point

Polar tetrahedral hydrides, the lightest member has an unexpectedly high boiling point

This is due to hydrogen bonding that exist among the smallest molecule with the most polar X—H bond.

Which Intermolecular forces are strongest?

Bonds are WAY stronger than forces Ionic>Ionic/Dipole>H>Dipole/Dipole>

Dipole/Induced> Induced/Induced The stronger the intermolecular forces

the higher the melting and boiling points Solids have highest intermolecular forces

followed by liquids and gases.

Laws of Thermodynamics1st Law of Thermodynamics (AKA Law of conservation

of Energy): Energy cannot be created or destroyed. It remains constant in the universe.

E = q + wE = change in system’s internal energyq = heatw = work

Heat: energy that flows into or out of a system because of difference in temperature between the system and its surrounding

Understanding Heats of Reaction

Thermodynamics: Science of the relationships between heat and other forms of energy

Thermochemistry: Study of heat absorbed or given off by chemical reactions.

Energy: Ability to do work (measured in Joules)

Work = Force (N) x Distance (m) Types of energy:

Kinetic =1/2 mv2

Potential = mgh

What is the Heat of Reaction?

Heat of Reaction (q)Before any reaction the system and its

surrounding are at the same temperature.

When the reaction starts, the temperature changes.

The value of q needed to return the system to the given temperature at the completion of the reaction is known as the heat of reaction.

Types of Reactions Exothermic Reaction (-q): Heat is given

off. Products contain less energy than reactants.

Endothermic Reaction(+q): Heat is absorbed. Reactants have less energy than products.

Measuring Heat Flow Measured in joules and calories Heat capacity – the amount of heat

needed to raise the temperature of an object exactly 1˚ C

q = C ∆ t Depends on mass and chemical make-up Specific heat (s)– the amount of heat it

takes to raise the temperature of 1g of the substance 1˚C

q = s x m x ∆t

Measuring Heat Flow Can be measured using a calorimeter The heat released by the system is equal

to the heat absorbed by the surrounding and vice versa

What is Enthalpy of Reaction?

Enthalpy (H): Extensive (meaning it depends on the amount of substance) property of a substance that can be used to obtain the heat absorbed or evolved in a chemical reaction.

All chemical reactions absorb or give off heat.

This change in energy is known as the change in enthalpy (heat content) of a system.

ΔH = H (products) – H (reactants)

The Second Law of Thermodynamics

The entropy (S) of the universe increases for any spontaneous process

ΔS universe = ΔS system + ΔS surroundings

Entropy: A measure of the degree of disorder Reactions are driven by the need for a

greater degree of disorder and the drive towards the lowest heat content

Reactions with negative ΔH’s are exothermic and those with positive ΔS’s are proceeding to greater disorder

Entropy Increases: When a gas is formed from a solid

CaCO3(s) CaO(s) + CO2(g) When a gas is evolved from a solution

Zn(s) + 2H+ H2(g) + Zn2+(aq) When the number of moles of gaseous

product exceeds the moles of gaseous reactant

2C2H6(g) + 7O2 4CO2(g) + 6H2O(g) When crystals dissolve in water

NaCl(s) Na+(aq) + Cl-(aq)

Changes in Enthalpy ΔH for an endothermic reaction is positive ΔH for an exothermic reaction is negative Changes in enthalpy are independent of

the path taken to change a system from the initial to final state

Heat absorbed or given off varies with the temperature

Standard enthalpies of formation are given at 25°C and 1 atm pressure (ΔH0)

Endothermic Reaction

Exothermic Reaction

Changes in Enthalpy Standard enthalpy of formation –

the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states at 25°C

These values are known

Calculating Enthalpy of a Reaction

Example: How much heat is liberated when

10.0 grams of CH4 (g) reacts with excess O2(g)?

CH4(g) + 2O2(g) → CO2 (g) + 2 H2O(l): ∆H = -890.3 kJ

Convert grams CH4 → moles of CH4 → kilojoules of heat

Changes in Free Energy of a System

Free energy – energy available to do work

The Gibb’s Free Energy Equation: ΔG = ΔH –TΔS The sign of ΔG can be used to

predict the spontaneity of a reaction

How factors in the equation affect ΔG : ΔG = ΔH –TΔS

ΔH ΔS ΔG Will it happen

Comment

Exothermic (-)

+ Always negative

Yes No exceptions

Exothermic (-)

- At lower temperatures

Probably At low temperature

Endothermic (+)

+ At higher temperatures

Probably At high temperatures

Endothermic (+)

- Never No No exceptions

Hess’s Law If a series of reactions are added

together, the enthalpy for the total reaction is the sum of the enthalpy changes for the individual steps

What is the ∆Hf of the following reaction?2C(graphite) + O2(g)→ 2CO(g)

2C(graphite) + 2 O2(g)→ 2 CO2(g) 2 CO2(g)→ 2CO (g) + O2(g)

Use the table to find ∆Hf 2C(graphite) + 2 O2(g)→ 2 CO2(g)

2(0) + 2(0) 2(-393.5 kJ)

(-787.0kJ) – 0 = -787.0 kJ

2 CO2(g)→ 2CO (g) + O2(g) 2 (-393.5kJ) 2 (-110.5)

+ 0 (-221.0kJ) – (-787.0kJ) = 566.0

kJ

∆Hf = -787.0 kJ + 566.0 kJ = -221.0 kJ