what changes occur during a reaction?
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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 - PowerPoint PPT PresentationTRANSCRIPT
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