Chapter 5/ 15 Flashcards
Energetics/thermochemistry (92 cards)
1
Q
Heat
A
The transfer of energy between objects of different temperature
- heat will spontaneous flow from an object of higher temp. to an object of lower temp.
- once two objects reach same temp., which is known as thermal equilibrium, no more energy will be transferred
2
Q
Enthalpy
A
Heat released or absorbed by a system at constant pressure
- is the heat content of a system
3
Q
Standard enthalpy of reaction
A
When the enthalpy change is measured under standard conditions
- pressure of 100 kPa
- temp. of 25 degrees Celsius (298K)
4
Q
Temperature
A
The average kinetic energy of the particles in a substance
- if temp. of a substance increases, average kinetic energy of particles is increasing
- absolute temp. is measured on Kelvin scale
5
Q
Kelvin scale
A
- absolute temp. (kelvin, K) is directly proportional to average kinetic energy of particles in a sample
- absolute 0 (0K is lowest temp. theoretically possible), temp. at which all particles have least amount of kinetic energy
6
Q
Kelvin - degrees Celsius equation
A
Temperature (K) = temperature (degrees Celsius) + 273.15
7
Q
System and surroundings
A
System= chemical reaction Surroundings= rest of the universe
8
Q
Open systems
A
Where matter and energy can move freely between system and surroundings
9
Q
Closed systems
A
Only energy is able to move between system and surroundings
10
Q
Relationship between system and surroundings
A
- energy that is lost from system is transferred to surroundings
- any energy gained by the system is transferred from surroundings to system
11
Q
Isolated system
A
- doesn’t exchange energy or matter with its surroundings
12
Q
Negative change in enthalpy
A
When heat is released from the system to the surroundings, enthalpy of the system decreases
- change in enthalpy is negative
- only exothermic reactions
13
Q
Positive change in enthalpy
A
When heat is transferred to a system from surroundings, enthalpy of system increases
- change in enthalpy is positive
- only endothermic reactions
14
Q
Law of conservation of energy
A
Energy cannot be created or destroyed; it’s converted from one form to another
- hence, total amount of energy in universe is constant
15
Q
Chemical reaction
A
Involves a transfer of heat between a system and its surroundings
- in majority of chemical reactions, heat is released from system to its surroundings (exothermic reactions)
16
Q
Exothermic reactions
A
- heat is released from system to its surroundings
- in aqueous solutions, cause an increase in temp. of reaction mixture; it will feel hotter
- warms the mixture, heat is lost to surroundings
eg. combustion and neutralisation reactions
17
Q
Energy profile for an exothermic reaction
A
- enthalpy change is negative because heat has been transferred from system to surroundings
- system has lost heat
- products of reaction have lower enthalpy than reactants, so are more energetically stable
18
Q
Thermochemical equations
A
Show enthalpy change of a reaction
19
Q
Endothermic reactions
A
Chemical reactions in systems that absorb heat energy from their surroundings
- in aqueous solutions cause a decrease in temp. of reaction mixture
- it feels colder
- reaction mixture is part of surroundings (not the system); system absorbs heat from its surroundings in an endothermic reaction
- take in heat; cools the mixture at first, and then heat is gained from the surroundings
20
Q
Energy profile for an endothermic reaction
A
- enthalpy change is positive because reactants gain heat from their surroundings
- products have higher enthalpy than reactants, and are less energetically stable
21
Q
Coffee cup calorimeter
A
- used to determine enthalpy changes
- a polystyrene cup w/ a lid and a thermometer
- reaction is carried out in the polystyrene cup, and temp. change of reaction mixture is recorded using thermometer
- once min. and max. temp. of reaction mixture are known, enthalpy change of reaction can be calculated
22
Q
Advantages of simple polystyrene cup as a calorimeter
A
- polystyrene is a good heat insulator, it reduces heat loss to the surroundings
- expanded polystyrene cup absorbs very little heat itself, no need to calculate heat absorbed by calorimeter
23
Q
Calculating enthalpy changes
A
q = m x c x (change in temp.)
q = heat absorbed or released in in J
c = specific heat capacity in J/g/K
m = mass of solution in g
change in temp. = change in temp. in degrees Celsius or K
24
Q
Specific heat capacity (c)
A
Amount of heat required to raise the temp. of one gram of a substance by 1 degree Celsius or 1 Kelvin
Units: J/g/degrees Celsius or J/g/K
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Specific heat capacity of certain substances
- metals
- water
Metals:
- lower specific heat capacity
- heat up quickly, but also lose heat quickly
Water:
- high specific heat capacity
- takes a lot of heat energy to increase temp. of water
- also retains that heat for a longer period of time
NB/ substances w/ higher specific heat capacities require more heat energy to increase their temp. and vice versa
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Enthalpy change of neutralisation
Enthalpy change when an acid and base react together to form one mole of water
Method:
- involves mixing known volumes and conc. of a strong acid and a strong base
- measuring temp. increase
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Method of measuring enthalpy change of neutralisation
- a measured volume of a strong alkali is placed into polystyrene cup- equal volume of strong acid is added
- temp. of reaction mixture increases until neutralisation is complete
- continued addition of acid produces a cooling effect, no further reaction takes place
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Standard enthalpy of combustion
The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions
- enthalpy changes of combustion are always negative- heat is released during combustion process
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Method for standard enthalpy of combustion
- known mass of alcohol is measured into a pre-weighed spirit burner and the alcohol is burned
- heat released increases temp. of a known volume of water in calorimeter
- temp. increase is measured for a certain time period, and experiment is then stopped
- spirit burner and its contents are re-weighed
- mass of alcohol burned to produce temp. increase is recorded, and molar enthalpy of combustion of alcohol can be calculated
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Percentage error equation
% error = ((experimental -theoretical))/ theoretical) x 100
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Limitations of calculating enthalpy changes in a school lab
- heat loss to surroundings
- incomplete combustion
- assumptions made about specific heat capacity and density of aqueous solutions
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State function
Value is independent of path taken to reach that specific value
- enthalpy (heat content of a system) is a state function
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Hess's law
States that, the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same
- change in enthalpy during a chemical reaction doesn't depend on whether reaction proceeds in one step or many steps
- used to calculate enthalpy changes for reactions that can't be determined experimentally
34
Enthalpy cycle
- a cycle of reactants being converted into products
- there are two routs, or paths, that a reaction can take
Direct route: conversion of reactants directly into products
Indirect route: conversion of reactants, via certain reaction intermediates, into products
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Standard enthalpy change of formation
Enthalpy change when one mole of a compound is formed from the elements in their standards states under standard conditions
- values indicate stability of compounds in relation to their elements
- more -ve the value, greater the stability of the compound
Standard enthalpy of formation of an element in its standard state is 0
- because, enthalpy change for formation of one mole of an element in its standard state from itself would be 0, as reactants and products would be the same
- hence, there wouldn't be a change in enthalpy
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Chemical reaction
Involves breaking of existing chemical bonds and making of new chemical bonds
- chemical bonds are results of electrostatic attractions between atoms or ions
- to overcome these attractions (to break the bond), energy must be absorbed
- bond breaking is endothermic (requires energy)
NB/ when new bonds are formed= energy is released
- bond making is exothermic
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When is a chemical reaction exothermic?
In a chemical reaction, when energy absorbed while breaking bonds is less than energy released when forming new bonds
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When is a chemical reaction endothermic?
In a chemical reaction, when energy absorbed while breaking bonds is more than energy released when forming new bonds
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Overall enthalpy change for a reaction
Difference between energy absorbed and energy released
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Bond enthalpy (H)
Energy required to break one mole of chemical bonds in the gaseous state
- Energy required to break a chemical bond
- also called bond dissociation energy, E
- bond enthalpy values are always positive- refer to bonds being broken (endothermic process)
- when new bonds are formed, energy released is equal to energy that was absorbed to break the bond
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Average bond enthalpy
Enthalpy change when one mole of bonds are broken in the gaseous state for the same bond in similar compounds
- same type of chemical bond in different compounds has different bond enthalpy values
- identical bonds in molecules w/ two (or more) types of bond also have different bond enthalpy values
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Experimental vs. average bond enthalpy values
Enthalpy changes for reactions calculated from average bond enthalpy data may differ from experimental values
- because average values taken from same bonds in a range of similar compounds
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Free radical
A highly reactive species due to presence of an unpaired electron
- unpaired electron is represented by a dot
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Earth's atmosphere
- earth has an oxygen-rich atmosphere
- atmosphere contains 2 forms of oxygen: di-oxygen and tri-oxygen (ozone)
- both of these allotropes help protect life on Earth's surface and in lower atmosphere by absorbing UV from the sun
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Early atmosphere of the earth
- lacked any significant level of oxygen- gave little protection against UV radiation
- this could have been a possible factor in rapid evolution, through mutation, of early life-forms
- but following appearance of photosynthesising organisms, levels of atmospheric O2 increased
- hence, came production of a protective, UV-absorbing layer of ozone
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Ozone layer
Stratosphere = 12-50km above surface of the Earth
- lower regions contain 90% ozone in atmosphere; conc. less than 10ppm
- stratospheric ozone is in a dynamic equilibrium w/ oxygen and is continually being formed and destroyed
- ozone levels are maintained naturally by a continuous cycle of synthesis and breakdown reactions by action of energy from UV-C and UV-B radiation
47
Ozone bonds vs. oxygen bonds
- ozone bonds are weaker to those in oxygen (hence, results in phenomena of ozone generation and depletion
- strong covalent bond in di-oxygen can be broken in the stratosphere (10-50km above Earth's surface) by UV w/ wavelength shorter than 242nm, producing 2 free oxygen radicals
48
Ozone generation
Bonds in di-oxygen are broken in stratosphere by UV radiation (wavelength shorter than 242nm) producing 2 free oxygen atoms
molecular oxygen ---> atomic oxygen
Oxygen atoms produced each have an unpaired electron- highly reactive free radicals, react w/ oxygen molecules to form ozone
Reaction removes high-energy UV radiation, prevents it from reaching Earth's surface
- results in this level of stratosphere having a higher temp. than lower region
49
Ozone depletion
Bonds in ozone molecules are weaker than double bonds in di-oxygen- broken by UV radiation of lower energy
- 2 reactions in this cycle depend on different wavelengths of UV radiation- effect is to remove higher energy radiation
- only longer wavelength, less damaging radiation, reaches Earth's surface
- breakdown of ozone requires UV radiation of shorter wavelength than breakdown of ozone
- strong double bond in oxygen is disrupted by Sun's high energy UV-C radiation to form free radicals
- these oxygen radicals can then react w/ an oxygen molecule to form ozone
- bonds in ozone, being weaker, can then be broken by less energetic UV-B radiation (of longer wavelength) to reform oxygen and an oxygen free radical
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UV radiation and the earth
Surface of the earth is protected by Ozone depletion and ozone generation from damaging effects of UV-B and UV-C radiation
- hence, majority of UV radiation reaching Earth's surface is the least harmful UV-A form
- ozone layer protects life on Earth from that radiation that would be most harmful to living tissues
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Ionic compound
Consists of oppositely charged ions held together in a lattice structure by strong electrostatic attractions
- requires energy to overcome electrostatic attractions between ions
- hence, bond breaking is endothermic
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Lattice enthalpy (lattice dissociation enthalpy)
The enthalpy change when one mole of a solid ionic compound breaks down to form gaseous ions under standard conditions
- amount of energy required to separate an ionic compound into its gaseous ions
- it is an endothermic process
NB/ It's directly proportional to product of ionic charges
- inversely proportional to distance between nuclei of ions
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What is the opposite of lattice enthalpy?
Lattice formation enthalpy
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Lattice formation enthalpy
The enthalpy change when one mole of an ionic solid is formed from gaseous ions
- would have same numerical value as lattice dissociation enthalpy, but opposite sign
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Magnitude of lattice enthalpy of an ionic compound
- a measure of the strength of the ionic bonds between the ions
- greater the value of the lattice enthalpy, more energy is required to overcome attraction between ions and stronger the bond
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Magnitude of lattice enthalpy depends on:
1. Charge of the ions (ionic charge)
- charge on ions has significant effect on electrostatic attraction between ions and magnitude of lattice enthalpy
- higher ionic charge = stronger electrostatic attraction between ions, leading to higher value of lattice enthalpy
2. Size of the ions (ionic radii)
- lattice enthalpy decreases as ionic radius of halide ion increases
- weaker electrostatic attractions between ions as ionic radius increases
NB/
- ionic charge increases + ionic radii decrease = lattice enthalpy increases
- ionic charge decreases + ionic radii increase = lattice enthalpy decreases
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Born-Haber cycle
An energy cycle that can be used to calculate the lattice enthalpy of an ionic compound
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Steps of the Born-Haber cycle
1. Enthalpy of formation- enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions
2. Enthalpy of atomisation- enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state
3. Bond dissociation energy: energy required to break 1 mole of bonds in its gaseous state
4. Ionisation energy: energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
5. Electron affinity: energy released when 1 mole of electrons are added to 1 mole of gaseous atoms
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Stable ionic compounds
- large negative values for their enthalpy of formation
| - as more energy is released, compound becomes more stable
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Process of dissolution of ionic compounds
Process of dissolving ionic compounds
1. Break-up of lattice structure; ionic solid is converted to gaseous ions
- enthalpy change for this process = lattice enthalpy
2. Once lattice structure has been broken down, separated gaseous ions are hydrated by surrounding water molecules
3. During hydration, ion-dipole forces are formed between gaseous ions and dipoles of water molecules
4. -ve dipoles are attracted to positively charged ions and +ve dipoles are attracted to negatively charged ions
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Enthalpy of hydration
Enthalpy change when one mole of gaseous ions dissolves in water to form a solution of infinite dilution
- change in enthalpy that occurs when an ion becomes hydrated
- formation of strong ion-dipole forces releases energy
- so, enthalpy of hydration values are always exothermic
General equation:
X+(g) --> X+(aq)
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A solution of infinite dilution
A solution that has a large excess of water
| - addition of more water wouldn't cause any more heat to be released or absorbed
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Enthalpy change of solution
Enthalpy change when 1 mole of solute dissolves to form a solution of infinite dilution
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Solubility of an ionic compound depends on:
- sign of enthalpy change of solution
| - magnitude of enthalpy change of solution
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Insoluble ionic compounds
Relatively high positive value for enthalpy of solution (endothermic)
- such values cause salt to be insoluble because combined enthalpy of hydration values of ions don't compensate for high lattice enthalpy value
- means enthalpy of solution value is endothermic, dissolving process is energetically favourable
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Soluble ionic compounds
- negative exothermic value for enthalpy of solution
- high enthalpy of hydration values compensation for endothermic enthalpy of lattice
- results in an exothermic enthalpy of solution value
- dissolving process is energetically favourable
Substances w/ slight positive (endothermic) values for enthalpy of solution can also be soluble
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Factors affecting enthalpy of hydration
1. Ionic radii
- as ionic radii increase, magnitude of enthalpy of hydration values decreases
- due to weaker ion-dipole forces produced as size of ion increases
2. Charge on ion (charge density of the ion)
- as ionic charge increases, enthalpy of hydration values become increasingly exothermic
- due to strength of ion-dipole force; higher the charge on the ion, stronger the force, more energy is released
Charge density of an ion increases w/ greater ionic charge and a decrease in ionic radius
- higher charge density results in a stronger ion-dipole force between ion and water molecule
- hence a greater, more negative value, for enthalpy of hydration
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Spontaneous reactions
For a reaction to be spontaneous:
- must be an associated increase in total entropy of system together w/ surroundings
- spontaneous change always increases total entropy of universe (2nd law of thermodynamics)
- as entropy increases, amount of energy in universe available in a form to do useful work decreases
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Entropy (S)
Refers to distribution of available energy among particles in a system
- more ways there is for energy to be dispersed, or spread out, greater the entropy
- in a spontaneous process, energy goes from a concentrated form to a dispersed form
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What happens in a spontaneous process?
Occurs without addition of energy, other than that required to overcome initial energy barrier
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Combustion of coal
- example of an increase in entropy for a spontaneous process
- coal is a conc. source of energy that is burned in power stations to generate electricity
- heat produced by combustion of coal heats water to produce steam, this is then used to drive a turbine and ultimately a generator
- steam then passes to cooling towers of power stations where it condenses to form water
- heat released when steam condenses is transferred to surrounding environment
Energy has gone from a conc. form in coal to dispersed form in surroundings
- once in dispersed form, energy is less able to 'do work' = low-quality energy
- energy's natural tendency is to spread out and become less 'useful' over time
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Relative entropy of solids, liquids and gases
Solids:
- fixed arrangements of particles
- lowest entropy
Gases:
- random arrangement of particles
- highest entropy, because energy can be distributed in more ways than in liquids and solids
Entropy of gas > liquid > solid under the same conditions
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Entropy change (ΔS)
- can be negative (−ΔS) or positive (+ΔS)
- in chemical reactions that involve solids or liquids being converted into gaseous products, there will be a large increase in entropy
- if there's a decrease in no. of moles of gas on the product side, decrease in entropy
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Gibbs free energy (G)
relates the energy that can be obtained from a chemical reaction to the change in enthalpy (ΔH), change in entropy (ΔS) and absolute temperature (T)
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The change in Gibbs free energy (ΔG)
a convenient way to take into account both the direct entropy change resulting from the transformation of the chemicals, and the indirect entropy change of the surroundings due to the gain/loss of heat energy
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Standard entropy
Standard entropy of a substance is the consequential entropy change from heating the substance from absolute 0 (0K) to the thermodynamic standard temperature, 298K
- a perfect crystal at 0K has a standard entropy of 0, hence, every other substance has a +ve standard entropy value
- standard entropy values can be used to to calculate standard entropy change (ΔS⦵)
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Standard entropy change (ΔS⦵)
ΔS⦵ = ΣΔS⦵ (products) − ΣΔS⦵ (reactants)
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Spontaneous process
occurs without the addition of energy, other than that required to overcome the initial energy barrier
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Non-spontaneous reaction
one that needs a constant input of energy to occur
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Second law of thermodynamics
states that the total entropy of the universe tends to increase
- tells us the direction of entropy change that favours a spontaneous chemical reaction or physical change taking place
- for a spontaneous process, the total entropy of the system and surroundings (ΔStotal) must increase
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Total entropy of the system and the surroundings
ΔS total = ΔS system + ΔS surroundings ≥ 0
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Why do some chemical reactions have a decrease in entropy but are still spontaneous?
- because the entropy of the surroundings increases to a much greater extent
- this gives an overall increase in entropy for the process
- increase in entropy of surroundings is usually as heat released from the system into the surroundings
- depending on temp. of the surroundings, heat released can have a small or a large effect on the entropy of the surroundings
- If the surroundings are cool, heat will have a large effect, but if surroundings are hot, heat will have a small effect
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Gibbs free energy change, G equation
ΔG⊖=ΔH⊖−TΔS⊖
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Gibbs free energy
The energy associated w/ a process that can be used to do work, specifically non-expansion work
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Work
Work done by a system can be:
1. Expansion work
- due to the change in the volume of a system by expansion or contraction
2. Non-expansion work
- work that isn't associated w/ expansion of a system
- 'useful work'
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Free energy changes in a reaction
- As a reaction proceeds, composition of the reaction mixture is constantly changing, as is the free energy
- position of equilibrium corresponds to a max. value of entropy and a min. value of Gibbs free energy
- once this point has been reached, reaction won't proceed any further – rates of the forward reaction and reverse reaction are now equal
- composition of an equilibrium mixture thus depends on value of ΔG⦵
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3 factors to consider when determining spontaneity of a reaction
- the sign of the ΔH
- the sign of the ΔS
- temp. at which reaction takes place
Thus, ΔG = ΔH – TΔS
ΔH = enthalpy change in kJ/mol
T = absolute temp. in K (kelvin)
ΔS = entropy change in J/K/mol
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Change in enthalpy vs. change in entropy
ΔH units = kJ/mol
ΔS units = J/K/mol
- necessary because unit for Gibbs free energy changes is kJ/mol
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Determining whether a reaction is spontaneous or not
1. If Gibbs free energy change, ΔG, is -ve, reaction is spontaneous
2. If Gibbs free energy change, ΔG, is +ve, reaction is non-spontaneous
3. If Gibbs free energy change, ΔG, is 0, reaction is at equilibrium
NB/ sign of the ΔG value indicates spontaneity of a reaction
- it gives no indication of the rate at which the reaction takes place
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Exothermic and endothermic reactions and spontaneity
1. Exothermic reactions w/ increase in entropy (-ve ΔH and +ve ΔS) are spontaneous regardless of temp.
2. Endothermic reactions w/ decrease in entropy (+ve ΔH and negative ΔS) are non-spontaneous at any temp.
In both of these cases, signs of change in enthalpy and change in entropy mean that ΔG is independent of temp.
1. Exothermic reactions w/ decrease in entropy are only spontaneous at low temp.
– when product of TΔS is small enough to give a -ve ΔG value when subtracted from -ve ΔH value
- At higher temp., product of TΔS, when subtracted from -ve ΔH value, gives a +ve ΔG value (non-spontaneous)
2. Endothermic reactions w/ increase in entropy are only spontaneous at high temp.
- If temp. is high enough, subtracting product of TΔS from ΔH value will give a -ve (spontaneous) ΔG value
- at lower temp., product of TΔS isn't large enough to cancel out +ve ΔH value, so the ΔG value is +ve (non-spontaneous)
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Gibbs free energy of formation (ΔG⦵f)
the change in free energy when one mole of a compound is formed from its elements in their standard states under standard conditions
- can be used to calculate standard Gibbs free energy change, ΔG⦵, for a reaction
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Standard Gibbs free energy change, ΔG⦵
- the change in Gibbs free energy for a reaction at a temperature of 298 K
Equation:
ΔG⦵ = ΣΔG⦵f (products) − ΣΔG⦵f (reactants)
Results:
- sign of ΔG⦵ value indicates stability of a compound relative to its constituent elements
- Compounds w/ a -ve free energy change are more thermodynamically stable than their constituent elements
- compounds w/ a +ve free energy change are less thermodynamically stable than their constituent elements
