Chapter 5/ 15 Flashcards
Energetics/thermochemistry
Heat
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
Enthalpy
Heat released or absorbed by a system at constant pressure
- is the heat content of a system
Standard enthalpy of reaction
When the enthalpy change is measured under standard conditions
- pressure of 100 kPa
- temp. of 25 degrees Celsius (298K)
Temperature
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
Kelvin scale
- 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
Kelvin - degrees Celsius equation
Temperature (K) = temperature (degrees Celsius) + 273.15
System and surroundings
System= chemical reaction Surroundings= rest of the universe
Open systems
Where matter and energy can move freely between system and surroundings
Closed systems
Only energy is able to move between system and surroundings
Relationship between system and surroundings
- energy that is lost from system is transferred to surroundings
- any energy gained by the system is transferred from surroundings to system
Isolated system
- doesn’t exchange energy or matter with its surroundings
Negative change in enthalpy
When heat is released from the system to the surroundings, enthalpy of the system decreases
- change in enthalpy is negative
- only exothermic reactions
Positive change in enthalpy
When heat is transferred to a system from surroundings, enthalpy of system increases
- change in enthalpy is positive
- only endothermic reactions
Law of conservation of energy
Energy cannot be created or destroyed; it’s converted from one form to another
- hence, total amount of energy in universe is constant
Chemical reaction
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)
Exothermic reactions
- 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
Energy profile for an exothermic reaction
- 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
Thermochemical equations
Show enthalpy change of a reaction
Endothermic reactions
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
Energy profile for an endothermic reaction
- enthalpy change is positive because reactants gain heat from their surroundings
- products have higher enthalpy than reactants, and are less energetically stable
Coffee cup calorimeter
- 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
Advantages of simple polystyrene cup as a calorimeter
- 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
Calculating enthalpy changes
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
Specific heat capacity (c)
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
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
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
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
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
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
Percentage error equation
% error = ((experimental -theoretical))/ theoretical) x 100
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
State function
Value is independent of path taken to reach that specific value
- enthalpy (heat content of a system) is a state function
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
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
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
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
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
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
Overall enthalpy change for a reaction
Difference between energy absorbed and energy released
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
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
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
Free radical
A highly reactive species due to presence of an unpaired electron
- unpaired electron is represented by a dot
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
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
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
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
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
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
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
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
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
What is the opposite of lattice enthalpy?
Lattice formation enthalpy
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
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
Magnitude of lattice enthalpy depends on:
- 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 - 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
Born-Haber cycle
An energy cycle that can be used to calculate the lattice enthalpy of an ionic compound
Steps of the Born-Haber cycle
- Enthalpy of formation- enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions
- Enthalpy of atomisation- enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state
- Bond dissociation energy: energy required to break 1 mole of bonds in its gaseous state
- Ionisation energy: energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
- Electron affinity: energy released when 1 mole of electrons are added to 1 mole of gaseous atoms
Stable ionic compounds
- large negative values for their enthalpy of formation
- as more energy is released, compound becomes more stable
Process of dissolution of ionic compounds
Process of dissolving ionic compounds
- Break-up of lattice structure; ionic solid is converted to gaseous ions
- enthalpy change for this process = lattice enthalpy - Once lattice structure has been broken down, separated gaseous ions are hydrated by surrounding water molecules
- During hydration, ion-dipole forces are formed between gaseous ions and dipoles of water molecules
- -ve dipoles are attracted to positively charged ions and +ve dipoles are attracted to negatively charged ions
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)
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
Enthalpy change of solution
Enthalpy change when 1 mole of solute dissolves to form a solution of infinite dilution
Solubility of an ionic compound depends on:
- sign of enthalpy change of solution
- magnitude of enthalpy change of solution
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
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
Factors affecting enthalpy of hydration
- 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 - 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
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
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
What happens in a spontaneous process?
Occurs without addition of energy, other than that required to overcome initial energy barrier
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
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
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
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)
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
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⦵)
Standard entropy change (ΔS⦵)
ΔS⦵ = ΣΔS⦵ (products) − ΣΔS⦵ (reactants)
Spontaneous process
occurs without the addition of energy, other than that required to overcome the initial energy barrier
Non-spontaneous reaction
one that needs a constant input of energy to occur
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
Total entropy of the system and the surroundings
ΔS total = ΔS system + ΔS surroundings ≥ 0
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
Gibbs free energy change, G equation
ΔG⊖=ΔH⊖−TΔS⊖
Gibbs free energy
The energy associated w/ a process that can be used to do work, specifically non-expansion work
Work
Work done by a system can be:
- Expansion work
- due to the change in the volume of a system by expansion or contraction - Non-expansion work
- work that isn’t associated w/ expansion of a system
- ‘useful work’
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⦵
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
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
Determining whether a reaction is spontaneous or not
- If Gibbs free energy change, ΔG, is -ve, reaction is spontaneous
- If Gibbs free energy change, ΔG, is +ve, reaction is non-spontaneous
- 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
Exothermic and endothermic reactions and spontaneity
- Exothermic reactions w/ increase in entropy (-ve ΔH and +ve ΔS) are spontaneous regardless of temp.
- 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.
- 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) - 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)
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
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
