Bond Energy

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Main Question or Discussion Point

"When we supply bond energy to two molecules that have bonded, the potential energy of the molecules increase causing the molecules to break bonds." -Is this the correct explanation regarding bond energy? If not, kindly explain this process to me, it's very confusing.
 

Answers and Replies

  • #2
Simon Bridge
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The bond energy is the amount of work you have to do to break the bond.
 
  • #3
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The bond energy is the amount of work you have to do to break the bond.
Doesn't the bond energy get converted to potential energy? If no, what does it get converted to?
 
  • #4
Simon Bridge
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The work that was done breaking the bond was stored as potential energy, yes.

A simple case would be the way everyone is bound to the earth by gravity.
A 1kg mass on the surface of the earth has -125.440MJ in gravitational potential energy ... notice the minus sign?
Any force you use to pull the mass away from the earth will need to add at least +125.44MJ to get the mass completely free.
More work than that and the 1kg will also have some kinetic energy after it is free.
If the force only did, say, 25MJ work, the new potential energy would be -100.44MJ (and the mass would be higher up).
The 25MJ of energy went into the gravitational field between the earth and the mass... you can get it back by letting the mass fall.

chemical bonding is sort of like that - only it uses electromagnetism instead of gravity and it happens on smaller scales (and there's quantum stuff.)
That help?
 
  • #5
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The OP asked about bond energy between two molecules not two atoms. These would be Van der Waal's forces, not covalent bonding. Am I misreading, or did the OP mean covalent bonds?
 
  • #6
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The work that was done breaking the bond was stored as potential energy, yes.

A simple case would be the way everyone is bound to the earth by gravity.
A 1kg mass on the surface of the earth has -125.440MJ in gravitational potential energy ... notice the minus sign?
Any force you use to pull the mass away from the earth will need to add at least +125.44MJ to get the mass completely free.
More work than that and the 1kg will also have some kinetic energy after it is free.
If the force only did, say, 25MJ work, the new potential energy would be -100.44MJ (and the mass would be higher up).
The 25MJ of energy went into the gravitational field between the earth and the mass... you can get it back by letting the mass fall.

chemical bonding is sort of like that - only it uses electromagnetism instead of gravity and it happens on smaller scales (and there's quantum stuff.)
That help?
Yes that helped very much. :D Just one last question sir, can I say that the potential energy is helping the molecule to move further away and stay there instead of banging into the other molecule?
 
  • #7
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The OP asked about bond energy between two molecules not two atoms. These would be Van der Waal's forces, not covalent bonding. Am I misreading, or did the OP mean covalent bonds?
Hey, I was talking about the intermolecular forces.
 
  • #8
Simon Bridge
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Van der Walls forces are electromagnetic in nature... with quantum stuff. ;)
https://en.wikipedia.org/wiki/Van_der_Waals_force.
"Van der Waals forces include attraction and repulsions between atoms, molecules, and surfaces, as well as other intermolecular forces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics.)"

can I say that the potential energy is helping the molecule to move further away and stay there instead of banging into the other molecule?
That would be like saying that gravitational potential energy helps falling objects avoid banging into the ground.

Objects that are bound to each other sort of hang around together.
At the molecular level - they are like dancers at a party: jiggling and twirling around all over the place but still associated with each other.
The jiggling is thermal - basically it is kinetic energy that holds objects, otherwise attracted to each other, apart.
There are also repulsive intermolecular forces (see link above).

This is also the scale where objects stop having surfaces the way we are used to ... so they cannot strike each other in the way that pool balls do.
 
  • #9
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Van der Walls forces are electromagnetic in nature... with quantum stuff. ;)
https://en.wikipedia.org/wiki/Van_der_Waals_force.
"Van der Waals forces include attraction and repulsions between atoms, molecules, and surfaces, as well as other intermolecular forces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics.)"


That would be like saying that gravitational potential energy helps falling objects avoid banging into the ground.

Objects that are bound to each other sort of hang around together.
At the molecular level - they are like dancers at a party: jiggling and twirling around all over the place but still associated with each other.
The jiggling is thermal - basically it is kinetic energy that holds objects, otherwise attracted to each other, apart.
There are also repulsive intermolecular forces (see link above).

This is also the scale where objects stop having surfaces the way we are used to ... so they cannot strike each other in the way that pool balls do.
So what force during state transition pulls molecules away from each other and then gets converted to potential energy? What basically is this work? There must be some force doing this work, what do we call it?
 
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  • #10
Simon Bridge
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You can identify the forces at play yourself... it's a good exercise since you need to get used to thinking about this.

The most common example of a state-change where intermolecular bonds are broken is the vaporization of water.
You've done it lots! So what do you do to make water go from liquid to vapour state?
What force are you applying?
 
  • #11
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You can identify the forces at play yourself... it's a good exercise since you need to get used to thinking about this.

The most common example of a state-change where intermolecular bonds are broken is the vaporization of water.
You've done it lots! So what do you do to make water go from liquid to vapour state?
What force are you applying?
Let me put it like this: An Ice cube is given enough heat to break down its bonds, this heat initially turns to kinetic energy which drags the molecules away from each other and then it quickly turns into potential or chemical energy. Now these molecules are further apart and the bonds have vanished due to the distance.
Does that make sense or am I still getting it all wrong? :p
 
  • #12
Simon Bridge
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That's not bad. I'll just tweak it a bit ... note: heat is a form of kinetic energy.
When you heat something on an element, you are transmitting heat by conduction and some radiation to the ice (via the pan) ... the heat itself is probably provided by electricity and heat moves from hot places to cold places by molecular jostling: which works by the electromagnetic force.

molecules of the solid ice start out in a crystal lattice structure - all juggling about in place.
The jiggling motion is heat - kinetic energy is constantly being exchanged for electromagnetic potential energy.
The more heat the more jiggling you get. The average speed of the jiggling about is the temperature.

From here you can see that too much jiggling will break the solid apart.
 
  • #13
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That's not bad. I'll just tweak it a bit ... note: heat is a form of kinetic energy.
When you heat something on an element, you are transmitting heat by conduction and some radiation to the ice (via the pan) ... the heat itself is probably provided by electricity and heat moves from hot places to cold places by molecular jostling: which works by the electromagnetic force.

molecules of the solid ice start out in a crystal lattice structure - all juggling about in place.
The jiggling motion is heat - kinetic energy is constantly being exchanged for electromagnetic potential energy.
The more heat the more jiggling you get. The average speed of the jiggling about is the temperature.

From here you can see that too much jiggling will break the solid apart.
So because this heat turns into electromagnetic potential energy, the temperature of the whole thing does not rise. Whereas when we supply heat to the unbonded molecules most of this kinetic energy on being used up turns to heat again and causes the temperature to rise. Right?

And because the attraction is continuous the molecules firstly have to gain enough force against the attraction to break the bond, and when this energy gets converted to potential energy the molecules are a good distance apart from each other to bond again. Right?

Can I say that the kinetic energy initially gets converted to potential energy because the other molecule tries to restrict its movement?
 
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  • #14
ogg
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I don't agree with most of the responses here. Kinetic Energy is the energy involved in motion and vibration is not simply motion (since with it, velocity varies periodically with time). Vibration is both kinetic and potential energy and usually the SUM of the two is the (constant) vibrational energy. Just as a planet's orbital energy has both potential and kinetic terms. Or just as two opposite charges have both P.E. and (probably) K.E. (depending on the system). Heat is a very difficult term, generally its definition is either formal or vague/murky. Mostly, from a practical perspective, heat is the flow of energy. Or even more practical, heat is what a thermometer measures changes of. Perhaps you can see by this that both heat and P.E. are not specific types of energy? So using them to explain only works as long as you're willing to keep the explanation quite superficial. I also have a serious problem understanding the original post: 1) if two molecules have bonded then there is one molecule, right? If the product molecule required the addition of more energy, (and that energy wasn't returned as kinetic, electronic, or vibrational (rotational, etc.) energy) then the product is less stable than the reactants. Some chemical reactions (eg simple acid-base reactions) nearly instantaneously form product, but many reactions (especially the more "interesting" ones) require some more energy to form a transition state (which means a state of higher energy, bond energy to be specific (bond energy being the electromagnetic interactions between all atoms, which involves quantum mechanical interactions)) which then rearranges to form product(s). The over-all reaction may cause cooling or heating and may be slow or fast (but generally chemical reaction steps are very very fast compared to what we can perceive with our natural senses). Energy may come from bond energy or kinetic (thermal) energy (as well as rotational, vibrational, and conformational energy) and be "used" to change atoms' positions or to change the electron clouds surrounding atoms or groups of atoms (usually both). Do bonds break "because of" P.E.? Well, sure - in fact one way we determine whether two atoms are bonded is by measuring their distance apart. While it varies depending on the atoms and electronic states involved, most if not all (I can't think of an example where this is not true) bond breaks require the atoms to separate (distance wise) which therefore requires motion, hence k.e. and that has to come from somewhere. It can come from k.e. (collisions) but as soon as the energy is transferred from velocity (deceleration) it goes somewhere (and we call that P.E. unless some other object rebounds with more k.e.) or it can come from one or more of the many types of P.E. (which include electrostatic, conformation, vibration, and rotation (and by the way, there are several kinds of rotation)). The OP seems to assume that the product molecule is in its lowest (P.E.) energy state, and that is NOT given in the original problem. A molecule may spontaneously decay (if it isn't the lowest energy for the system), so saying added P.E. "causes" a bond to break is only true in certain situations. But bond energy is P.E. by definition, so any change in bonds is a change in P.E. (even if the total P.E. remains constant for the entire system). BTW, the OP also has a problem in the wording: Two molecules form one molecule which then breaks up into (two?) other molecules (or ions). Both the bond (Van der Waals "bonds" are generally not considered "bonds" or "chemical bonds" even though they are bonds between chemicals..but that depends on the context.) making and bond breaking involves changes in P.E. but what the TOTAL P.E. of the 3 states is, is not stated and could be anything A>B>C or A<B<C or A>B<C or A=B<C ..I think that makes for 9 possibilities....
 
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  • #15
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I don't agree with most of the responses here. Kinetic Energy is the energy involved in motion and vibration is not simply motion (since with it, velocity varies periodically with time). Vibration is both kinetic and potential energy and usually the SUM of the two is the (constant) vibrational energy. Just as a planet's orbital energy has both potential and kinetic terms. Or just as two opposite charges have both P.E. and (probably) K.E. (depending on the system). Heat is a very difficult term, generally its definition is either formal or vague/murky. Mostly, from a practical perspective, heat is the flow of energy. Or even more practical, heat is what a thermometer measures changes of. Perhaps you can see by this that both heat and P.E. are not specific types of energy? So using them to explain only works as long as you're willing to keep the explanation quite superficial. I also have a serious problem understanding the original post: 1) if two molecules have bonded then there is one molecule, right? If the product molecule required the addition of more energy, (and that energy wasn't returned as kinetic, electronic, or vibrational (rotational, etc.) energy) then the product is less stable than the reactants. Some chemical reactions (eg simple acid-base reactions) nearly instantaneously form product, but many reactions (especially the more "interesting" ones) require some more energy to form a transition state (which means a state of higher energy, bond energy to be specific (bond energy being the electromagnetic interactions between all atoms, which involves quantum mechanical interactions)) which then rearranges to form product(s). The over-all reaction may cause cooling or heating and may be slow or fast (but generally chemical reaction steps are very very fast compared to what we can perceive with our natural senses). Energy may come from bond energy or kinetic (thermal) energy (as well as rotational, vibrational, and conformational energy) and be "used" to change atoms' positions or to change the electron clouds surrounding atoms or groups of atoms (usually both). Do bonds break "because of" P.E.? Well, sure - in fact one way we determine whether two atoms are bonded is by measuring their distance apart. While it varies depending on the atoms and electronic states involved, most if not all (I can't think of an example where this is not true) bond breaks require the atoms to separate (distance wise) which therefore requires motion, hence k.e. and that has to come from somewhere. It can come from k.e. (collisions) but as soon as the energy is transferred from velocity (deceleration) it goes somewhere (and we call that P.E. unless some other object rebounds with more k.e.) or it can come from one or more of the many types of P.E. (which include electrostatic, conformation, vibration, and rotation (and by the way, there are several kinds of rotation)). The OP seems to assume that the product molecule is in its lowest (P.E.) energy state, and that is NOT given in the original problem. A molecule may spontaneously decay (if it isn't the lowest energy for the system), so saying added P.E. "causes" a bond to break is only true in certain situations. But bond energy is P.E. by definition, so any change in bonds is a change in P.E. (even if the total P.E. remains constant for the entire system). BTW, the OP also has a problem in the wording: Two molecules form one molecule which then breaks up into (two?) other molecules (or ions). Both the bond (Van der Waals "bonds" are generally not considered "bonds" or "chemical bonds" even though they are bonds between chemicals..but that depends on the context.) making and bond breaking involves changes in P.E. but what the TOTAL P.E. of the 3 states is, is not stated and could be anything A>B>C or A<B<C or A>B<C or A=B<C ..I think that makes for 9 possibilities....
I had concerns with bonding such that in ice molecules which is broken and the ice turns into water.
Also, I've few queries:
1) Does bonding force reduce with distance?
2) Does the kinetic energy get converted to potential energy because the molecule it was bonded to is opposing its motion?
 
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  • #16
Ygggdrasil
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1) Does bonding force reduce with distance?
The potential energy of intermolecular bonds is often modeled as a Lennard-Jones potential, which has the following shape:
=Figure_B.jpg

(note that the interaction between water molecules involves hydrogen-bonding so it is probably more complicated than an L-J potential)

Remember that force (F) is given by the equation: ##F = –\frac{dV}{dr}##, so the force is given by the derivative (slope) of the potential energy function (V). At distance ##r_{eq}## where V is minimum, the atom experiences no net force (because attractive and repulsive forces cancel out). As the distance between the atoms increases from ##r_{eq}##, the atom experiences an attractive force that increases to some point, but then begins to decrease again to zero as the two atoms become very far apart.

2) Does the kinetic energy get converted to potential energy because the molecule it was bonded to is opposing its motion?
Unbound atoms have a higher potential energy (zero in the diagram above) than bound atoms (-ε in the diagram above). Therefore, heat gets converted to potential energy when you break an intermolecular bond.
 
  • #17
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Unbound atoms have a higher potential energy (zero in the diagram above) than bound atoms (-ε in the diagram above). Therefore, heat gets converted to potential energy when you break an intermolecular bond.
Had these molecules been unbound, this latent heat would have been absorbed by the molecules and got converted to heat energy thereby increasing the temperature. But because this heat energy is absorbed by bonded molecules it gets converted to potential energy, therefore the temperature does not rise.
Is that right?
 
  • #18
Ygggdrasil
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Had these molecules been unbound, this latent heat would have been absorbed by the molecules and got converted to heat energy thereby increasing the temperature. But because this heat energy is absorbed by bonded molecules it gets converted to potential energy, therefore the temperature does not rise.
Is that right?
Yes, I'd say that's fairly accurate (though the term latent heat is misused).

Normally, when you add heat to a substance, it increases the kinetic energy of the molecules in that substance. At some point, you reach a point where you cannot add more kinetic energy without disrupting bonds between molecules. At this point a phase transition occurs and added heat goes to breaking the intermolecular bonds between molecules in the substance (increasing their potential energy). During the process of the phase transition, the added heat is converted to potential energy—not kinetic energy—so the temperature does not rise. The total amount of heat that needs to be added to fully effect the phase transition is called the latent heat and is essentially equal to the potential energy difference between the bound molecules before the transition and the unbound molecules after the transition.
 
  • #19
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Yes, I'd say that's fairly accurate (though the term latent heat is misused).

Normally, when you add heat to a substance, it increases the kinetic energy of the molecules in that substance. At some point, you reach a point where you cannot add more kinetic energy without disrupting bonds between molecules. At this point a phase transition occurs and added heat goes to breaking the intermolecular bonds between molecules in the substance (increasing their potential energy). During the process of the phase transition, the added heat is converted to potential energy—not kinetic energy—so the temperature does not rise. The total amount of heat that needs to be added to fully effect the phase transition is called the latent heat and is essentially equal to the potential energy difference between the bound molecules before the transition and the unbound molecules after the transition.
So this heat energy initially gets converted to kinetic energy and then potential energy, isn't that potential energy somewhere keeping the molecule away from the attractive force of the other molecule?
 
  • #20
Ygggdrasil
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Yes. In order to separate two bound molecules, you must perform work in order to overcome the attractive force between the molecules. Energy from heat is used to overcome the attractive force and break the intermolecular bond.
 
  • #21
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Yes. In order to separate two bound molecules, you must perform work in order to overcome the attractive force between the molecules. Energy from heat is used to overcome the attractive force and break the intermolecular bond.
Imagine an ice cube kept at -10 degree celsius. When we supply heat to this ice cube, the temperature goes down until 0 degrees celsius where phase transition occurs.
My question- These molecules at -10 degree celsius are also bonded but the heat they get gets converted to kinetic energy without any restriction from the bond and increases the temperature. But when this is heat is supplied to molecules at 0 degree celsius the attractive force of the bonds comes into place and restricts the movement of the molecules, converting the kinetic energy into potential energy. Why does this not happen at -10 degree celsius?
 
  • #22
Ygggdrasil
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In a solid, the atoms/molecules that compose the solid are confined to a crystal lattice. The kinetic energy of the system manifests through vibrations of these atoms/molecules within the lattice. As the temperature rises and the energy of these vibrations increases, at some point the vibrations will get large enough to begin disrupting the lattice. This is the point at which melting occurs. The attractive forces are always in place, and what determines when melting occurs is when the energy of the vibrations in the solid approach the energy of the attractive interactions.
 
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  • #23
Borek
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force of the bonds comes into place and restricts the movement of the molecules
That's not entirely true. While the molecules are kept in place they oscillate - and oscillation is a form of motion.
 
  • #24
James Pelezo
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Just an add-on FYI on bonding and bonding energies in science ... Generally, there are three types of bonding forces known, Physical Bonds, Chemical Bonds and Nuclear Bonds...Physical Bonds account for 'states of matter'; i.e., solids, liquid or gas. Individual bonding energies are typically very weak and are easy to break with low-energy input. Changes in state do not affect the chemical make-up of the substance. Chemical bonds are valence level interactions between elements. Chemical bonding energy depends upon the elements involved and their elemental properties. For example, a Carbon - Carbon 'single' bond ( C-C ) bond energy ~350 Kj/mol (~heat and light energy from a common burning wax candle), a 'double bond' ( C=C ) bond energy ~ 600 Kj/mol and a 'triple bond' ( C≡C ) bond energy ~ 835 Kj/mol. The larger numbers are the stronger bonds. It is important to note that the energy given up on bond formation is exactly equal to the amount of energy needed to break the bond. Change elements bonding and the bonding energy of the combination changes. Nuclear bonds are interactions between subnuclear particles; i.e., protons and neutrons. Bonding energies for nuclear level interactions are commonly referred to as 'Binding Energy per Nucleon' and expressed as Mega-electron-volt (MeV) values. (1.0 MeV = 1.6 x 10-16 Kj). This means, the binding energy per nucleon equals the total bonding energy of all nucleons divided by the total number of neutrons and protons making up the nucleus. Nuclear bonding energies are also referred to as the 'Mass Defect' on formation of a nucleus. That is, Mass Defect is the amount of nuclear material converted into energy on formation of a nucleus. Mathematically, this is defined by the famous Einstein Equation ΔE = Δmc2 where Δm is the mass defect and c is the speed of light (3 x 108 m/s). Molar level Nuclear Binding Energies run in the 108th Kj/mol; e.g., atomic bomb levels. Quite a jump from a burning candle. Also, the energy of the sun is nuclear fusion where 'Heavy' Hydrogen atoms collide forming Helium atoms + lots of sunshine. Doc
 
  • #25
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In a solid, the atoms/molecules that compose the solid are confined to a crystal lattice. The kinetic energy of the system manifests through vibrations of these atoms/molecules within the lattice. As the temperature rises and the energy of these vibrations increases, at some point the vibrations will get large enough to begin disrupting the lattice. This is the point at which melting occurs. The attractive forces are always in place, and what determines when melting occurs is when the energy of the vibrations in the solid approach the energy of the attractive interactions.
That energy which disturbs the lattice gets converted to potential energy. RIght?
 

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