Equilibrium: Crash Course Chemistry #28

Summary

This episode of Crash Course Chemistry explores the concept of chemical equilibrium, comparing it to life's balance and emphasizing its importance in various chemical reactions. It explains how reactions can be reversible, with forward and reverse processes occurring simultaneously until they reach a state where their rates are equal. The video delves into the Haber process for ammonia production, illustrating how equilibrium can be manipulated to favor desired outcomes. It also introduces Le Châtelier's Principle, which predicts how equilibrium shifts in response to changes in concentration, temperature, or pressure. The episode concludes with a visual demonstration involving cobalt ions, showcasing the dynamic nature of equilibrium.

Takeaways

• 🧘 Life and nature both seek balance, which is called equilibrium in science.
• 🍰 Disruptions to balance can be natural and temporary, such as overpopulation of deer in a forest.
• 🍕 Enjoying indulgences like a hot pocket can fit into a balanced diet with proper balance.
• 🔁 Chemical reactions aren't always one-way; many are reversible and can reach a state of equilibrium.
• ⚖️ Chemical equilibrium occurs when the forward and reverse reactions happen at the same rate.
• 🧩 Chemists often work to disrupt equilibrium to produce desired chemicals in greater concentrations.
• 🌟 The Haber process for making ammonia is an example of a reaction that is actually an equilibrium, contrary to common simplifications.
• 🔍 Le Châtelier's Principle states that a system at equilibrium will adjust to minimize the impact of any stress, such as changes in concentration, temperature, or pressure.
• 💥 Fritz Haber successfully developed a process for synthesizing ammonia, despite Le Châtelier's failed attempt, with significant consequences for warfare and agriculture.
• 🌡️ Temperature changes can affect equilibrium by favoring endothermic or exothermic reactions depending on whether heat is added or removed.
• 🤔 Understanding and manipulating chemical equilibrium involves mathematical calculations that will be covered in future lessons.

Q & A

What is the concept of balance in the context of life and nature as described in the script?

The concept of balance in life and nature refers to maintaining equilibrium in various aspects such as financial stability, a balanced diet, work-life balance, and overall health. In nature, equilibrium refers to the state where a natural system is in balance, and when disrupted, nature tends to restore the balance, such as in the case of a deer population in a forest.

What is chemical equilibrium and why is it significant in the context of chemical reactions?

Chemical equilibrium is a state in a reversible chemical reaction where the forward reaction (reactants forming products) and the reverse reaction (products turning back into reactants) occur at the same rate, resulting in no net change in the concentration of reactants and products. It is significant because it represents a dynamic balance where reactions continue to occur but appear to be at a standstill to an observer.

How does the script explain the relationship between diet and equilibrium?

The script uses the example of enjoying an occasional hot pocket without destroying a balanced diet. It suggests that it's possible to indulge as long as you balance it out with healthier food choices, illustrating the concept of maintaining equilibrium in personal dietary habits.

What is the Haber process and why is it important in the context of chemical equilibrium?

The Haber process is a method for producing ammonia from nitrogen in the air and hydrogen gas. It's important in the context of chemical equilibrium because it involves a reversible reaction that can reach equilibrium. However, industrially, chemists use techniques to prevent equilibrium from being reached, to maximize ammonia production.

Who was Henry Louis Le Châtelier and what principle is named after him?

Henry Louis Le Châtelier was a French chemist born in Paris in 1850. He is known for formulating Le Châtelier's Principle, which states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

What is the significance of Fritz Haber in the development of the process to produce ammonia?

Fritz Haber was a German chemist who developed the Haber process for synthesizing ammonia from nitrogen and hydrogen. Despite the controversy surrounding his involvement in chemical warfare, his process revolutionized agriculture by enabling the production of chemical fertilizers, thus earning him the 1918 Nobel Prize in Chemistry.

How does the script illustrate the effect of concentration changes on chemical equilibrium?

The script explains that adding or removing substances from a reaction at equilibrium can shift the balance. For instance, adding more nitrogen to the Haber process at equilibrium would drive the reaction to the right, producing more ammonia, while removing ammonia would have a similar effect by reducing what's available to break down.

How does pressure affect the position of equilibrium in a chemical reaction?

Pressure changes can shift the equilibrium position. In reactions involving gases, increasing the pressure favors the side of the reaction with fewer gas molecules, while decreasing the pressure has the opposite effect. This is demonstrated in the script with the Haber process, where high pressure favors the formation of ammonia.

What role does temperature play in affecting chemical equilibrium?

Temperature changes can also affect the position of equilibrium. Endothermic reactions (which absorb heat) are favored at higher temperatures, while exothermic reactions (which release heat) are favored at lower temperatures. Adding heat to a reaction shifts the equilibrium to the left (towards reactants), while cooling shifts it to the right (towards products).

Can you provide an example from the script of how changing conditions can shift a chemical equilibrium?

The script provides the example of two cobalt ions in an aqueous solution that reach equilibrium and change color depending on the conditions. Adding hydrochloric acid pushes the equilibrium to the right (blue side), while adding water pushes it to the left (pink side). Changing the temperature also shifts the equilibrium, with higher temperatures favoring the blue side and lower temperatures the pink side.

What did the script imply about the balance between the positive and negative impacts of scientific discoveries?

The script implies that scientific discoveries, such as the Haber process, can have both positive and negative impacts. While the process led to the production of chemical explosives used in wars, it also enabled the production of chemical fertilizers that revolutionized agriculture and saved millions of lives.

Outlines

00:00

🌟 The Concept of Balance and Equilibrium

The first paragraph introduces the overarching theme of balance in various aspects of life, including financial, dietary, and health-related balances, as well as the scientific concept of equilibrium. It explains that disruptions to natural systems can lead to imbalances, but nature often restores balance, such as in the case of a deer population in a forest. The paragraph also touches on the idea that chemical reactions aren't always straightforward processes with a clear beginning and end, but can instead be part of a dynamic equilibrium where reactants and products continuously convert into each other. The concept of chemical equilibrium is introduced, where the forward and reverse reactions occur at the same rate, and it's noted that chemists often work to disrupt this equilibrium to produce desired chemicals, such as in the Haber process for making ammonia.

05:02

🔬 The Haber Process and Le Châtelier's Principle

This paragraph delves into the specifics of the Haber process, which is used to synthesize ammonia from nitrogen and hydrogen. It clarifies that this process is actually an equilibrium reaction, contrary to previous simplifications. The paragraph explains how the rates of the forward and reverse reactions adjust as the concentrations of reactants and products change, eventually reaching a state of equilibrium where the concentrations remain constant. The concept of Le Châtelier's principle is introduced, which states that a system at equilibrium will adjust to minimize the impact of any stress, such as changes in concentration, temperature, or pressure. The historical context of the Haber process and its inventor, Fritz Haber, is also discussed, including the moral complexities surrounding his work and its uses in warfare and agriculture.

10:02

🔄 Factors Affecting Chemical Equilibrium

The final paragraph explores the factors that can affect the position of chemical equilibrium, such as changes in the concentration of reactants or products, pressure, and temperature. It explains how increasing or decreasing these factors can shift the equilibrium position either to the right (favoring product formation) or to the left (favoring reactant formation). The paragraph also discusses the role of solids and liquids in reactions and how they are less affected by pressure changes compared to gases. The effects of temperature on endothermic and exothermic reactions are highlighted, showing how temperature adjustments can influence the direction of the reaction. The paragraph concludes with a practical application of these principles using the example of cobalt ions in an aqueous solution, demonstrating visually how equilibrium can be manipulated through the addition of hydrochloric acid or water, and by temperature changes.

Keywords

💡Equilibrium

Equilibrium refers to a state of balance where opposing forces or reactions are equally matched, resulting in no net change. In the context of the video, it is used to describe the balance in natural systems and chemical reactions. The script explains that when a natural system is disrupted, it tends to restore balance, and similarly, many chemical reactions seek a state of equilibrium where the forward and reverse reactions occur at the same rate.

💡Chemical Reactions

Chemical reactions are processes that transform one set of chemical substances into another. The video script discusses how people often perceive chemical reactions as having a clear start and end, but in reality, many reactions are reversible and can continue indefinitely. This concept is central to understanding equilibrium in chemistry, as it is exemplified by the forward and reverse reactions in the context of chemical equilibrium.

💡Forward Reaction

A forward reaction is the process where reactants are converted into products in a chemical reaction. The script uses this term to describe the initial part of a reversible reaction, where substances like nitrogen and hydrogen combine to form ammonia, as in the Haber process. This concept is fundamental to understanding how equilibrium is achieved, as the forward reaction is one half of the balancing act.

💡Reverse Reaction

A reverse reaction is the opposite of a forward reaction, where products are converted back into reactants. The video script explains that in a system at equilibrium, there is also a reverse reaction occurring simultaneously with the forward reaction. This is crucial for understanding how equilibrium is maintained, as the reverse reaction counterbalances the forward reaction.

💡Gibb's Free Energy

Gibb's free energy is a thermodynamic potential that measures the maximum reversible work that a thermodynamic system can perform at constant temperature and pressure. The script mentions it in the context of explaining why reactions are reversible when they can go forward or backward without extra energy being used. It's a key concept for understanding the spontaneity of reactions and their tendency to reach equilibrium.

💡Haber Process

The Haber process is an industrial method for the production of ammonia from nitrogen and hydrogen gases. The script discusses this process as an example of a chemical reaction that is typically thought of as one-way but actually exists as an equilibrium. The Haber process is significant in the video's narrative as it illustrates the practical applications and implications of understanding and manipulating chemical equilibrium.

💡Le Châtelier's Principle

Le Châtelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. The video script explains this principle through the actions of adding or removing substances, or changing temperature or pressure, which affect the equilibrium position. This principle is central to the video's theme of balance and change in chemical reactions.

💡Concentration

Concentration in chemistry refers to the amount of a substance present in a given volume or mass. The script discusses how changes in the concentration of reactants or products can disturb an equilibrium, causing the reaction to shift in a direction that restores balance. This is a key aspect of Le Châtelier's Principle and is essential for understanding how equilibrium can be manipulated.

💡Pressure

Pressure is the force applied per unit area. In the context of the video, it is discussed as a factor that can affect the position of equilibrium, especially in reactions involving gases. The script uses the Haber Process as an example, explaining how increasing pressure favors the formation of ammonia by shifting the equilibrium to the right, while decreasing pressure has the opposite effect.

💡Temperature

Temperature is a measure of the average kinetic energy of the particles in a system. The video script explains how temperature changes can affect chemical equilibrium, favoring endothermic reactions when heat is added and exothermic reactions when heat is removed. This is a crucial concept for understanding how equilibrium can be influenced by environmental conditions.

💡Endothermic Reactions

Endothermic reactions are chemical reactions that absorb energy from their surroundings, typically in the form of heat. The script mentions endothermic reactions in the context of temperature changes affecting equilibrium. When heat is added to an endothermic reaction, it favors the reaction proceeding in the direction that consumes heat, which often means producing more products.

💡Exothermic Reactions

Exothermic reactions are chemical reactions that release energy, usually in the form of heat, to their surroundings. The video script explains that exothermic reactions are favored at lower temperatures and that the heat they produce can shift the equilibrium position of a reaction. This concept is important for understanding how temperature can influence the direction in which a reaction proceeds.

Highlights

Life balance is analogous to chemical equilibrium, where balance is a natural tendency but can be disrupted and restored.

Equilibrium in nature often self-corrects, such as overpopulation of deer being regulated by lack of resources and predators.

Chemical reactions are not always linear with a clear end, often seeking balance through reversible reactions.

Forward and reverse reactions represent the dynamic nature of many chemical processes.

Chemical equilibrium is achieved when forward and reverse reactions occur at the same rate.

Chemists often work to prevent equilibrium to favor desired product concentrations.

The Haber process for ammonia production is a real-world example of manipulating equilibrium for industrial purposes.

Le Châtelier's Principle states that a system at equilibrium will adjust to minimize stress from changes in concentration, temperature, or pressure.

Adding or removing reactants or products can shift the equilibrium position, as demonstrated by the Haber process.

Pressure changes can significantly affect the position of equilibrium, especially in gaseous reactions.

Temperature adjustments can favor either endothermic or exothermic reactions, impacting the equilibrium position.

The industrial Haber Process operates under high pressure to favor ammonia production.

Fritz Haber's development of the ammonia production process had both positive and negative impacts on society.

Haber's continuous removal of ammonia from the reaction mixture prevents the system from reaching equilibrium.

Chemical equilibrium can be visually demonstrated with the color-changing cobalt ions reaction.

Equilibrium is not about remaining static but adjusting to maintain balance as conditions change.

Transcripts

Life is all about balance.

You want to have a balanced bank account, balanced diet, balance between work and play,

and a healthy inner ear to keep you from falling down all the time.

And lots of things can mess up that balance.

One bad night at the cheesecake factory and you might wake up with your finances, social life, and health out of whack.

In science, our word for balance is equilibrium. You've heard of it.

When the balance of some natural system gets disrupted, we say that it's out of equilibrium,

but that's not always a bad thing, because nature usually finds a way to restore the balance.

Like if a population of deer in a forest suddenly goes nuts

and there's too many of them, they'll run out of food and space,

and predators will move in and the population will usually fall back to where the habitat can sustain it.

Likewise, there's no reason that I can't enjoy the occasional hot pocket

without destroying my so-called balanced diet,

I just have to be sure to eat something that wasn't prepared in a salty lard factory to balance it out.

We tend to think of chemical reactions as having a beginning and an end:

you start with some reactants, some chemistry happens, and you end up with the products, and that's that.

But that's not that, because you don't usually end up with pure products.

Many chemical reactions seek balance too, and just like us, they have to work through the stress in order to find it.

[Theme Music]

It's common, because it's usually helpful, to think of chemical reactions as simple straightforward processes,

but in reality, many reactions never finish. Ever.

When reactants react to form products, what we usually mean when we talk about chemical reactions,

it's technically called the forward reaction.

But sometimes, there's an opposite reaction that occurs with the products changing back into the reactants.

We call this a reverse reaction.

Now, exactly what kind of reactions do this, and why,

is hard to describe without busting out some 200 level chemistry kung fu,

but it all goes back to our lesson about Gibb's free energy and reactions that occur spontaneously.

Reactions are reversible when they can go forward or backward without any extra energy being used.

Remember, this isn't just one reaction happening once, it's billions and trillions of reactions,

and some of them might be going one way while others are going another way.

And when the forward and reverse reactions occur at the same rate, that's called chemical equilibrium.

And even though we all like the idea of balance in our lives, in this case it's not always a desirable thing.

In fact, most chemists make their living using tricks to prevent chemical equilibrium,

to knock it out of whack, maximizing the concentrations of the chemicals they want to produce

at the expense of the balance that nature usually seeks.

For example, consider the Haber process for making ammonia from nitrogen in the air.

When we discussed it as a redox reaction a few weeks ago,

we described it as a one way, complete reaction because that's all we needed to understand right then.

But I was lying to you.

Sorry, I'd rather you hear it from me than out on the streets.

Turns out, the reaction really exists as an equilibrium.

As nitrogen and hydrogen react to form ammonia, the concentration of those gases drops,

making them less likely to collide and keep reacting, so the rate of the forward reaction slows down.

At the same time, the concentration of ammonia rises.

More ammonia molecules are available to break up into the reactant gases.

And surprise, the rate of the reverse reaction speeds up.

Eventually, those processes reach a point where they happen at the same rate.

At that point, there's no discernible change in the concentration of any gas.

Nitrogen and hydrogen keep combining to form ammonia,

while ammonia keeps breaking down into nitrogen and hydrogen at the same time.

Notice that I did not say that the reaction stops.

The reaction basically never stops, it's just that we don't notice any changes at that point,

because everything that happens in one direction is perfectly balanced out by what happens in the other direction.

That's why reactants like this are written with a double arrow, indicating that the process runs both directions.

Like if I stand on a balanced board, I never stop moving.

I just shift my body back and forth to compensate for the motion of the board.

When it moves left, I move right and vice versa, and by doing that, I'm able to stay upright, hopefully.

But sometimes there is, if you will, a disturbance in the force,

and as any good Jedi knows, those things must be put right.

Chemical equilibria can be disturbed by changes in the concentration of one or more substances

or by changes in temperature or pressure.

We describe these changes based on which way they force the equilibrium.

We say a change shifts the reaction to the right if it tends to make more products form

and to the left if it tends to make more reactants form.

This was all summarized nicely by a French chemist named Henry Louis Le Châtelier who was born in Paris in 1850.

Today, we call call his summary Le Châtelier's Principle,

and it says that if stress is placed on a system that is at equilibrium

the system will proceed in a direction that minimizes the stress.

That can happen in several ways.

For starters, changing the concentration of any substance in a reaction

causes it to proceed in whatever direction restores the former balance.

For example, once the Haber reaction is at equilibrium if you were to add more nitrogen gas to the mix

the existing hydrogen would have more nitrogen to react with.

Thus, it would begin forming more ammonia,

sending the reaction noticeably to the right until the reaction balances itself again.

Removing some of the ammonia would have the same effect.

There would be less ammonia available to break down,

so the formation of the ammonia would temporarily exceed the formation of nitrogen and hydrogen,

and the reaction would again shift to the right until the balance is restored.

Now here's an ironic side of irony served up for you.

In 1901, Le Châtelier tried to invent a process for fixing nitrogen to make ammonia from nitrogen and hydrogen.

But he gave up the effort after the experiment caused a huge explosion that killed one of his lab assistants.

8 years later, the German chemist, Fritz Haber,

probably the closest the world ever got to a literal evil genius, did just what Le Châtelier was trying to do.

And in contrast to Le Châtelier's despair at the death of his assistant,

Haber had no problem at all with the fact that his procedure now known as the Haber Process

was used to make chemical explosives which killed hundreds of thousands of people in World Wars I & II.

But just to prove that there aren't clear winners or losers in science,

the Haber Process also saved uncounted millions of lives

because the ability to fix nitrogen has made chemical fertilizers possible,

and basically revolutionized modern agriculture.

Haber was awarded the 1918 Nobel Prize for his discovery,

but 30 years later he was considered by many to be a war criminal

and spent his last years as an object of shame to his family and hatred for many others.

Le Châtelier, meanwhile despite his great contributions to our understanding of equilibrium,

considered his failure to find an efficient method for making ammonia to be the greatest blunder of his scientific career.

One of the key achievements that Haber made was

figuring out a way to continuously remove ammonia as it was being produced.

This allowed the forward reaction to keep going faster than the reverse reaction.

Basically preventing the process from reaching a state of equilibrium.

But changing concentrations is only one part or actually one third of the story.

There's also pressure to consider.

When you look at the equation for the Haber Process, for example,

you can see that 4 moles of gas react to form 2 moles of gas.

So the forward reaction decreases the volume while the reverse reaction increases it.

This is important.

Increasing the pressure will then put more stress on the high volume reactants than the low volume products,

so the reaction shifts to the right, producing more ammonia than it does at low pressure.

And indeed, the industrial Haber Process is done at two hundred atmospheres of pressure.

Decreasing the pressure, meanwhile, has the opposite effect.

The reaction precedes in the direction that raises the pressure back to where it was before.

In this case that's the break-down of ammonia.

Toward the left where there are 4 moles of gas.

Solids and liquids aren't affected by pressure as much,

so the more gasses that are present in the reaction, the greater effect any pressure change will have.

The third and final way to affect the position of equilibrium is with temperature changes.

It gets a bit complicated up in here,

but basically the thing to keep in mind is that endothermic reactions, which consume heat,

are favored if heat is added so higher temperatures tend to feed endothermic reactions.

And by the same token, exothermic reactions, which release heat fair better at low temperatures.

And since exothermic reactions produce heat, that heat tends to favor the reverse endothermic reactions.

So if heat is added to reaction, it forces the reaction back to the left.

But if heat is removed by cooling the reaction mixture, the reaction will proceed to the right.

Now, you may be wondering since battling equilibrium is what most chemists do for a living,

how they work out the specifics of it.

How do they determine how much of each substance is present, or how much they need to add,

or what the temperature or pressure should be when everything is constantly changing.

You probably won't be surprised to learn that there is math we can do to answer those questions.

We'll get into the numbers in our next lesson.

But right now, let's have some fun.

I've been talking about the Haber Process because it's a simple reaction, it's an important reaction.

And because it's one of the main reactions that Le Châtelier studied.

But it's not very interesting to watch.

A much more interesting reaction involves 2 different ions of cobalt that reach equilibrium in an endothermic reaction.

As you can see, one of the ions is pink in aqueous solution and the other is blue.

One stress I can put on this system is the addition of hydrochloric acid.

This increases the chloride ion concentration and pushes the reaction to the right, the blue side.

And if I add water, the reaction is pushed right back to the left, pink. Cool, right?

I can also stress the reaction by changing its temperature.

If I raise the temperature, it's like adding a reactant so the reaction will proceed to the right.

You use it up and turn blue again.

You can probably guess by now that if I lower the temperature

it's like taking away some reactants so the reaction proceeds to the left and turns pink.

We could make these colors go back and forth all day by making changes to the reaction,

but I think you're getting the idea.

Equilibrium isn't about staying the same all the time,

it's just about keeping your balance as your circumstances change.

And now it is time for me to go get a nice Greek salad to help balance out

the corn dogs that I may or may not have eaten this weekend.

But first thank you for watching this episode of Crash Course Chemistry.

If you paid attention, you learned that equilibrium is just a fancy word for balance,

which is a thing that chemical reactions need just as much as we do.

And that the way they achieve balance is to compensate for change, not just to stop completely.

You also learned and saw that chemical equilibrium can be affected by the concentration of substances

their temperatures, and if they're gasses, their pressure.

Finally, you learned about Le Châtelier's principle of chemical equilibrium

and about his failed attempt to do what Fritz Haber eventually did.

Invent an efficient process to produce ammonia.

This episode of Crash Course Chemistry was written by Edi Gonzales.

This script was edited by Blake de Pastino. And our chemistry consultant was Dr. Heiko Langner.

It was filmed, edited, and directed by Nicholas Jenkins. Our script supervisor, Katherine Green.

And our sound designer is Michael Aranda. And of course, our graphic team is Thought Cafe.