×

Warning

Please Login to Read More...

Learning from life!
The nature of energy

The production of energy is one of the most important and pervasive aspects of life. Think for a moment about all of the chemical reactions that we use to produce energy during the course of a typical day in our lives: We eat foods to produce the energy needed to maintain our biological function. We burn fossil fuels to produce the electrical energy that is central to our modern society, to produce heat for our homes, and to produce power for planes, trains, and automobiles. We use ice cubes to cool our drinks, and we use heat to convert raw dough into baked bread. Green plants absorb energy from the Sun to fuel the chemical reactions that lead to their growth. All of these examples illustrate the general point that chemical reactions involve energy. Some reactions, such as the burning of gasoline, release energy.

Others, such as the splitting of water into hydrogen and oxygen, require the addition of energy. Over 90 percent of the energy produced in our society comes from chemical reactions, mostly from the combustion of coal, petroleum products, and natural gas. The concept of matter has always been as easy notion to grasp because matter can be seen and felt. Our article of thermodynamics will also involve the concepts of energy, work, and heat. The study of energy and its transformations is known as thermodynamics. Today thermodynamics is enormously important in all areas of science and engineering, as we will see throughout this article.

Wind is a source of energy Turbines like these capture the wind's energy and turn it into another form of energy, such as electricity.
What is energy?

In physical science, energy is one of the most fundamental and universal concepts, but one that is remarkably difficult to define in a way that is meaningful to most people. This perhaps reflects the fact that energy is not a "thing" that exists by itself, but is rather an attribute of matter (and also of electromagnetic radiation) that can manifest itself in different ways. It can be observed and measured only indirectly through its effects on matter that acquires, loses, or possesses it.

Energy is another word for power. Energy makes things move. It makes machines work and makes living things grow. Energy is the capacity to cause change. In everyday life, energy is important because some forms of energy can be used to do work i.e to move matter against opposing forces, such as gravity and friction. If we put it in another way, energy is the ability to rearrange a collection of matter. For example, we expend energy to turn the pages of a book, and our cells expend energy in transporting certain substances across membranes. Although there are many specific types of energy, the two major forms are Kinetic energy and Potential energy.

The kinetic energy of an object is the energy that it possesses due to its motion
Forms of energy transfer.
Forms of energy

Kinetic energy is associated with the motion of an object, and its direct consequences are part of everyone's daily experience; the faster the ball you catch in your hand, and the heavier it is, the more you feel it. We see that the kinetic energy increases as the speed of an object increases. For example, a car moving at 50 miles per hour (mph) has greater kinetic energy that it does when it travels at a speed of 40 miles per hour. We also see that, for a given speed, the kinetic energy increases with increasing mass. Thus, a large sport-utility vehicle travelling at 55 mph has greater kinetic energy than a small sedan travelling at the same speed, because of the greater mass of the former. Atoms and molecules have mass and are in motion and therefore possess kinetic energy.

An object can also possess potential energy by virtue of its position relative to other objects. Potential energy is, in essence, "stored" energy that results from the attractions and repulsions an object experiences in relation to other objects. We are all familiar with instances in which the potential energy of an object is converted into kinetic energy. For example, consider a bicyclist at rest on the top of a steep hill, as shown in the adjacent figure. Because of the attraction of gravity, the potential energy of the bicyclist and her bicycle is greater at the top of the hill than it would be at the bottom. As a result, the bicycle easily proceeds down the hill with increasing speed. As it does so, the stored potential energy of the bicycle is converted into kinetic energy. Forces other than gravity can lead to potential energy. The most important attractions and repulsions in chemistry are the electrostatic forces between charged particles. Thus, an electron has potential energy when it is near a proton, because they attract one another.

One of our goals in chemistry is to relate the energy changes that we see in our microscopic world to the kinetic or potential energy or substances at the atomic or molecular level. For example, many substances, such as fuels, release energy when they react. The chemical energy of these substances is due to the potential energy stored in the arrangements of the atoms of the substance. Likewise, we will see that the energy a substance possesses because of its temperature (its thermal energy) is associated with the kinetic energy of the molecules in the substance.

The energy is transferred from the bat to the ball, causing the ball to move against the force of gravity
Transferring energy: Work and heat

We can illustrate the forms of energy transfer with a simple example. In the figure, the energy is transferred from the bat to the ball, causing the ball to move against the force of gravity. We see energy transferred in the form of heat. We will see that energy is transferred in two general ways: to cause the motion of an object against a force or to cause a temperature change. A force is any kind of push or pull exerted on an object. As we noted in the figure, the force of gravity pulls a bicycle from the top of a hill to the bottom. Energy used to cause an object to move against a force is called work. The work, w, that we do in moving objects against a force equals the product of the force, F, and the distance, d, that the object is moved:

W = F × d

Thus, we perform work when we lift an object against the force of gravity. If we define the object as the system, then we – as part of the surroundings – are performing work on that system, transferring energy to it.

In an open system, both matter and energy can be exchanged with the system's surroundings freely. In this example with the pot of water boiling, matter is leaving the system in the form of steam.

In a closed system, only energy can be exchanged between the system and the surroundings.
Systems and surroundings

Thermodynamics is the study of energy transfers that occur in molecules or collections of molecules. When we are discussing thermodynamics, the two words we often come across, those are – system and surroundings. The particular item or collection of items that we’re interested in (which could be something as small as a cell, or as large as an ecosystem) is called the system, while everything that's not included in the system we’ve defined is called the surroundings.

There are three types of systems in thermodynamics: open, closed, and isolated.
Open systems is one that freely allows energy and matter to be transferred in an out of a system. For example, boiling water on the stovetop example would be an open system, because heat and water vapor can be lost to the air. Another interesting example is, - "We", are an open system. Whether we think about it or not, we are constantly exchanging energy and matter with our surroundings. For instance, suppose that we eat a carrot, or lift a bag of laundry onto a table, or simply breathe out and release carbon dioxide into the atmosphere. In each case, you are exchanging energy and matter with your environment.

A closed system, on the other hand, can exchange only energy with its surroundings, not matter. It allows heat to be transferred from the stove to the water in the previous example with closed lid. Heat is also transferred to the surroundings, steam is not allowed to escape. Example of a closed system – a pressure cooker, it would approximate a closed system. An isolated system is one that cannot exchange either matter or energy with its surroundings. This system is completely sealed. Matter is not allowed to be exchanged with the surroundings. Heat cannot transfer to the surroundings. Example – A thermoflask is an isolated system.

Exchanges of energy that take place in living creatures must follow the laws of physics. In this regard, let's take a closer look at how the laws of thermodynamics (physical rules of energy transfer) apply to living beings like yourself.

Shown are two examples of energy being transferred from one system to another and transformed from one form to another. Humans can convert the chemical energy in food, like this ice cream cone, into kinetic energy (the energy of movement to ride a bicycle). Plants can convert electromagnetic radiation (light energy) from the sun into chemical energy.
The first law of thermodynamics

The First law of thermodynamics states that heat is a form of energy, and thermodynamic processes are therefore subject to the principle of conservation of energy. This means that heat energy cannot be created or destroyed. Put another way, the First law of thermodynamics states that energy cannot be created or destroyed. It can only be in a change form or be transferred from one object to another. If we start to look at examples, we'll find that transfers and transformations of energy take place around us all the time. For example:

  • Light bulbs transform electrical energy into light energy (radiant energy).
  • One pool ball hits another, transferring kinetic energy and making the second ball move.
  • Plants convert the energy of sunlight (radiant energy) into chemical energy stored in organic molecules.
  • We are transforming chemical energy from our last snack into kinetic energy as we walk, breathe, and move our finger to scroll up and down this page.

But why does the change occur in the first place? From everyday experience, it seems that some changes happen by themselves – that is, spontaneously – almost as if a force were driving them in one direction and not the other. Turn on a gas stove, for example, and the methane mixes with oxygen and burns immediately to yield carbon dioxide and water vapor. But those products will not remake methane and oxygen no matter how long they mix. A steel shovel left outside slowly rusts, but put a rusty one outside and it won’t become shiny. A cube of sugar dissolves in a cup of coffee after a few seconds of stirring, but stir for another century and the dissolved sugar won’t ever reappear as a cube.

Chemists speak of a process that occurs by itself as being spontaneous. Some absorb energy, whereas others release it. The principles of thermodynamics, which were developed almost 200 years ago to help utilize the power derived from steam, allow us to understand the nature of spontaneous change. Despite their narrow historical purpose, these principles apply, as far as we know, to every system in the universe!

Significantly, none of these transfers is completely efficient. Instead, in each scenario, some of the starting energy is released as thermal energy, when it's moving from one object to another, thermal energy is called by the more familiar name of heat. It's obvious that glowing light bulbs generate heat in addition to light, but moving pool balls do too (thanks to friction), as do the inefficient chemical energy transfers of plant and animal metabolism. To see why this heat generation is important, let’s stay tuned for the second law of thermodynamics.

The term spontaneous does not mean instantaneous or have anything to say about how long a process takes to occur, it means that, given enough time, the process will happen by itself. Many processes are spontaneous but slow – ripening, rusting, and aging.
The Second law of thermodynamics

According to the first law of thermodynamics, if energy is never created or destroyed, that means that energy can just be recycled over and over again, right? As per this statement, energy cannot be created or destroyed, but it can change from more-useful forms into less-useful forms. As it turns out, in every real-world energy transfer or transformation, some amount of energy is converted to a form that’s unusable (unavailable to do work). In most cases, this unusable energy takes the form of heat. Although heat can in fact do work under the right circumstances, it can never be turned into other (work-performing) types of energy with 100% efficiency. So, every time an energy transfer happens, some amount of useful energy will move from the useful to the useless category.

In a formal sense, a spontaneous change of a system, whether a chemical or physical change, or just a change in location, is one that occurs by itself under specified conditions, without a continuous input of energy from outside the system. The freezing of water, for example, is spontaneous for the system at 1 atm and -5 degree C. A spontaneous process such as burning or falling may need a little "push" to get started – a spark to ignite gasloline, – but once the process begins, it keeps going without the need for any continuous external input of energy.

In contrast, for a non-spontaneous change to occur, the surroundings must supply the system with a continuous input of energy. A book falls spontaneously, but it rises only if something else, such as a human hand (or a strong wind), supplies energy in the form of work. Under a given set of conditions, if a change is spontaneous in one direction, it is not spontaneous in the order. The term spontaneous does not mean instantaneous or have anything to say about how long a process takes to occur, it means that, given enough time, the process will happen by itself. Many processes are spontaneous but slow – ripening, rusting, and aging. Thus, a chemical reaction proceeding toward equilibrium is an example of a spontaneous change.

Entropy is the degree of randomness or disorder in a system
Entropy and the Second Law of Thermodynamics

Entropy is the degree of randomness or disorder in a system. The measure of the level of disorder in a closed but changing system, a system in which energy can only be transferred in one direction from an ordered state to a disordered state. Higher the entropy, higher the disorder and lower the availability of the system's energy to do useful work. Since we know that every energy transfer results in the conversion of some energy to an unusable form (such as heat), and since heat that does not do work goes to increase the randomness of the universe, we can state a biology-relevant version of the second law of thermodynamics: every energy transfer that takes place will increase the entropy of the universe and reduce the amount of usable energy available to do work (or, in the most extreme case, leave the overall entropy unchanged). In day-to-day life it manifests in the state of chaos in a household or office when effort is not made to keep things in order. In other words, any process, such as a chemical reaction or set of connected reactions, will proceed in a direction that increases the overall entropy of the universe.

To sum up, the first law of thermodynamics tells us about conservation of energy among processes, while the second law of thermodynamics talks about the directionality of the processes, that is, from lower to higher entropy (in the universe overall).

A whole-body calorimeter In this room-sized apparatus, a subject exercises while respiratory gases, energy input and output, and other physiological variables are monitored.
Do living things obey the laws of thermodynamics?

Organisms can be thought of as chemical machines that evolved by extracting energy as efficiently as possible from the environment. Such a mechanical view holds that all processes, whether they involve living or nonliving systems, are consistent with thermodynamic principles. Let's examine the first and second laws of thermodynamics to see if they apply to living systems.

Organisms certainly comply with the first law. The chemical bond energy in food is converted into the mechanical energy of sprouting, crawling, swimming, and countless other movements; the electrical energy of nerve conduction; the thermal energy of warming the body; and so forth. Many experiments have demonstrated that in all these energy conversions, total energy is conserved.

It may not seem as clear, however, that an organism, and even the whole parade of life, complies with the second law. Mature humans are far more complex than the simple egg and sperm cells from which they develop, and modern organisms are far more complex than the one-celled ancestral specks from which they evolved. Energy must be highly localized and macromolecules constrained to carry out myriad reactions that synthesize biopolymers from their monomers.

Are the growth of an organism and the evolution of life exceptions to the spontaneous tendency of natural processes to disperse their energy and increase molecular freedom? Does biology violate the second law? Not at all, if we examine the system and surrounding- For an organism to grow or a species to evolve, countless moles of oxygen and food molecules – carbohydrates, proteins, and fats – undergo exothermic combustion reactions to form many more moles of gaseous CO2 and H2O. The formation and discharge of these waste gases, as well as the heat released, represent a tremendous net increase in the entropy of the surroundings. Thus, the localization of energy and restrictions on molecular freedom apparent in the growth and evolution of organisms occur at the expense of a far greater dispersal of energy and freedom of motion in the Earth-Sun surroundings. When system and surroundings are considered together, the entropy of the universe, as always, increases.

To sum up, the high degree of organization of living things is maintained by a constant input of energy, and is offset by an increase in the entropy of the surroundings.

References

  • Silberberg – Chemistry – The Molecular Nature of Matter and Change
  • Brown LeMay Bursten
  • https://www.khanacademy.org/science/biology/energy-and-enzymes/the-laws-of-thermodynamics/a/the-laws-of-thermodynamics
  • https://www.livescience.com/50881-first-law-thermodynamics.html
  • http://www.sustainablebrands.com/news_and_views/new_metrics/joss-tantram/vitality-how-life-makes-best-2nd-law-thermodynamics