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0:01
The existence of atoms has been proposed since ancient times. Around 2500 years ago, Greek philosopher Democritus questioned himself, “could an object be divided into smaller and smaller pieces forever?” He believed that after a certain point we can’t divide more. So, there is a limit to the number of times an object could be divided. That indivisible single solid part was called ‘atomos’, so tiny that the human eye can’t detect them. According to his idea,
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everything is made up of this indivisible matter and comes in infinite amounts of varieties like different shapes and sizes. Differences in shape and size determines the properties of matter. During the period of Democritus, an Indian philosopher, Archarya Kanad, proposed the same concept: that an object cannot be divisible into smaller parts after a certain level and that the indivisible tiny matter comes in different varieties. He named that tiny smallest matter ‘anu’.
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These earlier philosophical ideas were mere speculations and there was no way to test them experimentally. The words ‘atomos’ and ‘anu’ were used in many philosophical books in later years, however these ideas remained inactive for a very long time. The atomic theory of matter was first proposed on a scientific basis by a British school teacher John Dalton in 1808. He proposed that all matter is made of atoms which are indivisible,
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with the smallest part being a solid sphere and massy particles. Every atom of a given element is the same as every other atom of that element. For example, in gold, there are gold atoms, and all gold atoms in that element have the same size, mass, and properties. He also noted that the atoms of one element differ from the atoms of all other elements, meaning that there are various types of atoms, with each type having different properties such as mass, chemical
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reactions, boiling points, or melting points. A sodium atom is not the same as a carbon atom. He briefly explained the law of conservation of mass and the law of constant composition at atom level. In 1789, Antoine Lavoisier introduced the law of the conservation of mass, which states that mass is neither created nor destroyed during a chemical reaction. For example, if we add 8 grams of milk with 8 grams of coffee powder, we would get 16 grams of
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coffee. Here, the mass remains the same, conserved after the chemical reaction. The same concept was explained by Dalton using atoms. He stated that atoms cannot be created or destroyed. Since atoms cannot be destroyed, the same mass will remain after a chemical reaction. Now, let’s take two hydrogen atoms and one oxygen atom. The atomic mass of hydrogen is 1.00784, and for oxygen 15.999. So, when combined, we will get a water molecule with a weight
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of 18.01468. This weight is the same for all water molecules. Hence, the total mass of all the atoms before the reaction must equal the total mass of all atoms after the reaction. Another important law that was easily explained by Dalton is the law of constant composition or law of definite proportions. This law states that, when different elements are involved in a chemical reaction, they are always combines by a fixed ratio. For example,
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1 gram of water always has 0.1 grams of hydrogen and 0.9 grams of oxygen, therefore the mass ratio of water is 1:9. This can be explained with the number of atoms. If we add 8 hydrogen atoms with 8 oxygen atoms, we will get 4 water molecules. Here, one water molecule contains 2 hydrogen and 1 oxygen atoms. The remaining 4 oxygen atoms will be left over. So, the water molecule's fixed ratio by atom is 2:1. These ratios are fixed and unchangeable.
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With the help of Dalton's atomic theory in later years, most of the atom's mass and many molecules' fixed composition ratios were found. However, he concluded that atoms are spherical, solid, indivisible, and the final ultimate particles. This statement was followed for many years until one of the greatest discoveries happened in science. From the 1850’s, scientists were studying cathode rays. Well, what are cathode rays? To produce
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cathode rays, the experiment starts with a low air pressure glass tube. An anode and a cathode are fixed inside the glass tube. Anodes and cathodes are called electrodes. The fixed anode and cathode are connected with a high voltage battery. The anode is connected to the positive side of the battery and the cathode is connected to the negative side of the battery. Normally, a complete electric circuit has a conductor which is a connection between the negative
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terminal and positive terminals. But here, they used electrodes. Electrodes are metal plates used to make contact in non-metallic parts and act as conductors. Here we can see low pressure gasses between the cathode and anode. Gasses are non-metallic matter. When switching on the battery, a high voltage current is passed through the cathode to the anode. From the cathode, a ray passes through the anode to the other end of
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the tube in a straight line. This experiment proved to be a nightmare to the scientists, who struggled to guess what these rays were actually made of. At that time, everyone was well aware that something was coming out from the cathode metal plate, but they didn't know what it actually was. So, they named that mysterious phenomenon as “cathode rays”. But in 1897, British physicist J. J. Thompson discovered what cathode rays actually are. At
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that time, physicists were well aware that the same charges repel each other and opposite charges attract each other. Thompson conducted the same experiment but he added another negative and positive plate within the path of the cathode ray. Once he placed the charged plates, he noticed that the path of cathode rays changed in the direction of the positive plate. He found out that the rays are nothing but negative particles,
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that’s why they attract to the positive plate direction. So, he concluded that if negative particles were coming from the cathode, the atom should contain negative particles. He named those particles as ‘electrons’. With the result of his many experimental observations, he proposed an atomic model called ‘the plum pudding model’, which had negatively charged electrons embedded within a positively charged soup, or, simply, electrons floating within a
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sea of positive charge. This is because opposite charges attract each other, and so electrons, which are negatively charged, stick with positively charged fluid. Negatively charged electrons and positively charged soup together make an atom of neutral charge. Even though Thompson was the first to discover a subatomic particle, physicists and chemists further investigated and questioned this model.
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In 1911, New Zealand born physicist, Ernest Rutherford, conducted a famous experiment with gold foil and alpha particles. During the radioactive process, positively charged alpha particles are produced. For now, it’s better to skip the makeup of alpha particles, because during the time of the experiment, physicists did not know about the composition of alpha particles. However, they were sure that alpha particles were positively charged particles.
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Rutherford’s experiment started with a very thin foil of pure gold, a radioactive source which produces alpha particles, and a circular screen with fluorescent zinc sulfide coating . Whenever alpha particles struck the screen, a tiny flash of light was produced at the collision point. We already know that gold foil is made of gold atoms. During the experiment, Rutherford noticed that most of the alpha particles penetrated into the
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gold foil without any deflection. A few alpha particles slightly deflected and very few particles deflected back at an almost 180-degree angle. With this experiment, he concluded that most of the space inside an atom is empty. That’s why alpha particles easily pass through gold foil without any deflection. Few particles were deflected because the positive charge in an atom is not uniformly distributed,
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unlike in J. J. Thompson’s atomic model. Very few alpha particles deflected back because the volume occupied by positively charged particles in an atom is very very small compared to the total volume of an atom. Here again, we need to consider that, because the same charges repel each other, when a positively charged alpha particle hits the positively charged particle of an atom, it is repelled and deflected.
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Over years of experiments, Rutherford proposed his atomic model. The positively charged particles and most of the mass of an atom is concentrated in an extremely small volume, which he called the nucleus, with negatively charged electrons surrounding it. He also claimed that those electrons revolve around the nucleus at a very high speed in a fixed circular path. This is similar to the planetary model. The Sun is the nucleus;
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planets are electrons. His atomic model gave a perfect understanding of atoms. But the real problem came with James Clerk Maxwell, who was well-known for his work on electromagnetism. According to his theory, an accelerating charged particle emits energy in the form of electromagnetic radiation. Electrons orbit around the nucleus in a circular path. Circular motion is an accelerating motion because circular motion
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continuously changes its direction or velocity. If a particle continuously changes its direction, it cannot be in constant motion. Therefore, circular motion is an accelerated motion. As we already saw, according to Maxwell, an accelerated charged particle emits electromagnetic radiation. So,the electron is a charged particle, and when it accelerates, it must continuously emit electromagnetic radiation. Within a fraction of a second it falls into the nucleus. Then,
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no atom can survive more than a trillionth of a second. So, there is no atom, no matter, no universe. Simply, Rutherford’s atomic model failed to explain the stability of an atom. In order to explain the stability of an atom, Niels Bohr came up with a new model. Bohr’s model of the atom comes with 4 postulates. The first one was to explain the stability of an atom. Even though an electron orbits in a circular path with the accelerated motion,
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it does not lose any energy because it moves in a fixed energy state or fixed circular path where electrons cannot lose energy — these are called orbits or shells. The first shell closest to the nucleus is n=1, the second shell n=2, the third shell n=3, and so on. They can also be called K, L, M, N, etc.. Electrons cannot occupy orbits between these discrete ones. The number of electrons in these fixed orbits is determined by the shells. The first inner shell
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can hold up to 2 electrons, second shell up to 8 electrons, the third shell up to 18 electrons and next level shells can hold even more. Bohr derived the mathematical formula for finding the number of electrons that can hold in different shells as 2n squared. For example, if we want to find the number of electrons that can be held in the fourth shell, we just need to apply n=4 in this formula. So, we get the answer: 32. It means the fourth shell can hold up to 32 electrons.
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We know that electrons orbit positively charged nuclei at different shells. But at what distance are they orbiting? This is where the second postulate of Bohr’s atomic model comes in: radius of orbit. He derived an equation to find the shell radius from the center. For our understanding, we will look at that equation in simple terms. Bohr’s atomic radius is equal to 0.53 n squared divided by z Angstrom. Here, n means shell number, Z is the atomic number of an atom or how
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many positively charged particles are present in an atom (simply, known as protons). For example, in a hydrogen atom, there is only one positively charged particle, so the atomic number of a hydrogen atom is 1. Angstrom is a unit to measure distance. 1 Angstrom = 10^-10 m or 0.1 nanometer. From this equation we can find at which distance the electron orbits. If we want to find the radius
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of the first shell of a hydrogen atom, we simply put the values in. The first shell of a hydrogen atom is n=1, and atomic number is z=1. So, we will get 0.53 Angstrom. For the second shell radius of the hydrogen atom, it will be 2.11 Angstrom. He even derived a formula for finding the speed at which electrons orbit: velocity = 2.186 * 10 ^ 6 (z divided by n). Similar to the radius of a shell, we can calculate the speed of electrons. Then, for the first shell of a hydrogen
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atom, the electrons orbit at the speed of 2.186 * 10 ^ 6 m/s. So, electrons orbit the nucleus in a fixed orbit at a fixed radius at fixed speed. And what about fixed energy? That is the third postulate of the Bohr atomic model. Once he derived a formula for velocity of an electron, then it was easy to find the total energy of an electron by adding the kinetic and potential energies of an electron. So, the formula for the total energy of an electron is -13.6 * (z squared
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divided by n squared ) electron volts. So, for the hydrogen atoms, the first shell’s electron has -13.6 electron volts, the second shell -3.4, the third shell -1.5, the fourth shell -0.85, the fifth shell -0.54, and so on. So, the energy of electrons in atoms is quantized. Quantized means the first shell has energy of -13.6 electron volts; it can’t be -13.7 or -13.5 as it is fixed. It will be the same for all hydrogen atoms which exist in our universe.
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But his fourth postulate states that electrons can lose or gain energy. How is that possible? To understand this, we need to travel back to the time of James Maxwell. He said that when charged particles move, they emit alternating electromagnetic radiation. This radiation is characterized by the properties, frequency, and wavelength. Radiation is nothing but electromagnetic waves. These waves come in different lengths — shorter
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to longer wavelengths. Regardless of the wavelength, in a vacuum, it always travels at the same speed — 300,000 km/s — and doesn’t require any medium to travel through. Based on its wavelengths, radiation is divided into different categories: radio waves, microwaves, infrared, visible light, ultraviolet, x rays and gamma rays. Humans can see only a certain range, 380 to 740 nanometers of wavelengths which we call visible light.
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As we already know, visible light has different colors because different wavelengths of visible light gives different colors. If we pass a beam of sunlight through a prism, we can see the output of the color range starting from red to violet. Each part of visible light's wavelengths has a color that we call the visible light continuous spectrum. So, ordinary white light consists of all the wavelengths in the visible range.
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But the mystery continued for a long time. When white light passes through gasses, the output doesn’t come with a continuous spectrum. It comes as a line spectrum. For example, when white light enters hydrogen gas, the outcome is only certain patterns of color such as red, green, blue, and violet. This means that something in the hydrogen atom absorbs the remaining wavelengths. But which part of the atom absorbs this radiation?
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Since the early decades, light has been known as waves and consists of energy. But later, the photoelectric effect proved that light is also a particle. They named it a photon. Of course, light is a wave but also expresses particle features, so there was no clear winner in the debate of whether light is a wave or particle. It has dual features, meaning electromagnetic radiation acts like both waves and particles, which we call wave-particle duality.
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Now, we come to the Bohr model of hydrogen atom. Electrons orbit in a natural orbit, but when they gain energy, they jump from a lower orbit to a higher orbit. What does that mean? How did Bohr explain that? As with the third postulate of Bohr’s atomic model, we know that the energy of the electron orbit is quantized. Electrons in the first orbit of a hydrogen atom have an energy of -13.6 electron volts, the second orbit -3.4, the third -1.5, and so on. Each atom contains an
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infinite number of orbits. At the “infinite” orbit level, the energy of an electron becomes zero. The energy of an electron doesn’t change with time because it is quantized. However, when electrons gain energy, they move from lower stationary states to higher state orbit. These electrons gaining energy are called excited electrons. Since the energy of an electron is quantized, how can it gain energy? And when electromagnetic radiation or light passes through hydrogen atoms,
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only red, green, blue and violet color wavelengths are emitted. How is this done? First, we need to know that all ranges of electromagnetic radiation have different levels of energy. Since an electron's energy is quantized, it absorbs the right amount of energy. For example, if it gains 10.2 ev, it moves from first orbit to second orbit. Because the first orbit of a hydrogen atom electron has an energy of -13.6 ev,
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when it gains +10.2 ev the energy of an electron becomes -3.4 electron volts. This -3.4-electron volt is an energy state of the second orbit of a hydrogen atom. With this energy, an electron can only exist in the second orbit. So, it moves to a second orbit. If it gains 13.06 electron volt energy, it directly jumps from first orbit to fifth orbit. However, these electrons are always trying to come into their stationary state. As a result,
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they release energy, photons, or electromagnetic radiation. But how does a hydrogen atom only emit red, green, blue, and violet visible light? Now, consider the electron returning from the second orbit to the first orbit. Here, the second orbit has energy of -3.4 electron volts. In order to return to the first orbit, it should release 10.2 electron volts. We can convert this energy into nanometers with our equation, then we will get a wavelength of 122 nanometers. So,
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what kind of photon is this? A 122 nanometer wavelength is nothing but the wavelength of ultraviolet photons. This is not visible light so it can’t be detected by human eyes. But there is also a hydrogen atom that emits visible light photons. We know that as the hydrogen emission line spectrum. If an electron in a hydrogen atom returns from the third orbit to the second orbit it emits energy of 1.9 electron volts ,this energy is a
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photon with a 654 nanometer wavelength. It is a visible light photon of a red color. Similarly, returning from the fourth to second orbit, it emits visible light photons of 488 nanometer wavelengths which are greenish blue. From the fifth to second orbit, it’s 435 nanometers, which is deep blue. Sixth to second orbit, it’s 412 nanometers, which is a violet color. With its energy difference, we can calculate all other possibilities, but these are only
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visible light photons emitted by hydrogen atoms. That’s why we get these four colors of visible light line spectrum when electromagnetic radiation passes through hydrogen atoms. It means that orbits are quantized, the energy is quantized, the velocity of an electron is quantized, the position of an electron is quantized and the radius of the electron’s orbit is quantized. This is the Bohr atomic model of the hydrogen atom. Even though
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he couldn’t explain atoms with more than one electron, this model was easier to understand. In 1917, protons were discovered by Ernest Rutherford. He also predicted that atoms should contain neutral particles which are responsible for a certain amount of atomic mass, but he failed to prove this. However, his student discovered them several years later. In 1932, the neutrally charged particle — the neutron — was discovered by Sir James Chadwick. So, all subatomic particles
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were discovered. An atom contains a nucleus, which is the combination of protons and neutrons, with electrons orbiting around the nucleus. Protons have positive charge, electrons have negative charge, and neutrons have no charge. The number of protons determines the kind of atom. In a neutral atom, the number of protons is equal to the number of electrons so that the charge is balanced out.
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Let’s take a hydrogen atom. It contains one proton and one electron called protium. If we add one neutron to it, nothing changes — it is still a hydrogen atom. But now the atomic mass doubles, and we call this variant of hydrogen “deuterium”. If we add one more neutron it becomes tritium, but still, it is hydrogen. These extra neutron atoms are called isotopes. All are hydrogen atoms, but with various masses. Even though it has different atomic mass,
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the chemical properties are still the same. But what about if we add extra protons? To balance out the charge, an extra electron from somewhere will come into it. Now, the atom gets completely different chemical properties and becomes a different element entirely. In this case, it becomes a helium atom. If we add one more proton and electron, it becomes lithium. A basic carbon atom always has six protons, six electrons and six neutrons. If we add together
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the number of protons and neutrons, we get a total of 12, so it is called carbon-12. An atom with six protons and seven neutrons is called carbon-13. An atom with six protons and eight neutrons called carbon-14. From here we can notice that all types of carbon atoms always have six protons. Of course, without six protons it can’t be a carbon atom. Even though the number of neutrons doesn’t look important, let’s look at a uranium atom.
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If a uranium atom contains 92 protons and 143 neutrons, it is called uranium-235. It is a very very rare isotope which is highly unstable and decays very quickly. If we add 3 more neutrons, it becomes uranium-238, but its lifetime is 4.5 billion years. These are very useful to understand the structure and behavior of atoms. Bohr’s model explains how electrons are orbiting around the nucleus, but the dual behavior of matter and the Heisenberg uncertainty principle
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questioned the Bohr model of the atom and so an even more suitable atomic model was required. In 1924, French physicist De Broglie proposed that every matter should also exhibit dual behavior like electromagnetic radiation, meaning both particle and wavelike properties. According to him, every object in motion has a wave character but differences depend on the mass. Even we humans have wave behavior, but the wavelength is too short which cannot be detectable. The subatomic
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particles are tiny and too small in mass, so their wave natures are experimentally detectable. In 1927, German physicist Warner Heisenberg proposed the uncertainty principle. It states that, it is impossible to determine simultaneously the exact position and exact velocity of an electron. For example, if we want to detect something, light is the key, and light means photons. But when a photon collides with an electron, it would change
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the energy of the electron. We already saw that in Bohr's model, when an electron absorbs a photon, it moves into a different energy state. In this scenario we would be able to calculate the position of the electron, but this will affect the velocity of the electron. The principle is only applicable for the motion of microscopic objects and negligible for macroscopic objects. Bohr’s atomic model determined the exact velocity and position of the electron. But
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the Heisenberg uncertainty principle is true, and it rules out the existence of a definite fixed path of electrons. So, according to the wave nature of matter and uncertainty principle, scientists introduced a new concept of electron orbit, that is ‘probability’. Classical mechanics, based on Isaac Newton's laws of motion, successfully describes the motion of all macroscopic objects, however it fails when applied to microscopic objects like electrons
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and atoms. The classical mechanical model ignores the wave-particle dual nature of matter, especially subatomic particles and the uncertainty principle. A branch of a new science called quantum mechanics emerged to account for the dual behavior of matter. Quantum mechanics was developed independently by Werner Heisenberg and Erwin Schrödinger in 1926. Quantum mechanics is, at its basis, founded on complex mathematical explanations. So, let’s dive into the quantum model of atoms.
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We can measure the location of a particle, but when it is in a wave form, how can we measure that location? If an electron is a wave, it doesn’t exist in a particular location, as a single electron is spread out throughout the entire atom. But how does an electron spread out through an entire atom? We already know that every matter in motion has a wave nature but it depends on their mass. For easy understanding, let’s take the De Broglie
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wave behavior formula : Lambda = Planck’s constant divided by mass times velocity. Here, Consider a baseball which travels 40 meters per second, and its weight 0.15 kgs. If we want to find out the wave nature of a baseball, we just need to put its values in this formula. So, we will get a wave nature of baseball that is just 10 ^ -34 meters. At this level, it is simply undetectable. But for an electron, it will be 10 ^ -10 meters because electrons
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have a mass of 9.1 * 10 ^ -31 kgs. Since the size of a hydrogen atom is 10 ^ -10 meters and it has only one electron, an electron can spread out throughout the entire atom. So, according to the quantum model, electrons orbiting the nucleus in a waveform make a cloud shape called electron clouds. Within that cloud, an electron can be anywhere. But at the time of your measurement, only you can know about
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the exact location of that electron. While trying to make measurements of an electron’s position, its wave nature collapses and becomes a particle because you disturbed its motion. Due to the Heisenberg uncertainty principle, it is impossible to know for a given electron both its position and velocity. Since electron velocity is related to electron energy, it is very necessary to find its velocity and position at the same time. But nature has prevented us from finding
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the exact location and speed of a subatomic particle at the same time. So, scientists come to the conclusion that, since we can’t find its exact location, we can make a probability prediction. If we understand what ‘probability’ is, we can understand the state of subatomic particles. For example, lock your friend in a dark room where a circle is drawn in the center of the room. Since it’s dark, he can’t see where the circle is. When you open the door, you notice
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that he was standing near the circle, the next time he was standing in the corner of the room, and each time you noted his location. After a hundred of these observations, you notice that most of the time he was standing near the circle. So, his average position is approximately around and near the circle. It is similar to an electron’s position in an atom. Schrodinger’s equation tells us that when you measure the position of an electron, most of the time it is
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observed at a particular distance from the nucleus. According to Niels Bohr, the first orbit radius is 0.53 Angstroms. The Schrodinger equation also says that 90 percent of the time, an electron can be measured at 0.53 Angstroms, but the remaining 10 percent of time it can be either below or above 0.53 Angstroms. This probability is called orbital. It is not a fixed orbit as Bohr said. It’s probability, and an electron can be anywhere
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around the nucleus at the same time, because of its wave nature. It means that finding an electron’s exact location is impossible. But the probability says that 90 percent of the time, an electron is to be a particular radius from the nucleus but can be found anywhere in this radius. We can only know an electron’s location when measuring it. But before measurement, it can be anywhere around the nucleus. This model is called the quantum model of an atom.
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Even though we can’t tell the electron's exact location, it is believed to be the final atomic model because this is a restriction of nature and humans can’t do anything with it. This is the beauty of the quantum world.
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