Introduction to Structure of Atom An atom is present at the most basic level in everything we see around us. In fact, every living organism is composed of atoms. Every non-living thing around you such as tables, chairs, water, etc is made up of matter. But the building blocks of matter are atoms. Therefore, living or non-living, everything is composed of atoms. Let us take a look at the structure of atom. Atoms Atom is a Greek word which means “indivisible.” The Greeks believed that matter can be broken down into very small invisible particles called atoms. Greek philosophers such as Democritus and John Dalton put forward the concept of the atom. Democritus explained the nature of matter. He also proposed that all substances are made up of matter. He stated atoms are constantly moving, invisible, minuscule particles that are different in shape, size, temperature and cannot be destroyed. Atomic Number Elements are the building blocks for all the matter you see in the world. Now, the question is how can you distinguish these elements? The answer is “Atomic Number.” Every element has their unique atomic number that helps to distinguish between two different elements. Let us study further to know about its significance. Atomic Number We know that an atom consists of electrons, protons, and neutrons. Atomic number is one of the fundamental properties of an atom. Each atom can be characterized by a unique atomic number. It is represented by the letter “Z.” The total number of protons present in the nucleus of an atom represents the atomic number of a particular atom. Every atom of a particular element is composed of the same number of protons and therefore has the same atomic number. However, atoms of different elements have unique atomic numbers that vary from one element to the other. An atom does not have any net charge and is thus electrically neutral. This means that the number of electrons will be equal to the number of protons present in an atom thereby making an atom electrically neutral. “Atomic Number = No. of Protons = No. of Electrons” For example, each atom of oxygen has 8 protons in their nucleus so the atomic number is 8. Similarly, the atomic number of carbon is 6 because all atoms of carbon have 6 protons in their nucleus. Atomic numbers are whole numbers because it is the total number of protons and protons are generally units of matter. It ranges from 1 to 118. It starts with hydrogen and ends with the heaviest known element Oganesson (Og). Theoretically, the atomic numbers can be increased if more elements are discovered. However, with the addition of more number of protons and neutrons, the elements become prone to radioactive decay. Later in the year 1808, John Dalton proposed the atomic theory and explained the law of chemical combination. By the end of 18th and the early 20th centuries, many scientists such as J.J Thomson, Gold stein, Rutherford, Bohr among others developed and proposed several concepts on “atom.” Atom is the smallest unit of matter that is composed of a positively charged centre termed as “nucleus” and the central nucleus is surrounded by negatively charged electrons. Even though an atom is the smallest unit of matter but it retains all the chemical properties of an element. For example, a silver spoon is made up of silver atoms with few other constituents. A silver atom obtains its properties from tiny subatomic particles that it is composed of. Atoms are further arranged and organized to form larger structures known as molecules. Atoms, molecules adhere to the general chemistry and physics rule even when they are part of living breathing human body. Let us study atoms and molecules to further understand how the atoms react, behave and interact in cells form an essential part of the living and nonliving world. Rutherford Atomic Model Rutherford proposed that an atom is composed of empty space mostly with electrons orbiting in a set, predictable paths around fixed, positively charged nucleus. Rutherford’s Atomic Model (Source Credit: Britannica) History The concept of atom dates back to 400 BCE when Greek philosopher Democritus first conceived the idea. However, it was not until 1803 John Dalton proposed again the idea of the atom. But at that point of time, atoms were considered indivisible. This idea of an atom as indivisible particles continued until the year 1897 when British Physicist J.J. Thomson discovered negatively charged particles which were later named electrons. He proposed a model on the basis of that where he explained electrons were embedded uniformly in a positively charged matrix. The model was named plum pudding model. However, J.J. Thomson’s plum pudding model had some limitations. It failed to explain certain experimental results related to the atomic structure of elements. A British Physicist “Ernest Rutherford” proposed a model of the atomic structure known as Rutherford’s Model of Atoms. He conducted an experiment where he bombarded αparticles in a thin sheet of gold. In this experiment, he studied the trajectory of the αparticles after interaction with the thin sheet of gold Bohr's model of hydrogen Key points Bohr's model of hydrogen is based on the non-classical assumption that electrons travel in specific shells, or orbits, around the nucleus. Bohr's model calculated the following energies for an electron in the shell, n: 1 𝐸𝑛 = − 2 𝑋13.6𝑒𝑉 𝑛 Bohr explained the hydrogen spectrum in terms of electrons absorbing and emitting photons to change energy levels, where the photon energy is ℎ𝑣 = ∇𝐸 = ( 1 𝑛𝑖2 − 1 𝑛𝑗2 ) 𝑋13. 𝑒𝑉 (Where 𝑛𝑖2 is for low energy and 𝑛𝑗2 is for high energy). Bohr's model does not work for systems with more than one electron. The planetary model of the atom At the beginning of the 20th century, a new field of study known as quantum mechanics emerged. One of the founders of this field was Danish physicist Niels Bohr, who was interested in explaining the discrete line spectrum observed when light was emitted by different elements. Bohr was also interested in the structure of the atom, which was a topic of much debate at the time. Numerous models of the atom had been postulated based on experimental results including the discovery of the electron by J. J. Thomson and the discovery of the nucleus by Ernest Rutherford. Bohr supported the planetary model, in which electrons revolved around a positively charged nucleus like the rings around Saturn—or alternatively, the planets around the sun. Image of Saturn and rings Many scientists, including Rutherford and Bohr, thought electrons might orbit the nucleus like the rings around Saturn. Image credit: Image of Saturn by NASA However, scientists still had many unanswered questions: Where are the electrons, and what are they doing? If the electrons are orbiting the nucleus, why don’t they fall into the nucleus as predicted by classical physics? How is the internal structure of the atom related to the discrete emission lines produced by excited elements? Bohr addressed these questions using a seemingly simple assumption: what if some aspects of atomic structure, such as electron orbits and energies, could only take on certain values? Quantization and photons By the early 1900s, scientists were aware that some phenomena occurred in a discrete, as opposed to continuous, manner. Physicists Max Planck and Albert Einstein had recently theorized that electromagnetic radiation not only behaves like a wave, but also sometimes like particles called photons. Planck studied the electromagnetic radiation emitted by heated objects, and he proposed that the emitted electromagnetic radiation was "quantized" since the energy of light could only have values given by the following equation: E_{\text{photon}}=nh\nuEphoton=nhν, where nnn is a positive integer, hhh is Planck’s constant—6.626 \times10^{-34}\,\text {J}\cdot \text s6.626×10−34J⋅s6, point, 626, times, 10, start superscript, minus, 34, end superscript, space, J, dot, s—and \nuν is the frequency of the light, which has units of \dfrac{1}{\text s}s1start fraction, 1, divided by, s, end fraction. As a consequence, the emitted electromagnetic radiation must have energies that are multiples of h\nuhν. Einstein used Planck's results to explain why a minimum frequency of light was required to eject electrons from a metal surface in the photoelectric effect. When something is quantized, it means that only specific values are allowed, such as when playing a piano. Since each key of a piano is tuned to a specific note, only a certain set of notes—which correspond to frequencies of sound waves—can be produced. As long as your piano is properly tuned, you can play an F or F sharp, but you can't play the note that is halfway between an F and F sharp. Atomic line spectra Atomic line spectra are another example of quantization. When an element or ion is heated by a flame or excited by electric current, the excited atoms emit light of a characteristic color. The emitted light can be refracted by a prism, producing spectra with a distinctive striped appearance due to the emission of certain wavelengths of light. Emission spectra of sodium, top, compared to the emission spectrum of the sun, bottom. The dark lines in the emission spectrum of the sun, which are also called Fraunhofer lines, are from absorption of specific wavelengths of light by elements in the sun's atmosphere. The side-by-side comparison shows that the pair of dark lines near the middle of the sun's emission spectrum are probably due to sodium in the sun's atmosphere. Image credit: From the Biodiversity Heritage Library For the relatively simple case of the hydrogen atom, the wavelengths of some emission lines could even be fitted to mathematical equations. The equations did not explain why the hydrogen atom emitted those particular wavelengths of light, however. Prior to Bohr's model of the hydrogen atom, scientists were unclear of the reason behind the quantization of atomic emission spectra. Bohr's model of the hydrogen atom: quantization of electronic structure Bohr’s model of the hydrogen atom started from the planetary model, but he added one assumption regarding the electrons. What if the electronic structure of the atom was quantized? Bohr suggested that perhaps the electrons could only orbit the nucleus in specific orbits or shells with a fixed radius. Only shells with a radius given by the equation below would be allowed, and the electron could not exist in between these shells. Mathematically, we could write the allowed values of the atomic radius as r(n)=n^2\cdot r(1)r(n)=n2⋅r(1)r, left parenthesis, n, right parenthesis, equals, n, start superscript, 2, end superscript, dot, r, left parenthesis, 1, right parenthesis, where nnn is a positive integer, and r(1)r(1)r, left parenthesis, 1, right parenthesis is the Bohr radius, the smallest allowed radius for hydrogen. He found that r(1)r(1)r, left parenthesis, 1, right parenthesis has the value \text{Bohr radius}=r(1)=0.529 \times 10^{-10} \,\text{m}Bohr radius=r(1)=0.529×10−10mB, o, h, r, space, r, a, d, i, u, s, equals, r, left parenthesis, 1, right parenthesis, equals, 0, point, 529, times, 10, start superscript, minus, 10, end superscript, space, m An atom of lithium shown using the planetary model. The electrons are in circular orbits around the nucleus. Image credit: planetary atomic model from Wikimedia Commons, CC-BY-SA 3.0 By keeping the electrons in circular, quantized orbits around the positively-charged nucleus, Bohr was able to calculate the energy of an electron in the nnnth energy level of hydrogen: E(n)=-\dfrac{1}{n^2} \cdot 13.6\,\text{eV}E(n)=−n21⋅13.6eVE, left parenthesis, n, right parenthesis, equals, minus, start fraction, 1, divided by, n, start superscript, 2, end superscript, end fraction, dot, 13, point, 6, space, e, V, where the lowest possible energy or ground state energy of a hydrogen electron—E(1)E(1)E, left parenthesis, 1, right parenthesis—is -13.6\,\text{eV}−13.6eVminus, 13, point, 6, space, e, V. Note that the energy is always going to be a negative number, and the ground state, n=1n=1n, equals, 1, has the most negative value. This is because the energy of an electron in orbit is relative to the energy of an electron that has been completely separated from its nucleus, n=\inftyn=∞n, equals, infinity, which is defined to have an energy of 0\,\text{eV}0eV0, space, e, V. Since an electron in orbit around the nucleus is more stable than an electron that is infinitely far away from its nucleus, the energy of an electron in orbit is always negative. Absorption and emission The energy level diagram showing transitions for Balmer series, which has the n=2 energy level as the ground state. The Balmer series—the spectral lines in the visible region of hydrogen's emission spectrum—corresponds to electrons relaxing from n=3-6n=3−6n, equals, 3, minus, 6 energy levels to the n=2n=2n, equals, 2 energy level. Bohr could now precisely describe the processes of absorption and emission in terms of electronic structure. According to Bohr's model, an electron would absorb energy in the form of photons to get excited to a higher energy level as long as the photon's energy was equal to the energy difference between the initial and final energy levels. After jumping to the higher energy level—also called the excited state—the excited electron would be in a less stable position, so it would quickly emit a photon to relax back to a lower, more stable energy level. The energy levels and transitions between them can be illustrated using an energy level diagram, such as the example above showing electrons relaxing back to the n=2n=2n, equals, 2 level of hydrogen. The energy of the emitted photon is equal to the difference in energy between the two energy levels for a particular transition. The energy difference between energy levels n_{high}nhighn, start subscript, h, i, g, h, end subscript and n_{low}nlow n, start subscript, l, o, w, end subscript can be calculated using the equation for E(n)E(n)E, left parenthesis, n, right parenthesis from the previous section: \begin{aligned} \Delta E &= E(n_{high})-E(n_{low}) \\ \\ &=\left( -\dfrac{1}{{n_{high}}^2} \cdot 13.6\,\text{eV} \right)-\left(\dfrac{1}{{n_{low}}^2} \cdot 13.6\,\text{eV}\right) \\ \\ &= \left(\dfrac{1}{{n_{low}}^2}-\dfrac{1}{{n_{high}}^2}\right) \cdot 13.6\,\text{eV} \end{aligned}ΔE=E(nhigh)−E(nlow)=(−nhigh21 ⋅13.6eV)−(−nlow21⋅13.6eV)=(nlow21−nhigh21)⋅13.6eV Since we also know the relationship between the energy of a photon and its frequency from Planck's equation, we can solve for the frequency of the emitted photon: \begin{aligned} h\nu &=\Delta E = \left(\dfrac{1}{{n_{low}}^2}\dfrac{1}{{n_{high}}^2}\right) \cdot 13.6\,\text{eV} ~~~~~~~~~~~~\text{Set photon energy equal to energy difference}\\ \\ \nu &= \left(\dfrac{1}{{n_{low}}^2}\dfrac{1}{{n_{high}}^2}\right) \cdot \dfrac{13.6\,\text{eV}}{h}~~~~~~~~~~~~~~~~~~~~~~\text{Solve for frequency}\end{aligned}hνν=ΔE=(nlow21−nhigh21 )⋅13.6eV Set photon energy equal to energy difference=(n low21−nhigh21)⋅h13.6eV Solve for frequency We can also find the equation for the wavelength of the emitted electromagnetic radiation using the relationship between the speed of light \text ccc, frequency \nuν, and wavelength \lambdaλlambda: \begin{aligned}\text c &=\lambda \nu ~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~ ~~~~~~~~~~~~~~\text{Rearrange to solve for }\nu . \\ \dfrac{\text c}{\lambda}&=\nu=\left(\dfrac{1}{{n_{low}}^2}\dfrac{1}{{n_{high}}^2}\right) \cdot \dfrac{13.6\,\text{eV}}{h}~~~~~~~~~~~~~~\text{Divide both sides by c to solve for }\dfrac{1}{\lambda}.\\ \\ \dfrac{1}{\lambda} &=\left(\dfrac{1}{{n_{low}}^2}\dfrac{1}{{n_{high}}^2}\right) \cdot \dfrac{13.6\,\text{eV}}{h\text c} \end{aligned}cλcλ1 =λν Rearrange to solve for ν .=ν=(nlow21−nhigh21)⋅h13.6eV Divide both sides by c to solve for λ1.=(nlow21−nhigh21 )⋅hc13.6eV [What about the Rydberg constant?] RR R=\dfrac{13.6\,\text{eV}}{h\text c}=1.1\times 10^7\,\dfrac{1}{m}R, equals, start fraction, 13, point, 6, space, e, V, divided by, h, c, end fraction, equals, 1, point, 1, times, 10, start superscript, 7, end superscript, space, start fraction, 1, divided by, m, end fraction RR \dfrac{1}{\lambda}=\left(\dfrac{1}{{n_{low}}^2}\dfrac{1}{{n_{high}}^2}\right) \cdot Rstart fraction, 1, divided by, lambda, end fraction, equals, left parenthesis, start fraction, 1, divided by, n, start subscript, l, o, w, end subscript, start superscript, 2, end superscript, end fraction, minus, start fraction, 1, divided by, n, start subscript, h, i, g, h, end subscript, start superscript, 2, end superscript, end fraction, right parenthesis, dot, R RRccE(1)E, left parenthesis, 1, right parenthesishh Thus, we can see that the frequency—and wavelength—of the emitted photon depends on the energies of the initial and final shells of an electron in hydrogen. What have we learned since Bohr proposed his model of hydrogen? The Bohr model worked beautifully for explaining the hydrogen atom and other single electron systems such as \text{He}^+He+H, e, start superscript, plus, end superscript. Unfortunately, it did not do as well when applied to the spectra of more complex atoms. Furthermore, the Bohr model had no way of explaining why some lines are more intense than others or why some spectral lines split into multiple lines in the presence of a magnetic field—the Zeeman effect. In the following decades, work by scientists such as Erwin Schrödinger showed that electrons can be thought of as behaving like waves and behaving as particles. This means that it is not possible to know both a given electron’s position in space and its velocity at the same time, a concept that is more precisely stated in Heisenberg's uncertainty principle. The uncertainty principle contradicts Bohr’s idea of electrons existing in specific orbits with a known velocity and radius. Instead, we can only calculate probabilities of finding electrons in a particular region of space around the nucleus. The modern quantum mechanical model may sound like a huge leap from the Bohr model, but the key idea is the same: classical physics is not sufficient to explain all phenomena on an atomic level. Bohr was the first to recognize this by incorporating the idea of quantization into the electronic structure of the hydrogen atom, and he was able to thereby explain the emission spectra of hydrogen as well as other one-electron systems. Copyright © Michael Richmond. This work is licensed under a Creative Commons License. Heisenberg's Microscope This lecture describes a hypothetical microscope that Werner Heisenberg used to explain his Uncertainty Principle. A good place to learn more about Heisenberg and his ideas is the AIP History of Heisenberg. Copyright © Michael Richmond. This work is licensed under a Creative Commons License.