On its most basic level, you can think of the periodic table as a simple listing of all of the different types of atoms we have discovered or created. As I write this, there are currently 118 elements listed in the periodic table. [An element is a pure substance made up of a single type of atom. Water (H2O) is not an element because it is made up of both oxygen and hydrogen atoms. Hydrogen is an element because it is made up of only a single type of atom… hydrogen atoms. Basically, each element corresponds to a type of atom.]
Elements (or atoms) are listed in the periodic table according to their atomic number. The atomic number is the number of protons an atom has in its nucleus. Remember, an atom is identified based on its number of protons.
|atom/element||atomic number||mass number||number of protons||number of neutrons||number of electrons||atomic mass (u)|
Copper has an atomic number of 29 and a mass number of 63. The atomic number is the number of protons in the nucleus, and the mass number is the number of protons and neutrons in the nucleus. This means that the copper atom has 29 protons and 34 neutrons in its nucleus (29 + 34 = 63). A neutrally charged copper atom will then have 29 electrons in orbit around its nucleus (29 − 29 = 0). Because almost all of an atom’s mass is in its nucleus, we can estimate an atom’s mass just by counting its protons and neutrons. Both protons and neutrons have masses approximately equal to 1 u (1 proton ≈ 1.0073 u, 1 neutron ≈ 1.0087 u, and 1 u ≈ 1.66 × 10-27 kg), so the mass of a copper atom will be approximately 63 u, and the mass of 1 mole (6.022 × 1023 atoms) of copper will have a mass of approximately 63 g.
I am being imprecise again. There is not just one type of copper atom. Copper can exist as 29 different isotopes, and two of those isotopes are stable. Copper-63 (63Cu) is the most common isotope. About 69% of the copper atoms on Earth are copper-63 (this percentage may vary in other solar systems and galaxies). It has a mass number of 63 (29 protons and 34 neutrons) and an atomic mass of 62.9296 u. Most of the remaining 31% of copper atoms on Earth are copper-65 (65Cu). It has a mass number of 65 (29 protons and 36 neutrons) and an atomic mass of 64.9278 u. Copper-65 has a greater mass than copper-63 because it has two more neutrons in its nucleus. One mole of copper will therefore have a combination of both copper-63 atoms and copper-65 atoms in it, and the “molar” mass of copper is considered to be 63.546 g/mol. However, the actual mass of a mole of copper will depend on the composition of the specific sample.
It is easy to get lazy and ignore the fact that most types of atoms can exist as several different isotopes because one of those isotopes is usually far more common than the others. 99.99% of the hydrogen atoms on Earth are 1H, so when we talk about the “hydrogen” atom, we usually mean 1H. 98.94% of the carbon atoms on Earth are 12C, 99.64% of the nitrogen atoms are 14N, 99.76% of the oxygen atoms are 16O, 75.76% of the chlorine atoms are 35Cl, 96.94% of the calcium atoms are 40Ca, and 91.75% of the iron atoms are 56Fe. All of the other elements listed in the table above have isotopes that are even more common (practically 100%). Copper and chlorine are slightly unusual in that they both have two relatively common isotopes.
The first modern periodic table was created in 1869. It listed 63 elements. The lightest element was hydrogen (atomic number 1) and the heaviest element was uranium (atomic number 92). There were many gaps in the table where individual elements had not yet been discovered, including fluorine (atomic number 9), neon (atomic number 10), argon (atomic number 18), krypton (atomic number 36), xenon (atomic number 54), radon (atomic number 86), and radium (atomic number 88). By 1940, all of the elements from atomic number 1-92, except promethium (atomic number 61) had been discovered and identified.
Nuclear research in the 1940s (including the Manhattan Project) that led to the first atomic bombs produced the first man-made elements. These heavier elements were created by bombarding atomic nuclei with protons and neutrons in nuclear reactors or particle accelerators (machines that use electromagnetic fields to accelerate particles close to the speed of light). Originally, it was believed that uranium was the heaviest known atom that occurred naturally, but it was later discovered that neptunium and plutonium (both originally synthesized and detected in the lab) also occurred naturally but in very small quantities. (Elements that occur “naturally” are created by nucleosynthesis in a star, during supernovas, or when natural but unstable isotopes decay radioactively.)
By 1996, the first 112 elements had been identified. Copernicium (atomic number 112) was finally named in February 2010. The elements 113-118 have all been synthesized and detected in the lab, but so far, only the discoveries of ununquadium (atomic number 114) and ununhexium (atomic number 116) have been officially confirmed in June 2011. Neither of those elements has been named by their discoverers yet.
As you have learned, the electrons in an atom must have specific and discrete energy levels. The lowest energy level that an electron can have is n = 1. The next lowest energy level is n = 2. An electron cannot have an energy level of n = 1.5. (This is a bit like saying that a cat can weigh 9 lb or 10 lb, but not 9.5 lb!) Although this may seem very, very strange (perhaps even impossible), it accurately describes the behavior of atoms and is one of the key differences between classical mechanics and quantum mechanics.
|energy level||electron shell||maximum number of electrons||number of atomic orbitals|
|n = 1||1||2||1|
|n = 2||2||8||4|
|n = 3||3||18||9|
|n = 4||4||32||16|
|n = 5||5||50||25|
|n = 6||6||72||36|
Only a specific number of electrons can exist with a specific energy level. For example, only two electrons can have an energy level of n = 1 and only eight electrons can have an energy level of n = 2. Sometimes it is convenient to think of the electrons with the same energy level as sharing an “electron shell” in the electron cloud surrounding an atom’s nucleus. The technical reason why the first electron shell can only hold two electrons is because only two quantum states can have an energy level of n = 1. But that reason is very abstract unless you have a deep understanding of quantum mechanics. There is a classical mechanical explanation I can give for why only a specific number of electrons can have the same energy level, but it is not the complete reason.
The higher the energy level of an electron, the farther away it is likely to be from the atom’s nucleus. Since electrons in the first electron shell have the lowest energy level (n = 1), they will be closest to the nucleus. However, while negatively charged electrons are strongly attracted to the positively charged nucleus, they are also strongly repelled by each other. This means that electrons will push each other away if they get too crowded together. Trying to add a third electron to the first electron shell would make the electron shell too crowded, so the electron gets pushed into the second electron shell with a higher energy level (n = 2). The second electron shell can hold more electrons before it gets too crowded because it is farther away from the nucleus and has more physical space to fill.
One way to represent the structure of an atom is with a Bohr diagram.
atomic number: 13
mass number: 27
neutrons: 27 − 13 = 14
While a Bohr diagram is easy to draw and conveys useful information about an atom’s structure, it can reinforce certain misconceptions. First, electrons do not orbit an atom’s nucleus in concentric circles. This is a holdover from when we used to model electrons like planets in a solar system. Second, electrons do not have specific positions at any given time. And third, electrons do not occupy electron shells. Electron shells are a construct that we created to help us think about atoms. To really understand atoms and atomic behavior, you need to think in terms of atomic orbitals.
The thirteenth electron of an aluminum atom actually occupies an atomic orbital that looks something like this. This is known as a “p-orbital.” This particular orbital has a quantum state of n = 2, l = 1, and m = 0; and is shaped a bit like a dumbbell, with one region in front of the nucleus and a second region behind the nucleus. Up to two electrons can occupy any one atomic orbital: one electron with a spin of s = +½ and one electron with a spin of s = -½. There is one atomic orbital with an energy level of n = 1, four atomic orbitals with an energy level of n = 2, and nine atomic orbitals with an energy level of n = 3 (see the table above). Atomic orbitals are sometimes referred to as “subshells.”
Having a detailed understanding of atomic orbitals is not important at this point, but it is important to recognize that the Bohr diagram grossly oversimplifies the structure of an atom and that the Bohr model of an atom (with electrons traveling in concentric, circular orbits) was replaced by a more modern model of an atom based firmly on quantum mechanics decades ago. Having some familiarity with atomic orbitals will also help you better understand the structure of the periodic table and valence electrons later in this unit.
Electrons in an atom can transition to a different energy level by gaining or losing energy, as long as there is room at that energy level for another electron. (Remember, the number of electrons that can have any given energy level is limited because no two electrons can share the same quantum state.) In general, electrons will occupy the lowest energy levels available. This is why, when drawing our Bohr diagram of the aluminum atom, we started by completely filling the first electron shell before moving on to the second electron shell. This electron configuration is called the ground state of the atom, and it represents the lowest energy state an atom can have.
But an electron does not have to occupy the lowest energy level available. If an electron absorbs the energy from a photon, it can get kicked up to a higher energy level even though there is still room at a lower energy level. In this Bohr diagram of an aluminum atom, there are five electrons with an energy level of n = 3 even though there is still enough room in the second electron shell for two more electrons with a lower energy level of n = 2. When an atom is not in the ground state, it is said to be in an excited state.
It takes energy to move an electron from a low energy level to a higher energy level. And with enough energy, an electron can be completely removed from an atom. This corresponds to an energy level of n = ∞ (infinity). The energy required to remove an electron from an atom is called the ionization energy. Removing an electron turns an atom into an ion because the number of protons and electrons is no longer equal and the atom now has a positive charge. It takes 9.59 × 10-19 J of energy to remove the thirteenth electron from an aluminum atom, turning it into an ion (Al+). Removing a second electron, turning the atom into the ion Al2+, then takes even more energy (3.02 × 10-18 J).
Fire is the rapid oxidation of a material through chemical reactions known as combustion. Combustion releases heat. This thermal energy can excite the atoms in the fire, causing electrons to jump to higher energy levels. (If the fire is hot enough, matter can even transition to a plasma state if electrons get completely separated from atomic nuclei.) Each time an electron drops from an excited state back down to the ground state, it releases energy in the form of a photon (light or energy particle). The amount of energy released by the transition from a high energy level down to a lower energy level determines the frequency and wavelength of the photon. The more energy, the higher the frequency and the shorter the wavelength (blue light has a wavelength of about 450 nm and red light has a wavelength of about 700 nm). This is where firelight comes from.
Imagine that the thirteenth electron of an aluminum atom has a quantum state of n = 2, l = 1, m = -1, and s = +½; and that this quantum state has an energy of E1. Now imagine that the unoccupied quantum state of n = 3, l = 0, m = 0, and s = -½ has an energy of E2. In order for the thirteenth electron to jump from the first state to this second state, it must absorb a photon with exactly E2 − E1 energy (the energy difference between the two states).
The thirteenth electron cannot absorb a photon with less energy. Absorbing a photon with less energy would give the electron too much energy to stay in the first state, but too little energy to reach the second state. The thirteenth electron cannot absorb a photon with more energy either. It would have too much energy to stay in the first state or stay in the second state. To jump into the second state, the electron must have the exact right amount of energy, E2. (Quantum mechanics is so picky!) Of course, the electron could jump into some third or fourth state if the energy of the photon was just right; after all, the second state is not the only unoccupied quantum state available. But an electron cannot be between states.
The frequency and wavelength of a photon determines the amount of energy it has. So the thirteenth electron of an aluminum atom in the first quantum state can only absorb light with specific frequencies and wavelengths. And the other electrons in the aluminum atom can only absorb light with other specific frequencies and wavelengths. Scientists can analyze the light absorbed by atoms to identify the type of atom. This is called atomic spectroscopy.
The first spectrum shows the light absorbed by an atom. The black lines are absorption lines. Those wavelengths correspond to the energy of photons that can be absorbed by electrons in the atom, enabling them to jump from the ground (lowest energy) state to an excited (higher energy) state. The second spectrum shows the light emitted by the same atom. The colored lines are emission lines. Those wavelengths correspond to energies that are emitted by electrons in the atom when they drop from an excited state back down to the ground state. Notice that the energy that an electron needs to absorb to jump to an excited state is the same amount of energy it releases when it drops back down to the ground state.
The atomic line spectrum of each atom is unique. It is basically like the fingerprint of an atom. Not only does atomic spectroscopy give scientists a powerful tool for identifying atoms (a similar technique is also used for molecules), but the fact that the atomic orbital model accurately predicts these atomic line spectra provides evidence of the quantum mechanical nature of the atom and its subatomic particles.
We have just touched the surface of what the periodic table can tell us. Next, you will learn how chemical properties of an atom can be predicted from the atom’s structure and how the periodic table is organized to group elements with common properties together.