Tuesday, November 27, 2012

Atoms Part 4A: Atoms and Chemistry – Atomic Orbitals and Bonding

In Atoms Part 2 and Part 3, we explored how the electrons in atoms emit photons of electromagnetic radiation by comparing different kinds of spectra from hydrogen atoms and from the Sun, essentially a huge ball of hydrogen.

We're not done with electrons in atoms - not only do they change energy states by moving into different orbitals, they interact with electrons in other nearby atoms. In a nutshell, this is what chemistry is.

Chemistry is All About Bonds

You may have heard of two basic kinds of bonds: ionic and covalent.  We will explore and compare them in the next article. They both achieve the same goal: atoms bond with each other because they want to find the lowest possible energy state. Chemical bonding helps atoms achieve this goal. In order to understand bonding, we first need to look at the atoms themselves and how their electrons are configured, and that is what we're about to do here.


Orbitals: Their Shapes and Their Electrons

In atoms, electrons occupy orbitals. An orbital is a cloud-like area around the atom's nucleus where an electron is likely to be found. To give you an idea of what these shapes look like, the first five atomic orbital shapes are shown below.

From left to right they are the 1s, 2s, 2px, 2py and 2pz orbitals. The 2px, 2py and 2pz are actually suborbitals, subtypes of the 2p orbital. All this will become clearer as we go on.

Orbitals play a crucial role in how atoms bond with each other. There is an orbital rule every atom must obey - only two electrons can occupy the same orbital, and only if their spins are opposite. So, the 1s and 2s orbitals can fit two electrons in them, and the 2px, 2py and 2pz orbitals, likewise, each can only fit in two electrons.

This rule is based on the Pauli exclusion principle, which says that no two electrons in an atom can share all four of the same quantum numbers. Quantum numbers are part of the quantum mechanical model that describes the atom. These numbers describe the energy, angular momentum, magnetic moment and spin of electrons. The shape of each electron orbital depends on the electron's angular momentum. The energy of the electron determines which orbital it occupies. The electron's magnetic moment is behind all magnetic phenomena. That leaves spin, and electrons can choose one of two possible spins. This means that two, and only two, electrons can share the same orbital (or suborbital) because each electron can have one of two possible quantum spins. All the other quantum numbers of these two electrons will be the same.

You can quickly find out how many electrons any atom has by looking at the periodic table:
The number in each square is the atomic number. That's the number of electrons (and protons) in each atom. Hydrogen [H] has only one electron while chlorine [Cl] has 17 electrons, for example.

Hydrogen just has one (ground state) orbital, 1s1. The superscript "1" tells you the number of electrons occupying the orbital, in this case just one. Lithium [Li] has three electrons so it fills up the 1s orbital (this orbital can fit two electrons) and sets one electron in the next highest energy 2s shell: 1s22s1. Phosphorus [P] with 15 electrons is written as 1s22s22p63s23p3. Notice that the 2p shell allows six electrons in it, while the 3s orbital allows 2 electrons. Remember that the 2p orbital is actually made up of three suborbitals (take a quick look at the previous orbital diagram, at the three suborbitals to the right). Each of these three suborbitals holds two electrons, one of each spin. The suborbitals keep the electron pairs separate. A way to visualize the p orbital is shown below. Each p suborbital (a separate colour) has two lobes and it lines up along one of three axes:
This is what a typical p orbital looks like, whether it is 2p, 3p, 4p and so on. As the number in front increases, the p orbital's energy increases, so a 3p orbital has more energy than a 2p orbital, for example. The overall shape stays the same. The atomic nucleus is in the middle where the lobes intersect.

Now let's compare the p orbitals of two atoms. Argon [Ar], a noble gas atom, has 18 electrons (1s22s22p63s23p6). Its outermost 3p orbital is full; it has 6 electrons in it. Another atom, phosphorus has 15 electrons (1s22s22p63s23p3). It has a 3p orbital that's only half filled, with three electrons. Argon's filled orbitals make it very stable. The more stable an atoms is, the lower its potential energy is. This allows argon to exist in a very low energy state and it likes to stay that way. It is a very nonreactive, or inert, atom. The half-filled 3p orbital in phosphorus, on the other hand, makes its electron configuration a bit less stable. This atom is reactive; it can adopt a lower energy argon-like stability if it attracts three electrons to itself. It can do this by forming bonds. This is how atomic orbitals make chemistry possible.

The d orbital is higher energy yet. It can hold up to ten electrons. It holds a maximum of two (opposite spin) electrons in each of five suborbitals. This orbital looks a bit like a fancy three-dimensional daisy, shown bottom middle, below:

Electron Orbitals Versus Electron Shells

You have probably heard of electron shells and valence shells before. For example, argon, a noble gas, has a full valence shell of eight electrons.

Talking about orbitals and shells can get a bit confusing, so let's compare the two. If you look at a typical Bohr diagram of argon, shown below, you will see rings, or shells as they are called, representing specific energies at which electrons may be found.

(commons:User:Pumbaa (original work by commons:User:Greg Robson); Wikipedia)

The maximum number of electrons per shell depends on the shell. The first or closest shell can hold two electrons. The second one can hold eight, the third can hold 18, the fourth can hold 32, and so on. Inner shells generally must be filled before outer ones. The outermost shell of electrons is called the valence shell. Notice that argon has eight electrons in its third shell, which is its valence shell. At first glance it looks like argon's valence shell is less than half full. It has eight out of 18 possible electrons in it. Shells consist of subshells. Subshells and orbitals are the same thing. So, a p subshell is a p orbital (and remember the p orbital can be further divided into three suborbitals). The number in front of the subshell indicates what the electron energy is. A 3p subshell belongs to shell n = 3 in a Bohr diagram. Argon has eight n = 3 shell electrons. The three p suborbitals account for six electrons in total, so where do the other two come from? From the 3s orbital - it also has n = 3 energy, and being an s orbital, it contains two electrons when filled.

Argon has eight electrons in its valence shell (n = 3), which consists of two subshells or orbitals: 3s and 3p. The 3s (two electrons) and 3p (six electrons) subshells make up the n = 3 electron shell (eight electrons). As I said, the n = 3 shell can hold even more electrons – 18 in total. There is another orbital available to electrons with n = 3 energy: the d orbital. The 3d orbital can hold up to ten electrons in it. Take another quick look at the daisy-shaped d orbital diagram above. Argon has no electrons in this orbital. Its 3s and 3p orbitals, however, are full and that makes it stable.

Transition Metals: An Exception to the Orbital-Filling Scheme

There are a few exceptions to the general orbital filling scheme. One exception you might come across is the transition metal group. Many gems such as rubies and sapphires, owe their brilliant colours to the uniquely spaced electron orbitals of transition metal atoms inside them. It is also why transition metals can form a variety of different oxidation states. For example, Iron [Fe] can form both Fe2+ and Fe3+ ions when it reacts with oxygen to form iron (II) oxide and iron (III) oxide. Vanadium [V] has even more oxidation states – four, each with its own attractive colour in solution.

Nickel [Ni] is a transition metal atom with ten more electrons than argon. It does not just go ahead and fill up the 3d orbital, however. It fills up the 4s orbital first.

(commons:User:Pumbaa (original work by commons:User:Greg Robson); Wikipedia)

In transition metals, the inner 3d orbital generally has more energy than the valence shell 4s orbital.

Noble Gas Orbital Writing Shortcut

A shortcut to writing out long orbital notations is to use a noble gas core. Noble gases (group 18 on the periodic table) have completely filled orbitals. Our example argon is one of them. A shortcut for phosphorus, for example, uses the next smallest noble gas core. Phosphorus has 15 electrons, five of which are valence electrons, confined in two more orbitals of electrons than neon has. Neon [Ne], with 10 electrons, is written as 1s22s22p6, so for phosphorus we can write [Ne]3s23p3. Nickel, above, can be written as [Ar]4s2 3d8.

Halogens (group 17 on the periodic table) are just the opposite of the noble gases. They have an outer shell of seven electrons rather than the stable eight of the noble gases. These elements want just one more electron very badly and that makes them highly reactive. They gain this electron by reacting with other elements.

Fluorine [F], shown as a Bohr diagram below, is a great example.

It's the most reactive element there is. It will attack even inert materials like glass, and it will even form compounds with the heavier noble gases, which are generally nonreactive. Once fluorine bonds with something, that bond is so strong that almost nothing can break it. This makes Teflon so perfect as a non-stick non-reactive coating. It's made of carbon bonded with fluorine.

Now that we understand what makes some atoms more stable than other ones, we can explore how atoms increase their overall stability by bonding with each other, in Atoms Part 4B – Ionic and Covalent Bonds.


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