Where Are the Electrons in Atoms

Experiments on the interactions of light with matter provide important information about the energy and location of electrons in atoms. 

See an animation of Atomic Line Spectra at http //hrookscole.com/ chemist rv/ioesten4 

Bohr assumed that atoms can exist only in certain energy states. 

In 1913, Niels Bohr introduced his model of the hydrogen atom. He proposed that the single electron of the hydrogen atom could occupy only certain energy levels. He referred to these energy levels as orbits and represented the energy difference between any two adjacent orbits as a single 

Continuous spectrum A spectrum that contains radiation distributed over all wavelengths 

---

You know already what makes up an atom—protons, neutrons, and electrons. The protons and neutrons make up the central core of an atom—the nucleus-—while the electrons form some sort of cloud around it. As chemists, we are concerned with the electrons in atoms and more importantly with the electrons in molecules chemists need to know how many electrons there are in a system, where they are, and what energy they have. Before we can understand the behaviour of electrons in molecules, we need to look closely at the electronic structure of an atom. Evidence first, theory later. 

We want to predict, as far as we can, where all the electrons in different molecules are to be found including the ones not involved with bonding. We want to know where the molecule can accommodate extra electrons and from where electrons can be removed most easily. Since most molecules contain many electrons, the task is not an easy one. However, the electronic structure of atoms is somewhat easier to understand and we can approximate the electronic structure of molecules by considering how the component atoms combine. 

The modern theory of the atom cannot tell you exactly where the electrons in atoms are placed. However, it does define regions in space called orbitals, where there is a 95 percent probability of finding an electron. The lowest energy orbital in any atom is called the Is orbital. In this MiniLab, you will simulate the probability distribution of the Is orbital by noting the distribution of impacts or hits around a central target point. 

Electrons in a bonding molecular orbitals are most likely to be found in the region between the two bonded atoms. Why does this arrangement favor bonding In a o- antibonding orbital, where are the electrons most likely to be found in relation to the nuclei in a bond ... 

Just as in the molecular case, and discussed in detail in Chapter 8, fhere is always a choice to be made between filling all the lowest levels with electron pahs and the alternative of allowing some of the higher energy levels to be occupied by electrons with parallel ins. 13.5 and 13.6 showed two extreme cases where all the electrons in a band were either all spin unpaired or all spin paired. An intermediate situation shown In 13.50 is also of importance, where not all of the spins are unpaired. An example of this type, s ch, in addition to being metallic is ms netic, is found in the body-centered cubk structure of elemental hon. There are about 1.5 unpahed electrons per atom. 13.51 depicts an alternative way of showing this result which... 

The decrease in atomic radius moving across the periodic table can be explained in a similar manner. Consider, for example, the third period, where electrons are being added to the third principal energy level. The added electrons should be relatively poor shields for each other because they are all at about the same distance from the nucleus. Only the ten core electrons in inner, filled levels (n = 1, n = 2) are expected to shield the outer electrons from the nucleus. This means that the charge felt by an outer electron, called the effective nuclear charge, should increase steadily with atomic number as we move across the period. As effective nuclear charge increases, the outermost electrons are pulled in more tightly, and atomic radius decreases. 

For H2 to be a stable molecule, the sum of the attractive energies must exceed the sum of the repulsive energies. Figure 9A shows a static arrangement of electrons and nuclei In which the electron-nucleus distances are shorter than the electron-electron and nucleus-nucleus distances. In this arrangement, attractive interactions exceed repulsive interactions, leading to a stable molecule. Notice that the two electrons occupy the region between the two nuclei, where they can interact with both nuclei at once. In other words, the atoms share the electrons in a covalent bond. 

---