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Problem : What is the net charge on an ion of aluminum that is isoelectronic with neon? With magnesium?
+3, +1
Problem : Applying the Aufbau Principle, write the electronic configuration of the sulfide anion, which has a charge of negative 2.The sulfide ion with a charge of negative 2 has the Argon Noble gas configuration:
Problem : Write the electronic configuration of atomic sulfur, keeping in mind Hund’s rule.
As seen in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group of atoms with a net charge (whether positive or negative) is called an ion. A positively charged ion is a cation while a negatively charged ion is an anion. In this section, we briefly look at some of the processes through which electrons are gained and lost in the formation of ions.
The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity.
The ionization energy is the energy it takes to fully remove an electron from the atom. Ionization energy is a property that varies predictably across the periodic table. Group I and II elements with few electrons in their outer shell have very low ionization energies, while ionization energies increase dramatically moving right along the periodic table. The octet rule gives a straightforward (albeit simplified) explanation of this trend: elements with few valence electrons (those on the left of the periodic table) readily give them up in order to attain a full octet within their inner shells.
When several electrons are removed from an atom, the energy that it takes to remove the first electron is called the first ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so on. In general, the second ionization energy is greater than first ionization energy. This is because the first electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted to the nucleus.
An atom’s electron affinity is the energy change in an atom when that atom gains an electron. The sign of the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its potential energy decreases, meaning that upon gaining an electron the atom gives off energy and the electron affinity is negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom’s electron affinity is positive. An atom with a negative electron affinity is far more likely to gain electrons.
Like ionization energy, electron affinity exhibits periodic trends, with electron affinities becoming increasingly negative from left to right. Remember, as the electron affinity of an atom becomes more negative, it becomes more likely for an atom to gain an electron.
Figure %: Comparing electron affinities of lithium (Group I), carbon (Group II), and fluorine (Group VII). Of these, only fluorine has a tendency to ionize to form anions because it has a very negative electron affinity.
An ionic bond is comprised of the electrostatic attraction of positively and negatively charges ions which holds them together. A common example of a compound held together by ionic bonds is table salt (NaCl), which consists of Na+ cations and Cl- anions held together in a solid crystal. The attractive force between positive and negative ions stabilizes the crystal.
It is important to remember that ionic bonds, unlike covalent bonds, are adirectional, meaning that ionic bonds occur between the ion and all other ions surrounding it. Hence ionic compounds do not occur as discrete units but as large aggregates.
Furthermore, when ionic compounds are placed in water or other polar solvents they dissociate into their component ions. When you encounter ionic compounds in the context of organic reactions, they will almost always occur as free ions in solution. We will see that in the context of organic chemistry, covalent bonding is far more important than ionic bonding.
Problem : What is the net charge on an ion of aluminum that is isoelectronic with neon? With magnesium?
+3, +1
Problem : Applying the Aufbau Principle, write the electronic configuration of the sulfide anion, which has a charge of negative 2.The sulfide ion with a charge of negative 2 has the Argon Noble gas configuration:
Problem : Write the electronic configuration of atomic sulfur, keeping in mind Hund’s rule.
The electrons in an atom fill up its atomic orbitals according to the Aufbau Principle; “Aufbau,” in German, means “building up.” The Aufbau Principle prescribes a few simple rules to determine the order atomic orbitals are filled with electrons:
The outermost shell of an atom is its valence shell, and the electrons in the valence shell are valence electrons. Valence electrons are the highest energy electrons in an atom and are therefore the most reactive. While inner electrons (those not in the valence shell) typically don’t participate in chemical bonding and reactions, valence electrons can be gained, lost, or shared to form chemical bonds. For this reason, elements with the same number of valence electrons tend to have similar chemical properties, since they tend to gain, lose, or share valence electrons in the same way. The Periodic Table was designed with this feature in mind. Each element has a number of valence electrons equal to its group number on the Periodic Table.
The electron configurations for the first and second row elements are shown in in simplified notation.
Our discussion of valence electron configurations leads us to one of the cardinal tenets of chemical bonding, the octet rule. The octet rule states that atoms become especially stable when their valence shells gain a full complement of valence electrons. For example, in above, Helium (He) and Neon (Ne) have outer valence shells that are completely filled, so neither has a tendency to gain or lose electrons. Therefore, Helium and Neon, two of the so-called Noble gases, exist in free atomic form and do not usually form chemical bonds with other atoms.
Most elements, however, do not have a full outer shell and are too unstable to exist as free atoms. Instead they seek to fill their outer electron shells by forming chemical bonds with other atoms and thereby attain Noble Gas configuration. An element will tend to take the shortest path to achieving Noble Gas configuration, whether that means gaining or losing one electron. For example, sodium, which has a single electron in its outer 3s orbital, can lose that electron to attain the electron configuration of neon. Chlorine, with seven valence electrons, can gain one electron to attain the configuration of argon. When two different elements have the same electron configuration, they are called isoelectronic.
When an atom only contains a single electron, its orbital energies depend only on the principle quantum numbers: a 2s orbital would be degenerate with a 2p orbital. However, this degeneracy is broken when an atom has more than one electron. This is due to the fact that the attractive nuclear force any electron feels is shielded by the other electrons. s-orbitals tend to be closer to the nucleus than p-orbitals and don’t get as much shielding, and hence become lower in energy. This process of breaking degeneracies within a shell is known as splitting. In general s orbitals become lowest in energy, followed by p orbitals, d orbitals, and so forth.
It is often convenient to depict orbitals in an orbital energy diagram, as seen below in . Such diagrams show the orbitals and their electron occupancies, as well as any orbital interactions that exist. In this case we have the orbitals of the hydrogen atom with electrons omitted. The first electron shell (n = 1) contains just the 1s orbital. The second shell (n = 2) holds a 2s orbital and three 2p orbitals. The third shell (n = 3) holds one 3s orbital, three 3p orbitals, and five 3d orbitals, and so forth. Note that the relative spacing between orbitals becomes smaller for larger n. In fact, as n gets large the spacing becomes infinitesimally small.
You will see such energy diagrams quite often in your continuing study of organic chemistry. Notice that all orbitals with the same n have the same energy. Orbitals with identical energies are said to be degenerate (not in the moral sense!). Electrons in higher level orbitals have more potential energy and are more reactive, i.e. more likely to undergo chemical reactions.
This chapter is intended as a review of concepts covered in more depth in general chemistry from the organic chemist’s point of view. Bear in mind that many of the topics here are emphasized with an organic bent.
An atom consists of a nucleus of protons and neutrons surrounded by electrons. Each of the elements in the periodic table is classified according to its atomic number, which is the number of protons in that element’s nucleus. Protons have a charge of +1, electrons have a charge of -1, and neutrons have no charge. Electrically, neutral atoms have the same number of electrons and protons, but they can have a varying number of neutrons. Within a given element, atoms with different numbers of neutrons are isotopes of that element. We will see that isotopes typically exhibit similar chemical behavior to each other.
Electrons have such little mass that they exhibit properties of both particles and waves. We know from Heisenberg’s Uncertainty Principle that it is impossible to know the precise location of an electron. Despite this limitation, there are regions around the atom where the electron has a high probability of being found. Such regions are referred to as orbitals.
For isolated atoms (meaning non-bonded), electrons reside in the atomic orbitals of those atoms. Atomic orbitals are classified according to a set of four quantum numbers which describe the energy, shape, and orientation of the orbital.
Principle Quantum Number (n): Indicates how far the orbital is from the nucleus. Electrons are farther away for higher values of n. By Coulomb’s Law we know that electrons which are closer to the positively charged nucleus are more powerfully attracted and thus have lower potential energies. Electrons of orbitals with higher values of n, being farther away from the nucleus, have greater potential energies. In a given atom, all the atomic orbitals with the same n are collectively known as a shell. n can take on integer values of 1 or higher (ex. 1, 2, 3, etc.).
Angular Momentum Quantum Number (l): Describes the shape of the orbital. The angular momentum number (or subshell) can be represented either by number (any integer from 0 up n-1) or by a letter (s, p, d, f, g, and then up the alphabet), with 0 = s, 1 = p, 2 = d, and so on. For example:
when n = 1, l can only equal 0; meaning that shell n = 1 has only an s orbital (l = 0).
when n = 3, l can equal 0, 1, or 2; meaning that shell n = 3 has s, p, and d orbitals.
s orbitals are spherical, whereas p orbitals are dumbbell-shaped. d orbitals and beyond are much harder to visually represent.
Magnetic Quantum Number (m): Gives the orientation of the orbital in space; in other words, the value of m describes whether an orbital lies along the x-, y-, or z-axis on a three-dimensional graph, with the nucleus of the atom at the origin. m can take on any value from –l to l. For our purposes, it is only important that this quantum number tells us that for each value of n there may be up to one s-orbital, three p-orbitals, five d– orbitals, and so on: The s orbital (l = 0) has one orbital, since m can only equal 0. That orbital is spherically symmetrical about the nucleus.
The p orbital (l = 1) has three orbitals, since m = -1, 0, and 1. These three orbitals lie along the x-, y-, and z-axes.
The d orbital (l = 2) has five orbitals, since m = -2, -1, 0, 1, and 2. It is far more difficult to describe the orientation of d orbitals, as you can see:
Spin Quantum Number (s): Tells whether a given electron is spin up (+1/2) or spin down (-1/2). Because the Pauli Exclusion Principle tells us that no two electrons of an atom can have the same set of quantum numbers, each orbital is limited to holding two electrons at most.
Atoms consist of protons and neutrons in the nucleus, surrounded by electrons that reside in orbitals. Since electrons are wave-like in behavior, it is impossible to determine the exact position of an electron. Instead, orbitals describe regions in space where electrons are likely to reside. Orbitals are classified according to the four quantum numbers that represent any one particular orbital’s energy, shape, and orientation. Electrons fill up these orbitals in a systematic fashion, with two electrons per orbital.
When considering the electron configuration of atoms it is useful to consider the valence electrons separately from the inner electrons, since much of the chemistry that elements undergo occurs as a result of the octet rule. The octet rule is the tendency for atoms to gain a full valence shell of electrons. For that reason, elements with similar valence shell configurations have similar chemical properties, giving rise to much of the periodicity of the Periodic Table.
Two such period properties are an atom’s ionization energy and its electron affinity, which are the energies involved when an atom loses and gains electrons, respectively. An atom’s ionization energy and electron affinity determine how easily that atom can lose or gain electrons and thereby form ions with a full valence shell. In gaining and losing electrons atoms also can become positively or negatively charged. When positive and negative ions interact, this gives rise to attractive forces that form the basis of ionic bonding.