Group composition and electronic structure
The oxygen subgroup includes five chemical elements collectively known as chalcogens. The most important representative of this subgroup is oxygen. It is followed by sulfur, selenium, tellurium, and polonium, which closes the group. The latter is an extremely unstable radioactive element that has no stable isotopes in nature. The most well-known natural polonium nuclide with a mass number of 210 has a half-life of only 138.4 days.
A fundamental similarity among all Group VIA elements is the structure of their outer energy level, which consistently contains exactly six electrons. Of this number, two electrons always remain unpaired in the ground state.
The general electronic formula of the outer level is written as $ns^2np^4$. Considering the distribution of electrons across orbitals as a graphic diagram, the picture is as follows: the $ns$-orbital contains one pair of electrons with opposite spins, and the $np$-orbitals contain four electrons (one paired pair and two unpaired electrons).
It is important to note the difference in complete electronic formulas. For oxygen and sulfur, the configuration includes the noble gas core of the preceding period and the outer level ($ns^2np^4$). However, for the heavier analogues — selenium and tellurium — a completely filled penultimate $d$-sublevel appears before the outer level, and their formula takes the form $(n-1)d^{10}ns^2np^4$.
Valence capabilities and excited states
An analysis of the electronic structure makes it possible to explain the valence states of chalcogens. For absolutely all Group VIA elements, the baseline and most characteristic state is the divalent state (II), which is directly related to the presence of two unpaired $p$-electrons.
Oxygen occupies a special place in the subgroup. Its valence can never exceed II. This is due to a fundamental reason: the oxygen atom lacks $d$-orbitals on its outer (second) energy level, meaning there are no free orbitals close in energy to the $2p$-sublevel. Consequently, electron unpairing is physically impossible. In this regard, oxygen shows a strong resemblance to fluorine.
Unlike oxygen, sulfur, selenium, tellurium, and polonium have vacant $d$-orbitals on their outer energy level. This enables them to transition into excited states by expending energy, unpairing their electrons and increasing their valence to IV and VI.
Let us examine this process using the sulfur atom as an example:
- Ground state — S(II): Electronic configuration $3s^2 3p^4$. There are two unpaired electrons. Examples of typical compounds in this state are hydrogen sulfide ($H_2S$) and carbon disulfide ($CS_2$).
- First excited state — S(IV): One electron from the $3p$-sublevel transitions to a vacant $3d$-orbital. This forms the $3s^2 3p^3 3d^1$ configuration with four unpaired electrons. Typical examples are sulfur dioxide ($SO_2$) and sulfur tetrachloride ($SCl_4$).
- Second excited state — S(VI): Further excitation leads to the transition of one electron from the $3s$-sublevel to the $3d$-sublevel. This forms the $3s^1 3p^3 3d^2$ configuration, providing six unpaired electrons. Examples of compounds are sulfur trioxide ($SO_3$) and sulfur hexafluoride ($SF_6$).
It must be emphasized as an important pattern: high valence states (IV and VI) remain stable only if the chalcogen atom is bonded to elements with high electronegativity — oxygen, fluorine, or chlorine. In interactions with other elements, sulfur, selenium, and tellurium can exhibit exclusively a valence of II.
Periodic trends
The chemical and physical properties of elements in the oxygen subgroup vary predictably with increasing atomic number. The most important parameter is electronegativity ($\chi$).
Oxygen has an extremely high electronegativity of 3.5, second only to fluorine in the entire Periodic Table. For the next element, sulfur, this value drops sharply to 2.6. The general trend is a gradual decrease in electronegativity down the group, although the difference between sulfur and selenium turns out to be very insignificant.
The second key trend is the change in metallic properties. Moving down the subgroup from top to bottom, from oxygen to polonium, non-metallic properties weaken, while metallic properties progressively strengthen. This allows the elements to be classified as follows:
- Non-metals: Oxygen (O), Sulfur (S), and Selenium (Se).
- Metalloid: Tellurium (Te), which already begins to exhibit dual characteristics.
- Metal: Polonium (Po), which completes the group as an element with pronounced metallic properties.
Briefly: answers to topical questions
What oxidation states are characteristic of the oxygen subgroup elements?
The sources describe the valence states of the oxygen subgroup elements, not oxidation states. All chalcogens are characterized by a divalent state.
- Oxygen (O) — exhibits a maximum valence of II due to the absence of vacant d-orbitals.
- Sulfur and analogues (S, Se, Te, Po) — possess vacant d-orbitals and are capable of transitioning into excited states, exhibiting valences of IV and VI.
High valence states are stable only in compounds with fluorine, oxygen, and chlorine.
What is the biological role of the macroelements oxygen and sulfur in the human body?
The macroelements oxygen and sulfur perform vital functions in supporting the body's life activities.
- Oxygen (O) — acts as an oxidizing agent in exothermic respiration reactions, being absorbed from the air. It is part of the organic substances that sustain life.
- Sulfur (Sulfur) — is always part of proteins, specifically the amino acids cystine and cysteine. The total mass of sulfur in the human body is about 140 g.
Which protein amino acids contain sulfur?
Sulfur is part of sulfur-containing proteinogenic amino acids:
- Cysteine (Cys) — contains a thiol (sulfhydryl) group in its side chain.
- Methionine (Met) — contains a thioether group and serves as a sulfur donor for the synthesis of cysteine.
The root "sulf" in a name does not always mean sulfur is present: for example, serine contains no sulfur.
What toxic inorganic compounds are formed by oxygen subgroup elements?
Oxygen subgroup elements form a number of highly toxic inorganic compounds, in particular, sulfur derivatives.
- Hydrogen sulfide ($H_2S$) — a colorless gas that is a potent poison. It is formed in the intestines during protein decay and requires detoxification in the liver.
- Sulfur dioxide ($SO_2$) — a toxic sulfurous gas with a pungent odor. It has a pronounced irritating effect on the mucous membranes of the respiratory tract, capable of causing asphyxia and acute pulmonary edema.
Class and exam questions
Why cannot oxygen exhibit a valence of IV or VI like sulfur?
The oxygen atom lacks d-orbitals on its outer energy level. Consequently, it has no vacant orbitals with suitable energy for electron unpairing, so its maximum valence is strictly limited to II.
Under what condition are sulfur and selenium compounds with higher valences stable?
Higher valence states (IV and VI) are stable exclusively in compounds with the most electronegative elements: fluorine, oxygen, and chlorine. With other partners, these elements exhibit only a valence of II.
How are the elements of the oxygen subgroup classified according to metallic properties?
From top to bottom, metallic properties intensify. Oxygen, sulfur, and selenium are typical non-metals, tellurium is classified as a metalloid, and polonium acts as a true metal.
What else to study on this topic
- Drawing up complete electronic formulas of oxygen and sulfur atoms
- Writing the abbreviated electronic formula for selenium and tellurium
- Graphic diagrams of electron distribution across quantum boxes
- Comparison of the properties of the oxygen atom and the fluorine atom
- Energy costs during the transition of an atom to an excited state
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Other topics of section "Chapter 10. Group VIA Elements"
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