Structure of energy levels and the first elements
The population of an atom always starts from the lowest energy levels. The first period features only hydrogen (H) and helium (He) since their single available 1s-sublevel is being filled. It is important to note that in a helium atom, this sublevel is located at a lower energy than in hydrogen. This is because the helium nucleus has a larger charge and attracts electrons more strongly.
At helium, the capacity of the first tier is completely exhausted. For the elements following it (starting from lithium), the filling of the second energy level begins.
The total capacity of each period is strictly governed by physical laws. For example, the second and third periods contain exactly 8 elements each. This is explained by the total capacity of the working sublevels: the s-sublevel holds a maximum of 2 electrons, and the p-sublevel holds 6. This is precisely why the short-period version of the Periodic Table distinguishes exactly 8 main groups.
Orbital filling and relation to valence
The chemical capabilities of an atom depend on unpaired electrons — those particles that occupy orbitals singly. The general rule states that the valence of an element is determined by the number of such electrons.
Let's examine the dynamics using elements that fill equivalent (degenerate) 2p-orbitals as an example:
- Nitrogen (N, Z=7): Configuration $1s^2 2s^2 2p^3$. Three electrons are placed singly in three different 2p-orbitals. Result: three unpaired particles and a valence of III (as in ammonia).
- Carbon (C, Z=6): In the ground state, it has two single electrons (formally divalent). However, upon energy absorption, one electron from the 2s-sublevel jumps to a vacant 2p-orbital. An excited state is formed with four unpaired electrons. In a free state, such an atom exists for an extremely short time, but during molecule synthesis, the energy cost of excitation is outweighed by the heat released from new chemical bonds. Therefore, carbon is stably tetravalent.
When free orbitals run out, electron pairing begins. Oxygen (Z=8) acquires one pair in the 2p-tier, leaving two single electrons (valence II). Fluorine (Z=9) has two pairs and only one unpaired electron (valence I). Neon (Z=10) has a fully completed $2p^6$ shell with all electrons paired, making it a chemically inactive gas with a valence of 0.
Classification of elements by blocks
Based on energy diagrams and the order of level filling, chemical elements are divided into blocks:
- s-elements: Located in groups IA and IIA. The formation of their outer s-sublevel is completed here.
- p-elements: Located in groups IIIA–VIIIA. Sequential filling of the outer p-orbitals takes place.
- d-elements: Form secondary subgroups (denoted with "B"); their penultimate d-tier is being filled.
Elements within the same group demonstrate marked chemical similarity. The reason lies in the identical population of the outer level, which dictates the common valence. At the same time, properties vary smoothly from period to period due to the increasing number of electron shells.
Group numbering is directly related to electrons. For groups III through VIII, the number corresponds to the sum of electrons in the two upper sublevels (e.g., s and p, or s and d). Groups IB and IIB are numbered according to the number of electrons in the outer s-sublevel.
Fourth period, d-block, and abbreviated formulas
The fourth period is notable for the appearance of the d-block. The energy sequence here is as follows: the 3d-sublevel requires more energy than 4s, but less than 4p.
The period opens with two s-elements — potassium and calcium, which fill the 4s shell. Next (from scandium to zinc) begins the zone of d-elements, where 10 electrons gradually occupy the 3d-sublevel. When it is completely filled (as in zinc), it becomes energetically stabilized and drops below the 4s-sublevel. Standard p-elements close the period.
For convenience, physicians and chemists use abbreviated notation. Deep layers (the electronic core) are denoted by the symbol of the previous noble gas in square brackets:
- Sodium (Na, Z=11): $[Ne]3s^1$ (one active electron shown).
- Silicon (Si, Z=14): $[Ne]3s^2 3p^2$ (two paired at 3s and two single at 3p according to Hund's rule).
- Manganese (Mn, Z=25): $[Ar]4s^2 3d^5$ (argon core and 7 particles on the upper sublevels).
- Bromine (Br, Z=35): $[Ar]4s^2 3d^{10} 4p^5$ (accounts for the fully closed d-tier).
Briefly: answers to topical questions
What is the maximum electron capacity of s-, p-, d-, and f-sublevels?
The maximum electron capacity of sublevels is determined by the general formula $2(2l + 1)$, where $l$ is the orbital quantum number. The capacity of any energy sublevel is always equal to twice the number of its orbitals.
| Sublevel | Number of orbitals | Maximum number of electrons |
|---|---|---|
| $s$ | 1 | 2 |
| $p$ | 3 | 6 |
| $d$ | 5 | 10 |
| $f$ | 7 | 14 |
What is the Pauli exclusion principle when placing electrons in orbitals?
The Pauli exclusion principle states that all electrons in an atom are in different states and are characterized by different sets of four quantum numbers. When choosing a specific orbital, three quantum numbers are fixed, and only the spin number can change.
Important placement rules follow from this:
- Orbital capacity — a single orbital holds no more than two electrons.
- Electron pair — two electrons in the same orbital with opposite spins.
This fundamental law is strictly observed, takes precedence over the aufbau principle, and ensures the individuality of the electronic structure of each chemical element.
What is Hund's rule for distributing electrons within a single sublevel?
Hund's rule states that the ground state of an atom corresponds to the filling of the maximum number of energetically equivalent orbitals by electrons. Furthermore, the spin quantum numbers (spins) of these electrons must be identical (parallel).
The physical meaning of this rule includes:
- Minimizing repulsion — placing electrons in different spatial orbitals increases the average distance between them.
- Energy reduction — decreasing inter-electron repulsion leads to a reduction in the total energy of the system.
As a result, the configuration with the maximum number of unpaired electrons turns out to be energetically the most favorable for the atom.
For quick reading of abbreviated formulas: mentally replace the noble gas symbol in square brackets (e.g., [Ne] or [Ar]) with the concept of a "ready-made foundation". Everything written to the right is the "working floors" (valence electrons) where all chemical reactions take place.
Class and exam questions
How are unpaired electrons related to the valency of an atom?
Valency is determined by the number of unpaired electrons — those occupying an orbital alone. For example, nitrogen has three such electrons, which is why it forms three chemical bonds.
Why does carbon exhibit a valence of IV if it has only 2 unpaired electrons in its ground state?
Upon receiving energy, a carbon atom transitions into an excited state: one electron jumps from the 2s-sublevel to an empty 2p-orbital. 4 unpaired electrons are formed. The energy spent on this transition is more than compensated for by the formation of chemical bonds.
What explains the similarity in chemical properties among elements of the same group?
Elements of the same group (e.g., lithium and sodium) have an absolutely identical population of the outer energy level. This provides them with a similar valence and identical chemical behavior.
What is the capacity of the s-, p-, and d-sublevels?
The s-sublevel can accept a maximum of 2 electrons, the p-sublevel 6 electrons, and the d-sublevel is designed for 10 electrons.
What else to study on this topic
- Complete energy diagrams of aluminum (Al) and sulfur (S)
- Features of the electronic structure of the copper (Cu) atom
- Compiling abbreviated formulas for phosphorus (P) and chlorine (Cl)
- Detailed distribution of electrons across 2p-orbitals (using oxygen and fluorine as an example)
- Physical justification for the numbering of secondary subgroups IB and IIB
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