Variation of atomic radii
Within the Periodic Table, the size of atoms changes cyclically. Moving across a period from left to right, the nuclear charge increases, and electrons are pulled more strongly toward the center. Because of this, atomic radii decrease. However, as the energy level fills up, electrons begin to repel each other more strongly, so the contraction slows down toward the end of the period.
When transitioning from a noble gas to the alkali metal of the next period, the radius increases sharply due to the appearance of a new electron shell. In main subgroups, moving from top to bottom, radii also increase because the number of energy levels is constantly growing. The graph of radius versus nuclear charge is sawtooth-shaped: alkali metals (lithium, sodium, potassium) are at the maxima, and noble gases are at the minima.
Ionization energy and electron affinity
The ability of an atom to interact with electrons is described by two key parameters that change in directions opposite to the change in radius:
- Ionization energy is the energy required to remove an electron and turn an atom into a cation. It increases across periods from left to right and decreases down groups from top to bottom.
- Electron affinity is the energy released when capturing an outside electron to form an anion. The trend is similar, but there are exceptions: alkaline earth metals and noble gases have no electron affinity at all.
Even in active alkali metals, removing an electron is much more difficult than attaching one. Forming a sodium cation requires the absorption of 495.8 kJ/mol, whereas forming an anion releases only 52.9 kJ/mol. That is why negatively charged sodium ions are extremely unstable and do not occur in complex substances.
Electronegativity and bond polarity
Electronegativity indicates an atom's ability to attract electron pairs of chemical bonds. The absolute value is calculated as the half-sum of the ionization energy and electron affinity. In practice, relative values calculated with respect to lithium are used.
Let us look at the values for the second period: from lithium (1.0), the index steadily rises through beryllium (1.6), boron (2.0), carbon (2.5), nitrogen (3.0), and oxygen (3.5) up to fluorine (4.0). In the third period, the trend continues: from sodium (0.9) to chlorine (3.0).
The difference in electronegativity between two bonded atoms determines the degree of polarity (ionic fraction) of the bond. The atom with the higher electronegativity receives a partial negative charge, and the one with the lower electronegativity receives a positive charge. For example, at a difference of 1.0, the pole charge reaches 0.22, which corresponds to 22% ionicity. If the electron pair completely transfers to one of the partners, the pole charges equal unity, and the bond is considered ionic. It should be remembered that electronegativity is not a strict constant, and estimates based on it are approximate.
Briefly: answers to topical questions
How do the metallic and reducing properties of elements change in groups and periods?
The metallic properties of elements increase when moving down groups from top to bottom.
- Group IV — the nonmetal carbon is replaced by the metal lead.
- Group V — the nonmetal nitrogen is replaced by the metal bismuth.
- Group VI — the nonmetal oxygen is replaced by the metal polonium.
Information on the variation of metallic properties within periods, as well as periodic trends in reducing properties, is not available in the sources.
How does the electronegativity of elements in main subgroups change when moving from top to bottom?
When moving down groups from top to bottom, the electronegativity of elements decreases. This pattern correlates directly with the decrease in ionization energy and electron affinity.
| Element | Electronegativity |
|---|---|
| Oxygen | 3,5 |
| Sulfur | 2,6 |
The difference in indices between certain elements when moving down may be insignificant (for example, between sulfur and selenium). Overall in the Periodic Table, the lowest electronegativity is observed in francium.
Radius and ionization energy work like a seesaw: when the atomic radius drops across a period, the ionization energy inevitably skyrockets.
Class and exam questions
Why does the atomic radius decrease across a period even though the number of electrons increases?
As the atomic number increases, the positive nuclear charge grows. It begins to pull electron shells much more strongly, compressing the atom.
Can noble gases form negative ions?
No. According to periodic trends, elements of group VIIIA completely lack electron affinity.
How to determine charges in a molecule using electronegativity?
The negative charge always shifts toward the atom with the higher electronegativity. The atom with the lower value becomes positively charged.
How do the radii of halogens change from fluorine to bromine?
Radii increase because these elements are in the same group, and the number of electron shells increases when moving downward.
What else to study on this topic
- The sawtooth dependence of atomic radius on nuclear charge and its extreme points.
- The ratio of ionization energy and electron affinity using the sodium atom as an example.
- The effect of electron repulsion on the slowdown of radius contraction toward the end of a period.
- Calculation of the theoretical dipole charge of a hydrogen chloride molecule using the electronegativity table.
- Reasons for the lack of electron affinity in group IIA elements.
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