Question

Explain how and why the sizes of atoms change across a period of the periodic table.

Short answer

Atomic size decreases across a period due to increased nuclear charge pulling electrons closer.

Step-by-step solution

Understand Atomic Size

The atomic size, or atomic radius, refers to the distance from the nucleus of an atom to the outermost electron. It helps us understand how large an atom actually is.

Examine the Periodic Table

A period in the periodic table is a horizontal row. As you move from left to right across a period, the atomic number of the elements increases, which means more protons and electrons are added.

Analyze Effective Nuclear Charge

As more protons are added to the nucleus across a period, the effective nuclear charge increases. This pulls the electrons closer to the nucleus, as the inner electrons do not shield the outer electrons effectively from the increased positive charge.

Determine the Impact on Atomic Size

With increased effective nuclear charge and poor shielding, the outermost electrons are more strongly attracted to the nucleus. This results in a decrease in the atomic size as you move from left to right across a period.

Key concepts

Periodic Table

The periodic table is like a map of the elements, organizing them based on their atomic number. Each element has a unique box in the table where information like its chemical symbol and atomic number is displayed. This organization allows you to see patterns and trends, including how atomic size changes across a period. A period is a horizontal row in the periodic table. Moving across a period from left to right, each successive element has an additional proton in the nucleus and usually an additional electron in the electron cloud. This structural arrangement highlights essential concepts like the effective nuclear charge and electron shielding, which influence atomic size trends.

Effective Nuclear Charge

Effective nuclear charge refers to the net positive charge experienced by an electron in a multi-electron atom. It's crucial because it affects how tightly electrons are held to the nucleus. As you move across a period in the periodic table, the number of protons in the nucleus increases, enhancing the effective nuclear charge. This is often represented by the formula: \[ Z_{eff} = Z - S \] where \( Z \) is the actual nuclear charge (number of protons), and \( S \) is the shielding constant, representing the average number of electrons between the nucleus and the electron being studied. As effective nuclear charge increases across a period, outer electrons experience a stronger pull towards the nucleus. This results in a decrease in the atomic radius, since the electrons are held more tightly.

Atomic Radius

The atomic radius is a way to measure the size of an atom, even though atoms are not solid spheres. It's defined as the distance from the nucleus to the outermost shell of electrons. Atomic radius becomes smaller as you move across a period from left to right on the periodic table. This change is primarily due to the increase in effective nuclear charge pulling the electrons closer to the nucleus. Although more electrons are added as you cross a period, they do not significantly outpace the increasing positive charge of the nucleus. Instead of occupying more space, the electron cloud contracts because of the stronger attraction to the nucleus. In summary, this means:

• Left to right within a period: atomic radius decreases
• Top to bottom within a group: atomic radius typically increases

Electron Shielding

Electron shielding occurs when inner electrons block or reduce the positive charge that valence (outermost) electrons feel from the nucleus. While moving across a period, electrons are added to the same energy level, not the inner levels, so the shielding effect remains relatively constant. The additional valence electrons do not significantly increase shielding like they might if they were added to new inner levels. This relatively constant electron shielding alongside an increasing number of protons means that the effective nuclear charge increases. Therefore, electron shielding is a vital component in understanding why the atomic radius decreases across a period. The outer electrons do not get much more protection from the stronger pull of the nucleus, leading to a compacted atomic structure.