Question

Explain why the atomic radius of Be is smaller than that of Li

Short answer

Be has a smaller atomic radius than Li because its higher nuclear charge pulls electrons closer.

Step-by-step solution

Understanding Atomic Radius

The atomic radius is the distance from the nucleus of an atom to the outer boundary of its surrounding cloud of electrons. This is generally measured in picometers (pm).

Identifying the Periodic Trend

In the periodic table, as we move from left to right across a period, the atomic radius tends to decrease. This is because additional protons in the nucleus create a stronger positive charge that pulls the electrons closer.

Locate Elements in the Periodic Table

Locate Li and Be in the periodic table. Both elements are in the same period (Period 2), but Be is to the right of Li.

Compare Nuclear Charge

Beryllium (Be) has a higher nuclear charge compared to Lithium (Li) because it has more protons in the nucleus. This increased positive charge more effectively pulls the electron cloud closer to the nucleus.

Understand Electron Shielding

Both Be and Li have electrons in the first and second energy levels (1s and 2s), but they are similar in shielding effect since they are close in the same period. However, the higher proton count in Be still dominates, reducing its radius.

Key concepts

Periodic Trends

The concept of periodic trends refers to patterns observed in the properties of elements across the periodic table. One such trend is the change in atomic radius as you move across a period from left to right. When considering atomic radius, it tends to decrease across a period. This is primarily driven by an increase in the positive charge of the nucleus as you add more protons to the atoms. In simpler terms:

• Across a period, atomic numbers increase.
• Nuclear attraction toward electrons strengthens.
• Electrons are pulled closer, decreasing atomic size.

This trend explains why beryllium (Be), located to the right of lithium (Li) in the same period, showcases a smaller atomic radius. The cumulative effect of increasing protons results in a stronger nuclear pull on the electron cloud, thereby reducing the size of the atom overall.

Nuclear Charge

Nuclear charge signifies the total positive charge within an atom's nucleus due to its protons. It plays a crucial role in determining the strength of attraction between the nucleus and the surrounding electron cloud. An increase in nuclear charge leads to:

• Greater attractive force on electrons.
• A stronger pull on the electron cloud towards the nucleus.
• Potential decrease in atomic radius.

In the case of beryllium versus lithium: - Beryllium (Be) contains 4 protons, while lithium (Li) has only 3. - The additional proton in Be increases the nuclear charge, thus exerting a stronger attraction on its electrons than Li. Therefore, a higher nuclear charge in Be compared to Li draws the electron cloud more tightly around the nucleus. This results in a smaller atomic radius for Be.

Electron Shielding

Electron shielding, or screening, describes how inner-shell electrons can partially block the attractive force between the nucleus and outer-shell electrons. Essentially, these inner electrons create a "shield," reducing the full effect of the nucleus's positive charge on the electrons in the outermost shell. When comparing Be and Li:

• Both possess similar electron configuration up to their valence shell (1s and 2s orbital).
• Shielding effect remains relatively constant across a period as electrons are added to the same energy level.

However, even though the shielding is similar for both elements: - The increased nuclear charge of Be overrides the slight shielding difference. - Resulting in a smaller atomic radius for Be as its nucleus more effectively pulls its valence electrons inwards. Understanding this balance between shielding and nuclear attraction is key to comprehending why beryllium's atomic radius is smaller than lithium's.