Question

With only minor irregularities, the melting points of the first series of transition metals rise from that of Sc to that of Cr and then fall to that of Zn. Give a plausible explanation for this phenomenon based on atomic structure.

Short answer

The trend in melting points from Scandium to Zinc in transition metals can be explained by the atomic structure, specifically the number of unpaired d-electrons involved in metallic bonding. The melting points increase as the number of unpaired d-electrons increases from Sc to Cr, strengthening the metallic bond. After Cr, the electrons start to pair up, meaning fewer metallic bonds and thus reduced melting points. Zinc has no unpaired electrons, hence has a lower melting point.

Step-by-step solution

Understanding Transition Metals

Transition metals are those found in Groups 3-12 on the Periodic Table. They have special properties due to the presence of electrons in their d orbital. It's important to understand their general atomic structure first which has an incompletely filled d orbital where one or two of their external energy level electrons are found.

Realizing the trend of melting points within the series

It's observed that the melting points of transition metals rise from Scandium (Sc) to Chromium (Cr), and then fall to Zinc (Zn). Thus, we can highlight a trend: Sc < Cr > Zn.

Explaining the Trend

The observed trend in melting points can be explained by the bond strength between atoms. As we move from Sc to Cr, the number of unpaired d-electrons increases. Unpaired d-electrons can form more and stronger metallic bonds, resulting in an increase in melting points. After reaching Cr, the number of unpaired d-electrons starts to decrease because the electrons start pairing up in the d-orbital, weakening the metallic bonds, and hence the melting points drop.

Case of Zinc

In the case of Zinc (Zn), its d-orbital is fully filled and there are no unpaired electrons to form metallic bonds. This explains its lower melting point as compared to the others.

Key concepts

Atomic Structure

Atomic structure is pivotal in understanding the melting points of transition metals. Transition metals are found in Groups 3-12 of the Periodic Table. Their distinct feature is their d-orbital electrons, which play a significant role in their chemical properties. Each transition metal possesses an incompletely filled d-orbital, which is the highest energy level occupied by these electrons. These d-orbitals participate in bonding and influence physical properties such as melting point.

The electrons in the d-orbitals of transition metals influence their bonding capabilities. This, in turn, affects their structures and melting points. In particular, the d-orbitals allow these metals to form electron clouds, leading to metallic bonds that contribute to their overall strength and stability. As the atomic number increases, the addition of electrons to the d-orbital can lead to variations in bonding characteristics and atomic structures.

Understanding how these electrons are arranged in the orbitals helps explain why certain trends, such as those in melting points, are observed across the series of transition metals.

Melting Points

The melting points of transition metals don't follow a simple pattern and can be quite fascinating. Generally, there's a rise from Scandium (Sc) to Chromium (Cr), followed by a fall to Zinc (Zn). This trend isn't random but can be explained by examining the nature of metallic bonds within these metals.

As we move from Sc to Cr, the transition metals exhibit an increase in the number of unpaired d-electrons. These unpaired electrons bolster the metallic bond strength between atoms, raising the energy required to break these bonds, and thus the melting point.

• Scandium: Low unpaired electron count, weaker metallic bonds, hence lower melting point.
• Chromium: Peak unpaired electron count, resulting in strong metallic bonds and higher melting point.

As we progress past Chromium toward Zinc, the scenario changes. The d-orbital electrons begin pairing up, diminishing the number of unpaired electrons. Consequently, the strength of metallic bonds weakens, which explains the falling trend in melting points from Chromium to Zinc.

• Zinc: All d-electrons are paired, very weak metallic bonds, and thus lower melting point.

D-Orbital Electrons

D-orbital electrons are crucial in dictating the physical characteristics of transition metals. The d-orbital can hold a maximum of 10 electrons, and their arrangement impacts how these metals bond. Having an incomplete d-orbital gives transition metals their unique properties, including varied melting points.

• Unpaired d-electrons contribute to stronger metallic bonds.
• Pairing of d-electrons results in weaker metallic bonds.

This concept is linked with the melting points of transition metals. For example, as the d-electrons start filling up, each new unpaired electron (from Sc to Cr) enhances the metal's ability to form strong metallic bonds, which require more heat energy to break, hence a higher melting point.

However, once the electrons begin pairing (past Cr in the series), the metallic bonds become weaker. The energy required to break these bonds decreases, thus lowering melting points. In the case of Zinc, the d-orbital is complete with paired electrons, resulting in the weakest metallic bonds within this series.