Metallic bonds are how metal atoms stick together, and they change as you look across the periodic table. This change is mostly due to three things: atomic size, charge density, and the number of delocalized electrons (which are electrons that can move around freely).

Moving Across a Period:

When you look at the periodic table from left to right, the strength of metallic bonds usually gets stronger.

This is because the nuclear charge (the positive charge from the nucleus) increases. This stronger nuclear charge pulls the positive metal ions and the delocalized electrons closer together.

For example, sodium has weaker metallic bonds than magnesium, and magnesium has weaker bonding compared to aluminum. As you add more delocalized electrons, the bonds become stronger.

Moving Down a Group:

When you look at a group (the columns in the periodic table), the strength of metallic bonds usually increases as you go down the group.

This happens because the atoms get bigger, which means the distance between the positive nucleus and the delocalized electrons also increases. Even though there are more delocalized electrons in a larger atom, this distance weakens the bond.

For example, lithium has weaker metallic bonds than cesium. Even though cesium has more delocalized electrons, the big distance between them and the nucleus makes the bond weaker.

Summary:

Here’s a simpler breakdown of how metallic bonds change:

• Across a Period:

  • The bond strength increases because the stronger nuclear charge pulls the electrons and ions closer.

• Down a Group:

  • The bonds may weaken due to larger atomic size, which means greater distance and less pull from the nucleus, even if there are more delocalized electrons.

Understanding these patterns is important. It helps you see how metallic bonds work and how they affect the properties of different metals.