Unit 9: Classification of Metals, Non-metals & Metalloids
Understanding the fundamental categories of elements and their bonding characteristics.
9.1 What are Metals? (Metallic Bonding)
Metals are a class of elements that are typically shiny, dense, and good conductors of heat and electricity. They occupy the majority of the periodic table, specifically the left-hand side and the center. The unique properties of metals arise from a special type of chemical bond known as metallic bonding.
In metallic bonding, metal atoms, which are generally electropositive (meaning they have a low ionization energy and readily lose electrons), release their valence (outermost) electrons. These electrons are no longer associated with any single atom and become delocalized, forming a "sea" of mobile electrons that surrounds a fixed lattice of positive metal ions (cations). The strong electrostatic attraction between the positive cations and the negative electron sea holds the entire structure together, giving metals their characteristic strength and integrity.
This model explains why metals are excellent conductors: the delocalized electrons are free to move and carry an electrical current. It also explains their malleability and ductility, as the layers of cations can slide past one another without breaking the metallic bond.
Solved Examples:
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Why is sodium (Na) considered a typical metal?
Solution: Sodium is in Group 1 of the periodic table. It has a low first ionization energy and readily loses its single valence electron to form a Na⁺ ion. These lost electrons form a delocalized sea, resulting in metallic bonding, which is characteristic of metals. -
What is the "electron sea" model?
Solution: The electron sea model is a simplified description of metallic bonding. It visualizes a metal as a regular array of positive ions (cations) immersed in a "sea" of mobile valence electrons that are shared among all the ions. -
Why does metallic character increase down a group in the periodic
table?
Solution: As you move down a group, the number of electron shells increases. This leads to a larger atomic radius and increased shielding of the valence electrons from the nucleus. Consequently, the valence electrons are held less tightly and are more easily lost, enhancing the element's electropositivity and metallic character.
9.2 What are Non-metals? (Covalent Bonding)
Non-metals are elements found on the upper right-hand side of the periodic table. Unlike metals, they are typically poor conductors of heat and electricity, are not shiny (lustrous), and are brittle in their solid form.
Non-metals are generally electronegative, meaning they have a strong attraction for electrons. Instead of losing electrons to form a delocalized sea, non-metal atoms tend to achieve a stable electron configuration by sharing their valence electrons with other atoms. This sharing of electrons creates strong, directional covalent bonds.
Because electrons in non-metals are localized in these covalent bonds (or held tightly by individual atoms), they are not free to move. This explains why non-metals are poor conductors of electricity. Examples of non-metals include carbon (C), nitrogen (N), oxygen (O), sulfur (S), and chlorine (Cl).
Solved Examples:
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Why does chlorine (Cl) exist as a diatomic molecule ($Cl_2$) with a covalent
bond?
Solution: A chlorine atom has 7 valence electrons and is highly electronegative. To achieve a stable octet, it is much more energetically favorable for two chlorine atoms to share a pair of electrons, forming a single covalent bond, than it is for one to lose an electron and the other to gain one. -
Explain why graphite, an allotrope of the non-metal carbon, can conduct
electricity.
Solution: This is a special exception. In graphite, each carbon atom is covalently bonded to three other carbons in flat layers. The fourth valence electron of each carbon atom is delocalized within the layer. These delocalized electrons are free to move along the layers, allowing graphite to conduct electricity. -
Why does metallic character decrease from left to right across a
period?
Solution: Moving across a period, the number of protons in the nucleus increases, leading to a stronger effective nuclear charge. This pulls the valence electrons closer and holds them more tightly, making them harder to lose. Elements become more electronegative and less likely to form metallic bonds, favoring covalent bonding instead.
9.3 What are Metalloids?
Metalloids, or semi-metals, are a unique group of elements that exhibit properties intermediate between those of metals and non-metals. They are found along the "staircase" line that separates metals from non-metals on the periodic table. Common metalloids include Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), and Tellurium (Te).
The most defining characteristic of metalloids is their ability to act as semiconductors. They are poor conductors of electricity at low temperatures (like non-metals) but become better conductors at higher temperatures (a property unlike most metals, whose conductivity decreases with temperature). This behavior makes them essential in the electronics industry for manufacturing transistors, computer chips, and other components.
Physically, metalloids are often brittle and can be shiny, but their chemical reactivity depends on the substance they are reacting with. They can behave as non-metals when reacting with metals, and as metals when reacting with non-metals.
Solved Examples:
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Silicon (Si) is the basis of modern electronics. Explain why its metalloid
properties are crucial for this role.
Solution: As a semiconductor, silicon's electrical conductivity can be precisely controlled. By introducing specific impurities (a process called doping), its conductivity can be modified to create the p-type and n-type materials that form transistors, the fundamental building blocks of all microchips. -
An element is a brittle solid, a semiconductor, and forms an oxide that is
weakly acidic. Classify this element.
Solution: The combination of being a brittle solid (non-metal property), a semiconductor (defining metalloid property), and forming a weakly acidic oxide (a tendency of non-metals and some metalloids) strongly suggests the element is a metalloid.
9.4 Further Classification of Metals (s-block, p-block, d-block)
Metals can be further categorized based on the orbital their valence electrons occupy, which corresponds to their position in the periodic table.
- s-Block Metals: These are the metals in Groups 1 (alkali metals) and 2 (alkaline earth metals). Their outermost valence electrons are in an s-orbital. They are highly reactive, electropositive, soft, and have low melting points. Examples: Sodium (Na), Calcium (Ca).
- p-Block Metals: These metals are located in Groups 13 to 16. Their outermost electrons occupy both s and p-orbitals. They are generally softer and have lower melting points than d-block metals but are less reactive than s-block metals. Examples: Aluminum (Al), Tin (Sn), Lead (Pb).
- d-Block Metals: Also known as the transition metals, these are found in Groups 3 to 12. Their valence electrons are in both s and d-orbitals. They are known for being hard, having high melting points and densities, forming colored compounds, having variable oxidation states, and often acting as good catalysts. Examples: Iron (Fe), Copper (Cu), Gold (Au).
The f-block metals (lanthanides and actinides) are another category, but they are less commonly discussed in introductory chemistry.
Solved Examples:
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Classify the following metals into s-block, p-block, or d-block: Potassium
(K), Zinc (Zn), and Aluminum (Al).
Solution:
- Potassium (K): Located in Group 1. Its valence electron is in the 4s orbital. It is an s-block metal.
- Zinc (Zn): Located in Group 12. Its valence electrons are in the 4s and 3d orbitals. It is a d-block metal.
- Aluminum (Al): Located in Group 13. Its valence electrons are in the 3s and 3p orbitals. It is a p-block metal. -
Why are d-block metals often good catalysts?
Solution: d-block metals are effective catalysts because they can exist in multiple oxidation states and have available d-orbitals. This allows them to readily donate and accept electrons from reactant molecules, providing an alternative reaction pathway with a lower activation energy.