In this lesson, we will explore the characteristics and relationships between metals and non-metals, particularly focusing on cations and anions. Students will understand why alkali metals cannot be found in their free state in nature and discuss the differences in ionization energies between alkali and alkaline earth metals.

We will examine the position of sodium, calcium, and magnesium in the periodic table, highlighting their properties and uses. Additionally, we will differentiate between soft and hard metals, using sodium and iron as examples, and describe the inertness of noble metals.

Moreover, we will identify the commercial value of silver, gold, and platinum, and compile important reactions involving halogens. Finally, we will name some elements that exist naturally in their uncombined form.

Materials around us come in various forms, from airplanes and trains to buildings and tools, largely due to the diverse properties of metals. Non-metals, on the other hand, can exist as gases, liquids, and soft or hard solids, typically found in the upper right positions of the periodic table. Examples of non-metals include carbon, nitrogen, phosphorus, oxygen, sulfur, and most halogens. These elements exhibit a wide range of chemical reactivities and can form various ionic and covalent compounds, many of which are either solid or gaseous.

Metals

Metals are elements (excluding hydrogen) that are electropositive, meaning they tend to lose electrons and form cations. Metals can be categorized based on their chemical reactivity:

• Very Reactive Metals: These include potassium, sodium, calcium, magnesium, and aluminum.
• Moderately Reactive Metals: Examples are zinc, iron, tin, and lead.
• Least Reactive or Noble Metals: This group includes copper, mercury, silver, and gold.

Physical Characteristics of Metals

Some key physical properties of metals are:

1. Most metals are solid at room temperature, with mercury being a notable exception.
2. They generally have high melting and boiling points, although alkali metals are exceptions to this rule.
3. Metals exhibit a shiny appearance (metallic luster) and can be polished.
4. They are malleable, meaning they can be hammered into thin sheets, and ductile, allowing them to be drawn into wires. They also produce a ringing sound when struck.
5. Metals are good conductors of heat and electricity.
6. They typically have high densities.
7. Most metals are hard, though sodium and potassium are softer.

Chemical Properties of Metals

Key chemical properties of metals include:

1. They easily lose electrons to form positive ions.
2. Metals readily react with oxygen to produce basic oxides.
3. They tend to form ionic compounds when reacting with non-metals.
4. Metals exhibit metallic bonding, which contributes to their unique properties.

Electropositive Character

Metals have a tendency to lose their valence electrons, a property known as electropositivity or metallic character. The ease with which a metal can lose its electrons determines its level of electropositivity. The number of electrons that a metal atom can lose is referred to as its valency. For instance, sodium can lose one electron to form a positive ion, giving it a valency of 1. Similarly, zinc can lose two electrons, resulting in a valency of 2.

Trends in Electropositivity

Electropositive character increases down a group in the periodic table because the size of the atoms increases. For example, lithium is less electropositive than sodium, which is in turn less electropositive than potassium. In contrast, electropositive character decreases across a period from left to right. This occurs because atomic sizes decrease due to an increase in nuclear charge, meaning that elements on the left side of a period are more metallic compared to those on the right.

Electropositivity and Ionization Energy

Electropositive character is closely linked to ionization energy, which depends on the size and nuclear charge of the atom. Smaller atoms with higher nuclear charges have higher ionization energies, making them less electropositive or metallic. This is why alkali metals, which have larger atomic sizes and lower ionization energies compared to alkaline earth metals, exhibit the highest metallic character.

For example, sodium has a relatively low first ionization energy, but magnesium has a high first ionization energy and an even higher second ionization energy. This means it is much more difficult to remove a second electron from the Mg⁺ ion due to the strong attraction from the nuclear charge, which also leads to a decrease in the size of the ion.

Overall, alkaline earth metals typically have higher ionization energies than alkali metals, reflecting their lower electropositive character.

Inertness of Noble Metals

Noble metals, which include gold, silver, and platinum, are part of the transition metals, characterized by the filling of their d-orbitals. These metals exhibit various oxidation states and are typically found in the fourth, fifth, and sixth periods of the periodic table.

The chemical behavior of the first transition series resembles that of active metals, with the notable exception of copper. Gold and silver, which belong to group 11, are relatively inactive because they do not easily lose electrons.

Silver is a lustrous white metal known for its excellent conductivity of heat and electricity. It is highly ductile and malleable, allowing it to be shaped easily. Silver’s polished surfaces reflect light well, and it forms a thin layer of oxide or sulfide, making it relatively unreactive. Under normal atmospheric conditions, silver is not affected by air, though it can tarnish in the presence of sulfur-containing compounds like hydrogen sulfide (H₂S). Due to its softness, silver is often alloyed with copper for use in coins, silverware, and jewelry. Additionally, silver compounds are utilized in photographic films and dental applications.

Gold is a yellow, soft metal renowned for being the most malleable and ductile of all metals. One gram of gold can be stretched into a wire over one and a half kilometers long. Gold is extremely nonreactive or inert, unaffected by atmospheric conditions or any single mineral acid or base. Its inertness makes it a popular ornamental metal and a choice for coinage, although it is usually alloyed with other metals, such as copper or silver, to enhance its durability.

Platinum is prized for its unique properties, including color, beauty, strength, flexibility, and resistance to tarnish. It is commonly used in jewelry, providing a secure setting for diamonds and other gemstones. An alloy of platinum, palladium, and rhodium serves as a catalyst in automobile catalytic converters, helping to convert toxic gases like carbon monoxide (CO) and nitrogen oxides (NO₂) into less harmful substances like carbon dioxide, nitrogen, and water vapor. Furthermore, platinum is used in the production of coatings for hard disk drives, fiber optic cables, and fiberglass-reinforced plastics, as well as in the manufacture of glass for liquid crystal displays (LCDs).

Non-Metals

Non-metals are elements that can form negative ions (anions) by gaining electrons, making them electronegative in nature. This property allows them to form acidic oxides. The valency of non-metals depends on the number of electrons they accept; for example, chlorine has a valency of 1 because it accepts one electron, while oxygen has a valency of 2 as it can accept two electrons.

Non-Metallic Character

The non-metallic character of an element is determined by its electron affinity and electronegativity. Smaller elements with a high nuclear charge tend to be more electronegative and have higher electron affinities, leading to stronger non-metallic characteristics. Generally, non-metallic character decreases down a group in the periodic table and increases across a period from left to right, reaching its peak with the halogens. Fluorine is recognized as the most non-metallic element.

Non-metals are found in Groups 14 (carbon), 15 (nitrogen and phosphorus), 16 (oxygen, sulfur, and selenium), and 17 (the halogens: fluorine, chlorine, bromine, and iodine) of the periodic table.

Physical Properties of Non-Metals

The physical properties of non-metals vary but generally include:

1. They can exist in all three states of matter: gases (like oxygen) at the top of the group, liquids, or solids.
2. Solid non-metals are typically brittle and break easily.
3. They are poor conductors of heat and electricity, with graphite being an exception.
4. Non-metals have a dull appearance, except for iodine, which is lustrous.
5. They are generally soft, except for diamond.
6. Non-metals usually have low melting and boiling points, with exceptions like silicon, graphite, and diamond.
7. They tend to have low densities.

Chemical Properties of Non-Metals

Key chemical properties of non-metals include:

1. Their valence shells are electron-deficient, leading them to readily accept electrons to achieve stability.
2. Non-metals form ionic compounds with metals and covalent compounds with other non-metals (e.g., CO₂, NO₂).
3. Generally, non-metals do not react with water.
4. They do not typically react with dilute acids, as non-metals are themselves electron acceptors.

Halogens

The halogens, found in Group 17 of the periodic table, include fluorine, chlorine, bromine, iodine, and astatine. Fluorine and chlorine exist as diatomic gases at room temperature, while bromine is a liquid and iodine is a solid due to increasing intermolecular forces as atomic size increases down the group.

Halogens have a valence shell electronic configuration of ns²np⁵, meaning they are only one electron short of a full outer shell. This enables them to readily accept an electron from metals or share an electron with other non-metals, forming ionic bonds with metals and covalent bonds with non-metals.

Oxidizing Properties of Halogens

All halogens act as oxidizing agents, with fluorine being the strongest and iodine the weakest. Fluorine can oxidize any halide ion (X⁻) in solution, converting itself to the fluoride ion (F⁻). Similarly, chlorine can displace bromide and iodide ions from their salt solutions, oxidizing them to form bromine and iodine.

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