With reference to the periodic table, name (a) a halogen element in the fourth period, (b) an element \(\operatorname{sim}\) ilar to phosphorus in chemical properties, (c) the most reactive metal in the fifth period, \((\) d) an element that has an atomic number smaller than 20 and is similar to strontium.

The halogens are part of the periodic table's Group 17 and possess unique characteristics. These elements, including fluorine, chlorine, bromine, iodine, and astatine, are known for their high reactivity. Particularly, bromine is the halogen located in the fourth period. Halogens are nonmetals and tend to form salts when they react with metals. For instance, the reaction between sodium and chlorine to form table salt (sodium chloride) is well-known.

Each halogen has seven electrons in its outermost shell, which explains why they tend to gain one electron to achieve a stable octet configuration. This electron affinity makes them strong oxidizing agents, capable of bleaching colors and disinfecting materials. Due to their reactivity, halogens are rarely found in a free state in nature; they're usually found combined with other elements, often metals, forming ionic compounds.

Understanding chemical properties is essential to predict how substances will interact with one another. These properties are based on the ability of an element to form new substances, either during a reaction or when the conditions around the element change. For example, elements that share the same group in the periodic table, like phosphorus and arsenic, have similar chemical properties and react in comparable ways.

Phosphorus is known for its role in energy transfer within cells, while arsenic, though it can be toxic, has uses in various alloys. Both form similar compounds such as oxides and exhibit allotropy, which is the ability of an element to exist in different physical forms while in the same physical state. Their positioning in the periodic table helps predict these properties and their potential reactions with other elements.

Reactive metals are found on the periodic table's left side, particularly in Group 1, known as the alkali metals, and include lithium, sodium, and potassium, among others. Potassium is highlighted as the most reactive metal in the fifth period due to its readiness to lose the single electron in its outer shell, resulting in a positive ion.

These metals are characterized by their softness, low density, and high reactivity, particularly with water, to form hydroxides and release hydrogen gas, a reaction that can be explosively exothermic. Their high reactivity means that in nature, they're never found in their pure elemental state but always in compounds. Safety precautions are a must when handling these metals, as they can react with moisture in the air or on the skin.

The alkaline earth metals make up Group 2 of the periodic table and include beryllium, magnesium, calcium, strontium, barium, and radium. These metals have two electrons in their outer shell, and as such, they readily give up these two electrons to form divalent cations. Magnesium, with an atomic number smaller than 20, is an example of an alkaline earth metal and is similar to strontium in its chemical behavior.

Alkaline earth metals are less reactive than alkali metals but are still quite reactive, and they're known for their role in biological systems—for instance, magnesium is essential for chlorophyll's function in plants. They're also used in various applications such as lightweight structural materials in the case of magnesium and in fireworks and flares for strontium and barium due to their ability to impart colors to flames.