General Science MCQs – Page 28 (Q271-280)

Fluorine, chlorine, bromine, iodine, and astatine together belong to the group of elements known as halogens, occupying group 17 of the periodic table and sharing seven valence electrons, just one short of a complete, stable outer shell, which drives their strong, characteristic tendency to gain an additional electron and form minus-one ions or share electrons through covalent bonds. The name halogen derives from Greek words meaning salt-producer, reflecting the group's tendency to react directly with metals to form salts, such as sodium chloride from sodium and chlorine. Reactivity generally decreases down the group as atomic size increases and electron attraction weakens.

Fluorine is generally regarded as the most electronegative element on the periodic table, meaning it has the strongest tendency among all elements to attract shared electrons toward itself within a chemical bond, a property that follows the general periodic trend of electronegativity increasing toward the upper right of the table, excluding the largely unreactive noble gases. This extreme electronegativity makes fluorine the single most reactive element overall, capable of forming compounds with nearly every other element, including some noble gases under special conditions. Linus Pauling developed the widely used electronegativity scale in the 1930s, with fluorine assigned the maximum value of 4.0.

Alkali metals, occupying group 1 of the periodic table, typically have one valence electron in their outermost shell, which they readily lose to form a stable, positively charged ion with a full outer shell resembling the nearest noble gas configuration. This single, loosely held valence electron explains why alkali metals are so highly reactive, particularly with water, and why reactivity increases down the group as atomic size grows and the outer electron becomes progressively easier to remove. This shared single-valence-electron configuration is the fundamental reason lithium, sodium, potassium, rubidium, and cesium all display remarkably similar chemical behavior despite their differing atomic sizes and masses.

Alkaline earth metals, occupying group 2 of the periodic table, typically have two valence electrons in their outermost shell, which they tend to lose to form stable, doubly positively charged ions, making them reactive, though generally somewhat less reactive than the neighboring alkali metals in group 1, which need to lose only a single electron. This group includes biologically vital elements like calcium, essential for bones and teeth, and magnesium, essential for chlorophyll and numerous enzyme functions, alongside beryllium, strontium, barium, and radioactive radium, all sharing this characteristic two-valence-electron configuration that defines their common chemical behavior.

Most noble gases, except helium, which has a full outer shell with just two electrons, typically have eight electrons in their outermost shell, satisfying what chemists call the octet rule, a stable, low-energy electron configuration that gives noble gases their characteristic chemical inertness and reluctance to form bonds with other elements. This octet configuration is the same stable target that other elements strive toward when forming chemical bonds, whether by losing, gaining, or sharing electrons, which is why the noble gas electron arrangement serves as a foundational reference point throughout the study of chemical bonding and reactivity in introductory and advanced chemistry alike.

Helium is considered chemically stable because its single outer electron shell is full with just two electrons, satisfying a special stability condition sometimes called the duet rule, an exception to the more common octet rule that applies to most other elements, which typically require eight outer electrons for a stable configuration. Because helium's shell has a maximum capacity of only two electrons, filling it completely requires far fewer electrons than heavier noble gases, yet the resulting stability is just as complete, which is why helium, like the other noble gases, essentially never forms chemical bonds under normal conditions.

The period number of an element in the periodic table generally corresponds to the number of electron shells, or energy levels, that its atoms occupy, with elements in period 1 having electrons in only one shell, period 2 elements having electrons across two shells, and so on up to period 7. As you move across a period from left to right, atomic number increases and additional electrons fill the same outermost shell, gradually changing the element's properties from highly reactive metals to nonmetals and finally to unreactive noble gases, before a new period begins with the addition of an entirely new outer shell.

The isotope of hydrogen that contains one proton and one neutron is commonly known as deuterium, sometimes called heavy hydrogen, roughly twice the mass of ordinary hydrogen, protium, which has no neutrons at all. Deuterium occurs naturally in small quantities and is used to make heavy water, employed as a neutron moderator in certain types of nuclear reactors, and it also serves as a valuable tracer in scientific research due to its distinct mass, detectable through techniques like mass spectrometry and nuclear magnetic resonance. Unlike tritium, deuterium is stable and not radioactive.

The isotope of hydrogen that contains one proton and two neutrons is commonly known as tritium, the heaviest naturally occurring hydrogen isotope, which unlike protium and deuterium is radioactive, undergoing beta decay with a half-life of about 12.3 years. Tritium is used in self-powered lighting applications, such as luminous exit signs and watch dials, where its radioactive decay causes a phosphor coating to glow continuously without requiring an external power source, and it also plays a role in thermonuclear fusion research and weapons, since deuterium-tritium fusion reactions release energy more readily than other fusion fuel combinations.

Carbon-14 is the radioactive isotope commonly used in archaeology to date organic remains, a technique called radiocarbon dating developed by Willard Libby in the late 1940s, based on the fact that living organisms continuously absorb carbon-14 from the atmosphere, but once they die, the isotope begins decaying at a known, steady rate with a half-life of about 5,730 years. By measuring the remaining ratio of carbon-14 to stable carbon-12 in a sample, scientists can estimate how long ago an organism died, making this technique valuable for dating wood, bone, charcoal, and other organic archaeological and geological materials up to roughly 50,000 years old.