General Science MCQs – Page 30 (Q291-300)

Helium is the lightweight, non-flammable noble gas commonly used to fill party balloons and airships, chosen specifically because its low density allows balloons to float in air while its position as a chemically inert noble gas makes it completely non-flammable and safe, unlike the much cheaper but dangerously explosive hydrogen once used for the same purpose, notably in the 1937 Hindenburg disaster. Helium is produced primarily as a byproduct of natural gas extraction, since it accumulates underground from the radioactive decay of uranium and thorium in rocks, and global helium supplies have periodically faced shortages due to this limited extraction source.

The common form of carbon used as pencil lead, mixed with clay to control hardness, is the allotrope graphite, whose layered atomic structure, with atoms bonded tightly within flat hexagonal sheets but only weakly attracted between separate layers, allows those layers to easily slide apart and rub off onto paper, leaving a visible gray mark. Despite the popular name pencil lead, graphite pencils have never actually contained the toxic metal lead, a naming holdover from centuries ago when graphite was mistakenly believed to be a form of that metal. Adjusting the ratio of graphite to clay produces pencils of varying hardness, labeled on a scale from soft to hard.

Zinc is the metal commonly coated onto iron or steel through a process called galvanizing to help prevent rusting, working by two complementary mechanisms: the zinc coating physically shields the underlying iron from exposure to oxygen and moisture, and even where the coating is scratched or damaged, zinc acts as a sacrificial anode, corroding preferentially before the iron because zinc is more reactive, thereby protecting the exposed steel underneath through a form of cathodic protection. Galvanized steel is widely used in outdoor construction, fencing, and automotive body panels, where long-term corrosion resistance is essential.

The standard atomic mass unit is defined based on exactly one-twelfth the mass of a single atom of carbon-12, the most common and stable isotope of carbon, chosen as the reference standard in 1961 to replace earlier, less consistent definitions based on hydrogen or oxygen. This definition assigns carbon-12 an atomic mass of exactly 12 atomic mass units, and all other elements' atomic masses are measured relative to this fixed reference point, expressed as a weighted average accounting for the relative abundance of each element's naturally occurring isotopes, which is why most atomic masses listed on the periodic table are not whole numbers.

Protons and neutrons found together in the nucleus of an atom are collectively referred to as nucleons, a term reflecting their shared role as the fundamental building blocks of atomic nuclei, held together by the strong nuclear force despite the mutual electrical repulsion between positively charged protons. The total number of nucleons in an atom's nucleus determines its mass number, while only the proton count, the atomic number, determines its elemental identity. Nucleons themselves are composed of even smaller fundamental particles called quarks, held together by the strong force mediated by particles called gluons, according to the standard model of particle physics.

Compared to protons and neutrons, the mass of an electron is generally considered nearly negligible, since an electron is roughly 1,836 times lighter than a proton, meaning electrons contribute only a tiny fraction of an atom's overall mass despite occupying most of its physical volume through their orbital motion around the nucleus. This is why atomic mass calculations essentially only need to account for the number of protons and neutrons in the nucleus, largely ignoring electron mass for practical purposes, even though electrons are entirely responsible for an atom's chemical bonding behavior and reactivity through their arrangement in outer shells.

For a given chemical element, the number of protons in its atoms is always fixed and characteristic of that specific element, since the atomic number, defined precisely as the proton count, is the defining property that distinguishes one element from every other element on the periodic table. While the number of neutrons can vary between different isotopes of the same element, and the number of electrons can change when an atom forms an ion, the proton count never changes without transforming the atom into an entirely different element altogether, a transmutation process that occurs naturally in radioactive decay or artificially in nuclear reactors.

As you move down a group in the periodic table, atomic size generally tends to increase, because each successive element in the group has an additional electron shell surrounding the nucleus, placing the outermost, valence electrons progressively farther from the positively charged nucleus and increasing the overall atomic radius. This trend explains why cesium, near the bottom of group 1, has a much larger atomic radius than lithium at the top of the same group, despite both sharing the same single-valence-electron configuration. This increasing atomic size down a group also generally correlates with increasing reactivity among metals like the alkali metals.

As you move from left to right across a period, the metallic character of elements generally tends to decrease, since atomic number and nuclear charge increase while electrons are added to the same outer shell, pulling that shell increasingly tightly toward the nucleus and making it progressively harder for atoms to lose electrons, a key metallic behavior. This trend explains why a period typically begins with highly reactive metals on the left, transitions through metalloids, and ends with reactive nonmetals and finally an unreactive noble gas on the far right, reflecting a steady shift from metallic, electron-losing behavior toward nonmetallic, electron-gaining behavior.

J. J. Thomson discovered the electron in 1897 through his experiments with cathode ray tubes, making it the first subatomic particle ever identified and definitively proving that atoms, previously thought to be indivisible, actually contain smaller internal components. This discovery led Thomson to propose the plum pudding model of the atom, envisioning negatively charged electrons scattered throughout a diffuse, positively charged sphere, a model later disproven and replaced by Ernest Rutherford's nuclear model following his 1911 gold foil experiment, which revealed that positive charge and most of an atom's mass are concentrated in a tiny, dense central nucleus rather than spread throughout the atom.