Atomic Structure Quiz & Flashcards

The number of protons, or atomic number, defines the element, while neutrons and electrons vary.

D orbitals have a cloverleaf shape, whereas s and p orbitals are spherical and dumbbell-shaped, respectively.

Increased nuclear charge pulls electrons closer, reducing atomic radius across a period.

The Aufbau principle states electrons occupy the lowest energy orbitals available.

An ion is charged due to gain or loss of electrons, while a neutral atom has no net charge.

Isotopes have the same number of protons but differ in neutron count, affecting the mass number.

Electrons are described by probabilistic wavefunctions, not fixed orbits, in the quantum model.

A p subshell can hold a maximum of 6 electrons, with 3 orbitals each holding 2 electrons.

The magnetic quantum number (m_l) specifies the orientation of an orbital in space.

The experiment revealed the atom's dense nucleus, not information about neutrons or electron orbits.

Fluorine has the highest electronegativity of all elements, which allows it to attract electrons strongly.

Atomic radius increases down a group due to additional electron shells.

Covalent bonds involve sharing electron pairs, unlike ionic and metallic bonds which involve transfer or pooling.

Frequency and wavelength are inversely proportional, as wavelength increases, frequency decreases.

Electron shielding reduces effective nuclear charge, increasing atomic size by lessening the pull on outer electrons.

The n=3 energy level can hold a maximum of 18 electrons (2 in s, 6 in p, and 10 in d orbitals).

The principal quantum number n=2 denotes the outermost electrons in the second period.

An anion is a negatively charged ion formed by gaining electrons.

The Pauli exclusion principle ensures that no two electrons in an atom have the same quantum numbers, maintaining unique states.

Isotopes have the same number of protons but vary in the number of neutrons, leading to different mass numbers.

Ionization energy is the required energy to remove an electron from an atom in its gaseous state.

Transition metals involve filling d orbitals, which are responsible for their characteristic properties.

Noble gas configuration refers to a complete outer electron shell, contributing to element stability.

A common misconception is that electrons orbit the nucleus in fixed paths, while they actually occupy probabilistic orbitals.

The fluoride ion (F-) has the same number of electrons as neon, making it isoelectronic.

Ionization energy decreases down a group as electron shielding increases and electrons are further from the nucleus.

Electrons, particularly valence electrons, are responsible for the chemical properties of an element.

Metallic bonds involve electron pooling, where electrons are free to move around metal ions.

Chlorine, being a halogen, readily forms an ionic bond by accepting an electron from sodium.

The principle introduces probabilistic positioning, making precise electron location and momentum impossible to determine simultaneously.

Noble gases are known for being unreactive due to their complete outer electron shells.

Helium has the smallest atomic radius due to having only one electron shell which is held tightly by the nucleus.

Nuclear charge affects electron affinity as a stronger charge attracts additional electrons more effectively.

Sodium, an alkali metal, easily loses an electron to form a cation due to its low ionization energy.

Bohr's model incorrectly describes electrons as being in fixed orbits, not probabilistic orbitals.

An s orbital can hold a maximum of 2 electrons with opposite spins.

Zinc has a complete d subshell (d10) in its ground state configuration.

Electronegativity generally increases across a period due to increased nuclear charge attracting electrons more strongly.

An f subshell can hold up to 14 electrons, with complex shapes unlike spherical or dumbbell orbitals.

Emission occurs when an electron drops to a lower energy level, releasing energy as light.

Transition metals have variable d orbital occupancy, allowing them to exhibit multiple oxidation states.

The transition from a higher to a lower energy level, like n=2 to n=1, releases energy as light.

An excited state has electrons in higher than normal energy levels due to energy absorption.

Alkali metals have one valence electron, making them highly reactive and prone to losing this electron.

The electron cloud refers to probabilistic locations of electrons around the nucleus, not fixed paths or static positions.

Argon has a filled 3p subshell in its electron configuration, unlike the other elements listed.

Carbon is most likely to form covalent bonds due to its ability to share electrons with other atoms.

Noble gases have a complete outer shell, making them stable and resulting in high ionization energy.

Helium has an electron configuration of 1s², filling its only electron shell.

Calcium's outermost electrons occupy the 4s subshell, unlike the other elements listed.