Periodic Trends Quiz & Flashcards

Argon has the smallest atomic radius because it is at the end of the period with the highest effective nuclear charge.

Metallic character increases down Group 1 as atoms more readily lose electrons.

The electron enters a higher energy p orbital, making it easier to remove than from the filled s orbital in beryllium.

Both ionization energy and electronegativity increase with greater effective nuclear charge and smaller atomic radius.

Fluorine is the most reactive non-metal due to its high electronegativity and small atomic radius.

Electron affinity becomes more negative across a period as atoms more readily accept additional electrons.

Atomic radius is determined by both nuclear charge and electron shielding, affecting electron attraction to the nucleus.

Halogens have the highest electron affinities because they are one electron short of a stable noble gas configuration.

Melting points decrease down the group as the metallic bonds weaken with increased atomic size.

Osmium is known for having the highest density among transition metals due to its tightly packed atomic structure.

Noble gases have a full valence electron shell, making them stable and not inclined to react.

Reactivity increases down the group due to decreasing ionization energy, making it easier to lose electrons.

Atomic size decreases across a period due to increasing effective nuclear charge attracting electrons closer to the nucleus.

Chlorine has the highest electronegativity in Period 3, as it is a halogen with a strong ability to attract electrons.

Effective nuclear charge increases across a period as protons are added, enhancing electron attraction.

Ionization energy decreases down a group because outer electrons are farther from the nucleus and more shielded.

Fluorine, when it forms an ion, has a small ionic radius due to its high effective nuclear charge attracting electrons closely.

Metallic character increases down a group as atoms more easily lose electrons and exhibit metallic properties.

Stable electron configurations, such as filled or half-filled subshells, can cause irregularities in electron affinity trends.

Sodium has a larger atomic radius with a single valence electron that is less tightly held than in magnesium.

Sulfur is most likely to form a negative ion because it is a non-metal with high electronegativity, seeking to gain electrons.

Electronegativity decreases down Group 17 as atomic size increases, reducing the ability to attract electrons.

Ionization energy decreases down Group 2 because the outer electrons are more shielded and further from the nucleus.

Aluminum has the greatest metallic character as it is a metal, while the others are nonmetals or metalloids.

Increased nuclear charge across a period pulls electrons closer to the nucleus, decreasing atomic radius.

The smaller atomic radius in halogens allows them to attract electrons more effectively, increasing reactivity.

Atomic radius increases down Group 15 due to the addition of electron shells.

Krypton, as a noble gas, has the highest first ionization energy due to its stable electron configuration.

Electronegativity decreases down Group 2 as atomic size increases, reducing the ability to attract bonding electrons.

Partially filled d orbitals allow transition metals to lose different numbers of electrons, leading to variable oxidation states.

Ionization energy generally increases across a period for nonmetals due to increasing nuclear charge.

Density generally increases across a period as atomic mass increases and atomic volume decreases.

There is a decrease from nitrogen to oxygen due to electron repulsion in the p orbital, making oxygen's electron easier to remove.

Anions have a larger ionic radius than their atomic radius due to increased electron-electron repulsion after gaining electrons.

Nitrogen has the highest electronegativity due to its small atomic size and high effective nuclear charge.

Melting points generally increase across a period due to stronger bonding forces in metals and covalent networks.

Fluorine has the highest electron affinity due to its high electronegativity and small atomic radius.

Electronegativity increases across a period as the effective nuclear charge increases, enhancing the attraction for electrons.

Alkali metals have lower ionization energy, making it easier to lose their single valence electron and react.

Transition metals are known for high melting and boiling points due to strong metallic bonding.

Atomic radius in transition metals varies irregularly due to d-orbital electron filling, which affects shielding and nuclear charge.

Oxygen is a non-metal and is more likely to gain electrons to form anions rather than lose them to form cations.

Increased nuclear charge decreases atomic radius by attracting electrons more strongly towards the nucleus.

The electron in the p orbital of Group 13 elements is easier to remove than the s orbital electron in Group 2 elements.

Lithium has the smallest first ionization energy because it only has one electron in its outer shell, which is easily removed.

Atomic radius increases among noble gases down the group as additional electron shells are added.

Neon has the highest ionization energy in Period 2 due to its stable noble gas electron configuration.

Silicon is a metalloid, exhibiting properties of both metals and nonmetals, unlike the other listed elements.