Metals, Non-metals, Ions, and Periodic Trends

Flashcards covering the formation of ions (cations and anions), charges based on periodic table groups, characteristics of transition metals, and the definitions and periodic trends of ionization energy and electron affinity.

Cations

Positive ions formed when metals lose valence electrons, resulting in more protons than electrons.

Anions

Negative ions formed when non-metals gain electrons, resulting in more electrons than protons.

Metals (ion formation)

Lose valence electrons to become cations with a positive charge, aiming for the nearest noble gas configuration by moving 'backwards' on the periodic table.

Non-metals (ion formation)

Gain electrons to become anions with a negative charge, aiming for the nearest noble gas configuration by moving 'forwards' on the periodic table.

Group 1A elements (and Hydrogen) Charge

Always form a +1 charge by losing one valence electron.

Group 2A elements Charge

Always form a +2 charge by losing two valence electrons.

Group 3A elements (e.g., Aluminum) Charge

Mostly form a +3 charge by losing three valence electrons.

Group 7A elements (Halogens) Charge

Always form a -1 charge by gaining one electron.

Group 6A elements Charge

Always form a -2 charge by gaining two electrons.

Group 5A elements Charge

Always form a -3 charge by gaining three electrons.

Group 4A elements (e.g., Carbon)

Energetically unfavorable to lose or gain four electrons; they typically achieve stability by sharing electrons in covalent compounds.

Noble Gas Configuration

The stable electron arrangement of a noble gas (typically an octet of eight valence electrons), which other atoms strive to achieve by losing or gaining electrons.

Transition Metals (D-block)

Form cations by losing both s and d valence electrons; they often have variable charges and do not strictly follow the octet rule.

Ionization Energy (IE)

The energy required to remove an electron from a single atom in the gaseous state; low IE values indicate an atom loses electrons easily, favoring cation formation.

Electron Affinity (EA)

The energy released when an electron is added to a single atom in the gaseous state; high EA values indicate an atom gains electrons easily, favoring anion formation.

Alkali Metals Reactivity

Highly reactive due to having the lowest ionization energies on the periodic table, meaning they most readily lose an electron.

Halogens Reactivity

Highly reactive due to having the highest electron affinities (excluding noble gases), meaning they most readily gain an electron.

Shielding Effect

Inner shell electrons reduce the attractive force of the nucleus on outer valence electrons, making them easier to remove in larger atoms.

Ionization Energy Trend on Periodic Table

Generally increases from the bottom-left to the top-right of the periodic table; alkali metals have very low IE, while noble gases have extremely high IE.

Electron Affinity Trend on Periodic Table

Generally increases from the bottom-left to the top-right of the periodic table, with halogens having the highest values and noble gases having zero EA.