Introductory Chemistry: Periodic Table, Bonding, and Basic Concepts Flashcards

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Alkali metals

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A group of elements in Group 1 (except hydrogen) such as $Li$, $Na$, and $K$. They have one valence electron and commonly form $+1$ ions; they are highly reactive metals.

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Alkaline earths

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Elements in Group 2 like $Mg$ and $Ca$. They have two valence electrons and typically form $+2$ ions; they are reactive but less so than alkali metals.

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Valence electrons

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Electrons in the outermost energy level of an atom that determine chemical bonding and reactivity. Valence count helps predict ionic charges and bonding behavior.

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Core electrons

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Electrons occupying inner energy levels that do not usually participate in bonding. They screen the nucleus and affect effective nuclear charge felt by valence electrons.

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Cation

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A positively charged ion formed when an atom loses one or more electrons (e.g., $Na^+$). Metals commonly form cations during ionic bond formation.

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Anion

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A negatively charged ion formed when an atom gains electrons (e.g., $Cl^-$). Nonmetals commonly form anions during ionic bond formation.

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Transition metals

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Elements in Groups 3–12 that often have variable oxidation states and form colored ions. Examples include $Fe$, $Cu$, and $Zn$, the latter frequently found as $+2$.

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Representative elements

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Elements in Groups 1–2 and 13–18 showing predictable valence behavior based on group number. They include many common metals, nonmetals, and metalloids.

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Lanthanides

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The top row of the inner transition metals, occupying the $f$-block. They are typically rare-earth metals with similar chemical properties.

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Actinides

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The bottom row of the inner transition metals, many of which are radioactive. They include elements like uranium and thorium used in nuclear chemistry.

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Noble gases

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Group 18 elements such as $He$, $Ne$, and $Ar$ that have full valence shells and are chemically inert under most conditions. They are stable and rarely form compounds.

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Electronegativity

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A measure of an atom's ability to attract electrons in a bond. It increases toward the upper right of the periodic table, with fluorine being the most electronegative.

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Metalloids

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Elements with properties intermediate between metals and nonmetals, such as silicon and germanium. They can conduct electricity weakly and are important semiconductors.

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Diatomic elements

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Elements that naturally exist as molecules of two atoms: $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, $I_2$. Memorize these seven for basic chemistry.

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Ionic bond

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A bond formed by electron transfer from a metal to a nonmetal producing oppositely charged ions that attract electrostatically. Example: $NaCl$ forms from $Na^+$ and $Cl^-$.

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Covalent bond

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A bond formed by sharing electron pairs between atoms, typically between nonmetals. Molecules like $H_2$ and $CO_2$ are held together by covalent bonds.

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Polar covalent

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A covalent bond in which electrons are shared unequally due to differences in electronegativity, producing partial charges (δ+ and δ-). For example, the $H-F$ bond is polar.

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Nonpolar covalent

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A covalent bond with equal or nearly equal sharing of electrons, often between identical atoms or atoms with very similar electronegativity (e.g., C–H is generally nonpolar).

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Allotropes

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Different structural forms of the same element that have distinct properties, such as carbon's diamond and graphite forms. Allotropes can vary in conductivity and hardness.

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Oxidation state

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A formal charge assigned to an atom in a compound representing electron loss or gain relative to the elemental state. Transition metals often display multiple oxidation states.