Introductory Chemistry: Periodic Table, Bonding, and Basic Concepts Summary & Study Notes

🔬 Periodic Table Overview

The periodic table organizes elements by increasing atomic number and recurring chemical properties. Elements are arranged in rows (periods) and columns (groups or families); members of the same group share similar valence electron counts and chemical behavior.

⚛️ Groups & Periods

Group 1 (also Group 1A) contains alkali metals such as $H$ (hydrogen, a nonmetal), $Li$, $Na$, $K$, $Rb$ and $Cs$. These alkali metals have one valence electron and commonly form $+1$ ions. Group 2 (alkaline earth metals) like $Be$, $Mg$, $Ca$, $Sr$, $Ba$ have two valence electrons and commonly form $+2$ ions. Transition metals occupy Groups 3–12 and often exhibit variable oxidation states.

🧪 Representative vs. Inner Transition Elements

Groups 1–2 and 13–18 are called representative elements, while Groups 3–12 are transition metals. The lanthanides and actinides are the inner transition series, typically shown as two rows below the main table.

⚖️ Ions, Valence, and Core Electrons

Valence electrons are the electrons in the outermost shell and determine chemical reactivity. Core electrons occupy inner shells and do not typically participate in bonding. Metals tend to lose valence electrons to form cations (positive ions), while nonmetals tend to gain electrons to form anions (negative ions).

🌊 Electronegativity & Periodic Trends

Electronegativity increases toward the upper right of the periodic table (highest at fluorine). Metals are typically electropositive (tend to give up electrons), while nonmetals are electronegative (tend to attract electrons).

🧭 Metals, Nonmetals, Metalloids

To the lower left of the diagonal line on most periodic tables lie metals: good electrical and thermal conductors, malleable and ductile. To the upper right lie nonmetals: poor conductors and often insulating. Elements along the diagonal (e.g., silicon and germanium) are metalloids, showing intermediate conductivity.

🔗 Bonding: Ionic vs. Covalent

Ionic bonds form by electron transfer between a metal and a nonmetal, producing oppositely charged ions (e.g., $Na^+$ and $Cl^-$) that attract electrostatically. Covalent bonds form when atoms (usually nonmetals) share electrons to form molecules (e.g., $H_2$, $CO_2$). Identifying a metal + nonmetal pair is a quick ionic indicator; two nonmetals usually imply covalent bonding.

↔️ Polarity and Dipoles

When two atoms share electrons unequally due to differing electronegativities, the bond is polar covalent and the molecule has a dipole with a partial negative (δ-) at the more electronegative atom and partial positive (δ+) at the less electronegative atom. A bond is often considered polar if the electronegativity difference is ≳ $0.5$ and nonpolar if it's < $0.5$. Molecules with only C and H are typically nonpolar.

🧱 Diatomic Elements & Allotropes

Seven common diatomic elements exist in their standard states: $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, and $I_2$. Allotropes are different structural forms of the same element (e.g., carbon as diamond and graphite). Diamond is an electrical insulator but an excellent thermal conductor; graphite conducts electricity along planes.

✅ Quick Classification Tips

• Metal + nonmetal → likely ionic (e.g., $MgO$, $LiCl$).
• Nonmetal + nonmetal → likely covalent (e.g., $CO_2$, $H_2O$, $F_2$).
• Group 1 → typically $+1$; Group 2 → typically $+2$; Group 17 (7A) → typically $-1$; Group 16 (6A) → typically $-2$.
• Transition metals often have variable charges (e.g., $Fe^{2+}$, $Fe^{3+}$; $Cu^+$ or $Cu^{2+}$).

Review these concepts and memorize key group behaviors, diatomic elements, and typical oxidation states to prepare for introductory chemistry assessments.