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Example Chemistry flashcards
Define a chemical bond.
A force of attraction between atoms that holds them together in molecules or compounds. Involves the transfer or sharing of valence electrons.
What is the difference between ionic and covalent bonds?
Ionic bonds: electrostatic attraction between oppositely charged ions (electrons transferred). Covalent bonds: atoms share electron pairs (electrons remain shared between nuclei).
What does electronegativity measure?
An atom's ability to attract shared electrons in a chemical bond. Higher electronegativity = stronger pull on electrons. Measured on the Pauling scale (0–4).
How do you determine if a covalent bond is polar or nonpolar?
Compare electronegativity values of bonded atoms. Difference > 0.4 = polar (unequal sharing). Difference ≤ 0.4 or identical atoms = nonpolar (equal sharing).
What is a Lewis structure and what does it show?
A diagram representing valence electrons in atoms/molecules using dots and lines. Lines = bonding pairs; dots = lone pairs. Shows electron distribution without 3D shape.
Explain VSEPR theory and how it predicts molecular geometry.
Valence Shell Electron Pair Repulsion theory: electron pairs (bonding and lone) repel each other and position as far apart as possible to minimize repulsion. This arrangement determines molecular shape.
What is hybridization and why does it occur?
Mixing of atomic orbitals to form new hybrid orbitals suited for bonding. Occurs because hybrid orbitals have better geometry and overlap than unhybridized orbitals, lowering energy.
Determine the hybridization of a carbon atom bonded to four single bonds (sp³ example).
Carbon has 4 bonding regions and 0 lone pairs. One s orbital and three p orbitals mix to form 4 sp³ hybrid orbitals, arranged in tetrahedral geometry (109.5° angles).
What is the difference between sigma (σ) and pi (π) bonds?
σ bonds: head-on overlap of orbitals along internuclear axis; free rotation possible; allows single bonds. π bonds: lateral overlap above/below axis; no rotation; occurs in double/triple bonds alongside σ.
How do you predict bond strength order in N₂, N₂⁺, and N₂⁻ using bond order?
Bond order = (bonding electrons − antibonding electrons)/2. N₂ (BO=3) > N₂⁺ (BO=2.5) > N₂⁻ (BO=2.5). Higher BO = stronger, shorter bond. N₂ is strongest and shortest.
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