Understand why and how atoms bond together with this essential chemical
bonding flashcard deck. Chemical bonding is the foundation of molecular
chemistry — mastering it explains the properties of every substance
you encounter, from water to DNA to metals...
Understand why and how atoms bond together with this essential chemical
bonding flashcard deck. Chemical bonding is the foundation of molecular
chemistry — mastering it explains the properties of every substance
you encounter, from water to DNA to metals.
This deck covers ionic bonding, covalent bonding (polar and nonpolar),
metallic bonding, Lewis dot structures, electronegativity and bond
polarity, VSEPR theory and molecular geometry, hybridization, and
intermolecular forces including hydrogen bonding, dipole-dipole
interactions, and London dispersion forces. Essential for AP Chemistry,
GCSE, A-Level, and university general chemistry courses.
Ready to test yourself?
Flip through all 25 cards in interactive study mode.
A force of attraction that holds two or more atoms together. Bonds form when atoms achieve a more stable (lower energy) arrangement, usually by completing their outer electron shell.
A bond formed by the transfer of electrons from a metal to a nonmetal, creating oppositely charged ions (cation + anion) that attract each other. Example: Na⁺Cl⁻ (table salt).
A bond formed by the sharing of electrons between two nonmetal atoms. Both atoms achieve a full outer shell by sharing. Example: H₂, H₂O, CO₂.
Polar: Electrons shared unequally because atoms have different electronegativities (e.g., H₂O, HCl). Nonpolar: Electrons shared equally between identical or similar atoms (e.g., H₂, O₂, CH₄).
Bonding in metals where positively charged metal ions are surrounded by a "sea" of delocalized electrons. Explains metals' conductivity, malleability, and ductility.
Diagrams showing valence electrons of atoms as dots around the element symbol. Used to visualize bonding and lone pairs in molecules. Each dot represents one valence electron.
Atoms tend to form bonds until they are surrounded by 8 electrons in their outer shell (like noble gases). Hydrogen is an exception — it only needs 2 electrons.
Single bond: 1 shared pair (2 electrons) — e.g., H–H. Double bond: 2 shared pairs (4 electrons) — e.g., O=O. Triple bond: 3 shared pairs (6 electrons) — e.g., N≡N. Stronger and shorter as bond order increases.
The tendency of an atom to attract shared electrons. Difference < 0.4 = nonpolar covalent. Difference 0.4–1.7 = polar covalent. Difference > 1.7 = ionic bond.
Valence Shell Electron Pair Repulsion theory — electron pairs around a central atom repel each other and arrange themselves to be as far apart as possible, determining molecular shape.
Bent (V-shaped). Central oxygen has 2 bonding pairs and 2 lone pairs. Bond angle ≈ 104.5°. Polar molecule due to asymmetric shape and O–H bond polarity.
Trigonal pyramidal. Central nitrogen has 3 bonding pairs and 1 lone pair. Bond angle ≈ 107°. Polar molecule.
Tetrahedral. Carbon has 4 bonding pairs, no lone pairs. Bond angle = 109.5°. Nonpolar molecule — symmetrical shape cancels out bond dipoles.
Linear. Carbon has 2 double bonds, no lone pairs. Bond angle = 180°. Nonpolar overall — polar C=O bonds cancel out due to symmetry.
The mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies. sp³ (tetrahedral), sp² (trigonal planar), sp (linear).
A strong intermolecular attraction between a hydrogen atom bonded to F, O, or N and the lone pair of another F, O, or N atom. Explains water's high boiling point and surface tension.
The weakest intermolecular force — temporary dipoles caused by random electron movement attracting neighboring atoms/molecules. Present in ALL molecules. Stronger in larger, heavier molecules.
Intermolecular attractions between the positive end of one polar molecule and the negative end of another. Stronger than London forces but weaker than hydrogen bonds.
Stronger IMFs = higher boiling point, melting point, and viscosity. Order of strength: London forces < dipole-dipole < hydrogen bonds < ionic interactions.
A covalent bond where both electrons in the shared pair come from the same atom. Example: NH₄⁺ (ammonium ion) — nitrogen donates both electrons to H⁺.
The distance between the nuclei of two bonded atoms. As bond order increases (single→double→triple), bond length decreases and bond strength increases.
When a molecule cannot be represented by a single Lewis structure — the actual structure is an average (resonance hybrid) of multiple structures. Example: benzene (C₆H₆), ozone (O₃).
Sigma bond: Head-on orbital overlap — forms first in any bond (all single bonds are sigma). Pi bond: Side-on orbital overlap — forms in addition to sigma in double and triple bonds.
Bent shape + high electronegativity of O + 2 lone pairs create strong hydrogen bonds. This gives water an abnormally high boiling point, specific heat capacity, and surface tension.
Intramolecular: Forces within a molecule (chemical bonds — ionic, covalent, metallic). Intermolecular: Forces between separate molecules (LDF, dipole-dipole, hydrogen bonds). IMFs are much weaker.
