Bonding Flashcards

Interactive active recall cards for ionic, covalent, and metallic structures, molecule shapes, polarity, and forces.

Ionic & Metallic Bonding: read all 4 questions and answers

Define ionic bonding.
The electrostatic attraction between oppositely charged ions in a giant lattice structure.
State three physical properties of giant ionic compounds.
High melting and boiling points, electrical conductivity only when molten or aqueous, and solubility of many ionic salts in polar solvents (depending on lattice vs hydration enthalpies).
Define metallic bonding.
The electrostatic attraction between a lattice of positive metal ions and a surrounding sea of delocalised electrons.
Why are metals malleable and ductile?
The layers of positive metal ions can slide over each other without breaking the metallic bond, as the sea of delocalised electrons is free to move with them.

Covalent & Macromolecular Structures: read all 4 questions and answers

Define a covalent bond.
The electrostatic attraction between shared pairs of electrons and the nuclei of the bonded atoms.
What is a coordinate (dative) covalent bond?
A covalent bond where both of the shared electrons are donated by the same atom.
Give two examples of macromolecular structures and describe their bonding.
Diamond (each C atom is bonded tetrahedrally to 4 other C atoms in a 3D giant lattice) and Graphite (each C atom is bonded trigonal planarly to 3 other C atoms in hexagonal sheets, with a delocalised pi system extending along and across each sheet and weak van der Waals forces between sheets).
Why does graphite conduct electricity but diamond does not?
Graphite has delocalised electrons free to move along its layers, whereas in diamond all outer electrons are localised in fixed covalent bonds.

Shapes of Molecules (VSEPR): read all 5 questions and answers

State the fundamental principle of VSEPR theory.
Electron pairs around a central atom repel each other and arrange themselves as far apart as possible to minimise repulsion.
Compare the relative strength of repulsion between lone pairs and bonding pairs of electrons.
Lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion.
What is the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs?
Tetrahedral, 109.5°.
What is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pair?
Pyramidal, 107° (the lone pair repels bonding pairs, reducing the bond angle by approximately 2.5°).
What is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs?
Non-linear (bent), 104.5°.

Electronegativity & Polarity: read all 4 questions and answers

Define electronegativity.
The power of an atom to attract the electron density in a covalent bond towards itself.
State the trend in electronegativity across a Period and down a Group.
Electronegativity increases across a Period (nuclear charge increases, shielding remains constant) and decreases down a Group (atomic radius and shielding increase).
What makes a covalent bond polar?
An unequal sharing of electrons due to a difference in electronegativity between the two bonded atoms.
Why can a molecule with polar bonds (e.g. CO2) be non-polar overall?
The molecule is highly symmetrical, so the individual bond dipoles cancel each other out.

Intermolecular Forces: read all 3 questions and answers

State the three types of intermolecular forces in order of increasing strength.
London forces (induced dipole-dipole) < permanent dipole-permanent dipole forces < hydrogen bonds.
Under what conditions does hydrogen bonding occur?
When a hydrogen atom is covalently bonded to a highly electronegative atom (N, O, or F) that has at least one lone pair of electrons.
Why is ice less dense than liquid water?
In ice, hydrogen bonds hold water molecules in a rigid, open hexagonal lattice structure, putting the molecules further apart.