Topic 2 Flashcards
Isoelectronic ions
Ions of different elements with the same electronic configuration (but different ionic radii)
Covalent bond
The electrostatic attraction between 2 nuclei of 2 atoms and the shared pair of electrons between them
Polar covalent bond
Each atom does not have an equal attraction to the bonding electron pair, so there are positive and negative partial charges at the ends of the molecule
Dative covalent bond
An atom donates its lone pair of electrons to another atom/ion, forming a covalent bond in which the pair of electrons comes from one atom
Electronegativity
The tendency of an atom to attract a bonding pair of electrons
Hydrogen bond
The electrostatic attraction between the partial positive charge of a hydrogen atom and the partial negative charge of a lone pair of electrons on an atom of a neighbouring molecule
Metallic bonding
The strong electrostatic attraction between the metal cations and sea of delocalised electrons
Dipole
When two charges of equal magnitude but opposite signs are separated by a small distance
Ionic bond
The strong electrostatic attraction between oppositely charged ions
Factors affecting strength of ionic bonding
Greater charge of anion and cation.
Smaller ionic radius (less shells).
-greater charge density, higher lattice energy, stronger bond
Polyatomic
An ion with more than one atom
Trend in boiling points of G4-H molecules
Boiling point increases down group 4.
-london forces
Lower than G5/6/7 because G4-H molecules are non-polar so only have london forces.
Trend in boiling points of G5/6/7-H molecules
Decreases then increases down the group.
O,F,N form hydrogen bonds- high boiling points.
For rest, boiling point increases because of london forces.
Boiling points vs chain length
Boiling point increases as chain length increases
-london forces
Boiling points vs branching
Branching reduces the contact surface area
-reduces the number of london forces that can form
