Bonding and Chemical Interactions (3) Flashcards
Electronic geometry
describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs.
(CH4 , NH3 and H2O all have tetrahedral electronic geometry but differ in their molecular shapes.)
Molecular geometry
describes the spatial arrangement of only the bonding pairs of electrons.
(CH4 is tetrahedral , NH3 is pyramidal and H2O is bent)
Cation
atoms that lose the electrons
meTals lose electrons to become caTions = + charge
Anion
atoms that gain electrons
(Nonmetals gain electrons to become aNions = Negative (-) charge ).
Formal charge
Valence electrons - bonds - lines
Sigma bonds
result of head to head overlap.
Bond allow for free rotation
Pi bonds
result of the overlap of two parallel electron cloud densities.
Do not allow free rotation because the electron densities of the orbitals are parallel and cannot be twisted.
Periods
rows
Families
column
Effective nuclear charge (Zeff)
a measure of the net positive charge experienced by the outermost electrons
increases left to right
up and right
Atomic and Ionic Radii
refers to the size of a neutral element, while an ionic radius is dependent on how the element ionizes based on its element type and group number
Increases right to left
down and left
Ionization energy (IE)
the energy required to remove an electron from gaseous species.
Removing electrons from an atom requires the input of heat, which makes it an endothermic processes.
Mg –> Mg+ + e- first ionization energy
Mg+ –> Mg2+ + e- second ionization energy
Electron affinity
refers to the energy dissipated by gaseous species when it gains energy. Exothermic process
BEAR (Mneomnic)
B (basicity) E(electonegativity, EA
and IE)
A(acidity) R (radius)
Incomplete octet
elements are stable with fewer than 8 electrons in their valence shell: Hydrogen (2) Helium (2) Lithium (2) Beryllium (4) Boron (6)
Expanded octet
any elements in period 3 and greater can hold more than 8 electrons including:
Phosphorus (10)
Sulfur (12)
Chlorine (14)
Bond order
As bond order increases, bond strength increases, bond energy increases, bond length decreases
triple bond –> bond strength ↑ , bond energy ↑ , bond length ↓
Bond energy
the energy required to break a bond by separating its components into their isolated, gaseous atomic states.
Range of electronegativies
nonpolar bonds –> 0 to 0.5
polar bonds –> 0.5 to 1.7
ionic bonds –> higher than 1.7
Dipole moment Equation
p = qd
q= charge d= distance
VESPER Theory (molecular geometry)
only bonding electron pairs aka bonds
2 ——–> Linear
3 ———> Triagonal planar Bent
4 ———> Tetrahedral Trigonal pyramidal, Bent
5———->Trigonal Bipyramidal Seesaw, T-shaped, Linear
6———-> Octahedral Square pyramidal, square
planar, T-shaped, Linear
Electronic geometry vs Molecular geometry
CH4, NH3, and H2O all have tetrahedral molecular geometry but differ in their molecular shapes.
CH4 is tetrahedral, NH3 is trigonal pyramidal, H2O is bent as molecular shape.
London dispersion force (van der Waals force)
weakest interactions but are present in all atoms and molecules.
Size ↑ , London dispersion force ↑
Dipole-Dipole interaction
occurs between oppositely charged ends of polar molecules, stronger than London forces. Present in solid and liquid phases but no in the gas phase.
Hydrogen bonding
Occurs in FON (phone)
the strong form of dipole-dipole interaction. Strongest intramolecular force and intermolecular.
FON (phone)
Fluorine
Oxygen
Nitrogen
