Chapter 4 Flashcards
When a main group element forms a cation
,it tends to lose all its valence
electrons
When a transition metal forms a cation
it tends to lose its outermost s
electrons first, then sometimes lose one or two additional d electrons
from the next-to-outermost shell.
The bond length is determined by
the distance at which the lowest potential energy is achieved
Breaking a chemical bond
takes energy (endothermic process)
Forming a chemical bond
releases energy (exothermic process)
A pure covalent bond is
a covalent bond are identical (diatomic)
A polar covalent bond is
when one atom has a partial negative charge and the other has a partial positive charge
Electronegativity is
how much an atom “wants” electrons.
In a polar bond, the more electronegative atom
pulls the electrons closer, creating a positive (δ+) and a negative (δ-) charge on the atoms.
The greater the difference
in electronegativity
The more polar it is
Electronegativity
increase across a period, and a
decrease down a group
Electron Affinity
measures the energy change when an electron is added to an atom
The lattice energy (∆Hlattice ) of an ionic compound
is
a measure of the strength of the attraction between its positive and negative ions
Exceptions to the octet rule:
1. Odd-electron molecules
only get each atoms as close
as possible to a full octet
Exceptions to the octet rule:
2. Electron-deficient molecules (the central atom doesn’t have an octet)
a central atom from group 2 or 13
Exceptions to the octet rule:
3. Hypervalent molecules (the central atom has more than an octet)
from period 3 and higher,
Formal charge =
valence electrons – # lone pair electrons – ½ # bonding electrons
The shorter the bond..
The stronger it is
Bond energy
decreases down a group
increases across a period (left to right)
increases with the number of bonds
Bond length
increases down a group
decreases across a period (left to right)
decreases with the number of bonds
Bond length Exception:
H-F is longer than
H-H, but stronger
Enthalpy change
Δ𝐻 = bonds broken-bonds formed
Electron-pair geometry
includes all electron pairs
(lone pairs and/or bonds)
Molecular structure
Only includes bonds and
atoms (no lone pairs)
lone pair-lone pair
lone pair-bonding pair
bonding pair-bonding pair
most repulsion
second most
less repulsion
lone pair
triple bond
double bond
single bond
most space
second most
third most
least space
Polar covalent bonds connect two atoms with differing electronegativities is called
bond dipole moment
A bond dipole moment is
represented with an arrow pointing towards the more EN atom.
f the overall dipole moment is
zero, the molecule is..
If the overall dipole moment is
non-zero, the molecule is…
non-polar.
polar.
he strength of a covalent bond depends on
the extent of overlap of the
orbitals involved.
A sigma bond (σ bond) is
a covalent bond where the electron density is focused along the axis that connects the nuclei of the two atoms.
A pi bond (π bond) is
a type of covalent bond formed when two p orbitals overlap side by side. The overlap happens on either side of the line connecting the two nuclei, and there is no electron density along that line.
which bond is stronger
π bonds are generally weaker than σ bonds for the same pair of atoms,
because they have less orbital overlap
sp hybridization
LINEAR
This process took one s orbital and one
p orbital to form two sp orbitals
An sp orbital has 50% s character, and
50% p character
sp2 hybridization
TRIGONAL PLANAR
This process took one s
orbital and two p orbitals to
form three sp2 orbitals
An sp2 orbital has 33% s
character, and 66% p
character
sp3 hybridization
TETRAHEDRAL
This process took one s orbital
and three p orbitals to form
four sp3 orbitals
An sp3 orbital has 25% s
character, and 75% p character
Hybrid orbitals overlap to form
Unhybridized orbitals overlap to
form
σ bonds.
π bonds.
