Bonding and molecular structure Flashcards
Schrödinger’s equation
Used to show that electrons behave in a wave-like manner.
Used to show the probability of finding an electron at a point in space.
Only applicable to Hydrogen.
Principal quantum number (n)
Represents the energy and size of the atomic orbital.
Orbitals within the same “n” are in the same energy level.
Only integer values are used.
The angular momentum quantum number (l)
Represents different shapes with different orbits.
Restricted by “n” as l = 0 to (n-1)
“n” values and their effect on the l value and the corresponding orbit:
n = 1
n = 2
n = 3
n = 4
⠀
⠀
l = 0 = s (sharp) orbital
l = 0,1 = s and p (principal) orbitals
l = 0,1,2 = a,p,d (diffuse)
l = 0,1,2,3 = s,p,d,f (fundamental)
The magnetic quantum number (mₗ)
Represents the orientation of orbitals in space.
-l to +l
if n = 2, l = 1, and mₗ = -1, 0, 1
The spin quantum number (mₛ)
Each electron has an opposite spin.
Values are -1/2 and 1/2.
s orbital
Spherical, the larger the value of n, the bigger the orbital.
p orbital
Bi-lobed / dumbbell-shaped, they have 3 different orientations (called degenerate orbitals).
py, px, pz
Energy conversion
As the value for n gets higher, the energy levels converge.
Periodicity as you go across the PT:
Atomic radii
Electronegativity
⠀
Decreases (stronger nucleus charge)
Increases (stronger nucleus charge)
Ionic bonding
Positively and negatively charged ions held together in a lattice.
Involves a transfer of electrons and then they stay attracted due to strong, opposite charges.
Lewis structures
Draw 3 random ones out.
// go revise topic 14.1
Formal charge
Number of valence electrons - 1/2 x bonding electrons - number of non-bonding electrons
The most likely structure is the one that is most correct (ie has the correct charge on the molecule and also has the - or + charge on the most or least electronegative atom)
Polarity
A bond is polar if its centres of positive and negative charge do not coincide.
Hybridisation
The overlap of different orbitals.
It causes the s orbital to raise up to the p orbital causing sp¹/²/³ orbitals to formed
Sigma bonds (σ)
Caused by the axial overlap of orbitals.
Can be done with two s orbitals, one s orbital and one p orbital, or two p orbitals
Single bonds
Pi bonds (π)
Caused by the side-to-side overlap of orbitals.
Occurs only with two p orbitals.
Double bonds, triple bonds.
NH₃ hybridisation
Has sp³ hybridisation because the lone pair can form a dative covalent bond so it behaves
Amine
A nitron bonded to a carbon
Amide
A nitrogen bonded to the same carbon that an oxygen is also double bonded to.
R-C=O
⠀⠀⠀\N
Imine
Nitrogen has a double bond with the carbon(s)
R=N-R
Why hydrogen bonding occurs so easily and strongly
It is the only atom present in molecules for which the only electrons are valence electrons;
Thus:
* Hydorgen’s electrons are located between the bonded atoms, so the positive charge is relatively exposed, attracting a negative charge.
The relative length of hybrid “s” orbitals.
As the “s” character of a hybrid orbital increases, the electrons are held closer to the nucleus: sp hybrid orbitals hold electrons closer to the nucleus than sp³ hybrid orbitals
Strength of acids formed by hybridised carbons
Acidic character is judged by the stability of the anion formed after releasing a proton.
As the “s” character increases, the easier it is to remove a proton, the carbanion is stronger and more stable, and it is more acidic (it holds tighter to the lone pairs)
Imines:
⠀
* hybridisation
* Geometry
* σ bond overlap
* π bond overlap
* Lone pair location
* Basic in comparison to amines
Imines:
⠀
* Both C & N - sp2 hybridised
* Both planar geometry
* σ from sp2 orbital overlap
* π from remaining p orbital overlap
* Lone pair is in an sp2 orbital;
* Imine is less basic than amines (electron pair less available)
Nitriles:
⠀
* hybridisation
* Geometry
* σ bond overlap
* π bond overlap
* Lone pair location
* Basic in comparison to amines
Nitriles:
⠀
* Both C & N - sp hybridised
* Both linear geometry
* σ from sp orbital overlap
* 2π from two 2p orbital overlaps
* Lone pair is in an sp orbital;
* Nitrile is less basic than amines and nitriles (electron pair less available)
Imines:
⠀
* hybridisation
* Geometry
* σ bond overlap
* π bond overlap
* Lone pair location
* Basic in comparison to amines
Imines:
⠀
* Both C & N - sp2 hybridised
* C - planar geometry, N - bent
* σ from sp2 orbital overlap
* π from remaining p orbital overlap
* Lone pair is in an sp2 orbital;
* Imine is less basic than amines (electron pair less available)
Alcohol:
- Geometry:
- Lone pair availability:
- Bond polarity
Alcohol:
- C, tetrahedral; O, bent
- in sp³ orbitals, OH can be a base
- O is more electronegative than C/H
Ketones and aldehydes:
- Geometry:
- hybridisation
- Lone pair availability
- Location of lone pairs
- Formation of σ bond
- Formation of π bond
Alcohol:
* C = trigonal planar
* Both have sp² hybridisation
* sp² orbitals, less available.
* Lone pairs are in 2 sp² orbitals on O
* sp² orbitals overlap
* p orbitals on both C and O overlap
Macromolecule
A molecule comprised of many molecules
Intramolecular interactions
Influence the shape and stability of a macromolecule:
Ionic bonding, covalent bonding
Intermolecular interactions
Affect physiochemical properties and are the basis for selective molecular recognition between molecules:
Dispersion, dipole-dipole, hydrogen
Dispersion forces
Caused when one molecule is near another molecule.
Since electrons are always moving, at any given point, the electron density will be closer to one atom out of the two atoms at any specific point.
This will cause a temporary dipole to be induced in a nearby molecule and cause a very weak, basic attractive force to exist
Dipole-dipole forces
Occurs between polar molecules.
Molecules line up to satisfy intermolecular attraction: atoms that are δ+ in one molecule will line up next to atoms that are δ⁻ in another molecule.
This will form a relatively strong attraction that is weaker than hydrogen bonding but stronger than dispersion forces
Constructive interference
Ideal for bonding: the + and - waves are aligned and there are electrons in the molecular orbital (a sigma bond (σ))
Destructive interference
Not ideal for bonding: the + and - waves are aligned and there are no electrons in the molecular orbital (an antibonding sigma bond (σ*))
Electrons in σ orbitals: what does it do to the species?
Stabiilise a species
Electrons in σ* orbitals: stabilises or destabilise a species?
Destabilise a species
When can a species not exist
If both σ and σ* bonds are fully occupied and there is no net stabilisation
When can a species exist
If the σ bond contains more electrons than the σ* bond, and there is net stabilisation
Key parts to an energy diagram regarding bonding and antibonding orbitals
1) π/π* 2px & π/π* 2py are degenerate (same);
2) The π/ π∗ energy gap is less than the σ/σ* gap – more efficient overlap than pz orbitals;
3) electrons are filled from the bottom up, filling the bonding orbitals first.
Paramagnetism
Occurs as a species is attracted to a magnetic field due to having unpaired electrons
Diamagnetism
Occurs as a species is weakly repelled by a magnetic field due to having only paired electrons
