Bonding Flashcards
1
Q
Orbitals
A
- Orbitals are representations of the area
of space occupied by an electron (the
likelihood of finding an electron in a
given space)
2
Q
Pauli Exclusion Principle:
A
- A maximum of two electrons can
occupy any one orbital (must have opposite charge)
3
Q
Carbon
A
- Atomic number 6
- Got 4 valence electron
- Every stable covalent bond must have 2 electrons
4
Q
(Hund’s Rule)
A
- For orbitals of the same energy a single electron is added to each before pairing begins.
–This is to avoid e to e replusion
5
Q
Valence-Shell Electron-Pair Repulsion
A
- In a molecule all groups that are attached to a single atom should be as far away from each other in space as possible to minimize electron repulsions.
6
Q
Methane (CH2)
A
C has 4 valence electron
2 will bond in 1S and 2S
2 will go to 2px, 2py and H will bond here with their electron
Two of carbon’s valence electrons are in a filled 2s orbital and can’t participate in bonding. We have the 2px and 2py orbitals, each with 1 electron, which can covalently
bond with two hydrogens
7
Q
Hybridization
A
- The atomic orbital of carbon mix to form a set of four hybrid orbital that we call sp3 orbitals
They are equal in energy and are directed to the corner of a regular tetrahedron (109.5 degree) - 2s, 2px,2py, 2pz will turn to 4 sp3 when hybridized (when you combine 4 orbitals you will end up with 4 new orbitals)
8
Q
Bonds
A
- When orbitals overlap, electrons will be shared by both atoms. This will form MO (molecular orbitals)
- The number of AOs or hybrids is equal to the number of MOs
- There will be an MO lower in energy to a constituent orbital (bonding MO) and the other higher in energy (anitibonding MO)
- A covalent bond is formed whenever two electrons occupy a molecular orbital
- There need to be two orbitals overlapping to be called sigma bond is formed)
- Anti-bonding orbital has no electron in a stable compounds but it still exist and IS able to accept e under certain circumstances
- Having e in anti-bonding orbitals make compounds unstable.
9
Q
Nitrogen
A
- Got 5 valence electrons
- In sp3 hybrid orbital the electron is placed in each orbital until they all have one then the remaining electron is paired up in one of the orbital
- The filed orbital is called lone pair (push down on bonds)
– NH3 got 110.6 degree because of lone pair - Lone pair can interact with vacant orbital (since the e both originally belonged to nitrogen but are now shared this mean that the N is now electron deficient and DEVELOPS a Positive charge)
10
Q
Oxygen
A
- Got 6 valence electrons
- It got two lone pairs (filled sp3) and two singly occupied orbitals that can overlap
- H2O got 104.5 degree
11
Q
Fluorine
A
- Got 7 valence electrons
- 3 lone pairs and F will only be able to form one covalent bond
12
Q
Ethylene
A
- 3 atoms bonded to carbon: needs to stay as far as possible from each other (120degree)
- A trigonal planar does that and keep them at 120 degree
- An ethylene is two carbons bound together and have only 2H on each
13
Q
Hybridization
A
- sp3 will not give us a trigonal planar arrangement so an alternative hybridization is required
- We need 3 hybrid orbital so we mix 3 atomic orbital (2s+2px+2py) and we get 3 sp2 hybrid orbital (gonna get an equilateral triangle 120 degree)
- Each sp2 will allow for forming three single bonds)
- The remaining orbital is carried unchanged (2pz) and it will remain perpendicular to the plane formed by the sp2
14
Q
Hybridization (2pz)
A
- Orbital that exist in perpendicular planes cannot interact with one another (stay independent of the hybrid orbital)
- Form a double bond (pi bond) with another C in C2H4 as they are next to each other and perpendicular to the rest)
- Double bonds are more reactive than single bonds
15
Q
Double bonding
A
- Form from the mixing of two Pz orbitals and they form two set of new MO
- The lower energy MO is called bonding pi orbital
- The higher energy (vacant) is called an anti bonding pi-orbital
