Molecules of life
No multiple choice questions
Lecture 1 Mayer 24-02-2020 Chapter 1
Organic compounds: C, (N, H, O)
An organic compounds: everything else
Same structure of molecules show similarities in function
Structure of an atom:
• Consists of:
o Nucleus:
▪ Protons +
▪ Neutrons
o Electrons –
• Number of protons = number of electrons
• Electrons are important for chemistry
Periodic table:
• Top left corner: atomic number = number of protons
• Number under: 1 AMU
• 12 C & 13 C → difference in number of neutrons
Carbon dating: 14C 6 protons, 8 neutrons
Electrons are distributed in shells around the core: bohr model
• First shell: 2e-
• Second shell: 8e-
• Third shell: 18e-
• Fourth shell: 32e-
Electrons have wave like properties
Shells contain subshells: atomic orbitals
First shell Second shell Third shell
Atomic orbitals S S, P S, P, D
Number of atomic 1 1, 3 1, 3, 5
orbitals
Maximum number 2e 2, 6 e 2, 6, 10 e
of electrons
Core electrons: electrons in the inner shell
Valence electrons: electrons in the outmost shell
Relative energy of atomic orbitals: 1s < 2s < 2p < 3s < 3p < 3d
Three rules about ground state electron configuration
1. Aufbau principle: electron goes into the atomic orbital with the lowest energy
2. Pauli exclusion principle: no more than two electrons can be in an atomic
orbital
3. Hund’s rule: an electron goes into an empty degenerate orbital rather than
pairing up
, Atom Name of Atomic 1s 2s 2px 2py 2pz
element number
H Hydrogen 1 ↑
He Helium 2 ↑↓
Li Lithium 3 ↑↓ ↑
Be Beryllium 4 ↑↓ ↑↓
B Boron 5 ↑↓ ↑↓ ↑
C Carbon 6 ↑↓ ↑↓ ↑ ↑
N Nitrogen 7 ↑↓ ↑↓ ↑ ↑ ↑
O Oxygen 8 ↑↓ ↑↓ ↑(↓) ↑(↓) ↑(↓)
2px, 2py, 2pz have same energy level → chance that two electrons in one p shell is equal
Only the most right row in the periodic table is stable because it has a closed shell
Octet rule: atoms are most stable when their outmost shell is filled
Hydrogen is very reactive because it can lose an electron or gain an electron
• Lose: proton
• Gain: hydride ion
Electrostatic potential map: you can see if an element is positive or
negative
Ionic interaction / ionic band: combination of ions to adhere to the
octet rule
Sharing electrons results in covalent bonds
• Polar bonds: higher electronegativity
o H-Cl & C=O
• Apolar bonds: dissolves in water
o C-H & H-H
Electronegativity: increases → & ↑
Continuum of bonds:
• Ionic bonds: > 2.0
• Polar bonds: < 2.0 & > 0.5
• Apolar bonds: < 0.5
1 electron means that the molecule can only form one bond
Lone pair: a electron pair that does not have to form a bond with another molecule
• H: 1 bond
• C: 4 bonds
• N: 3 bonds, 1 lone pair
• O: 2 bonds, 2 lone pairs
• F, Br, Cl, I: 1 bond, 3 lone pairs
, Formal charge: number of valence electrons – (number of lone pair electrons + number
of bonds) (formele lading)
Lecture 2 Mayer 25-02-2020 Chapter 8.1-8.5
3D structures of molecules
• Kekulé: perspective
• Ball and stick model
• Electrostatic potential map
S orbitals are spherical
• 2s orbital has a node: place in the orbital where are no
electrons
• Two s orbitals form a sigma bond
Orbitals are conserved – molecular orbitals (MOs)
• Number of atom orbitals combined = the number of molecular
• σ*- antibonding orbitals
• If orbitals are conserved they cannot be destroyed
P orbital resemble dumbbells
• Two p orbitals form a pi bond
• Plates: no electron density
No multiple choice questions
Lecture 1 Mayer 24-02-2020 Chapter 1
Organic compounds: C, (N, H, O)
An organic compounds: everything else
Same structure of molecules show similarities in function
Structure of an atom:
• Consists of:
o Nucleus:
▪ Protons +
▪ Neutrons
o Electrons –
• Number of protons = number of electrons
• Electrons are important for chemistry
Periodic table:
• Top left corner: atomic number = number of protons
• Number under: 1 AMU
• 12 C & 13 C → difference in number of neutrons
Carbon dating: 14C 6 protons, 8 neutrons
Electrons are distributed in shells around the core: bohr model
• First shell: 2e-
• Second shell: 8e-
• Third shell: 18e-
• Fourth shell: 32e-
Electrons have wave like properties
Shells contain subshells: atomic orbitals
First shell Second shell Third shell
Atomic orbitals S S, P S, P, D
Number of atomic 1 1, 3 1, 3, 5
orbitals
Maximum number 2e 2, 6 e 2, 6, 10 e
of electrons
Core electrons: electrons in the inner shell
Valence electrons: electrons in the outmost shell
Relative energy of atomic orbitals: 1s < 2s < 2p < 3s < 3p < 3d
Three rules about ground state electron configuration
1. Aufbau principle: electron goes into the atomic orbital with the lowest energy
2. Pauli exclusion principle: no more than two electrons can be in an atomic
orbital
3. Hund’s rule: an electron goes into an empty degenerate orbital rather than
pairing up
, Atom Name of Atomic 1s 2s 2px 2py 2pz
element number
H Hydrogen 1 ↑
He Helium 2 ↑↓
Li Lithium 3 ↑↓ ↑
Be Beryllium 4 ↑↓ ↑↓
B Boron 5 ↑↓ ↑↓ ↑
C Carbon 6 ↑↓ ↑↓ ↑ ↑
N Nitrogen 7 ↑↓ ↑↓ ↑ ↑ ↑
O Oxygen 8 ↑↓ ↑↓ ↑(↓) ↑(↓) ↑(↓)
2px, 2py, 2pz have same energy level → chance that two electrons in one p shell is equal
Only the most right row in the periodic table is stable because it has a closed shell
Octet rule: atoms are most stable when their outmost shell is filled
Hydrogen is very reactive because it can lose an electron or gain an electron
• Lose: proton
• Gain: hydride ion
Electrostatic potential map: you can see if an element is positive or
negative
Ionic interaction / ionic band: combination of ions to adhere to the
octet rule
Sharing electrons results in covalent bonds
• Polar bonds: higher electronegativity
o H-Cl & C=O
• Apolar bonds: dissolves in water
o C-H & H-H
Electronegativity: increases → & ↑
Continuum of bonds:
• Ionic bonds: > 2.0
• Polar bonds: < 2.0 & > 0.5
• Apolar bonds: < 0.5
1 electron means that the molecule can only form one bond
Lone pair: a electron pair that does not have to form a bond with another molecule
• H: 1 bond
• C: 4 bonds
• N: 3 bonds, 1 lone pair
• O: 2 bonds, 2 lone pairs
• F, Br, Cl, I: 1 bond, 3 lone pairs
, Formal charge: number of valence electrons – (number of lone pair electrons + number
of bonds) (formele lading)
Lecture 2 Mayer 25-02-2020 Chapter 8.1-8.5
3D structures of molecules
• Kekulé: perspective
• Ball and stick model
• Electrostatic potential map
S orbitals are spherical
• 2s orbital has a node: place in the orbital where are no
electrons
• Two s orbitals form a sigma bond
Orbitals are conserved – molecular orbitals (MOs)
• Number of atom orbitals combined = the number of molecular
• σ*- antibonding orbitals
• If orbitals are conserved they cannot be destroyed
P orbital resemble dumbbells
• Two p orbitals form a pi bond
• Plates: no electron density
