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CHM 2210 - Organic Chemistry I
Learning Objectives – Chapter 1, Chapter 2, and Chapter 3
- Chapter 1 -
1. Draw constitutional isomers for a given molecular formula
Constitutional isomers: compounds with the same molecular formula but =/=
structures
2. Define and give examples of covalent bonds, valence electrons, lone pairs, formal
charges, and the octet rule
• covalent bondsà pair of elec. SHARED btwn 2 atoms (attractive forces btwn
+ nuclei & - elect./repulsive forces btwn both + nuclei and – elec. Keep the
bond at optimal length)
• valence electrons à elec. In the utermost region of atoms that enters into
the formation of chemical bonds.
• lone pairs à a pair of valence electrons that are not shared with another
atom in a covalent bond
• octet rule à toms tend to form compounds in ways that give them eight
valence electrons/ full valence shell
• formal charges = valence elec. – lines (bonds) – dots (lone electrons)
-F.C: 1 extra val. electron
+ F-C: “lacking” 1 val. e.
3. Draw Lewis structures of neutral and charged species
1) Draw the individual atoms using dots to represent the valence electrons.
2) Put the atoms together so they share pairs of electrons to make complete
octets.
, 2
4. Describe the relationship between electronegativity, inductive effects, and the
three main types of bonding (nonpolar covalent, polar covalent, and ionic)
The +. à + polar bond
• Induction: pull of electron density
5. Determine the number of carbon atoms and hydrogen atoms in a bond-line
structure
6. Describe the shape and phase of s and p atomic orbitals and the process of filling
orbitals with electrons
Filling orbitals:
- Aufbau Principle: start at lowest E.level à up
- Pauli Exclusion P.: 2 elec. In = orbital à =/= directions
- Hund’s Rule:
CHM 2210 - Organic Chemistry I
Learning Objectives – Chapter 1, Chapter 2, and Chapter 3
- Chapter 1 -
1. Draw constitutional isomers for a given molecular formula
Constitutional isomers: compounds with the same molecular formula but =/=
structures
2. Define and give examples of covalent bonds, valence electrons, lone pairs, formal
charges, and the octet rule
• covalent bondsà pair of elec. SHARED btwn 2 atoms (attractive forces btwn
+ nuclei & - elect./repulsive forces btwn both + nuclei and – elec. Keep the
bond at optimal length)
• valence electrons à elec. In the utermost region of atoms that enters into
the formation of chemical bonds.
• lone pairs à a pair of valence electrons that are not shared with another
atom in a covalent bond
• octet rule à toms tend to form compounds in ways that give them eight
valence electrons/ full valence shell
• formal charges = valence elec. – lines (bonds) – dots (lone electrons)
-F.C: 1 extra val. electron
+ F-C: “lacking” 1 val. e.
3. Draw Lewis structures of neutral and charged species
1) Draw the individual atoms using dots to represent the valence electrons.
2) Put the atoms together so they share pairs of electrons to make complete
octets.
, 2
4. Describe the relationship between electronegativity, inductive effects, and the
three main types of bonding (nonpolar covalent, polar covalent, and ionic)
The +. à + polar bond
• Induction: pull of electron density
5. Determine the number of carbon atoms and hydrogen atoms in a bond-line
structure
6. Describe the shape and phase of s and p atomic orbitals and the process of filling
orbitals with electrons
Filling orbitals:
- Aufbau Principle: start at lowest E.level à up
- Pauli Exclusion P.: 2 elec. In = orbital à =/= directions
- Hund’s Rule:
