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What is the primary concept of organic chemistry?
What is the primary concept of organic chemistry?
Study of the structure, properties, composition, reactions and preparations of carbon-containing compounds.
What is the 'octet rule'?
What is the 'octet rule'?
An atom will attempt to gain, lose, or share valence electrons to have 8 electrons in its outer shell, making it more stable.
A single bond is a sigma bond.
A single bond is a sigma bond.
True
Select all the possible types of hybrid orbitals.
Select all the possible types of hybrid orbitals.
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What is the number of sp3 hybrid orbitals that are formed from a 2s and three 2p atomic orbitals?
What is the number of sp3 hybrid orbitals that are formed from a 2s and three 2p atomic orbitals?
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Rotation around a C=C bond in ethylene is unrestricted.
Rotation around a C=C bond in ethylene is unrestricted.
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A triple bond comprises of one sigma and three pi bonds.
A triple bond comprises of one sigma and three pi bonds.
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What is the hybridization of carbon in ethane?
What is the hybridization of carbon in ethane?
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Electronegativity is a measure of an atom's attraction for electrons in a nucleus.
Electronegativity is a measure of an atom's attraction for electrons in a nucleus.
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Carbon-carbon bonds are nonpolar.
Carbon-carbon bonds are nonpolar.
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Carbon-hydrogen bonds are polar.
Carbon-hydrogen bonds are polar.
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A bond formed between two atoms of different electronegativities is considered polar.
A bond formed between two atoms of different electronegativities is considered polar.
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The 8- symbol indicates an electron rich atom.
The 8- symbol indicates an electron rich atom.
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A polar molecule has at least one polar bond.
A polar molecule has at least one polar bond.
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An 'inductive effect' is the pull of electron density through sigma bonds, caused by electronegativity differences of atoms.
An 'inductive effect' is the pull of electron density through sigma bonds, caused by electronegativity differences of atoms.
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An inductive effect is the pull of electron density through pi bonds.
An inductive effect is the pull of electron density through pi bonds.
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What is the term used to describe the combination of two or more atomic orbitals to form new orbitals with the same shape and energy?
What is the term used to describe the combination of two or more atomic orbitals to form new orbitals with the same shape and energy?
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Increased percent s- character results in a longer bond length.
Increased percent s- character results in a longer bond length.
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What does the symbol 8+ indicate in a molecule?
What does the symbol 8+ indicate in a molecule?
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The concept of '_____' relates to a partial separation of charge in a molecule.
The concept of '_____' relates to a partial separation of charge in a molecule.
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Study Notes
Organic Chemistry Definition
- Study of the structure, properties, composition, reactions, and preparations of carbon-containing compounds, including hydrocarbons and compounds containing other elements (H, N, O, halogens, phosphorus, silicon, sulfur, etc.)
- Most compounds contain a hydrogen-carbon bond (American Chemical Society)
Reaction Example ("Dehydration")
- A similar dehydration reaction occurs regardless of the molecule, involving the loss of water (H₂O)
- This reaction is seen in the synthesis of natural products (e.g., vitamin A and patchouli alcohol)
Additional Course Information
- Lectures will correspond to the textbook.
- Chapter 1 of the textbook will be reviewed (covering atoms, molecules, bonding, etc.)
- Chapter 2 will not be fully covered; some basic concepts will be reviewed later.
Atomic Structure 1
- The nucleus contains protons (positive charge) and neutrons (no charge).
- The electron cloud surrounds the nucleus and contains electrons (negative charge).
Atomic Structure 2
- Atomic number = number of protons
- Mass number = number of protons + neutrons
- In a neutral atom, the number of protons equals the number of electrons.
- The atomic weight of an element is the weighted average of the masses of its isotopes, reported in atomic mass units (amu).
Ions
- Cations are positively charged ions with fewer electrons than protons.
- Anions are negatively charged ions with more electrons than protons.
Isotopes
- Isotopes are atoms of the same element with different numbers of neutrons.
- Isotopes have different mass numbers.
Periodic Table
- Elements in the same row are similar in size.
- Elements in the same column have similar electronic and chemical properties
- Carbon is located in the second row, group 4A
Atomic Orbitals
- s orbitals - spherical electron density; lower energy than other orbitals in the same shell.
- p orbitals - dumbbell shape; contains a node (no electron density) at the nucleus; higher energy than s orbitals in the same shell.
Periodic Table (Row 1)
- The first shell has only one orbital (1s orbital).
- Each shell can hold a maximum of two electrons.
- Hydrogen (H) and Helium (He) are in the first row.
Periodic Table (Row 2)
- Each element has a single 2s orbital and three 2p orbitals.
- The second shell can hold a total of eight valence electrons.
Bonding
- Bonding is the joining of two atoms in a stable arrangement.
- Atoms gain, lose, or share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table
- Ionic bonding - electrons are transferred.
- Covalent bonding - electrons are shared between two nuclei.
- Covalent bonds occur between atoms with similar electronegativity.
- Compounds with covalent bonds are called molecules.
Covalent Bonding 1
- Covalent bonds occur with elements with similar electronegativities (nonmetals).
Covalent Bonding 2
- Hydrogen forms one covalent bond.
- When two hydrogen atoms bond, each achieves a filled valence shell of two electrons.
Valence Electrons
- Second-row elements have a maximum of 8 electrons.
- Atoms with 1-4 valence electrons form 1-4 bonds, respectively.
- Atoms with 5 or more valence electrons form enough bonds to complete their outer shell.
Nonbonded Electrons
- Unshared electrons are also called lone pairs.
Lewis Structures
- Lewis structures are electron dot representations of molecules.
- Solid lines represent two-electron covalent bonds.
- Organic chemistry often involves covalent bonds between carbon atoms and other atoms (H, N, O, halogens) , forming single and multiple bonds.
Formal Charge
- Formal charge is the charge assigned to an atom in a Lewis structure.
- Calculated by subtracting the number of unshared electrons and one half the bonding electrons, from the number of valence electrons.
Electron Ownership
- Electrons "owned" by atoms are determined by the bonds and lone pairs of electrons.
Isomers
- Isomers are different molecules with the same molecular formula.
- Some types of isomers are constitutional isomers
Exceptions to the Octet Rule
- Elements in Groups 2A and 3A can sometimes have fewer or more than 8 valence electrons.
- Elements in the third row can have more than 8 valence electrons.
Resonance
- Some molecules are better represented by a combination of Lewis structures (resonance structures).
- A double-headed arrow separates resonance structures.
- Resonance structures differ only in the arrangement of electrons, not atoms.
- The true structure is a composite of resonance structures (a resonance hybrid).
- Some atoms can delocalize electrons to achieve more stability.
- Rules of Resonance:
- Resonance structures differ only in the position of multiple bonds and nonbonded electrons
- Resonance structures must have the same number of unpaired electrons
- Resonance structures must be valid Lewis structures for the molecule. Hydrogen must have two electrons, a second row element can have no more than eight.
Curved Arrow Notation
- Curved arrows show the movements of electron pairs
- Tail of arrow points at the electron pair in the molecule while the head points to the position where the electron pair is moving/shifting.
Atoms Without Octets
- Some resonance structures can show an atom with less than 8 valence electrons.
- A second-row element cannot have more than 8 valence electrons in a resonance structure
Occurrences of Resonance
- Resonance structures can occur when a lone pair is on an atom directly bonded to a double bond.
- The position of double bonds and lone pairs can differ in the resonance forms.
Resonance Hybrids
- The resonance hybrid is a combination of all resonance structures.
- If two resonance structures are different, the hybrid will resemble the “better” resonance structure (major contributor).
Determining Molecular Shape and Geometry 1
- Two factors define a molecule's structure: bond length and bond angle.
- Bond length decreases across a row in the periodic table and increases down a column.
Determining Molecular Shape and Geometry 2
- Table of Average Bond Lengths (providing data for various bonds, including H-H, C-H, C-F, etc.).
Molecular Geometry
- Valence Shell Electron Pair Repulsion (VSEPR) theory - the most stable molecular arrangement keeps groups as far away from each other as possible.
- Number of groups surrounding an atom determines its geometry
- 2 groups- linear
- 3 groups - trigonal planar
- 4 groups - tetrahedral
Two Groups Around an Atom
- Linear
Three Groups Around an Atom
- Trigonal planar
Four Groups Around an Atom
- Tetrahedral
Drawing Three-Dimensional Structures
- Solid lines: Bonds in the plane
- Wedges: Bonds in front of the plane of the paper
- Dashed lines: Bonds behind the plane of the paper
Equivalent Representations for Methane (3-D Orientation)
- Multiple representations of a molecule are considered valid if they share the same connections.
Wedges and Dashed Wedges
- Wedges and dashed wedges represent connections of atoms that are in the front and back of the plane of the drawing, respectively.
A Nonbonded Electron Pair as a Group
- In ammonia, one of four groups surrounding the central N is a lone pair of electrons to determine geometry
- Molecular shape of ammonia is trigonal pyramidal, group geometry is tetrahedral.
3-D Structure of Water
- In water, two of four groups surrounding the central O atom are lone pairs
- Group geometry is tetrahedral, molecular shape is bent.
Varying Bond Angles
- Bond angle differs from expected tetrahedral structure due to repulsion of lone pairs.
Summary: Predicting Geometry Based on Number of Groups
- Table summarizing geometries for specific number of groups.
Language of Organic Chemistry
- Condensed structure
- Skeletal structure
- Hybridization
Drawing Organic Molecules—Condensed Structures
- Omitting bond lines and showing similar groups bonded to the same atom in parentheses.
Examples of Condensed Structures
- Examples of molecules in condensed structure.
Condensed Structures with C=O
- Condensed structures for molecules with a carbon-oxygen double bond
Skeletal Structures 1
- Assume a carbon atom at the intersection or end of lines in skeletal structure; there are enough hydrogens for each carbon to be tetravalent.
Skeletal Structures 2
- Interpret the skeletal structures to determine the number of atoms and lone pairs
Skeletal Structures with Charged Carbon Atoms
- A charge on a carbon atom takes a hydrogen atom's place.
- Negatively charged carbon atoms have one lone pair.
- Positively charged carbon atoms have none.
Lone Pairs on Heteroatoms
- Lone pairs (on heteroatoms), often omitted in skeletal structures, can be inferred by using the formal charge.
Orbitals and Bonding: Hydrogen
- 1s orbital overlap forms a σ bond between H atoms.
- The σ bond is cylindrically symmetrical.
Orbitals and Bonding: Methane
- Carbon in methane uses four sp³ hybrid orbitals and creates four sigma bonds.
- sp³ hybrid orbitals have one large and one small lobe.
- The sp³ hybrid orbitals are oriented toward the vertices of a tetrahedron, creating four covalent bonds (one hydrogen per bond)
Hybrid Orbitals
- Hybridization is the combining of atomic orbitals to form new hybrid orbitals.
- Four sp³ orbitals form to allow four covalent bonds
- Four sp³ orbitals are oriented toward the corners of a tetrahedron.
Shape and Orientation of sp³ Hybrid Orbitals
- sp³ orbitals have one large and one small lobe.
- Shape of sp³ orbitals is tetrahedral
Bonding Using sp³ Hybrid Orbitals
- Each carbon bond in CH4 is formed by an overlap of an sp³ hybrid orbital on carbon with a 1s orbital of hydrogen.
Other Hybridization Patterns
- Different hybridization forms for two and three groups
- 2 groups- sp
- 3 groups - sp2
Determining Hybridization
- The number of groups around a given atom correspond to the number of orbitals that must be hybridized
Hybrid Orbitals of NH3 and H2O
- Hybrid orbital depictions of ammonia and water.
Hybridization and Bonding in Ethane
- Drawing and explanation of bonding in ethane, illustrating the sp³ overlap.
- Rotation occurs around a σ bond.
Energy Levels and Orbitals
- Diagram showing energy levels and orbitals of different electronic groups from periodic table
Ethane, CH3-CH3
- Rotation can occur around the C-C σ bond.
Hybrid Orbitals in Ethylene
- Each carbon is trigonal planar and sp² hybridized.
- sp² hybridization has three equivalent sp² hybrid orbitals and an unhybridized p orbital.
σ and π Bonds in Ethylene
- Overlap of sp² hybrid orbitals forms σ bonds (between atoms.)
- Overlap of unhybridized p orbitals creates π bonds. The π bond is above and below the molecule.
No Free Rotation in Ethylene
- Rotation about the C=C double bond is restricted
- requires breaking the π bond while reforming it, a significant energy cost.
sp Hybrid Orbitals
- sp hybridization of a carbon atom in an example like acetylene, shows two sp-orbital hybridization and two unhybridized p-orbitals.
Acetylene (Ethyne)
- Each carbon has two sp hybrid orbitals and two unhybridized 2p orbitals.
- The C-H bonds and the C-C bond are σ bonds.
Triple Bonds
- The overlap of two sets of 2p orbitals create two π bonds. The σ bond forms from the overlap of sp-hybridized orbitals
- The triple bond is composed of one σ and two π bonds.
Summary of Bonding in Acetylene 1 & 2
- Summaries of bonding characteristics (number of groups bonded to C, hybridization, bond angle) in various carbon compounds (ethane, ethylene, and acetylene).
Bond Length and Bond Strength
- The number of electrons between two nuclei is directly related to bond length and strength.
- Triple bonds are shorter and stronger than double bonds, double bonds are shorter and stronger than single bonds.
- The bond strength increase and the bond length decrease as the number of bonds between two nuclei increase.
Carbon-Hydrogen Bonds 1 & 2
- The length and strength of C-H bonds vary based on the hybridization of the carbon atom.
- Tables of bond lengths and strengths for C-H and C-C bonds in different compunds (ethane, ethylene, acetylene).
Percent s-Character
- Percent s-character of different types of hybridized orbitals (sp, sp², sp³).
- Increased s-character corresponds to increased bond strength, decreased length.
Electronegativity
- Electronegativity values for various elements shown in a table format (Figure 1.12) Explaining what Electronegativity is and why it is so important.
Bond Polarity
- Electronegativity difference indicates if electrons are shared equally or unequally.
- Non-polar bonds - electrons are shared equally.
- Polar bonds - electrons are shared unequally.
Nonpolar Bonds
- Carbon-carbon bonds and carbon-hydrogen bonds are usually considered nonpolar because of the similar electronegativities of the atoms involved.
Polar Bonds
- Atoms with different electronegativities form polar bonds.
- The element with higher electronegativity pulls electrons closer, creating a partial negative charge (δ-).
Depicting Polarity
- Drawing bond dipoles to show the direction of charge shift in a polar bond.
Polarity of Molecules
- Procedure to determine the net polarity of a molecule: identify all polar bonds and their directions; determine bond geometry and decide if bond dipoles cancel or reinforce each other.
Nonpolar Molecules
- Nonpolar molecules either have only nonpolar bonds or have polar bonds that cancel each other out.
Electrostatic Potential for Polar and Nonpolar Molecules
- Visualization of molecular dipoles using electrostatic potential plots.
Polarity Can "Indicate" Acidity
- Polarity can be used as an indication of acidity (referring to Chapter 2 contents).
Inductive Effects
- Electron pull through sigma bonds caused by electronegativity difference (inductive effect)
- More electronegative atoms stabilize high electron density regions.
- The inductive effect becomes stronger as the electronegativity of the atom increases, and the closer it is to the site of the negative charge.
Inductive Effects in Trifluoroethanol
-The acidity of Trifluoroethanol is higher than that of ethanol due to the stabilizing effect of electronegative fluorine atoms on the conjugate base
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Description
Test your knowledge on the foundational concepts of Organic Chemistry in this quiz based on Chapter 1 of the textbook. You'll explore topics such as nomenclature, reactions, and stereochemistry, enhancing your understanding of carbon-containing compounds. Get ready to recognize functional groups and their interactions!
