GOB Chemistry Exam 2 Study Guide: Covalent Compounds, Organic Chemistry, Thermodynamics, and Chemical Reactions

Getting Covalent Compounds into Shape

VSEPR Theory and Molecular Geometry

The shape of covalent compounds is determined by the arrangement of atoms and electron pairs around a central atom. The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict these shapes.

• VSEPR Theory: Electron pairs around a central atom repel each other, resulting in specific molecular geometries.

• Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Wedge-and-Dash Notation: Used to represent three-dimensional structures in two dimensions, with wedges indicating bonds coming out of the plane and dashes indicating bonds going behind the plane.

• Lewis Structures: Diagrams showing the arrangement of atoms and electron pairs in a molecule.

• Example: Methane ($CH_4$) has a tetrahedral shape due to four bonding pairs around the central carbon atom.

Electronegativity and Molecular Polarity

Molecular polarity depends on both the polarity of individual bonds and the overall shape of the molecule.

• Electronegativity: The ability of an atom to attract electrons in a bond. Differences in electronegativity create polar covalent bonds.

• Bond Polarity: A bond is polar if the atoms have different electronegativities; nonpolar if they are the same.

• Molecular Polarity: Determined by the vector sum of bond dipoles and the molecular geometry.

• Example: Water ($H_2O$) is polar due to its bent shape and the difference in electronegativity between hydrogen and oxygen.

Representing the Structures of Organic Compounds

Lewis, Condensed, and Skeletal Structures

Organic compounds can be represented in several ways to convey different levels of structural detail.

• Lewis Structures: Show all atoms, bonds, and lone pairs.

• Condensed Structures: Group atoms together, omitting some bonds for simplicity (e.g., $CH_3CH_2OH$).

• Skeletal Structures: Lines represent carbon-carbon bonds; hydrogens attached to carbons are usually omitted.

• Example: Ethanol can be drawn as a Lewis structure, condensed as $CH_3CH_2OH$, or as a skeletal structure.

Alkanes: The Simplest Organic Compounds

Alkanes are saturated hydrocarbons containing only single bonds between carbon atoms.

• Saturated Hydrocarbons: Only single bonds; maximum number of hydrogen atoms.

• Unsaturated Hydrocarbons: Contain double or triple bonds; fewer hydrogen atoms.

• First Eight Straight-Chain Alkanes: Methane, ethane, propane, butane, pentane, hexane, heptane, octane.

• Cycloalkanes: Alkanes arranged in a ring structure.

• General Formula: For straight-chain alkanes: $C_nH_{2n+2}$; for cycloalkanes: $C_nH_{2n}$.

• Example: Butane ($C_4H_{10}$) vs. cyclobutane ($C_4H_8$).

Families of Organic Compounds—Functional Groups

Functional Groups in Organic Molecules

Functional groups are specific groups of atoms within molecules that determine the chemical properties of those molecules.

• Common Functional Groups: Alcohols (-OH), carboxylic acids (-COOH), amines (-NH_2), ketones (C=O), aldehydes (CHO).

• Unsaturated Hydrocarbons: Alkenes (C=C), alkynes (C≡C), aromatics (benzene ring).

• Example: Ethanol contains an alcohol functional group.

Nomenclature of Simple Alkanes

Alkanes are named according to IUPAC rules, which provide systematic names based on the number of carbon atoms and branching.

• Branched-Chain Alkanes: Named by identifying the longest chain and the position of branches (substituents).

• Example: 2-methylpropane is a branched-chain alkane.

Isomerism in Organic Compounds

Types of Isomerism

Isomers are compounds with the same molecular formula but different structures or spatial arrangements.

• Structural Isomers: Differ in the connectivity of atoms.

• Conformational Isomers: Differ by rotation around single bonds.

• Cis-Trans Isomers: Occur in cycloalkanes and alkenes due to restricted rotation; cis (same side), trans (opposite sides).

• Chiral Centers: Carbon atoms bonded to four different groups, leading to non-superimposable mirror images (enantiomers).

• Example: 1,2-dichloroethene has cis and trans isomers.

Thermodynamics

Heat, Energy, and Spontaneity

Thermodynamics studies the flow of energy and the spontaneity of chemical reactions.

• Spontaneity: Determined by the Gibbs free energy change ($riangle G$).

• Exergonic Reactions: Release energy; $riangle G < 0$.

• Endergonic Reactions: Absorb energy; $riangle G > 0$.

• Calorimetry: Technique to measure heat changes in chemical reactions.

• Energy Content in Food: Calculated using nutritional information and conversion factors.

• Example Equation: $riangle G = riangle H - T riangle S$

Chemical Reactions: Kinetics

Reaction Rates and Catalysts

Kinetics examines the speed of chemical reactions and the factors that influence it.

• Activation Energy: Minimum energy required for a reaction to occur.

• Factors Affecting Rate: Temperature, concentration of reactants, presence of a catalyst.

• Enzymes: Biological catalysts that increase reaction rates by lowering activation energy.

• Example Rate Equation: $ext{Rate} = k[ ext{Reactant}_1][ ext{Reactant}_2]$

Overview of Chemical Reactions

Types and Classification of Reactions

Chemical reactions are classified by the changes that occur in reactants and products.

• Synthesis: Two or more substances combine to form one product.

• Decomposition: One substance breaks down into two or more products.

• Exchange (Single/Double Replacement): Atoms or ions are exchanged between compounds.

• Reversible vs. Irreversible Reactions: Reversible reactions can proceed in both directions; irreversible reactions go to completion.

• Combustion of Hydrocarbons: Hydrocarbon reacts with oxygen to produce carbon dioxide and water. General Equation: $C_xH_y + O_2 ightarrow CO_2 + H_2O$

• Organic vs. General Chemical Equations: Organic equations focus on changes in organic molecules; general equations apply to all chemical reactions.

Oxidation and Reduction

Redox Reactions in Inorganic and Organic Chemistry

Oxidation-reduction (redox) reactions involve the transfer of electrons between substances.

• Oxidation: Loss of electrons or increase in oxidation state.

• Reduction: Gain of electrons or decrease in oxidation state.

• Identifying Redox Changes: In inorganic reactions, track electron transfer; in organic reactions, look for changes in bonds to oxygen or hydrogen.

• Example: Oxidation of ethanol to acetaldehyde.

Organic Addition Reactions to Alkenes

Hydrogenation and Hydration of Alkenes

Alkenes undergo addition reactions where small molecules add across the double bond.

• Hydrogenation: Addition of hydrogen ($H_2$) to an alkene, converting it to an alkane.

• Hydration: Addition of water ($H_2O$) to an alkene, forming an alcohol.

• General Reaction: $RCH=CHR' + H_2 ightarrow RCH_2CH_2R'$ (hydrogenation)

• General Reaction: $RCH=CHR' + H_2O ightarrow RCH(OH)CH_2R'$ (hydration)

• Example: Ethene ($C_2H_4$) reacts with $H_2$ to form ethane ($C_2H_6$).