GOB Chemistry Exam 2 Study Guide: Covalent Compounds, Organic Structures, Thermodynamics, Kinetics, and 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 molecular shapes based on the repulsion between electron pairs.

• VSEPR Theory: Electron pairs around a central atom arrange themselves to minimize repulsion, determining the molecule's geometry.

• 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: Show the arrangement of atoms and electron pairs, serving as the basis for predicting molecular shapes.

• Example: Methane ($CH_4$) is tetrahedral; water ($H_2O$) is bent.

Electronegativity and Molecular Polarity

Molecular polarity depends on both the polarity of individual bonds and the overall shape of the molecule. Electronegativity differences between atoms create polar covalent bonds.

• Electronegativity: The tendency of an atom to attract electrons in a bond. Fluorine is the most electronegative element.

• Bond Polarity: A bond is polar if the atoms have different electronegativities; the greater the difference, the more polar the bond.

• Molecular Polarity: Determined by both bond polarities and molecular geometry. Symmetrical molecules may be nonpolar even if they contain polar bonds.

• Example: Carbon dioxide ($CO_2$) is nonpolar due to its linear shape, while water ($H_2O$) is polar due to its bent shape.

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: Use lines to represent bonds between carbon atoms; hydrogens attached to carbons are usually omitted.

• Conversion: Practice converting between these representations for clarity and efficiency.

• Example: Ethanol can be written as Lewis, condensed ($CH_3CH_2OH$), or skeletal structure.

Alkanes: The Simplest Organic Compounds

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

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

• Unsaturated Hydrocarbons: Contain double or triple C–C bonds.

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

• Molecular Formulas: Straight-chain alkanes: $C_nH_{2n+2}$; cycloalkanes: $C_nH_{2n}$.

• Drawing Structures: Practice drawing skeletal and condensed structures from molecular formulas.

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 and reactions of those molecules.

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

• Identification: Recognize functional groups in molecular structures.

• 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 the structure of the molecule.

• Branched-Chain Alkanes: Identify the longest carbon chain, number the chain, and name substituents as prefixes.

• Example: 2-methylpropane (isobutane).

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 isomers have substituents on the same side, trans on opposite sides.

• Chiral Centers: Carbon atoms bonded to four different groups; molecules with chiral centers can exist as enantiomers.

• Example: But-2-ene 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 ($riangle G$).

• Equation: $\triangle G = \triangle H - T\triangle S$

• Exergonic vs. Endergonic: Exergonic reactions release energy ($\triangle G < 0$), endergonic reactions absorb energy ($\triangle G > 0$).

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

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

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.

• Reaction Energy Diagram: Shows energy changes during a reaction, including activation energy and overall energy change.

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

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

• Example: The decomposition of hydrogen peroxide is faster with the enzyme catalase.

Overview of Chemical Reactions

Types and Classification of Reactions

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

• Synthesis Reaction: Two or more substances combine to form one product ($A + B \rightarrow AB$).

• Decomposition Reaction: One substance breaks down into two or more products ($AB \rightarrow A + B$).

• Exchange Reaction: Parts of two compounds are exchanged ($AB + CD \rightarrow AD + CB$).

• Combustion Reaction: Hydrocarbon reacts with oxygen to produce carbon dioxide and water ($C_xH_y + O_2 \rightarrow CO_2 + H_2O$).

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

• Balancing Equations: Ensure the same number of each atom on both sides of the equation.

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

• Example: In the reaction $2Na + Cl_2 \rightarrow 2NaCl$, sodium is oxidized, chlorine is reduced.

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 to form an alkane.

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

• General Reaction: $RCH=CH_2 + H_2 \rightarrow RCH_2CH_3$ (hydrogenation)

• General Reaction: $RCH=CH_2 + H_2O \rightarrow RCH(OH)CH_3$ (hydration)

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

Type of Reaction	General Equation	Example
Synthesis	$A + B \rightarrow AB$	$2H_2 + O_2 \rightarrow 2H_2O$
Decomposition	$AB \rightarrow A + B$	$2H_2O \rightarrow 2H_2 + O_2$
Exchange	$AB + CD \rightarrow AD + CB$	$NaCl + AgNO_3 \rightarrow NaNO_3 + AgCl$
Combustion	$C_xH_y + O_2 \rightarrow CO_2 + H_2O$	$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Additional info: Some context and examples have been inferred and expanded for completeness and clarity, as the original notes were brief and fragmented.