General Chemistry Review

Significant Figures and Unit Conversions

Understanding significant figures and unit conversions is essential for accurate chemical calculations.

• Significant Figures: The number of meaningful digits in a measured or calculated quantity. Rules include counting all nonzero digits, zeros between nonzero digits, and trailing zeros in decimals as significant.

• Unit Conversions: Use conversion factors to change units (e.g., meters to centimeters: ).

• Example: Convert 8.11 meters to centimeters: .

Density and Solution Calculations

Density relates mass and volume, and is used to solve for unknowns in solution problems.

• Density Formula:

• Example: If ethanol has a density of , the mass of is .

Temperature Conversions

Temperature can be converted between Celsius, Kelvin, and Fahrenheit using standard equations.

• Celsius to Kelvin:

• Celsius to Fahrenheit:

Atoms, Molecules, and Compounds

Atomic Structure and Isotopes

Atoms consist of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons.

• Isotope Symbol: , where is mass number, is atomic number.

• Example: Potassium with 22 neutrons: ; symbol: .

Ions and Ionic Compounds

Ions are charged species formed by gaining or losing electrons. Ionic compounds are formed from cations and anions.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Example: and combine to form .

Empirical and Molecular Formulas

Formulas represent the composition of compounds.

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms in a molecule.

Stoichiometry and Chemical Reactions

Mole Concept and Molar Mass

The mole is a counting unit for atoms and molecules. Molar mass is the mass of one mole of a substance.

• Avogadro's Number: particles/mol

• Molar Mass: Sum of atomic masses in grams per mole.

• Example: Molar mass of :

Balancing Chemical Equations

Balanced equations have equal numbers of each atom on both sides.

• Example:

Limiting Reactant and Yield

The limiting reactant is consumed first and determines the amount of product formed.

• Theoretical Yield: Maximum amount of product possible.

• Percent Yield:

Solutions and Concentrations

Types of Solutions and Electrolytes

Solutions are homogeneous mixtures. Electrolytes conduct electricity when dissolved in water.

• Strong Electrolyte: Completely dissociates in solution (e.g., NaCl).

• Weak Electrolyte: Partially dissociates (e.g., acetic acid).

• Nonelectrolyte: Does not dissociate (e.g., sugar).

Concentration Units

Concentration expresses the amount of solute in a given amount of solvent.

• Molarity (M):

• Example: of in solution:

Acids, Bases, and pH

Acid-Base Definitions and Reactions

Acids donate protons (H+), bases accept protons.

• Bronsted-Lowry Acid: Proton donor.

• Bronsted-Lowry Base: Proton acceptor.

• Conjugate Acid-Base Pairs: Differ by one proton.

pH Calculations

pH measures the acidity of a solution.

• pH Formula:

• Example: ,

Organic Chemistry Fundamentals

Hydrocarbons: Alkanes, Alkenes, and Alkynes

Hydrocarbons are compounds containing only carbon and hydrogen.

• Alkanes: Saturated hydrocarbons with single bonds ().

• Alkenes: Unsaturated hydrocarbons with at least one double bond ().

• Alkynes: Unsaturated hydrocarbons with at least one triple bond ().

• Example: Cyclohexane (alkane), 1-hexene (alkene).

Isomerism and Chirality

Isomers have the same molecular formula but different structures.

• Structural Isomers: Different connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

• Chirality: A molecule is chiral if it is not superimposable on its mirror image. Chiral centers are typically carbon atoms with four different groups attached.

Cis-Trans (Geometric) Isomerism

Cis-trans isomerism occurs in alkenes due to restricted rotation around the double bond.

• Cis Isomer: Similar groups on the same side of the double bond.

• Trans Isomer: Similar groups on opposite sides.

Carbohydrates and Amino Acids

Monosaccharides and Polysaccharides

Carbohydrates are classified by the number of sugar units.

• Monosaccharides: Simple sugars (e.g., glucose).

• Polysaccharides: Long chains of monosaccharides (e.g., starch, cellulose).

• Fischer and Haworth Projections: Ways to represent carbohydrate structures.

Amino Acids and Proteins

Amino acids are the building blocks of proteins. They contain an amino group, carboxyl group, hydrogen, and a unique side chain attached to a central carbon.

• Classification: Amino acids can be acidic, basic, neutral, primary, secondary, or quaternary based on their side chains and structure.

• Example: L-alanine and β-alanine are structural isomers.

Practice Problems and Applications

Sample Calculations and Problem Types

• Unit conversions (e.g., meters to centimeters)

• Significant figures in calculations

• Density, mass, and volume relationships

• Temperature conversions (Celsius, Kelvin, Fahrenheit)

• Stoichiometry: moles, molar mass, limiting reactant

• Balancing chemical equations

• Solution concentrations and dilutions

• Acid-base reactions and pH calculations

• Identifying isomers, chiral centers, and types of hydrocarbons

• Carbohydrate and amino acid structure identification

Sample Table: Types of Electrolytes

Additional info:

• Some questions reference Fischer and Haworth projections, which are important for understanding carbohydrate stereochemistry.

• Questions on chirality and isomerism are foundational for organic chemistry, especially in the context of biological molecules.

• Practice problems cover both general and introductory organic chemistry concepts, suitable for a comprehensive final exam review.