Final Exam Study Guide - Summary General Chemistry I PDF

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Final exam study guide for Mary Tam's class....

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Final Exam Study Guide This is a BRIEF summary of topics discussed in class. It is meant to help you identify the BASIC topics in first semester general chemistry. Each item is only considered as a jumping off point—you are responsible to review and study the in-depth details of each. Chapter 1  Definition of chemistry  Properties/changes and states of matter  The scientific method  Units, Measurement, SI prefixes and conversion factors  Scientific Notation, Significant Figures, including mathematical rules  Accuracy vs. precision Chapter 2  Matter and substances  Laws: mass conservation, definite composition, multiple proportions  Dalton’s Atomic Theory  Experiments and Results of: Thompson, Rutherford, Millikan  The atom and its particles  The periodic table  Names and Formulas of chemical compounds, acids/bases Chapter 3  Mole, molar mass, Avagadro’s Number and conversions  Mass percent calculations  Determining empirical and molecular formulas  Writing balanced chemical equations  Stoichiometry calculations—the mole ratio  Limiting reaction problems, calculating reactant in excess  Theoretical, actual and percent yields Chapter 4  Water as a Solvent—it’s polar nature, how ionic/covalent compounds dissolve in water  Electrolyte vs. non-electrolytes  Precipitation reactions and solubility rules  Writing net ionic equations and all that goes with it  Acid/base reactions/equations  Oxidation/Reduction reactions, oxidation numbers  Types of redox reactions  Solution chemistry—molarity and calculations

Chapter 5  Gas pressure and units/conversion

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Kinetic Molecular Theory of Gases All the gas laws—Boyle’s, Charles’s, Avagadro’s, Amonton’s and the IDEAL gas law— know trends Gas Behavior at standard conditions Application of the IDEAL gas law—calculating density, molar mass Dalton’s Law and partial pressures Graham’s Law, effusion and diffusion Real Gases—deviation from the ideal gas laws

Chapter 6 Forms of energy—heat and work, transfer to system/surroundings  Law of conservation of energy—1st Law of Thermodynamics  Units of energy and conversions  Endothermic vs. exothermic  Enthalpy and H  Stoichiometry of thermochemical reactions  Hess’s Law  Standard Heat of Reactions, Horxn and Heats of Formation, Hof  Calorimetry (both coffee-cup and bomb) along with calculations  Heat capacity and specific heat, calculate q  Chapter 7  Light, the electromagnetic spectrum, wave and particle properties  Planck’s idea  Calculations of wavelength, frequency, energy, transitions  Bohr Model of the Atom  Spectral analysis and energy states of the hydrogen atom  Wave-particle duality of electrons  Heisenberg Uncertainty Principle  Schrodinger’s Equation and Theory  Quantum Mechanics and quantum numbers  Atomic orbitals, shapes, and electron probability Chapter 8  Pauli Exclusion Principle and magnetic spin  Diamagnetic vs. paramagnetic  Aufbau Principle and Hund’s Rule  Break up periodic table in s, p, d and f block orbitals  Orbital Diagrams and Electron configurations  Unusual configurations, those of cations and anions  Periodic trends

Chapter 9  Ionic vs. covalent bonding  Drawing lewis structures of individual atoms  Lattice energy

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Bonding and chemical change Determine bond strength and strength Electronegativity and bond polarity Polar covalent bonds, partial bond charges

Chapter 10  Lewis structures of molecules  Resonance, octet rule and exceptions  Formal charges  VESPER Theory and molecular shape  Regions of electron density  Determine electron group geometry, molecular geometry and bond angles  Polarity and dipoles Chapter 11  Valence bond theory  Orbital hybridization  Orbital overlap in single, double, and triple bonds  Types of bonding  Isomerism Chapter 12  States of matter  Chemical vs. physical properties  Types of forces and strength  Dispersion, dipole-dipole, ion-dipole, hydrogen bonding, ionic bonding  Relationship between intermolecular forces and boiling point  Phase changes and Phase diagrams Chapter 13  IMFs in pure substances  IMFs in solutions  IMFs and relationship to boiling point  The solution process  The many units of concentration  Colligative properties of solutions

Chapter 16  Different types of rates and how to calculate  Orders of reactions and calculations  Integrated rate laws  Half lives

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Effect of temperature and concentration on rate of reaction Reaction mechanisms Catalysis

Chapter 17 The concept of equilibrium  Equilibrium constant and the reaction quotient  Kc and Kp  The direction of the reaction, Q vs. K  Solving equilibrium problems, ICE charts  Le Chatelier’s Principle  Chapter 18 Acids and Bases, conjugate pairs  Water and the pH scale  Calculating pH, pOH, [H+], [OH-] of strong acids or bases  Calculating pH, pOH, [H+], [OH-], Ka, Kb of weak acids/bases through ICE charts  Relationship between pH, pOH, [H+], [OH-], Ka, Kb  Polyprotic acids  Acid/base properties of salt solutions  Trends in acid strength  Chapter 19  Acid/Base Buffer systems  How a buffer works  Henderson-Hasselbalch equation  Buffer capacity and buffer range  Acid/Base titration curves  Indicators  Solubility of ionic compounds, Ksp calculations  Effect of the common ion, pH, and ligands on the solubility of an ionic compound Chapter 20  Entropy and its calculations  Enthalpy and its calculations  Gibbs free energy and its calculations  Equilibrium, direction of reaction determined through spontaneity

Chapter 21  Redox reactions and the agents  Oxidation numbers  Voltaic and electrolytic cells  Cell potential  Free energy and electrical work  Batteries

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Using electrical energy to drive a non spontaneous reaction—current, charge, mass

Chapter 23  Radioactive decay and types of particles  Nuclear stability  Kinetics of radioactive decay  Units of radiation...