GENCHEM 221 Lecture Notes Summary

Kinetic Energy ( $E_k$ )
- Energy of motion.
Potential Energy ( $E_p$ )
- Energy of position or energy at rest.
- Conservation of energy: $\Delta PE = \Delta KE$
- Energy is conserved in systems; as potential energy increases (e.g., lifting a weight), kinetic energy is converted when the weight falls.

Density
- Definition: mass per unit volume.
- Example: 91 cm³ = 91 mL or mass (g) over volume (mL).
- Intensive Property
- A property that does not change with the quantity of the substance. Example: Specific heat remains unchanged regardless of the amount of substance.
- Heat
- Amount of heat that flows to and from an object.
- Temperature
- Measures the average kinetic energy of the particles in a substance. The scales used are:
  - Fahrenheit (°F)
  - Celsius (°C)
  - Kelvin (K)
- Conversions
- Relationships:
  - $F = \frac{9}{5}C + 32$
  - $C = K - 273.15$

Significant Figures
- Rules:
- Multiplication/Division: Keep the same number of significant figures as the measurement with the least.
- Addition/Subtraction: Keep the same number of decimal places as the measurement with the least.
- Rounding Rules:
- If the digit removed is 5 followed by zeros, the preceding number increases by 1 if it is odd and remains unchanged if even.
- Example: 17.75 → 17.8 (rounds up), 17.65 → 17.6 (remains the same).
Precision vs. Accuracy
- Precision: How close measurements are to each other.
- Accuracy: How close measurements are to the actual value.
- Errors:
- Systematic Error: Consistently higher or lower values due to a faulty measuring device.
- Random Error: Unpredictable variations causing scatter around an average value.

Pure Substances:
- Consist of only one type of atom or molecule.
  - Elements: Cannot be broken down (e.g., H, Na, Ar, Fe).
  - Compounds: Chemical combinations of elements (e.g., NaCl).
Mixtures:
- Contain two or more substances not chemically bonded.
  - Homogeneous Mixtures: Uniform composition (e.g., solutions). Also known as solutions, and when water is the solvent, they are called aqueous solutions.
  - Heterogeneous Mixtures: Non-uniform composition (e.g., fruit salad).
Separation Techniques:
- Filtration: Based on particle size differences.
- Crystallization: Based on solubility differences.
- Distillation: Based on volatility differences.
- Chromatography: Based on solubility differences.

Law of Mass Conservation:
- Total mass remains constant during a chemical reaction.
Law of Definite Composition:
- A compound consists of the same elements in fixed ratios by mass regardless of sample size.
Law of Multiple Proportions:
- If two elements can form more than one compound, the mass ratios of the second element that combine with a fixed mass of the first can be expressed as ratios of small whole numbers.

Structure of Atoms:
- Protons ( $p^+$ ): Charge +1, mass approximately 1 amu, found in the nucleus.
- Neutrons ( $n^0$ ): No charge, mass approximately 1 amu, found in the nucleus.
- Electrons ( $e^{-}$ ): Charge -1, negligible mass, found in electron clouds around the nucleus.
Isotopes:
- Atoms of the same element with the same number of protons but different numbers of neutrons.

Ionic Bonds:
- Formed by electron transfer from metals to nonmetals.
- Metal atoms lose electrons, become cations, while nonmetals gain electrons, becoming anions.
- Strength of ionic bonding is influenced by charge and size of the ions.
Covalent Bonds:
- Formed through the sharing of electrons between nonmetal atoms.
- Governed by the concept of electronegativity; the ability of an atom to attract shared electrons.
- Bond strength and length are inversely related; shorter bonds are generally stronger.

Molecular Compounds: Basic units consist of molecules formed by covalently bonded atoms.
Binary Compounds: Consist of two elements.
- Ionic Compounds: Contains a metal and a nonmetal. The cation is named first.
- Covalent Compounds: Consist of two nonmetals; use prefixes for the number of atoms (e.g., CO₂ is carbon dioxide).

Acids: Produce $H^+$ ions in solution. The strength of acids varies based on their dissociation in water.
Bases: Produce $OH^-$ ions in solution. Strong bases dissociate completely; weak bases dissociate partially.

Precipitation Reactions: Occur when two soluble ionic compounds react to produce an insoluble product (precipitate).
Types of Chemical Equations:
- Molecular Equation: Shows all reactants and products in their molecular form.
- Ionic Equation: Shows all species as ions except insoluble solids and pure liquids.
- Net Ionic Equation: Shows only the ions that participate in the reaction, omitting spectator ions.

Titration: A technique to determine the concentration of an unknown solution by reacting it with a standard solution. The endpoint is indicated by a color change or pH change.
Rate of Reaction: Factors affecting it include concentration, temperature, and surface area.
Order of Reaction: Can be determined by observing changes in concentration and change in the rate of reaction.

Enthalpy Change ( $\Delta H$ ): Heat content change during a reaction.
- $\Delta H = H<em>{products} - H</em>{reactants}$
Calorimetry: Measuring heat changes associated with chemical reactions.
- *C: Specific heat capacity, $C = \frac{q}{m \Delta T}$
Laws of Thermodynamics:
- The first law states that energy cannot be created or destroyed, it only changes form.

Redox Reactions: Involves the transfer of electrons between substances. Oxidation is loss, reduction is gain (OIL RIG).
Oxidation States: Used to keep track of electrons in chemical reactions.

Ideal Gas Law: Combines the relationships between pressure (P), volume (V), temperature (T), and the number of moles (n) into the equation: $PV = nRT$
Kinetic Molecular Theory: Describes the behavior of ideal gases which are in constant random motion.

Dynamic Equilibrium: When the rate of the forward reaction equals the rate of the reverse reaction.
Equilibrium Position: Concentration of reactants/products remains constant over time.