General Chemistry Study Notes: Reactions in Aqueous Solution, Periodicity, and Ionic Bonding

Chapter 4: Reactions in Aqueous Solution

Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed. Each side of the equation must have the same number of atoms for each element.

• Steps to Balance:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms (never change subscripts).

  4. Check your work.

• Example: Balance the reaction of hydrogen and oxygen to form water: $2H_2 + O_2 \rightarrow 2H_2O$

Types and Characteristics of Chemical Reactions

Chemical reactions can be classified by their characteristics and the changes they produce.

• Precipitation Reactions: Two aqueous solutions form an insoluble solid (precipitate).

• Acid-Base (Neutralization) Reactions: An acid reacts with a base to produce water and a salt.

• Oxidation-Reduction (Redox) Reactions: Electrons are transferred between species, changing oxidation states.

Solution Stoichiometry

Solution stoichiometry involves calculations with solutions, using molarity to relate volume and moles.

• Molarity (M): The concentration of a solution, defined as moles of solute per liter of solution. $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution: Adding solvent to decrease concentration. The number of moles of solute remains constant. $M_1V_1 = M_2V_2$

Titrations

Titration is a technique to determine the concentration of a solution by reacting it with a standard solution.

• Key Steps:

  1. Add titrant to analyte until the reaction reaches the endpoint (often indicated by a color change).

  2. Use the balanced equation to relate moles of titrant and analyte.

Precipitation Reactions and Predicting Precipitates

Precipitation reactions occur when two soluble salts form an insoluble product.

• Predicting Precipitates: Use solubility rules to determine if a solid forms.

• Example: Mixing $AgNO_3$ and $NaCl$ forms $AgCl$ precipitate: $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$

Net Ionic Equations

Net ionic equations show only the species that change during the reaction.

• Steps:

  1. Write the balanced molecular equation.

  2. Write the complete ionic equation (all strong electrolytes as ions).

  3. Cancel spectator ions to get the net ionic equation.

• Example: $Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)$

Neutralization Reactions: Complete and Net Ionic Equations

Neutralization involves an acid and a base forming water and a salt.

• Example: $HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$

• Net Ionic Equation: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$

Identifying Oxidation States

Oxidation states (numbers) indicate the degree of oxidation of an atom in a compound.

• Rules:

  • Free elements: 0

  • Monatomic ions: charge of the ion

  • Oxygen: usually -2

  • Hydrogen: +1 with nonmetals, -1 with metals

  • Sum in a compound: 0; in a polyatomic ion: equals the ion charge

Identifying Reaction Types

• Precipitation: Formation of an insoluble solid

• Acid-Base: Transfer of H+ ions

• Oxidation-Reduction: Change in oxidation states

Chapter 5: Periodicity and Electronic Structure of Atoms

Periodic Trends and Properties

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Atomic Radius: Size of an atom; decreases across a period, increases down a group.

• Nuclear Effective Charge (Zeff): Net positive charge experienced by valence electrons; increases across a period.

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and electrons.

• Principal Quantum Number (n): Energy level (n = 1, 2, 3, ...)

• Angular Momentum Quantum Number (l): Shape of orbital (l = 0 to n-1)

• Magnetic Quantum Number (ml): Orientation of orbital (-l to +l)

• Spin Quantum Number (ms): Electron spin (+1/2 or -1/2)

Shapes of s and p Orbitals

• s Orbitals: Spherical shape

• p Orbitals: Dumbbell-shaped, oriented along x, y, or z axes

Pauli Exclusion Principle and Hund’s Rule

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

Electron Configurations

Electron configurations show the arrangement of electrons in an atom.

• Standard Notation: List orbitals in order of filling (e.g., 1s2 2s2 2p6).

• Abbreviated Notation: Use noble gas core (e.g., [Ne] 3s2 3p4).

• Valence Electrons: Electrons in the outermost shell, important for chemical reactivity.

Chapter 6: Ionic Compounds and Periodic Trends

Ion Electron Configurations

Ions form when atoms gain or lose electrons to achieve noble gas configurations.

• Cations: Lose electrons; configuration reflects loss from highest energy level.

• Anions: Gain electrons; configuration reflects addition to valence shell.

• Example: $Na: 1s^2 2s^2 2p^6 3s^1 \rightarrow Na^+: 1s^2 2s^2 2p^6$

Periodic Properties of Ions

• Ionic Radius: Cations are smaller, anions are larger than their parent atoms.

• Isoelectronic Series: Ions with the same electron configuration; size decreases with increasing nuclear charge.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• Trend: Increases across a period, decreases down a group.

• Equation: $M(g) \rightarrow M^+(g) + e^-$

Electron Affinity

Electron affinity is the energy change when an electron is added to a neutral atom.

• Trend: Generally becomes more negative across a period (atoms more likely to gain electrons).

• Equation: $X(g) + e^- \rightarrow X^-(g)$

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to noble gases.

• Application: Explains the formation of most ionic and covalent compounds.