General Chemistry Study Notes: Solutions, Atomic Structure, and Ionic Bonding

Chapter 4: Reactions in Aqueous Solution

Solution Composition and Concentration

Understanding the composition and concentration of solutions is fundamental in chemistry, especially for reactions in aqueous media.

• Solvent: The major component of a mixture; in aqueous solutions, water is the solvent.

• Solute: The minor component dissolved in the solvent.

• Solution: A homogeneous mixture of two or more substances.

• Molarity (M): The concentration of a solution, defined as moles of solute per liter of solution.

  • Formula: $M = \frac{\text{mol of solute}}{\text{L of solution}}$

  • To find moles: $\text{mol} = M \times L$

  • To find volume: $L = \frac{\text{mol}}{M}$

• Mass Percent (%): The mass of solute divided by the mass of solution, multiplied by 100.

  • Formula: $\text{Mass \%} = \frac{\text{g of solute}}{\text{g of solution}} \times 100$

  • Alternatively: $\text{Mass \%} = \frac{\text{g of solute}}{100\,\text{g solution}}$

• Dilution Formula: Used to calculate the concentration or volume after dilution.

  • Formula: $M_1V_1 = M_2V_2$

Electrolytes and Types of Compounds in Solution

Electrolytes are substances that dissociate into ions in aqueous solution, enabling the solution to conduct electricity.

• Electrolytes: Substances that produce ions in solution.

  • Strong electrolytes: Completely dissociate into ions (e.g., strong acids, strong bases, soluble salts).

  • Weak electrolytes: Partially dissociate into ions.

• Acids: Compounds containing hydrogen that ionize to produce H+ in solution (e.g., $\text{HNO}_3$ is a strong acid).

• Bases: Compounds containing hydroxide ions (OH-) (e.g., $\text{Ba(OH)}_2$ is a strong base).

• Salts: Ionic compounds formed from the reaction of an acid and a base; cation comes from the base, anion from the acid.

Types of Chemical Equations

Chemical reactions in solution can be represented in different ways to highlight various aspects of the reaction.

• Molecular Equation: Shows all reactants and products as compounds.

• Total Ionic Equation: Shows all strong electrolytes as dissociated ions; solids, liquids, and gases remain as compounds.

• Net Ionic Equation: Shows only the species that undergo change; spectator ions are omitted.

Example (Acid-Base Neutralization):

• Molecular: $2\text{HNO}_3(aq) + \text{Ba(OH)}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{Ba(NO}_3)_2(aq)$

• Total Ionic: $2\text{H}^+(aq) + 2\text{NO}_3^-(aq) + \text{Ba}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{Ba}^{2+}(aq) + 2\text{NO}_3^-(aq)$

• Net Ionic: $2\text{H}^+(aq) + 2\text{OH}^-(aq) \rightarrow 2\text{H}_2\text{O}(l)$

Calculating Ion Concentrations in Strong Electrolyte Solutions

Strong electrolytes dissociate completely, so the concentration of each ion can be determined from the formula and the initial concentration.

• Example: For $0.15\,M$ $\text{Ba(NO}_3)_2$:

  • Dissociation: $\text{Ba(NO}_3)_2 \rightarrow \text{Ba}^{2+} + 2\text{NO}_3^-$

  • $[\text{Ba}^{2+}] = 0.15\,M$

  • $[\text{NO}_3^-] = 2 \times 0.15\,M = 0.30\,M$

Types of Reactions in Aqueous Solution

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

  • Example: $\text{Na}_2\text{S}(aq) + \text{Fe(NO}_3)_2(aq) \rightarrow 2\text{NaNO}_3(aq) + \text{FeS}(s)$

• Acid-Base Neutralization: Acid reacts with base to produce water and a salt.

• Redox Reaction: Involves transfer of electrons (oxidation and reduction).

Naming Compounds and Writing Formulas

• -ic & -ate: Used for acids with polyatomic ions ending in -ate (e.g., nitric acid from nitrate).

• -ous & -ite: Used for acids with polyatomic ions ending in -ite (e.g., nitrous acid from nitrite).

Stoichiometry with Standard Solutions

Stoichiometric calculations involve using balanced equations and solution concentrations to determine quantities of reactants or products.

Redox Reactions and Reactivity Series

• Oxidizing Agent: Causes oxidation of another species; is itself reduced.

• Reducing Agent: Causes reduction of another species; is itself oxidized.

• Reactivity Series: Ranks metals by their tendency to be oxidized; more active metals can reduce ions of less active metals.

  • Example: $\text{Ca}(s) + \text{Zn}^{2+}(aq) \rightarrow \text{Zn}(s) + \text{Ca}^{2+}(aq)$

Chapter 5: Periodicity & Electronic Structure of Atoms

Electromagnetic Radiation and Atomic Spectra

Light and other forms of electromagnetic radiation exhibit both wave-like and particle-like properties.

• Speed of Light: $c = 3.00 \times 10^8\,\text{m/s}$

• Relationship: $\lambda \nu = c$

  • $\lambda$ = wavelength (m)

  • $\nu$ = frequency (Hz)

• Photon Energy: $E_{\text{photon}} = h\nu = \frac{hc}{\lambda}$

  • $h$ = Planck's constant ($6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$)

Einstein’s Theory and the Photoelectric Effect

Einstein explained the photoelectric effect by proposing that light consists of photons, each with energy $E = h\nu$.

• Electrons are ejected from a metal surface only if the photon energy exceeds a threshold value.

de Broglie Wave Theory

All moving particles have an associated wavelength, though it is only significant for very small particles like electrons.

• de Broglie wavelength: $\lambda = \frac{h}{mv}$

  • $m$ = mass (kg), $v$ = velocity (m/s)

Quantum Numbers and Atomic Orbitals

Quantum numbers describe the properties of atomic orbitals and the electrons within them.

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

• Angular Momentum Quantum Number (l): Subshell shape (l = 0 to n-1)

• Magnetic Quantum Number (ml): Orientation in space (ml = -l to +l)

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

Electron Configuration Principles

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

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

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

Writing Electron Configurations

• Use quantum numbers, subshell designations, noble gas core notation, and orbital-filling diagrams.

• For d-block elements, electrons in d subshells are not counted as valence electrons.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

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

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electron Affinity: Energy change when an atom gains an electron; trends vary across the periodic table.

Chapter 6: Ionic Compounds—Periodic Trends and Bonding Theory

Ionic Bonding and Electron Configuration

Ionic bonding involves the transfer of electrons from metals to nonmetals, resulting in the formation of ions with noble gas configurations.

• Write ground state electron configurations for main group and transition metal ions.

• Determine the number of unpaired electrons in ions.

  • Example: Mn [Ar] 4s2 3d5; Mn2+ [Ar] 3d5; Mn4+ [Ar] 3d3

Periodic Trends in Ionic Compounds

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

• Ionization Energy: Energy required to remove an electron from an atom or ion.

• Electron Affinity: Energy change when an atom gains an electron.

The Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a noble gas electron configuration (eight valence electrons).

Born-Haber Cycle and Lattice Energy

The Born-Haber cycle is a thermochemical cycle used to calculate the lattice energy of an ionic compound.

• Lattice Energy: The energy required to separate one mole of an ionic solid into its gaseous ions.

• Factors affecting lattice energy:

  • Charge of ions (higher charge = higher lattice energy)

  • Size of ions (smaller ions = higher lattice energy)

Born-Haber Cycle Steps (for NaCl as example):

1. Sublimation of Na(s) to Na(g)

2. Ionization of Na(g) to Na+(g)

3. Dissociation of Cl2(g) to Cl(g)

4. Electron affinity: Cl(g) to Cl-(g)

5. Formation of NaCl(s) from Na+(g) and Cl-(g)

Sum of all steps equals the enthalpy of formation of the ionic compound.

Type of Reaction	Definition	Example
Precipitation	Formation of an insoluble solid from two aqueous solutions	Na2S(aq) + Fe(NO3)2(aq) → 2NaNO3(aq) + FeS(s)
Acid-Base Neutralization	Acid reacts with base to form water and a salt	2HNO3(aq) + Ba(OH)2(aq) → 2H2O(l) + Ba(NO3)2(aq)
Redox	Transfer of electrons between species	Ca(s) + Zn2+(aq) → Zn(s) + Ca2+(aq)

Additional info: For more details on solubility rules, strong acids/bases, and the full reactivity series, consult your textbook's appendices or reference tables.