Understanding Solutions & Molarity

Ever wondered why salt disappears in water but oil doesn't? It's all about solutions! A solution has two key parts: the solute (what gets dissolved) and the solvent (what does the dissolving). In aqueous solutions, water is always the solvent.

Solutions can exist in different states of matter. Usually, the solution takes on the state of the solvent (like salt water being liquid). However, solutions can be solid too—like hydrogen in platinum metal!

When substances dissolve in water, some conduct electricity and others don't. Electrolytes can conduct electricity and come in two types: strong electrolytes (like salts and strong acids) completely ionize and make a bright light in a conductivity test, while weak electrolytes only partially ionize and produce a dim light. Nonelectrolytes like sugar dissolve but don't conduct electricity at all.

Quick Tip: Remember "like dissolves like" - polar substances (like water and sugar) dissolve other polar substances, while nonpolar substances dissolve other nonpolar substances.

Molarity measures concentration as moles of solute per liter of solution. When given a molarity value, you can use it as a conversion factor since it connects two units.

Solution Properties & Calculations

Need to make a less concentrated solution from a stock solution? That's dilution! When you add water to a concentrated solution, you're decreasing its molarity without changing the total moles of solute present. The magic formula is M₁V₁ = M₂V₂, where M is molarity and V is volume in liters.

Solution solubility depends on several factors:

• Structure and polarity - "like dissolves like" means polar substances dissolve in polar solvents
• Pressure (for gases) - higher pressure above a solution means more gas dissolves (Henry's Law)
• Temperature - generally, solids become more soluble at higher temperatures, while gases become less soluble (which is why cold fizzy drinks stay fizzy longer!)

Solutions can be unsaturated (could dissolve more solute), saturated (at maximum solute capacity), or supersaturated (holding more solute than normally possible under special conditions).

Remember This: When working with solution stoichiometry problems, always start by writing a balanced chemical equation!

Colligative properties like boiling point elevation and freezing point depression depend on the number of particles in solution. The van't Hoff factor (i) tells you how many ions a compound breaks into when it dissolves. Sugar doesn't dissociate so its i=1, while table salt (NaCl) splits into two ions for an i=2.

Acids & Bases Fundamentals

Acids and bases are everywhere - from the citrus in your breakfast to the soap you wash with! Acids have a pH below 7, taste sour, turn litmus paper red, conduct electricity, and can react with metals to produce hydrogen gas. Bases have a pH above 7, taste bitter, turn litmus blue, feel slippery, conduct electricity, and are often found in household cleaners.

Scientists have defined acids and bases in three important ways:

• Arrhenius definition: Acids increase H⁺ concentration in water; bases increase OH⁻ concentration
• Brønsted-Lowry definition: Acids donate protons (H⁺); bases accept protons
• Lewis definition: Acids accept electron pairs; bases donate electron pairs

The Brønsted-Lowry definition is particularly useful because it's more inclusive than Arrhenius (not all bases contain hydroxide). A substance that can act as either an acid or base depending on the reaction is called amphiprotic (like water or bicarbonate ion HCO₃⁻).

Got It? Try this: HCO₃⁻ + H⁺ → H₂CO₃ (acting as a base) and HCO₃⁻ + OH⁻ → H₂O + CO₃²⁻ (acting as an acid)

When acids and bases react, they form conjugate acid-base pairs. A conjugate acid forms when a base accepts H⁺, while a conjugate base forms when an acid donates H⁺. The driving force behind these reactions is always the movement of hydrogen ions!

Strong vs. Weak Acids & Bases

Not all acids and bases are created equal! The difference between strong and weak is all about how completely they ionize in water.

Strong acids completely ionize in water and are strong electrolytes. There are just 6 you need to remember: HCl, HBr, HI, HNO₃, HClO₄, and H₂SO₄ (though only the first hydrogen in sulfuric acid is considered strong). We write their dissociation with a one-way arrow: HCl → H⁺ + Cl⁻.

Weak acids only partially ionize and are weak electrolytes. We write their dissociation with a two-way arrow showing equilibrium: HC₂H₃O₂ ⇌ H⁺ + C₂H₃O₂⁻. If an acid isn't one of the six strong ones, assume it's weak!

Similarly, strong bases completely ionize $like NaOH → Na⁺ + OH⁻$. Most soluble hydroxides $-OH$ from Group 1 and heavy Group 2 elements are strong bases. Weak bases only partially ionize, like ammonia: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. If it's a base but not a hydroxide, it's probably weak.

Pro Tip: The strength of acids and bases is measured by Ka and Kb values. The higher the Ka, the stronger the acid; the higher the Kb, the stronger the base.

When analyzing diagrams of acids with equal molarity, look at how much has dissociated. More dissociation (more products) means a stronger acid with a higher Ka value.

Understanding pH & Calculations

Water is amazing - it can act as both an acid and a base (amphiprotic), and it can even react with itself in a process called auto-ionization: H₂O + H₂O → H₃O⁺ + OH⁻.

The pH scale measures how acidic or basic a solution is based on the concentration of hydrogen ions. Mathematically: pH = -log([H⁺]) where $H⁺$ is the molar concentration. In pure water at 25°C, the concentration of H⁺ is 1.0 × 10⁻⁷ M, giving a neutral pH of 7.0.

The relationship between $H⁺$ and $OH⁻$ is defined by the water ionization constant: Kw = $H⁺$$OH⁻$ = 1.0 × 10⁻¹⁴. This gives us several useful conversion formulas:

• pH + pOH = 14.00
• $H⁺$ × $OH⁻$ = 1.0 × 10⁻¹⁴
• From pH to $H⁺$: $H⁺$ = 10⁻ᵖᴴ
• From $H⁺$ to pH: pH = -log([H⁺])

Wait, what? Yes, pH can actually be less than 0 or greater than 14 for highly concentrated solutions, though you won't encounter this often in everyday life.

For strong acids, calculating pH is straightforward because they completely ionize. A 0.01 M solution of HCl will have $H⁺$ = 0.01 M. For strong bases, any hydroxide in solution contributes directly to $OH⁻$.

Remember that a strong acid and a weak acid can have the same pH if their $H⁺$ concentrations are equal - pH only measures hydrogen ion concentration, not the acid's strength.

Neutralizations & Titrations

When acids and bases meet, they undergo neutralization reactions - a type of double replacement reaction that typically produces water and a salt. For example: HCl + NaOH → NaCl + H₂O.

When writing net ionic equations for neutralization reactions, remember to only split strong electrolytes into ions. Weak acids and bases should remain intact in your equation. For example, when barium hydroxide reacts with acetic acid (HC₂H₃O₂), the net ionic equation is: HC₂H₃O₂ + OH⁻ → C₂H₃O₂⁻ + H₂O.

Titration is a laboratory technique where you add a solution of known concentration to a solution of unknown concentration until the reaction is complete. By measuring how much of the known solution was needed, you can calculate the unknown concentration using stoichiometry.

Titrations work for reactions between:

• Strong acid and strong base
• Strong acid and weak base
• Strong base and weak acid

The equivalence point occurs when both acid and base are completely reacted according to the mole ratio in the chemical equation. We detect this point using pH meters or chemical indicators that change color.

Laboratory Connection: Titrations are commonly used to determine the concentration of an unknown acid or base, or to find the molar mass of an unknown substance.