Solutions, Mixtures, Acids, Bases, and pH

Solutions

• Definition: Homogeneous mixtures where components are dissolved and permanently and equally distributed.
• Two Parts of a Solution:
  • Solute: The chemical(s) that dissolve into the solution.
    • Example: Sugar in sugar water.
  • Solvent: The chemical that the solute(s) dissolve into.
    • Example: Water in sugar water (water is often called the "universal solvent").
  • Combined: Sugar (solute) + Water (solvent) $\rightarrow$ Sugar Solution.
  • Example: Vinegar in water (vinegar is a solution itself, implying acetic acid is a solute in water).

Hydrophobic Substances

• Definition: "Water-fearing" chemicals that do not dissolve in water.
• Characteristics: Unable to form a solution when mixed with water.
• Molecular Structure: Include molecules that only contain non-polar covalent bonds.
• Examples: Any type of oil, cinnamon, gasoline.
• Solubility: Hydrophobic molecules/substances can dissolve into each other.
  • Example: Vegetable oil mixed with olive oil.

Hydrophilic Substances

• Definition: "Water-loving" chemicals that can dissolve in water and form a solution.
• Molecular Structure: Include molecules with one or more polar covalent bonds and/or ionic bonds.
• Solubility: Hydrophilic molecules/substances can dissolve into each other.

Homogeneous Mixtures

• Definition: Mixtures where the components/chemicals are equally distributed.
  • Example: Salt water (salt/$H_2O$ mixture).

Types of Homogeneous Mixtures

A) Suspensions

• Definition: When the components of a mixture are temporarily distributed equally.
• Characteristic: Over time, suspensions eventually become a heterogeneous mixture (components separate).
• Examples:
  • Salad Dressing: A mix of oil, vinegar, and chunks initially appears homogeneous when shaken but separates into layers (e.g., oil and vinegar) with chunks settling over time.
  • Blood: When first drawn, blood appears homogeneous (components evenly distributed) due to constant mixing in the body.
    • However, after a minute of sitting, it separates into layers: plasma, white blood cells, and red blood cells.

C) Colloids (or "Gels")

• Definition: A cross between suspensions and solutions; homogeneous mixtures where a compound is partially suspended and partially dissolved.
• Example: Jello.
  • Jello contains gelatin, which is a protein (also known as collagen).
  • Collagen is the most abundant protein in the body, found in bones, nails, cartilage, skin, etc.
  • Gelatin Molecule Characteristics:
    • Part of the molecule is hydrophilic at all temperatures.
    • Another part is hydrophilic at high temperatures but becomes hydrophobic at low temperatures.

Acids

• Definition: Chemicals that produce hydrogen cations ($H^+$) when mixed into a solution.
• $H^+$ in Solution: The presence of $H^+$ ions makes a solution acidic.
• Example 1: Hydrochloric Acid (HCl)
  • When HCl is mixed into water ($H_2O$), its ionic bond breaks.
  • It forms hydrogen ions ($H^+$) and chloride ions ($Cl^-$).
  • Equation: $HCl \xrightarrow{H_2O} H^+ + Cl^-$
• Example 2: Sulfuric Acid ($H<em>2SO</em>4$)
  • When mixed into water, $H<em>2SO</em>4$ breaks apart and produces $H^+$ions.

Buffers

• Definition: Chemicals that react with hydrogen ($H^+$) and hydroxide ($OH^-$) ions, thereby removing them from a solution.
• Function: When a buffer is added to a solution, it makes the solution neutral, or less acidic or basic.
• Example: Bicarbonate ($HCO_3^-$)
  • $HCO_3^-$ is the main ingredient in antacids, also known as "baking soda."
  • Scenario 1: Bicarbonate added to an acidic solution
    • The bicarbonate reacts with excess $H^+$ ions.
    • Reaction: $HCO<em>3^- + H^+ \rightarrow H</em>2O + CO_2$
    • This reaction removes $H^+$ from the solution, making it less acidic (e.g., pH changes from 3 to 6 or 7).
  • Scenario 2: Bicarbonate added to a basic solution
    • The bicarbonate reacts with excess $OH^-$ ions.
    • Reaction: $HCO<em>3^- + OH^- \rightarrow H</em>2O + CO_3^{2-}$
    • This reaction removes $OH^-$ from the solution, making it less basic.
  • Analogy: Buffers act like a "chemical sponge" that gets rid of $H^+$ and $OH^-$ ions, maintaining pH stability.

Bases

• Definition: Chemicals that produce hydroxide anions ($OH^-$) when mixed into water ($H_2O$).
• $OH^-$ in Solution: The presence of $OH^-$ ions makes a solution "basic" (also known as "alkaline").
• Example 1: Sodium Hydroxide (NaOH)
  • Also known as "lye."
  • When NaOH is mixed into water ($H_2O$), it separates into $Na^+$ and $OH^-$ ions floating around.
• Example 2: Ammonia ($NH_3$)
  • When $NH<em>3$ is added to water, it forms $NH</em>4^+$ and $OH^-$ ions.
  • Reaction: $NH<em>3 + H</em>2O \rightarrow NH_4^+ + OH^-$
  • This makes the solution basic.

pH Scale

• Definition: A measure of how acidic or basic a solution is.
• Range: The pH scale ranges from 0 to 14.
• pH Values and Characteristics:
  • Acidic Solutions: Have a pH between 0 and 6.9.
    • Contain hydrogen ions ($H^+$).
    • Increasing $H^+$ concentration corresponds to decreasing pH (more acidic).
  • Basic (Alkaline) Solutions: Have a pH between 7.1 and 14.
    • Contain hydroxide ions ($OH^-$).
    • Increasing $OH^-$ concentration corresponds to increasing pH (more basic).
  • Neutral Solutions: Have a pH of 7.0.
    • Have equal amounts of $H^+$ and $OH^-$ ions, or none at all.
• Visual Representation:
  • $\text{Acidic (0)} \longleftrightarrow \text{Neutral (7.0)} \longleftrightarrow \text{Basic (14)}$
  • Increasing $H^+$ concentration goes towards 0.
  • Increasing $OH^-$ concentration goes towards 14.