Yours

13.2: Solutions- Homogeneous Mixtures

CHEM 114: Introductory Chemistry · LibreTexts

0%

Contents

    Text

    Size
    100%
    Theme
    Typeface
    Line length
    Spacing

    From LibreTexts, under CC BY-NC-SA 4.0, as the source states it; its equations are written in TeX, which the PDF shows typeset. This page at LibreTexts.

    Chapter 1
    1. Chapter 1: The Chemical World
    2. 1.1: The Scope of Chemistry
    3. 1.2: Chemicals Compose Ordinary Things
    4. 1.3: Hypothesis, Theories, and Laws
    5. 1.4: The Scientific Method: How Chemists Think
    6. 1.5: A Beginning Chemist: How to Succeed
     •  Chapter 2
    1. Chapter 2: Measurement and Problem Solving
    2. 2.1: Taking Measurements
    3. 2.2: Scientific Notation: Writing Large and Small Numbers
    4. 2.3: Significant Figures: Writing Numbers to Reflect Precision
    5. 2.4: Significant Figures in Calculations
    6. 2.5: The Basic Units of Measurement
    7. 2.6: Problem Solving and Unit Conversions
    8. 2.7: Solving Multistep Conversion Problems
    9. 2.8: Units Raised to a Power
    10. 2.9: Density
    11. 2.10: Numerical Problem-Solving Strategies and the Solution Map
    12. 2.E: Measurement and Problem Solving (Exercises)
     •  Chapter 3
    1. Chapter 3: Matter and Energy
    2. 3.1: In Your Room
    3. 3.2: What is Matter?
    4. 3.3: Classifying Matter According to Its State: Solid, Liquid, and Gas
    5. 3.4: Classifying Matter According to Its Composition
    6. 3.5: Differences in Matter: Physical and Chemical Properties
    7. 3.6: Changes in Matter: Physical and Chemical Changes
    8. 3.7: Conservation of Mass: There is No New Matter
    9. 3.8: Energy
    10. 3.9: Energy and Chemical and Physical Change
    11. 3.10: Temperature: Random Motion of Molecules and Atoms
    12. 3.11: Temperature Changes: Heat Capacity
    13. 3.12: Energy and Heat Capacity Calculations
    14. 3.E: Exercises
     •  Chapter 4
    1. Chapter 4: Atoms and Elements
    2. 4.1: Experiencing Atoms at Tiburon
    3. 4.2: Indivisible: The Atomic Theory
    4. 4.3: The Nuclear Atom
    5. 4.4: The Properties of Protons, Neutrons, and Electrons
    6. 4.5: Elements: Defined by Their Numbers of Protons
    7. 4.6: Looking for Patterns: The Periodic Law and the Periodic Table
    8. 4.7: Ions: Losing and Gaining Electrons
    9. 4.8: Isotopes: When the Number of Neutrons Varies
    10. 4.9: Atomic Mass: The Average Mass of an Element’s Atoms
      •  Chapter 5
    1. Chapter 5: Molecules and Compounds
    2. 5.1: Sugar and Salt
    3. 5.2: Compounds Display Constant Composition
    4. 5.3: Chemical Formulas: How to Represent Compounds
    5. 5.4: A Molecular View of Elements and Compounds
    6. 5.5: Writing Formulas for Ionic Compounds
    7. 5.6: Nomenclature: Naming Compounds
    8. 5.7: Naming Ionic Compounds
    9. 5.8: Naming Molecular Compounds
    10. 5.9: Naming Acids
    11. 5.10: Nomenclature Summary
    12. 5.11: Formula Mass: The Mass of a Molecule or Formula Unit
      •  Chapter 6
    1. Chapter 6: Chemical Composition
    2. 6.1: How Much Sodium?
    3. 6.2: Counting Nails by the Pound
    4. 6.3: Counting Atoms by the Gram
    5. 6.4: Counting Molecules by the Gram
    6. 6.5: Chemical Formulas as Conversion Factors
    7. 6.6: Mass Percent Composition of Compounds
    8. 6.7: Mass Percent Composition from a Chemical Formula
    9. 6.8: Calculating Empirical Formulas for Compounds
    10. 6.9: Calculating Molecular Formulas for Compounds
      •  Chapter 7
    1. Chapter 7: Chemical Reactions
    2. 7.1: Grade School Volcanoes, Automobiles, and Laundry Detergents
    3. 7.2: Evidence of a Chemical Reaction
    4. 7.3: The Chemical Equation
    5. 7.4: How to Write Balanced Chemical Equations
    6. 7.5: Aqueous Solutions and Solubility: Compounds Dissolved in Water
    7. 7.6: Precipitation Reactions: Reactions in Aqueous Solution That Form a Solid
    8. 7.7: Writing Chemical Equations for Reactions in Solution: Molecular, Complete Ionic, and Net Ionic Equations
    9. 7.8: Acid–Base and Gas Evolution Reactions
    10. 7.9: Oxidation–Reduction Reactions
    11. 7.10: Classifying Chemical Reactions
    12. 7.11: The Activity Series
     •  Chapter 8
    1. Chapter 8: Quantities in Chemical Reactions
    2. 8.1: Climate Change: Too Much Carbon Dioxide
    3. 8.2: Stoichiometry
    4. 8.3: Making Molecules: Mole-to-Mole Conversions
    5. 8.4: Making Molecules: Mass-to-Mass Conversions
    6. 8.5: Limiting Reactant, Theoretical Yield, and Percent Yield
    7. 8.6: Limiting Reactant, Theoretical Yield, and Percent Yield from Initial Masses of Reactants
    8. 8.7: Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction
    Chapter 9
    1. Chapter 9: Electrons in Atoms and the Periodic Table
    2. 9.1: Blimps, Balloons, and Models of the Atom
    3. 9.2: Light: Electromagnetic Radiation
    4. 9.3: The Electromagnetic Spectrum
    5. 9.4: The Bohr Model: Atoms with Orbits
    6. 9.5: The Quantum-Mechanical Model: Atoms with Orbitals
    7. 9.6: Quantum-Mechanical Orbitals and Electron Configurations
    8. 9.7: Electron Configurations and the Periodic Table
    9. 9.8: The Explanatory Power of the Quantum-Mechanical Model
    10. 9.9: Periodic Trends: Atomic Size, Ionization Energy, and Metallic Character
     •  Chapter 10
    1. Chapter 10: Chemical Bonding
    2. 10.1: Bonding Models and AIDS Drugs
    3. 10.2: Representing Valence Electrons with Dots
    4. 10.3: Lewis Structures of Ionic Compounds: Electrons Transferred
    5. 10.4: Covalent Lewis Structures: Electrons Shared
    6. 10.5: Writing Lewis Structures for Covalent Compounds
    7. 10.6: Resonance: Equivalent Lewis Structures for the Same Molecule
    8. 10.7: Predicting the Shapes of Molecules
    9. 10.8: Electronegativity and Polarity: Why Oil and Water Don’t Mix
     •  Chapter 11
    1. Chapter 11: Gases
    2. 11.1: Extra-Long Straws
    3. 11.2: Kinetic Molecular Theory: A Model for Gases
    4. 11.3: Pressure: The Result of Constant Molecular Collisions
    5. 11.4: Boyle’s Law: Pressure and Volume
    6. 11.5: Charles’s Law: Volume and Temperature
    7. 11.6: Gay-Lussac's Law: Temperature and Pressure
    8. 11.7: The Combined Gas Law: Pressure, Volume, and Temperature
    9. 11.8: Avogadro’s Law: Volume and Moles
    10. 11.9: The Ideal Gas Law: Pressure, Volume, Temperature, and Moles
    11. 11.10: Mixtures of Gases: Why Deep-Sea Divers Breathe a Mixture of Helium and Oxygen
    12. 11.11: Gases in Chemical Reactions
     •  Chapter 12
    1. Chapter 12: Liquids, Solids, and Intermolecular Forces
    2. 12.1: Interactions between Molecules
    3. 12.2: Properties of Liquids and Solids
    4. 12.3: Intermolecular Forces in Action: Surface Tension and Viscosity
    5. 12.4: Evaporation and Condensation
    6. 12.5: Melting, Freezing, and Sublimation
    7. 12.6: Types of Intermolecular Forces: Dispersion, Dipole–Dipole, Hydrogen Bonding, and Ion-Dipole
    8. 12.7: Types of Crystalline Solids: Molecular, Ionic, and Atomic
    9. 12.8: Water: A Remarkable Molecule
     •  Chapter 13
    1. Chapter 13: Solutions
    2. 13.1: Prelude - Tragedy in Cameroon
    3. 13.2: Solutions: Homogeneous Mixtures
    4. 13.3: Solutions of Solids Dissolved in Water: How to Make Rock Candy
    5. 13.4: Solutions of Gases in Water: How Soda Pop Gets Its Fizz
    6. 13.5: Solution Concentration: Mass Percent
    7. 13.6: Solution Concentration: Molarity
    8. 13.7: Solution Dilution
    9. 13.8: Solution Stoichiometry
    10. 13.9: Freezing Point Depression and Boiling Point Elevation: Making Water Freeze Colder and Boil Hotter
    11. 13.10: Osmosis: Why Drinking Salt Water Causes Dehydration
      •  Chapter 14
    1. Chapter 14: Acids and Bases
    2. 14.1: Sour Patch Kids and International Spy Movies
    3. 14.2: Acids: Properties and Examples
    4. 14.3: Bases: Properties and Examples
    5. 14.4: Molecular Definitions of Acids and Bases
    6. 14.5: Reactions of Acids and Bases
    7. 14.6: Acid–Base Titration: A Way to Quantify the Amount of Acid or Base in a Solution
    8. 14.7: Strong and Weak Acids and Bases
    9. 14.8: Water: Acid and Base in One
    10. 14.9: The pH and pOH Scales: Ways to Express Acidity and Basicity
    11. 14.10: Buffers: Solutions That Resist pH Change
     

    Template:HideTOC

    Learning Objectives
    • Learn terminology involving solutions.
    • Explain the significance of the statement "like dissolves like."
    • Explain why certain substances dissolve in other substances.

    The major component of a solution is called the solvent. The minor component of a solution is called the solute. By major and minor we mean whichever component has the greater or lesser presence by mass or by moles. Sometimes this becomes confusing, especially with substances with very different molar masses. However, here we will confine the discussion to solutions for which the major component and the minor component are obvious.

    A spoon is pouring sugar into a glass of water on a wooden surface.
    Figure \(\PageIndex{1}\): Making a saline water solution by dissolving table salt (NaCl) in water. The salt is the solute and the water the solvent. (CC-BY-SA 3.0; Chris 73).

    Solutions exist for every possible phase of the solute and the solvent. Salt water, for example, is a solution of solid \(\ce{NaCl}\) in liquid water, while air is a solution of a gaseous solute (\(\ce{O2}\)) in a gaseous solvent (\(\ce{N2}\)). In all cases, however, the overall phase of the solution is the same phase as the solvent. Table \(\PageIndex{1}\) lists some common types of solutions, with examples of each.

    Table \(\PageIndex{1}\): Types of Solutions
    Solvent Phase Solute Phase Example
    gas gas air
    liquid gas carbonated beverages
    liquid liquid ethanol (C2H5OH) in H2O (alcoholic beverages)
    liquid solid salt water
    solid gas H2 gas absorbed by Pd metal
    solid liquid Hg(ℓ) in dental fillings
    solid solid steel alloys
    Example \(\PageIndex{1}\): Sugar and Water

    A solution is made by dissolving 1.00 g of sucrose (\(\ce{C12H22O11}\)) in 100.0 g of liquid water. Identify the solvent and solute in the resulting solution.

    Solution

    Either by mass or by moles, the obvious minor component is sucrose, so it is the solute. Water—the majority component—is the solvent. The fact that the resulting solution is the same phase as water also suggests that water is the solvent.

    Exercise \(\PageIndex{1}\)

    A solution is made by dissolving 3.33 g of \(\ce{HCl(g)}\) in 40.0 g of liquid methyl alcohol (\(\ce{CH3OH}\)). Identify the solvent and solute in the resulting solution.

    Answer

    solute: \(\ce{HCl(g)}\)

    solvent: \(\ce{CH3OH}\)

    Like Dissolves Like

    A simple way to predict which compounds will dissolve in other compounds is the phrase "like dissolves like". What this means is that polar compounds dissolve polar compounds, nonpolar compounds dissolve nonpolar compounds, but polar and nonpolar do not dissolve in each other.

    Even some nonpolar substances dissolve in water but only to a limited degree. Have you ever wondered why fish are able to breathe? Oxygen gas, a nonpolar molecule, does dissolve in water—it is this oxygen that the fish take in through their gills. The reason we can enjoy carbonated sodas is also due to a nonpolar compound that dissolves in water. Pepsi-cola and all the other sodas have carbon dioxide gas, \(\ce{CO_2}\), a nonpolar compound, dissolved in a sugar-water solution. In this case, to keep as much gas in solution as possible, the sodas are kept under pressure.

    This general trend of "like dissolves like" is summarized in the following table:

    Table \(\PageIndex{2}\): Summary of Solubilities
    Solute (Polarity of Compound) Solvent (Polarity of Compound) Dominant Intermolecular Force Is Solution Formed?
    Polar Polar Dipole-Dipole Force and/or Hydrogen Bond yes
    Non-polar Non-polar Dispersion Force yes
    Polar Non-polar   no
    Non-polar Polar   no
    Ionic Polar Ion-Dipole yes
    Ionic Non-polar   no

    Note that every time charged particles (ionic compounds or polar substances) are mixed, a solution is formed. When particles with no charges (nonpolar compounds) are mixed, they will form a solution. However, if substances with charges are mixed with other substances without charges, a solution does not form. When an ionic compound is considered "insoluble", it doesn't necessarily mean the compound is completely untouched by water. All ionic compounds dissolve to some extent. An insoluble compound just doesn't dissolve in any noticeable or appreciable amount.

    What is it that makes a solute soluble in some solvents but not others?

    The answer is intermolecular interactions. The intermolecular interactions include London dispersion forces, dipole-dipole interactions, and hydrogen bonding (as described in Chapter 10). From experimental studies, it has been determined that if molecules of a solute experience the same intermolecular forces that the solvent does, the solute will likely dissolve in that solvent. So, NaCl—a very polar substance because it is composed of ions—dissolves in water, which is very polar, but not in oil, which is generally nonpolar. Nonpolar wax dissolves in nonpolar hexane, but not in polar water.

    A clear glass containing a layered liquid, with yellow liquid on top and colorless liquid below.
    Figure \(\PageIndex{2}\): Water (clear liquid) and oil (yellow) do not form liquid solutions. (CC BY-SA 1.0 Generic; Victor Blacus)
    Example \(\PageIndex{2}\): Polar and Nonpolar Solvents

    Would \(\ce{I2}\) be more soluble in \(\ce{CCl4}\) or \(\ce{H2O}\)? Explain your answer.

    Solution

    \(\ce{I2}\) is nonpolar. Of the two solvents, \(\ce{CCl4}\) is nonpolar and \(\ce{H2O}\) is polar, so \(\ce{I2}\) would be expected to be more soluble in \(\ce{CCl4}\).

    Exercise \(\PageIndex{2}\)

    Would \(\ce{C3H7OH}\) be more soluble in \(\ce{CCl4}\) or \(\ce{H2O}\)? Explain your answer.

    Answer

    \(\ce{H2O}\), because both experience hydrogen bonding.

    Example \(\PageIndex{3}\)

    Water is considered a polar solvent. Which substances should dissolve in water?

    1. methanol (\(\ce{CH3OH}\))
    2. sodium sulfate (\ce{Na2SO4}\))
    3. octane (\(\ce{C8H18}\))
    Solution

    Because water is polar, substances that are polar or ionic will dissolve in it.

    1. Because of the OH group in methanol, we expect its molecules to be polar. Thus, we expect it to be soluble in water. As both water and methanol are liquids, the word miscible can be used in place of soluble.
    2. Sodium sulfate is an ionic compound, so we expect it to be soluble in water.
    3. Like other hydrocarbons, octane is nonpolar, so we expect that it would not be soluble in water.
    Exercise \(\PageIndex{3}\): Toluene

    Toluene (\(\ce{C6H5CH3}\)) is widely used in industry as a nonpolar solvent. Which substances should dissolve in toluene?

    1. water (H2O)
    2. sodium sulfate (Na2SO4)
    3. octane (C8H18)
    Answer

    Octane (\(\ce{C8H18}\)) will dissolve. It is also non-polar.

    Summary

    • Solutions are composed of a solvent (major component) and a solute (minor component).
    • “Like dissolves like” is a useful rule for deciding if a solute will be soluble in a solvent.

    13.2: Solutions- Homogeneous Mixtures is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.