21.2 Chemical Properties, Physical/Chemical Changes & Mixtures

Overview & Exam Relevance

Competency 008 (Physical and Chemical Properties) of the TExES Core Subjects EC-6 Science subject exam also covers chemical properties of matter, differentiating between physical and chemical changes, applying the Law of Conservation of Mass, and classifying and separating mixtures and solutions. Elementary science curriculum standards across Texas require educators to guide students from simple sensory observations of materials to rigorous experimental investigations where students identify evidence of chemical change and design physical separation systems.

In the elementary TEKS, students observe how heating and cooling change materials, describe and separate simple mixtures, and recognize that some mixtures keep the physical properties of their ingredients (such as iron filings and sand). Grade 5 students also explore solutions and the conservation of matter. Identifying evidence of chemical change is formally emphasized in middle school, but the 391 expects you to understand and teach the distinction.

On the TExES 391 exam, test questions routinely target common instructional traps: mistaking dissolving for a chemical change, assuming matter is "lost" when gas escapes during combustion or effervescence, confusing phase changes with chemical synthesis, and failing to select the correct sequential physical separation technique for complex multi-part mixtures.

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Chemical Properties of Matter

A chemical property is an intrinsic characteristic of a substance that describes its potential to undergo a chemical transformation that alters its molecular composition and creates new substances. Unlike physical properties, chemical properties cannot be observed or measured without changing the chemical identity of the substance.

KEY CHEMICAL PROPERTIES
│
├── Reactivity ────────► Tendency to chemically combine with other substances (e.g., metals with acids)
├── Flammability ──────► Capability of catching fire and burning in presence of oxygen
├── Acidity / Basicity ─► Concentration of hydrogen ions; measured on the pH scale (0 to 14)
└── Oxidation ─────────► Reaction with oxygen/loss of electrons (e.g., iron rusting, apple browning)

1. Reactivity

Reactivity refers to the relative ease and speed with which a chemical substance interacts chemically with other specific substances. For example, alkali metals such as potassium or sodium exhibit violent reactivity when exposed to water, rapidly generating hydrogen gas and caustic hydroxides, whereas noble gases such as helium and argon exhibit near-zero chemical reactivity due to stable valence electron configurations.

2. Flammability & Combustibility

• Flammability: The ability of a substance to ignite and burn readily in the presence of an ignition source at relatively low temperatures (such as ethanol, gasoline, or paper).
• Combustibility: The measure of how easily a substance burns when subjected to sustained high temperatures. Combustion is an intensely exothermic reaction between a fuel and an oxidant (typically atmospheric oxygen, $\text{O}_2$), yielding carbon dioxide, water vapor, ash, and thermal energy.

3. Acidity, Basicity & The pH Scale

The pH scale measures the concentration of dissociated hydrogen ions ($[\text{H}^+]$) in an aqueous solution, spanning logarithmically from 0 to 14:

• Neutral Solutions ($\text{pH} = 7$): Pure distilled water at $25^\circ\text{C}$ exhibits a neutral balance where hydrogen ion concentration equals hydroxide ion concentration ($[\text{H}^+] = [\text{OH}^-]$).
• Acids ($\text{pH} < 7$): Substances that release hydrogen ions ($[\text{H}^+]$) in water.
  • Characteristics: Taste sour (e.g., citric acid in lemons, acetic acid in vinegar), conduct electricity in solution, react with active metals (such as zinc and magnesium) to liberate flammable hydrogen gas ($\text{H}_2$), and react with carbonates (such as baking soda) to liberate carbon dioxide gas ($\text{CO}_2$).
  • Litmus Paper Response: Acids turn blue litmus paper red (mnemonic: Acid turns red, Danger ahead).
• Bases / Alkaline Substances ($\text{pH} > 7$): Substances that release hydroxide ions ($[\text{OH}^-]$) or accept protons in water.
  • Characteristics: Taste bitter, feel slippery or soapy to the touch (due to saponification of skin oils), and act as powerful cleaning agents (e.g., sodium hydroxide in drain openers, household ammonia, sodium hypochlorite bleach, sodium bicarbonate baking soda).
  • Litmus Paper Response: Bases turn red litmus paper blue (mnemonic: Base turns Blue).
• Neutralization: Mixing an acid with a base in stoichiometric proportions produces a neutral salt and water: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$. For example: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$.

4. Oxidation & Rusting

Oxidation occurs when a chemical element combines with oxygen or loses electrons during a reaction:

• Corrosion and Rusting: When iron ($\text{Fe}$) is exposed to moisture and oxygen over time, it undergoes slow chemical oxidation to form hydrated iron(III) oxide (rust):

$4\text{Fe} + 3\text{O}_2 + 6\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$

The iron(III) hydroxide gradually loses water to form hydrated iron(III) oxide ($\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$), the reddish-brown rust students see.

Rust is a distinct chemical substance with different properties than metallic iron: it is brittle, porous, non-conductive, flaky reddish-brown, and significantly less magnetic.

• Tarnish: Silver surfaces reacting with atmospheric sulfur compounds to produce dark silver sulfide ($\text{Ag}_2\text{S}$).
• Enzymatic Browning: When fruit flesh (such as sliced apples, bananas, or potatoes) is cut and exposed to atmospheric oxygen, polyphenol oxidase enzymes catalyze the oxidation of phenols into brown melanin pigments.

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Physical Changes versus Chemical Changes

Distinguishing physical changes from chemical changes represents one of the most foundational and frequently tested concepts on the TExES science exam.

PHYSICAL VS. CHEMICAL CHANGES AT A GLANCE

┌─────────────────────────────────────────┐   ┌─────────────────────────────────────────┐
│            PHYSICAL CHANGE              │   │             CHEMICAL CHANGE             │
│  - Form/State changes                   │   │  - Bonds broken and formed              │
│  - Molecular identity UNCHANGED         │   │  - NEW substance(s) created             │
│  - No new substances formed             │   │  - Distinct new properties              │
│  - Usually reversible by physical means │   │  - Difficult/impossible to reverse      │
│  E.g., Melting ice, dissolving salt     │   │  E.g., Burning wood, rusting iron       │
└─────────────────────────────────────────┘   └─────────────────────────────────────────┘

1. Physical Changes

A physical change modifies the physical form, shape, size, or state of matter without altering its internal chemical structure or molecular identity. The atoms remain bonded together in their original molecular formulas:

• Changes of State: Melting ice into liquid water, freezing water into ice, evaporating ethanol, boiling water into steam, and condensing humidity onto a cold window are all physical changes. In every phase, the substance remains chemically $\text{H}_2\text{O}$.
• Mechanical Alterations: Cutting paper, chopping wood, crushing an aluminum soda can, grinding a mineral crystal into powder, and shredding cloth change dimensions and surface area but alter zero chemical bonds.
• Dissolving Solutes (Crucial Exam Concept!): When table salt ($\text{NaCl}$) dissolves in liquid water, the ionic crystal separates into hydrated sodium ions ($\text{Na}^+$) and chloride ions ($\text{Cl}^-$) surrounded by water molecules. No new chemical substance is formed. The water and salt retain their original chemical identities and can be completely separated by physical evaporation. Dissolving sugar or salt in water is always a physical change!

2. Chemical Changes (Chemical Reactions)

A chemical change (or chemical reaction) occurs when chemical bonds within reactant substances are broken and new bonds are formed, reorganizing atoms into one or more entirely new products possessing distinct physical and chemical properties. Chemical changes cannot be undone by simple physical separation techniques such as filtering, boiling, or magnetic sorting.

3. The Five Diagnostic Indicators of Chemical Reactions

Elementary science teachers must instruct students to search for empirical evidence signaling that a chemical reaction has taken place:

1. Unexpected Color Change: An abrupt shift in color that cannot be explained by simple mechanical mixing or dilution (e.g., clear vinegar and colorless phenolphthalein turning bright magenta upon adding sodium hydroxide; an iron nail developing reddish-brown rust; sliced green apples turning brown).
2. Production of a Gas (Effervescence / Bubbling): The spontaneous release of gas bubbles without the application of external heat. For example, mixing baking soda (sodium bicarbonate) with vinegar (acetic acid) spontaneously produces vigorous bubbling of carbon dioxide gas ($\text{CO}_2$). Note: Bubbles produced when water is heated to boiling are merely physical vapor, not evidence of a chemical reaction.
3. Formation of a Precipitate: An insoluble solid that spontaneously forms and settles out of a solution when two clear liquid solutions are mixed together. For example, mixing aqueous solutions of potassium iodide and lead(II) nitrate instantly produces a bright yellow precipitate of solid lead(II) iodide.
4. Temperature Change (Thermal Energy Transfer):
   • Exothermic Reactions: Release thermal energy into the surrounding environment, causing the reaction vessel to feel warm or hot (e.g., combustion of wood, chemical hand warmers, acid-base neutralization).
   • Endothermic Reactions: Absorb thermal energy from the immediate environment, causing the temperature to drop and the reaction vessel to feel cold to the touch (e.g., baking soda reacting with vinegar or citric acid feels cool; photosynthesis absorbs light energy). Caution: a temperature change alone does not prove a chemical reaction. Instant cold packs get cold because ammonium nitrate dissolves in water, an endothermic physical process.
5. Emission of Light, Sound, or Odor: Chemical bonds releasing energy as electromagnetic photons (e.g., glow sticks undergoing chemiluminescence; fireworks exploding with brilliant flashes and loud concussive sounds) or yielding distinctive new volatile molecules (e.g., food rotting to emit a foul sulfurous odor; baking bread producing fragrant aromas).

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Comparison: Physical Changes versus Chemical Changes

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Chemical Reactions in Daily Life

The 391 framework lists examples of chemical reactions that occur in daily life:

Photosynthesis and cellular respiration are essentially reverse processes: the products of one are the reactants of the other.

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The Law of Conservation of Mass

Formulated by French chemist Antoine Lavoisier in 1789, the Law of Conservation of Mass states that in any closed system, matter is neither created nor destroyed during physical changes or chemical reactions. The total mass of the starting substances (reactants) must exactly equal the total mass of the resulting substances (products):

$\sum \text{Mass of Reactants} = \sum \text{Mass of Products}$

CONSERVATION OF MASS: BALANCED CHEMICAL REACTION

   Methane Gas     +     Oxygen Gas       ───►    Carbon Dioxide   +     Water Vapor
      CH4          +        2 O2          ───►         CO2         +        2 H2O
  1 Carbon atom         4 Oxygen atoms            1 Carbon atom        4 Hydrogen atoms
  4 Hydrogen atoms                                2 Oxygen atoms       2 Oxygen atoms
  ────────────────────────────────────            ───────────────────────────────────
  Total: 1 C, 4 H, 4 O (Mass = 80 g)              Total: 1 C, 4 H, 4 O (Mass = 80 g)

1. Atomic Balancing in Chemical Equations

Chemical equations represent the conservation of mass symbolically. Atoms are merely rearranged into new configurations:

• Subscripts: The small numbers printed below elemental symbols (such as the "2" in $\text{H}_2\text{O}$) indicate the fixed number of atoms within that specific molecule. Subscripts can never be altered when balancing equations because changing subscripts creates an entirely different compound.
• Coefficients: The whole numbers placed in front of chemical formulas (such as the "2" in $2\text{H}_2\text{O}$) indicate how many units of that molecule participate in the reaction.
• In the combustion of methane ($\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$):
  • Reactants contain: 1 Carbon atom, 4 Hydrogen atoms, and 4 Oxygen atoms.
  • Products contain: 1 Carbon atom, 4 Hydrogen atoms, and 4 Oxygen atoms.
  • Total mass before combustion precisely matches total mass after combustion.

2. Open versus Closed Systems in Elementary Inquiry

A pervasive misconception among elementary students is that burning wood, fizzing antacid tablets, or evaporating puddles causes matter to "disappear" or be destroyed. This misconception stems from conducting investigations in open systems:

• Open System: Both matter and energy can freely exchange with the surrounding environment. When wood burns in an open fireplace, solid wood and oxygen react to produce carbon dioxide, water vapor, and lightweight ash. Because the gaseous products escape into the atmosphere unweighed, the remaining ash has far less mass than the original log, leading novice students to conclude incorrectly that mass was destroyed.
• Closed System: Matter cannot enter or escape across the boundaries. To empirically demonstrate the Law of Conservation of Mass in the classroom, teachers must use closed systems:
  • Classroom Investigation: Place vinegar in an Erlenmeyer flask, place baking soda in an uninflated rubber balloon, and stretch the balloon neck tightly over the flask opening without mixing. Measure the total mass of the apparatus on a balance. Tip the balloon to dump the baking soda into the vinegar. Intense effervescence occurs, inflating the balloon with carbon dioxide gas. When the reaction ceases, the balance registers the exact identical mass as before mixing, proving that when gaseous products are trapped, mass is perfectly conserved.

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Classification of Matter: Pure Substances, Mixtures & Solutions

All matter can be classified based on whether its composition is chemically fixed or physically variable.

TAXONOMY OF MATTER
│
├── Pure Substances (Definite chemical composition; cannot separate physically)
│   ├── Elements ──► Composed of one atom type (e.g., Au, O2, Cu)
│   └── Compounds ─► Two or more elements chemically bonded in fixed ratio (e.g., H2O, NaCl)
│
└── Mixtures (Physical combinations of substances; variable ratios; separate physically)
    ├── Heterogeneous Mixtures ──► Non-uniform distribution; distinct visible phases
    │   ├── Coarse Mixtures ─────► Visible macroscopic parts (e.g., trail mix, salad, granite)
    │   ├── Suspensions ─────────► Large particles that settle over time (e.g., muddy water)
    │   └── Colloids ────────────► Medium particles that scatter light (e.g., milk, fog)
    └── Homogeneous Mixtures ────► Uniform distribution throughout; single phase
        └── Solutions ───────────► Molecular dispersal (e.g., saltwater, brass, air)

1. Pure Substances versus Mixtures

• Pure Substance: Matter that possesses a constant, unalterable chemical composition and uniform properties throughout. Pure substances consist solely of either individual elements or definite compounds. They can only be separated or transformed via chemical reactions.
• Mixture: A physical combination of two or more pure substances in which each substance retains its individual chemical identity and physical properties. Mixtures exhibit variable composition (e.g., you can make saltwater with 5 grams or 20 grams of salt per liter) and can always be separated by physical methods without breaking chemical bonds.

2. Heterogeneous Mixtures

A heterogeneous mixture is non-uniform in composition. Its individual components remain physically segregated, creating distinct macroscopic or microscopic phases:

• Coarse Mixtures: Obvious discrete components that can be sorted by hand (e.g., trail mix with nuts and raisins, a tossed garden salad, gravel, concrete).
• Suspensions: Fluid mixtures containing relatively large solid particles that do not dissolve and will gradually settle to the bottom due to gravity if left undisturbed (e.g., muddy pond water, unhomogenized Italian salad dressing, calamine lotion). Suspensions can be separated by gravity sedimentation or filtration.
• Colloids: Heterogeneous mixtures containing intermediate-sized particles that remain permanently suspended and do not settle out over time (e.g., homogenized milk, mayonnaise, fog, gelatin). Colloids scatter light beams, a phenomenon known as the Tyndall effect.

3. Homogeneous Mixtures (Solutions)

A homogeneous mixture (or solution) is completely uniform in composition throughout. Its constituent particles are dispersed at the molecular or ionic level, exhibiting a single continuous physical phase:

• Aqueous Solutions: Liquid mixtures where water acts as the solvent (e.g., salt water, dissolved sugar water, tea, vinegar).
• Solid Solutions (Alloys): Homogeneous metallic mixtures melted together and cooled (e.g., brass is an alloy of copper and zinc; bronze is copper and tin; steel is iron and carbon).
• Gaseous Solutions: Homogeneous mixtures of gases (e.g., clean air is a gaseous solution consisting of approximately 78% nitrogen, 21% oxygen, 0.9% argon, and trace carbon dioxide and water vapor).

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Methods for Separating Mixtures Based on Physical Properties

Every physical method for separating mixtures exploits a specific difference in the physical properties of the constituent substances. Elementary teachers must help students identify which physical property provides the mechanism for separation.

SEPARATION TECHNIQUES AND EXPLOITED PHYSICAL PROPERTIES
│
├── Magnetism ────────► Magnetic attraction (separates ferromagnetic iron from sand)
├── Sieving / Screening ► Particle size of dry solids (separates gravel from fine sand)
├── Filtration ───────► Particle size and solubility (separates insoluble sand from saltwater)
├── Decanting ────────► Density difference / settling (pours liquid off settled sediment)
├── Evaporation ──────► Boiling point difference (recovers solid salt from water)
├── Distillation ─────► Boiling point difference (collects purified condensed liquid)
└── Chromatography ───► Solubility and capillary affinity (separates ink pigments)

1. Filtration

• Physical Property Exploited: Particle size and solubility.
• Method: The mixture is poured through a porous barrier (such as filter paper supported by a glass funnel). The liquid and dissolved molecular solutes (the filtrate) pass freely through microscopic pores, while insoluble solid particles larger than the pores (the residue) are trapped on the paper.
• Classroom Example: Separating insoluble sand from a mixture of sand and saltwater.

2. Evaporation & Crystallization

• Physical Property Exploited: Boiling point differences between a non-volatile dissolved solid solute and a volatile liquid solvent.
• Method: The solution is heated in an open evaporating dish or allowed to stand at room temperature. The liquid solvent vaporizes into the atmosphere, leaving dry, solid solute crystals behind.
• Classroom Example: Boiling or evaporating saltwater to recover solid white sodium chloride crystals.

3. Distillation

• Physical Property Exploited: Differences in the boiling points of miscible liquids, or separating a liquid solvent from dissolved solids when the solvent itself must be preserved.
• Method: The liquid mixture is heated in a distillation flask to the boiling point of the more volatile substance. The vapor rises, enters a water-cooled condenser tube, cools, condenses back into a pure liquid (the distillate), and drips into a receiving flask.
• Classroom Example: Purifying fresh drinkable water from ocean seawater, or separating ethanol (boiling point $78^\circ\text{C}$) from water (boiling point $100^\circ\text{C}$).

4. Magnetic Separation

• Physical Property Exploited: Magnetism (specifically ferromagnetism).
• Method: A permanent bar magnet wrapped in plastic wrap is passed over a dry mixture. Ferromagnetic particles (iron, nickel, cobalt) are attracted to the magnet, leaving non-magnetic materials behind.
• Classroom Example: Extracting iron filings from a mixture of iron filings, sulfur powder, and sand.

5. Sieving / Screening

• Physical Property Exploited: Particle size among dry solid particles.
• Method: The dry mixture is placed into wire mesh screens with calibrated opening sizes. Smaller particles pass through the mesh, while larger particles are retained.
• Classroom Example: Separating coarse gravel from fine beach sand, or sifting flour in baking.

6. Settling & Decanting

• Physical Property Exploited: Density and solubility.
• Method: In a suspension, denser insoluble solid particles settle to the container bottom under gravity. The clear supernatant liquid is carefully poured (decanted) off the top without disturbing the sediment.
• Classroom Example: Allowing muddy pond water to settle in a jar and gently pouring off the clear surface water, or using a separatory funnel to drain dense vinegar from lighter vegetable oil.

7. Paper Chromatography

• Physical Property Exploited: Differential solubility in a mobile solvent versus molecular adhesion to a stationary medium (capillary action).
• Method: A spot of ink or plant extract is placed near the bottom of a strip of absorbent filter paper. The paper strip is suspended with its bottom edge touching a liquid solvent (water or alcohol). As the solvent wicks upward via capillary action, it dissolves and carries pigment molecules along. Pigments that are highly soluble in the solvent and weakly bound to paper fibers travel rapidly toward the top, while less soluble pigments move slowly, separating into distinct colored bands.
• Classroom Example: Separating black water-soluble marker ink into its constituent cyan, magenta, and yellow dyes.

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Comparison Table: Separation Methods for Mixtures

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Classroom Instructional Strategies & Scenario Application

The 5E Mixture Separation Design Challenge

To prepare elementary students for fifth-grade TEKS assessments, educators should transition from simple teacher demonstrations to student-driven inquiry challenges where students analyze physical properties and formulate multi-step separation plans.

Classroom Context: Ms. Lin presents her 5th-grade science class with a beaker containing a complex dry mixture of four components: iron filings, coarse gravel, fine sand, and table salt ($\text{NaCl}$). The challenge is to separate all four substances into pure, dry piles while documenting the physical property governing each procedural step.

Step-by-Step Student Separation Sequence:

┌───────────────────────────────────────────────────────────────┐
│ 1. DRY MAGNETIC SEPARATION                                    │
│    Pass a plastic-wrapped magnet over the dry mixture.        │
│    ► ISOLATES: Iron filings (exploits Ferromagnetism).        │
└───────────────────────────────┬───────────────────────────────┘
                                │ (Remaining: Gravel, Sand, Salt)
                                ▼
┌───────────────────────────────────────────────────────────────┐
│ 2. DRY SIEVING / SCREENING                                    │
│    Pour remaining mixture through a wire mesh screen.         │
│    ► ISOLATES: Coarse gravel on mesh (exploits Particle Size).│
└───────────────────────────────┬───────────────────────────────┘
                                │ (Remaining: Fine Sand and Salt)
                                ▼
┌───────────────────────────────────────────────────────────────┐
│ 3. DISSOLUTION IN WATER                                       │
│    Add warm water and stir thoroughly until salt dissolves.   │
│    ► Salt forms aqueous solution; sand remains insoluble.     │
└───────────────────────────────┬───────────────────────────────┘
                                │
                                ▼
┌───────────────────────────────────────────────────────────────┐
│ 4. GRAVITY FILTRATION                                         │
│    Pour suspension through filter paper seated in a funnel.   │
│    ► ISOLATES: Insoluble sand on paper (Particle Size).       │
└───────────────────────────────┬───────────────────────────────┘
                                │ (Filtrate: Dissolved Saltwater)
                                ▼
┌───────────────────────────────────────────────────────────────┐
│ 5. THERMAL EVAPORATION                                        │
│    Heat saltwater filtrate in an evaporating dish on hotplate.│
│    Water vaporizes into steam (exploits Boiling Point).       │
│    ► ISOLATES: Pure, dry table salt crystals in dish.         │
└───────────────────────────────────────────────────────────────┘

Pedagogical Significance: Notice why procedural order is essential! If students add water in Step 1 before removing the iron filings, the iron will begin to rust (a chemical change), and separating wet iron filings from wet sand becomes extraordinarily difficult. Conducting magnetic extraction and sieving dry before introducing water preserves the physical identities of all four materials.