21.1 Physical Properties of Matter, States & Density

Overview & Exam Relevance

Competency 008 of the TExES Core Subjects EC-6 Science exam (Subject Exam 904) assesses your mastery of the physical properties of matter, states of matter, and the thermodynamic processes governing phase changes. In elementary education, physical science provides the empirical entry point for scientific inquiry. The elementary TEKS move from sorting materials by observable properties in the primary grades to measuring and comparing properties such as mass, volume, temperature, magnetism, physical state, and relative density (sinking and floating) by Grade 4. Grade 5 adds solubility and whether materials conduct or insulate thermal and electrical energy. Calculating density with $D = m/V$ begins in middle school, but the 391 expects you to know it as teacher-level content.

On the TExES 391 exam, you must demonstrate more than rote recall of definitions. You will be evaluated on your pedagogical capacity to design hands-on investigations, guide students in using laboratory measurement tools accurately, identify and remediate pervasive student misconceptions (such as confusing mass with weight, or assuming that density changes when an object is cut in half), and interpret scientific representations such as heating and cooling curves.

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Atomic Structure & Particulate Nature of Matter

All matter in the universe is particulate. Matter is formally defined as anything that possesses mass and occupies volume (takes up three-dimensional space). Light, sound, magnetic fields, and thoughts are forms of energy or abstract phenomena; they are not matter because they lack mass and volume.

THE COMPOSITION OF MATTER
│
├── Atom ──────────► Smallest unit of an element retaining chemical identity
│   ├── Protons ───► Positive charge (+1); 1 amu; located in central nucleus
│   ├── Neutrons ──► Zero charge (0); 1 amu; located in central nucleus
│   └── Electrons ─► Negative charge (-1); ~0 amu; move in electron cloud
│
├── Element ───────► Pure substance composed of only one type of atom (e.g., Fe, O, Au)
├── Molecule ──────► Two or more atoms chemically bonded together (e.g., O2, H2O)
└── Compound ──────► Two or more DIFFERENT elements chemically combined in fixed ratios (e.g., H2O, NaCl)

1. Subatomic Architecture

Every atom consists of a dense central core termed the atomic nucleus, enveloped by a diffuse region of rapid particle motion known as the electron cloud:

• Protons: Subatomic particles carrying a unit positive electrical charge ($+1$) and an atomic mass of approximately $1\text{ atomic mass unit (amu)}$. The number of protons in the nucleus represents the atomic number ($Z$) and determines the elemental identity of the atom (e.g., every atom with 6 protons is carbon, regardless of neutron or electron count).
• Neutrons: Electrically neutral subatomic particles ($0\text{ charge}$) residing in the nucleus with protons, possessing a mass of approximately $1\text{ amu}$. The total number of protons and neutrons is the atom's mass number ($A = \text{protons} + \text{neutrons}$), which accounts for nearly all of its mass.
• Electrons: Tiny subatomic particles carrying a unit negative electrical charge ($-1$) and an extremely small mass (about 1/1836 the mass of a proton, treated as negligible in basic mass calculations). Electrons occupy quantized energy levels or orbitals within the electron cloud. The electrons in the outermost unfilled energy level are termed valence electrons. Valence electrons govern how atoms interact, bond, and participate in chemical reactions.

2. Elements, Molecules, and Compounds

• Element: A pure chemical substance that cannot be broken down into simpler substances by ordinary chemical means. Each element consists of atoms containing the identical number of nuclear protons. Currently, 118 elements are recognized, arranged systematically on the Periodic Table.
• Molecule: An electrically neutral entity formed when two or more atoms share electrons via covalent bonds. A molecule can consist of identical atoms (such as diatomic oxygen gas, $\text{O}_2$) or different atoms (such as methane, $\text{CH}_4$).
• Compound: A pure substance composed of two or more different chemical elements chemically united in a definite, unalterable stoichiometric proportion (such as water, $\text{H}_2\text{O}$, or sodium chloride, $\text{NaCl}$). Crucially, the physical and chemical properties of a compound differ completely from those of its constituent elements. For instance, elemental sodium ($\text{Na}$) is an explosive, soft alkali metal that reacts violently with moisture, and chlorine ($\text{Cl}_2$) is a suffocating toxic green gas; when chemically bonded in a 1:1 ratio, they produce sodium chloride (table salt), an edible crystalline compound vital for animal physiology.

3. Organization of the Periodic Table

The Periodic Table of Elements organizes elements sequentially by increasing atomic number (number of protons) and groups them according to recurring periodic trends in electronic configuration:

• Periods: The seven horizontal rows of the table. The period number corresponds directly to the number of occupied electron shells or principal energy levels in an unexcited atom.
• Groups (Families): The eighteen vertical columns. Elements in the same group possess the same number of valence electrons, resulting in remarkably similar chemical properties, bonding tendencies, and reactivity patterns (e.g., Group 1 Alkali Metals are fiercely reactive with water; Group 18 Noble Gases possess full valence octets and are chemically inert).
• Metals: Positioned to the left and center of the periodic table. Metals are typically solid at room temperature (with the notable exception of liquid mercury), lustrous (shiny), malleable (capable of being hammered into thin sheets without shattering), ductile (capable of being drawn into thin wires), dense, and excellent conductors of both heat and electricity.
• Nonmetals: Positioned on the upper right side of the table (plus hydrogen on the top left). Nonmetals lack metallic luster, are brittle in solid form, have lower densities, and act as thermal and electrical insulators.
• Metalloids (Semiconductors): Straddle the diagonal stair-step boundary separating metals from nonmetals (including silicon, germanium, arsenic, and boron). Metalloids exhibit intermediate properties; for example, silicon possesses metallic luster but is brittle, and conducts electricity under specific thermal conditions, making it indispensable for microelectronics.

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Measurable & Observable Physical Properties

A physical property is an observable or measurable characteristic of a substance that can be evaluated without altering the substance's chemical identity or molecular composition.

Intensive versus Extensive Physical Properties

Elementary science educators must understand the foundational distinction between extensive and intensive properties:

• Extensive Properties: Depend directly on the quantity or sample size of matter present. If you double the amount of substance, extensive properties double. Examples include mass, volume, length, and total thermal energy.
• Intensive Properties: Inherent to the nature of the substance itself and completely independent of sample size. Whether you analyze a single drop of pure water or an entire swimming pool of pure water, intensive properties remain identical. Examples include density, melting point, boiling point, solubility, electrical conductivity, and color. Intensive properties serve as scientific "fingerprints" to identify unknown substances.

1. Mass versus Weight

A pervasive misconception tested on the TExES exam is the conflation of mass and weight:

• Mass: The measure of the actual quantity of matter contained within an object, directly proportional to the total number of atoms it comprises. Mass is invariant; an object maintains the identical mass whether measured on Earth, on the Moon, or floating in deep space. Mass is measured in grams ($\text{g}$) or kilograms ($\text{kg}$) using a triple-beam balance or a calibrated digital pan balance.
• Weight: The measure of the downward gravitational force exerted on an object's mass by a planetary body ($W = mg$, where $m$ is mass and $g$ is gravitational acceleration). Because weight is a force, it fluctuates depending on local gravitational field strength. On the Moon, where gravity is approximately one-sixth ($1/6$) that of Earth, an astronaut's weight decreases to one-sixth of their terrestrial weight, while their mass remains completely unchanged. Weight is measured in Newtons ($\text{N}$) or pounds using a spring scale.

2. Volume & The Water Displacement Method

Volume is the quantity of three-dimensional space an object occupies. Volume measurement depends on the geometric regularity of the sample:

• Regular Rectangular Solids: Volume is determined mathematically using a metric ruler to measure linear dimensions, applying the formula:

$V = \text{length} \times \text{width} \times \text{height}$

Expressed in cubic centimeters ($\text{cm}^3$) or cubic meters ($\text{m}^3$).

• Liquids: Measured directly in calibrated glassware, preferably a graduated cylinder. The liquid's volume must be read at eye level at the very bottom of the meniscus (the curved liquid surface created by surface tension and adhesive forces between the liquid and container walls) to avoid parallax error. Expressed in milliliters ($\text{mL}$) or liters ($\text{L}$).
• Irregular Solids (Water Displacement): For objects with irregular geometries (such as rocks, metal bolts, or mineral samples), volume cannot be computed by simple geometric formulas. Instead, Archimedes' water displacement principle is utilized:
  1. Pour a known baseline volume of water into a graduated cylinder and record the initial volume ($V_{\text{initial}}$).
  2. Gently slide the submerged solid down the tilted glass wall to avoid splashing or air bubbles.
  3. Record the new, elevated water volume ($V_{\text{final}}$).
  4. Calculate sample volume: $\Delta V = V_{\text{final}} - V_{\text{initial}}$. (Conversion equivalence: $1\text{ mL} = 1\text{ cm}^3$ of solid volume).

3. Density & Relative Buoyancy

Density is the physical ratio of an object's mass to its volume, describing how tightly packed matter is within a given spatial boundary:

$D = \frac{m}{V}$

Density is expressed in units of grams per cubic centimeter ($\text{g/cm}^3$) for solids, or grams per milliliter ($\text{g/mL}$) for liquids.

• Relative Density in Water: Liquid water at $4^\circ\text{C}$ has an established standard density of $1.0\text{ g/cm}^3$ (or $1.0\text{ g/mL}$).
  • If an object has a density greater than $1.0\text{ g/cm}^3$ ($D > 1.0$), it sinks in water because downward gravitational pull exceeds upward buoyant force.
  • If an object has a density less than $1.0\text{ g/cm}^3$ ($D < 1.0$), it floats in water until the mass of displaced water equals the total mass of the object.
  • If an object has a density exactly equal to $1.0\text{ g/cm}^3$ ($D = 1.0$), it exhibits neutral buoyancy, remaining suspended at any depth.
• Density as an Intensive Property: Novice students frequently predict that cutting a block of wood or paraffin wax in half will halve its density. Teachers must emphasize that cutting an object in half divides both mass and volume by two simultaneously, leaving the mathematical ratio ($m/V$) completely unchanged.

4. Magnetism

Magnetism is an intrinsic physical property involving the attraction of materials to magnetic fields. In elementary classrooms, students must understand that not all metals are magnetic:

• Ferromagnetic Materials: Strongly attracted to permanent magnets. The primary ferromagnetic elements are iron ($\text{Fe}$), nickel ($\text{Ni}$), and cobalt ($\text{Co}$), along with alloys containing them, such as steel.
• Non-Magnetic Metals: Metals that do not experience magnetic attraction under classroom conditions include aluminum, copper, brass (copper-zinc alloy), lead, silver, and gold. A standard TExES question involves presenting a set of metallic objects (such as an iron nail, an aluminum beverage can, a copper penny, and a brass paper fastener) and requiring candidates to identify which items can be sorted using a bar magnet.

5. Thermal & Electrical Conductivity

Conductivity describes how readily energy flows through a material:

• Thermal Conductivity: The rate at which thermal energy (heat) conducts through a substance via molecular collisions. Metals (such as copper, aluminum, and cast iron) are superior thermal conductors, which is why cookware is fabricated from metals. Thermal insulators (such as wood, plastic, styrofoam, fiberglass, and cork) impede heat flow because their atomic structures lack freely moving electrons, making them ideal for pot handles and thermal drink containers.
• Electrical Conductivity: The ability of a substance to allow electric charges (electrons) to flow freely through an electrical circuit. Conductors include metals (especially copper and silver) and aqueous ionic solutions (saltwater). Electrical insulators (such as rubber, vinyl, glass, dry air, and porcelain) tightly bind their valence electrons, preventing electric current from leaking or causing shocks.

6. Solubility

Solubility is the physical ability of one substance (the solute) to dissolve completely into another substance (the solvent) to create a homogeneous mixture called a solution:

• Solute: The substance that is dissolved and dispersed at the molecular scale (e.g., sucrose sugar, table salt).
• Solvent: The dissolving medium present in greater volume (e.g., water, ethanol). Water is designated the "universal solvent" because its polar molecular geometry allows it to dissolve an exceptionally wide variety of ionic compounds and polar covalent substances.
• Factors Affecting Dissolution Rate of Solids in Liquids:
  • Temperature: Elevating solvent temperature increases molecular kinetic energy, accelerating solvent-solute molecular impacts and breaking crystal lattices faster.
  • Surface Area (Particle Size): Crushing a solid into a fine powder dramatically increases exposed surface area, allowing more solvent molecules to contact solute particles simultaneously.
  • Agitation (Stirring): Mechanically disperses solute particles away from crystal surfaces, maintaining steep concentration gradients.
• Saturation: A solution is saturated when it contains the maximum concentration of dissolved solute possible at a given temperature; any additional solute added will precipitate to the bottom.

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States of Matter & Kinetic Molecular Theory

The Kinetic Molecular Theory (KMT) provides the fundamental model explaining the physical states of matter:

1. All matter is composed of microscopic particles (atoms or molecules).
2. These particles are in constant, random, ceaseless motion.
3. The average kinetic energy of the particles is directly proportional to the absolute temperature (measured in Kelvin) of the substance.
4. The physical state of a substance is determined by the competition between the thermal kinetic energy of the particles (which tends to disperse them) and the intermolecular attractive forces pulling them together.

PARTICLE ARRANGEMENT ACROSS THE FOUR STATES OF MATTER

   SOLID                 LIQUID                  GAS                  PLASMA
┌─────────┐           ┌─────────┐            ┌─────────┐           ┌─────────┐
│ ● ● ● ● │           │  ●  ●   │            │ ●       │           │ ⊕ ⊖   ⊕ │
│ ● ● ● ● │           │ ●  ●  ● │            │      ●  │           │   ⊖ ⊕   │
│ ● ● ● ● │           │   ●  ●  │            │  ●      │           │ ⊕   ⊖ ⊕ │
└─────────┘           └─────────┘            └─────────┘           └─────────┘
Vibrating in        Sliding past each       Rapid, random,        Ionized gas with
fixed lattice       other; fluid            far apart             free electrons

1. Solid

In a solid, intermolecular attractive forces dominate over kinetic energy. Particles are packed tightly together in rigid, fixed geometric patterns (crystalline lattices) or closely bound irregular networks (amorphous solids). The particles cannot slide or translate from place to place; they can only vibrate in place around fixed equilibrium positions. Consequently, solids have a definite (fixed) shape and a definite (fixed) volume, and are virtually incompressible.

2. Liquid

In a liquid, thermal kinetic energy is sufficient to partially overcome rigid intermolecular forces. Particles remain in close physical contact with one another, but they possess enough translational kinetic energy to slide, roll, and flow past one another. Consequently, a liquid possesses a definite (fixed) volume, but an indefinite (variable) shape; it conforms fluidly to the contours of whatever container it occupies, filling from the bottom up. Liquids exhibit very low compressibility.

3. Gas

In a gas, thermal kinetic energy completely overwhelms intermolecular attractions. Particles are separated by immense empty distances relative to their microscopic size. Gas particles move in rapid, random, straight-line trajectories until colliding elastically with one another or the container walls. Consequently, gases have indefinite (variable) shape and indefinite (variable) volume; they spontaneously expand to fill the entirety of any container. Because of the vast empty spaces between particles, gases are highly compressible.

4. Plasma

Plasma is a superheated, highly energetic state of matter formed when gas particles absorb immense thermal or electrical energy. The energetic collisions strip electrons away from atomic nuclei, creating a roiling mixture of free negative electrons and positively charged ions. While plasma behaves like an electrically conductive fluid similar to gas in having indefinite shape and volume, its charged constituents make it responsive to electromagnetic fields. Plasma is the most abundant state of matter in the visible universe, composing stars, lightning bolts, the solar wind, the aurora borealis, and commercial neon lighting tubes.

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Comparison of States of Matter

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Phase Changes & Thermal Energy Dynamics

A phase change (or transition of state) is a reversible physical transformation of a substance from one state of matter to another, driven by the addition or removal of thermal energy.

PHASE TRANSITION TERMINOLOGY

         ┌────────────── Sublimation (Endothermic) ─────────────┐
         │                                                       ▼
    ┌─────────┐         Melting (Endo)         ┌──────────┐   Vaporization (Endo)   ┌─────────┐
    │  SOLID  │ ─────────────────────────────► │  LIQUID  │ ──────────────────────► │   GAS   │
    │         │ ◄───────────────────────────── │          │ ◄────────────────────── │         │
    └─────────┘         Freezing (Exo)         └──────────┘    Condensation (Exo)   └─────────┘
         ▲                                                       │
         └────────────── Deposition (Exothermic) ────────────────┘

1. Phase Transition Terminology

• Melting: Solid → Liquid. Thermal energy is absorbed (endothermic), causing particles to vibrate vigorously until breaking free from lattice constraints.
• Freezing (Solidification): Liquid → Solid. Thermal energy is released (exothermic); particles decelerate and intermolecular attractions lock them into fixed lattice sites. Pure water freezes at $0^\circ\text{C}$ ($32^\circ\text{F}$).
• Vaporization: Liquid → Gas. Thermal energy is absorbed (endothermic). Vaporization manifests in two distinct physical modes:
  • Evaporation: A slow surface phenomenon occurring at any temperature below the boiling point. High-energy molecules at the liquid surface overcome intermolecular forces and escape into the vapor phase.
  • Boiling: A rapid transition occurring throughout the entire body of the liquid when vapor pressure equals surrounding atmospheric pressure. Pure water boils at $100^\circ\text{C}$ ($212^\circ\text{F}$) at 1 atmosphere of pressure.
• Condensation: Gas → Liquid. Thermal energy is released (exothermic). Gaseous particles lose kinetic energy upon contacting cooler surfaces, coalescing into liquid droplets (e.g., dew on morning grass, water droplets on the exterior of an iced beverage cup).
• Sublimation: Solid → Gas directly without passing through a liquid intermediate (endothermic). Observed in solid carbon dioxide ("dry ice") at room temperature, mothballs (naphthalene), and freeze-drying processes.
• Deposition: Gas → Solid directly without passing through an intermediate liquid phase (exothermic). An example is subfreezing water vapor depositing directly as crystalline frost onto automobile windshields on winter mornings.

2. Heating and Cooling Curves: Latent Heat & Temperature Plateaus

A critical concept tested on science certification exams is the thermodynamic behavior of a substance during phase transitions. When a pure substance (like ice) is heated at a steady rate, a plot of temperature versus time yields a heating curve:

HEATING CURVE FOR WATER

Temp (°C) ▲
          │                                           Gas (Steam) heating
   100°C ─┼─ - - - - - - - - - - - - - ┌─────────────────────────►
          │                           │ ◄── BOILING PLATEAU (Liquid -> Gas)
          │                          /      (Temp constant at 100°C)
          │                         / ◄───── Liquid water warming
          │                        /
     0°C ─┼─ - - - - - ┌──────────┘ ◄─────── MELTING PLATEAU (Solid -> Liquid)
          │           /                      (Temp constant at 0°C)
          │          / ◄──────────────────── Solid ice warming
          │         /
          └────────┴───┴──────────┴───────────┴──────────► Time / Heat Added

• Sloped Regions (Single States): During the segments where only solid, only liquid, or only gas exists, added thermal energy directly increases the average kinetic energy of the molecules. Consequently, the thermometer registers a steady rise in temperature.
• Flat Plateaus (Phase Transitions): During melting ($0^\circ\text{C}$) and boiling ($100^\circ\text{C}$), the temperature line remains completely flat and horizontal, despite continuous thermal input from the heat source. Why does temperature not rise during melting or boiling?
  • Scientific Explanation: Temperature is exclusively a measure of average translational kinetic energy. During a phase change, added thermal energy is absorbed as latent heat (latent heat of fusion during melting; latent heat of vaporization during boiling). This energy is completely consumed to overcome and break intermolecular electrostatic attractions (such as hydrogen bonds in water) and increase potential energy, rather than increasing molecular kinetic speed. Only when the last droplet of liquid has vaporized does added heat resume raising kinetic energy, causing the temperature of steam to rise above $100^\circ\text{C}$.

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Classroom Instructional Strategies & Scenario Application

Overcoming Common Student Misconceptions

Exemplary Classroom Inquiry Scenario

Classroom Context: Mr. Alvarez is leading a 5th-grade extension investigation. (The Grade 5 TEKS stop at relative density, sinking and floating; calculating $D = m/V$ comes in middle school, so this task stretches advanced students and models the teacher-level content the 391 expects.) Each cooperative team is given a mystery cube ($2.0\text{ cm} \times 2.0\text{ cm} \times 2.0\text{ cm}$) made of an unknown solid material and tasked with identifying the material from a reference table of physical properties.

Investigation Steps:

1. Volume Calculation: Students measure dimensions with a metric ruler: $V = 2.0\text{ cm} \times 2.0\text{ cm} \times 2.0\text{ cm} = 8.0\text{ cm}^3$.
2. Mass Measurement: Students calibrate a triple-beam balance, place the cube on the pan, and record a mass of $21.6\text{ grams}$.
3. Density Calculation: Students apply the formula $D = m/V = 21.6\text{ g} / 8.0\text{ cm}^3 = 2.7\text{ g/cm}^3$.
4. Testing Additional Properties: Students test with a bar magnet (no magnetic attraction observed) and insert the cube into a simple circuit with a battery and light bulb (the bulb illuminates, confirming electrical conductivity).
5. Data Deduction: Consulting a standard physical properties table, students compare their findings:
   • Iron ($D = 7.9\text{ g/cm}^3$, magnetic, conductor)
   • Aluminum ($D = 2.7\text{ g/cm}^3$, non-magnetic, conductor)
   • Acrylic Plastic ($D = 1.2\text{ g/cm}^3$, non-magnetic, insulator) Students definitively deduce that their mystery specimen is aluminum, integrating multiple physical properties through empirical evidence.