7.1 States of Matter & Physical Changes

States of Matter & Kinetic Particle Theory

Quick Summary: All physical objects in the universe are composed of matter—defined as anything that has mass and occupies volume (takes up space). Matter exists in three fundamental physical states: solid, liquid, and gas. The Kinetic Particle Theory of Matter explains that all matter consists of minute particles (atoms, molecules, or ions) that are in continuous motion. The physical state of a substance depends on the balance between the kinetic energy of its particles and the attractive forces drawing them together.

The Kinetic Particle Theory of Matter

To understand how matter behaves, scientists rely on the kinetic particle theory. This model makes four key assertions:

1. Composition: All matter is built from extremely small particles.
2. Continuous Motion: Particles are constantly moving or vibrating; they possess kinetic energy (energy of motion).
3. Inter-particle Spaces: Empty spaces exist between particles.
4. Attractive Forces: Forces of attraction (intermolecular forces) exist between neighbouring particles, holding them together with varying strength depending on temperature and state.

+-------------------------------------------------------------------------+
|                         STATES OF MATTER SUMMARY                        |
+-------------------+--------------------+--------------------------------+
| State             | Particle Packing   | Motion & Energy                |
+-------------------+--------------------+--------------------------------+
| Solid             | Tightly packed in  | Vibrates in fixed positions;   |
|                   | regular lattice    | low kinetic energy             |
| Liquid            | Closely packed in  | Slides past each other;        |
|                   | random order       | moderate kinetic energy        |
| Gas               | Widely separated;  | Moves rapidly in all directions|
|                   | random layout      | high kinetic energy            |
+-------------------+--------------------+--------------------------------+

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Particle Arrangements in Solids, Liquids, and Gases

1. Solids

In a solid, particles are packed very closely together in a fixed, regular pattern known as a lattice structure. The attractive forces between solid particles are exceptionally strong, preventing them from moving out of their assigned positions. As a result, solid particles can only vibrate to and fro about fixed points.

• Shape & Volume: Solids possess a definite (fixed) shape and a definite volume.
• Compressibility & Density: Solids cannot be compressed easily because the particles are already touching. They generally have high density.
• Real-World Examples: Ice blocks, sugar crystals, copper wires, iron nails, granite rocks.

2. Liquids

In a liquid, particles remain close to one another, but they are arranged randomly without a rigid lattice. The attractive forces between liquid particles are moderate—strong enough to keep the particles contained within a bulk liquid, but weak enough to permit individual particles to slide and flow past one another.

• Shape & Volume: Liquids have a definite volume but no fixed shape; they take the shape of whatever container they occupy.
• Compressibility & Density: Liquids are virtually incompressible under normal conditions because particle spacing remains minimal. They have moderate to high density.
• Real-World Examples: Water, coconut oil, seawater, liquid mercury, fruit juice.

3. Gases

In a gas, particles are separated by vast distances relative to their size. The attractive forces between gas particles are negligible or practically non-existent. Gas particles possess high kinetic energy and move rapidly and randomly in all directions in straight paths until colliding with each other or the walls of their container.

• Shape & Volume: Gases have no fixed shape and no definite volume; they expand completely to fill every available space in a sealed container.
• Compressibility & Density: Gases are highly compressible because large empty spaces exist between particles. They have very low density.
• Real-World Examples: Oxygen, nitrogen, carbon dioxide, water vapour, methane.

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Comparison of Physical Properties Across States

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Scientific Inquiry: Hypotheses, Variables and Fair Tests

Science Target 2 of the official CPEA standards is Problem-solving and communication skills and strategies. It asks you to form and test a hypothesis and interpret results, predict or describe the outcome in advance of an event, design procedures to obtain information, collect and record data using tables, charts and diagrams, and analyse data to form conclusions. Those skills are examined on the Science paper and assessed again in your project.

The Vocabulary of an Investigation

What Makes a Test "Fair"

A fair test changes one variable only and holds every other variable constant. If you change two things at once, you cannot tell which one caused the result.

Worked example, taken from the CXC Science specimen paper. Five grams of sugar are added to each of several beakers of water and stirred at the same rate, and a thermometer is later placed in each beaker.

• Independent variable: the water temperature in each beaker.
• Dependent variable: how quickly the sugar dissolves.
• Controlled variables: the mass of sugar (5 g in every beaker), the volume of water, and the stirring rate.

Notice why the paper tells you the stirring rate was the same: stirring speeds up dissolving, so if one beaker were stirred harder the test would no longer be fair. Examiners include that detail deliberately, and questions often ask why a particular quantity was kept the same.

Expected conclusion: sugar dissolves faster in hotter water, because the particles move faster and collide with the sugar more often.

Predicting Before You Measure

A prediction states the expected outcome and the reason: "The sugar will dissolve fastest in the hottest beaker, because warmer particles move faster." A prediction without a reason earns less credit than one with a reason, and the same rule applies to your project write-up.

Recording and Communicating Results

Record measurements in a table with clear headings and units, then display them in a bar chart or line graph so a pattern is visible. Finish by stating whether the results supported the hypothesis. A result that disproves your hypothesis is still a valid result — science does not require you to be right, only to be honest about what the data show.

Phase Changes & Thermal Expansion

Key Concept: A phase change (or state change) is a reversible physical transformation from one state of matter to another. Phase changes occur when thermal energy is added to or removed from a substance. Adding thermal energy increases particle kinetic energy, overcoming attractive forces. Removing thermal energy decreases particle kinetic energy, allowing attractive forces to draw particles closer together.

Physical Phase Transformations

Heating or cooling a substance can cause it to change state without altering its chemical identity (water remains H₂O whether it exists as solid ice, liquid water, or gaseous steam).

               +-----------------------------------+
               |           SUBLIMATION             |
               |  (Solid directly turns to Gas)    |
               +-----------------------------------+
                                 ^ 
                                 |
          +----------+  Melting  +----------+ Evaporation/ +----------+
          |          |---------->|          | ------------>|          |
          |  SOLID   |           |  LIQUID  |  Boiling     |   GAS    |
          |          |<----------|          | <------------|          |
          +----------+ Freezing  +----------+ Condensation +----------+
                                 | 
                                 v
               +-----------------------------------+
               |            DEPOSITION             |
               |  (Gas directly turns to Solid)    |
               +-----------------------------------+

1. Melting (Solid to Liquid)

When a solid is heated, its particles absorb thermal energy and vibrate more vigorously. At a specific temperature called the melting point (0°C for pure ice at sea level), the particles gain sufficient kinetic energy to break free from their rigid lattice arrangements, transitioning into a liquid.

2. Freezing / Solidification (Liquid to Solid)

When a liquid is cooled, thermal energy is removed. The particles slow down, losing kinetic energy. As temperature reaches the freezing point, attractive forces pull the particles back into a fixed lattice, forming a solid.

3. Evaporation vs. Boiling (Liquid to Gas / Vaporisation)

• Evaporation: The gradual transition of a liquid into a gas at temperatures below the boiling point. Evaporation occurs only at the liquid surface, where energetic surface particles absorb ambient thermal energy, break free, and escape as vapour. Evaporation causes a cooling effect (e.g., sweating cools human skin in hot Caribbean climates).
• Boiling: Rapid vaporisation occurring throughout the entire body of liquid at a specific, fixed temperature called the boiling point (100°C for pure water at standard atmospheric pressure). Bubbles of gas form within the bulk liquid and rise to the surface.

4. Condensation (Gas to Liquid)

When a gas is cooled, its particles lose kinetic energy and slow down. As particles come closer together, attractive forces pull them into the liquid state (e.g., water droplets forming on the exterior of a cold drink glass or dew forming on leaves overnight).

5. Sublimation & Deposition

• Sublimation: Direct change from solid to gas without passing through the liquid phase (e.g., solid carbon dioxide / dry ice, mothballs / naphthalene, solid iodine when heated).
• Deposition: Direct change from gas to solid without forming a liquid (e.g., frost formation on cold surfaces on high mountains).

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Thermal Expansion and Contraction

When substances undergo temperature changes without changing state, their dimensions alter through thermal expansion and thermal contraction.

The Mechanism of Expansion and Contraction

• Heating (Thermal Expansion): Heating increases the kinetic energy of particles. Particles vibrate or move more rapidly, colliding with greater force and pushing slightly further apart. This increase in inter-particle spacing causes the entire volume of the material to expand.
• Cooling (Thermal Contraction): Cooling decreases particle kinetic energy. Particles slow down, collide less forcefully, and move closer together, causing the volume to contract.

Real-World Engineering Applications in the Caribbean

1. Expansion Joints on Bridges and Roadways: Metal expansion joints with interlocking teeth are built into concrete bridges to prevent cracking or buckling as the bridge expands during hot tropical days and contracts during cooler nights.
2. Overhead Electrical Power Lines: Copper and aluminium power cables are strung loosely with deliberate sag. During hot weather, thermal expansion causes them to expand and sag further; if strung too tightly in cold weather, thermal contraction could snap them.
3. Liquid-in-Glass Thermometers: Clinical and weather thermometers utilise the thermal expansion of coloured alcohol or mercury inside a narrow capillary tube to measure temperature accurately.
4. Bimetallic Strips in Appliances: Thermostats in iron boxes, refrigerators, and electric kettles use two bonded metals (e.g., brass and iron) with different expansion rates. Heating bends the strip to break an electrical contact, providing automated temperature control.

Mixtures, Pure Substances & Separation Techniques

Definition: Matter can be classified based on chemical composition into pure substances (elements and compounds) and mixtures (homogeneous and heterogeneous).

                                    MATTER
                                      |
             +------------------------+------------------------+
             |                                                 |
       PURE SUBSTANCES                                     MIXTURES
   (Fixed composition & properties)                 (Variable physical combination)
             |
     +-------+-------+                                 +-------+-------+
     |               |                                 |               |
  ELEMENTS       COMPOUNDS                        HOMOGENEOUS     HETEROGENEOUS
(Single atom)  (Chemically bound)                (Uniform mixture)(Visible phases)
 e.g. Iron,     e.g. Water (H₂O),                 e.g. Saltwater,  e.g. Sand+water,
 Oxygen         Sodium Chloride                   Air, Alloys      Oil+water

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Pure Substances vs. Mixtures

Pure Substances

A pure substance consists of only one type of matter with uniform chemical composition throughout. It exhibits constant, sharp physical properties (such as exact melting and boiling points).

• Elements: Fundamental substances composed of only one type of atom that cannot be broken down into simpler substances by chemical means (e.g., Gold [Au], Oxygen [O₂], Carbon [C], Copper [Cu]).
• Compounds: Substances formed when two or more different elements are chemically combined in fixed mass ratios (e.g., Pure Water [H₂O], Carbon Dioxide [CO₂], Table Salt [NaCl]). Compounds can only be separated into constituent elements via chemical reactions.

Mixtures

A mixture consists of two or more substances physically mixed together in any variable proportion without chemical bonding. Each component retains its individual chemical identity and physical properties.

• Homogeneous Mixtures (Solutions): Mixtures with a uniform composition throughout; individual components are invisible to the naked eye (e.g., dissolved sugar in water, clean air, brass alloy).
• Heterogeneous Mixtures: Mixtures with a non-uniform composition featuring visibly distinct phases or components (e.g., oil and water, muddy water, sea sand mixed with seashells).

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Laboratory Separation Techniques for Mixtures

Because components of physical mixtures retain their distinct physical properties (particle size, solubility, boiling point, density, magnetic behaviour), physical separation methods can be employed.

1. Sieving

• Principle: Separates solid particles based on differences in particle size using a mesh or perforated barrier.
• Process: Smaller particles pass through the sieve mesh, while larger solid particles are retained on top.
• Applications: Separating fine sand from coarse gravel in construction; sieving flour in Caribbean bakeries to remove clumps.

2. Filtration

• Principle: Separates an insoluble solid from a liquid using a porous filter medium (e.g., filter paper or fine cloth).
• Process: The mixture is poured into a funnel lined with filter paper. The liquid passes through the microscopic pores as the filtrate, while the solid particles are trapped on the paper as the residue.
• Applications: Filtering ground coffee beans from brewed coffee; removing suspended silt from river water.

3. Evaporation

• Principle: Separates a dissolved soluble solid from a liquid solvent by exploiting differences in volatility/boiling points.
• Process: The solution is heated in an evaporating dish until the liquid solvent boils off completely as gas, leaving behind dry crystals of the solid solute.
• Applications: Extracting sea salt from ocean brine in solar salt pans (e.g., Great Inagua in the Bahamas).

4. Magnetic Separation

• Principle: Separates magnetic substances (iron, steel, nickel, cobalt) from non-magnetic materials.
• Process: A magnet is moved over the mixture, attracting magnetic particles while leaving non-magnetic material behind.
• Applications: Separating iron filings from sulfur powder; recovering scrap steel in recycling facilities.

5. Simple Distillation

• Principle: Separates a liquid solvent from a dissolved solid (or two miscible liquids with widely different boiling points) when both components must be recovered.
• Process: The solution is boiled in a flask. Gaseous vapour rises, enters a water-cooled Liebig condenser, cools down, and condenses back into pure liquid called the distillate.
• Applications: Desalination of seawater into drinkable fresh water; production of distilled water for laboratories.

6. Paper Chromatography

• Principle: Separates small quantities of dissolved pigments or dyes based on their differing solubilities in a moving solvent.
• Process: A spot of mixture (e.g., ink) is applied to absorbent chromatography paper dipped in solvent. Highly soluble components travel faster up the paper, producing distinct coloured bands.