Chapter 2: Matter and States of Matter (Vocabulary Flashcards)

Matter

• Matter is defined as any material that has mass and occupies space. It is made up of small particles such as atoms and molecules and includes both living and nonliving things (e.g., soil, plants, rocks, water, wood, clothing).
• Chemistry is the study of matter, focusing on the properties of different types of matter and how they behave, including interactions with energy.
• Matter can be categorized by type and by its physical and chemical properties; the study often involves understanding elements, compounds, and how substances react to form new ones.

Energy

• Energy is the part of the universe that has the ability to do work.
• Types mentioned include physical energy and kinetic energy; energy changes are involved in chemical changes.
• Nearly all changes that matter undergoes involve absorption or release of energy; energy exchanges accompany physical and chemical processes.

Physical and Chemical Properties

• Properties are used to identify chemical substances and describe matter, including physical properties and chemical properties.
• Physical properties are observable or measurable without changing the substance’s identity or composition (e.g., color, odor, state of matter: solid, liquid, gas).
• Density is a physical property:
  ho = rac{m}{V} where mass is involved and volume is the space occupied; this is a physical property because it does not change the substance’s identity.
• Melting point, boiling point, and freezing point are physical properties describing phase transitions of a substance.
• Viscosity is a physical property describing resistance to flow (e.g., molasses vs. water).
• Chemical properties describe how a substance interacts with other substances to form new substances (e.g., reactivity).
• Chemical changes involve transforming one or more substances into new substances with different properties.
• Example: hydrogen gas (H₂) and oxygen gas (O₂) react to form water (H₂O). The resulting substance has different properties (density, phase tendencies, etc.).
• Evidence of chemical change includes: color change, temperature change, odor, production or emission of light, formation of a gas, or formation of a new solid (precipitate).
• Example: rusting of iron is an oxidation reaction where iron forms iron oxide in the presence of oxygen and moisture; the resulting substance has a different color, composition, and properties.
• Important caveat: in a chemical change, the original substances are converted to new substances; in a physical change, the identity of the substance remains the same, though its form may change (e.g., melting, freezing, bending).

Classification of Matter

• Matter can be classified as a pure substance or a mixture.
• Pure substances consist of only one type of substance and have the same chemical and physical properties throughout.
• Elements are pure substances and are the building blocks of everything; they cannot be broken down into simpler substances by ordinary chemical means. Elements are made up of atoms (e.g., hydrogen,
  nitrogen, oxygen).
• Compounds are pure substances formed by chemically combining two or more elements; they have a consistent composition throughout (e.g.,
  water, NaCl).
• A compound can be separated into simpler substances only by chemical means (e.g., electrolysis of water to separate H₂ and O₂).
• Mixtures consist of two or more substances physically combined and can have variable composition.
• Mixtures can be homogeneous (uniform composition, visually indistinguishable components) or heterogeneous (nonuniform composition, components visibly distinct).
• Examples:
  • Homogeneous: air (mixture of nitrogen, oxygen, water vapor, CO₂, etc.), seawater (salts dissolved in water), etc.
  • Heterogeneous: trail mix, fruit salad, salsa (visible components like tomatoes, candies, fruits).
• In mixtures, components can often be separated by physical methods (filtration, evaporation, chromatography). In compounds, separation requires chemical processes.
• When separated, components retain their individual properties.
• Homework example (chapter 2): classify these substances — water (H₂O), air, helium (He), trail mix, carbon dioxide (CO₂):
  • water: compound (pure substance) with formula ext{H}_2 ext{O}
  • air: homogeneous mixture
  • helium: element (atom: He)
  • trail mix: heterogeneous mixture
  • carbon dioxide: compound (pure substance) with formula ext{CO}_2

States of Matter and Phase Concepts

• Matter exists in three main states: solid, liquid, and gas (often liquid and gas are called with their common terms and states; vapor is the gaseous form of a substance that is typically a liquid at room temperature).
• States of matter are physical descriptions of a substance and relate to how molecules are arranged and how they behave.
• Water is a classic example: it can be a solid (ice), a liquid (water), or a gas (water vapor) under appropriate conditions.
• Solids:
  • Definite shape and definite volume.
  • Particles are tightly packed and usually arranged in fixed positions (crystalline lattices) or in fixed but disordered arrangements (amorphous).
  • Crystalline solids have orderly, repeating lattice structures (example: table salt, NaCl).
  • Amorphous solids lack a long-range order (example: wax, glass-like materials, or napthalene in some forms).
  • Typical solids are rigid and may be hard or brittle; they require thermal energy to overcome intermolecular interactions to melt.
  • Examples: table salt (NaCl) is crystalline; wax is amorphous. Metals like copper have distinctive properties such as color (reddish), luster, and high thermal and electrical conductivity.
• Liquids:
  • Definite volume but no definite shape; they take the shape of their container.
  • Molecules are less tightly packed than in a solid and can flow; they remain in contact with each other but are not in fixed positions.
  • Volume remains constant when in a container of fixed size; shape adapts to the container.
  • Example: wine in a bottle remains 750 mL whether in the bottle or decanter; its shape changes with the container.
• Gases:
  • No definite shape or volume; they expand to fill the available space.
  • Molecules are far apart and move freely, with minimal interactions among them.
  • Gases adopt the shape and volume of their container (e.g., room, classroom, etc.).
• Visual comparisons (molecular level): solids have tightly packed particles; liquids have more space and can flow; gases have particles far apart and move rapidly.

Phase Changes (Physical Changes)

• A phase change is a physical change that alters appearance but not identity or composition.
• Common phase changes:
  • Melting: solid to liquid.
  • Freezing: liquid to solid.
  • Boiling/Evaporation: liquid to gas.
  • Condensation: gas to liquid.
  • Sublimation: solid to gas (direct, skipping liquid).
  • Deposition: gas to solid (direct).
• Sublimation example: dry ice is solid CO₂; it can sublimate directly to CO₂ gas at room temperature, appearing to disappear without a liquid phase.
• Deposition example: formation of frost (gas to solid) occurs when water vapor deposits as ice directly.
• Water illustrates phase changes—melting, boiling, and condensation—though sublimation is less common for water under everyday conditions.

Solids, Liquids, and Gases: Detailed Comparisons

• Solids
  • Definite shape and volume; particles in fixed positions; strong intermolecular interactions; little to no translational movement (mostly vibrational).
  • Crystalline solids: ordered, repeating lattice structure (e.g., table salt NaCl).
  • Amorphous solids: disordered arrangement with irregular packing (e.g., wax, many polymers).
  • Examples of properties: rigid structure, density, brittleness, melting behavior.
• Liquids
  • Definite volume but no definite shape; able to flow and adapt to container shape; particles less tightly packed than in a solid but still in contact.
  • Molecules have enough energy to overcome some intermolecular forces, enabling flow.
  • Examples of properties: viscosity, surface tension (not detailed here but related to flow and packing).
• Gases
  • No definite shape or volume; expand to fill container; molecules far apart and move rapidly, with minimal intermolecular interactions.
  • Gases can be compressed or expanded, showing high compressibility and fluidity.

Real-World Connections and Implications

• Copper is a good conductor of heat and electricity, which explains the common use of copper wiring and piping in homes.
• Sodium (Na) and chlorine (Cl) are both highly reactive on their own; sodium metal reacts violently with water and chlorine gas is highly toxic; yet, when combined chemically to form table salt (NaCl), the substance is essential and widely used in food, illustrating how chemical changes can radically alter properties.
• Phase changes and state behavior are fundamental to technology and everyday life (cooling, heating, distillation, separation processes).
• Chromatography and filtration are practical methods for separating mixtures; chromatography is a physical separation technique applicable to many fields (e.g., lab analysis, drug testing).

Summary of Key Concepts and Formulas

• Matter classification: pure substances (elements or compounds) vs mixtures (homogeneous vs heterogeneous).
• Elements: building blocks; cannot be broken down by chemical means; atoms are the basic unit.
• Compounds: chemically bound combinations of two or more elements; consistent composition; can be broken into elements only by chemical processes (e.g., electrolysis of water).
• Mixtures: physically combined; components retain their properties; separable by physical means.
• Physical properties: observable without changing identity (e.g., color, odor, state, density, melting/boiling points, viscosity).
• Chemical properties: describe how a substance reacts to form new substances.
• Evidence of chemical change: color change, temperature change, odor, light, gas formation, precipitate formation.
• Phase changes:
  • ext{Solid}
    ightarrow ext{Liquid}
    ightarrow ext{Gas} (melting, boiling).
  • Sublimation: ext{Solid}
    ightarrow ext{Gas}.
  • Deposition: ext{Gas}
    ightarrow ext{Solid}.
• States of matter differ in molecular arrangement and energy; solids are tightly packed, liquids have some mobility, gases move freely and fill space.
• Density:
  ho = rac{m}{V}.
• Common chemical formulas mentioned:
  • Water: ext{H}_2 ext{O}
  • Sodium chloride: ext{NaCl}
  • Hydrogen gas: ext{H}_2
  • Oxygen gas: ext{O}_2
  • Carbon dioxide (dry ice context): ext{CO}_2
• Notable example reactions:
  • Formation of water: ext{2H}2 + ext{O}2
    ightarrow ext{2H}_2 ext{O}
  • Electrolysis of water: ext{2H}2 ext{O} ightarrow ext{2H}2 + ext{O}_2
• Sample volumes: liquids can be measured by volume; a typical example uses 750\,\mathrm{mL} of liquid in a bottle or decanter.

Quick Reference for Homework Questions (Chapter 2 Part 1)

• Water (H₂O): Compound (Pure substance) – formula ext{H}_2 ext{O}
• Air: Homogeneous mixture
• Helium (He): Element
• Trail mix: Heterogeneous mixture
• Carbon dioxide (CO₂): Compound (Pure substance)

Note on Phase Change Examples Mentioned

• Sublimation example: dry ice (solid CO₂) sublimates to CO₂ gas at room temperature.
• Deposition example: frost formation (gas to solid) under suitable conditions.
• Water can exhibit solid, liquid, and gaseous states under everyday conditions; sublimation is less common for water under normal conditions but is discussed as a concept.

Next: The lecture continues with part two, which will delve deeper into energy calculations and further aspects of matter and energy interactions.