Properties, Changes, and Classifications of Matter

Fundamentals and States of Matter

• Matter: Anything that has mass and takes up space.

• Law of Conservation of Matter: Matter can neither be created nor destroyed; it can only change forms.

• Substance (Pure Substance): Matter that has a uniform and definite composition and identical properties throughout.

• States of Matter: Matter commonly exists in three primary physical states, distinguished by their microscopic and macroscopic properties:

  • Solid: Particles are held tightly in a fixed pattern, vibrating in place. Solids possess a definite shape and a definite volume, and are virtually incompressible.

  • Liquid: Particles are held close together but are arranged randomly and are free to move past one another. Liquids possess a definite volume but a variable (indefinite) shape, taking the shape of their container. They are virtually incompressible.

  • Gas: Particles are very far apart, moving very fast and in random directions. Gases possess an indefinite shape and an indefinite volume, expanding to completely fill whatever container holds them.

  • Vapor: The gaseous state of a substance that is normally a solid or liquid at room temperature.

  • Plasma: A high-energy state of matter that results when gas particles become ionized (charged) and are broken apart into smaller charged particles; naturally occurring in stars.

• Physical Properties of Matter: Characteristics that can be observed or measured without changing the sample's chemical composition.

  • Examples include: color, odor, luster, ductility, and density.

  • Extensive Properties: Physical properties that depend on the amount of substance present (e.g., mass, volume, length).

  • Intensive Properties: Physical properties that do not depend on the amount of substance present; these can be used to identify an unknown substance (e.g., density, boiling point, melting point, color, luster, ductility).

• Chemical Properties of Matter: The ability or inability of a substance to combine with or change into one or more new substances (reactivity or lack of reactivity).

  • The Periodic Table is organized according to families of elements that display similar chemical reactions.

  • Chemical properties cannot always be determined visually; experimentation is required to observe whether a substance reacts with another (e.g., water does not react with glass).

  • Chemical properties of a substance can change when the substance undergoes a change in state.

Physical and Chemical Changes

• Physical Changes: Processes that cause no change in the fundamental composition or chemical identity of the substance.

  • Examples of physical alteration: tearing, chopping, bending, crushing.

  • Phase Changes: Transition from one state of matter to another (e.g., melting, boiling, freezing, condensing) are physical changes.

  • Dissolving: Dissolving a substance in water is classified as a physical change because when the water is evaporated or removed, the original substance re-forms.

• Chemical Changes: A process where one or more substances are transformed into entirely new substances with different chemical compositions (also referred to as a chemical reaction).

  • The resulting products possess distinct physical and chemical properties compared to the original reactants.

  • Chemical Reaction Representation: $\text{Reactants} \rightarrow \text{Products}$

  • Key terminology signaling chemical change: decompose, explode, rust, oxidize, corrode, tarnish, ferment, burn, rot.

• Law of Conservation of Mass: Mass is conserved during any chemical reaction or physical process. The total mass involved remains constant.

  • $\text{mass of reactants} = \text{mass of products}$

• Check-Up: Chemical vs. Physical Properties:

  • Sugar dissolving in water: Physical Property

  • Platinum not reacting with oxygen at room temperature: Chemical Property

  • The boiling point of a certain alcohol being $78\,^{\circ}\text{C}$: Physical Property

  • Sugar fermenting to form alcohol: Chemical Property

• Check-Up: Chemical vs. Physical Changes:

  • Ripping a piece of paper: Physical Change

  • Combining two clear substances to form a white chunky solid: Chemical Change

  • Dissolving a blue crystal in water to form a green solution: Physical Change

  • Dissolving a white solid in water causing the beaker to get very hot: Chemical Change

Elements, Compounds, and Percent Composition

• Pure Substances: Matter with constant composition and properties throughout, categorized into elements or compounds.

  • Elements: Pure substances that cannot be broken down or separated into simpler substances by physical or chemical means.

  • There are $90$ naturally occurring elements; the remaining elements are synthesized in laboratory settings.

  • There are $81$ stable elements; the remaining elements are radioactive.

  • Elements exist in various physical states under standard temperature and pressure.

  • Elements are organized into the Periodic Table of Elements based on periodic trends in physical and chemical properties.

  • Out of over $100$ elements, $10$ account for approximately $90\%$ of all matter on Earth.

  • Earth's crust is composed primarily of five elements: Oxygen ($\text{O} = 49.5\%$), Silicon ($\text{Si} = 25.7\%$), Aluminum ($\text{Al} = 7.5\%$), Iron ($\text{Fe} = 4.7\%$), and Calcium ($\text{Ca} = 3.4\%$).

  • The human body is composed primarily of Oxygen ($\text{O} = 65\%$).

  • Compounds: Pure substances formed by two or more different elements chemically combined in a fixed ratio.

  • There are over $50\text{ million}$ known compounds, with approximately $100{,}000$ new compounds developed or discovered each year.

  • Chemical formulas use periodic symbols to depict compound composition:

    • Table Salt: Contains $1$ part sodium and $1$ part chlorine ($\text{NaCl}$).

    • Water: Contains $2$ parts hydrogen and $1$ part oxygen ($\text{H}_2\text{O}$).

    • Separating water ($\text{H}_2\text{O}$) yields its component elements: hydrogen gas and oxygen gas.

  • The chemical and physical properties of a compound differ completely from those of its constituent elements.

  • Compounds exhibit constant mass proportions (Law of Definite Proportions): Table salt ($\text{NaCl}$) always consists of $39.3\%$ sodium by mass and $60.7\%$ chlorine by mass, regardless of sample size.

• Check-Up: Element or Compound Classification:

  • Aluminum foil: Element

  • Carbon dioxide gas: Compound

  • Rust: Compound

  • Helium gas: Element

• Percent Composition by Mass:

  • Formula: $\%\,\text{by Mass} = \frac{\text{mass of element}}{\text{mass of compound}} \times 100$

  • Worked Example 1: In a $50.0\,g$ sample of methane ($\text{CH}_4$), $37.4\,g$ of carbon are present.

  • Percentage of Carbon: $\%\,\text{Carbon} = \frac{37.4\,g}{50.0\,g} \times 100 = 74.8\%$

  • Percentage of Hydrogen: $\text{Mass of Hydrogen} = 50.0\,g - 37.4\,g = 12.6\,g$

  • $\%\,\text{Hydrogen} = \frac{12.6\,g}{50.0\,g} \times 100 = 25.2\%$

  • Worked Example 2: $2.00\,g$ of hydrogen is combined with $16.0\,g$ of oxygen to produce water.

  • Total Mass of Water Compound: $2.00\,g + 16.0\,g = 18.00\,g$

  • Percentage of Hydrogen: $\%\,\text{Hydrogen} = \frac{2.00\,g}{18.00\,g} \times 100 = 11.1\%$

  • Percentage of Oxygen: $\%\,\text{Oxygen} = \frac{16.0\,g}{18.00\,g} \times 100 = 88.9\%$

  • Definite Composition Rule: If two samples do not have identical percent compositions, they cannot be the same compound, even if composed of the same constituent elements.

Mixtures and Methods of Separation

• Mixture: A physical combination of two or more pure substances in which each pure substance retains its individual chemical properties. Mixtures can consist of elements combined with elements, elements combined with compounds, or compounds combined with compounds.

• Heterogeneous Mixture: A mixture that does not blend smoothly throughout, in which individual constituent substances remain visually or physically distinct. Samples taken from different locations in the mixture yield different properties due to uneven particle distribution.

• Homogeneous Mixture (Solution): A mixture that exhibits a completely constant composition throughout and always exists in a single phase.

  • Examples of homogeneous mixtures (solutions): Air in a classroom, saline water, metal alloys, pool water, tap water, Windex, paint.

• Check-Up: Pure Substance vs. Mixture & Detailed Classification:

  • Apple Juice: Mixture | Homogeneous

  • Carbon Dioxide ($\text{CO}_2$): Pure Substance | Compound

  • Helium: Pure Substance | Element

  • Oxygen ($\text{O}_2$): Pure Substance | Element

  • Brass: Mixture | Homogeneous (Alloy)

  • Concrete: Mixture | Heterogeneous

  • Pond Water: Mixture | Heterogeneous

  • Methane ($\text{CH}_4$): Pure Substance | Compound

• Submicroscopic Representation Classifications:

  • Diagram A: Compound

  • Diagram B: Element

  • Diagram C: Mixture of Compounds

  • Diagram D: Mixture of Elements

• Physical Methods for Separating Mixtures:

  • Filtration: Separation based on particle size. Uses a porous barrier (such as filter paper in a filter funnel) to separate an insoluble solid from a liquid or to separate different sizes of solids. (Example: Separating sand from water; sand remains as residue, water passes through as filtrate).

  • Distillation: Separation based on differences in boiling points. A liquid mixture is heated until the substance with the lowest boiling point vaporizes into steam, passes through a cooling condenser, and is collected in pure liquid form. (Example: Purifying water by separating pure liquid water from dissolved solid salts).

  • Chromatography: Separation based on the relative affinity (attraction) of components for two distinct phases. Components dissolved in a liquid or gas (mobile phase) travel across a fixed solid substrate (stationary phase) at different speeds based on solubility and attraction. (Example: Separating ink dyes; separating plant pigments like carotenoids, xanthophylls, chlorophyll a, chlorophyll b, and anthocyanin from leaf extracts using propanone solvent).

  • Crystallization: Separation based on solubility and crystal formation. A separation technique that results in the formation of solid particles of a substance from a solution containing the dissolved substance upon solvent evaporation. (Example: Formation of rock candy sugar crystals as water evaporates).

Iron Oxidation Case Study and Comprehensive Module Review

• Iron Oxidation Case Study:

  • Mass Comparison: The mass of rusted iron is greater than the mass of the initial iron powder because iron reacts chemically with oxygen gas in the air, combining to form rust ($\text{Fe}_2\text{O}_3$).

  • Change Classification: Oxidation of powdered iron is a chemical change because a chemical reaction converts iron and oxygen into a new oxidized compound.

• Comprehensive Review Statements:

  • Matter is anything that has mass and takes up space.

  • Compounds can be broken down into elements by chemical changes.

  • The three common physical states of matter are solid, liquid, and gas.

  • Solids and liquids are virtually incompressible; they have definite volumes.

  • Liquids have definite volumes but variable shapes.

  • Matter in the gaseous state has no shape and fills completely whatever container holds it.

  • Acetone is highly flammable because it reacts easily with oxygen gas in the atmosphere; this is an example of a chemical property of acetone.

  • An element cannot be broken down into simpler substances by chemical means.

  • A pure sample of an element contains only one kind of atom.

  • A solution that is bright green in color exhibits a physical property. When concentrated ammonia is added to it, turning it dark blue and forming a solid, a chemical change occurs.

  • Often, elements combine chemically together to form a compound.

  • The Law of Conservation of Mass states that mass is never created or destroyed in any physical or chemical change.

  • The properties of a compound are usually very different from those of the elements it contains.

  • A pure substance has a definite/constant composition.

  • A heterogeneous mixture can be defined as something that has variable composition.

  • A homogeneous mixture is more commonly referred to as a solution. If you had a salt water solution, you could collect the salt through the process of crystallization.

  • A mixture of salt and sand would be an example of a heterogeneous mixture. You can collect the sand through the process of filtration.