Elements & Compounds

Matter can be classified into two broad categories: pure substances and mixtures. A pure substance has a fixed, definite chemical composition and uniform properties throughout any sample — it cannot be separated into simpler components by physical means such as filtration or distillation. Pure substances are further divided into two sub-categories: elements and compounds, distinguished by whether their particles consist of one type of atom or more than one type joined by chemical bonds. An element is a pure substance made entirely of one kind of atom — atoms that all share the same atomic number (number of protons). There are 118 confirmed elements, each listed on the periodic table with a unique chemical symbol (e.g., oxygen, O; iron, Fe; gold, Au). Elements cannot be broken down into simpler substances by ordinary chemical reactions; they represent the most fundamental chemical building blocks of matter. Some elements exist as diatomic molecules (H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂) under standard conditions, yet they are still classified as elements because only one type of atom is present. A compound is a pure substance formed when two or more different elements are chemically bonded together in a fixed mass ratio. The properties of a compound are entirely different from those of its constituent elements — for example, water (H₂O) is a liquid at room temperature, yet its components hydrogen (H₂) and oxygen (O₂) are both colourless gases. Compounds can only be separated into their elements by chemical processes such as electrolysis or thermal decomposition, not by physical methods. This distinction — chemical bonding and fixed composition — is the defining feature that separates compounds from mixtures of elements.

Key takeaways
  • Defining an Element
  • Defining a Compound
  • The Role of Chemical Bonding
  • Fixed Composition and the Law of Definite Proportions
  • Decomposition of Compounds
  • Summary Classification Table

Defining an Element

An element contains only one type of atom, identified by its atomic number (Z). All atoms of a given element have the same number of protons. For example, every atom of carbon has Z = 6. Even if carbon forms a molecule like C₆₀ (buckminsterfullerene), it remains an element. The periodic table organises all 118 known elements by increasing atomic number and groups them by chemical behaviour (metals, non-metals, metalloids). Elements are represented by a one- or two-letter chemical symbol: the first letter is always capitalised and the second (if present) is always lowercase — e.g., Na (sodium), Ca (calcium), Cl (chlorine).

Defining a Compound

A compound forms when atoms of two or more different elements are joined by chemical bonds (ionic, covalent, or metallic) in a definite, fixed ratio by mass. This ratio is described by the compound's chemical formula. For instance, water is always H₂O — two hydrogen atoms bonded to one oxygen atom — giving a hydrogen-to-oxygen mass ratio of approximately 1:8. Sodium chloride (NaCl) is an ionic compound in which Na⁺ and Cl⁻ ions are present in a 1:1 ratio. The compound's physical and chemical properties (melting point, reactivity, colour, etc.) are entirely distinct from those of its component elements.

The Role of Chemical Bonding

The crucial distinction between a compound and a mixture of elements is chemical bonding. When sodium (a reactive silver metal) and chlorine (a toxic yellow-green gas) combine chemically, they form sodium chloride (common table salt — a white crystalline solid safe to consume). Energy is released or absorbed during bond formation, and the resulting compound has its own unique enthalpy of formation (ΔH°f). In contrast, physically mixing sodium and chlorine powders would still leave them as separate elements with their individual properties, classifying the result as a mixture, not a compound.

Fixed Composition and the Law of Definite Proportions

Every sample of a given compound always contains its constituent elements in the same proportion by mass, regardless of the sample's origin or size. This is Proust's Law of Definite Proportions (1799). For water (H₂O): molar mass = 2(1.008) + 16.00 = 18.016 g mol⁻¹. The mass fraction of hydrogen = (2 × 1.008) / 18.016 ≈ 0.1119 (11.19 %) and oxygen = 16.00 / 18.016 ≈ 0.8881 (88.81 %). Any pure water sample, from any source in the universe, must contain exactly these proportions. A mixture of hydrogen and oxygen gases does not obey this constraint — the ratio can be anything.

Decomposition of Compounds

Because a compound's elements are chemically bonded, physical separation methods (filtration, evaporation, distillation) cannot recover the individual elements. Chemical decomposition is required. Electrolysis of water: 2H₂O(l) → 2H₂(g) + O₂(g) uses electrical energy to break O–H covalent bonds. Thermal decomposition of calcium carbonate: CaCO₃(s) → CaO(s) + CO₂(g) (requires T > 840 °C) breaks ionic/covalent bonds by supplying thermal energy. These reactions confirm that the compound's composition is fixed and its separation is a chemical, not physical, change.

Summary Classification Table

Pure Substance → Element: one atomic number, cannot be chemically decomposed, e.g., Cu, S₈, O₂. Pure Substance → Compound: two or more elements in fixed mass ratio, joined by chemical bonds, separable only by chemical means, e.g., H₂O, NaCl, CO₂, C₁₂H₂₂O₁₁. Key tests: (1) Does the substance have a fixed melting/boiling point? (Yes → pure substance.) (2) Is it listed on the periodic table or composed of identical atoms? (Yes → element.) (3) Can it be broken into simpler substances by chemical reaction? (Yes → compound.) (No → element.)

The Key Equation

Mass fraction of element X=nX×MXMcompound\text{Mass fraction of element } X = \frac{n_X \times M_X}{M_{\text{compound}}}
  • n_X — Number of atoms of X per formula unit (dimensionless): The subscript of element X in the chemical formula, e.g., 2 for H in H₂O.
  • M_X — Molar mass of element X (g mol⁻¹): The mass of one mole of atoms of element X, taken from the periodic table.
  • M_{\text{compound}} — Molar mass of the compound (g mol⁻¹): Sum of the molar masses of all atoms in one formula unit: M = Σ(nᵢ × Mᵢ).
  • \text{Mass fraction} — Mass fraction of element X (dimensionless (0 to 1)): Multiply by 100 to obtain percentage by mass. For a pure compound this value is constant regardless of sample size (Law of Definite Proportions).

Worked Example

Problem. Carbon dioxide (CO₂) is a compound of carbon and oxygen. (a) Identify it as an element or compound and justify your answer. (b) Calculate the percentage by mass of carbon in CO₂. (c) A student claims that a 22.0 g sample of CO₂ contains 8.0 g of carbon. Is this consistent with CO₂ being a pure compound? (Relative atomic masses: C = 12.01, O = 16.00)

%C in CO2=12.0144.01×100≈27.29%\% \text{C in CO}_2 = \frac{12.01}{44.01} \times 100 \approx 27.29\%

Answer. CO₂ is a compound (two elements chemically bonded). Carbon makes up 27.29% of CO₂ by mass. A 22.0 g sample should contain 6.00 g of carbon, not 8.0 g — so the student's claim is inconsistent with CO₂ being a pure compound.

Figure
Diagram
Classification of Pure SubstancesPURE SUBSTANCEFixed composition; uniform propertiesELEMENTOne type of atom (one Z)COMPOUND2+ elements; chemical bonds; fixed ratioExamples of ElementsCuCopper (Z=29)O₂Oxygen (Z=8)AuGold (Z=79)Examples of CompoundsH₂OWater (H & O bonded)NaClSalt (Na⁺ & Cl⁻ ionic)CO₂Carbon dioxide (C & O)Elements: only physical change possible. Compounds: require chemical change to separate.
Origins · 1 minJoseph Louis Proust and how this was discovered
JL
Joseph Louis Proust
1754–1826 · French

Proust formulated the Law of Definite Proportions (also called Proust's Law) in 1799, demonstrating through careful quantitative analysis that every compound always contains its constituent elements in the same fixed ratio by mass regardless of the sample's source or preparation method. Working in Madrid, he analysed copper carbonate produced by different routes and consistently obtained the same mass ratios, directly proving that compounds have a fixed composition — a cornerstone concept distinguishing compounds from mixtures. His work provided crucial empirical support for Dalton's subsequent atomic theory. Proust conducted many of his landmark experiments in Spain while working at the Royal Artillery College in Segovia and later in Madrid. He famously engaged in a prolonged scientific debate (1799–1807) with fellow French chemist Claude Louis Berthollet, who argued that compounds could have variable composition. Proust ultimately prevailed: careful experimentation showed that apparent variations were due to impurities, not genuine variability in compound composition.

  • Particulate Nature of Matter — Understanding that all matter is composed of discrete particles (atoms, molecules, ions) provides the atomic-level foundation for defining elements (one atom type) and compounds (chemically bonded different atoms).
  • Mixtures: Homogeneous vs Heterogeneous — The classification of matter distinguishes pure substances (elements and compounds) from mixtures. Contrasting compounds with mixtures highlights that fixed composition and chemical bonding define a compound, while a mixture has variable composition and retains the properties of its components.
  • Law of Definite Proportions — Proust's law is the quantitative expression of fixed composition in compounds. It explains why every pure sample of a compound contains elements in the same mass ratio — a direct consequence of the fact that chemical bonds join atoms in specific integer ratios.
  • Physical vs Chemical Changes — Compounds can only be separated into their elements by chemical changes (e.g., electrolysis, thermal decomposition), whereas mixtures are separated by physical changes. Distinguishing these two types of change is essential for understanding the difference between compounds and mixtures.
  • Law of Conservation of Mass — When elements combine to form a compound or a compound is decomposed, the total mass is conserved. This law underpins the quantitative relationships in chemical formulas and reinforces why compounds have fixed mass ratios.
  • Separation Techniques — Physical separation techniques (distillation, filtration, chromatography) work for mixtures but cannot separate the elements of a compound. Recognising this limitation reinforces the defining role of chemical bonding in compounds.
  • States of Matter: Solid, Liquid, Gas — Elements and compounds each have characteristic melting and boiling points — fixed values that confirm their purity. State behaviour at the particle level connects directly to how atoms and molecules of elements and compounds are arranged and interact.
  • Chromatography (Detailed) — Chromatography can confirm whether a substance is a pure compound (single spot/peak) or a mixture (multiple spots/peaks), providing a practical analytical technique to verify purity and therefore assist in the classification of matter.