Atoms, Molecules, and Ions

Introduction to the Nature of Matter

Understanding the composition and structure of matter is fundamental to chemistry and cell biology. Atoms, molecules, and ions form the basis of all substances, including those found in living cells.

• Key Question: What is matter made of, and what holds it together?

• Example: Gold can be subdivided from ingots to nuggets to dust, but the smallest unit retaining its identity is the atom.

Dalton’s Atomic Theory

Foundational Concepts of Atomic Structure

Dalton’s Atomic Theory, developed in the early 19th century, provided the first modern explanation of the nature of matter.

• All matter is made up of atoms. Atoms are indivisible particles that cannot be created or destroyed in chemical reactions.

• Atoms of a given element are identical in mass and properties, but differ from atoms of other elements.

• Compounds are formed by the combination of atoms of different elements in fixed, small, whole-number ratios.

• Chemical reactions involve rearrangement of atoms; atoms themselves are not changed.

• Example: Water (H2O) always contains two hydrogen atoms and one oxygen atom.

Radioactivity

Types and Properties of Radiation

Radioactivity is the spontaneous emission of radiation by certain unstable atomic nuclei. It was first observed by Henri Becquerel in 1896.

• Alpha (α) particles: Helium nuclei; positively charged.

• Beta (β) particles: Electrons; negatively charged.

• Gamma (γ) rays: Electromagnetic radiation; no charge.

• Separation by charge: When passed through an electric field, α and β particles are deflected in opposite directions, while γ rays are unaffected.

Atomic Models

Historical Development of Atomic Structure

The understanding of atomic structure has evolved over time, with several key models proposed.

• Plum Pudding Model (J.J. Thomson, ~1900): Atoms consist of a positively charged sphere with negatively charged electrons embedded within it.

• Limitations: This model could not explain results from later experiments, such as Rutherford’s gold foil experiment.

• Example: The atom was visualized as a 'pudding' of positive charge with 'plums' of electrons scattered throughout.

Structure of the Atom

Subatomic Particles and Their Properties

Atoms are composed of three principal subatomic particles: protons, neutrons, and electrons.

• Proton (p+): Positively charged, located in the nucleus.

• Neutron (n0): Neutral, located in the nucleus.

• Electron (e-): Negatively charged, found in the space surrounding the nucleus.

• Relative masses: Proton ≈ Neutron >> Electron (electron mass is much smaller).

Symbols of Elements and Isotopes

Notation and Atomic Mass

Each element is represented by a unique symbol, and isotopes are distinguished by their mass numbers.

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 (A=12, Z=6), Carbon-13 (A=13, Z=6).

Average Atomic Mass and Isotopic Abundance

Calculating Atomic Mass

The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Formula:

• Example (Carbon):

• Carbon-12: 98.892% abundance, mass = 12.00000 amu

• Carbon-13: 1.108% abundance, mass = 13.00335 amu

• Calculation:

amu

The Periodic Table

Organization and Classification of Elements

The periodic table arranges elements by increasing atomic number and groups elements with similar properties into columns.

• Periods: Horizontal rows; elements have the same number of electron shells.

• Groups: Vertical columns; elements have similar chemical properties and the same number of valence electrons.

• Main group elements: Groups 1A–8A (representative elements).

• Transition elements: Groups 1B–8B.

• Inner transition elements: Lanthanides and actinides (bottom two rows).

• Metals: Good conductors, malleable, ductile, shiny, tend to lose electrons.

• Nonmetals: Poor conductors, brittle, various colors, tend to gain electrons.

• Metalloids: Exhibit properties intermediate between metals and nonmetals.

Molecules and Chemical Formulas

Types of Chemical Formulas

Molecules are groups of atoms bonded together. Chemical formulas represent the composition of molecules.

• Molecular formula: Shows the exact number of atoms of each element in a molecule (e.g., H2O).

• Structural formula: Shows how atoms are bonded (e.g., H–O–H for water).

• Empirical formula: Shows the simplest whole-number ratio of atoms (e.g., CH2O for glucose).

• Diatomic molecules: Seven elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

Ions and Ionic Compounds

Formation and Properties of Ions

Ions are charged particles formed by the gain or loss of electrons. Ionic compounds are composed of cations and anions held together by electrostatic forces.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Ionic bond: The electrostatic attraction between oppositely charged ions.

• Example: Sodium chloride (NaCl) consists of Na+ and Cl- ions arranged in a crystal lattice.

Properties of Ionic Compounds

• Formed by electron transfer between metals and nonmetals.

• Solid at room temperature with high melting points.

• Conduct electricity when dissolved in water (as ions are free to move).

• Have a regular geometric structure (crystal lattice).

Predicting and Naming Ions

Monatomic Ions and Their Charges

The charge of a monatomic ion can often be predicted from its position in the periodic table.

• Main group metals: Form cations with a charge equal to their group number (e.g., Na+, Mg2+).

• Main group nonmetals: Form anions with a charge equal to group number minus 8 (e.g., Cl-, O2-).

• Transition metals: May form more than one cation with different charges (e.g., Fe2+, Fe3+).

Naming Binary Compounds

• Binary ionic compounds: Name the cation first, then the anion (ending in -ide).

• Transition metals: Indicate the charge with Roman numerals (e.g., iron(II) chloride).

• Binary covalent compounds: Use prefixes to indicate the number of atoms (e.g., carbon dioxide, dinitrogen tetroxide).

Hydrates and Molecular Compounds

Hydrates

Hydrates are compounds that contain water molecules within their crystal structure.

• Example: CuSO4·5H2O is copper(II) sulfate pentahydrate.

• Anhydrous: A compound without water.

Acids, Anions, and Oxyanions

Acids and Their Anions

Acids are substances that release hydrogen ions (H+) in water. Their corresponding anions are named based on the acid and the number of oxygen atoms present.

• Oxyanions: Polyatomic ions containing oxygen (e.g., sulfate SO42-, nitrate NO3-).

• Naming: The suffix -ate is used for the most common oxyanion, -ite for one fewer oxygen.

• Prefixes: Per- (one more oxygen), hypo- (one fewer oxygen).

• Example Table:

• Hydrogen-containing anions: Prefix 'hydrogen' or 'bi-' is used (e.g., hydrogen carbonate or bicarbonate, HCO3-).

Additional info: These foundational chemical concepts are essential for understanding the molecular basis of cell structure and function in cell biology.