Atoms, Elements, and the Periodic Table

Introduction to Chemical Symbols and the Periodic Table

The study of biology requires a foundational understanding of chemistry, particularly the structure and properties of atoms and molecules. The periodic table organizes all known elements, each represented by a unique chemical symbol.

• Element: A pure substance consisting of only one type of atom, identified by its chemical symbol (e.g., H for hydrogen, O for oxygen).

• Periodic Table: A tabular arrangement of elements by increasing atomic number, grouping elements with similar chemical properties.

• There are 92 naturally occurring elements; a few are essential for life (C, H, O, N make up ~96% of living mass).

Atomic Structure and Isotopes

Atomic Models and Isotopes

Atoms are composed of subatomic particles: protons, neutrons, and electrons. The number of protons defines the element, while the number of neutrons can vary, resulting in isotopes.

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

• Atomic Mass (A): Sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons (e.g., Carbon-12 and Carbon-14).

• Neutral Atom: Number of protons equals number of electrons.

Example: Carbon-12 (12C) has 6 protons and 6 neutrons; Carbon-14 (14C) has 6 protons and 8 neutrons. Both have 6 electrons if neutral.

Electron Orbitals and Valence Shells

Electron Configuration and Chemical Reactivity

Electrons occupy specific energy levels or orbitals around the nucleus. The outermost shell, or valence shell, determines an atom's chemical reactivity.

• Electron Orbitals: Regions around the nucleus where electrons are likely to be found.

• Valence Shell: The outermost electron shell; atoms are most stable when this shell is full.

• Atoms with incomplete valence shells are more chemically reactive, seeking to gain, lose, or share electrons to achieve stability.

Molecules and Compounds

Definitions and Classification

Molecules and compounds are formed by the combination of atoms through chemical bonds.

• Molecule: Two or more atoms held together by covalent bonds (e.g., O2, H2O).

• Compound: A substance composed of two or more different elements in a fixed ratio (e.g., NaCl, H2O).

Example: O2 is a molecule but not a compound; NaCl and H2O are both molecules and compounds.

Moles, Atomic Mass Units, and Avogadro's Number

Quantifying Atoms and Molecules

Because atoms and molecules are extremely small, chemists use the mole as a counting unit to relate mass to number of particles.

• Atomic Mass Unit (amu): A standard unit of mass for atoms and molecules; 1 amu ≈ mass of 1 proton or 1 neutron.

• Mole (mol): The amount of substance containing as many entities (atoms, molecules) as there are atoms in 12 g of carbon-12.

• Avogadro's Number: entities per mole.

• Moles allow conversion between atomic/molecular scale and laboratory scale (grams).

Example: 1 mole of water (H2O) has a mass of 18 g and contains molecules.

Chemical Bonds

Covalent, Ionic, and Hydrogen Bonds

Atoms form chemical bonds to achieve stable electron configurations. The main types of bonds are covalent, ionic, and hydrogen bonds.

• Covalent Bond: Atoms share pairs of electrons. Can be single, double, or triple bonds.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., O2).

• Ions: Atoms or molecules with a net electric charge due to loss or gain of electrons.

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

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

• Ionic Bond: Electrostatic attraction between oppositely charged ions (e.g., Na+ and Cl- in NaCl).

• Hydrogen Bond: Weak attraction between a hydrogen atom (partially positive) and an electronegative atom (e.g., O or N) in another molecule.

Example: In water, hydrogen bonds form between the hydrogen of one molecule and the oxygen of another, contributing to water's unique properties.

Chemical Equations and Reactions

Reactants, Products, and Equilibrium

Chemical reactions involve the transformation of reactants into products, often represented by chemical equations.

• Reactant: Substance present at the start of a reaction.

• Product: Substance formed as a result of a reaction.

• Equilibrium: State in which the forward and reverse reactions occur at equal rates.

Example Equation:

Double arrows () indicate equilibrium.

Redox (Oxidation-Reduction) Reactions

Electron Transfer and Energy Movement

Redox reactions involve the transfer of electrons between atoms, resulting in changes in oxidation states and often the movement of energy.

• Oxidation: Loss of electrons by an atom, molecule, or ion.

• Reduction: Gain of electrons by an atom, molecule, or ion.

• Oxidation and reduction always occur together (redox pair).

Example: Rusting of Iron

• Iron is oxidized (loses electrons), oxygen is reduced (gains electrons).

Example: Formation of NaCl

• Sodium (Na) is oxidized (loses an electron), chlorine (Cl) is reduced (gains an electron).

In biological systems, redox reactions are essential for processes such as cellular respiration and photosynthesis, where energy is transferred via electron movement.

Summary Table: Types of Chemical Bonds

Key Terms and Definitions

• Element: Substance that cannot be broken down by chemical means.

• Atom: Smallest unit of an element retaining its properties.

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

• Molecule: Two or more atoms covalently bonded.

• Compound: Substance of two or more elements in fixed ratio.

• Mole: 6.02 × 1023 entities (Avogadro's number).

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the summary tables and explicit definitions of key terms.