Structure of matter

A fundamental understanding of biological processes requires knowing how matter is constructed. Starting with atoms and their subatomic particles, this article will explore the bonding mechanisms that connect atoms to form molecules, from simple to highly complex structures.

Atoms with the same number of protons exhibit the same chemical behavior and are classified as an element. The periodic table of elements arranges these elements according to their atomic number (the number of protons and, in a neutral atom, electrons). An element's properties can be predicted from its position in the periodic table; elements in the same column form a group with similar chemical behaviors.

Substances, which are collections of molecules in a gaseous, liquid, or solid state, rarely exist as pure substances and are usually found as mixtures. Based on their chemical and physical properties, these mixtures can be separated using various methods, some of which are presented here.

For a long time, atoms were considered the smallest particles of matter (from the Greek "atomos," meaning indivisible). Today, we know that atoms are composed of even smaller, fundamental particles. Many properties of chemical substances can be explained by the structure of these atoms.

Building blocks of atoms

• Definition: Atoms are made up of smaller, subatomic particles. The three most relevant to health care professionals are discussed here.
• Types: protons, neutrons, and electrons
• Properties
  • Mass: Neutrons and protons each have a mass of approximately 1 u.
    • Unit: The atomic mass unit (u) is defined as 1/12 the mass of a single carbon-12 atom.
    • Relative atomic mass: Because the mass of an atom in "u" is defined relative to the mass of a carbon-12 atom, it is a relative, not an absolute, mass.
  • Charge: The electron has a negative charge of -1e, while a proton has a positive charge of +1e, which is equal in magnitude. Neutrons are uncharged.
    • Unit: e (elementary charge) is the smallest discrete unit of electrical charge observed in nature
  • Spin (physics)
    • Electrons: Electrons are in constant motion and also possess an intrinsic angular momentum known as spin.
    • Nucleus: An atomic nucleus can also have a net spin. If the total number of protons and neutrons (mass number) is odd, the nucleus has a magnetic dipole moment, causing it to behave like a tiny magnet. This phenomenon is the basis for magnetic resonance imaging (MRI).

Chemists approximate the relative atomic mass of an element by its mass number, which is the sum of its protons and neutrons.

Atoms

Before the discovery of subatomic particles, atoms were considered the fundamental components of matter. They consist of a nucleus and a surrounding shell, with the diameter of the shell being up to 150,000 times larger than that of the nucleus.

• Structure: Atoms consist of a positively charged nucleus and a negatively charged electron shell.
• Properties
  • Charge: In their neutral ground state, atoms have no net charge.
  • Shape: nearly spherical
• Classification: Atoms are systematically arranged in the periodic table of the elements (see below).
  • Element: The term "element" refers to atoms that have the same number of protons in the nucleus and thus the same chemical properties.
  • Nomenclature: Elements are designated by symbols, which are usually derived from their Latin or Greek names: e.g., Calcium = Ca; Oxygen (Lat. oxygenium) = O; Silver (Lat. argentum) = Ag.

Atomic nucleus

Atomic nuclei consist of protons and neutrons, which are collectively called nucleons. The number of protons defines the element, but the number of neutrons can vary. Atoms of the same element that have different numbers of neutrons are called isotopes.

• Structure
  • Number of protons: The identity of an atom or element is determined by its number of protons, which is called the atomic number (Z).
  • Number of neutrons: In lighter elements, the number of neutrons is often similar to the number of protons, but it can vary. In heavier stable elements, neutrons increasingly outnumber protons.
    • Isotopes: Forms of an element that have the same number of protons but different numbers of neutrons.
      • Nomenclature: To specify a particular isotope (or nuclide), the mass number (total number of protons and neutrons) is written as a superscript and the atomic number as a subscript, both preceding the element symbol.
        • e.g., ¹¹₆C, ¹²₆C, ¹⁴₆C, ¹⁶₈O
      • Stability: Isotopes of an element have different stabilities. For example, carbon has three natural isotopes: ¹²C and ¹³C are stable, while ¹⁴C is unstable.
        • Proton excess: A nucleus with too many protons relative to neutrons is often unstable. It may undergo radioactive decay to convert a proton into a neutron, such as through β⁺ (positron) emission.
        • Neutron excess: While many larger nuclei require an excess of neutrons for stability, too many neutrons can also be energetically unfavorable. Such an isotope may decay by converting a neutron into a proton via β⁻ decay.
      • Radioactive decay: Unstable nuclides decay by emitting radiation; they are therefore called "radioactive." For details, see: "Ionizing radiation".
      • Isotope mixtures: Most elements exist naturally as a mixture of several isotopes. The atomic mass listed on the periodic table is the weighted average of the masses of these naturally occurring isotopes.
        • Calculation: The average atomic mass of an element is calculated by summing the products of each isotope's mass and its natural abundance.
• Properties
  • Charge
    • The atomic nucleus is positively charged.
    • The number of protons determines its total charge and is therefore equal to the atomic number (Z).
    • In a neutral atom, the positive charge of the nucleus is balanced by the negative charge of the electron shell.
  • Mass
    • Although the nucleus is very small, it contains nearly all the mass of an atom.
    • Its relative mass is approximately equal to the number of nucleons, known as the mass number (A).

Nuclear forces and stability

• Strong nuclear force: the force that holds protons and neutrons together in the nucleus
  • It is the strongest of the four fundamental forces but acts only over very short distances (within the nucleus)
  • It overcomes the powerful electrostatic repulsion between positively charged protons, ensuring nuclear stability
• Mass defect (or mass deficit): the difference between the predicted mass of a nucleus (sum of its individual protons and neutrons) and its actual, measured mass
  • The actual mass of a nucleus is always less than the sum of the masses of its individual components.
  • This "missing" mass is converted into energy that binds the nucleons together (nuclear binding energy)
• Nuclear binding energy: the energy released when nucleons bind together to form a nucleus
  • It is equivalent to the energy required to break the nucleus apart into its constituent protons and neutrons
  • Calculated using Einstein's mass-energy equivalence principle: E = mc²
    • E = nuclear binding energy
    • m = mass defect
    • c = speed of light in a vacuum
  • A higher binding energy per nucleon corresponds to a more stable nucleus.
    • Iron (Fe) has one of the highest binding energies per nucleon, making it one of the most stable elements.

Radioactive decay

• Definition: the spontaneous process by which an unstable atomic nucleus loses energy by emitting radiation; this process, also known as nuclear transmutation, results in the formation of a different nuclide
• Half-life (t₁/₂): the time required for half of the radioactive nuclei in a sample to undergo decay; it is a constant value for each specific isotope
• Modes of decay
  • Alpha (α) decay: emission of an alpha particle (a helium nucleus, ⁴₂He), which consists of two protons and two neutrons
    • The mass number (A) decreases by 4.
    • The atomic number (Z) decreases by 2.
    • Example: ²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂He
  • Beta (β) decay: emission of a beta particle from the nucleus
    • Beta-minus (β⁻) decay: a neutron is converted into a proton, and an electron (β⁻ particle) is emitted
      • The mass number (A) remains unchanged.
      • The atomic number (Z) increases by 1.
      • Example: ¹⁴₆C → ¹⁴₇N + e⁻
    • Beta-plus (β⁺) decay (positron emission): a proton is converted into a neutron, and a positron (β⁺ particle, the antiparticle of an electron) is emitted
      • The mass number (A) remains unchanged.
      • The atomic number (Z) decreases by 1.
      • Example: ²²₁₁Na → ²²₁₀Ne + e⁺
  • Gamma (γ) decay: emission of high-energy photons (gamma rays) from an excited nucleus
    • Often occurs after α or β decay, as the daughter nucleus transitions from a high-energy state to a more stable, lower-energy state
    • Both the mass number (A) and the atomic number (Z) remain unchanged

Isotopes in medicine
Many radioactive isotopes are used in medicine for both therapeutic and diagnostic purposes. Radioactive iodine (¹³¹I), for example, is administered because it selectively accumulates in the thyroid gland. In radioiodine therapy, the emitted radiation generates free radicals that destroy the DNA of nearby cells, leading to their death. This makes it an effective treatment for conditions like thyroid cancer. Because this destructive effect also impacts healthy tissue, its diagnostic use in the radioiodine uptake test has become less common.

Electron shell

Electrons occupy the electron shell, where their collective negative charge balances the positive charge of the nucleus. Various models have been developed to describe the structure of the electron shell and the behavior of electrons within it.

• Structure: The shell is composed of electrons. In a neutral atom, the number of electrons equals the number of protons in the nucleus.
• Properties
  • Charge: The electron shell is negatively charged.
  • Mass: Although the shell accounts for most of an atom's volume, its mass is negligible compared to the nucleus.
• Models for electron distribution: Electrons are negatively charged particles that repel one another and can only occupy the shell under specific conditions described by quantum mechanics. These conditions are simplified in the following models.
  • Bohr model: This early model depicts electrons moving around the nucleus in fixed circular orbits called shells. Starting from the nucleus and moving outward, the shells are designated as follows:
    • K-shell (n=1): holds up to 2 electrons
    • L-shell (n=2): holds up to 8 electrons
    • M-shell (n=3): holds up to 18 electrons
    • N-shell (n=4): holds up to 32 electrons
  • Atomic orbital model: This model describes the probability of finding an electron in a three-dimensional region of space within a given shell.
    • Basic assumption: Electrons do not move in fixed orbits but occupy regions of high probability called atomic orbitals (AOs).
    • Atomic orbital types: Orbitals are distinguished by their shape, denoted by letters.
      • s-orbitals (spherical)
      • p-orbitals (dumbbell-shaped)
      • d-orbitals (more complex shapes, e.g., cloverleaf)
    • Orbital occupation
      • Electron configuration: Describes the distribution of electrons among the atomic orbitals.
        • Orbital notation: nℓ^x format (e.g., 1s²)
          • n = principal quantum number (shell), ℓ = orbital type (s, p, d, f), and x = number of electrons in that orbital
      • Rules
        • Aufbau principle: Electrons fill the lowest-energy orbitals first. The common filling sequence is 1s → 2s → 2p → 3s → 3p → 4s → 3d.
        • Hund's rule: When filling a subshell with multiple orbitals of the same energy (e.g., the three p-orbitals), each orbital gets one electron before any orbital gets a second.
        • Pauli exclusion principle: Once all orbitals in the subshell are half-filled, they begin to fill with a second electron of the opposite spin. Each atomic orbital can hold a maximum of two electrons.
    • Example: The electron configuration of manganese (Mn, atomic number 25) is 1s²2s²2p⁶3s²3p⁶4s²3d⁵.
    • Exceptions and ionic behavior
      • Neutral‑atom exceptions include chromium (Cr) and copper (Cu).
        • Chromium (Cr), atomic number 24
          • Expected: 1s²2s²2p⁶3s²3p⁶4s²3d⁴
          • Actual: 1s²2s²2p⁶3s²3p⁶4s¹3d⁵ (achieves a more stable, half-filled d-subshell)
        • Copper (Cu), atomic number 29
          • Expected: 1s²2s²2p⁶3s²3p⁶4s²3d⁹
          • Actual: 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰ (achieves a more stable, fully-filled d-subshell)
      • Ion formation (transition metals): In ions/compounds, s electrons are frequently removed before d electrons.
        • Example: Manganese (Mn), atomic number 25
          • Neutral atom (Mn): 1s²2s²2p⁶3s²3p⁶4s²3d⁵
          • Cation (Mn²⁺): 1s²2s²2p⁶3s²3p⁶3d⁵ (The two electrons from the 4s orbital are removed first)
• Valence electrons: The electrons in the outermost shell of an atom. These electrons are involved in forming chemical bonds with other atoms.
• Octet rule: Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with a completely filled outermost shell. For most elements relevant to medicine, this means having eight valence electrons. Hydrogen is an exception, achieving stability with two electrons.

Ions

When an atom or molecule gains or loses electrons, the number of electrons no longer equals the number of protons, resulting in a net-charged particle called an ion. The formation of ions requires an input of energy. This energy is called ionization energy for electron removal and electron affinity for electron gain.

• Definition: An ion is a charged atom or molecule.
• Classification: Ions are classified by their net charge.
  • Anions: ions with a negative charge
  • Cations: ions with a positive charge
• Formation
  • By gaining or losing electrons
  • By gaining or losing protons (H⁺)
  • Ionization energy: the energy required to remove one or more electrons from a neutral atom to form a cation
    • This energy increases with each successive electron removed from an atom.
    • The ionization energy decreases as the distance between the electron and the nucleus increases.
  • Electron affinity: The energy change that occurs when an electron is added to a neutral atom to form an anion. This process can either release or require energy.
• Reactivity: Atoms form ions to achieve a more stable electron configuration.
  • Stabilizing influences
    • Noble gas configuration: Ions that have the same electron configuration as a noble gas (a completely filled valence shell) are particularly stable.
      • Common stable ions, such as Na⁺, Mg²⁺, and Cl⁻, have a noble gas configuration.
    • Resonance stabilization: Ions are more stable if their charge can be delocalized, or spread out, over multiple atoms, as seen in many polyatomic anions.
    • Resonance (chemistry): If a molecule or ion can be represented by two or more valid Lewis structures that differ only in the placement of electrons, its true structure is a hybrid of these forms. This delocalization of electrons, known as resonance, increases stability.
  • Generally, reactivity increases with the magnitude of the ion's charge.

Mass spectrometry

• Definition: An analytical technique used to determine the mass-to-charge ratio (m/z) of ions, allowing for the determination of a sample's elemental composition and molecular structure.
• Process
  1. Ionization: The sample is vaporized and then ionized, typically by bombarding it with electrons to create positively charged ions.
  2. Acceleration: The newly formed ions are accelerated by an electric field into a magnetic field.
  3. Deflection: In the magnetic field, the ions are deflected based on their m/z ratio. Lighter ions and more highly charged ions are deflected more than heavier and less charged ions.
  4. Detection: A detector measures the abundance of ions at each m/z ratio, generating a mass spectrum.
• Application: can be used to identify the isotopic composition of an element, as different isotopes will have different masses and thus produce distinct peaks on the spectrum

In the periodic table of the elements (PTE), over one hundred elements are organized by criteria such as atomic number and electron configuration. It serves as an invaluable tool for predicting the properties of elements.

Organizational criteria in the PTE

The structure of the periodic table is based on the atomic structure of the elements. The following criteria determine its organization:

• Order of elements: Elements are arranged in order of increasing atomic number (number of protons).
• Periods (rows): Each row corresponds to a principal quantum number n. A new row starts when electrons first occupy a higher principal shell
• Groups (columns): elements in the same column have the same number of valence electrons and similar chemical behavior
  • Main group elements: valence electrons occupy s or p orbitals (groups 1–2 and 13–18)
    • Groups 1 and 2: valence electrons fill s-orbitals (one for group 1, two for group 2)
    • Groups 13-18: valence electrons fill p-orbitals
  • Transition metals (d-block): electrons fill d-orbitals (Groups 3–12)
  • Lanthanides/actinides (f‑block): f orbitals fill in periods 6 and 7
• Structure of periods: Because lower-numbered shells have fewer orbital types, the first few periods are shorter than later ones.

In the periodic table, columns are called groups and rows are called periods. The atomic number increases from top to bottom within a group and from left to right across a period.

The PTE is structured into blocks: s‑block (Groups 1–2), p‑block (Groups 13–18), d‑block (transition metals), and f‑block (lanthanides/actinides).

Trends in the periodic table

Due to the systematic arrangement of the elements, their properties exhibit predictable trends across the periodic table. The most important trends are summarized below.

Classification of elements

• Metals: located on the left side and in the center of the periodic table
  • Properties
    • Characterized by low ionization energies and low electronegativity
    • They are typically lustrous, malleable, ductile, and excellent conductors of heat and electricity
    • They tend to lose electrons to form cations
  • Examples: alkali metals (group 1), alkaline earth metals (group 2), transition metals, lanthanides, and actinides
• Nonmetals: located on the upper right side of the periodic table
  • Properties
    • Exhibit high ionization energies, high electron affinities, and high electronegativity
    • They are generally brittle in solid form, dull, and are poor conductors of heat and electricity
    • They tend to gain electrons to form anions
  • Examples: halogens (group 17), noble gases (group 18), as well as elements like carbon, nitrogen, oxygen, and sulfur
• Metalloids (semimetals): located in a staircase pattern between the metals and nonmetals
  • Properties
    • They possess properties that are intermediate between those of metals and nonmetals
    • Their reactivity and electronic behavior depend on the elements they are bonded to
  • Examples: include boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), and tellurium (Te)

Metallic character is a collection of properties, not a single quantifiable metric, so the distinction between metals, metalloids, and nonmetals is not always sharp.

Element groups and their properties

Elements within the same group of the periodic table share similar properties.

The most important elements in medicine

Elements typically do not occur as isolated atoms but combine to form larger units, such as molecules. This occurs via four basic types of chemical bonds, though the distinctions between them are not always sharp.

Ionic bond

An ionic bond typically forms between an element with low electronegativity (a metal) and an element with high electronegativity (a nonmetal).

• Bonding force: electrostatic attraction between oppositely charged ions
• Bond properties: non-directional (acts equally in all directions)
• Product: salts
  • Structure: To maximize attraction and minimize repulsion, cations surround anions and vice versa, forming a highly ordered, three-dimensional structure called a crystal lattice.
  • Appearance: form well-defined crystals, are typically hard and brittle, and are often colorless
  • Physical properties
    • Have high melting and boiling points due to strong lattice interactions
    • Many dissolve in water, dissociating into individual ions.
    • They do not conduct electricity in solid form but do conduct when molten or dissolved in water, as the ions are then free to move.

Atomic bond

An atomic bond, more commonly called a covalent bond, typically forms between two nonmetallic elements.

• Bonding force: the sharing of valence electrons between two atoms to form stable, shared molecular orbitals
• Theoretical explanation: Two main theories, valence bond theory and molecular orbital theory, describe covalent bond formation.
• Bond properties
  • Highly directional, forming bonds at specific angles and giving molecules a defined shape
  • Forms discrete molecules rather than extensive lattices
  • Polarization of bonds: If the two bonded atoms have significantly different electronegativities, the electrons are shared unequally. The more electronegative atom attracts the electrons more strongly, gaining a partial negative charge (δ‑), while the other atom gains a partial positive charge (δ+).
    • These polarized bonds are the basis for weak intramolecular and intermolecular interactions.
• Product: molecules
  • Polarity of molecules
    • Arises when different regions (functional groups) have unequal electron density or when bonds within the molecule are highly polar.
    • A molecule with an asymmetrical distribution of charge due to polar bonds is called a dipole. The magnitude of this charge separation is described by the dipole moment.

Partial charges represent real, measurable distributions of electron density within molecules (e.g., from spectroscopy or computational electron-density maps). They are distinct from formal charges and oxidation numbers, which are bookkeeping conventions used to track electrons in Lewis structures and redox chemistry.

Valence shell electron pair repulsion model/VSEPR model

• Definition: a model used to predict the three-dimensional geometry of a molecule based on the principle that electron pairs in the outer (valence) shell of a central atom repel each other
• Rule: Valence electron pairs arrange themselves to maximize their distance from one another.
  • The space required by electron pairs follows this general trend: lone pair > triple bond > double bond > single bond.
  • Lone pairs exert stronger repulsive forces than bonding pairs, which compresses bond angles between adjacent bonding pairs.

Molecular geometries according to the VSEPR model

Based on these rules, the following geometries can be predicted for molecules with a central atom (Z), bonded atoms (L), and lone pairs (E).

Covalent bonds in valence bond theory

This theory describes a covalent bond as the overlap of atomic orbitals from two atoms. A key concept is hybridization, where an atom's native atomic orbitals (like s and p) mix to form new, identical hybrid orbitals that are optimally arranged to form bonds.

• Types of covalent bonds
  • σ-bond (sigma bond)
    • Formed by the direct, head-on overlap of orbitals (s-s, s-p, or p-p).
    • The electron density is concentrated along the axis connecting the two nuclei.
    • All single bonds are σ-bonds.
  • π-bond (pi bond)
    • Formed by the sideways overlap of unhybridized p-orbitals.
    • The electron density is concentrated above and below the axis connecting the two nuclei.
    • A double bond consists of one σ- and one π-bond; a triple bond consists of one σ- and two π-bonds.

Atomic orbitals that are already filled with two electrons generally do not participate in bonding and are called lone pairs.

Representation of molecules

Lewis structures are a simple way to represent covalent bonds and lone pairs of electrons in a molecule.

• How to draw a Lewis structure
  1. Determine the total number of valence electrons for all atoms in the molecule.
  2. Arrange the element symbols to show how they are connected, usually with the least electronegative atom in the center.
  3. Draw single bonds between the central atom and surrounding atoms. Subtract the electrons used from the total.
  4. Distribute the remaining electrons as lone pairs on the outer atoms first to satisfy the octet rule, then place any leftover electrons on the central atom.
  5. If the central atom does not have an octet, move lone pairs from outer atoms to form double or triple bonds.

If a molecule's bonding can be represented by multiple valid Lewis structures, the actual structure is a resonance hybrid of these forms. This phenomenon is called mesomerism or resonance.

Molecules with unpaired electrons are called radicals. They are typically very reactive. Nitric oxide (NO) is a medically important radical.

Special case: coordinate bond

A coordinate bond (or dative bond) is a type of covalent bond where both shared electrons are donated by the same atom.

• Definition: a covalent bond in which one atom (the ligand) provides both bonding electrons to another atom (the central atom) that provides an empty orbital
• Product: a coordination complex
• Structure: consists of a central atom that can accept electron pairs and one or more ligands that can donate electron pairs
  • Central atom: typically a metal ion (often a transition metal) with available empty orbitals
    • Coordination number: the number of coordinate bonds formed by the central atom
    • Coordination polyhedron: the spatial arrangement of the ligands around the central atom (e.g., tetrahedral, square planar, octahedral)
  • Ligands: molecules or ions with lone pairs of electrons that can be donated (e.g., H₂O, NH₃, Cl⁻)
    • Denticity: the number of binding sites a single ligand uses to attach to the central atom (e.g., monodentate, bidentate)
      • Chelating ligands: ligands that bind to a central atom via two or more sites, like a claw (from Greek "chele")
• Properties
  • Highly directional, forming discrete molecules with specific geometries
  • The total charge of the complex is the sum of the charges of the central atom and its ligands.
  • Complexes involving transition metals are often strongly colored.

Coordination complexes play a crucial role in biology and medicine. They are found in the active sites of enzymes (e.g., iron-sulfur clusters), are crucial for transporting substances (e.g., the heme group in hemoglobin), and are used as drugs (e.g., cisplatin as a chemotherapy agent).

Cisplatin is a square planar platinum(II) complex used as a cytostatic agent. It works by binding to DNA (primarily at guanine bases), forming cross-links that distort the DNA structure. This disruption inhibits DNA replication and transcription, ultimately triggering apoptosis (programmed cell death) in rapidly dividing cancer cells. Other platinum-based drugs with similar mechanisms but different side-effect profiles include carboplatin and oxaliplatin.

Metallic bond

A metallic bond forms between atoms of metallic elements.

• Bonding force: Valence electrons are delocalized from their parent atoms and form a mobile "sea of electrons" that surrounds a fixed lattice of positive metal cations.
• Bond properties: non-directional and extends throughout the entire metal structure
• Product: metals and alloys
  • Easily deformable (malleable and ductile) because the layers of metal cations can slide past one another without disrupting the bonding provided by the electron sea
  • Excellent conductors of electricity and heat due to the mobility of the delocalized electrons

In a metallic bond, the valence electrons are delocalized and exist as a so-called electron sea.

Intermolecular forces

Weak electrostatic interactions can exist between different molecules or within different parts of a large molecule. If these interactions are due to attractive forces between dipoles, they are broadly categorized as van der Waals forces.

• Dipole: a molecule or part of a molecule with a separation of charge, resulting in a partially positive end and a partially negative end
  • Property: Dipoles can interact with each other electrostatically, much like tiny magnets.
  • Formation: In a polar bond, electrons are drawn toward the more electronegative atom, giving it a partial negative charge (δ‑) and leaving the less electronegative atom with a partial positive charge (δ+).
  • Dipole moment: a measure of a molecule's overall polarity
• Dipole interactions
  • Relatively weak electrostatic forces between molecules
  • They are crucial for determining the physical properties of substances (like boiling point) and the three-dimensional structures of large biomolecules (e.g., proteins, DNA)

Hydrogen bonds

A hydrogen bond is a particularly strong type of dipole-dipole interaction. It is fundamental to the properties of water and the structure of many biological molecules.

• Bonding force: electrostatic attraction between a partially positive hydrogen atom that is covalently bonded to a highly electronegative atom (N, O, or F) and a nearby electronegative atom
• Bond properties
  • Partners: a hydrogen-bond donor (the H atom bonded to N, O, or F) and a hydrogen-bond acceptor (a lone pair on another N, O, or F atom)
  • Geometry: typically linear and asymmetrical, with the H-atom remaining closer to the atom it is covalently bonded to
  • Consequence: leads to strong spatial ordering of molecules, significantly influencing the properties of a substance

Example: water

The macroscopic properties of water are a direct result of extensive hydrogen bonding between its molecules.

• Water molecules (H₂O)
  • Valency: Oxygen has a valency of 2 (H-O-H).
  • Geometry: the molecule is bent, with a bond angle of about 104.5°
• Density anomaly of water: Liquid water reaches its maximum density at 4°C. Ice is less dense than liquid water, which is why it floats.
• Water as a solvent: Water is an excellent solvent for polar and charged substances ("like dissolves like").
• Hydrophobic effect: Nonpolar molecules do not dissolve well in water and are repelled by it.
• Hydration: When an ionic substance dissolves in water, the ions are surrounded by water molecules, forming hydration shells. The partially negative oxygen atoms of water orient toward cations, while the partially positive hydrogen atoms orient toward anions.

Hydrogen bonds are vital in biological systems. They stabilize the secondary (α-helices and β-sheets) and tertiary structures of proteins and hold together the two strands of the DNA double helix. Because cellular environments are aqueous, the properties of water and its interactions are central to all biological processes.

Mixtures

Elements and compounds rarely exist as pure substances and are usually found as mixtures of several components.

• Heterogeneous: a mixture in which the components are not uniformly distributed and can be visually distinguished
  • Mixture of solids: a mix of different solid particles (e.g., sand and gravel)
  • Suspension: solid particles dispersed in a liquid (e.g., muddy water)
  • Foam: gas bubbles dispersed in a liquid (e.g., whipped cream)
  • Emulsion: a mixture of two immiscible liquids (e.g., oil and water)
  • Aerosol: solid or liquid particles suspended in a gas
    • Smoke: solid particles in a gas
    • Fog: liquid droplets in a gas
• Homogeneous: a mixture in which the components are uniformly distributed at a molecular level and cannot be visually distinguished
  • Alloy: a solid mixture of metals (e.g., brass, an alloy of copper and zinc)
  • Solution: a solid, liquid, or gas dissolved in a liquid solvent (e.g., salt dissolved in water)
  • Gas mixture: a mixture of gases (e.g., air)

Separation of mixtures

Various methods are used to separate mixtures based on the physical properties of their components.