Chemistry glossary
Browse existing editorial definitions and examples, then follow deeper reading in the subject’s native course. Shared senses can belong to more than one subject. Historical source definitions and tables remain in the 242-entry Physics source glossary.
Chemistry printables · Chemistry interactive and source resources
207 of 207 editorial senses
Absorbance
Chemistry · Definition
Absorbance is a logarithmic measure of how much a sample reduces transmitted light at a specified wavelength.
It depends on the ratio of incident to transmitted intensity rather than simply being the fraction absorbed.
For example, absorbance 1 corresponds to transmission of one-tenth of the incident intensity.
Acid dissociation constant
Chemistry · Definition
An acid dissociation constant describes the equilibrium for proton donation by an acid in the specified solvent.
A larger value indicates greater proton-donation extent under comparable conditions.
For example, two weak acids at the same analytical concentration can produce different hydronium concentrations because their constants differ.
Acid–base indicator
Chemistry · Definition
An acid–base indicator has differently colored acid and conjugate-base forms.
Its color change occurs over a pH range, so selection must fit the titration's equivalence region.
For example, an indicator changing in a basic range can be suitable for some weak-acid–strong-base titrations.
Activation energy
Chemistry · Definition
Activation energy is the energy barrier associated with a reaction pathway in an introductory energy-profile model.
It concerns the pathway and is distinct from the reactant-to-product energy change.
For example, a thermodynamically favored reaction can remain slow when its activation barrier is large.
Alcohol functional group
Chemistry · Definition
An alcohol contains a hydroxyl group attached to a saturated carbon atom.
This group can participate in hydrogen bonding, while the surrounding carbon structure also affects properties.
For example, ethanol is an alcohol that mixes readily with water.
Aldehyde
Chemistry · Definition
An aldehyde contains a carbonyl carbon bonded to at least one hydrogen.
This distinguishes it from a ketone, whose carbonyl carbon has two carbon-containing attachments.
For example, ethanal has a terminal carbonyl group with a hydrogen attached to its carbonyl carbon.
Aliphatic compound
Chemistry · Definition
Aliphatic compounds are organic compounds classified as nonaromatic.
They can contain open chains or nonaromatic rings and can be saturated or unsaturated.
For example, hexane and cyclohexane are both aliphatic hydrocarbons.
Aliquot
Chemistry · Definition
An aliquot is a measured portion taken from a sample.
In analysis, a representative aliquot lets a measured result be related to the original sample's total amount.
For example, a 10-milliliter portion of a uniform solution can be an aliquot of a larger prepared sample.
Alkane
Chemistry · Definition
An alkane is an acyclic saturated hydrocarbon with carbon–carbon single bonds.
Related saturated ring compounds are called cycloalkanes.
For example, propane is an alkane containing three carbon atoms.
Alkene
Chemistry · Definition
An alkene is a hydrocarbon containing a carbon–carbon double bond.
The double bond affects geometry and can create distinct arrangements that a formula alone does not specify.
For example, ethene is the simplest alkene.
Alkyl group
Chemistry · Definition
An alkyl group is a carbon–hydrogen substituent corresponding to removal of a hydrogen from an alkane.
It can form a branch on a larger organic structure.
For example, a methyl group is a one-carbon alkyl group attached through its carbon.
Alkyne
Chemistry · Definition
An alkyne is a hydrocarbon containing a carbon–carbon triple bond.
In a simple localized model, that triple bond contains one sigma bond and two pi bonds.
For example, ethyne has two carbons joined by a triple bond.
Alloy
Chemistry · Definition
An alloy is a metallic material containing more than one element.
Added atoms can replace host atoms or occupy spaces between them, changing properties such as hardness.
For example, carbon atoms in steel affect how easily iron atoms can move relative to one another.
Amine
Chemistry · Definition
An amine is an organic derivative of ammonia with one or more hydrogen atoms replaced by carbon-containing groups.
A nitrogen lone pair can allow many amines to act as bases.
For example, methylamine can accept a proton to form its conjugate acid.
Amphoteric species
Chemistry · Definition
An amphoteric species can behave as either an acid or a base.
An amphiprotic species specifically can donate or accept a proton in Brønsted reactions.
For example, bicarbonate can accept a proton to form carbonic acid or donate one to form carbonate.
Analyte
Chemistry · Definition
An analyte is the substance being identified or measured in an analysis.
Its amount or concentration is the target of the measurement rather than necessarily the largest component of the sample.
For example, an unknown acid is the analyte in an acid–base titration.
Anion
Chemistry · Definition
An anion is a negatively charged ion.
It has more electrons than needed to balance its positive nuclear charge.
For example, a chlorine atom gaining one electron forms a chloride anion.
Anode
Chemistry · Definition
The anode is the electrode where oxidation occurs.
This definition applies to both galvanic and electrolytic cells.
For example, a zinc electrode releasing zinc ions and electrons is an anode in the zinc–copper galvanic model.
Aqueous solution
Chemistry · Definition
An aqueous solution is a solution in which water is the solvent.
The dissolved substance may originate as a solid, liquid, or gas, so aqueous does not mean that a solid has melted.
For example, sodium chloride dissolved in water is aqueous even though the pure salt is solid.
Aromaticity
Chemistry · Definition
Aromaticity describes a special stability associated with particular cyclic, conjugated electron arrangements.
Having a ring alone does not make a compound aromatic.
For example, benzene is aromatic while cyclohexane is not.
Arrhenius equation
Chemistry · Definition
The Arrhenius equation models how a rate constant depends on absolute temperature and an activation-energy parameter.
Within the applicable model, increasing temperature raises the rate constant for a positive activation energy.
For example, a larger activation-energy parameter can make a reaction's rate constant more sensitive to a temperature change.
Atomic orbital
Chemistry · Definition
An atomic orbital is a quantum state used to describe an electron in an atom.
Its spatial description gives a probability distribution rather than a definite path around the nucleus.
For example, an s orbital has a spherical spatial probability distribution, and one orbital can accommodate at most two electrons with opposite spins.
Aufbau principle
Chemistry · Definition
The Aufbau principle describes constructing a ground-state electron configuration by occupying lower-energy orbitals before higher-energy ones.
Orbital energies and some observed configurations require more care than a simple fixed filling diagram suggests.
For example, helium fills its 1s orbital before lithium begins occupying 2s.
Avogadro’s constant
Chemistry · Definition
Avogadro’s constant connects a number of specified entities to an amount in moles.
Its exact value is 6.02214076 × 10²³ per mole.
For example, multiplying two moles of helium atoms by this constant gives their number of atoms.
Beer–Lambert law
Chemistry · Definition
The Beer–Lambert law relates absorbance to concentration, light-path length, and molar absorptivity under suitable conditions.
Comparisons require a controlled wavelength, consistent path length, and an appropriate calibration.
For example, doubling concentration doubles absorbance in the law's linear range if the other factors remain fixed.
Bond dipole
Chemistry · Definition
A bond dipole describes a bond's separation of positive and negative charge.
Its direction and magnitude depend on the distribution of electron density.
For example, an O–H bond has a dipole because its electron density is drawn toward oxygen.
Bond enthalpy
Chemistry · Definition
A bond enthalpy is the enthalpy required to break a specified bond in the gas phase.
Tables often give average values, so estimates for complex reactions are approximate.
For example, a reaction estimate adds the costs of bonds broken and subtracts the energy released by bonds formed.
Bond order
Chemistry · Definition
Bond order describes the degree of bonding between two atoms in a specified bonding model.
Localized single, double, and triple bonds have orders one, two, and three, while delocalized models can give fractional values.
For example, equivalent bonds in a resonance description can have bonding intermediate between a single and double bond.
Brønsted acid and base
Chemistry · Definition
A Brønsted acid donates a proton, and a Brønsted base accepts a proton.
These roles describe a particular reaction, so some substances can act in either role.
For example, water accepts a proton from hydrogen chloride but donates one to ammonia.
Buffer
Chemistry · Definition
A buffer contains appreciable amounts of a weak acid and its conjugate base, or the corresponding weak-base pair.
Those partners consume small added amounts of strong base or acid, limiting the pH change within the buffer's capacity.
For example, an acetic acid–acetate mixture can resist modest additions of acid or base.
Buffer capacity
Chemistry · Definition
Buffer capacity describes how much acid or base a buffer can absorb while keeping its pH change within a specified limit.
It depends on the amounts of both buffer partners as well as their ratio.
For example, a more concentrated buffer can generally absorb more added reagent than a dilute buffer with the same initial pH.
Calorimetry
Chemistry · Definition
Calorimetry determines heat transfer from measured thermal changes in a specified setup.
The calculation accounts for the measured material and any supplied calorimeter heat capacity using an energy balance.
For example, warming solution can reveal heat released by a reaction when the reaction and surroundings are correctly distinguished.
Carbonyl group
Chemistry · Definition
A carbonyl group contains a carbon atom double-bonded to oxygen.
Its surrounding attachments distinguish classes such as aldehydes, ketones, and carboxylic acids.
For example, the carbonyl carbon in acetone is attached to two carbon-containing groups.
Carboxylic acid
Chemistry · Definition
A carboxylic acid contains a carboxyl group with a carbonyl and a hydroxyl attached to the same carbon.
Its acidic proton can be transferred to a base, forming a carboxylate conjugate base.
For example, acetic acid forms acetate after proton donation.
Catalyst
Chemistry · Definition
A catalyst participates in a reaction pathway and is regenerated overall.
It can increase the reaction rate by providing a different pathway without changing the reaction's equilibrium constant at fixed temperature.
For example, a catalyst can help a mixture reach equilibrium sooner without changing the final equilibrium composition.
Cathode
Chemistry · Definition
The cathode is the electrode where reduction occurs.
It may deposit metal or host another reduction without necessarily gaining mass.
For example, a copper electrode receiving electrons that reduce dissolved copper ions is a cathode.
Cation
Chemistry · Definition
A cation is a positively charged ion.
Its total positive nuclear charge exceeds its total electron charge.
For example, calcium commonly forms a cation with charge +2 by losing two electrons.
Cell potential
Chemistry · Definition
Cell potential is an electrical potential difference associated with a specified cell reaction and conditions.
A reduction-potential table supplies reference values for half-reactions, and voltage is not multiplied when equation coefficients are scaled.
For example, balancing twice as many transferred electrons changes reaction amount bookkeeping but does not double the cell voltage.
Chemical equilibrium
Chemistry · Definition
Chemical equilibrium is a dynamic state in which forward and reverse reaction rates are equal.
Macroscopic composition remains constant in a closed system under fixed conditions, but reactant and product concentrations need not be equal.
For example, a reversible reaction can reach equilibrium with much more reactant than product.
Chirality
Chemistry · Definition
Chirality is the property of an object that cannot be superimposed on its mirror image.
Molecules related as nonsuperimposable mirror images are enantiomers and can interact differently with other chiral species.
For example, a tetrahedral carbon bonded to four different groups can produce a pair of enantiomers.
Chromatography
Chemistry · Definition
Chromatography separates components through their differing interactions with mobile and stationary phases.
The measured movement or retention depends on the chosen materials and conditions.
For example, two dissolved dyes can travel different distances on a chromatographic strip.
Colligative property
Chemistry · Definition
Colligative properties depend chiefly on the number of dispersed solute particles in an appropriate dilute-solution model.
Familiar examples include boiling-point elevation and freezing-point depression.
For example, an ideally dissociated salt can affect freezing temperature more than the same molal amount of a nondissociating solute, though real concentrated solutions require additional care.
Collision model
Chemistry · Definition
The collision model explains reaction in terms of encounters between particles.
Effective encounters require sufficient energy and an appropriate arrangement, so not every collision produces reaction.
For example, increasing temperature can increase the fraction of collisions able to cross an energy barrier.
Common-ion effect
Chemistry · Definition
The common-ion effect is an equilibrium response caused by adding a species already present in the relevant equilibrium.
In the simple dissolution model, adding a shared ion lowers the amount of solid that can dissolve at fixed temperature.
For example, added chloride can reduce silver chloride's solubility when other reactions are negligible.
Complex ion
Chemistry · Definition
A complex ion is a charged coordination species containing a central atom or ion and bound ligands.
Ligand binding can change properties such as color and the amounts of free metal ions in solution.
For example, ammonia ligands can bind to a copper ion to form a differently colored complex.
Concentration cell
Chemistry · Definition
A concentration cell uses the same redox couple under differing activities in its two half-cells.
The concentration difference can produce a potential even though the standard half-cell potentials cancel.
For example, an ideal metal-ion concentration cell can operate in the direction that reduces the ion-concentration imbalance.
Condensation
Chemistry · Definition
Condensation is a phase change from gas to liquid.
It releases energy to the surroundings as particles form more closely associated liquid arrangements.
For example, water vapor condensing on a cooler surface forms liquid droplets.
Conjugate acid–base pair
Chemistry · Definition
A conjugate acid–base pair consists of species differing by one proton.
Losing a proton forms an acid's conjugate base, while gaining one forms a base's conjugate acid.
For example, ammonium and ammonia form a conjugate acid–base pair.
Core electron
Chemistry · Definition
Core electrons occupy inner shells and are usually less directly involved in bonding than valence electrons.
They help shield outer electrons from the nucleus's full positive charge.
For example, sodium's ten inner electrons form its neon-like core beneath its outer electron.
Corrosion
Chemistry · Definition
Corrosion is deterioration of a material through chemical or electrochemical reactions with its environment.
Metal corrosion commonly involves oxidation coupled to a reduction process.
For example, rust formation involves iron oxidation together with reactions involving oxygen and water.
Coupled reactions
Chemistry · Definition
Coupled reactions are linked through a mechanism or energy-transfer connection so that one can help drive another.
Adding favorable and unfavorable free-energy values is meaningful only when the processes are physically or chemically connected.
For example, a shared intermediate can connect an unfavorable transformation to a sufficiently favorable step.
Covalent bond
Chemistry · Definition
A covalent bond involves shared electron density between atoms.
The sharing can be approximately equal or unequal depending on the atoms and their environment.
For example, the bond in a hydrogen molecule shares electrons between two identical atoms.
Covalent network solid
Chemistry · Definition
A covalent network solid contains atoms joined by covalent bonds throughout an extended structure.
Changing or melting that structure requires more than separating intact small molecules.
For example, diamond contains a three-dimensional network of carbon–carbon bonds.
Crystal lattice
Chemistry · Definition
A crystal lattice describes the repeating arrangement of structural sites in a crystalline solid.
The occupants can be ions, atoms, or molecules depending on the material.
For example, an ionic lattice places opposite charges near one another throughout a sodium chloride crystal.
Current efficiency
Chemistry · Definition
Current efficiency is the fraction of transferred charge that produces the specified electrode process.
Competing reactions or lost material can make the observed yield lower than the ideal charge-based prediction.
For example, an 80 percent efficiency means only 0.80 of the total charge contributes to the selected deposition process.
Deionized water
Chemistry · Definition
Deionized water has had most dissolved ions removed.
It is not necessarily free of every nonionic impurity or microorganism.
For example, using deionized water can reduce unintended ions in a prepared analytical solution.
Diamagnetism
Chemistry · Definition
Diamagnetism is a weak response opposing an applied magnetic field.
In the introductory classification, species with all electrons paired are described as diamagnetic when no stronger magnetic behavior dominates.
For example, neon has a filled electron configuration and is diamagnetic.
Diffusion
Chemistry · Definition
Diffusion is the net spreading of particles associated with random motion and concentration differences.
Individual particles move in many directions even as the overall distribution becomes more uniform.
For example, a gas initially concentrated in one part of a container spreads throughout the available space.
Dilution
Chemistry · Definition
Dilution lowers solute concentration by adding solvent while conserving the solute amount when no reaction or loss occurs.
Concentration calculations use the final solution volume rather than simply the volume of solvent added.
For example, increasing a solution's volume from 100 to 200 milliliters halves its concentration if its solute amount is unchanged.
Dipole–dipole forces
Chemistry · Definition
Dipole–dipole forces involve interactions between permanent molecular dipoles.
Favorable orientations place relatively positive and negative regions near one another.
For example, neighboring hydrogen chloride molecules can attract through their permanent dipoles.
Distillation
Chemistry · Definition
Distillation separates mixture components using differences in their vaporization behavior.
Vapor is collected and condensed, with composition depending on the mixture and conditions.
For example, a vapor can be enriched in the more volatile component of a liquid mixture.
Effective nuclear charge
Chemistry · Definition
Effective nuclear charge describes the net nuclear attraction experienced by an electron after shielding is considered.
It depends on nuclear charge and the electron distribution rather than proton count alone.
For example, increasing effective attraction across a main-group period generally pulls valence electrons closer to the nucleus.
Effusion
Chemistry · Definition
Effusion is the escape of gas particles through an opening small enough that particle–particle collisions within the opening are negligible.
Under comparable conditions, lighter gases effuse faster on average than heavier gases.
For example, helium effuses faster than argon at the same temperature in the ideal model.
Electrochemical cell
Chemistry · Definition
An electrochemical cell separates oxidation and reduction into electrode processes.
A galvanic cell delivers electrical energy from a favored reaction, while an electrolytic cell uses supplied electrical energy to drive an unfavored direction.
For example, a zinc–copper galvanic model transfers electrons through the external wire while ions carry charge through the solutions.
Electrolysis
Chemistry · Definition
Electrolysis uses an external electrical source to drive an electrode reaction in an otherwise unfavorable direction.
The products depend on the specified half-reactions and medium, including possible competition from water.
For example, reduction of metal ions at a cathode can produce an electroplated metal layer.
Electrolyte
Chemistry · Definition
An electrolyte forms mobile ions when dissolved and can allow a solution to conduct electric current.
Strong and weak electrolytes differ in the extent to which dissolved species form ions under the stated conditions.
For example, dissolved sodium chloride supplies sodium and chloride ions to water.
Electron affinity
Chemistry · Definition
Electron affinity describes the energy change associated with adding an electron to a gaseous atom.
Tables may use different sign conventions, so the stated definition and sign must be checked.
For example, adding an electron to a chlorine atom releases energy under the usual energy-change convention.
Electron configuration
Chemistry · Definition
An electron configuration describes how an atom's or ion's electrons occupy shells and subshells.
A ground-state configuration places electrons in the lowest-energy arrangement consistent with the electron-filling rules.
For example, neutral helium has two electrons in its 1s subshell.
Electron delocalization
Chemistry · Definition
Electron delocalization means that an electronic description extends across multiple atoms rather than being confined to one localized bond or atom.
It can affect bond equivalence and charge distribution.
For example, nitrate's bonding electrons are described across several nitrogen–oxygen positions.
Electron domain
Chemistry · Definition
An electron domain is a region of electron density around an atom in the VSEPR model.
A bond of any order counts as one domain, and a lone pair counts as another.
For example, two bonds and two lone pairs give water's oxygen four domains and a steric number of four.
Electron shielding
Chemistry · Definition
Electron shielding is the reduction in nuclear attraction experienced by an electron because other electrons also repel it.
Inner electrons generally shield an outer electron more effectively than electrons in the same shell.
For example, adding an occupied inner shell helps explain why atomic size usually increases down a periodic-table group.
Electron subshell
Chemistry · Definition
A subshell is a set of atomic orbitals with the same principal and angular-momentum quantum numbers.
The s, p, d, and f subshells contain one, three, five, and seven orbitals, respectively.
For example, a p subshell can accommodate six electrons because each of its three orbitals can hold two.
Electronegativity
Chemistry · Definition
Electronegativity describes an atom's tendency to attract shared electron density in a bond.
It helps predict bond polarity but is not the same as ionization energy or electron affinity.
For example, oxygen attracts shared electrons more strongly than hydrogen in an O–H bond.
Elementary reaction
Chemistry · Definition
An elementary reaction is one step in a proposed mechanism rather than the sum of multiple steps.
For a specified elementary step, the rate-law concentration factors follow its reacting-particle event.
For example, a bimolecular elementary collision can have a rate proportional to the concentrations of its two reactants.
Empirical formula
Chemistry · Definition
An empirical formula gives the simplest whole-number ratio of atoms in a compound.
It is obtained by comparing amounts in moles of each element rather than comparing their masses directly.
For example, both ethene and propene have the empirical formula CH₂ despite having different molecular formulas.
Endothermic process
Chemistry · Definition
An endothermic process absorbs heat into the chosen system.
Under constant-pressure conditions with only pressure–volume work, its enthalpy change is positive.
For example, melting ice absorbs heat while the phase change occurs.
Enthalpy
Chemistry · Definition
Enthalpy is a state function defined from internal energy, pressure, and volume.
At constant pressure with only pressure–volume work, an enthalpy change equals the heat transferred to the system.
For example, a negative reaction enthalpy indicates heat release for the written reaction under those conditions.
Entropy
Chemistry · Definition
Entropy is a state function related to the number of microscopic arrangements compatible with a macroscopic state.
Changes depend on both matter and energy distribution, so visual messiness alone is not a reliable measure.
For example, an ideal gas expanding at fixed temperature has more accessible positional arrangements and increased entropy.
Equilibrium constant
Chemistry · Definition
An equilibrium constant is the value of the reaction quotient at equilibrium for a specified reaction and temperature.
Its magnitude indicates equilibrium composition preference rather than reaction speed.
For example, a large constant often indicates a product-favored equilibrium, but the starting composition still determines the immediate net direction.
Equivalence point
Chemistry · Definition
The equivalence point is the condition where titrant and analyte have reacted in the balanced equation's stoichiometric ratio.
Its pH depends on the species remaining and is not always seven.
For example, a weak-acid–strong-base titration has a basic equivalence-point solution under the usual model.
Ester
Chemistry · Definition
A carboxylic ester has a carbonyl carbon bonded to an oxygen that also connects to a carbon-containing group.
It differs from a carboxylic acid because that linking oxygen does not carry the acid's proton.
For example, ethyl acetate is an ester containing an ethyl group attached through oxygen.
Ether
Chemistry · Definition
An ether has an oxygen atom bonded to two carbon-containing groups.
Its oxygen can accept hydrogen bonds, but an ether lacking an O–H group cannot donate one in the usual introductory model.
For example, dimethyl ether has an oxygen joining two methyl groups.
Evaporation
Chemistry · Definition
Evaporation is vaporization occurring at a liquid's surface.
It can occur below the boiling point as some particles escape into the gas phase.
For example, liquid water can evaporate at room temperature and leave less liquid behind.
Exothermic process
Chemistry · Definition
An exothermic process transfers heat from the chosen system to its surroundings.
Under constant-pressure conditions with only pressure–volume work, its enthalpy change is negative.
For example, freezing water releases heat to the surroundings.
Faraday’s constant
Chemistry · Definition
Faraday’s constant is the magnitude of electric charge carried by one mole of electrons.
It is approximately 96485 coulombs per mole and connects measured charge to the amount of electrons transferred.
For example, depositing one mole of a metal from ions with charge +2 requires two moles of electrons in the ideal model.
Filtration
Chemistry · Definition
Filtration separates suspended particles from a fluid using a barrier that retains appropriate particle sizes.
It generally does not remove individually dissolved ions or molecules.
For example, a filter can separate an undissolved precipitate from the surrounding solution.
Formal charge
Chemistry · Definition
Formal charge is electron bookkeeping that assigns half of each bonding pair to each bonded atom.
It compares that assigned electron count with the neutral atom's valence-electron count and is not a direct measurement of actual charge distribution.
For example, all formal charges in an ion's Lewis structure must sum to the ion's total charge.
Formula unit
Chemistry · Definition
A formula unit specifies the simplest whole-number ratio of ions in an ionic compound.
It is a counting unit for an extended solid rather than an isolated molecule.
For example, one calcium chloride formula unit represents one calcium ion and two chloride ions.
Functional group
Chemistry · Definition
A functional group is a characteristic atom arrangement associated with patterns of chemical behavior in organic compounds.
Its effect depends on the rest of the molecule as well as the group's identity.
For example, an alcohol's hydroxyl group can support hydrogen bonding.
Gibbs free energy
Chemistry · Definition
Gibbs free energy combines enthalpy and entropy contributions to a system at specified temperature and pressure.
For a reaction at constant temperature and pressure, a negative actual reaction free-energy change favors its forward direction, while standard and actual values must be distinguished.
For example, a favorable reaction can still be slow if its pathway has a large activation barrier.
Half-reaction
Chemistry · Definition
A half-reaction describes either the oxidation or reduction part of a redox process.
It includes electrons explicitly so that charge and electron transfer can be balanced.
For example, zinc forming a zinc ion and two electrons is an oxidation half-reaction.
Heat capacity
Chemistry · Definition
Heat capacity is the heat required to raise a specified object's temperature by one temperature unit under stated conditions.
It refers to the entire object and differs from specific heat capacity, which is expressed per unit mass.
For example, doubling the amount of the same material approximately doubles its heat capacity.
Heisenberg uncertainty principle
Chemistry · Definition
The Heisenberg uncertainty principle places an intrinsic limit on simultaneously specifying a particle's position and momentum.
It is a property of quantum states rather than merely a problem caused by poor measuring instruments.
For example, atomic electrons are described with probability distributions instead of exact classical trajectories around a nucleus.
Henderson–Hasselbalch equation
Chemistry · Definition
The Henderson–Hasselbalch equation relates a buffer's pH to an acid's pKa and the ratio of conjugate-base to acid activities.
Concentration ratios approximate activity ratios only when the model's conditions are appropriate.
For example, equal effective amounts of an acid and its conjugate base give pH approximately equal to pKa.
Henry’s law
Chemistry · Definition
Henry’s law relates a gas's dissolved concentration to its partial pressure over the solution at fixed temperature in an appropriate dilute model.
The proportionality constant depends on the gas, solvent, temperature, and the convention used to write the relation.
For example, increasing carbon dioxide's partial pressure can increase its dissolved concentration when the model applies.
Hess’s law
Chemistry · Definition
Hess’s law states that an overall enthalpy change is independent of the reaction path.
Balanced step equations can be added while their enthalpy changes are added with matching reversals and scaling.
For example, reversing a step changes the sign of its enthalpy before it is included in a sum.
Hund’s rule
Chemistry · Definition
Hund’s rule describes the lowest-energy arrangement within a set of degenerate orbitals.
Electrons occupy those orbitals singly with parallel spins before pairing.
For example, nitrogen's three 2p electrons occupy three different p orbitals in its ground state.
Hydrocarbon
Chemistry · Definition
A hydrocarbon is a compound containing only carbon and hydrogen.
Its carbon skeleton may contain chains, rings, or multiple bonds.
For example, methane and ethene are both hydrocarbons despite having different bonding structures.
Hydrogen bond
Chemistry · Definition
Hydrogen bonding is a directional interaction involving hydrogen bonded to a sufficiently electronegative atom and an appropriate electron-rich acceptor.
In the introductory model, N, O, and F commonly provide the relevant partners.
For example, water molecules hydrogen bond when one molecule's O–H group interacts with another molecule's oxygen lone pair.
Hydronium
Chemistry · Definition
Hydronium is the H₃O⁺ ion used in the introductory model of proton behavior in water.
Aqueous H⁺ notation is shorthand for a solvated proton rather than a bare proton moving independently in the liquid.
For example, a Brønsted acid can transfer a proton to water and form hydronium.
ICE table
Chemistry · Definition
An ICE table organizes initial concentrations, stoichiometric changes, and equilibrium concentrations.
Its change entries follow the balanced equation and do not imply equal changes for every species.
For example, producing two units of a product for each reactant consumed requires twice the concentration-change entry for that product.
Ideal gas law
Chemistry · Definition
The ideal gas law relates pressure, volume, amount, and absolute temperature in an idealized gas model.
It treats particles as having negligible volume and no intermolecular attractions.
For example, at fixed amount and temperature, doubling the container volume halves the ideal gas pressure.
Induced dipole
Chemistry · Definition
An induced dipole is a separation of charge produced when an external electric field distorts a particle's electron distribution.
Induction can be caused by an ion, a permanent dipole, or a fluctuating dipole, so the term is broader than London dispersion.
For example, a polar molecule can induce a dipole in a neighboring nonpolar molecule.
Intermolecular force
Chemistry · Definition
Intermolecular forces are interactions between molecules.
They affect properties such as boiling point and solubility and are distinct from the bonds within each molecule.
For example, vaporizing water separates molecules while largely preserving their O–H bonds.
Ion
Chemistry · Definition
An ion is an atom or group of atoms with a net electric charge.
Ordinary ion formation changes electron count without changing the element's proton count.
For example, a sodium atom losing one electron becomes a sodium ion with charge +1.
Ion–dipole force
Chemistry · Definition
An ion–dipole force is an interaction between an ion and a polar molecule.
The molecule orients its oppositely charged region toward the ion.
For example, water's oxygen end points toward a dissolved sodium cation.
Ionic bonding
Chemistry · Definition
Ionic bonding is electrostatic attraction among oppositely charged ions.
An ionic solid normally forms an extended lattice rather than separate ion-pair molecules.
For example, sodium chloride consists of a repeating array of sodium and chloride ions.
Ionic strength
Chemistry · Definition
Ionic strength summarizes dissolved-ion concentrations with weighting by the square of each ion's charge number.
It helps describe deviations between ionic activities and concentrations.
For example, at equal concentration a divalent ion contributes more to ionic strength than a monovalent ion.
Ionization energy
Chemistry · Definition
Ionization energy is the energy required to remove an electron from a gaseous atom or ion.
Removing an electron requires energy, and successive removals can reveal a transition from valence to core electrons.
For example, sodium's second ionization requires much more energy than its first because it removes a core electron.
Isotope
Chemistry · Definition
Isotopes are atoms of the same element with different numbers of neutrons.
They have the same proton number but different masses.
For example, carbon-12 and carbon-13 each have six protons but have six and seven neutrons, respectively.
Ketone
Chemistry · Definition
A ketone contains a carbonyl carbon bonded to two carbon-containing groups.
The carbonyl contributes polarity and can accept hydrogen bonds in the introductory model.
For example, acetone is a ketone with a methyl group on each side of its carbonyl carbon.
Kinetic molecular theory
Chemistry · Definition
Kinetic molecular theory connects gas pressure and temperature to particle motion and collisions.
In its ideal form, particles move randomly, occupy negligible volume, and undergo elastic collisions without intermolecular attractions.
For example, hotter ideal gas particles have a greater average translational kinetic energy.
Le Châtelier’s principle
Chemistry · Definition
Le Châtelier’s principle predicts the qualitative response when a system initially at equilibrium is disturbed.
A justified explanation distinguishes the immediate disturbance from the subsequent reaction toward a new equilibrium.
For example, adding a reactant can favor its net consumption, while temperature changes can also alter the equilibrium constant.
Lewis acid and base
Chemistry · Definition
A Lewis acid accepts an electron pair, and a Lewis base donates an electron pair to form a bonding interaction.
This definition includes some acid–base interactions without proton transfer.
For example, a nitrogen lone pair can be donated when ammonia binds to an electron-pair acceptor.
Lewis structure
Chemistry · Definition
A Lewis structure represents valence electrons using bonds and lone pairs.
Its construction requires an electron count, an atom arrangement, and a final check against the total available electrons and charge.
For example, water's Lewis structure has two O–H bonds and two lone pairs on oxygen.
Ligand
Chemistry · Definition
A ligand is an ion or molecule that binds to a central atom or ion in a coordination complex.
In the common introductory model, it donates an electron pair to the central metal.
For example, water molecules can act as ligands around a metal ion.
Limiting reactant
Chemistry · Definition
The limiting reactant is consumed first under the specified complete-reaction model and limits the product amount.
Identifying it requires comparing available amounts with the balanced equation's ratios.
For example, abundant hydrogen cannot produce more water after the available oxygen has been consumed.
London dispersion forces
Chemistry · Definition
London dispersion forces arise from fluctuating electron distributions and the dipoles they induce in neighboring particles.
They occur between all atoms and molecules, including nonpolar ones.
For example, dispersion attracts neighboring argon atoms even though they have no permanent molecular dipoles.
Lone pair
Chemistry · Definition
A lone pair is a pair of valence electrons assigned to one atom rather than to a bond in a Lewis model.
Lone pairs influence geometry and can participate in interactions or new bonds.
For example, oxygen's two lone pairs help explain water's bent molecular shape.
Mass fraction
Chemistry · Definition
Mass fraction is a component's mass divided by the total mass of the sample.
Multiplying that fraction by 100 gives its mass percentage.
For example, 5 grams of a solute in 100 grams of solution corresponds to a mass fraction of 0.05 and a mass percentage of 5 percent.
Mass spectrum
Chemistry · Definition
A mass spectrum displays detected ions according to their mass-to-charge ratio and relative signal.
In this lesson's elemental model, singly charged monatomic ions make peak positions correspond numerically to isotope masses.
For example, two isotope peaks can be combined with their relative abundances to determine a weighted average atomic mass.
Metallic bonding
Chemistry · Definition
Metallic bonding involves metal centers held together by electrons delocalized through the solid.
The mobile electronic charge helps explain electrical conductivity and the ability of many metals to deform.
For example, copper remains bonded as its atoms shift when a wire is bent.
Mixture
Chemistry · Definition
A mixture contains two or more substances whose proportions can vary.
Its components retain their chemical identities until a chemical reaction changes them.
For example, water and ethanol can form mixtures with many different compositions.
Mobile phase
Chemistry · Definition
The mobile phase is the fluid that carries a sample through a chromatographic system.
A component's movement depends on its interactions with this phase and the stationary phase.
For example, a solvent can carry dissolved dyes along a chromatographic strip.
Molality
Chemistry · Definition
Molality is the amount of solute in moles divided by the solvent mass in kilograms.
It uses solvent mass rather than solution volume and differs from molarity.
For example, one mole of solute dissolved in two kilograms of water has molality 0.50 mole per kilogram.
Molar absorptivity
Chemistry · Definition
Molar absorptivity describes a species' absorption strength at a specified wavelength in the Beer–Lambert model.
It depends on the absorbing species and measurement conditions, while concentration and path length are separate factors.
For example, selecting a wavelength near an absorption maximum can give a larger absorbance for the same concentration and path length.
Molar mass
Chemistry · Definition
Molar mass is the mass of one mole of a specified substance.
It is commonly expressed in grams per mole and connects a sample's mass to its amount of substance.
For example, water has a molar mass of approximately 18.02 grams per mole.
Molarity
Chemistry · Definition
Molarity is the amount of a solute in moles divided by the final solution volume in liters.
It describes concentration rather than the total amount of solute.
For example, 0.10 mole dissolved to make 0.50 liter of solution gives a molarity of 0.20 mole per liter.
Mole
Chemistry · Definition
A mole is an amount of substance containing exactly 6.02214076 × 10²³ specified entities.
The entities can be atoms, molecules, ions, or formula units, so the counted entity must be named.
For example, one mole of oxygen molecules contains two moles of oxygen atoms.
Mole fraction
Chemistry · Definition
Mole fraction is a component's amount in moles divided by the mixture's total amount in moles.
The fractions of all components sum to one.
For example, one mole of one gas mixed with three moles of another gives the first gas a mole fraction of 0.25.
Molecular formula
Chemistry · Definition
A molecular formula gives the actual number of each kind of atom in a molecule.
It can be a whole-number multiple of the empirical formula.
For example, glucose has molecular formula C₆H₁₂O₆ and empirical formula CH₂O.
Molecular orbital
Chemistry · Definition
A molecular orbital is an electronic state described over a molecule rather than assigned exclusively to a single atom.
Molecular-orbital models can describe bonding, antibonding, and nonbonding electron arrangements.
For example, occupying a bonding orbital between two atoms can contribute to holding those atoms together.
Molecular polarity
Chemistry · Definition
Molecular polarity concerns a molecule's overall separation of charge.
Both bond polarity and molecular geometry determine whether individual bond dipoles cancel.
For example, carbon dioxide has polar bonds but is nonpolar overall because its linear bond dipoles cancel.
Nernst equation
Chemistry · Definition
The Nernst equation relates a cell's potential to temperature and the reaction quotient for its written reaction.
It distinguishes the current composition-dependent potential from the standard reference potential.
For example, approaching chemical equilibrium brings the ideal cell's actual potential toward zero.
Net ionic equation
Chemistry · Definition
A net ionic equation shows the species that undergo the chemical change.
It removes spectator ions while preserving atom and charge balance.
For example, precipitation of silver chloride can be written using aqueous silver and chloride ions forming solid silver chloride.
Nitrile
Chemistry · Definition
A nitrile contains a carbon atom triple-bonded to nitrogen.
The nitrile carbon is part of the functional group and affects systematic naming.
For example, acetonitrile contains a methyl group attached to a nitrile carbon.
Octet rule
Chemistry · Definition
The octet rule is a useful model in which many main-group atoms form arrangements with eight valence-shell electrons.
It has important exceptions, including hydrogen's duet and electron-deficient or odd-electron species.
For example, methane's Lewis structure gives carbon an octet through four bonds.
Orbital hybridization
Chemistry · Definition
Hybridization is a localized bonding model that combines atomic-orbital descriptions on an atom.
It can help describe bond directions but is not a statement that isolated atomic orbitals physically mix as a timed step.
For example, an sp³ model describes four tetrahedrally directed orbitals around carbon in methane.
Organic compound
Chemistry · Definition
Organic compounds comprise a broad class of carbon-based compounds, often built around carbon–carbon and carbon–hydrogen structures.
Some carbon-containing substances, such as carbon dioxide and carbonates, are conventionally classified as inorganic.
For example, ethanol is organic while sodium carbonate is usually classified as inorganic.
Oxidation
Chemistry · Definition
Oxidation is electron loss, represented by an increase in oxidation number.
It occurs together with reduction so that electrons are conserved in the overall reaction.
For example, a zinc atom becoming a zinc ion with charge +2 undergoes oxidation.
Oxidation number
Chemistry · Definition
An oxidation number is a formal electron-accounting value assigned using specified rules.
Changes in oxidation number help identify oxidation and reduction but are not always actual atomic charges.
For example, oxygen usually has oxidation number −2 in water while hydrogen has +1.
Oxidizing agent
Chemistry · Definition
An oxidizing agent causes another species to lose electrons.
The agent accepts those electrons and is itself reduced.
For example, a copper ion acts as an oxidizing agent when it accepts electrons from zinc metal.
Paramagnetism
Chemistry · Definition
Paramagnetism is attraction to an applied magnetic field associated with unpaired electrons in the introductory model.
The relevant electron arrangement must be considered rather than inferred from atom count alone.
For example, molecular oxygen has unpaired electrons and is paramagnetic.
Partial charge
Chemistry · Definition
A partial charge describes an uneven distribution of electron density within a bond or molecule.
It is marked by a delta symbol and does not necessarily represent a full ionic charge.
For example, oxygen carries a partial negative charge in water while its hydrogens carry partial positive charges.
Partial pressure
Chemistry · Definition
A gas's partial pressure is the pressure it would exert alone in the mixture's volume at the same temperature.
In an ideal mixture, the component partial pressures add to the total pressure.
For example, a component with mole fraction 0.20 contributes 0.20 of the total ideal pressure.
Passivation
Chemistry · Definition
Passivation occurs when a surface layer reduces further reaction of the underlying material.
It can make an idealized redox prediction differ from the observed reaction rate or cell behavior.
For example, a protective oxide layer can slow a metal's further oxidation.
Pauli exclusion principle
Chemistry · Definition
The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers.
Consequently, an orbital can contain at most two electrons, and those electrons must have opposite spins.
For example, the two electrons in helium's 1s orbital are represented by opposing arrows.
Percent ionization
Chemistry · Definition
Percent ionization is the fraction of an initially supplied acid or base that forms the relevant ions, multiplied by 100.
It depends on concentration and conditions as well as the species' equilibrium constant.
For example, a weak acid can have a larger ionized fraction after dilution even though its hydronium concentration decreases.
pH
Chemistry · Definition
pH is the negative base-ten logarithm of hydronium activity.
In suitable dilute models, activity is approximated by molar concentration relative to the standard concentration.
For example, lowering hydronium concentration by a factor of ten raises pH by one unit.
Phase change
Chemistry · Definition
A phase change alters a substance's physical state without changing its chemical identity.
It involves a change in particle arrangement and interactions and can transfer energy without raising temperature during a pure-substance transition at fixed pressure.
For example, vaporizing liquid water separates water molecules into a gas.
Photoelectron spectroscopy
Chemistry · Definition
Photoelectron spectroscopy uses photons to remove electrons and measure their binding energies.
Peak position indicates how tightly electrons are bound, while peak intensity can indicate the relative population of a subshell in the simplified model.
For example, an inner-shell electron normally produces a peak at higher binding energy than a valence electron.
Photon
Chemistry · Definition
A photon is a quantum of electromagnetic radiation.
Its energy increases with frequency and decreases as wavelength increases.
For example, a blue-light photon carries more energy than a red-light photon.
Pi bond
Chemistry · Definition
A pi bond has bonding electron density on opposite sides of the line joining two nuclei.
In a localized model, it results from sideways overlap of appropriately oriented orbitals.
For example, a carbon–carbon double bond contains one sigma bond and one pi bond.
pKa
Chemistry · Definition
pKa is the negative base-ten logarithm of an acid's dissociation constant.
A smaller pKa means a stronger acid under comparable conditions.
For example, when an acid and its conjugate base have equal activities, their buffer's pH equals that acid's pKa.
pOH
Chemistry · Definition
pOH is the negative base-ten logarithm of hydroxide activity.
At a specified temperature, pH plus pOH equals the corresponding pKw, which is approximately 14 at 25 °C.
For example, increasing hydroxide activity tenfold lowers pOH by one unit.
Polarizability
Chemistry · Definition
Polarizability describes how readily a particle's electron distribution can be distorted.
Larger, more diffuse electron clouds are often more polarizable, strengthening dispersion interactions when other factors are comparable.
For example, iodine is more polarizable than fluorine.
Polymer
Chemistry · Definition
A polymer is a large molecule built from many linked repeating units.
Chain length, branching, and interactions between chains influence its material properties.
For example, polyethylene contains long chains of repeating carbon-based units.
Polyprotic acid
Chemistry · Definition
A polyprotic acid can donate more than one proton per molecule in successive steps.
Each step has its own equilibrium constant, so the protons need not dissociate to the same extent.
For example, phosphoric acid can lose three protons through separate equilibria.
Pre-equilibrium approximation
Chemistry · Definition
The pre-equilibrium approximation treats a fast reversible mechanism step as close to equilibrium before a slower step proceeds.
Its equilibrium relation can replace an intermediate concentration in the proposed rate law.
For example, a fast association can express an intermediate's concentration in terms of the reactants that form it.
Precipitate
Chemistry · Definition
A precipitate is a solid phase that forms from species in a solution.
Whether it forms depends on the ion concentrations and the relevant equilibrium conditions.
For example, silver ions and chloride ions can produce a silver chloride precipitate.
Pure substance
Chemistry · Definition
A pure substance has one chemical identity and a fixed composition.
It may be an element or a compound, so containing more than one element does not make it a mixture.
For example, pure water contains hydrogen and oxygen in a fixed ratio.
Qualitative analysis
Chemistry · Definition
Qualitative analysis seeks the identities or presence of substances rather than their numerical amounts.
Individual observations often require supporting evidence because different species can produce similar signals.
For example, a supplied flame-color observation can suggest an ion's identity without establishing it uniquely.
Rate constant
Chemistry · Definition
A rate constant is the proportionality factor in a specified rate law.
Its value depends on conditions such as temperature, and its units depend on the law's overall order.
For example, a first-order rate constant has units of inverse time when concentration is measured per unit volume.
Rate law
Chemistry · Definition
A rate law relates a reaction rate to reactant concentrations and a rate constant under stated conditions.
For an overall reaction, its concentration exponents are established from evidence rather than copied from the balanced equation.
For example, doubling one reactant concentration doubles the rate when the reaction is first order in that reactant and other conditions are fixed.
Reaction half-life
Chemistry · Definition
A reaction half-life is the time needed for a reactant's concentration to fall to half its starting value under the specified model.
A first-order reaction has a half-life independent of its starting concentration, while other orders generally do not.
For example, consecutive equal time intervals halve a first-order reactant's concentration repeatedly.
Reaction intermediate
Chemistry · Definition
A reaction intermediate is produced in one mechanism step and consumed in another.
It cancels from the overall equation and differs from a catalyst, which is present before it is consumed and then regenerated.
For example, a short-lived species formed in step one and used in step two can be an intermediate.
Reaction mechanism
Chemistry · Definition
A reaction mechanism proposes a sequence of elementary steps that adds to the net reaction.
It must also be consistent with observed kinetic evidence, though agreement does not prove that the proposal is unique.
For example, an intermediate can form in one step and be consumed in a later step.
Reaction order
Chemistry · Definition
Reaction order describes the concentration dependence in a rate law.
An order in one reactant is its exponent, and overall order is the sum of the concentration exponents.
For example, a second-order dependence on one reactant makes quadrupling the rate follow from doubling that reactant's concentration under fixed conditions.
Reaction quotient
Chemistry · Definition
The reaction quotient describes a mixture's current composition using the equilibrium-expression form.
It uses species activities, approximated by concentrations or partial pressures in suitable models, with powers from balanced coefficients.
For example, a quotient below the equilibrium constant favors net forward reaction toward equilibrium.
Reaction rate
Chemistry · Definition
Reaction rate describes how quickly a reactant or product quantity changes with time.
Rates expressed using different species must account for their stoichiometric coefficients.
For example, a product's concentration–time slope can be used to determine its rate of formation.
Real-gas behavior
Chemistry · Definition
Real-gas behavior includes effects of particle attractions and particle volume that the ideal gas law neglects.
Those effects compete, so actual pressure can fall below or rise above the ideal prediction under different conditions.
For example, attraction can lower pressure while excluded particle volume can raise it under strong compression.
Redox reaction
Chemistry · Definition
A redox reaction combines oxidation of one species with reduction of another.
Electron loss and gain must balance in the net equation along with atoms and total charge.
For example, zinc reducing copper ions involves zinc oxidation and copper-ion reduction.
Reducing agent
Chemistry · Definition
A reducing agent causes another species to gain electrons.
The agent supplies those electrons and is itself oxidized.
For example, zinc acts as a reducing agent when it transfers electrons to copper ions.
Reduction
Chemistry · Definition
Reduction is electron gain, represented by a decrease in oxidation number.
It is paired with oxidation in an electron-transfer reaction.
For example, a copper ion with charge +2 gaining two electrons to form copper metal undergoes reduction.
Relative isotope abundance
Chemistry · Definition
Relative isotope abundance is the fraction or percentage of an element's atoms belonging to a particular isotope.
It concerns the number of atoms rather than their fraction of the sample's mass.
For example, an isotope present in 25 percent of the atoms contributes a factor of 0.25 to the weighted average atomic mass.
Resonance in bonding
Chemistry · Definition
Resonance uses multiple Lewis contributors with the same atom positions to describe delocalized electrons.
The actual species has one electronic structure rather than switching back and forth between the drawings.
For example, nitrate's equivalent N–O bonds are described with several resonance contributors.
Reversible reaction
Chemistry · Definition
A reversible reaction can proceed in both forward and reverse directions under the modeled conditions.
Its net direction depends on the current composition and temperature.
For example, a product-rich starting mixture can react backward toward equilibrium.
Salt bridge
Chemistry · Definition
A salt bridge provides an ionic conducting path between separated half-cells.
Ion movement helps prevent bulk charge buildup while electrons move through the external circuit.
For example, bridge anions can move toward a half-cell where metal cations are being produced.
Saturated solution
Chemistry · Definition
A saturated solution has the dissolved concentration corresponding to equilibrium with the undissolved phase under the stated conditions.
Microscopic dissolution and precipitation can continue at equal rates when the solid is present.
For example, additional solid may remain undissolved after a saturated salt solution is reached at fixed temperature.
Sigma bond
Chemistry · Definition
A sigma bond has bonding electron density along the line joining two nuclei.
It can arise from end-to-end overlap in a localized orbital model.
For example, the carbon–carbon bond in ethane is a sigma bond.
Solubility
Chemistry · Definition
Solubility describes how much of a substance can dissolve in a specified solvent under stated conditions.
It depends on the solute, solvent, temperature, and other relevant conditions and is distinct from how fast dissolution occurs.
For example, two salts can have different equilibrium solubilities in water at the same temperature.
Solubility product
Chemistry · Definition
A solubility product is an equilibrium constant for a specified dissolution reaction.
Its expression includes dissolved-species activities and omits the pure solid.
For example, dissolving one formula unit that releases two anions gives an anion activity squared in the expression.
Solute
Chemistry · Definition
A solute is a component dispersed in a solvent to form a solution.
It need not be a solid, and a solution need not be liquid.
For example, carbon dioxide is a gaseous solute in carbonated water.
Solvation
Chemistry · Definition
Solvation is the interaction of solvent particles with dispersed solute particles.
When the solvent is water, the process is called hydration.
For example, oriented water molecules form hydration shells around dissolved ions.
Solvent
Chemistry · Definition
A solvent is the component used as the dissolving medium in a solution.
It is often present in the greater amount, though the designation also depends on the mixture's context.
For example, water is the solvent in a dilute aqueous sugar solution.
Specific heat capacity
Chemistry · Definition
Specific heat capacity is the heat required per unit mass for a one-unit temperature increase under stated conditions.
It connects a mass and measured temperature change to heat transfer when the model applies.
For example, water requires about 4.18 joules per gram for a one-kelvin temperature rise near room temperature.
Spectator ion
Chemistry · Definition
A spectator ion appears unchanged on both sides of the complete ionic equation.
It is omitted from the net ionic equation because it does not undergo the represented reaction.
For example, sodium ions can be spectators when two other dissolved ions form a precipitate.
Spectroscopy
Chemistry · Definition
Spectroscopy studies interactions between matter and electromagnetic radiation.
The absorbed or emitted frequencies can reveal particular energy changes in a species.
For example, infrared absorption often corresponds to changes in molecular vibration.
Standard enthalpy of formation
Chemistry · Definition
Standard enthalpy of formation is the enthalpy change for forming one mole of a substance from its elements in their standard reference forms.
The substance's phase and specified temperature matter.
For example, an element in its standard reference form has zero standard formation enthalpy by convention.
Standard molar entropy
Chemistry · Definition
Standard molar entropy is the molar entropy of a substance in its specified standard state at a stated temperature.
It is an absolute-property value rather than an entropy of formation.
For example, an element can have a positive standard molar entropy even though its standard formation enthalpy is zero.
Standard state
Chemistry · Definition
A standard state is a defined reference for a substance's thermodynamic properties.
It specifies activities or reference forms and does not by itself fix the temperature at 25 °C.
For example, gases commonly use a one-bar reference pressure while the table states temperature separately.
Stationary phase
Chemistry · Definition
The stationary phase is the material that remains in place in a chromatographic system.
Components differ in how strongly they interact with it relative to the mobile phase.
For example, stronger stationary-phase retention can make one dye move less far than another under the same conditions.
Stoichiometry
Chemistry · Definition
Stoichiometry uses a balanced chemical equation to relate amounts of reactants and products.
Equation coefficients provide mole ratios rather than direct mass ratios.
For example, forming water from hydrogen and oxygen consumes two moles of hydrogen molecules per mole of oxygen molecules.
Strong acid or base
Chemistry · Definition
A strong acid transfers protons essentially completely to water in the applicable aqueous-solution model.
A dissolved strong hydroxide base dissociates into its constituent ions and supplies hydroxide; it does not have to react with water to generate that hydroxide.
Strength is distinct from concentration and, for a solid base, from how much can dissolve.
For example, dissolved sodium hydroxide supplies sodium and hydroxide ions even when the prepared solution is dilute.
Sublimation
Chemistry · Definition
Sublimation is a phase change directly from solid to gas without an intervening liquid phase.
Whether it occurs depends on the substance and pressure–temperature conditions.
For example, solid carbon dioxide sublimes at ordinary atmospheric pressure.
Thermal equilibrium
Chemistry · Definition
Thermal equilibrium means that bodies in thermal contact have the same temperature and no net heat transfer between them.
It does not mean that they contain equal amounts of energy or have the same mass.
For example, a small metal sample and a large water bath can reach the same temperature.
Thermodynamic favorability
Chemistry · Definition
Thermodynamic favorability describes whether a process has a driving force in the stated direction under specified conditions.
It does not determine how quickly the process occurs.
For example, a reaction with a favorable free-energy change can persist without visible progress when a large kinetic barrier blocks it.
Thiol
Chemistry · Definition
A thiol contains an S–H group attached to a carbon-containing structure.
Its bonding and properties differ from those of the corresponding alcohol, so replacing oxygen with sulfur does not preserve every interaction.
For example, ethanethiol is the sulfur-containing analogue of ethanol.
Titrant
Chemistry · Definition
A titrant is the reagent added in measured amounts during a titration.
Its concentration and reaction stoichiometry connect the delivered volume to the analyte amount.
For example, a standardized base solution can be the titrant for an unknown acid.
Titration
Chemistry · Definition
Titration measures an analyte by its reaction with a measured amount of another reagent.
The balanced reaction and the observed endpoint must be connected to the equivalence condition.
For example, an acid's amount can be inferred from a measured volume of a base with known concentration.
Titration endpoint
Chemistry · Definition
A titration endpoint is the observed signal used to stop or interpret a titration.
It approximates the equivalence point but can differ because of the chosen indicator or measurement conditions.
For example, an indicator's persistent color change may mark an endpoint near the steep region of a titration curve.
Transition state
Chemistry · Definition
A transition state is the high-energy configuration at a barrier along a reaction pathway.
It is not a stable intermediate that accumulates as a separately observable species in the basic model.
For example, the top of an elementary step's energy-profile peak represents its transition state.
Valence electron
Chemistry · Definition
Valence electrons are the electrons most directly involved in an atom's bonding and chemical behavior.
For main-group atoms, they occupy the highest occupied principal shell.
For example, neutral chlorine has seven valence electrons, helping explain its tendency to form a chloride ion.
Vaporization
Chemistry · Definition
Vaporization is a phase change from liquid to gas.
It requires energy to separate particles against their attractive interactions without ordinarily breaking the molecules' internal bonds.
For example, boiling liquid water produces water vapor while preserving the water molecules.
Vector
Physics 1 · Physics 2 · Chemistry · Definition
A vector is a quantity with a magnitude and a direction, represented by an arrow or by components along chosen axes.
Vectors add by adding corresponding components rather than simply adding their magnitudes.
For example, an eastward displacement followed by a northward displacement produces a resultant pointing northeast.
The zero vector has zero magnitude and no unique direction.
VSEPR
Chemistry · Definition
VSEPR is a model for predicting molecular geometry from electron domains around a central atom.
Bonding domains and lone pairs arrange to reduce repulsion, and a multiple bond counts as one domain for the basic geometry classification.
For example, water has four electron domains around oxygen but a bent molecular shape.
Weak acid or base
Chemistry · Definition
A weak acid or base reacts only partially with water and establishes an equilibrium.
Its equilibrium constant describes strength, while analytical concentration describes how much was prepared per volume.
For example, acetic acid leaves both acid molecules and its conjugate base in water.