Glossary
Anion
A negatively charged ion, formed when an atom gains one or more electrons.
Example:
When a chlorine atom gains an electron, it forms a chloride anion (Cl⁻).
Atom
The smallest unit of an element that retains the chemical identity of that element. Atoms are composed of protons, neutrons, and electrons.
Example:
A single atom of gold is still gold, even though it's too small to see.
Atomic radius
A measure of the size of an atom, typically defined as half the distance between the nuclei of two identical atoms that are bonded together.
Example:
As you move down a group on the periodic table, the atomic radius generally increases due to the addition of new electron shells.
Aufbau principle
States that electrons fill atomic orbitals of the lowest available energy levels before occupying higher energy levels.
Example:
According to the Aufbau principle, electrons will fill the 1s orbital before moving to the 2s orbital.
Average atomic mass
The weighted average of the masses of all naturally occurring isotopes of an element. It is calculated based on the mass and relative abundance of each isotope.
Example:
The average atomic mass of chlorine (35.45 amu) is closer to 35 than 37 because the isotope chlorine-35 is more abundant than chlorine-37.
Avogadro's number
The number of constituent particles (atoms, molecules, ions, etc.) per mole of a substance, approximately 6.022 x 10^23 mol⁻¹.
Example:
There are Avogadro's number of carbon atoms in exactly 12 grams of carbon-12.
Bohr model
A model of the atom where electrons orbit the nucleus in discrete energy levels or shells. Electrons can jump between these levels by absorbing or emitting specific amounts of energy.
Example:
The Bohr model successfully explained the line spectrum of hydrogen by proposing that electrons exist in quantized energy states.
Cation
A positively charged ion, formed when an atom loses one or more electrons.
Example:
When a sodium atom loses an electron, it forms a sodium cation (Na⁺).
Chromatography
A laboratory technique for separating mixtures. The mixture is dissolved in a fluid (mobile phase) and passed through a structure holding another material (stationary phase), separating components based on differential partitioning.
Example:
Forensic scientists use chromatography to separate and identify different pigments in a pen ink sample.
Core electrons
The electrons in an atom that are not valence electrons; they are located in the inner shells and are not typically involved in chemical bonding.
Example:
In a sodium atom, the 1s², 2s², and 2p⁶ electrons are core electrons, while the 3s¹ electron is a valence electron.
Coulomb's law
States that the force between two charged particles is directly proportional to the product of their charges and inversely proportional to the square of the distance between them.
Example:
The strong attraction between a positively charged proton and a negatively charged electron in an atom is explained by Coulomb's law.
Covalent bonds
Chemical bonds formed by the sharing of electrons between two atoms. They typically form between two nonmetal atoms.
Example:
In a water molecule (H₂O), hydrogen and oxygen atoms are connected by covalent bonds as they share electrons.
Dalton's theory
An early atomic theory proposing that all matter is composed of indivisible atoms, atoms of a given element are identical, and chemical reactions involve the rearrangement of atoms.
Example:
Dalton's theory laid the foundation for modern chemistry by proposing that elements are made of discrete units called atoms.
Dimensional analysis
A problem-solving method that uses the fact that any number or expression can be multiplied by one without changing its value. It is used for converting units.
Example:
To convert 5 kilometers to meters, you would use dimensional analysis by multiplying 5 km by the conversion factor (1000 m / 1 km).
Distillation
A separation technique that purifies liquids by heating them to create vapor, then cooling the vapor to condense it back into a liquid, separating components based on boiling points.
Example:
To separate alcohol from water, a chemist might use distillation to take advantage of their different boiling points.
Effective nuclear charge
The net positive charge experienced by an electron in a multi-electron atom. It is influenced by the actual nuclear charge and the shielding effect of inner electrons.
Example:
Valence electrons experience a lower effective nuclear charge than core electrons because the core electrons shield them from the full positive charge of the nucleus.
Electron affinity
The energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion. It can be positive or negative.
Example:
Chlorine has a high electron affinity because it readily accepts an electron to achieve a stable noble gas configuration, releasing energy.
Electron configuration
The arrangement of electrons in an atom's orbitals. It describes how electrons are distributed among the various atomic orbitals.
Example:
The electron configuration of oxygen is 1s²2s²2p⁴, indicating how its 8 electrons are distributed.
Electron orbitals
Regions around the nucleus of an atom where electrons are most likely to be found. Each orbital can hold a maximum of two electrons with opposite spins.
Example:
The 1s electron orbital is a spherical region closest to the nucleus, capable of holding two electrons.
Electron-electron repulsion
The repulsive force between electrons in an atom due to their like negative charges. This repulsion influences electron shielding and orbital energies.
Example:
The electron-electron repulsion between electrons in the same orbital helps explain why they must have opposite spins.
Electronegativity
A measure of the tendency of an atom to attract a bonding pair of electrons in a chemical bond. It generally increases across a period and decreases down a group.
Example:
Fluorine has the highest electronegativity, meaning it strongly pulls shared electrons towards itself in a bond.
Electrons
Negatively charged subatomic particles that orbit the nucleus of an atom. They are involved in chemical bonding and determine an atom's chemical properties.
Example:
The flow of electrons through a wire creates an electric current.
Empirical formula
The simplest whole-number ratio of atoms in a compound. It is determined from experimental data.
Example:
The molecular formula for glucose is C₆H₁₂O₆, but its empirical formula is CH₂O, representing the simplest ratio of atoms.
Filtration
A separation technique used to separate insoluble solid particles from a liquid or gas by passing the mixture through a filter medium.
Example:
When making coffee, the coffee grounds are separated from the liquid coffee by filtration.
Formula units
The empirical formula of an ionic compound, representing the simplest whole-number ratio of ions in the compound.
Example:
For sodium chloride, NaCl, one formula unit represents one sodium ion and one chloride ion.
Gas
A state of matter characterized by both variable shape and variable volume. Particles in a gas are far apart and move randomly and rapidly.
Example:
The air inside a balloon is a gas; it expands to fill the balloon and has no fixed shape.
Heterogeneous mixtures
Mixtures that have a non-uniform composition, where the components are not evenly distributed and can often be visibly distinguished.
Example:
A bowl of cereal with milk is a heterogeneous mixture because you can clearly see the individual cereal pieces and the milk.
Homogeneous mixtures
Mixtures that have a uniform composition and appearance throughout. The components are evenly distributed and indistinguishable.
Example:
Sweet tea is a homogeneous mixture because the sugar is completely dissolved and evenly distributed, making every sip taste the same.
Hund's rule
States that when electrons occupy degenerate orbitals (orbitals of the same energy), they will first fill each orbital singly with parallel spins before pairing up.
Example:
When filling the 2p orbitals, Hund's rule dictates that each of the three 2p orbitals will receive one electron before any orbital gets a second electron.
Ionic bonds
Chemical bonds formed by the electrostatic attraction between oppositely charged ions. They typically form between a metal and a nonmetal through the transfer of electrons.
Example:
Sodium chloride (NaCl) is held together by ionic bonds where sodium donates an electron to chlorine, forming Na⁺ and Cl⁻ ions.
Ionic radius
The radius of an ion. Cations are smaller than their parent atoms, while anions are larger.
Example:
When a sodium atom loses an electron to form a cation, its ionic radius becomes significantly smaller than its atomic radius.
Ionization energy
The minimum energy required to remove one electron from a neutral atom in its gaseous state. Successive ionization energies increase as more electrons are removed.
Example:
It takes a relatively low ionization energy to remove the single valence electron from a sodium atom, making it reactive.
Isotopes
Atoms of the same element that have the same number of protons but different numbers of neutrons. This results in different mass numbers for the atoms.
Example:
Carbon-12 and Carbon-14 are isotopes of carbon; both have 6 protons, but Carbon-14 has 8 neutrons while Carbon-12 has 6.
Law of definite proportions
States that a given chemical compound always contains its component elements in fixed ratios by mass, regardless of the source or method of preparation.
Example:
No matter where you find it, water will always be composed of hydrogen and oxygen in a 1:8 mass ratio, illustrating the law of definite proportions.
Liquid
A state of matter characterized by a fixed volume but a variable shape. Particles in a liquid are close together but can move past one another.
Example:
Water in a glass is a liquid; it takes the shape of the glass but maintains its volume.
Mass spectroscopy
An analytical technique that measures the mass-to-charge ratio of ions. It is used to determine the elemental composition of a sample or the masses of particles and molecules.
Example:
Scientists use mass spectroscopy to identify unknown compounds in a forensic sample by analyzing their unique fragmentation patterns.
Matter
Anything that has mass and occupies space. It exists in various states and forms the physical world around us.
Example:
A block of ice, a cloud of steam, and the air you breathe are all examples of matter.
Metalloids
Elements that have properties intermediate between those of metals and nonmetals. They are often semiconductors.
Example:
Silicon is a metalloid that is crucial in the electronics industry due to its semiconducting properties.
Metals
Elements that are typically good conductors of heat and electricity, malleable, ductile, and tend to form positive ions (cations).
Example:
Copper is a metal widely used in electrical wiring due to its excellent conductivity.
Mixtures
Combinations of two or more substances that are not chemically bonded together. They can be separated by physical means.
Example:
Salt dissolved in water forms a mixture because the salt and water molecules retain their individual identities and can be separated by evaporation.
Molar mass
The mass in grams of one mole of a substance. It is numerically equal to the atomic mass (for elements) or formula mass (for compounds) in atomic mass units (amu).
Example:
The molar mass of oxygen gas (O₂) is approximately 32.00 g/mol, meaning 32 grams of O₂ contains one mole of O₂ molecules.
Mole
A unit of measurement used in chemistry to express amounts of a chemical substance. One mole is defined as exactly 6.022 x 10^23 entities (atoms, molecules, ions, etc.).
Example:
If you have a mole of water molecules, you have enough water to fill a small swimming pool with individual molecules.
Neutrons
Neutral (no charge) subatomic particles found in the nucleus of an atom. They contribute to the atom's mass but not its charge.
Example:
Deuterium, an isotope of hydrogen, has one neutron in its nucleus in addition to its single proton.
Nonmetals
Elements that are generally poor conductors of heat and electricity, brittle, and tend to form negative ions (anions) or share electrons.
Example:
Sulfur is a nonmetal that exists as a yellow, brittle solid at room temperature and is a poor conductor.
Nonpolar covalent bonds
Covalent bonds in which electrons are shared equally between two atoms because they have similar or identical electronegativities.
Example:
The bond between two oxygen atoms in an O₂ molecule is a nonpolar covalent bond as both atoms have equal pull on the shared electrons.
Nuclear charge
The total positive charge of the atomic nucleus, determined by the number of protons. It is the primary attractive force on electrons in an atom.
Example:
As you move across a period, the increasing nuclear charge pulls the electrons closer to the nucleus, leading to a decrease in atomic radius.
Nucleus
The dense, positively charged center of an atom, composed of protons and neutrons. It contains most of the atom's mass.
Example:
The nucleus of a uranium atom is unstable and can undergo radioactive decay.
Pauli exclusion principle
States that no two electrons in the same atom can have the same set of four quantum numbers. This means an atomic orbital can hold a maximum of two electrons, and they must have opposite spins.
Example:
The Pauli exclusion principle explains why a single orbital can only accommodate two electrons, one spinning up and one spinning down.
Periodic table
A tabular arrangement of chemical elements, organized by atomic number, electron configuration, and recurring chemical properties. It allows for the prediction of element behavior.
Example:
By looking at the periodic table, you can quickly identify that elements in the same column, like fluorine and chlorine, share similar chemical properties.
Photoelectric effect
The emission of electrons when electromagnetic radiation, such as light, hits a material. The energy of the light must exceed a certain threshold for electrons to be emitted.
Example:
Solar panels convert sunlight into electricity by utilizing the photoelectric effect to release electrons from a semiconductor material.
Photoelectron spectroscopy (PES)
An experimental technique that measures the kinetic energy of electrons emitted from a material when it is exposed to X-rays or UV light. It provides information about electron binding energies and electron configurations.
Example:
A Photoelectron spectroscopy spectrum can show distinct peaks corresponding to the core and valence electrons in an atom, revealing their energy levels.
Polar covalent bonds
Covalent bonds in which electrons are shared unequally between two atoms due to a difference in electronegativity, creating partial positive and negative charges.
Example:
The bond between oxygen and hydrogen in a water molecule is a polar covalent bond because oxygen attracts the shared electrons more strongly.
Polarity
A property of a molecule where there is an unequal distribution of charge, resulting in a partial positive end and a partial negative end.
Example:
Water is a polar molecule because oxygen attracts electrons more strongly than hydrogen, creating a slight negative charge on oxygen and slight positive charges on hydrogen.
Protons
Positively charged subatomic particles found in the nucleus of an atom. The number of protons determines the atomic number and the identity of an element.
Example:
Every hydrogen atom has exactly one proton in its nucleus.
Pure substances
Matter that has a uniform and definite composition. This includes elements and compounds.
Example:
Distilled water is a pure substance because it consists only of H₂O molecules with a fixed composition.
Solid
A state of matter characterized by a fixed shape and fixed volume. Particles in a solid are tightly packed and vibrate in fixed positions.
Example:
An ice cube is a solid because it maintains its shape and volume regardless of the container it's in.
Valence electrons
The electrons in the outermost shell of an atom. These electrons are primarily involved in chemical bonding and determine an atom's chemical properties.
Example:
Carbon has four valence electrons, which allows it to form four bonds and create diverse organic compounds.