CSET Science: Chemistry flashcards
CSET Science: Chemistry · Build recall with a focused review session.
32 cards. Read the term and the definition here. The deck at the top of this page is optional.
- Atomic number
- The number of protons in an atom's nucleus; it uniquely identifies the element and, in a neutral atom, equals the number of electrons.
- Isotope
- An atom of the same element with a different number of neutrons, giving a different mass number but nearly identical chemical properties.
- Average atomic mass
- The weighted average of the masses of an element's naturally occurring isotopes, weighted by their fractional abundances; reported in atomic mass units.
- Aufbau principle
- Electrons fill atomic orbitals in order of increasing energy, beginning with 1s and proceeding through 2s, 2p, 3s, 3p, 4s, 3d and beyond.
- Hund's rule
- Within a subshell, electrons occupy separate orbitals with parallel spins before any orbital receives a second electron, minimizing electron-electron repulsion.
- Valence electrons
- The electrons in an atom's outermost principal energy level; they determine bonding behavior and group-related chemical properties.
- Ionization energy
- The energy required to remove one electron from a gaseous atom or ion; it generally increases across a period and decreases down a group.
- Electronegativity
- A measure of an atom's ability to attract shared electrons in a bond; fluorine is highest, and values increase across a period and decrease down a group.
- Atomic radius trend
- Atomic radius decreases left to right across a period as effective nuclear charge rises, and increases down a group as additional shells are added.
- Ionic bond
- An electrostatic attraction between oppositely charged ions, typically formed when a metal transfers electrons to a nonmetal with a large electronegativity difference.
- Covalent bond
- A bond formed when two atoms share one or more pairs of electrons; it may be nonpolar or polar depending on the electronegativity difference.
- Metallic bonding
- Bonding in metals in which positively charged lattice ions are held together by a sea of delocalized valence electrons, explaining conductivity and malleability.
- VSEPR theory
- A model predicting molecular shape by arranging electron domains around a central atom as far apart as possible; lone pairs repel more than bonding pairs.
- Polarity of molecules
- A molecule is polar when its bond dipoles do not cancel; symmetric geometries such as linear CO2 or tetrahedral CH4 are nonpolar despite polar bonds.
- Intermolecular forces
- Attractions between molecules, including London dispersion, dipole-dipole, and hydrogen bonding; they are weaker than covalent bonds but govern boiling point and solubility.
- Hydrogen bonding
- A strong dipole attraction between a hydrogen bonded to N, O, or F and a lone pair on a nearby N, O, or F atom; it raises boiling points of water and alcohols.
- Mole concept
- One mole contains 6.022e23 representative particles (Avogadro's number); the molar mass of a substance is its mass in grams per mole.
- Molarity
- Concentration expressed as moles of solute divided by liters of solution; dilution obeys M1V1 = M2V2 because moles of solute are conserved.
- Limiting reactant
- The reactant that is completely consumed first and determines the maximum product formed; found by comparing mole-to-coefficient ratios.
- Percent yield
- The actual product mass divided by the theoretical mass predicted from stoichiometry, multiplied by 100; it reflects reaction efficiency and losses.
- Empirical formula
- The simplest whole-number ratio of atoms in a compound; it is derived by converting percent composition to moles and dividing by the smallest value.
- Ideal gas law
- PV = nRT relates pressure, volume, moles, and absolute temperature; R is 0.08206 L atm/(mol K) or 8.314 J/(mol K).
- Combined gas law
- P1V1/T1 = P2V2/T2 applies to a fixed amount of gas; temperature must be expressed in kelvin in all gas calculations.
- Dalton's law
- The total pressure of a gas mixture equals the sum of the partial pressures of its component gases, each behaving independently in the container.
- Colligative properties
- Properties depending on the number of dissolved particles rather than identity: vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
- Acid and base definitions
- Arrhenius acids yield H+ and bases yield OH- in water; Bronsted-Lowry acids donate protons and bases accept them; Lewis acids accept electron pairs.
- pH scale
- pH = -log[H3O+]; values below 7 are acidic, 7 is neutral at 25 C, and above 7 is basic; each unit represents a tenfold change in hydronium concentration.
- Strong versus weak acids
- Strong acids ionize completely in water, while weak acids ionize only partially and establish an equilibrium described by Ka; smaller pKa means stronger acid.
- Le Chatelier's principle
- A system at equilibrium shifts to partially offset a change in concentration, pressure, or temperature; catalysts speed attainment but do not shift the equilibrium position.
- Equilibrium constant
- K expresses the ratio of product to reactant activities at equilibrium; K greater than 1 favors products, and K depends on temperature only.
- Oxidation and reduction
- Oxidation is loss of electrons and increase in oxidation number; reduction is gain of electrons and decrease in oxidation number; they occur together in redox reactions.
- Endothermic versus exothermic
- Endothermic reactions absorb heat and have positive enthalpy change; exothermic reactions release heat and have negative enthalpy change, often accompanied by bond formation.