Unit 1 - Structure and Properties of Matter Flashcards
Electromagnetic radiation
Radiant energy that exhibits wavelike behavior and travels through space at the speed of light in a vacuum.
Wavelength
The distance between two consecutive peaks of troughs in a wave.
Frequency
The number of waves (cycles) per second that pass a given point in space.
Planck’s constant
The constant relating the change in energy for a system to the frequency of the electromagnetic radiation absorbed or emitted; equal to 6.626 x 10^-34 Js.
Quantization
The concept that energy can occur only in discrete units called quanta.
Ground state
The lowest possible energy state of an atom or molecule.
Dual nature of light
The statement that light exhibits both wave and particulate properties.
Diffraction
The scattering of light from a regular array of points or lines, producing constructive and destructive interference.
Continuous spectrum
A spectrum that exhibits all the wavelengths of visible light.
Line spectrum
A spectrum showing only certain discrete wavelengths.
Standing wave
A stationary wave as on a string of a musical instrument; in the wave mechanical model, the electron in the hydrogen atom is considered to be a standing wave.
Wave function
A function of the coordinates of an electron;s position in 3D space that describes the properties of the electron.
Orbital
A specific wave function for an electron in an atom. The square of this function gives the probability distribution for the electron.
Heisenberg uncertainty principle
A principle stating that there is a fundamental limitation to how precisely both the position and momentum of a particle can be known at a given time.
Probability distribution
The square of the wave function indicating the probability of finding an electron at a particular point in space.
Principal quantum numbers (n)
The quantum number relating to the size and energy of an orbital; it can have any positive integer value.
Angular momentum number (l)
The quantum number relating to the shape of an atomic orbital, which can assume any integral value from 0 to n-1 for each value of n.
Magnetic quantum number (ml)
The quantum number relating to the orientation of an orbital in space relative to the other orbitals with the same l quantum number. It can have integral values between l and -l, including zero.
Subshell
A set of orbitals with a given azimuthal quantum number.
Node
An area of an orbital having zero electron probability.
Degenerate orbitals
A group of orbitals with the same energy.
Electron spin quantum number (ms)
A quantum number representing one of the two possible values for the electron spin; either +1/2 or -1/2.
Pauli exclusion principle
In a given atom no two electrons can have the same set of four quantum numbers.
Ionization energy
The minimum energy required to remove the most loosely held electron from an atom or an ion at its gaseous, ground state.
Electron affinity
The energy change associated with the addition of an electron to a gaseous atom.
Bond energy
The energy required to break a given chemical bond.
Ionic bonding
The electrostatic attraction between oppositely charged ions.
Ionic compound
A compound that results when a metal reacts with a nonmetal to form a cation and an anion.
Coulomb’s law
E = 2.31 x 10^-19 * [(Q1Q2)/r], where E is the energy of interaction between a pair of ions, expressed in joules; r is the distance between the ion centers in nm; and Q1 and Q2 are the numerical ion charges.
Bond length
The distance between the nuclei of the two atoms connected by a bond; the distance where the total energy of a diatomic molecule is minimal.
Covalent bonding
A type of bonding in which electrons are shared by atoms.
Polar covalent bonding
A covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other.
Electronegativity
The tendency of an atom in a molecule to attract shared electrons to itself.
Dipole moment
A property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
Isoelectronic ions
Ions containing the same number of electrons.
Lattice energy
The energy change occurring when separated gaseous ions are packed together to form an ionic solid.
Localized electron (LE) model
A model that assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Valence shell electron-pair repulsion (VESPR) model
A model whose main postulate is that the structure around a given atom in a molecule is determined principal by minimizing electron-pair repulsions.
Hybridization
A mixing of the native orbitals on a given atom to form special atomic orbitals for bonding.
Hybrid orbitals
A set of atomic orbitals adopted by an atom in a molecule different from those of the atom in the free state.
Condensed states of matter
Liquids and solids.
Intermolecular forces
Relatively weak interactions that occur between molecules.
Dipole-dipole attraction
The attractive force resulting when polar molecules line up so that the positive and negative ends are close to each other.
Hydrogen bonding
Unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly electronegative atom.
London dispersion forces
The forces, existing among noble gas atoms and nonpolar molecules, that involve an accidental dipole that induces a momentary dipole in a neighbor.
Aufbau principle
The principle stating that as protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals.
Hund’s rule
The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli exclusion principle in a particular set of degenerate orbitals, with all unpaired electrons having parallel spins.
Intramolecular forces
The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as intramolecular attractions.
