Ch. 2 - Atomic Structure and Interatomic Bonding Flashcards
Atomic number (Z)
Number of protons in the nucleus.
Isotope
Atoms of the same element with a different number of neutrons, and therefore a different atomic masses.
Atomic weight
Weighted average of the atomic masses of an atom’s naturally occurring isotopes.
Atomic mass unit (amu)
Defined as 1/12 of the atomic mass of the most common isotope of carbon, carbon 12. Can be used in calculations of atomic weight.
Mole
6.023e23 atoms or molecules in one ____ of a substance.
Quantum mechanics
A set of principles and laws that govern systems of atomic and subatomic entities.
Bohr Atomic Model
Electrons are assumed to revolve around the atomic nucleus in discrete orbitals and the position of any particular electron is more or less well defined in terms of its orbital.
Wave mechanical model
The electron is considered to exhibit both wave-like and particle-like characteristics. The electron is no longer a particle moving in discrete orbital, rather position is considered to be the probability of an electron being at various locations around the nucleus. Electron cloud.
Replaced Bohr model.
Quantum numbers
Used to characterize every electron in an atom.
Principal quantum
n = 1, 2, 3, 4, 5…
Related to the distance of an electron from the nucleus, or its position.
Second quantum number
l = s, p, d, f
Related to the shape of the electron subshell.
Third quantum number
mL
Determines the number of energy states for each subshell.
For s, there is a single energy state. For p, there are 3 energy states. For d, there are five. For f, there are seven.
Fourth quantum number
mS
Spin moment associated with each electron. Must be oriented either up or down. +1/2 or -1/2
Electron states
Values of energy that are permitted for electrons.
Pauli exclusion principle
Each electron state can hold no more than two electrons, and they must have opposite spins. The s subshell can hold 2 electrons, p can hold 6, d can hold 10, and f can hold 14. Electrons fill up the lowest possible energy states in the electron shells and subshells, two per state.
Ground state
When all electrons occupy the lowest possible energies.
Electron configuration
Represents the manner in which the energy states of an atom are occupied.
Valence electrons
Electrons that occupy the outermost shell. These electrons participate in bonding between atoms to form atomic and molecular aggregates.
Stable electron configuration
When the valence electron shell is completely filled. An atom can achieve this state based on its default electron configuration or by gaining or losing electrons to form charged ions, or by sharing electrons with other atoms.
Noble gases
Elements with their outermost shell full. They are inert (unreactive).
Electropositive
Elements that are capable of giving up their few valence electrons to become positively charged ions. Mostly metals.
Electronegative
Atoms situated in the right side of the periodic table will readily accept electrons to form negatively charged ions, or sometimes share electrons with other atoms.
This property increases from left to right and from bottom to top.
Bonding forces
The magnitude of the attractive force varies with distance. When the valence shells to two atoms begin to overlap, strong repulsive forces come into play.
Bonding energy
A state of equilibrium exists when the centers of the two atoms are separated by ro, which is approximately 0.3nm for most atoms. Once in this position, it will take a lot of energy to separate the atoms or push them closer together. The energy at this point is known as the bonding energy. This is the minimum point in the potential energy curve.
Primary bond
Ionic
Covalent
Metallic
Ionic bonding
Type of bonding found in compounds composed of both metallic and nonmetallic elements (element in horizontal extremes of table). Metallic atoms easily give up their valence electrons to nonmetallic elements. All atoms then acquire a stable/inert gas configuration. Atoms also become ions (ie gain an electrical charge).
NONDIRECTIONAL
Common in ceramics.
Relatively large bonding energies.
Nondirectional
The magnitude of the bond is equal in all directions around an ion.
Covalent bonding
Type of bonding in which stable electron configurations are assumed by sharing electrons between adjacent atoms. Two atoms bonded in this way will each contribute at least one electron to the bond. The electrons may be considered to belong to both atoms.
DIRECTIONAL, such that the bond is between specific atoms and may only exist in the direction between one atom and another that participated in electron sharing.
Common in nonmetallic molecules (F2, H2), molecules containing dissimilar atoms (H2O, HNO3), and elemental solids (diamond).
Bonding energies can be strong (diamond) or weak (bismuth).
Polymers typically have this bond.
Metallic bonding
Type of bonding found in metals and their alloys. Valence electrons are not bound to any particular atom within the solid, but rather are free to drift throughout the entire metal. “Sea of electrons”. Nonvalence electrons form ion cores. Free electrons act as glue to hold ion codes together.
NONDIRECTIONAL.
Bonding energies can be weak or strong.
Sea of electrons is what gives metals good conductivity.
Secondary bonds (van der Waals bonds)
Weak bonds in comparison to primary bonds. Exists between virtually all molecules or atoms. Arises from atomic or molecular dipoles. Bonding results from coulombic attraction between the positive end of one dipole and the negative region of an adjacent one.
Fluctuating induced dipoles
Polar molecule-induced dipole bonds
Permanent dipole bonds
Dipole
Exists whenever there is some separation of positive and negative positions of an atom or molecule.
Hydrogen bonding (permanent dipole bond)
Special type of secondary bonding, found to exist between some molecules that have hydrogen as one of the constituents. Bonded hydrogen acts as a bare proton. Highly positive and therefore capable of strong attractive forces with the negative end of an adjacent molecule.
This type of bonding is what causes ice to be less dense than water.
Induced dipole bonds
A dipole can produce a displacement of the electron distribution of an adjacent molecule or atom, which indices the second one also to become a dipole that is then weakly attracted or bonded to the first.
Type of van der Waals bonding.
Polar molecule-induced dipole bonding
A polar molecule is a molecule with a permanent dipole moment that exists by virtue of an asymmetrical arrangement of positively and negatively charged regions.
This type of molecule can induce dipoles in adjacent nonpolar molecules and a bind will form as a result of the attractive forces between the two molecules. This bond is stronger than those in fluctuating induced dipoles.
