Concepts (Mid Term II) Flashcards
Z Effective (Zeff)
The net positive charge pulling e- towards the nucleus. The stronger the pull on the outermost e-, the higher the Zeff.
Isoelectric Series
A groups of ions that have the same number of e-
Electron Affinity (EA)
The energy released when a gaseous atom gains an e- to form an anion. DeltaE = EA
Inverse of IE
Ionization Energy (IE)
The amount of energy required (added) to remove an e- from an isolated atom or molecule (gaseous)
Eel = k(-1)Zeff / d
Inverse of EA
Lattice Energy
The amount of energy released when gas-phase ions combine to form a mole of a crystalline solid smaller distance = larger charge
larger distance = smaller charge
Eel = K ((Q+ * Q-)/d)
Electronegativity
An atom’s tendency to attract electrons towards itself within a covalent bond
Small radius, a large Zeff = a lot of electronegativity
Photoelectric Spectroscopy
an experimental technique that measures the relative energies of electrons in atoms and molecules.
Electrostatic potential energy (Eel)
the energy a charged particle has b/c of its position relative to another charged particle
DIRECTLY proportional to the product of the charges of the particle and INVERSLEY proportional to the distance between them
Also called coulumbic attraction
Binary Molecular Compounds
Two non mentals
Binary Ionic Compounds
Metal and nonmetal
Polyatomic Ion
Molecular ions composed of two or more atoms bonded by covalent bonds and acting as a single unit, but unlike molecules, they have a net charge on them
Oxoacids
A compound composed of oxoanions bonded to H+ ions
Binary Acids
Hydrogen combined with a nonmetal element
Allotropes
Different molecular forms of the same element
Resonance
A characteristic of e- distribution in which two or more equivalent Lewis structures can be drawn for one compound
Electron Pair Delocalization
The spreading out of e- density over several atoms
reduces the e- potential energy and lowers the energy of the molecule –> resonance stabilization
Bond Length
the distance between the centers (nuclei) of two covalently bonded atoms
Bond Order
The number of bonds between atoms; 1 for a single bond, 2 for a double bond, 3 for a triple bond
For resonance structures, take an average
Bond Energy/Bond Strength
the amount of energy needed to break a bond (same number as energy needed to make the bond)
As the bond order increases, the bond strength increases as atoms in bonding come closer to each other while bond length decreases.
Formal Charge
The value calculated for an atom in a molecule or polyatomic ion by determining the difference between the number of valence e- in the free atom and the sum of the lone-pair e- plus half the e- in the atom’s bonding pair
Free Radical
an atom, ion, or molecule with unpaired valence e-
Generally makes things very reactive
Expanded Octets
When atoms may have more than 8 e- associated with them. Only applicable to 3rd row elements
Bond Angle
The angle (in degrees) shown by lines joining the centers of two atoms to a third atom to which they are chemically bonded
VSEPR
Valence Shell Electron Pair Repulsion Theory
A model predicting the arrangement of valence e- lairs around a central atom that minimizes their mutual repulsion to produce the lowest energy orientation
Based on fundamental chemical principle that e- have neg charges and repel each other
Steric Number
The sum of the number of atoms bonded to a central atom plus the number of lone pairs of e- on the central atom
Bond Dipole
Separation of electrical charge created when atoms w different electronegativities form a covalent bond
An unequal distribution of e- between two atoms produces a partial neg charge on one end of the bon and a partial positive charge on the other
Permanent Dipole
Permanent separation of electrical charge in a molecule resulting from unequal distributions of bonding and/or lone pairs of e-
Dipole Moment
a measure of the degree to which a molecule aligns itself in an applied electric field, a quantitative expression of the polarity of a molecule
Valence Bond Theory
A quantum mechanics-based theory of bonding that assumes covalent bonds from when half-filled orbitals on different atoms overall or occupy the same region in space
A chemical bond forms when the atomic orbitals of two atoms overlap
Sigma Bond
A covalent bond in which the highest e- density is between the two atoms along the internuclear axis
Single bonds
Pi Bond
A covalent bond in which e- density is greatest above and below the internuclear axis
double bond
Hybridization
In valence bond theory, the mixing of atomic orbitals to generate new sets of orbitals that may form sigma bonds with other atoms
Hybrid Atomic Orbital
In valence bond theory, one set of equivalent orbitals about an atom created when specific atomic orbitals are mixed
