Glossary of chem test 5 chap 9 and 10
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- Valence electrons
- are the outer shell electrons of an
atom. The valence electrons are the electrons that
participate in chemical bonding. group number= #of valence e-
- Electron configurations of Ions of the Representative Elements
- These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital.
Electron configurations can predict stable ion formation
- covalent bond
- is a chemical bond in which two or more electrons are shared by two atoms.
- lengths of covalent bonds
- Triple bond < Double Bond < Single Bond; go in order of atomic radii size (dec right, inc down)
- In many cases the electrons are shared, but not equally.
These bonds are called polar-covalent and are considered to be partially positively charged and partially negatively charged. EX: HCl
- the relative ability of a bonded atom to
attract the shared electrons in a chemical bond; inc right, dec down like bond energy;
-will most readily GAIN an e-
-most negative e- affinity
-greatest attraction for an e- in a covalent bond
-most polar; highest diploe
- formal charge
- the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure; formula: FC= (#valence e-)-(#nonbonding e-) - 1/2(#bonding e-); it is the charge an atom would have if the bonding electrons were shared equally; The formal charges must add up to equal the actual charge on
the species; the lowest one is correct although zeros overide negatives; IF THE SING ON THE MOLECULE IS NEGATIVE AND 2 ANS YIELD THE SAME RESULTS, YOU LOOK FOR THE MOST ELECTRONEG OF THE TWO. IF SIGN OF MOLECULE IS POS, YOU LOOK FOR LEAST ELECTRONEG OF THE TWO
- the use of two or more Lewis
structures to represent a particular molecule
or ion; connected by a
double-headed arrow. The true structure is a
blend of the resonance structures; Resonance indicates that electron delocalization occurs in the molecule; one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.
- Are covalent bonds weaker than ionic bonds?
- In covalent bonding, electron sharing leads to strong, localized bonds. But most compounds with covalent bonds are gases, liquids, or low-melting solids. Ionic compound are high-melting solids; Within a molecule there are strong intramolecular bonds between two atoms; Between molecules there are weak intermolecular forces;
The latter are what determine melting and boiling points.
- bond energy
- The enthalpy change required to break a particular bond in one mole of gaseous molecules; Single bond < Double bond < Triple bond; same as electronegativity (inc right, dec down)
- the relationship between bond length, bond strength, and bond energy is :
- The shorter the bond length,
the greater the bond strength,
the greater the bond energy!
The longer the bond length,
the lower the bond strength,
the lower the bond energy!
- Valence Shell Electron-Pair Repulsion
- Each group of valence electrons around a
central atom is located as far away from the
others as possible in order to minimize
An electron group may be a single, double,
or triple bond or a lone pair.
- molecular shape
- is defined by the relative positions of the atomic nuclei. If there are no lone pairs of electrons the molecular shape is the same as the electron-group arrangement.
- Valence Bond Theory
- a covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons
occupies the region between the nuclei.
Usually this means that each bonding orbital
should contain one electron. The electrons
must have opposite spins; The bond strength depends on the attraction of nuclei for
the shared electrons, so the greater the orbital overlap,
the stronger the bond.
- is the mixing of atomic orbitals in an atom to generate a set of new orbitals called hybrid orbitals; The total number of orbitals doesn’t change.
Covalent bonds are formed by:
Overlap of hybrid orbitals with atomic orbitals or
Overlap of hybrid orbitals with other hybrid orbitals
- sigle bond
- sigma bond
- dbl bond
- sigma + pi
- triple bond
- sigma + 2pi
- Delocalized electrons
- are not confined between two adjacent bonding atoms, but actually extend over three or more atoms; Structures with delocalized electrons always have greater stability than similar structures without delocalized electrons; resonance indicates that this occurs.
- Molecular orbital theory
- bonds are formed from interaction of atomic orbitals to form molecular orbitals.
- bonding molecular orbital
- has lower energy and greater stability than the atomic orbitals from which it was formed
- antibonding molecular orbital
- has higher energy and lower stability than the atomic orbitals from which it was formed
- distance between elements in the central atom plane. for trigonal bipyramidal, its 120. for octahedral, its 90
- distance from atom in central atom plane, to a point above or below that plane. in trigonal bipyramidal, its 90. in octahedral, its 90.
- the higher the bond order, the more stable the species. formula: 1/2(bonding e- - antibonding e-); energy inc as you go up. the top one (antibonding one w/ the star) has more energy and less stability than the lower one.
- exceptions to octet rule
- S, P, B, Be, Xe, N
- when you dont know what the central atom is...
- make it carbon
- atomic radiius
- (determines lenghts of covalent bonds)
-most readily lose and e-
-least affinity for e-
- intramolecular bonds
- Within a molecule (between 2 atoms) there are strong intra bonds.
- Between molecules there are weak intermolecular forces.
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