Chemical Bonding Flashcards
The * refers to a description of chemical bonding through Lewis symbols and Lewis structures in accordance with a particular set of rules.
1. Electrons, especially those of the *, play a
fundamental role in chemical bonding.
2. In some cases, electrons are * from one atom to another. Positive and negative ions are formed and attract each other through *
called * bonds.
Lewis theory
outermost (valence) electronic shell
transferred
electrostatic forces, ionic
Lewis theory:
3. In other cases, one or more pairs of electrons are shared between atoms. A bond
formed by the * of electrons between atoms is called a * bond.
4. Electrons are transferred or shared in such a way that each atom acquires an especially * electron configuration. Usually this is a *, one with eight outer-shell electrons, or an *.
sharing, covalent
stable: noble gas configuration
octet
Octet rule–In forming bonds, * elements gain, lose, or share electrons to achieve a stable electron configuration with * valence electrons.
main group, eight
same electronic configuration as a noble gas, or *
closed shell configuration
- shells,which has * electrons, are less stable than closed shells
open, unpaired
Bonds are * of * between atoms.
electrostatic forces, attraction
Chemical Bonds and Electronegativity
Bonds are electrostatic forces of attraction between atoms.
Start with 2 concepts
• * – atoms acquire a special stability when orbitals are completely filled
• * between the two bonded atoms.
Octet rule
Electronegativity difference
In the * of an element, valence electrons are represented by * placed around the chemical symbol of the element.
Lewis symbol, dots
A * is a * of Lewis symbols that depicts the * or * of electrons in a chemical bond.
Lewis structure, combination
transfer, sharing
- –an idealized type of bonding based on the attraction between
metal ions and a * of their valence electron.
In general, metal atoms are relatively large and have a small number of outer electrons that are shielded from the * by filled inner levels.
• Thus they lose their electrons comparatively easily (* IE) but do not gain them very readily (* EA).
• Thus metal atoms are in a pool of evenly distributed sea of electrons that flow among the metal ions and attracts them together. The electrons in metallic
bonding are delocalized and move freely throughout the piece of metal.
Metallic bonding
delocalized “sea”
nuclear charge
low IE
small or positive EA
Ionic bond–a bond formed on the basis of * that exist between *. The ions are formed from atoms by a * of one or more electrons
Ionic bonds exist between atoms that have a high *
electrostatic forces
oppositely charged ions
transfer
EN difference.
Ions are * atom or group of atoms (poly-atomic ion)
- atoms that lose electron(s) become *
- atoms that gain electron(s) become *
electrically charged
cations
anions
Binary ionic compounds consist of * cations and anions.
Ternary ionic compounds consist of * and * ions.
monatomic
monatomic, polyatomic
Cations are * than their parent atom.
• The outermost electron(s) vacates its (their) orbital, resulting an a * of
electron-electron repulsion and a * effective inward pull of the nucleus.
smaller
decrease, greater
The increase in electrons * the electron-electron repulsion and thus the
electrons are more spread out in space.
increases
- is the quantity of energy released in the formation of one mole of a * from its * ions.
Lattice energy, crystalline ionic solid
separated gaseous
The * is an important indication of the * of ionic interactions
and is a major factor in influencing * (mesh) of
ionic compounds.
• Lattice energies can be calculated through a * , which is a series
of steps from elements to ionic compounds for which all the enthalpies are known except the lattice energy.
lattice energy, strength
melting points, hardness, and solubilities
Born-Haber cycle
A covalent bond is a bond formed between two or more atoms by a * of electrons
• Covalent bonds exist between atoms with high *
• Both would want to keep their electrons and are not willing to give up their
electrons. The only way they can get each others electrons and attain stability
through the octet rule is by sharing electrons.
sharing
EN values.
A * is a pair of electrons involved in covalent bond formation.
A * is a pair of electrons found in the * of an atom and not involved in bond formation.
bond pair
lone pair, valence shell
The sharing of a * pair of electrons between bonded atoms produces a single covalent bond.
• In a double covalent bond, * pairs of electrons are shared between bonded atoms. The bond is represented by a double-dash sign. E.g., CO2.
• In a triple covalent bond, * pairs of electrons are shared between the bonded atoms. It is represented by a triple-dash sign. E.g., HCN
• In a * , electrons shared between two atoms are contributed by just * of the atoms. As a result, the bonded atoms exhibit
* . E.g., NH4+
single two three coordinate covalent bond - one formal charges
Those with EN difference (DEN) of zero are considered * . Equal sharing of electrons. E.g., Cl2, H2.
• The maximum possible DEN occurs in * Considered as *. Complete transfer of electrons.
• All the rest are partly covalent and partly ionic
• The greater the DEN, the more * the bond. This will have great relevance to
polarity of molecules, and in turn to the phases of matter.
pure covalent
CsF (or FrF).
purely ionic.
polar
In a polar covalent bond a * exists between the * of positive and
negative charge in the bond.
• The * sharing of the electrons leads to a * negative charge (d−) on the more nonmetallic element, and a corresponding partial positive charge (d+)
on the more metallic element
separation, centers
unequal
partial
An * depicts the electron charge distribution in a
molecule. The color red is used to represent the region with the most *
charge, and blue represents the most * charge.
electrostatic potential map
negative, positive
Multiple covalent bonds are formed most readily by *
C, N, O, P, and S atoms.
- A * atom is bonded to two or more atoms.
* A * atom is bonded to just one other atom.
central
terminal
• Hydrogen atoms are always * atoms.
• Central atoms are generally those with the * .
• Carbon atoms are always * atoms.
• Except for the very large number of chain-like organic molecules, molecules
and polyatomic ions generally have compact, symmetrical structures.
terminal
lowest electronegativity
central
- is the number of outershell (valence) electrons in an isolated atom
minus the number of electrons assigned to that atom in a Lewis structure.
Formal charge
The sum of the formal charges in a Lewis structure must equal * for a neutral molecule and must equal the magnitude of the * for a polyatomic ion.
• Where formal charges are required, they should be as * as possible.
• Negative formal charges usually appear on the * atoms; positive formal charges, on the * .
• Structures having formal charges of the same sign on adjacent atoms are *.
zero, charge
small
most electronegative, least electronegative atoms
unlikely
- occurs when two or more plausible Lewis structures can be written for a species. The true structure is a * of these different contributing structures.
• The possible Lewis structures are the * (or resonance forms).
• The actual molecule is a resonance * , something like an average of the resonance forms.
• The need for resonance structures is the result of *
Resonance composite or hybrid resonance structures hybrid electron-pair delocalization
Exceptions the the Octet Rule
Odd-Electron Species
* are highly reactive molecular fragments containing unpaired electrons.
*Most of them have a central atom from an odd-numbered group, such as *
*The presence of an unpaired electron causes odd-electron species to be *
e.g., NO, NO2
Free radicals
N and Cl
paramagnetic
Incomplete Octets
Gaseous compounds containing * as the central atom are often electron
deficient (> 8 in the central atom).
• e.g., BeCl2 and BF3
Electron-deficient molecules often attain an octet in reactions by forming additional
bonds. e.g., BF3 + NH3 → BF3NH3
Be and B
- is a term used to describe Lewis structures in which
certain atoms in the *or higher period of the periodic table appear to require * electrons in their valence shells.
e.g., SF6, PCl5, H2SO4, (SO4)2-, SO3, SO2
Expanded valence shell
third
10 or 12
S and P can accommodate a maximum of * electrons, and I as many as *. These atoms expand their valence shells to form more bonds and *
formal charge.
12, 14
minimize
- is one-half the difference between the numbers of electrons in bonding
and in antibonding molecular orbitals in a covalent bond.
Bond order
- is the distance between the centers of two atoms joined by a covalent bond which be approximated as the
sum of the covalent radii of the two atoms.
Bond length (bond distance)
- , D, is the quantity of energy required to break one mole of covalent bonds in a gaseous species, usually expressed in kJ/mol.
Bond-dissociation energy
An * is the average of bond-dissociation energies for a number of different species containing a particular covalent bond
average bond energy
Bond energies of higher order bonds are * than those of lower bond orders.
• Double bonds have higher bond energies than do single bonds between the
same atoms, but they are not twice as large.
• Triple bonds are stronger still, but their bond energies are not three times as
large as single bonds between the same atoms.
greater
Enthalpy of reactions involving gases can be approximated from bond energies
Delta H◦rxn = DeltaH (bond breakage) + DeltaH (bond formation)
≈ summation of BE(reactants) + summation og BE(products)
read
Process of reaction (enthalpy):
gaseous reactants → gaseous atoms → gaseous products
weak bonds (reactants)→ strong bonds (products)
Delta H < 0 exothermic
strong bonds (reactants) → weak bonds (products)
Delta H > 0 endothermic
