Chemical Bonding

Question 1

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.

Answer 1

Lewis theory
outermost (valence) electronic shell
transferred
electrostatic forces, ionic

Question 2

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 *.

Answer 2

sharing, covalent
stable: noble gas configuration
octet

Question 3

Octet rule–In forming bonds, * elements gain, lose, or share electrons to achieve a stable electron configuration with * valence electrons.

Answer 3

main group, eight

Question 4

same electronic configuration as a noble gas, or *

Answer 4

closed shell configuration

Question 5

• shells,which has * electrons, are less stable than closed shells

Answer 5

open, unpaired

Question 6

Bonds are * of * between atoms.

Answer 6

electrostatic forces, attraction

Question 7

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.

Answer 7

Octet rule

Electronegativity difference

Question 8

In the * of an element, valence electrons are represented by * placed around the chemical symbol of the element.

Answer 8

Lewis symbol, dots

Question 9

A * is a * of Lewis symbols that depicts the * or * of electrons in a chemical bond.

Answer 9

Lewis structure, combination

transfer, sharing

Question 10

• –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.

Answer 10

Metallic bonding
delocalized “sea”

nuclear charge
low IE
small or positive EA

Question 11

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 *

Answer 11

electrostatic forces
oppositely charged ions
transfer
EN difference.

Question 12

Ions are * atom or group of atoms (poly-atomic ion)

• atoms that lose electron(s) become *
• atoms that gain electron(s) become *

Answer 12

electrically charged
cations
anions

Question 13

Binary ionic compounds consist of * cations and anions.

Ternary ionic compounds consist of * and * ions.

Answer 13

monatomic

monatomic, polyatomic

Question 14

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.

Answer 14

smaller

decrease, greater

Question 15

The increase in electrons * the electron-electron repulsion and thus the
electrons are more spread out in space.

Answer 15

increases

Question 16

• is the quantity of energy released in the formation of one mole of a * from its * ions.

Answer 16

Lattice energy, crystalline ionic solid

separated gaseous

Question 17

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.

Answer 17

lattice energy, strength
melting points, hardness, and solubilities

Born-Haber cycle

Question 18

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.

Answer 18

sharing

EN values.

Question 19

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.

Answer 19

bond pair

lone pair, valence shell

Question 20

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+

Answer 20

single
two 
three 
coordinate covalent bond - one 
formal charges

Question 21

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.

Answer 21

pure covalent
CsF (or FrF).
purely ionic.
polar

Question 22

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

Answer 22

separation, centers
unequal
partial

Question 23

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.

Answer 23

electrostatic potential map

negative, positive

Question 24

Multiple covalent bonds are formed most readily by *

Answer 24

C, N, O, P, and S atoms.

Question 25

• A * atom is bonded to two or more atoms.

* A * atom is bonded to just one other atom.

Answer 25

central

terminal

Question 26

• 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.

Answer 26

terminal
lowest electronegativity
central

Question 27

• is the number of outershell (valence) electrons in an isolated atom
  minus the number of electrons assigned to that atom in a Lewis structure.

Answer 27

Formal charge

Question 28

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 *.

Answer 28

zero, charge
small
most electronegative, least electronegative atoms
unlikely

Question 29

• 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 *

Answer 29

Resonance 
composite or hybrid
resonance structures
hybrid
electron-pair delocalization

Question 30

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

Answer 30

Free radicals
N and Cl
paramagnetic

Question 31

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

Answer 31

Be and B

Question 32

• 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

Answer 32

Expanded valence shell
third
10 or 12

Question 33

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.

Answer 33

12, 14

minimize

Question 34

• is one-half the difference between the numbers of electrons in bonding
  and in antibonding molecular orbitals in a covalent bond.

Answer 34

Bond order

Question 35

• 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.

Answer 35

Bond length (bond distance)

Question 36

• , D, is the quantity of energy required to break one mole of covalent bonds in a gaseous species, usually expressed in kJ/mol.

Answer 36

Bond-dissociation energy

Question 37

An * is the average of bond-dissociation energies for a number of different species containing a particular covalent bond

Answer 37

average bond energy

Question 38

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.

Answer 38

greater

Question 39

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)

Answer 39

read

Question 40

Process of reaction (enthalpy):

Answer 40

gaseous reactants → gaseous atoms → gaseous products

Question 41

weak bonds (reactants)→ strong bonds (products)

Answer 41

Delta H < 0 exothermic

Question 42

strong bonds (reactants) → weak bonds (products)

Answer 42

Delta H > 0 endothermic