Chemical Bonds

Front 1

what is a bond

Back 1

electrostatic (stationary or slow moving) force btwn charged objects

Front 2

trends in Coulombic potential for ionic substances

Back 2

as distance (btwn nuclei) increases, PE decreases, as distance decreases, PE increases

Front 3

what does a bond represent in a PE diagram

Back 3

represents the min PE (repulsion raises E)

Front 4

when do bonds occur btwn nuclei

Back 4

if the E of the aggregate is less than the E of the separated atoms

Front 5

what do the energy terms in the energy vs internuclear distance diagram represent

Back 5

E terms involved are the PE that results from attractions and repulsions among charged particles and KE caused by motions of electrons

Front 6

what does a zero reference point in energy correspond to in internuclear distance

Back 6

this reference pt is defined for atoms at infinite separation

Front 7

why does the left side of the graph rise so steeply in terms of E

Back 7

at very short distances this happens bc of the great importance of internuclear repulsive forces at these distances

Front 8

what does the minimum E pt on the graph correspond to in terms of internuclear distance

Back 8

bond length, the bond energy corresponds to the difference between the zero ref. pt and the lowest pt

Front 9

why is the bonded H2 molecule more stable than 2 separate H atoms

Back 9

because in the H2 molecule the electrons reside primarily in the space btwn the 2 nuclei, where they are attracted simultaneously by both protons, this positioning leads to their greater stability (PE is lowered bc of incr. attractive force in the area)

Front 10

is the KE of the electrons also affected by individual atoms forming a molecule

Back 10

yes, the KE of the electrons is affected as well (the E graph shows a combination of KE and PE, actually)

Front 11

how can you think of bonds in terms of forces

Back 11

attraction btwn protons and electrons of different atoms pulls the protons toward each other, there is a balance btwn the attractive forces and the proton-proton and electron-electron repulsive forces

Front 12

covalent bonds

Back 12

electrons are shared by nuclei

Front 13

polar covalent bonds

Back 13

situated btwn covalent and ionic bonding extremes, in polar covalent bonds, the electrons are pulled more to one atom (unequal sharing)

Front 14

δ

Back 14

symbol for partial (fractional) charge, followed by + or -

Front 15

what property describes the different affinities of atoms for the electrons in a bond

Back 15

electronegativity

Front 16

dipole moment (aka dipolar)

Back 16

molecule that has a center of positive charge and negative charge

Front 17

debye (D)

Back 17

SI units of Coulomb-meter (C m), dipole moment measurement, determined by bond length distance x electron charge, 1 D = 3.336x10 -30

Front 18

3 types of molecules with polar bonds but no resulting dipole moment

Back 18

linear (2 identical bonds), planar (3 identical 120° bonds), tetrahedral (4 identical 109.5° bonds)

Front 19

electronegativity value range for nonpolar bonds

Back 19

Δχ btwn 0 and 0.6

Front 20

electronegativity value range for polar bonds

Back 20

Δχ btwn 0.6 and 1.6

Front 21

electronegativity value range for ionic bonds

Back 21

generally greater than 1.6 (exception - HF is polar, even though Δχ is 1.8)

Front 22

bond energy (dissociation energy)

Back 22

E req'd to break a bond btwn 2 atoms in gas phase (always a positive quantity)

Front 23

formula for dipole moment

Back 23

μ (dipole moment) = δ (partial charge) x d (dist. btwn charges)

Front 24

μ is what type of quantity

Back 24

μ (dipole moment) is a vector quantity, with magnitude and direction (indicated pointing toward NEG charge, + forms tail of arrow)

Front 25

what are bond dissociation energies used for

Back 25

can be used to calculate thermodynamic properties, such as enthalpy (heats of formation and heats of reaction)

Front 26

formula for calculating bond dissociation energies

Back 26

ΔH = ΣBE (reactants) - ΣBE (products) (NB one of few times where we see intial minus final)

Front 27

As the number of bonds between two carbon atoms increases, which one of the following decreases - bond energy, bond length, number of electrons btwn atoms

Back 27

bond length (bond energy increases)

Front 28

When can coulomb's law be applied

Back 28

only when the charged bodies are much smaller than the distance separating them and therefore can be treated approximately as point charges. When combined with principles of quantum physics, Coulomb's law helps describe the forces that bind electrons to an atomic nucleus, that bind atoms together into molecules, and that hold together solids and liquids

Front 29

another name for bond length

Back 29

equilibrium internuclear distance