Chapter 8 and 9.1- Bonding Flashcards
3 different terms effectively meaning the same thing…
Define:
1. Dipole
2. Dipole Moment
3. Polar Bond
- Something with 2 poles
- A polar molecule (has a positive end and a negative end)
- Polar bond (bond with unequal electron sharing)
Formal charge and which formal charges are best
Group number - (dots + lines)
Best structure = closest to formal charge of 0 on ALL ATOMS, and if tied best = most negative charge on the most electronegative element
Lewis structures:
1. Which element makes up the center?
2. Where do hydrogens go in a compound?
- Least electronegative ones, carbon if it is present
- If the center has been filled, the most electronegative ions
Free radical
element with an unpaired lone electron
When determining shapes using VSEPR Theory single, double, and triple bonds all count as…
single bonds
How are dipole moments drawn? (individual AND overall)
Overall: Arrows with perpendicular lines at the bottom pointing towards the negative end
INdividual: For EACH polar bond use that weird little squiggle guy for positive and negative parts
Heat of reaction of bonds
Sum of bonds broken - sum of bonds formed (in kj/mol)
Breaking bonds is always…
Forming bonds is always…
(exothermic/endothermic)
Breaking bonds is endothermic
Forming is exothermic
Bond order
Number of lines/number of locations
Tells us if lines are single double or triple, and therefore the energy length and energy. Higher values mean more attraction and shorter lengths, whereas the opposite is true for lower values.
When is a bond ionic?
metal-nonmetal OR polyatomic ions
Lone electron pairs repel other domains (more/less) than atoms (VSEPR)
more
What are the electronegativity differences for nonpolar, polar, and ionic bonds respectively? (ex. Fluorine is 1.9 units more electronegative than Oxygen)
Nonpolar: 0.3 or below
Polar: 0.3-1.7
Ionic: 1.7-3.3
What are the most electronegative elements?
Cl, N, O, F (in increasing order)
Charge (almost) always creates a bigger difference than energy level
Difference between ionic, covalent, and metallic bonds?
Ionic: metal-nonmetal, moreso “stealing” of electrons rather than sharing, held together via opposite charge creating electrostatic attraction
Covalent: nonmetal-nonmetal, sharing of electrons to create stability
Metallic: like an electron “soup,” electrons don’t stay with any particular atom, using the force of attraction between the delocalized electrons and cations to hold the bond together
As covalent bond interatomic distance decreases energy (increases/decreases)
increases until a point of maximum stability and lowest energy
Pulling the atoms too close will result in the force of repulsion between nuclei to overcome interatomic electron-nucleus attraction
What is the purpose of resonance structures?
- To provide stability by sharing the uneven formal charge between atoms
- The stability of a structure (closeness to 0 on everything) determines the time spent in the structure
What are the exceptions to the octet rules?
- 3rd period and below NONMETALS can use the d orbitals and have up to 6 bonds
- Boron can have 3 bonds
- Beryllium can have 2 bonds
- Hydrogen can have 1 bond
Define domain
area of an electron (lone pairs, single bonds, double bonds, and tripe bonds all count as one)
Name the bonds that make up single, double, and triple bonds
single: sigma bond
Double: 1 sigma bond, 1 pi bond
triple: 1 sigma bond, 2 pi bonds
Polar vs. nonpolar covalent bonds
Polar - Unequal sharing due to diff. electronegativity, creating separate geometric centers of average positive and negative charge
Nonpolar - Equal pull/electronegativity of molecules
When asked about lattice energy, the related equation is…
Coulomb’s law (relating charge first, and if the same then distance to force of attraction/energy holding the lattice together) (multiply the absolute value of charges of the atoms (do not take into account the # of atoms))
How might orbitals change as an atom goes from a free state to a component of a molecule?
Hybridization of orbitals due to shifting electrons to arrange them in the manner with the least energy when existing free vs. in a molecule
What energy level is 2sp3 closest to? 2sp2? 2sp1?
2sp3 is closest to 2p, being 75% p character, with 2sp2 being 67% p character and 2sp being in the middle with 50% p and 50% s character
Which orbitals are atomic? Which are molecular?
Standard orbitals are atomic whereas hybrid orbitals (ex. sp3) are molecular
When do Lewis diagrams/atoms have a lone pair before a single electron in each valence orbital?
In the case of the carbon family hybridizing its electrons in sp3, promoting all electrons into equal degenerate orbitals
Breaking bonds are (exothermic/endothermic)
endothermic
In determining bond length for covalent vs. ionic bonds…
Covalent: Prioritize whether bonds are single, double, or triple, then radius
Ionic: Ionic “bond lengths” should be determined by charge then radius (charge overcomes distance)
Define covalent networks
Covalent molecules in a network (think crystal lattice w/o the whole + or - charges)
Most often made by Carbon or Silicon due to ability to bond to 4 molecules
Is C-H polar?
nO
Can halogens form double bonds?
non
What is a sigma bond? a pi bond?
- A sigma bond is a direct sharing of electrons via overlapping orbitals
- A pi bond is a pair of parallel p orbitals that do not overlap and are perpendicular to the sigma bonds (attraction above AND below a sigma is 1 pi bond)
Formula for determining net bond energy
Sum of REACTANTS - sum of products
Does electronegativity have an impact on bond length?
No
