2 - Bonding, IMFs, and Thermodynamics

strength of chemical bonds depend on: (2)

1. more electrons shared = stronger bond
2. shorter distance b/w atoms = stronger bond

increasing bond strength = (increasing/decreasing) bond dissociation energy

increasing

covalent bonds

formed b/w atoms with HIGH electronegativity (nonmetals with nonmetals)

what are some properties of compounds with covalent bonds? (3)

1. not conductors
2. insulators
3. rigid

metallic bonds

formed b/w atoms with low electronegativity (metals with metals)

what are some properties of compounds with metallic bonds? (4)

1. delocalized electrons = conduction
2. 'cloud of electrons' donated among all atoms
3. malleable
4. conductors

coordinate covalent bonds

formed b/w atoms with lone pairs and electron deficient species

what are some properties of compounds with coordinate covalent bonds? (3)

1. Electrons are DONATED from nucleophile (similar to Lewis acids and bases)
2. compounds are easily dissociated (coordinate complexes)
3. rigid and directional

which are more than the other? coordinate covalent and covalent bonds...
1. polarity
2. bond length
3. stronger

1. coordinate covalent bond more polar
2. coordinate covalent longer
3. covalent stronger

ionic bonds

formed between particles of opposite charges

what are some properties of compounds with covalent bonds? (4)

1. non-conductors (electrons localized)
2. ions (cations/anions) dissociate in aqueous solution as electrolytes
3. insulators
4. brittle

intermolecular forces (IMFs)

force produced when particles of opposite charge attract each other (TWO different molecules that might have intramolecular bonds/forces within themselves)

strength of IMFs depends on: (2)

1. larger charges = strong attractive force
2. particle size is negligible due to large distance

types of IMFs: (5)

1. ion-dipole
2. dipole-dipole
3. dipole-induced dipole
4. london dispersion
5. H-bonds

ion-dipole forces:

IMF produced b/w ions and polar molecules

dipole-dipole forces

IMF produced b/w polar molecules

how to increase dipole-dipole forces: (2)

1. increase the charge magnitude
2. increase the length (of dipole) = reduce attractions between individual dipoles

dipole-induced dipole forces:

IMF produced between polar and apolar molecules

how does polar-induced dipole forces occur?

polar molecule attracts the electron density of an apolar molecule. Therefore very easily CLEAVED

how to increase polar-induced dipole forces? (2)

1. INCREASE SIZE for grater separation b/w 2 polar molecules -> reduce dipole moment
2. INCREASE CHARGE (increase pos and eg charge in apolar molecule -> charge separation -> will act more polar)

london dispersion forces

IMF produced by temporary and small collisions by small dipoles -> electron cloud gets deformed

hydrogen bond

IMF involving hydrogen bonds between VERY similar molecules

what are the hydrogen bond acceptors? (3)

nitrogen, oxygen, fluorine

rate the bond and IMF strengths in increasing order:
dipole-dipole, metallic, LDW, ionic, hydrogen, covalent,

LDW < dipole-dipole < H-bond < metallic < ionic < covalent

what forces do van der waals contain? (3)

LDW, dipole-dipole, H-bonds

does boiling point (increase/decrease) with increasing IMF?

increase

does melting point (increase/decrease) with decreasing IMF?

decrease

does vapor pressure (increase/decrease) with increasing IMF?

decrease

does viscosity (increase/decrease) with decreasing IMF?

increase

enthalpy (H)

the energy stored in chemical bonds/any attractive force

how is change in enthalpy (ΔH) described as in terms of products and reactants?

ΔH is the nrg difference b/w reactants and products

how does ΔH look for exothermic rxns? why

ΔH < 0 bc nrg is RELEASED

how does ΔH look for endothermic rxns? why?

ΔH > 0 bc nrg is put in

what is ΔHf and why is it 0?

amount of nrg associated w/ forming ONE MOLE of compound (in standard state). ΔHf is 0 bc rxns are in EQ in its standard state

hess' law

the idea that the TOTAL enthalpy of a rxn is the sum of the ΔH values of the processes. reverse rxn's direction changes the ΔH sign

how to increase entropy (S) of a rxn? (2)

1. increase # of particles (2 mols have more S than 1)
2. changing the phase of the rxn (solid to liq, liq to gas)

gibbs free nrg (ΔG)?

energy available to do work. negative in spontaneous rxn, positive in endothermic rxn

what is gibbs free nrg formula?

ΔG = ΔH - TΔS
- ΔH = change in enthalpy
- ΔS = change in entropy
- T = temp (Kelvin)

Is dephosphorylation of ATP spontaneous? (say in terms of ΔG = ΔH - TΔS)

- ΔH = pos (b/c bonds are BROKEN)
- ΔS = pos (b/c more particles now (P + ADP vs. ATP)
- T = if temp is HIGH, ignore ΔH (-); if temp low, ignore S (+)

ATP dephosphorylation is only spontaneous in HIGH TEMPERATURE

how to determine hybridization? (formula)

hybridization = (atoms bonded) + (lone pairs - count as PAIRS)

What is the formal charge on the central oxygen atom of ozone, O3?

A. -2
B. +1
C. -1
D. 0

B.
think monkey think

In which of the following does hydrogen have a partial negative charge?

A. H2O
B. NH3
C. BH3
D. CH4

C.

Which of the following pairs of molecules have the same shape?

A. BF3 and NH3
B. H3O+ and PH3
C. H2O and PH3
D. PH3 and CH4

B.
H3O+ (has 3 groups + 1 lone pair = sp3)
PH3 (has 3 groups + 1 lone pair = sp3)
both make trigonal pyramidal (sp3 w/ 1 lone pair)

How many of which type of orbitals are used to create the hybrid orbitals on carbon in a molecule of ethylene (C2H4)?

A. One s and one p orbital
B. Two s and two p orbitals
C. One s and three p orbitals
D. One s and two p orbitals

D.
C2H4 has sp2 hybridization (1 s and 2 p) dumba*s

How many electrons are found in π bonds in acetylene (C2H2)?

A. 4
B. 0
C. 6
D. 2

A.
HC≡CH

How many electrons are shared in the bond between sodium and chlorine in a molecule of NaCl?

A. 1
B. 0
C. 2
D. 3

B.
NaCl is an IONIC compound

Which of the following describes the orbital geometry of an sp2 hybridized atom?

A. Tetrahedral
B. Trigonal planar
C. Bent
D. Linear

B.
sp - linear
sp2 - trigonal planar
sp3 - tetrahedral
A bent geometry is generally associated with an sp3 hybridized central atom that has two lone pairs of electrons.

Rank the average C—O bond length from shortest to longest for CO, CO2, and CO3 2-.

A. CO3 2- < CO2 < CO
B. CO < CO2 < CO3 2-
C. CO3 2- < CO < CO2
D. CO < CO3 2- < CO2

B.
triple bond ALWAYS shorter than double. double bond ALWAYS shorter than single bond

Which of the following orbitals is NOT found in the carbon-carbon bonds of propene?

A. Pure p orbitals
B. sp3 hybrid orbitals
C. Pure s orbitals
D. sp2 hybrid orbitals

C.

What molecular shape AND angle does sp hybridization and NO lone pairs make?

linear (180°)

What molecular shape AND angle does sp2 hybridization and NO lone pairs make?

trigonal planar (60°)

What molecular shape AND angle does sp3 hybridization and NO lone pairs make?

tetrahedral (109.5°)

What molecular shape AND angle does sp2 hybridization and ONE lone pair make?

bent (~109.5°)

What molecular shape AND angle does sp3 hybridization and ONE lone pairs make?

trigonal pyramidal (<109.5°)

What molecular shape AND angle does sp3 hybridization and TWO lone pairs make?

bent (~109.5°)