Kinetics

Presence of a catalyst

Concentration of reactants

Color of reactants

Temperature at which reaction occurs

It works for both one-step mechanisms and multi-step mechanisms

None of the above

It needs experimental data

It can be used for the slow elementary step to write rate laws

Add inert gas

Remove some solvent

Increase quantity each component

Reduce size container

Endothermic step

Exothermic step

Intermediate step

Rate-determining step

The polar nature of NaCl reduces its tendency to vaporize.

The larger molar mass of I2 increases its vapor pressure.

Covalent bonds within I2 provide greater volatility.

The strong ionic bonds in NaCl require more energy to break into gas phase than London dispersion forces in I2.

The rate would decrease by a factor of 4.5.

The rate would double.

The rate would increase by a factor of 4.5.

The rate would remain unchanged.

Transition states occur at every mechanistic step whereas intermediates only form during reversible steps of mechanisms.

Intermediates increase overall energy barriers while transition states serve as temporary stable structures within reactions.

Both intermediates and transition states can be isolated under appropriate conditions if done quickly enough due to their stability levels being comparable.

An intermediate represents an actual chemical species that exists momentarily, while a transition state symbolizes an energetic peak between reactants and products.

N2

2NO

H2

N2O2

Covalent network solid

Molecular solid

Ionic solid

Metallic solid

The half-life halves as well since doubling concentration implies twice as much reacting substance every given period.

The half-life increases fourfold as it’s directly proportional to initial concentration squared for first-order kinetics systems.

The half-life doubles because each successive concentration requires double time to reduce by half due to exponential decay nature of first-order kinetics.

The half-life remains constant because it does not depend on initial concentration for first-order reactions.