CHE2B Midterm 1 Review

Internal energy (U)

a microscopic energy contained in a substance.

Internal energy (U) =

Thermal energy + Chemical Energy

Thermal energy

• The energy that results in temperature

• associated with the random molecular motion

• Kinetic Energy

• Higher T = faster molecular motions = higher kinetic energy = higher thermal E

Thermal energy ________ with temperature

increases

Chemical Energy

• associated with chemical bonds and intermolecular attractions

• potential energy

Chemical Energy changes when:

• a chemical reaction occurs (e.g. bond breaking resulting in increase in energy)

• A physical change occurs (e.g. the state of matter changes; i.e. Solid → Liquid → vapor)

Universe =

System + surroundings

The system and the surroundings are connected via:

• Matter - Flow of molecules across boundary.

• Heat – Transfer of energy from high T to low T.

• Work – Force acting through a distance.

Open System

a system that can exchange matter with the surroundings

• Energy exchange as heat or work is also allowed

Example of an open system

a beaker of hot coffee because:

• Matter Exchange: Water Vapor is escaping from the beaker

• Heat Exchange: Heat flows to the surroundings

Closed system

a system that does not allow the flow of matter with surroundings

• The transfer of energy as heat or work is allowed.

Example of a closed system

a sealed flask of hot coffee or a balloon of helium because:

• When a balloon expands, gas inside does work on the surroundings and loses energy.

• Heat will flow out of the balloon if the outside temperature is lower.

Isolated system

a system that cannot exchange any heat, work, or matter with the surroundings

Example of an isolated system

hot coffee in a thermos

Heat (q)

• Internal energy transferred between a system and its surroundings as a result of a temperature difference.

• “flows” from hotter to colder

Calorie (cal)

The quantity of heat required to change the temperature of one gram of water by one degree Celsius

Joule (J)

SI unit for heat

1 cal =

4.184 J

Heat and Law of Conservation of energy

• Heat gained by a system is lost by its surroundings.

• Energy is neither created nor destroyed

q_system (q_sys or q) and q_surroundings (q_surr) are defined as positive when

heat is gained

q_system (q_sys or q) and q_surroundings (q_surr) are defined as negative when

heat is lost

q_sys =

-q_surr

Heat flows from the surroundings to the system

q_sys(q) = +

q_surr = -

Heat flows from the system to the surroundings:

q_sys(q) = -

q_surr = +

Heat Capacity (C)

the quantity of heat required to change of a substance by 1^oC

True or false: Different Materials heat differently

True; 100 g of Water heats much slower than 100 g of metal.

An extensive property is

proportional to the system size (moles, grams)

• e.g. Volume, mass, U, heat capacity

An intensive property

does not depend on the system size (moles, grams)

• e.g. temperature, pressure, density

Chemical reactions

involve breaking and formation of chemical bonds, resulting in a change in chemical energy

If chemical energy decreases upon reaction,

• released energy is converted to thermal energy, and temperature increases

• the thermal energy is eventually released to the surroundings as Heat

Heat of reaction (q_rxn)

The quantity of heat absorbed by a system from its surroundings when a chemical reaction occurs within the system.

Exothermic Reaction

a reaction that releases heat

In an Exothermic Reaction:

Chemical energy is converted to thermal energy

• The system temperature increases

• Heat is released to surroundings.

• qrxn: Negative

Endothermic Reaction

a reaction that absorbs heat

In an Endothermic Reaction

Thermal energy is converted to chemical energy

• The system temperature decreases

• Heat is absorbed by the system

• q_rxn: Positive

Bomb calorimeter is

a constant volume calorimeter used in measuring the heat of combustion reaction.

• It can withstand large pressure

• The heat released by the reaction is absorbed by the calorimeter

• Temperature of calorimeter changes due to the heat

Coffee-Cup Calorimeter

a calorimeter used in measuring the heat of solution reaction at the constant pressure.

• The reaction mixture is in the inner cup.

• The outer cup is for thermal insulation.

• The heat from the solution reaction is absorbed by the solution in the cup.

Work

• the product of an external force acting on an object and the displacement the force has caused.

• Ex: Gas formed pushes against the atmosphere.

Pressure-Volume Work

a type of work done by the external pressure through a volume change. (W= -PΔV)

Gas expands against constant external pressure of P

• ΔV (+)

• w = -PΔV (-)

• The system does work on the surroundings

Gas is compressed by an external constant pressure of P

• ΔV (-)

• w = -PΔV (+)

• The surroundings does work on the system

The state of a system can be described

by a set of variables (e.g. pressure, temperature, number of moles, composition)

Any property that has a unique value for a specified state of a system is said to be a __________

state function.

Ex: 100g of water at 293.15 K and 1.00 atm; Density has a unique value (d = 0.99820 g/mL)

A state function ________ depend how the state was established

does not

Intensive state functions

Pressure, temperature

Extensive state functions

Volume, Internal Energy

Path Dependent Functions are

properties which depend on how the system changes from the initial state to the final state (path).

Example of path functions

Heat and work; Amount of work done by the system does depend on how many plates you remove at once.

The First Law of Thermodynamics (Law of Conservation of Energy)

Change in the internal energy of the closed system (∆𝑈) is equal to the amount of heat supplied to the system (𝑞), plus the amount of work done on the system (𝑤).

If we know q and w, then can calculate ∆U.

• Combustion reaction at constant volume (∆V=0) → ∆U=qv

• “The energy of an isolated system is constant.” q=0, w=0 → ∆U=0

Enthalpy (H)

• a type of energy defined as: H = U +PV

• Is an extensive state function

• Consider enthalpy change during a reaction at constant P.

At constant P:

• ∆H = ∆U + P∆V

• Since ∆U = q_p - P∆V, ∆H = (q_p - PΔV) + PΔV = q_p

• ΔH = q_p

• the change in enthalpy is the heat gained or lost

The fusion and the vaporization are bond breaking processes and are thus ________

endothermic. q_p > 0

The enthalpy change of 1 mole of a substance during the fusion is called

the (molar) enthalpy of fusion (in kJ/mol)

• ΔH_fus = q_p > 0

The enthalpy change of 1 mole of a substance during the vaporization is called _______.

the (molar) enthalpy of vaporization (in kJ/mol)

• ΔH_vap = q_p > 0

As heat is added to an ice cube at a constant rate, the water changes from a solid to a liquid, and finally to a gas. A plot of temperature versus heat added for a substance is called _________.

the heating curve

Standard States

• Gas: Pure gas at 1 atm

• Liquid: pure liquid at 1 atm

• Solid: pure solid at 1 atm

• Solute: At a concentration of 1 M (mole/liter)

• Temperature is 298.15 K unless stated otherwise

ΔH^o is proportional to the amounts of reactants and products. This means that if ΔH^o is 180 kJ when 1 mole each of N2 and O2 are reacted,

• ΔH^o is halved if only ½ mole each of N2 and O2 are reacted.

• ΔH^o changes sign when a process is reversed.

Hess’s Law

If a process occurs in steps (even hypothetically), the enthalpy change for the overall process is the sum of the enthalpy changes for the individual steps.

The standard enthalpy of formation (∆H_f^o) is the

enthalpy change in the formation of one mole of a substance in the standard state from the reference forms of its elements in their standard states.

The standard enthalpy of formation of a pure element in its reference form is ___.

0

If there are multiple allotropes, the ______ is usually the most stable form at the standard state and 25^oC.

reference form

Compounds with _____ tend to be stable because the enthalpy of the compound is lower relative to that of constituent elements and its decomposition needs energy input as heat.

ΔH° < 0

Compounds with ______ tend to be less stable because

they give off heat (energy) upon decomposition.

ΔH° > 0

Spontaneous Process

• a process that is capable of proceeding in a given direction without outside intervention.

• May be fast or slow

Ex: Combustion reactions because once the reaction starts, it continues on its own. (fast)

• Iron rust spontaneously in air (very slow)

Nonspontaneous processes

will not occur without outside intervention

If a forward process is spontaneous, the reverse process is always ___________.

nonspontaneous

Endothermic reactions (melting of ice) can be __________ (spontaneous or nonspontaneous)?

spontaneous

Exothermic reactions (freezing of water) can be ___________ (spontaneous or nonspontaneous)?

nonspontaneous

Mixing of two ideal gases _______ produce or absorb heat, but is spontaneous.

does NOT

Entropy (S)

measures the probability of the state (randomness)

Entropy is _____ when there is no randomness. Entropy has a ______ value otherwise.

zero; positive

The more random the system is

the higher the entropy.

Entropy is an (extensive or intensive) state function?

extensive

Third Law of Thermodynamics:

The entropy of a pure perfect crystal at 0 K is zero.

Entropy increase when:

• there is an increase in temperature

• there is phase changes due to the trend S_solid < S_liquid << S_vapor

ΔS = q_rev/ T holds true only if:

• The process is reversible (The system is in equilibrium during the entire process.)

• There is no temperature change during the process

Reversible processes:

• Melting at melting point

• Vaporization at 1 atm at normal boiling point

Irreversible processes:

• Melting at T > T_f (freezing point)

• Reaction not in an equilibirum

The Second Law of Thermodynamics

All spontaneous process produce an increase in the entropy of the universe. ΔS_univ > 0

The reaction is:

• Spontaneous if ΔS_univ > 0

• Nonspontaneous of ΔS_univ < 0

• In equilibrium if ΔS_univ = 0

A reaction proceeds spontaneously in the forward direction if

ΔG = ΔH - TΔS < 0

ΔG < 0, the process is

spontaneous

ΔG > 0, the process is

nonspontaneous

ΔG = 0, the process is

at equilibrium

If ΔH > 0, ΔS > 0

the reaction is spontaneous at high T.

If ΔH < 0, ΔS > 0,

the reaction is always spontaneous

If ΔH < 0, ΔS < 0

the reaction is spontaneous at low T

If ΔH > 0, ΔS < 0

the reaction is always nonspontaneous

Entropy (S) is an _______ state function while standard molar entropy (S^o) is an _______ state function.

extensive; intensive

The standard Gibbs free energy of formation, ΔG_f^o is

the change in Gibbs free energy for a reaction in which 1 mole of a substance in its standard state is formed from the reference forms of its elements in their standard states.

Molar entropy (randomness) of a gas _______ with _______ pressure P (in atm) at ________ temperature.

increases; decreasing; constant

What direction of spontaneous change is: ΔG < 0

Forward Reaction (product favored)

What direction of spontaneous change is: ΔG > 0

Reverse reaction (reactant favored)

What direction of spontaneous change is: ΔG = 0

System is at equilibrium

ΔH^o > 0 (endothermic), K_eq _______ with increasing T.

increases

ΔH^o < 0 (exothermic), K_eq _______ with increasing T.

decreases

Electric and magnetic fields propagate as _____ through a vacuum or through a medium.

waves

Black-body radiation

• All non-reflective objects (black bodies) emit light (Black-body radiation).

• A black-body at room temperature appears black as most of the energy it radiates is infrared.

• A heated metal emits visible light.

• As the temperature becomes higher, the color of the emitted light shifts from red to blue.

• The spectrum of black-body radiation is peaked at a characteristic frequency that shifts to shorter wavelength with increasing temperature.

Constructive interference

crest meets crest; amplitude of the waves add together