Gen. Chem.2 2nd

Thermodynamics – branch of physics hat deals the relationship between heat, temperature, and other forms of energy

Thermochemistry – deals with the energy changes involve in chemical reaction

Work – the created motion in the presence of an opposing force

Energy – the ability to carry out work

Temperature – describes the coldness and hotness of an object

Heat – flow of thermal energy transferred from a hotter object to a colder object that is in contact with it

THE UNIVERSE

• The system pertains to the process of chemical reaction of interest, while anything else outside the system is a part of the surroundings. Together, system and surroundings make up the universe.

Types of Systems

• Isolated system – both matter and energy cannot be exchanged with the surroundings

• Closed system – energy can be exchanged with the surroundings, but matter is prevented from exchanging with the surroundings

• Open system – both matter and energy an be exchanged with the surrounding

State and Path Functions

• State function – feals with the current condition and does not concern how this condition is attained

• Path function – describe how to arrive at the current condition

Kinetic and Potential Energy

• Kinetic energy – the energy in motion; affected by the mass and velocity;

• Potential Energy – the energy of the system which is affected by its position or composition

- attractive forces between each atom’s nucleus and electrons

- repulsive forces between nuclei and between electrons;

- intermolecular forces between molecules

Internal Energy (U)

• The sum of the kinetic energy and potential energy;

Energy Change of Reaction

Every molecule has a stored energy in it. During a chemical reaction, bond-breaking and bond-forming processes occur, and each of these processes has an accompanying energy. This energy can either be absorbed or released by the system.

Exothermic and Endothermic Processes

Exothermic – processes release energy in the form of heat to the surroundings. In these processes, products have lower reactants.

• accompanied by increase in temperature (or warming) of the surroundings

• Combustion – the reaction with oxygen to produce small molecules such as CO2 and H2O

• Phase Changes – Freezing, Condensation, Deposition

• Dissolution – some ionic compounds dissociate into ions in water release energy. Example is sodium hydroxide. The heat released from sodium hydroxide solutions, especially the concentrated ones, can cause burns. Thus, one must be cautious in preparing these solutions.

• Bond Formation – two unstable atoms may forma stable molecule

• Electron Gain – Electron affinity is the energy released when a gaseous atom accepts an electron

Endothermic – processes absorb energy in the form of heat from the surroundings. In these processes, reactants have lower energy than the products.

• Accompanied by lowering of the temperature (cooling) of the surroundings.

• Dissolution – can sometimes be endothermic. For example, dissolving table salt in water can yield a cold solution as a result of energy absorption during the dissociation of sodium chloride into ions

• Phase Changes – Fusion, Vaporization, Sublimation

• Bond Dissociation – energy is the energy needed to break a bond join two atoms

• Ionization – the energy needed to remove an electron from a gaseous atom

First Law of Thermodynamics

Energy Transformation

• In a flat iron, the electrical energy is converted into thermal energy, which his useful in ironing fabrics.

• A wind mill has a turbine in which the mechanical energy created by oscillation of the propeller is converted into electrical energy that can be used as source of electricity.

First Law of Conservation of Energy

• It is not possible for energy to arise from nothing of be destroyed

• “the energy of an isolated system is constant”

• Known as Law of Conservation of Energy

• “Energy can neither be created nor destroyed”

Internal Energy (U)

• The total energy possessed by a molecule and is consisting of its kinetic energy and potential energy

• The internal energy is a state function – a function that deals with the current condition, no matter how the condition is achieved.

Consider the reaction for the formation of water:

2H_2(g) + O_2(g)> 2H₂O_(l)

Internal energy change (ΔU) – the property where the internal energy of thesubstances in the reaction is not a measurable quantity

• The difference of the final state and the initial state. In the case of the given reaction, it is the internal energy of the product less than that of the reactants:

• ΔU = U_f - U_i

ΔU_sys = -ΔU_surr OR ΔU_sys + ΔU_surr = 0

• this supports that the energy of the universe is constant

• The heat released by the surroundings is absorbed by the system, and the heat absorbed by the surroundings is released by the system.

ΔU = q + w

• The mathematical way of stating the First Law of Thermodynamics

• Internal energy of a system can be altered by doing work, and (2) supplying heat

Heat (q)

(+) heat is absorbed by the system

(-) heat is released by the system

Work (w)

(+) work is done to the system

(-) work is done by the system

Units for Work and Energy

Joule (J)

• The SI unit for heat, work, and internal energy

• Named after James Presscot Joule

• Equal to

James Presscot Joule

• An English physicist who formulated the law of conservation of energy

Calorie

• 1 cal = 4.184 J

• Another unit for energy

• This amount of heat corresponds to the heat needed to raise the temperature of 1 g of water by 1°C

Example:

A gas sample expands in a vessel, exerting a work of 22.4kJ. If 14.6kJ of that heat is absorbed by the gas, calculate the internal energy for this expansion.

ΔU = q + w

= 14.6kJ + (-22.4kJ)

= -7.8kJ

Work (w is exerted by the system, hence (-)

Heat (q) is absorbed by the system, hence (+)

l

Example:

A gas sample in a vessel has an internal energy of

-1420 J. If 250 J of heat is added in the system, calculate the work here.

ΔU = q + w

w = ΔU - q

= -1420 J + 250 J

= -1170 J

Heat (q) is absorbed by the system, hence (+)

Expansion Work

• The amount of work accompanying expansion, a process in which volume is increased.

w = - PΔV = - P(V_f - V_i)

• The work done by a system as it expands against a constant external force (P)

• The negative sign indicates that work is done by the system during expansion

Example:

Calculate the work associated with compression of a gas under a constant pressure of 843 kPa from 50.0 L to 24.0 L.

w

The unit is the same as joule.

Enthalpy of Reaction

Enthalpy

• The sum of the internal energy (U) and the product of the pressure and volume of the system.

• H = U + PV

Change in Enthalpy

• (, at constant pressure

Energy Transformation

• Just like in internal energy, the value of absolute enthalpy (H) cannot be measured experimentally. Instead, the enthalpy change (ΔH) is the property that can be directly determined.

At constant pressure, the enthalpy of a process is equal to the heat (q). This is the reason why these quantities can be used interchangeably.

• ΔH = q_p

The enthalpy of reaction (Δ_rH)

• The difference between the enthalpies of product and enthalpies of reactants

Δ_rH = H_products - H_reactants

Enthalpy change values can be referred to as heat.

The ΔH of a reaction or transformation is used to assess whether the reaction if endothermic or exothermic:

Exothermic: ΔH< 0 (negative value)

Endothermic: ΔH> 0 (positive value)

Enthalpy change is an extensive property (depends on the amount of substances involved in a process)

Molar enthalpy change is the enthalpy change accompanying a process undergone by 1 mole of a substance. Sometimes, the word “molar” is omitted in referring to this quantity.

Thermodynamic Standard State

• Simply standard state, is the physical state of a pure substance in its most stable form at 1atm and 25°C

• Gases – must be at 1 atm

• Solutions – must have a concentration of 1M

• Water, mercury, chlorine, graphite

Standard Enthalpy of Formation (Δ_fH°)

• Standard heat of formation

• The enthalpy change involved in the formation of 1 mole of a substance from elements in their standard states.

Elements in their standard states have zero (Δ_fH°)

A substance of different physical states has different standard heat of formation values.

Standard Enthalpy of Reaction (Δ_rH°)

• Standard heat of reaction

• The enthalpy change involved in a chemical reaction at standard conditions (1atm and 25°C).

• Equal to the sum of the standard heats of formation of products less the sum of the standard heats of formation of reactants.

Thermochemical Equation

• A chemical equation that indicates the energy, particularly the heat or enthalpy change, that accompanies the reaction.

Two ways of writing thermochemical equation:

1. By indicating the heat or reaction after the equation

2. By writing the heat of reaction as a reactant or product

Exothermic Reactions

• The sign of the standard heat of reaction is negative.

CH_4(g) + 2O_2(g)→ CO_2(g) + 2H₂O_(l)Δ_rH° = -890kJ/mol

• Alternatively, the amount of energy can be written as one of the products since the energy is released in the process. Note that the negative sign is no longer indicated.

CH_4(g) + 2O_2(g) → CO_2(g) + 2H₂O_(l) + 890kJ/mol

Endothermic Reactions

• The sign of the standard heat of reaction is positive.

NaCl_(s)→Na⁺_(aq) + Cl^-_(aq)Δ_rH° = 3.88 kJ/mol

• Another way to write the equation is by writing the amount of energy as one of the reactants since energy is absorbed in the process.

NaCl_(s)+ 3.88 kJ/mol→Na⁺_(aq) + Cl^-_(aq)

Mole Reaction

CH_4(g) + 2O_2(g) → CO_2(g) + 2H₂O_(l)Δ_rH° = -890 kJ/mol

• In the given thermochemical equation, the value Δ_rH° = -890 kJ/mol corresponds to890kJ of energy released for 1 mole of reaction. A mole of reaction is equivalent to the number of moles (NOT number of molecules) of a substance in a reaction as indicated in the balanced equation.

• 1 mole of reaction is equivalent to 1 mole of CH_4(g), 2 moles of O_2(g), 1 mole of CO_2(g), and 2 moles of H₂O_(l).

Physical states of the substances must be indicated in a thermochemical equation since a substance of a different states have different heats of formation.

CH_4(g) + 2O_2(g) → CO_2(g) + 2H₂O_(l)Δ_rH° = -890 kJ/mol

CH_4(g) + 2O_2(g) → CO_2(g) + 2H₂O_(g)Δ_rH° = -802 kJ/mol

If the coefficients of the balanced equation are multiplied by a number, the heat of reaction is also multiplied by the same number.

Example:

Sublimation is the process of transforming a solid into a gas. Thus, the thermochemical equation is given by:

I_2(S)→I_2(g)Δ_subH° = 62.4 kJ/mol

or

I_2(S)+ 62.4 kJ/mol→ I_2(g)

Heat of Neutralization of an Acid

Heat is directly proportional to the mass of an object and the change in temperature.

q = heat

m = mass

C_p = specific heat

ΔT = change in temperature

= final T – initial T

Calorimetry

• Study that deals with measurement of heats of processes, be it physical or chemical.

Calorimeter

• An instrument that is usually used to measure heat

Neutralization

• The reaction occurring between an acid and a base to form water and salt

Hess Law

• The enthalpy of reaction is the sum of the enthalpy changes of all steps leading to this reaction.

• It is another way to calculate hess law

Properties of Enthalpy Change

• If a chemical reaction is reverse, the Δ_rH changes its sign

• If coefficients of a balanced equation are multiplied by a certain factor, the Δ_rH is also multiplied by the same factor.

• If two equations are combined, the overall Δ_rH° is the sum of Δ_rH° values corresponding to these equations.

Chemical Kinetics

• A subdiscipline of chemistry that deals with the speed or rate of chemical reactions.

Rate of Reaction (v)

• Describes how fast or slow a chemical reaction is

• Rusting of iron, evolution of CO2 from a mixture of baking soda and vinegar

• The change in molar concentration per unit change in time

For a reaction: A → B

rate of consumption of A

rate of formation of B

For a reaction with stoichiometric coefficients other than 1:

The rate if reaction is given by:

Factors Affecting the Rate of Reaction

• Concentration

• Temperature

• Presence of a catalyst

Concentration of the Reactants

• The reaction takes place at higher rate when the reactant concentration is higher.

• n is constant

Temperature

• The reaction speeds up as the temperature is increased.

• For example, enzymatic processes that cause spoilage of food can be prolonged by storing the food in a refrigerator.

Presence of Catalyst

• A catalyst is a substance that speeds up a reaction without itself being consumed.

Rate Law

• The material statement that shows the relationship between the rate of reaction and the concentration of the reaction

• The exponent, n, is the order of the reaction. The proportionality constant, k, is called the rate constant.

Reaction Order

• Usually a whole number (0, 1, 2, 3,…); fractional orders are observed in some reactions

• Determined experimentally and is not based on the reactant’s stoichiometric coefficients.

Example:

Consider the reaction given by:

2NO_2(g) → N₂O_4(g)

Experiment shows that the reaction is of order 2. Write the rate law for the reaction.

Solution: NO₂ is the sole reactant in the reaction. The order 2 must be the exponent for the molar concentration of NO­₂ in the rate law. Thus the rate law is given by:

Differential Rate Law

• An equation that relates the rate to the concentration of the reactant.

Integrated Rate Law

• Show s the concentration of the reactant remaining at a given time. It is derived from a differential rate law that is simplified by employing techniques in calculus.

Rate Law of a Reaction

• Reactions are governed by a certain rate law. Performing kinetic experiments give us data that help us determine the order with respect to each reactant. In turn, the rate law for the reaction can be deduced.

Order of Reactions

• Zero-order reaction (n = 0)

• First-order reaction (n=1)

• Second-order reaction (n = 2)

• Etc.

Zero-Order Reaction

Consider the following reaction:

A zero-order reaction has the following differential rate law:

The equation implies that the rate does not vary with reactant concentration.

The corresponding integrated rate law is given by:

Where:

[A] = concentration of the reactant at any time (t)

[A]₀ = initial concentration of the reactant (i.e. before the reaction starts)

k = rate constant

• Unit for k is concentration per unit time (e.g.

First-Order Reaction

Consider the following reaction:

A first-order reaction has the following differential rate law:

Looking at the equation, doubling the concentration of the reactant doubles that rate if reaction.

The corresponding integrated rate law is given by:

Where:

[A] = concentration of the reactant at any time (t)

[A]₀ = initial concentration of the reactant (i.e. before the reaction starts)

e = mathematical constant

The unit for k is inverse time (e.g. )

Second-Order Reaction

Consider the following reaction:

A first-order reaction has the following differential rate law:

The equation suggests that doubling the concentration of the reactant increases the rate of reaction by a factor of 4.

The corresponding integrated rate law is given by:

Where:

[A] = concentration of the reactant at any time (t)

[A]₀ = initial concentration of the reactant (i.e. before the reaction starts)

The unit for k is inverse conservation times inverse time (e.g. )