PRAYER TO THE HOLY SPIRIT

Breathe in me O Holy Spirit,

that my thoughts may all be
holy.
Act in me O Holy Spirit,
that my work, too, may be holy.
Draw my heart O Holy Spirit,
that I love but what is holy.
Strengthen me O Holy Spirit,
to defend all that is holy.
Guard me, then, O Holy Spirit,
that I always may be holy.

St. Augustine, pray for us.

With one mind and one heart, intent
upon God.
ATTENDANCE CHECKING
In an orderly manner, starting from the Boys,
then, Girls:

[Student Code]
[Last Name]
Say “Present”
GENERAL CHEMISTRY 2
SUBJECT TEACHER:
ENGR. REYFORD P. ORETA, CHE, RCHT
 KINETICS
§ EXPRESSING THE REACTION RATE
§ THE RATE LAW AND ITS COMPONENTS
§ INTEGRATED RATE LAWS: CONCENTRATION CHANGES
OVER TIME
§ THEORIES OF CHEMICAL KINETICS
§ REACTION MECHANISMS: THE STEPS FROM REACTANT TO
PRODUCT
§ CATALYSIS: SPEEDING UP A REACTION
CHEMICAL KINETICS
The study of how fast that change occurs, focuses on the
reaction rate, the change in the concentrations of
reactants (or products) as a function of time.
Faster reaction (higher rate), the reactant concentration decreases quickly,
where as in slower reaction (lower rate), it decreases slowly.
 FACTORS THAT AFFECT THE REACTION
RATE

1. Concentration: molecules must collide to react. The

higher the concentration of reactants, the greater the
reaction rate. Reaction rate is proportional to the
concentration of reactants.

Rate ∝ collision frequency ∝ concentration

2. Physical state: molecules must mix to collide. The

more finely divided a solid or liquid reactant, the greater
its surface area, the more contact it makes with other
reactant, and the faster the reaction occurs.
Substances must mix in order for particles to collide.
 FACTORS THAT AFFECT THE REACTION
RATE
3. Temperature: molecules must collide with enough
energy. The higher the temperature, the greater the
reaction rate.
§ Frequency of collisions. At higher temperature,
collisions occur more frequently, and so more
molecules react:
Rate ∝ collision frequency ∝ temperature
§ Energy of collisions. At higher temperature, more
sufficiently energetic collisions occur, and so more
molecules react
Rate ∝ collision energy ∝ temperature
REACTION RATE
Reaction rate is measured in terms of the changes in
concentrations of reactants or products per unit time.
Reactant concentrations decrease while product concentrations
increase.

For the general reaction: A → B

𝒄𝒉𝒂𝒏𝒈𝒆 𝒊𝒏 𝒄𝒐𝒏𝒄𝒆𝒏𝒕𝒓𝒂𝒕𝒊𝒐𝒏 𝒐𝒇 𝑨 𝒄𝒐𝒏𝒄 𝑨𝟐 &𝒄𝒐𝒏𝒄 𝑨𝟏 𝚫 (𝐜𝐨𝐧𝐜 𝐀)
Rate = - =- =-
𝒄𝒉𝒂𝒏𝒈𝒆 𝒊𝒏 𝒕𝒊𝒎𝒆 𝒕𝟐 & 𝒕𝟏 𝚫𝐭

The negative sign is used because the concentration of A is decreasing.

Concentration: [ ], indicate concentration in moles per liter.

𝚫 [𝐀]
Rate = -
𝚫𝐭
unit: mol / L . s
AVERAGE, INSTANTANEOUS AND INITIAL REACTION RATE
REACTION RATE IN TERMS OF REACTANT AND PRODUCT
CONCENTRATIONS

aA + bB → cC + dD

Where a, b, c and d are coefficients of the balanced equation:

𝟏 𝚫 [𝐀] 𝟏 𝚫 [𝐁] 𝟏 𝚫 [𝐂] 𝟏 𝚫 [𝐃]
Rate = - =- = =
𝒂 𝚫𝐭 𝒃 𝚫𝐭 𝒄 𝚫𝐭 𝒅 𝚫𝐭

Example: Reaction between hydrogen and iodine to form

hydrogen iodide:

H2(g) + I2(g) → 2HI(g)

Rate = ?
Sample Problem

Hydrogen gas has a nonpolluting combustion product

(water vapor). It is used as a fuel aboard the space shuttle
and in earthbound cars with prototype engines:

2H2(g) + O2(g) → 2H2O(g)

a) Express the rate in terms of changes in [H2], [O2] and

[H2O] with time.
b) When [O2] is decreasing at 0.23 mol/L.s, at what rate is
[H2O] increasing?
REACTION RATE IN TERMS OF REACTANT AND PRODUCT
CONCENTRATIONS

aA + bB → cC + dD

Where a, b, c and d are coefficients of the balanced equation:

𝟏 𝚫 [𝐀] 𝟏 𝚫 [𝐁] 𝟏 𝚫 [𝐂] 𝟏 𝚫 [𝐃]
Rate = - =- = =
𝒂 𝚫𝐭 𝒃 𝚫𝐭 𝒄 𝚫𝐭 𝒅 𝚫𝐭

Example: Reaction between hydrogen and iodine to form

hydrogen iodide:

H2(g) + I2(g) → 2HI(g)

Rate = ?
 RATE LAW AND ITS COMPONENTS: REACTION ORDERS

§ A reaction has an individual order “with respect to” or “in”

each reactant.
Reaction: A → products
§ If the rate doubles when [A] doubles, the rate depends on [A]1
and the reaction is first order with respect to A.
§ If the rate quadruples when [A] doubles, the rate depends on
[A]2 and the reaction is second order with respect to [A].
§ If the rate does not change when [A] doubles, the rate does
not depend on [A], and the reaction is zero order with respect
to A.
 RATE LAW AND ITS COMPONENTS: REACTION ORDERS

Rate = k [A]m [B]n …

§ k is the proportionality constant or the rate constant
§ Exponents m and n are reaction orders, how the rate is affected by reactant
concentration

§ A reaction has an individual order “with respect to” or “in” each

reactant, and an overall order, the sum of the individual orders.

A → products

§ First Order. The reaction is first order overall if the rate is

directly proportional to [A].

Rate = k[A]1 = [A]

 RATE LAW AND ITS COMPONENTS: REACTION ORDERS

Rate = k [A]m [B]n …

§ k is the proportionality constant or the rate constant
§ Exponents m and n are reaction orders, how the rate is affected by reactant
concentration

§ A reaction has an individual order “with respect to” or “in” each

reactant, and an overall order, the sum of the individual orders.

A → products

§ Second Order. The reaction is second order overall if the rate

is directly proportional to the square of [A].

Rate = k[A]2
 RATE LAW AND ITS COMPONENTS: REACTION ORDERS

Rate = k [A]m [B]n …

§ k is the proportionality constant or the rate constant
§ Exponents m and n are reaction orders, how the rate is affected by reactant
concentration

§ A reaction has an individual order “with respect to” or “in” each

reactant, and an overall order, the sum of the individual orders.

A → products

§ Zero Order. The reaction is zero order overall if the rate is not
dependent on [A].

Rate = k[A]0 = k(1) = k

 INDIVIDUAL AND OVERALL REACTION ORDERS

2NO(g) + 2H2(g) → N2(g) + 2H2O(g):

Rate = k[NO]2[H2]

§ The reaction is second order with respect to NO, first order

with respect to H2 and third order overall.

§ Note that the reaction is first order with respect to H2 even

though the coefficient for H2 in the balanced equation is 2.

Reaction orders must be determined from experimental data and

cannot be deduced from the balanced equation.
Problem: Determining Reaction Orders from Rate Laws

For each of the following reactions, use the given rate law
to determine the reaction order with respect to each
reactant and overall order:

(a) 2NO(g) + O2(g) → 2NO2(g) rate = k[NO]2[O2]

(b) CH3CHO(g) → CH4(g) + CO(g) rate = k[CH3CHO]3/2

(c) H2O2(aq) + 3I−(aq) + 2H+(aq) →I3−(aq) + 2H2O(l)

rate = k[H2O2][I−]
 DETERMINING REACTION ORDERS BY CHANGING THE
REACTANT CONCENTRATIONS
Reaction of oxygen and nitrogen monoxide, a key step in the formation of acid rain and in
the industrial production of nitric acid.

O2(g) + 2NO(g) → 2NO2(g)

Rate = k [O2]m [NO]n

 DETERMINING THE RATE CONSTANT

Reaction of oxygen and nitrogen monoxide, a key step in the

formation of acid rain and in the industrial production of nitric acid.

O2(g) + 2NO(g) → 2NO2(g)

Rate = k [O2][NO]2
STEPS FOR STUDYING THE KINETICS OF REACTION
 INTEGRATED RATE LAWS:
CONCENTRATION CHANGES
OVER TIME
• FOR GENERAL FIRST ORDER REACTION, A → B

𝚫[𝐀]
• RATE = -
𝚫𝐭

• RATE = K [A]
DETERMINING THE REACTANT CONCENTRATION AFTER
A GIVEN TIME

Sample Problem:

At 1000 degree C, cyclobutene (C4H8) decomposes to two molecules of ethylene (C2H4), in

first order reaction with the very high rate constant if 87 s -1.

A. The initial C4H8 concentration is 2.00 M. What is the concentration after 0.010 s?
B. What fraction of C4H8 has decomposed in this time?
 REACTION HALF LIFE

• THE HALF-LIFE (T1/2) IS THE TIME IT TAKES A GIVEN

REACTANT CONCENTRATION TO REACH HALF ITS INITIAL
VALUE.
DETERMINING THE HALF-LIFE OF A FIRST-ORDER
REACTION
ARRHENIUS EQUATION
Derived by Svante Arrhenius (1889), an equation to express the exponential relationship
between temperature and the rate constant.

!𝑬𝒂
𝒌= 𝑨𝒆 𝑹𝑻
§ k is the rate constant
§ e is the base of natural logarithms
§ T is the absolute temperature
§ R is the universal gas constant
§ A, frequency factor, A = pZ (collision frequency Z and the orientation
probability factor, p)

𝒌𝟐 𝑬𝒂 𝟏 𝟏
𝒍𝒏 = −
𝒌𝟏 𝑹 𝑻𝟏 𝑻𝟐
 DETERMINING THE ENERGY OF ACTIVATION

Problem

The decomposition of hydrogen iodide 2HI(g) → H2(g) + I2(g), has rate constant of 9.5x10-9
L/mol s at 500 K and 1.10x10-5 L/mol s at 600 K. Find the Ea.