Energy & Chemistry
ENERGY is the capacity to
do work or transfer heat.
HEAT is the form of energy
that flows between 2
objects because of their
difference in temperature.
Other forms of energy —
• light
• electrical
• kinetic and potential
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Energy & Chemistry
• Burning peanuts
supply sufficient
energy to boil a cup
of water.
• Burning sugar
(sugar reacts with
KClO3, a strong
oxidizing agent)
2
Energy & Chemistry
• These reactions are PRODUCT
FAVORED
• They proceed almost completely
from reactants to products, perhaps
with some outside assistance.
3
Thermodynamics
• Thermodynamics is the science of heat
(energy) transfer.
Heat energy is associated with
molecular motions.
Heat transfers until thermal equilibrium is
established.
4
Directionality of Heat Transfer
• Heat always transfer from hotter object to
cooler one.
• EXOthermic: heat transfers from SYSTEM to
SURROUNDINGS.
T(system) goes down
T(surr) goes up
5
Directionality of Heat Transfer
• Heat always transfer from hotter object to
cooler one.
• ENDOthermic: heat transfers from
SURROUNDINGS to the SYSTEM.
T(system) goes up
T (surr) goes down
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USING ENTHALPY
Consider the formation of water
H2(g) + 1/2 O2(g) --> H2O(g) + 241.8 kJ
Exothermic reaction — heat is a “product” and ∆H
= – 241.8 kJ
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USING ENTHALPY
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Making liquid H2O from H2 +
O2 involves two exothermic
steps.
H2 + O2 gas
H2O vapor
Liquid H2O
The Concept of Equilibrium
Chemical equilibrium occurs when a
reaction and its reverse reaction
proceed at the same rate.
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The Concept of Equilibrium
• As a system
approaches
equilibrium, both the
forward and reverse
reactions are
occurring.
• At equilibrium, the
forward and reverse
reactions are
proceeding at the
same rate.
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12
A System at Equilibrium
Once equilibrium is
achieved, the
amount of each
reactant and
product remains
constant.
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A System at
Equilibrium
Concentrations become constant
Rates become equal
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Depicting Equilibrium
In a system at equilibrium, both the forward and
reverse reactions are running simultaneously. We
write the chemical equation with a double arrow:
What Does the Value of K Mean?
• If K >> 1, the reaction
is product-favored;
product predominates
at equilibrium.
• If K << 1, the reaction is
reactant-favored;
reactant predominates
at equilibrium.
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The Reaction Quotient (Q)
• To calculate Q, one substitutes the initial concentrations
on reactants and products into the equilibrium
expression.
• Q gives the same ratio the equilibrium expression gives,
but for a system that is not at equilibrium.
If Q = K,
the system is at equilibrium.
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If Q > K,
there is too much product and the
equilibrium shifts to the left.
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If Q < K,
there is too much reactant, and the
equilibrium shifts to the right.
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Le Châtelier’s Principle
“If a system at equilibrium is disturbed by a change in
temperature, pressure, or the concentration of one of the
components, the system will shift its equilibrium position
so as to counteract the effect of the disturbance.”
Systems shift from “Q” towards “K”.
Stresses include:
temperature, pressure, volume, or concentration of a
system will result in predictable and opposing changes in the
system in order to achieve a new equilibrium state.
Equilibrium
and
Le Chatelier’s Principle
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Chemical Equilibrium
Reversible Reactions:
A chemical reaction in which the products
can react to re-form the reactants
Chemical Equilibrium:
When the rate of the forward reaction equals the rate of
the reverse reaction and the concentration of
products and reactants remains unchanged
2HgO(s)  2Hg(l) + O2(g)
Arrows going both directions (  ) indicates equilibrium in a chemical equation
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LeChatelier’s Principle
When a system at
equilibrium is placed
under
stress, the system will
undergo a change in such
a way as to relieve that
stress.
Henry Le Chatelier
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Le Chatelier Translated:
When you take something away from a
system at equilibrium, the system shifts in
such a way as to replace what you’ve
taken away.
When you add something to a system at
equilibrium, the system shifts in such a way
as to use up what you’ve added.
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LeChatelier Example #1
A closed container of ice and water at
equilibrium. The temperature is raised.
Ice + Energy  Water
The equilibrium of the system shifts to the
_____ right to use up the added energy.
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LeChatelier Example #2
A closed container of N2O4 and NO2 at
equilibrium. NO2 is added to the container.
N2O4 (g) + Energy

2 NO2 (g)
The equilibrium of the system shifts to the
left
______
to use up the added NO2.
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LeChatelier Example #3
A closed container of water and its vapor at
equilibrium. Vapor is removed from the system.
water + Energy  vapor
The equilibrium of the system shifts to the
right
_______
to replace the vapor.
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LeChatelier Example #4
A closed container of N2O4 and NO2 at
equilibrium. The pressure is increased.
N2O4 (g) + Energy  2 NO2 (g)
The equilibrium of the system shifts to the
leftto lower the pressure, because
_______
there are fewer moles of gas on that side of
the equation.
What Happens When More of a Reactant
Is Added to a System?
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The Effect of Changes in Temperature
Co(H2O)62+(aq) + 4 Cl(aq)
CoCl4 (aq) + 6 H2O (l)
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Catalysts increase the rate of both the
forward and reverse reactions.
Equilibrium is achieved faster, but the
equilibrium composition remains
unaltered.
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1.4 Potential Energy
Diagrams
Enthalpy
Chemical energy is most commonly converted to heat, we use the symbol, ΔH to
symbolize a change in energy available as heat. The symbol is sometimes read as
“delta H” or an enthalpy change.
Representing Combustion Reactions
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1.4 Potential Energy Diagrams
Energy of Chemical Bonds
Bond Energy
Endothermic
Exothermic
Representing the ΔH value associated with a chemical reaction.
Text Page 43
Review of Exothermic
► Reactants
Ep is higher than
Products Ep.
► Now, we must consider the
activation energy (the energy
needed so that the reactants
bonds will break and reform to
make product)
Review of Endothermic
► Reactants
Ep is lower
than Products Ep.
► Need to add more
energy to the system
for the forward
reaction to take place.
► Still need to consider
activation energy
Activated Complex
► Is
the short-lived, unstable structure formed during a
successful collision between reactant particles.
► Old bonds of the reactants are in the process of breaking,
and new products are forming
► Ea is the minimum energy required for the activation
complex to form and for a successful reaction to occur.
Activated Complex
► Is
the short-lived, unstable structure formed during a
successful collision between reactant particles.
► Old bonds of the reactants are in the process of breaking,
and new products are forming
► Ea is the minimum energy required for the activation
complex to form and for a successful reaction to occur.
Fast and slow reactions
► The
smaller the activation energy, the faster the reaction
will occur regardless if exothermic or endothermic.
► If there is a large activation energy needed, that means
that more energy (and therefore, time) is being used up for
the successful collisions to take place.
Practice:
1.
The following hypothetical reaction has an Ea of 120kJ
and a ΔH of 80kJ
2a + B  2C + D




Draw and label a potential energy diagram for this reaction.
What type of reaction is this?
Calculate the activation energy for the reverse reaction.
Calculate the ΔH for the reverse reaction.
2.
Analyze the
activation energy
diagram below.
►
What is the Ea for the
forward reaction? For
the reverse reaction?
What is the ΔH for the
forward reaction? For
the reverse reaction?
What is the energy of
the activated complex?
►
►
Answer to #2
►
The activation energy (Ea) for the forward reaction is shown by (a):
Ea (forward) = H (activated complex) - H (reactants) = 400 - 100 = 300 kJ mol-1
The activation energy (Ea) for the reverse reaction is shown by (b):
Ea (reverse) = H (activated complex) - H (products) = 400 - 300 = 100 kJ mol-1
The enthalpy change for the reaction is shown by (c):
H = H (products) - H (reactants) = 300 - 100 = +200 kJ mol-1
for the forward and reverse reaction.
Concentration of Solutions
►
►
Molarity: number of moles of solute in one liter of
solution
 Symbol is “M”
 If you have 1 mole of NaOH (40g) dissolved in enough
solvent (water) to make 1 liter, you will have 1 M
NaOH or one molar NaOH
 Molarity (M) = amount of solute (moles)
volume of solution (L)
Molality: number of moles of solute per kilogram of
solvent
 Symbol is “m”
 If you have one-half a mole of NaOH (20g) dissolved
in 1 Kg of water gives 0.5 m NaOH or one-half molal
of NaOH
 Molality (m) =
moles solute
mass of solvent (Kg)
Electrolyte: A solution or substance in solution
consisting of various chemicals that can carry
electric charges.
► Strong Electrolyte: Solutions conduct electricity
well.
 Ionizes completely: contains many ions
because of complete dissociation. (HCl)
 NaCl
Na+ (aq) +
(s)
Na+
Cl-
Cl-
Na+
Na+ Cl-
Cl-
(aq)
► Non-Electrolytes:
Solutions do not conduct
electricity
► Does not ionize: substance exists as dissolved
molecules in solution. Like sugar.
► Weak Electrolytes: Solutions conducts poorly
 Partially ionized: solution contains only a few
ions.
 Appears that only some of the substance has
dissociated or ionized. Equilibrium!!
+
 CH3COOH (aq) CH3COOCH
COO
+
H
3
(aq)
H+(aq)
CH3COOH
► Arrhenius
Acids and Bases
Acid: produces H+ ions, in aqueous solution
 H+ and H3O+ are the same thing, H can’t really lose its only e HCl(g) + H2O(l)
H3O+ (aq) + Cl- (aq)
► Arrhenius
Base: produces OH- ions, in aqueous solution
 referred to as alkaline
 KOH(s) + H2O(l)
K+(aq) + OH-(aq)
► B.L.
Acid: molecule or ion that is a proton donor
 a proton is H+
 HCl
+
NH3
NH4+
+
Cl-
► B.L.
Base: molecule or ion that is a proton acceptor
 H2O
+ NH3
NH4+
+ OH-
► Acids
Strength of Acids and Bases
and Bases are considered strong if they
completely ionize to form strong electrolytes
► Strong Acids: HCl, HNO3, H2SO4, HBr, HI, HClO4
 they easily lose their hydrogen ion
► Weak Acids: HF, HC2H3O2 (CH3COOH)
 They want to keep their hydrogen until someone
wants it more than them
► Strong Bases: those that are made up of hydroxide and
group 1 or 2 metal.
► Weak Bases: substances that do not contain hydroxide,
but rather generate hydroxide ions when reacting with
water
► Strong
acids and bases can be present in low
concentrations
► To measure concentration of a solution of an acid we use
Molarity, looking at the amount of H+ in a given volume.
► To measure strength of an acid we use pH, which measures
the amount of H+ ions.
► pH
of 4 could be a low concentration of a
strong acid or a high concentration of a weak
acid.
Indicators
► Acid-base
indicators are compounds whose
colors are sensitive to pH. (like
phenolphthalein)
► Indicators colors change in the presence of and
acid or base depending on the indicator
► Some indicators change color at low pH and
some at high pH
► Neutralization occurs when [H+] = [OH-]
pH and pOH
A neutral solution has a pH of 7.
pH scale : ranges from 0 to 14, where 7 is
neutral.
The lower the number, the more acidic (0 6)
The higher the number, the more basic (814)
pH: the measure of acidity
pOH: the measure of alkalinity or how
basic it is
pH + pOH = 14
pH = -log[H3O+]
pOH = -log[OH-]
Periodic Atomic Properties of the
Elements
Periodic law states that certain sets of physical
and chemical properties recur at regular
intervals when the elements are arranged
according to increasing atomic number
•Consider atomic radii: distance between the
nuclei of two atoms
•The distance between the nucleus and the
outer edge of the electron cloud
EOS
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
52
Ionic Radii
•The ionic radius of each
ion is the portion of the
distance between the nuclei
occupied by that ion
•If the size of an atom is
determined by the
outermost electrons, what
happens if you remove or
add an electron?
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Definitions
► Ionization
energy is the energy required
to remove an electron from a ground
state atom in the gaseous state
• to remove an electron, energy must be supplied to
overcome the attraction of the nuclear charge (endothermic,
always +)
► Electron
affinity is the energy change that
occurs when an electron is added to a
Take note: these refer to
gaseous atom
 How much an atom ‘likes’ Electrons (+ or -)
► the more negative it is the higher the EA
► Electronegativity
atoms Electronegativity
refers to molecules
is the ability of an
atom in a molecule to attract atoms to
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
54
Zeff & Shielding
 The order in which electrons are
assigned to subshells in an atom, as
well as other properties are because
of effective nuclear charge (Zeff)
 Shielding: electrons closer to the
nucleus screen or shield the effect of
nuclear charge on valence electrons
► the
number of shielding electrons
increases when you reach the end of the
period and go on to the next period.
► Shielding increases in steps as you start
a new period or go down a group
►Video
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
55
Atomic Radii
►
Atomic Radii decrease as atomic numbers
increase in an given period (going across).
 A proton and electron are added so the effective
nuclear charge increases because each proton has
more of an effect than each additional electron
►
►
As that attraction between the nucleus and electrons
increases, and the atomic radius decreases
Atomic Radii increase going down
 In going from top to bottom of a group, the valence
electrons are assigned to orbitals with increasingly
higher values of n (prin. Quantum number)
►
The underlying electrons requires some space, so the
electrons of the outer shell must be further (Your are
adding energy levels)
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Atomic Radii
►Zeff effective
nuclear charge: the nuclear charge experienced by
a particular electron in a multi-electron atom
 Adding protons Increases the attraction of the nucleus and pulls the
electron cloud closer to the nucleus resulting in a smaller
atomic radius
 (coulomb’s Law the force between two charged
particles is related to the product of their charges and
the distance between them.)
►Atomic
radii of transition metals trend a little differently
►Exceptions in atomic radii also exist in the lanthanide and
actinide series because of how the f subshells are uniquely filled
by electrons
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Transition Metal Atomic
Trends
►
►
From left to right across a period, the
radii initially decrease, then size
remains almost the same, then slightly
increases toward the end.
The small increase in atomic radii is
because of the d subshell is filled with
electrons and thus the ele-eletron
repulsions cause the size to increase
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Atomic Radii Properties
►
►
The increased number of energy
levels (n) increases the distance
over which the nucleus must pull
and therefore reduces the
attraction for electrons
Full energy levels provide
shielding between the nucleus
and valence electrons, so you
see an increase in shielding as
the level gets full
Illustration
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Ionic Radii
Cations are smaller than the
atoms from which they are
formed
– the nucleus attracts the remaining
electrons more strongly
Anions are larger than the
atoms from which they are
formed
– the greater number of electrons
repel more strongly
Think of the proton/electron ratio, -as electrons
are lost, the ratio of p+/e- increases and so the
electrons are held closer vv.
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Isoelectronic Configurations
Isoelectronic species are elements that all have
the same number of electrons
For isoelectronic species, the greater the
nuclear charge, the smaller the species
Effective nuclear charge
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Atomic and Ionic Radii
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Ionization Energy
Ionization energy is the energy required to remove an
electron from a ground state atom in the gaseous state
•to remove an electron, energy must be supplied to overcome the attraction of the nuclear
charge (endothermic, always +)
Continual removal of electrons increases ionization
energy greatly
+ + e–
–1
B

B
I
=
801
kJ
mol
It gets harder
with each
B+  B+2 + e– I = 2427 kJ mol–
electron removed
1 +2
+3 + e– I = 3660 kJ mol–
B

B
Because the
1
nucleus has a stronger
B+3  B+4 + e– I = 25,025 kJ
pull
–1
+4
mol
Illustration
B  B+5 + e– I = 32,822 kJ
EOS
mol–1
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
63
Ionization energy
► First
ionization energy- energy is increased with each
successive removal because the electron is being
removed from an increasingly positive ion
 The remaining electrons are held more tightly
 Notice the large jump at the 3rd level for Mg.
 There is a large increase as you remove electrons from
lower (inner) energy subshells
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First Ionization Energies
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Ionization energy
► Ionization
energy increases as atomic number
increases in any given period
 Zeff increases the attraction of the nucleus and holds the
electrons more tightly
group II to III, IE drops because the p electrons do not
penetrate the nuclear region as well as s electrons so aren’t as
tightly held
► Drop in IE also occurs between V & VI because of increased
repulsion created by the first pairing of electrons, that is stronger
than the increase in Zeff, lowering the energy required to remove
the electron
► Exceptions:
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
66
Ionization energy
► Ionization
energy decreases as atomic number increases
down a column or group
 The increased number of energy levels (n) increases the
distance over which the nucleus must pull, reducing the
attraction for electrons
 A full energy level provides some shielding between the nucleus
and valence electrons
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Electron Affinity
Electron affinity is the energy
change that occurs when an
electron is added to a gaseous
atom
-How much an atom ‘likes’
Electrons (+ or -)
-the more negative it is the higher the EA
(energy is flowing out of the system)
Electron affinities are
expressed as negative because
the process is exothermic
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Electronegativity
measure of the attraction of an atom for the pair of outer
shell electrons in a covalent bond with another atom
►A
 Pattern is same as electron affinity for same reasons
 Both are attraction nucleus has for electrons, one in forming an
ion (EA) and one in forming a molecule (EN)
 Fluorine is the most electronegative. The closer it is to fluorine,
the more electronegative it is.
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Metals, Nonmetals, and
Metalloids
Metals have a small number of electrons in their
valence shells and tend to form positive ions
Except for hydrogen
and helium, all sblock elements are
metals
All d- and f-block
elements are metals
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Metals, Nonmetals, and
Metalloids
Atoms of a nonmetal generally have larger
numbers of electrons in their valence shell than do
metals, and many tend to form negative ions
Nonmetals are all pblock elements and
include hydrogen and
helium
Metalloids have
properties of both
metals and nonmetals
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Metals
► Metals
react by losing electrons
 A loosely held electron will result in a more reactive metal
 This is tied directly to ionization energy
 With an increased # of energy levels (n), comes increased
distance from the nuclear attraction and thus a more loosely held
electron available for reactions
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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Non-metals
► Non-metals
tend to gain electrons, a strong nuclear
attraction will result in a more reactive non-metal
► This means that an atom with the highest Zeff and the least
number of energy levels should be the most reactive
nonmetal (F) because its nucleus exerts the strong pull
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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A Summary of Periodic Trends
Chapter 8: Electron Configurations, Atomic Properties and the
Periodic Table
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