AN INTRODUCTION TO
CHEMISTRY
Science
2009 – 2010 Academic Decathlon
A Brief History of Chemistry
In this section, we will cover:
Chemistry prior to the Scientific Revolution
Antoine Lavoisier and the Birth of Modern Chemistry
Chemistry After Lavoisier
Ten Independent Research Topics, including Mixing Metals and
Radioactivity
Chemistry Prior to the Scientific
Revolution


Gold → copper → tin
and bronze
Iron:
 Meteorites?
 Mixed
with carbon to
form steel

Glass and pottery:
decoration, utility
IRT: Mixing Metals to Make Bronze





Bronze: 90% copper, also arsenic, tin, antimony,
lead
First used by Sumerians (3600 BCE)
Used for weapons, decoration
Methods: open casting, “lost-wax”
Superior Chinese alloys → effective defense
IRT: The Use of Dyes and Preservatives



Cave paintings and Egyptian tombs → Roman
Empire, Phoenicians, Minoan Crete
Woad, indigo, oxides of mercury, Tyrian purple
Mummy wrappings, stained glass, linen and
hemp
IRT: Alchemy and the “Philosopher’s
Stone”



Transmutations: base metals into gold
Practiced as a science from 331 BCE to roughly
300 CE
Philosopher’s Stone:
 Transmutations
 Elixir of
Life
IRT: Gunpowder and Fireworks



Saltpeter, charcoal, sulfur
Invented by Chinese before 1100 CE
Roger Bacon recipe: Opus Tertium
 Sent to


Pope
Rockets, projectiles → cannons
Battle of Crecy: 1346 CE
IRT: Early Thinkers on the Nature of
Matter

Aristotle:
 Ideas
from Plato (used term
“element”) et al
 Four properties: hot, cold,
wet, dry
 Four elements: fire, air, water,
earth
 Fifth element: ether

Democritus:
 Small
discrete particles
 Properties of these “atoms”?
Antoine Lavoisier and the Birth of
Modern Chemistry

Notable chemists 16th-early 19th
century:


Johann Baptista van Helmont,
Robert Boyle, Joseph Black, Henry
Cavendish, Joseph Priestly
Antoine Lavoisier: coherent
gathering of current theories
(nature of air, oxidation, water,
matter)

Involved in French Revolution,
targeted by Jacobins
IRT: The “Living Tree” Experiment




Johann Baptista van
Helmont
Living systems
Tree growing out of
“water onely” [sic]
Tree weight vs. soil
weight
IRT: Antoine Lavoisier and His Role and
Fate in the French Revolution




Born to a wealthy lawyer, studied accounting
and law
President of a bank, member of the Ferme
Generale (private tax collection agency)
Supported the new regime during/after
revolution
Targeted and executed → links to chemistry and
old regime
IRT: Madame Lavoisier





Marie Anne Pierrette Paulz married Antoine
Lavoisier in 1771
Father was in the Ferme Generale
Learned chemistry and English to assist in lab
Arrested and held for 65 days by Jacobins in
power
Remarried in 1805, then divorced, died alone
Chemistry After Lavoisier







Henri Becquerel: radioactivity
Pierre and Marie Curie: radioactive decay
J.J. Thompson: electron
Ernest Rutherford: atomic nucleus
James Chadwick: neutron
Niels Bohr: electron orbitals
Frederick Soddy: isotopes
IRT: Radioactivity and Nuclear
Structure




Henri Becquerel: radioactive
decay with photographic plates,
1896
Pierre and Marie Curie:
radioactivity and two new
elements (polonium, radium),
1898
Ernest Rutherford: alpha particles
and atomic structure, 1920
James Chadwick: neutron, 1932
Chemistry After Lavoisier







Albert Einstein: photoelectric effect
Louis de Broglie and Erwin Shroedinger:
quantum energy relationships
Periodic Table
Inter-/Intramolecular Forces
Dipoles
Heat, Work, Temperature
Reactants, Products, Chemical Kinetics
IRT: The Periodic Table and Associated
Periodicity





Dmitri Mendeleev
Repeating properties
among elements
Issues with ordering by
weight
Re-measuring and
skipping positions
helped
Henry Moseley: ordering
by atomic numbers
Wrap-Up



Both times of peace and war brought about
advancements in chemistry
Antoine Lavoisier and those like him were vital
to the development of modern chemistry
Chemistry since Lavoisier has developed rapidly
across many fields
The Structure of Matter
In this section, we will cover:
Atomic Theory and Structure
Chemical Bonding and Intermolecular Forces
Molecular Models
Nuclear Chemistry
Ten Independent Research Topics, including Electronegativity and
Fission and Fusion Reactions
Atomic Theory and Atomic Structure




Atomic structure dictates element chemical
behavior
Positive, negative, neutral particles
Weight of one atom determined by weighing
many atoms
Mass spectrometers: accuracy
IRT: Mass Spectrometry


Separates and
measures
compounds
Main components:
 Ion
source
 Mass analyzer
 Detector

Curved magnet or
cycling magnetic
field
Atomic Theory and Atomic Structure:
Mass and Isotopes





Atomic number: protons
Atomic mass: protons + neutrons
Same element with different numbers of
neutrons: isotopes
Carbon: atomic standard (12 amu)
Weighted averages:
 (isotope A
abundance x isotope A weight) +
(isotope B abundance x isotope B weight)
IRT: Properties and Importance of
Commonly Recognized Isotopes
 21
H (Deuterium):
 Tracer
isotope
 Fusion reaction with tritium
 146
C:
 Radiocarbon dating
 Climate change studies
 6027
Co:
 Highly radioactive:
kills cancer cells and bacteria
 Examines steel components
Atomic Theory and Atomic Structure:
Electrons





Absorption or emission
spectrum: determining
structure of an atom
Bohr Model of the atom:
fixed orbits
Quantum Mechanical
Model: non-fixed orbits
Electron clouds: orbits (s
and p)
Orbital shapes determine
bonding behaviors
IRT: Wave and Particle Nature of the
Electron and Photon



All matter exhibits both wave and particle
properties
Light as a particle: photoelectric effect
Electrons as energy: Davisson-Germer
experiment
Atomic Theory and Atomic Structure:
The Periodic Table


Number of orbitals determine period
Across a row (period):
 Atomic radius decreases
 Ionization energy
increases
 Electron affinity increases
Atomic Theory and Atomic Structure:
The Periodic Table

Down a column (group or family):
Atomic radius increases
 Ionization energy decreases
 Electron affinity decreases

IRT: Electronegativity





One atom’s attraction of electrons from the
adjacent atom to which it is chemically bonded
Higher value = greater attraction
Increases up a group and across a period
Fluorine → most strongly electronegative
Values predict “winners”
Chemical Bonding and Intermolecular
Forces: Intramolecular Forces

Ionic:
 Electron
transfer
 NaCl

Covalent:
 Sharing
electrons
 CH4

Metallic:
 Electron sea
 Brass
Chemical Bonding and Intermolecular
Forces: Intermolecular Forces


Van der Waals force: uneven distribution of
positive and negative charges (temporary or
permanent)
Hydrogen bonds: strongly electromagnetic
atom bonded to hydrogen on another
molecule
IRT: The Importance of Hydrogen
Bonding in Living Systems


DNA contains
hydrogen,
oxygen and
nitrogen
Hydrogen bonds
in DNA create
its double helix
structure
Chemical Bonding and Intermolecular
Forces: Effects and Properties of Bonds

Solid structures:
 Ionic
lattice
 Covalent network or molecular solid



Translational motion
Strength of force determines state at room
temperature
Uneven bonds are polar
Molecular Models: Lewis Structures


G.N Lewis (18751946)
Lewis Structures
 Dots represent
electrons
 Valence electrons
(bonding)
 Bonding pairs and
non-bonding
(“lone”) pairs
Valence Bonds and Hybridization

Single bond
 One overlap
between orbitals

Double-bond or
triple-bond
 Multiple overlaps

Hybridization
 Different orbital
shapes combine to
form a new shape
IRT: The Formation of Molecular Orbitals


Orbitals are electron waves in particular
positions and shapes
Sigma (s) orbitals
 Overlap concentrated
along an imaginary
connecting line

Pi (p) orbitals
 Overlap concentrated
away from connecting line
IRT: The Formation of Molecular Orbitals




N2: one sigma and two pi bonds
O2: one sigma and one pi bond
F2: one sigma bond
CO2: one sigma and one pi bond for each oxygen
atom
Molecular Models: VSEPR Models



Valence Shell Electron Pair Repulsion model
Three dimensions
Molecular geometry (tetrahedron, linear, et al)
IRT: The Resonance Concept Model



Explains bond properties in mathematically uneven
bonds
Sharing and distributing electrons to satisfy the octet
O3 and SO3
Molecular Models: Oxidation States




Assigned based on electron loss/gain
H2O: H = +1 O = -2
Sum of oxidation numbers in neutral molecular
equals zero
Sum of oxidation numbers in charged molecule
equals total charge
Molecular Models: Dipole Moments and
Polarity

Dipole moment
 Lack
of symmetry
 Bond dipoles do not
cancel each other out

Polar molecules
polarity → strong
van der Waals forces
 Stronger bonds
 Higher boiling and
melting points
 High
Nuclear Chemistry

Radioactive atoms
 Unstable
nuclei (varying ratios of neutrons to
protons)
 Regain stability through various pathways



Alpha decay: loss of helium nucleus
Beta decay: neutron → proton
Positron decay: proton → neutron
IRT: Decay Equations and Predicting
Products of Decay – Alpha

Alpha decay
 Very
large nuclei
 Atoms of bismuth and
those larger

Sample:
 23892U
→
234 Th
90
+
4 He2+
2
IRT: Decay Equations and Predicting
Products of Decay – Beta and Positron

Beta (beta-minus) decay:
 Too
many neutrons
 Sample:
 3 2H

→ 31He + electron + antineutrino
Positron (beta-plus) decay:
 Too
many protons
 Sample:
 104C
→ 105B + positron + neutrino
IRT: Alpha Bombardment Reactions



Ernest Rutherford: 1919
Nuclear transformations can be caused by
bombardment (including alpha bombardment)
Example:
 42He
+ 147N → 178O + 11H
IRT: Fission and Fusion Reactions

Example fission of uranium-235:
+ neutron → 13454Xe80 + 10038Sr62 + neutron +
neutron
 Products vary (typically amu of 130 and 100 plus 2-3
neutrons)
 23592U143

Hydrogen-2 and Hydrogen-3 fusion:
→ 42He2 + neutron
 Not yet feasible for large-scale power
 21H1 + 31H1
Wrap-Up




Various notations and models are used to
express and explain atomic structure and bonds
Bonds vary in composition, type, structure and
polarity
Lewis and VSEPR models help visually express
molecular orientation and geometry
Nuclear chemistry involves radioactivity and
decay reactions of various types
States of Matter
In this section, we will cover:
Gases, Liquids and Solids
Phase Diagrams
Solutions
Four Independent Research Topics, including Carbon Dioxide and
Raoult’s Law
Gases: Laws of Ideal Gases




Boyle’s Law: P x V = a constant (C)
Charles’ Law: V/T = a constant (D)
Combination: PV/T = CD
Tracking changes:
 (P1V1)/T1 =
(P2V2)/T2
IRT: Partial Pressures and Correction
of Gas Volumes Collected Over Water


Gas proportions in mixtures → expressed in
mole fractions
Dalton’s Law:
 Mole

fraction A = Pressure of A / Total Pressure
Gas container over water
 Water vapor pressure relies only on
temperature
 Total pressure – water vapor pressure = gas pressure
Gases: Kinetic Molecular Theory

Four major assumptions about ideal gases:
1. A pure gas consists of tiny, identical molecules
2. The molecules move very rapidly in all directions
but at different speeds
3. No forces of repulsion or attraction exist between
the molecules
4. Gas pressure is a result of collisions of the
molecules with the walls of the container (no loss
of energy)
Gases: Particle Speed






Average molecule speed (u) determines
frequency of collisions with given side length (l)
Momentum change from collisions determines
force
Molecule mass = m
Force = (mu2)/l
Number of molecules = N
Pressure = (1/3)((Nmu2)/V) or PV = (1/3)Nmu2
Gases: Avogadro’s Law



Number of molecules determines gas behavior
Mass → less important
Given temperature, pressure and volume →
same number of molecules
Gases: Volume and Mass of One Mole

One mole:
 Number of
molecules in a volume of 22.4 liters at 1
atmosphere pressure at 273 K
OR
 Number of atoms in 12 grams of carbon-12


Avogadro’s number: 6.022 x 1023 molecules
Molar mass is g/mol
Gases: Root Mean Square Speed

Average single molecule’s speed:
u

Root mean square speed of one mole:
u

= sqrt((3kT)/m)
= sqrt((3RT)/M)
R is the Boltzmann constant recomputed for
one mole of gas (“universal molar gas constant”)
IRT: The Behavior of Gases Under
Extreme Conditions


High pressure, low volume and low temperature
→ gases do not behave ideally
Van der Waals’ formula to predict non-ideal gas
properties:
= ((nRT)/(V-nb)) – ((n2a)/V2)
 a and b: correction values for volume and molecular
attraction (smaller → more ideal)
P

Large van der Waals values make for ideal
refrigerator coolants
Gases: The Ideal Gas Equation



For one mole: pressure x volume = R (universal
molar gas constant) x temperature (in Kelvin)
For n number of moles:
Related to the combination of Boyle’s and
Charles’ Laws
Gases: Relative Rates of Diffusion and
Effusion




Diffusion: gas spreading out from a source
Effusion: gas escaping from a small hole
Impossible to determine in non-vacuum
environment
Relative speeds can be determined
 Heavier
(more massive) molecules move slower
Liquids

Intermediate between gas and solid:
 Some
intermolecular forces, translational motion
 Moderate degree of order
Liquids



Long-range ordering (depends on qualities of
liquid)
Water is more ordered than other liquids like
octane (stronger forces)
Intermediate density (between gas and solid)
Solids


Solids are highly ordered
Types:
 Ionic
lattice
 Covalent network
 Molecular
 Metallic

Some substances exist in multiple forms
(allotropes)
Solids

Carbon: many different
bonding arrangements
 Graphite: stable at room
temperature
 Diamond: formed when
graphite is under high
pressure
 Can


be created in labs
Particle size affects structure
Closely-packed particles
have strong bonds
Solids: Properties of Metals

Simple metallic structures:
 Body-centered
cubic (shown)
 Cubic closest packed
 Hexagonal closest packing

Properties of metals:
 Lustrous
 Good
conductors of heat and electricity
 Sonorous
 Malleable
 Ductile
Phase Diagrams: Concepts




1. Constructed assuming a sealed container
2. Dynamic transfer
3. Equilibrium
4. Vapor (gas) present at any temperature
Phase Diagrams: Features
Phase Diagrams: Water


Backward-sloping
line between solid
and liquid states
Gives ice and
liquid water
unique properties
IRT: Carbon Dioxide

Liquid CO2:
difficult to
observe
 High
pressure
and low
temperature

Supercritical CO2:
industrial solvent
Solutions: Concepts



Solubility: how much of a solute will dissolve
Concentration: relative amounts of solute in a
solution
Physical properties: some occur when solutions
are formed
Solutions: Types and Factors

“Like dissolves like”:
 Water
(polar) with salt or sugar
 Octane (non-polar) with vegetable oil


Strong reaction with water: hydration
Solubility : the relationship between
intermolecular forces and forces trying to break
molecules apart
Solutions: Solubility Rules
I.
II.
III.
IV.
V.
VI.
Common compounds of group I and ammonium are soluble
Nitrates, acetates and chlorates are soluble
Binary halogens (not F) are soluble with metals, except Ag,
Hg(I) and Pb
Sulfates are soluble, except barium, strontium, calcium,
lead, silver and mercury
Except for the first rule, carbonates, hydroxides, oxides,
silicates and phosphates are insoluble
Most sulfides are insoluble except calcium, barium,
strontium, magnesium, sodium, potassium and ammonium
Solutions: Aqueous Solutions

Maximum dissolved solute:
saturated solution


Ions combining in solution to
form insoluble particles →
precipitates


Lowering temperature can
bring crystals out of solution
Stalactites and stalagmites
Compounds with O-H bonds
dissolve in water (glucose)
Solutions: Organic Solvents



Often contain only carbon and hydrogen
Used for grease and oil removal
Toxic to humans
 Disposed


by burning
Recent developments → modern soap and
detergent: interact with non-polar molecules
but are water-soluble
Supercritical fluids: solvents?
Solutions: Expressing Concentration

Percent Composition
 X grams of

a solute in Y grams of solvent (usually 100)
Molarity
 Moles of
solute per liter of solution
 Used in scientific applications

Molality
 Moles of
solute per kilogram of solvent
 Mole fraction
 Tracks colligative properties
IRT: Raoult’s Law and Colligative
Properties: Salts


Physical properties of a solution are relative to
number of moles of solute
Salts in water create larger than expected
changes
 NaCl
in water has twice the effect: two moles of
ions per mole of NaCl
 CaCl2: three moles of ions per mole of CaCl2
 Salts lower freezing point of water → deicing roads
 NaCl
is harmful to the environment so calcium
magnesium acetate has been proposed (et al)
IRT: Raoult’s Law and Colligative
Properties: Distillation of Water


Vapor above a solution is pure solvent
Distillation seeks to capture this vapor (in a
water-based solution) to collect drinking water
 Easier to scale up,
less setup and maintenance, less
waste

Reverse osmosis is the most viable alternative
 Water
is pressurized and pumped through
membranes that filter out impurities
 Lower energy needs, lower discharge water
temperature, purer output, smaller physical area
Wrap-Up



Gases, liquids and solids each have unique
properties that govern their behavior
Phase diagrams illustrate the transitions
between and conditions of these three states
These behaviors and conditions are important
in determining how substances will interact and
what the products of those interactions
(solutions) will be
Reactions
In this section, we will cover:
Acid-Base, Precipitation and Redox Reactions
Electrochemistry
Stoichiometry
Equilibrium
Kinetics
Thermodynamics
Five Independent Research Topics, including Electroplating and
Hess’ Law
Types of Reactions

Synthesis (combination)
A

Decomposition
A

+ B → C or 2Na + Cl2 → 2NaCl
→ B + C or 2H2O2 → 2H2O + O2
Double replacement
+ CD → AD + CB or Pb(NO3)2(aq) + 2KI(aq) →
PbI2(s) + 2KNO3(aq)
 AB
Types of Reactions

Single replacement


With metal: M + BC → MC + B
 Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
With non-metal: N + BC → BN + C
 Cl2 + 2KBr → 2KCl + Br2
Types of Reactions

Combustion
 Reactant
+ O2
 CH4 + 2O2 → 2H2O + CO2
 Produces heat and
sometimes light

Properties of substances
involved dictate the type of
reaction that will occur
Acid-Base Reactions: Theories

Arrhenius
H+
 Bases yield OH NH3: basic but with no OH Acids yield

Brønsted-Lowry
H+
 Bases receive H+
 Explains NH3 (it receives H+)
 Acids donate

Water can be an acid or base: amphoteric
Acid-Base Reactions: pH





pH = -log[H3O+]
0-14 scale
Below 7 is acidic, above 7 is basic
Exactly seven is neutral (like pure water)
All acidic and basic solutions have both acids
and bases in them
Acid-Base Reactions: Titrations


Titration: acids and
bases mixed together
and measured as they
interact
Endpoint or
equivalence point:
moles of acid and
base are equal
 Colored
indictor
shows this point
Acid-Base Reactions


Acids can be diprotic or triprotic
Double replacement reaction:
+ base → salt + water
 Salt product can be acidic, basic or neutral
 acid

Stronger acids transfer more hydrogen ions to
water
IRT: Acid-Base Reactions and Salts

Salt ions can interact with water: hydrolysis
 Can

produce basic, acidic or neutral solutions
Basic salt (sodium acetate) in water
 Weak

acetic acid in a basic solution
Acidic salt (ammonium chloride) in water
 Ammonia

(weak base) in an acidic solution
Neutral salt (sodium chloride) in water
 No reaction,
neutral solution
Precipitation Reactions

A type of double
replacement
reaction
 Two solutions
mixed → one of the
products comes out
of solution as a solid

Spectator ions: ions
not forming
precipitates
Precipitation Reactions: Example

Balanced reaction equation:
 AgNo3(aq)

+ NaCl(aq) → AgCl(s) + NaNO3(aq)
With ions separated:
+ NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(s)
+ NO3-(aq) + Na+(aq)
 Ag+(aq)

Net reaction with no spectators:
 Ag+(aq)
+ Cl-(aq) → AgCl(s)
IRT: Precipitates

Mercury
 Harmful
to people and the environment
 Industries have reduced output

Atmospheric particulates
 Harmful
inside the lungs
 Can be brought out of solution as precipitate

Silver
 Used
in solution to develop photographs
 Can be reclaimed and used for other purposes
Oxidation-Reduction Reactions



Oxidation: loss of electrons
Reduction: addition of electrons
Oxidation number
 Equal
to the number of electrons that must be added
or subtracted to make an element neutral
 Can be positive, negative or neutral
Oxidation-Reduction Reactions

Rules of oxidation states:
 Group
I elements are all +1
 Oxygen is -2
 Neutral atoms are 0, neutral compounds add up to 0
 Polyatomic ions must add up to the total charge

Electrons are conserved
 All

freed electrons must be used
Balanced equation example:
 Cu
+ 2Ag+ → Cu2+ + 2Ag
Electrochemistry: Terms



Electrochemistry
uses redox
reactions
Electroplating
(including
chromeplating)
Voltage: tendency
of electrons to leave
or join an atom
(cell potential)
Electrochemistry: Voltage

Voltage = potential of oxidation – potential of
reduction
 Positive values proceed
forward
 Negative values proceed in reverse

Cu + 2Ag+ → Cu2+ + 2Ag
→ Cu2+ + 2e- E˚ = -0.34 V
 Ag+ + 1e- → Ag
E˚ = +0.80 V
 (-0.34 + 0.80) = +0.46 V
 Spontaneous reaction
 Cu
Electrochemistry: Galvanic
Electrochemical Cell & Electrolysis


J. F. Daniell 1836
Earliest reliable
battery
 Anode:
oxidation
 Cathode: reduction

Electrolysis:
nonspontaneous
reaction with voltage
applied
IRT: Electroplating

Auto industry
 Chrome plating:

hardness, corrosion-/wear-resistance
Aerospace industry
 Gold
plating: non-reactive protection, reflectivity
 Platinum, palladium, nickel, copper, silver and
rhodium

Faraday: electric charge on one mole of electrons
 One
Faraday = 96,500 coulombs of charge
 Voltage used x coulombs needed = energy in
kilojoules
IRT: The Nernst Equation



Connects cell potentials to free energy changes in
chemical reactions
E = E˚ – RT ln Q/nF or E = E˚ – (0.0592 log Q)/n
Example:
+ Cu2+(aq) → Zn2+(aq) + Cu(s)
 +1.10 under standard conditions, n for Zinc is 2
 E = +1.10 – 0.0592/2 log [Zn2+]/[Cu2+]
 Equal concentrations of reactants and products
yields standard value (+1.10)
 Zn(s)
Stoichiometry


Balanced equations that keep track of
substances
Stoichiometry preserves ratios of substances
 Same principle used
in cooking and recipe
conversion

Applies to ion charges and redox reactions
Stoichiometry



Stoichiometry is used to determine yields
Limiting reactant: the substance in a reaction
that will determine how much one can yield
Example: 2H2 + O2 → 2H2O with 12g of H2 and
32g of O2
 12g/2
M = 6 moles of hydrogen
 32g/32 M = 1 mole of oxygen
 4 moles of hydrogen left over
 Oxygen is the limiting reactant
Equilibrium

Reactions do not always go in just one direction
 Forward

and reverse at same rate: equilibrium
Equilibrium constant: K = ([C][D])/([A][B])
 Ka
– Acids, Kb – Bases, Ksp – Precipitates, Kp –
Pressures of gases, Kc – Solutions and
concentrations
 If K > 1, there is more product in the end
 If K < 1, there is more reactant
Equilibrium

Conversion from Kc to Kp value:
 Kp


= Kc(RT)Δn
Smaller values of Ka and Kb mean weaker acids
and bases
Ksp indicates how much solid will ionize and
solubility of insoluble substances
 Small
Ksp values indicate a precipitate will form
Kinetics



Kinetics: how fast reactions happen and what
affects that rate
Rate law: algebraic equation determined by
concentrations and their effect on reaction rates
Rate is determined by change in concentration
over time
 Instantaneous rate can
be determined on a graph
Kinetics

Collision model:
conditions affect rate of
collisions (i.e. rate of
reaction)
 Increasing
temperature
increases rate
 Higher concentration
increases rate

Activation energy: energy
needed to activate the
reaction
Kinetics

Catalysts lower the
required activation
energy
 Catalyzed
reactions
require less energy and
are faster

Rates of chemical
reactions in the human
body use catalysts
called enzymes
Thermodynamics: Concepts




Thermochemistry measures energy changes in
chemical reactions
Thermodynamics: energy and temperature are
related to particle motion
System + surroundings = universe
State functions: volume, energy content and
pressure
Thermodynamics: Heat and Reactions
Exothermic reactions give
off heat
Endothermic reactions
absorb heat
Thermodynamics: First Law

Enthalpy: the energy content given off or taken
in by a chemical reaction (symbol H)
 Enthalpy
is a state function
 Directly proportional to the moles of a chemical
present


Heat of formation: enthalpy change during
formation of a compound
Measured by calorimetry
IRT: Hess’ Law



Germain Hess (1802-1850)
Heat energy in a chemical reaction is the same
no matter the number of steps
Unknown enthalpy values can be calculated
using other known enthalpy values
 If
ΔH is known for the formation of CO2, and for
the oxidation of CO to CO2, then ΔH for the
formation of CO can be calculated
Thermodynamics: Second Law

Entropy: energy associated with disorder
 State function
(symbol S)
 Smaller values indicate greater order


Whether or not a chemical reaction will occur
relies on both enthalpy and entropy
Gibbs Free Energy (state function, symbol G)
 ΔG
= ΔH – TΔS (T is temperature in Kelvin)
 Signs of terms determine spontaneity of reactions
Relationship of Change in Free Energy to
Equilibrium Constants and Electrode Potentials

Free energy to equilibrium constants:
 ΔG˚

= -RTlnK
Free energy to cell potential:
 ΔG˚
= -nFE˚cell
ΔG˚
K
E˚cell
Reaction under standard-state conditions
Negative
>1
Positive
Favors products
0
1
0
Equilibrium
Positive
<1
Negative Favors reactants
Wrap-Up




There are several categories of reactions, all of
which have different sub-categories (acid-base,
precipitation, redox)
Studies of electrochemistry (et al) have led to
industrial advances
Stoichiometry is invaluable to scientific work
An understanding of equilibrium, kinetics and
thermodynamic is vital to understanding how
and why reactions proceed as they do