AN INTRODUCTION TO CHEMISTRY Science 2009 – 2010 Academic Decathlon

In this section, we will cover:Chemistry prior to the Scientific RevolutionAntoine Lavoisier and the Birth of Modern ChemistryChemistry After LavoisierTen Independent Research Topics, including Mixing

Metals and Radioactivity

A Brief History of Chemistry

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

In this section, we will cover:Atomic Theory and StructureChemical Bonding and Intermolecular ForcesMolecular ModelsNuclear ChemistryTen Independent Research Topics, including

Electronegativity and Fission and Fusion Reactions

The Structure of Matter

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 2

1H (Deuterium): Tracer isotope Fusion reaction with tritium

146C:

Radiocarbon dating Climate change studies

6027Co:

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 (1875-1946)

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

High polarity → strong van der Waals forces

Stronger bonds Higher boiling and

melting points

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

90Th  +  4

2He2+

IRT: Decay Equations and Predicting Products of Decay – Beta and Positron Beta (beta-minus) decay:

Too many neutrons Sample:

32H → 3

1He + electron + antineutrino Positron (beta-plus) decay:

Too many protons Sample:

104C → 10

5B + positron + neutrino

IRT: Alpha Bombardment Reactions

Ernest Rutherford: 1919 Nuclear transformations can be caused

by bombardment (including alpha bombardment)

Example: 4

2He + 147N → 17

8O + 11H

IRT: Fission and Fusion Reactions

Example fission of uranium-235: 235

92U143 + neutron → 13454Xe80 + 100

38Sr62 + neutron + neutron

Products vary (typically amu of 130 and 100 plus 2-3 neutrons)

Hydrogen-2 and Hydrogen-3 fusion: 2

1H1 + 31H1 → 4

2He2 + neutron Not yet feasible for large-scale power

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

In this section, we will cover:Gases, Liquids and SolidsPhase DiagramsSolutionsFour Independent Research Topics, including Carbon

Dioxide and Raoult’s Law

States of Matter

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

molecules2. The molecules move very rapidly in all

directions but at different speeds3. No forces of repulsion or attraction exist

between the molecules4. 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 KOR

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 = sqrt((3kT)/m)

Root mean square speed of one mole: u = 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: P = ((nRT)/(V-nb)) – ((n2a)/V2) a and b: correction values for volume and

molecular attraction (smaller → more ideal) 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. Common compounds of group I and ammonium are soluble

II. Nitrates, acetates and chlorates are solubleIII. Binary halogens (not F) are soluble with metals,

except Ag, Hg(I) and PbIV. Sulfates are soluble, except barium, strontium,

calcium, lead, silver and mercuryV. Except for the first rule, carbonates, hydroxides,

oxides, silicates and phosphates are insolubleVI. Most sulfides are insoluble except calcium,

barium, strontium, magnesium, sodium, potassium and ammonium

Solutions: Aqueous Solutions

Maximum dissolved solute: saturated solution Lowering temperature

can bring crystals out of solution

Ions combining in solution to form insoluble particles → precipitates 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

In this section, we will cover:Acid-Base, Precipitation and Redox ReactionsElectrochemistryStoichiometryEquilibriumKineticsThermodynamicsFive Independent Research Topics, including

Electroplating and Hess’ Law

Reactions

Types of Reactions

Synthesis (combination) A + B → C or 2Na + Cl2 → 2NaCl

Decomposition A → B + C or 2H2O2 → 2H2O + O2

Double replacement AB + CD → AD + CB or Pb(NO3)2(aq) +

2KI(aq) → PbI2(s) + 2KNO3(aq)

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 Acids yield H+

Bases yield OH-

NH3: basic but with no OH-

Brønsted-Lowry Acids donate H+

Bases receive H+

Explains NH3 (it receives H+) 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:

acid + base → salt + water Salt product can be acidic, basic or neutral

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:

Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) →

AgCl(s) + NO3-(aq) + Na+(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 Cu → Cu2+ + 2e- E˚ = -0.34 V Ag+ + 1e- → Ag E˚ = +0.80 V (-0.34 + 0.80) = +0.46 V Spontaneous reaction

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: Zn(s) + 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)

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˚cellReaction under standard-state

conditionsNegative

>1 Positive

Favors products

0 1 0 EquilibriumPositive <1 Negati

veFavors 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