LBS 171 REVIEW OUTLINE
INTRODUCTORY INFORMATION
Chemistry: the study of matter and its changes
Proust’s Law of Definite Proportions: any sample of a pure compound will have the same
elemental composition by mass.
DALTON

Law of Multiple Proportions: if elements A and B form two different compounds,
the ratio of the compounds’ mass ratios is a whole number or a simple fraction

Atomic Theory of Matter:
Elements are composed of tiny particles called atoms
Atoms of the same element have the same mass and the same chemical
properties
Atoms combine in whole number ratios to form compounds
Chemical reactions do not change the atoms, only the way they are
combined
ATOMS

Composed of:
-Protons (p+) -Neutrons (n) -Electrons (e-)

Element defined by number of p+ (z)

In a neutral atom, number of p+ = number of e
Mass number (A) = # p+ + # n

Atoms of same element with different A are isotopes

Atomic mass units: 12 atomic mass units (u) is the mass of one atom of 12C
THE MOLE AND ITS MEASURES

1 mole = amount of substance that has as many entities as 12 grams of 12C

Avigadro’s Number = 6.022 * 1023 = number of entities in 1 mole

6.022 * 1023 u = 1 gram

Molar mass of an element: mass of 1 mole of atoms in grams

Molar mass of a compound = sum of molar masses of all individual atoms in the
compound

Molar Mass (M)
o Used to measure moles of a substance in solution
o M = mol solute / L solution

In dilutions: M1V1 = M2V2
MICELLANEOUS MEASURES

Density: Mass of sample / volume of sample

Percent yield: Experimental mass of product / Theoretical mass of product
STOICHIOMETRY

“measuring elements”

Must use a balanced chemical equation

Balancing equations:
1.) Never change subscripts
2.) Use coefficients in front of chemicals to balance atoms that appear only
once per side
3.) Use coefficients (fractions are OK) to balance atoms that appear more than
once per side
4.) Multiply/divide to get the smallest whole number ratio

Sometimes, an excess of one reactant is used
o Limiting reactant: used up first. Reaction stops when limiting reactant in
gone
o Excess reactant: some remains after reaction stops
o Hint: if you are given a way to find moles of all starting materials, it is
probably a limiting reactant problem
ATOMIC STRUCTURE
THE e- IS CRITICAL IN CHEMISTRY

Sharing of e- between atoms- chemical bonding

Rearrangement of e- between atoms- chemical reactivity

Unpaired e- - magnetism

e- travel in a solid- conductivity
CATHODE RAY EXPERIMENT (late 1800’s)



http://www.lightandmatter.com/html_books/4em/ch01/figs/deflect.png
First evidence of eCathode ray attracted to a (+) electrical field and deflected by a (-) field, showed
that the particles had a (-) charge
ATOMIC MODELS

J.J. Thompson- late 1800’s proposed “plum pudding” model, hard e- buried in
spongy (+) material

Rutherford- 1911 gold foil experiment
http://abyss.uoregon.edu/~js/images/rutherford_exp.gif
o α (+) particles shot through gold foil in front of phosphorescent screen
o Observed: one large glowing spot opposite the α emitter and several tiny
flickering spots all over screen
o Shows that the atom is mostly empty space and has small, hard, (+)
charged nucleus
RYDBERG

Passed current through tube of low density H2 gas

Emission spectrum had four distinct colored lines for a reason unknown to him

Emperical equation for wavelength of each of the colored lines in H emission
spectrum: 1/λ = -R(1/ni2 – 1/nf2) where ni = 3,4,5,6, nf = 2, R = 1.097 * 107 m-1
ELECTROMAGNETIC RADIATION






http://www.hull.ac.uk/chemistry/spectroscopy/terms/wave.jpg
Frequency (ν) = number of peaks that pass by a given point in 1 second (1/s = Hz)
Wavelength (λ) = distance from peak to peak, usually in nm
c = λν, c = speed of light = 3.00 * 108 m/s
Ephoton = hν, h = Plank’s Constant = 6.626 * 10-34 J*s
Spectrum
http://acept.la.asu.edu/PiN/rdg/color/spectrum.gif
o Gamma rays: high E, very toxic, produced by black hole collisions and during
radioactive nuclear decay
o X-rays: high E, toxic in high doses, produced when high E e- collide with
metal target
o UV and visible: caused by e- dropping to lower E levels
o Infrared: emitted by vibrating chemical bonds
o Microwaves: produced by circulating electrical field, causes molecular
rotation
o Radio waves: caused by oscillating electrical fields
NIEL BOHR

Explained why only four lines were observed in H emission spectrum

e- orbits nucleus in circular path, only certain orbits with specific energies are
allowed

Electrical or thermal E can promote e- to higher E orbits

Light is emitted when e- drop down in orbits

E of light emitted corresponds only to the E gaps between the orbits that the etravels

Equation for energy of an e- in an H atom: E = -Rhc/n2 where n is orbit number
(principle quantum number as mentioned below)

From above equation: ΔE = Efinal – Einitial = -Rhc(1/nfinal2 – 1/ninitial2)

Energy of an e- has the same magnitude as the E it takes to remove that e- from
the atom (ionize)
QUANTUM NUMBERS
SCHRÖDINGER

Described electron as a wave not a particle

Created the wave function (ψ), which supports that the E of the e- is quantized

Ψ2 is related to the probability of finding the e- within a given region of space

To solve the wave function, the three integer quantum numbers-n, , and m-are
needed
PRINCIPLE QUANTUM NUMBER

Principle (n) = 1,2,3,…,

Primary factor in determining energy of an e
Defines size of an orbital- as n increases, so does the e-‘s average distance from
the nucleus

Two or more e- may have the same n

n = number of subshells in a shell, n2 = number of orbitals in a shell
ANGULAR MOMENTUM QUANTUM NUMBER

Angular (l) = 0,1,2,3,…,n-1


Determines subshell in which an e- resides
Each number corresponds to a different orbital shape or orbital type
Value of l Corresponding Subshell Label
0
1
2
3

s
p
d
f
Value of lalso specifies number of planar nodes in the orbital in any given
subshell
MAGNETIC QUANTUM NUMBER

Magnetic (ml) = any integer between - land l

Related to the orientation in space of the orbitals within a subshell

Number of m values for a given  equals the number of orbitals within that
subshell = 2 l+ 1
PAULI EXCLUSION PRINCIPLE

Principle: no two e- in a single atom may have the same exact set of quantum
numbers

4th QN-Spin Quantum Number (ms)
o Only permissible values are ½ and –½
o Therefore, only two e- can reside in a single orbital
ELECTRON CONFIGURATION
e- CONFIGURATION

Listing of e- populations in various subshells of an atom
Example: Fluorine 1s22s22p5

Represents ground state- e- placement for lowest atom E

Orbitals are still present when empty

e-/e- repulsion can occur, giving orbitals of a given n different E

Subshells are filled with e- from lowest E to highest E

Diagonal rule is used as a guide to subshell E ordering
http://www.explorelearning.com/ELContent/gizmos/ELScience_Deliverable/Explorat
ionGuides/images/EL_MSCH_ElectronConfig1.gif
o Some exceptions to diagonal rule
o e- may jump subshells to fill a nearly full subshell
o e- may jump subshells to half fill subshells
NOBLE GAS SHORTHAND

Noble gas configurations end at np6

Shorthand begins at previous noble gas

Only valence e- are seen in this notation- they are most critical for chemical
reactivity and physical properties

Examples: Ca [Ar]4s2, I [Kr]5s24d105p5
ENERGY LEVEL DIAGRAM

Each line represents an orbital

Orbitals are listed and filled in order of increasing E (using diagonal rule)

Orbitals of a subshell are labeled with n and spdf

Hund’s Rule (aka Macho Men on a Bus Rule): most stable e- configuration
maximizes the number of unpaired e
All attractive magnetic properties arise from unpaired eo If all e- are paired: diamagnetic
o If unpaired e- are present: paramagnetic (“magnets” like iron are
paramagnetic)
IONS
VALENCE ELECTRONS

Mendelev proposed first periodic table, grouping elements with similar reactivity
into columns

Elements in the same column have the same valence e- configuration and react
similarly
o Alkali Metals
-Na [Ne]3s1, K [Ar]4s1, Rb [Kr]5s1, Cs [Xe]6s1
-All ns1, very likely to lose 1 e- to obtain more stable noble gas
configuration in a reaction
-2Na + 2H2O  2NaOH + H2, 2K + 2H2O  2KOH + H2
o Halogens
-F [He]2s22p5, Cl [Ne]3s23p5, etc.
-All ns2np5, very likely to gain 1 e- to obtain more stable noble gas
configuration in a reaction
IONS



Ion: an atom or group of atoms that has lost or gained one or more electrons so
that it is no longer electrically neutral
Cation (+), Anion (-)
Main group (s & p blocks) elements tend to form ions in order to achieve noble
gas configuration

o Na ([Ne]3s1)  e- + Na+ ([Ne])
o N ([He]2s22p1) + 3e-  N3- ([Ne])
Transition metals (d block) form cations, but noble gas configuration is usually
not achieved
o Fe ([Ar]3d64s2)  2e- + Fe2+ ([Ar]3d6)
IONIC COMPOUNDS

Ionic compounds are overall electrically neutral compounds formed by the
combination of anion and cation
o Li+ N3-  Li3N
o Co2+ Cl-  CoCl2
o Cr4+ O2-  CrO2

Naming binary ionic compounds
o Cation gets element name and comes first
o Anion is element name ending in –ide and follows cation name
o No numeric prefixes are used
o Transition metal cation charges are indicated with a roman numeral
o Examples: Li3N – lithium nitride, CrO2 – chromium (IV) oxide
PERIODIC TRENDS AND CHEMICAL PROPERTIES
ATOMIC RADIUS

Atomic radius (r): size of atom

Going down a column, r increases
o More filled e- shells
o Larger amounts of e-/e- repulsion
o Atom size swells

Going across row, r decreases
o Filled inner e- shells screen out the full attractive effect of the (+) nucleus
o Going L to R, number of p+ increases while the inner filled e- shells
remain unchanged (thus screening effect does not change either)
o Outermost e- then feel a greater attraction to nucleus and atomic radius
decreases

Exception to trend: last column of d block elements are slightly larger than
expected – filled d subshell and greater e-/e- repulsion
IONIC RADIUS

Cation is smaller than its parent atom- fewer e-, less e-/e- repulsion

Anion is larger than its parent atom- more e-, more e-/e- repulsion

Isoelectronic cations have the same number of eo Example: Na+, Mg2+, Al3+
o Follow atomic radius trend by parent atom
IONIZATION ENERGY

Ionization energy (IE): energy required to remove an e- from a free atom in the
gas phase, measured in kJ/mol

IE1 represents first ionization energy, IE2 represents second ionization energy
o Li (g)  Li+ (g) + e- IE1 = 513 kJ/mol
o Li+ (g)  Li2+ (g) + e- IE2 = 7298 kJ/mol
o IE generally increases after each ionization- every time an atom loses an
e-, the remaining e- are pulled in closer to the nucleus and are held in with
a greater force

Going down a column, IE1 decreases
o Outermost e- is farther away from nucleus, making it easier to remove
o Explains increasing reactivity LiNaKRbCs

Going across row, IE1 increases
o r decreases from L to R
o outermost e- gets closer to nucleus and is harder to remove

Exception to trend: slightly easier to remove outermost e- from O than from N, as
doing so will create a half filled valence 2p shell in O
ELECTRON AFFINITY

e- affinity: E change that occurs when a neutral gas phase atom accepts an e
X + e-  Xo If X- is more stable than X, electron affinity will be (-), energy will be
released when X- is formed
o If X- is less stable than X, electron affinity will be (+), energy is required
to form X
Halogens have the most negative electron affinity, as accepting an e- creates noble
gas configuration

Noble gasses have the highest positive electron affinity, as accepting an e- breaks
noble gas configuration

There is no clear trend for electron affinity
ELECTRONEGATIVITY

Electronegativity (X): the ability of an atom, when in a compound, to attract e- to
itself

Critical to types of chemical bonding, intermolecular forces, and physical
properties

Going down a column, X decreases. Going across a row, X increases.

F has highest X = 4.0, Cs has lowest X = 0.7

Difference in X between two atoms (ΔX) determines type of bonding present
between them
o For metal non-metal compounds
 ΔX Є [0, 0.5) – covalent

 ΔX Є [0.5, 1.5) – polar covalent
 ΔX Є [1.5, 3] – ionic
o For non-metals only
 ΔX Є [0, 0.5) – covalent
 ΔX Є [0.5, 2] – polar covalent
Bonding
o Covalent: equal sharing of eo Polar covalent: e- attracted towards more X atom
o Ionic: total transfer of bonding e- from less X atom to more X atom
LEWIS DOT STRUCTURE
BASICS

Each valence e- depicted as a dot

Each bond, depicted as a dash, contains two valence e- shared between atoms

Works best for s and p block non-metal compounds

Octet rule: 2nd row elements want 8 valence e- to achieve noble gas configuration,
3rd row elements and higher with available d orbitals can expand octet
TECHNIQUE (taken from La Duca R. (2006). Drawing Lewis structures.)
1.) Count total valence electrons (VE)
2.) Place least X atom in center (unless it is H), and arrange others symmetrically
around central
3.) Connect outer atoms to central using single bonds (2 e- each)
4.) Place remaining e- as lone pairs around outer atoms to satisfy octet rule
5.) Place any remaining e- around central atom as lone pairs or single dot, follow
octet rule for 2nd row elements
6.) Compute formal charges (FC) = (original # VE) – (# bonds) – 2(# lone pairs)
Sum of individual atomic FC’s must equal overall charge
7.) If central atom has (+) FC, shift lone pairs from outer atoms in to form
multiple bonds, follow octet rule for 2nd row elements
8.) If FC cannot be completely minimized, leave (-) FC on more X atom
RESONANCE

Resonance is a means of representing the bonding when a single Lewis structures
fails to give an accurate picture- the true structure is a hybrid of the resonance
structures

Resonance structures differ only in the assignment of electron pair positions,
never in their atom positions

There is at least one multiple bond in every resonance structure

Molecules with resonance are more stable than expected
BOND ORDER

Bond Order (B.O.) = total # of bonds / # of connections

Applies to one resonance structure at a time

Can be used as a relative measure of bond length

o Double bonds are shorter than single bonds and longer than triple bonds
for the same set of atoms
o The higher the B.O. for each connection, the shorter the bond
Can be used as a relative measure of bond energy
o Triple bonds are stronger than double bonds, which are stronger than
single bonds between the same atoms
o The higher the B.O. for each connection, the stronger the bond
NAMING CHEMICAL COMPOUNDS
NAMING MOLECULAR COMPOUNDS

Non-metals only, not acids

Name less electronegative atom first

Name more electronegative atom second using –ide ending

Use numeric prefixes to indicate how many of each atom

Examples: H2O – dihydrogen monoxide, S4N4 – tetrasulfur tertranitride
NAMING IONIC COMPOUNDS

Binary ionic compounds
o Name metal first, then non-metal
o d block element- indicate charge as Roman numeral
o anion (non-metal) name, -ide ending

Polyatomic ions
o Charged groupings of non-metal atoms held together with covalent bonds
o Do not break down further in aqueous soution
o In compound name, place polyatomic ion name in its appropriate place
(cation/anion)
o Examples: CuSO4 – copper (II) sulfate, Na2CO3 – sodium carbonate,
Ca(ClO3)2 – calcium chlorate
o Common polyatomic ions (taken from La Duca R. (2006). Polyatomic
ions.
Acetate CH3COOAmmonium NH4+
Carbonate CO32Dichromate Cr2O72Chlorate ClO3
Hydroxide
OHChromate CrO42Nitrate
NO3Cyanide CN
Permanganate MnO4Cyanate CNOPhosphate
PO43Sulfate
SO42o Other anions can be named with the following rules (ibid)
 Removing or adding O atoms without changing the charge
 Adding an O atom to “ate” ion  per…ate
 Losing an O atom from “ate” ion  …ite
 Losing 2 O atoms from “ate” ion  hypo…ite


Adding an H+ to anions to form other ions
 Adding an H+ to an anion  hydrogen (anion name)
 Adding 2H+ to an anion  dihydrogen (anion name)
Substituting an S atom for an O atom without changing the
charge: +S, -O  thio(anion name)
NAMING ACIDS

Acid can be recognized as any neutral compound whose only cations are H+
Parent anion ending
Acid name
-ide
Hydro__ic acid
-ate
__ic acid
-ite
__ous acid
MOLECULAR STRUCTURE
VSEPR-VALENCE SHELL ELECTRON PAIR REPULSION

Theory: groups of e- on a central atom try to maximize the distance between
themselves

Predicts molecular geometry of main group p-block compounds

One e- group is any of the following: a single bond, a double bond, a triple bond, a
lone pair, a single lone electron

General shapes and bond angles
e groups
0 lone pairs (basis
1 lone pair
2 lone pairs
3 lone pairs
shape)
2
Linear 180°
3
Trigonal planar 120°
Bent
4
Tetrahedral 109°
Pyramidal
Bent
5
Bipyramidal 90°, 120°, See-saw
T-shaped
Linear
180°
6
Octahedral 90°
Square Pyramidal Square Planar
-
MOLECULAR POLARITY
POLARITY

Polarity: the ability of a substance to respond to an applied electric field
o Polar will respond to an electric field
o Nonpolar will not respond to an electric field

Polarity’s significance
o Polarity has a large effect on solubility
 Polar molecules can dissolve in polar solvents
 Nonpolar molecules can dissolve in nonpolar solvents
o Polarity also has a great effect on reactivity, especially in organic
chemistry
HOW TO DETERMINE POLARITY

Both of the following conditions must be true for a molecule to be polar
o Polar covalent bonds must be present
o Molecule must be asymmetric

Determine polarity of bonds using ΔX, for molecular compounds – polar covalent
if ΔX ≥ 0.5

To determine is molecule is asymmetric:
1.) Determine molecular shape
2.) Along polar bonds, draw vectors toward more electronegative atoms
3.) Using ΔX as relative vector values, apply vectors to the molecule
-If central atom moves, molecule is asymmetric
-If central atom does not move, molecule is symmetric
CHEMICAL BONDING
VALENCE BOND THEORY

Bonds form due to overlap of valence orbitals between two atoms

Each atom in bond supplies one orbital, with one e- in it, for overlap (covalent
bonding)
HYBRIDIZATION

Linus Pauling
o Knew that in CH4, simply overlapping an H with each of C’s 2p orbitals
would create 90° angles
o Realized that a C atom at its ground state has only two unpaired e-, not the
four that are needed to allow for four bonds
o Proposed orbital hybridization theory to describe bonding in CH4 and
others

Orbital hybridization
o Equivalent hybrid orbitals can be created by mixing the s, p, and d orbitals
on an atom
o The number of hybrid orbitals is always the same as the number of atomic
orbitals that are mixed to create the hybrid orbital set
o Hybridization by number of e- groups:
e- groups Hybridization Left over orbitals
2
sp
2p
2
3
sp
p
3
4
sp
(5d)
5
sp3d
4d
3 2
6
sp d
3d
o In molecules with two or more e- on the central atom, single bonds, σ
only, are formed by the overlap of a hybridized orbital of the central atom
and an orbital of a terminal atom
o Multiple bonds consist of 1 σ bonds and one or more π bonds. The π bonds
are formed by the overlap of unhybridized orbitals on two different atoms
ORGANIC CHEMISTRY
INTRO

Out of 15mil known chemical compounds, 14.9mil are organic

Because of a C base, many compounds can be made
o C is small and makes good orbital overlap
o C can make four bonds
o C can link into long chains or rings
o There is a 3-D geometry around C

Organic compounds make up all living material

Organic compounds are used to make medicines and polymers
NAMING
1.) Find longest uninterrupted C chain  parent name, consider functional groups
(discussed below)
2.) Number the chain from the end closest to any substituents (non-Hydrogen
groups) (Minimize sum of number positions of substituents and functional
groups/multiple bonds of alkenes/alkynes)
3.) State the position number of the substituent, name substituent as a prefix
4.) combine prefix and parent name
ISOMERS

Isomers have the same chemical formula but different structures

True isomers have unique names

Degree of Unsaturation (DOU) = (2+2(C)+N-H-X) / 2 = number of rings and/or π
bonds possible in any isomer of a given formula

Cis –trans isomers
o Cis isomers have identical groups attached to each C of a multiple bond on
the same side of the bond
o Trans isomers have identical groups attached to each C of a multiple bond
on opposite sides of the bond

Regioselective reactions favor the production of one isomer over all others
CHIRALITY

Enantiomers: isomers with handedness (non-superimposable images)
o Have same physical properties (except for rotation of plane polarized
light)
o Different enantiomers will react with other chiral molecules differently

Chiral molecules
o Have handedness
o Virtually all amino acids and sugars are chiral
o A molecule must have four different things attached to a single C atom to
be chiral (these C’s are stereocenters)
o Number of stereoisomers possible = 2n, n = number of stereocenters
The reactions in the following organic chemistry sections taken from
La Duca R. (2006). Organic reactivity
ALKANES






Have C-C and C-H single bonds only
General Formula: CnH2n+2
Saturated hydrocarbons
Name ends in -ane
Relatively unreactive
Reactions
o Halogenation: Alkane + Halogen (light)  Alkyl halide
o Combustion: Alkane + O2  CO2 + H2O
ALKENES





Each has a C=C double bond
General Formula: CnH2n
Name ends in –ene, indicate number position of double bond in carbon chain
More reactive than alkanes
Reactions
o Hydrogenation: Alkene + H2 (Pt catalyst)  Alkane
o Hydrohalogenation: Alkene + HBr  Alkyl Halide
ALKYNES





Each has a C≡C triple bond
General Formula: CnH2n-2
Name ends in –yne, indicate number position of triple bond in carbon chain
Similar reactivity to alkenes, slightly more reactive – more e- in triple bond
Reactions
o Hydrogenation: Alkyne + H2 (Pt catalyst) alkene (cis)
o Halogenation: Alkyne + X2  alkene dihalide (trans)
AROMATICS




Ring with alternating C-C single and C=C double bonds
Simplest – C6H6 benzene
Delocalized π overlap by unhybridized p orbitals on either side of the flat ring
makes the molecule more stable and less reactive than alkenes
Reactions
o Halogenation: Aromatic + X2 (FeX3 catalyst) aromatic halide
o Alkylation: Aromatic + RX (AlCl3 catalyst) alkyl aromatic
o Nitration: Aromatic + HNO3 (H2SO4) nitro-substituted aromatic
ALCOHOLS





Name ends in –ol, indicate number position of alcohol group in carbon chain
Primary Alcohols
o Have one C attached to the C attached to the alcohol group
o Oxidation: Primary alcohol + K2Cr2O7(aq)  Aldehyde
Secondary alcohols
o Have two C’s attached to the C attached to the alcohol group
o Oxidation: Secondary alcohol + K2Cr2O7  Ketone
Tertiary alcohols
o Have three C’s attached to the C attached to the alcohol group
o Cannot be oxidized
General reactions
o Deprotonization: alcohol + alkali metal  metal alkoxide
o Ester Synthesis: metal alkoxide + alkyl halide  ether
ALDEHYDES


Name ends in –al, C of aldehyde group is carbon-1 in carbon chain
Reactions
o Reduction: aldehyde + H-  primary alcohol
o Oxidation: Aldehyde + K2Cr2O7  carboxylic acid
KETONES



Name ends in –one, indicate number of C of ketone group in carbon chain
Cannot be oxidized
Reduction: Ketone + H-  secondary alcohol
CARBOXYLIC ACIDS


Name ends in –oic acid, C of carboxylic acid group is carbon-1 in carbon chain
Reactions
o Deprotonization: carboxylic acid + base  metal carboxylate
o Ester Synthesis: carboxylic acid + alcohol (strong acid catalyst) ester
ESTERS
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Naming: split RCO2– portion and the –R’ portion, –R’ portion is named normally
replacing ending with –yl. The acid part is named by replacing -oic ending of acid
with –oate. Example: CH3CH2CO2CH3 is named methyl propanoate
Saponification: ester (base catalyst)  carboxylic + alcohol
AMINES
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Protonation: amine + acid  alkylammonium salt
Condensation: amine + carboxylic acid  amide
POLYMERS
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Polymers are built from monomers, smaller repeating unit of a polymer
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Polymers are used in plastics and other technological materials
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Bipolymers make up cellulose, proteins, and genetic material
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Polymerization reactions
o Addition
 monomers add to each other with no stable byproducts
 Substitted polyalkanes via free radical initiator catalyst
o Condensation
 Two different monomers interact
 Stable small molecule byproduct is formed
 Dicarboxylic acid + diamine  polyamide
 Condensation of amino acids  polypeptides  proteins
THERMOCHEMISTRY
THERMOCHEMISTRY
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Study of heat changes during physical or chemical processes
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Thermal E: E associated with random motion of atoms and molecules
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Law of Conservation of Energy
o “First law of thermodynamics”
o States that total E in the Universe is constant
o E can neither be created nor destroyed, just changed from one form to
another
HEAT
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Heat: transfer of thermal E between two objects at different temperatures
System: the chemicals and their container
Surroundings: everything else
Universe (Ū): system + surroundings
Exothermic reactions release heat to surroundings
Endothermic reactions absorb heat from surroundings
HEAT TRANSFER
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Heat transfer (q)
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Enthalpy (H): heat transfer in an open system (constant pressure system)
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Change in enthalpy: ΔH = Hfinal - Hinitial (unit: kJ / mol)
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In an exothermic reaction:
A  B + heat
HA > HB
ΔH = HB – HA < 0
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In an endothermic reaction:
A + heat  B
HB > HA
ΔH = HA – HB > 0
COMPUTING HEAT CHANGES
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Heat change (q): q = mcsΔT
o m = mass
o cs = specific heat: amount of heat E required to raise the temperature of 1
g of substance by 1°C (unit: J / (g*°C))
o ΔT = Final temperature – initial temperature
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Phase changes
o There is no temperature change (ΔT = 0) during phase change
o All added heat E increases motion of molecules
o Heat change in phase change: q = nΔH
 n = number of moles of substance
 ΔH is constant and unique for each phase change of each
compound, i.e. ΔHfusion (H2O) = 6.01 kJ/mol
 Fusion: solid to liquid, freezing: liquid to solid,
vaporization: liquid to gas, condensation: gas to liquid,
sublimination: solid to gas
 If a reaction is reversed, ΔH is of the same magnitude but of
opposite sign (+, -) e.g. ΔHfusion(H2O) = -ΔHfreezing(H2O)
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Thermal equilibrium: a condition in which the system and its surroundings are at
the same temperature and heat transfer stops
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o Objects in contact at different temperatures will transfer heat to reach
thermal equilibrium
o On a macroscopic level, once thermal equilibrium is reached, no further
temperature change occurs throughout the entire system
Standard enthalpy change of formation (ΔHf°): enthalpy change when 1 mol of a
compound is made directly from its parent elements at standard conditions (25°C,
1 atm, 1.0M)
o Unit: kJ/mol
o ΔHreaction° = ∑ΔHf°(products) - ∑ΔHf°(reactants)
o In a reaction: aA + bB  cC + dD,
ΔHreaction° = [cΔHf°(C) + dΔHf°(D)] – [aΔHf°(A) + bΔHf°(B)]
o ΔHf° of pure elements in their natural state, such as O2, is always 0
o Hess’s Law: when reactants are converted to products, the ΔH is the same,
whether the reaction occurs in one step, or a series of steps
o Using Hess’s law, we can find ΔHf° for compounds that cannot be made
from their elements, such as DNA
ΔH of reactions
o Reactions can be performed in water within a calorimeter (an object that
loses very little energy, such as a Styrofoam container)
o Water will absorb/release heat and increase/decrease temperature
o Total energy change = 0 = qreaction + qwater + qcalorimeter (calorimeter
absorbs/releases such little heat that it can usually be assumed that
qcalorimeter = 0)
So: - (mass of water)(4.184J/g°C)(Tfinal – Tinitial) = qreaction