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 atomsAtoms of the same element have the same mass and the same chemical propertiesAtoms combine in whole number ratios to form compoundsChemical 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 solutiono 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 subscripts2.) Use coefficients in front of chemicals to balance atoms that appear only

once per side3.) 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 usedo Limiting reactant: used up first. Reaction stops when limiting reactant in

goneo Excess reactant: some remains after reaction stopso 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 e-

Cathode 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.gifo α (+) particles shot through gold foil in front of phosphorescent screeno Observed: one large glowing spot opposite the α emitter and several tiny

flickering spots all over screeno 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/nf

2) 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.gifo Gamma rays: high E, very toxic, produced by black hole collisions and during

radioactive nuclear decayo X-rays: high E, toxic in high doses, produced when high E e- collide with

metal targeto UV and visible: caused by e- dropping to lower E levelso Infrared: emitted by vibrating chemical bondso Microwaves: produced by circulating electrical field, causes molecular

rotationo 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 e-

travels 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/nfinal

2 – 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 s1 p2 d3 f

Value of lalso specifies number of planar nodes in the orbital in any given subshell

MAGNETIC QUANTUM NUMBER Magnetic (ml) = any integer between - land 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/ExplorationGuides/images/EL_MSCH_ElectronConfig1.gif

o Some exceptions to diagonal ruleo e- may jump subshells to fill a nearly full subshello 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 e-

o If all e- are paired: diamagnetico 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

similarlyo 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 configurationo 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 cationo Li+ N3- Li3No Co2+ Cl- CoCl2

o Cr4+ O2- CrO2

Naming binary ionic compoundso Cation gets element name and comes firsto Anion is element name ending in –ide and follows cation nameo No numeric prefixes are usedo Transition metal cation charges are indicated with a roman numeralo 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- shellso Larger amounts of e-/e- repulsiono Atom size swells

Going across row, r decreaseso Filled inner e- shells screen out the full attractive effect of the (+) nucleuso 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 e-

o 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/molo Li+ (g) Li2+ (g) + e- IE2 = 7298 kJ/molo 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 decreaseso Outermost e- is farther away from nucleus, making it easier to removeo Explains increasing reactivity LiNaKRbCs

Going across row, IE1 increaseso r decreases from L to Ro 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- X-

o 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

Bondingo Covalent: equal sharing of e-

o Polar covalent: e- attracted towards more X atomo 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 central3.) Connect outer atoms to central using single bonds (2 e- each)4.) Place remaining e- as lone pairs around outer atoms to satisfy octet rule5.) Place any remaining e- around central atom as lone pairs or single dot, follow

octet rule for 2nd row elements6.) Compute formal charges (FC) = (original # VE) – (# bonds) – 2(# lone pairs)

Sum of individual atomic FC’s must equal overall charge7.) If central atom has (+) FC, shift lone pairs from outer atoms in to form

multiple bonds, follow octet rule for 2nd row elements8.) 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-metalo d block element- indicate charge as Roman numeralo anion (non-metal) name, -ide ending

Polyatomic ionso Charged groupings of non-metal atoms held together with covalent bondso Do not break down further in aqueous soutiono 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 chlorateo Common polyatomic ions (taken from La Duca R. (2006). Polyatomic

ions.Acetate CH3COO- Ammonium NH4

+

Carbonate CO32- Dichromate Cr2O7

2-

Chlorate ClO3- Hydroxide OH-

Chromate CrO42- Nitrate NO3

-

Cyanide CN- Permanganate MnO4-

Cyanate CNO- Phosphate PO43-

Sulfate SO42-

o 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 shape)

1 lone pair 2 lone pairs 3 lone pairs

2 Linear 180° - - -3 Trigonal planar 120° Bent - -4 Tetrahedral 109° Pyramidal Bent -5 Bipyramidal 90°, 120°,

180°See-saw T-shaped Linear

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 fieldo Nonpolar will not respond to an electric field

Polarity’s significanceo 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 presento 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 shape2.) Along polar bonds, draw vectors toward more electronegative atoms3.) 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 hybridizationo Equivalent hybrid orbitals can be created by mixing the s, p, and d orbitals

on an atomo The number of hybrid orbitals is always the same as the number of atomic

orbitals that are mixed to create the hybrid orbital seto Hybridization by number of e- groups:

e- groups Hybridization Left over orbitals2 sp 2p3 sp2 p4 sp3 (5d)5 sp3d 4d6 sp3d2 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 overlapo C can make four bondso C can link into long chains or ringso There is a 3-D geometry around C

Organic compounds make up all living material Organic compounds are used to make medicines and polymers

NAMING1.) 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 prefix4.) 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 chiralo 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 halideo 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) Alkaneo 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 halideo Alkylation: Aromatic + RX (AlCl3 catalyst) alkyl aromatico 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 groupo Oxidation: Primary alcohol + K2Cr2O7(aq) Aldehyde

Secondary alcoholso Have two C’s attached to the C attached to the alcohol groupo Oxidation: Secondary alcohol + K2Cr2O7 Ketone

Tertiary alcoholso Have three C’s attached to the C attached to the alcohol groupo Cannot be oxidized

General reactionso Deprotonization: alcohol + alkali metal metal alkoxideo 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 alcoholo 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 carboxylateo Ester Synthesis: carboxylic acid + alcohol (strong acid catalyst) ester

ESTERS

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

Protonation: amine + acid alkylammonium salt Condensation: amine + carboxylic acid amide

POLYMERS Polymers are built from monomers, smaller repeating unit of a polymer Polymers are used in plastics and other technological materials Bipolymers make up cellulose, proteins, and genetic material 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 Study of heat changes during physical or chemical processes Thermal E: E associated with random motion of atoms and molecules Law of Conservation of Energy

o “First law of thermodynamics”o States that total E in the Universe is constanto E can neither be created nor destroyed, just changed from one form to

another

HEAT 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 Heat transfer (q) Enthalpy (H): heat transfer in an open system (constant pressure system) Change in enthalpy: ΔH = Hfinal - Hinitial (unit: kJ / mol) In an exothermic reaction:

A B + heatHA > HB

ΔH = HB – HA < 0 In an endothermic reaction:

A + heat B HB > HA

ΔH = HA – HB > 0

COMPUTING HEAT CHANGES Heat change (q): q = mcsΔT

o m = masso 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

Phase changeso There is no temperature change (ΔT = 0) during phase changeo All added heat E increases motion of moleculeso 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)

Thermal equilibrium: a condition in which the system and its surroundings are at the same temperature and heat transfer stops

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/molo 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 0o 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 stepso 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 temperatureo 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