Significant Figures

Ruler A – 4.40 cm

Ruler B – 4.4 cm

Add 1 more decimal than the smallest measurement can measure.

Type of 0 Significant?0.00003 placeholder no3,0000 placeholder no3,0003 trapped yes3.0000 precise yes

Adding sig figs – do not count the # of sig figs! Count the # of decimal places.

% error = I value acc - valueexpI x 100% valueacc

micro 106 m = 1 m

nano n 109 nm = 1 m

pico p 1012pm = 1 m

Substance Density Liq. Water 1.00 g/mLGas at STP mass of 1 mol/22.4L

Average atomic mass = % x mass + % x mass +….. (isotope 1) (isotope 2)

1 mole = 6.022 x 1023 particles. (Avogadro's number)

1 12C atom = 12.00 amu (atomic mass)

1 mole 12C atom = 12.00 g (molar mass)

A = mass number Z = atomic numberA = number of protons + neutrons

Z = number of protonsA – Z = # of neutrons

Isotopes= atoms with same number of protons but different numbers of neutronsIons= electrons varyIsomer – different arrangements if atoms within a moleculesPROTONS, Atomic # define the element!

Name Formulamethane CH4

ethane C2H6

propane C3H8

butane C4H10

pentane C5H12

hexane C6H14

heptane C7H16

octane C8H18

nonane C9H20

decane C10H22

….ene CnH2n

…yne CnH2n-2

Covalent prefixesNumber Prefix Number Prefix

1 mono- 6 hexa-

2 di- 7 hepta-

3 tri- 8 octa-

4 tetra- 9 nona-

5 penta- 10 deca-

MnO4- permanganate

S2O3-2 thiosulfate

CrO4-2 chromate

Cr2O7-2 dichromate

O2-2 peroxide

ClO- hypochlorite

ClO2- chlorite

ClO3- chlorate

ClO4- perchlorate

C2O4-2 oxalate

Fe(CN)6-4 hexacyanoferrate(II)/ ferrocyanide

Fe(CN)6-3 hexacyanoferrate(III)/ ferricyanide

BO3-3 borate

Hg2+2 mercury (I) ion

Molecular formula = the way a molecule really appears in natureEmpirical formula = smallest whole number ratio of elements

= = whole number always! Glucose has a molecular formula (MF) of C6H12O6

Glucose’s empirical formula (EF) is CH2O

Glucose’s molar mass (MM) is 180 g/mol

Glucose’s empirical mass (EM) is 30 g/mol

EF X 6 = MF; EM X 6 = MM for glucose!

Solving EF problems:1. If given percentage, drop % and use g. (assume 100 g)2. if given mass, just leave it!3. Convert g of each element to moles to get mole ratio.4. Divide by smallest to get lowest whole number ratio.5. if dividing by smallest does not give whole number ratio, then multiply all by a whole number:.2,.4,.6,.8 x 5.25, .75 x 4.33, .67 x 3

Diatomic molecules - BrNClHOFOther Molecular Elements

o Phosphorus--one form P4o Sulfur—often S8o Carbon--4 allotropes, diamond and graphite,

buckyballs, graphene, all covalent networks

Other compounds to memorize1. Hydrogen Peroxide – H2O2

2. Ammonia – NH3

REACTIONS TO MEMORIZE:1. Synthesis reactions of metal oxides with water: .

ex: Na2O + H2O 2NaOH

CaO +H2O Ca(OH)22. Synthesis reactions of nonmetal oxides with water:

ex: SO3 + H2O H2SO4

N2O5 + H2O 2HNO33. Decomposition of Binary Compounds ex: 2H2O 2H2 + O2

MgS Mg + SH2O2 H2O + O2

4. Decomposition of Metal Carbonates: ex: CaCO3 CaO + CO2

Na2CO3 Na2O + CO25. Decomposition of Metal Hydroxides:

ex: Ca(OH)2 CaO + H2O

2NaOH Na2O + H2O

6. Decomposition of Metal Chlorates: ex: 2KClO3 2KCl + 3O2

Mg(ClO3)2 MgCl2 + 3O27. Decomposition of Metal Nitrates:

ex: 2NaNO3 2NaNO2 + O2Ba(NO3)2 Ba(NO2)2 + O2

8. Decomposition of acids: certain acids can decompose to form nonmetal oxides and water.ex: H2CO3 CO2 + H2O

H2SO3 SO2 + H2O

9. Combustion: elements in a compound separate and combine with oxygen to make most reasonable cmpds.ex: ABC + O2 AO + BO + CO2

C2H5OH + 3 O2 2 CO2 + 3 H2O CS2 + 3 O2 CO2 + 2SO2

10. Single Replacement: A metal replaces a metal or a nonmetal replaces a nonmetal IF the element doing the replacing is higher on the activity series. Hydrogen replaces/is replaced by metals, though.Ex: MgCl2 + 2Na 2 NaCl + Mg

Cl2 + 2 NaI 2 NaCl + I2

2 HCl + Mg MgCl2 + H2

11. Double replacement: Only works if 1 of following is produced: 1. Precipitate 2. Gas 3. Water

If the following is made It decomposes into…NH4OH H2O + NH3

H2SO3 H2O + SO2

H2CO3 H2O + CO2

All compounds containing group 1 metals, acetates, nitrates, or ammonium are soluble.

BaSO4, PbI2, and AgCl are common insoluble substances!

There is no such thing as an oxide ion in solution! It will quickly become a hydroxide ion!

Preparing Solutions:1. Always use volumetric flask.2. Add solute or concentrated solution.3. Fill to the line with deionized water.4. Shake.

McVc = MdVdMc = molarity of concentrated solution (used to make new soln)Vc = molarity of concentrated solution (used to make new soln)Md = concentration of solution producedVd = volume of solution produced = vol. of concentrated soln +

volume of water!

Leo says Ger!!!!Losing Electrons is OxidationGaining Electrons is Reduction!EX: 2 AgNO3 + Zn Zn(NO3)2 + 2 Ag

2 Ag+ + Zn Zn+2 + 2 Ag Silver is reduced; zinc is oxidized2 moles of electrons are transferred (not 4!) Half reactions: 2 Ag+ + 2 e- 2 Ag Zn Zn+2 + 2e-

Half reactions are often easy!If not single replacement, use….

All Oranges Have Citrus Energy!1. Separate into half reactions. Leave out spectators.2. A = balance atoms (except H and O)

3. O = balance oxygens (add H2O for acid rxns; add 2 x OH- for bases)

4. H = balance hydrogens (add H+ for acid rxns; add H2O for bases)

5. C = balance charge (by adding electrons)

6. E = balance electrons (by multiplying half reactions)

Assigning Oxidation Numbers:1. In free elements, each atom has an oxidation number of 0.2. For ions consisting of a single atom, the oxidation number is

equal to the charge on the ion.3. For binary ionic compounds, assign oxidation numbers

(charges) as you always have!For covalent molecules, or polyatomic ions, assign oxidation numbers in the order shown below.

F-O-H, closest to F, farthest from F4. Fluorine is always assigned first an oxidation number of –1

when in a compound.5. Next, assign oxygen. The oxidation number of oxygen is -2 in

most compounds.Exceptions: Oxygen in peroxides is -1 Ex. H2O2 Na2O2

6. Then, assign hydrogen. The oxidation number of hydrogen is usually +1, but may be –1 when combined with a metal.

For example: H in hydrides is -1 Ex. NaH7. For compounds in which both atoms cannot have the

oxidation number which is equal to the charge the element commonly has, one closest to fluorine “wins”.

HEAT Total amount of kinetic energy of a sample

q Joules

TEMPERATURE A measure of average kinetic energy of a sample

T oC, K

SPECIFIC HEAT The amount of energy required to change the temp of 1 g of substance by 1oC

C J/goCcal/goCJ/moloC

HEAT CAPACITY

The amount of energy required to change the temp of a given substance by 1oC

none J/oC

cal/oC

1 mL H2O = 1 g H2O Enthalpy = heat

Amount of energy to change temp of 1oC of H2O = 1 cal4.184 J = 1 cal 1000 cal = 1 kcal = 1 Cal

200 g water at 70oC 100 g water at 70oCSame T

Same KEavg of particlesSame speed of particles

Same specific heat of waterMore heat energy (q) Less heat energy (q)

Heat of formation ∆Hf Energy required to form an element from its standard

stateHeat of fusion ∆Hfus Energy required to melt a

substance or RELEASED when frozen

Heat of vaporization

∆Hvap Energy required to boil a substance or RELEASED when

condensedHeat of reaction ∆Hrxn Energy required/released

during a reaction (per mole)Heat of combustion

∆Hcomb Energy required/released during a comb. reaction (per

mole)Ionization energy IE Energy required to remove 1

electron from an atomLattice Energy U Energy required to break 1

mole of ions in a crystal lattice into gaseous ions

Standard entropy ∆S Measure of disorder compared to that of a solid

crystal at 0KGibbs Free Energy of formation

∆Gf Measure of spontaneity; ability to do work on

surroundings as compared to elements in standard state

Gibbs free energy of reaction

∆Grxn Measure of spontaneity; ability to do work on

surroundings

Ways to Calculate the Enthalpy of reaction, ∆Hrxn:1. ∆Hf (products) - ∆Hf (reactants)2. bonds broken – bonds made3. sum of ∆Hrxn for reactions that add up to new reaction4. measure amount of energy gained/lost by water. Divide

by moles of 1 reactant reacted.

Person Experiment Discovered/ProposedDalton - Atomic Theory – 5 postulates, incorrect about all atoms of a given element

identical and atoms are indivisibleMendeleev Designed Periodic Table; left holes for 3 missing elements and predicted their

propertiesThomson Cathode rays All matter contains electrons (and therefore protons); determined mass:charge

of electron; Proposed plum pudding model of atom (p+/e- spread throughout)Millikan Oil drop experiment Determined exact mass and charge of electronRutherford Gold foil experiment Discovered positively charged, very dense nucleusEinstein Photoelectric Effect photonsPlanck - Calculated “size” of photon – Planck’s constantBohr Line emission spectrum of

(excited) hydrogen gasExistence of energy levels/ quantized energy states of electrons

De Broglie - De Broglie equation; wavelength of any moving objectSchrodinger - Schrodinger’s equation; calculates probability of finding electron in a given

region (orbital!) within an atom by treating electron as probability wave functionHeisenberg - Uncertainty Principle; the greater the precision in measuring a small object’s

location, the greater the uncertainty in measuring its velocity and vice versa (can’t know an electron’s location and velocity simultaneously)

Excited electrons have gained energy and jumped to a higher energy level. They possess more energy.

They fall back down to a lower energy state and must release energy in the form of 1 quantum/photon, E = hν

Only certain sized photons (lines of frequency/wavelength) are emitted so each element has its own distinct line emission spectrum. This is due to the existence of quantized energy states in atoms.

Pauli Exclusion Principle No 2 electrons in the same atom can have the same 4 quantum numbers; thus no 2 electrons can be in the same energy level/sublevel/orbital AND have same spin; 2 electrons in same orbital must have opposite spins

Aufbau Rule Electrons fill up orbitals from lowest energy to highest energy (this may not be in numerical order! See aufbau box below) 2 is higher than 1; d is higher than p, etc….

Hund’s Rule (Bus Rule!) If 2 equal energy orbitals are available, electrons each go to separate orbitals (with same spin) before pairing up 2p: ____ _____ _____

Sublevel orbitals within Shape of orbitals Picture # of electrons in an orbital

# of electrons in sublevel

s s Spherical 2 2

p px, py, pz Dumbbell shaped 2 6

d dxy, dxz, dyz, dz2, dx2 –y2 4 lobed 2 10

f 7 different f’s 8 lobed/too complex Don’t even try 2 14g 9 different g’s !!!!! AAAAH! 2 18

Quantum numbers describe where a given electron is in an atom.

1st quantum # Energy level 1, 2, 3, 4….2nd quantum # sublevel s, p, d, f3rd quantum # Orbital/orientation px, py, pz,for example

4th quantum # spin

Aufbau box: (add arrows)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f 5g

6s 6p 6d 6f 6g 6h

7s 7p 7d 7f 7g 7h 7i

8s 8p 8d…….

Example of orbital notation:

C: ____ ____ _____ _____ ______

1s 2s 2pExamples of electron configuration : Mg: (you do it!)K: 1s22s22p63s23p64s1

Mn: 1s22s22p63s23p64s23d5

Ag: 1s22s22p63s23p64s23d104p65s14d10

Pb: (you do it!)

# unpaired electrons = #electrons alone in orbitalC has 2 unpaired electronsK has 1, Ag has 1, Pb has 2, Mn has 5, Mg has 0

Unpaired electrons make element paramagnetic.Drawing Lewis Structures for COVALENT compounds:

1. The atom with lowest EN is the central atom.2. Count the total number of valence electrons.

MAGIC NUMBER! (adjust for ions)3. Place one sigma bond (a pair of electrons)

between each pair of bonded atoms.4. Subtract the total number of valence electrons

used for bonds from the MAGIC NUMBER. 5. Place lone pairs about each terminal atom

except for hydrogen atoms. Subtract the number of lone pairs from the sum.

6. If e- are still left at this point, assign them to the central atom. If the central atom is from the third or a higher period, it can accommodate more than four pairs of electrons.

7. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to pi bond pairs.

Which one has higher Coulombic attraction? Why?

IONIC COVALENTMetal + nonmetal

Metal +polyatomic ionPoly ion + nonmetalPoly ion + poly ion

Nonmetal + nonmetalWithin poly ion

Strongest type of bond Weaker bondElectrostatic attraction

between oppositely charged ions

Electrostatic attraction between nuclei and

shared electronsElectrons transferred,

creating ionsElectrons shared – NO

IONS!Forms crystal lattice

And thuscrystals

A few form covalent networks (SiO2, , most

form individual moleculesFormula units molecules

Higher MP, Higher BP(must break bond to

change phase)

Lower MP, lower BP(only break IMFS NOT

bonds to change phase)Lattice Energy Bond energy

Higher charge, smaller size= stronger bond

Higher bond order, smaller atom = stronger bond

Forms polar and nonpolar bonds

Large EN difference Medium (polar) or Small (nonpolar) EN difference

If soluble, forms electrolytes when

dissolved

Nonelectrolytes, may dissolve but won’t ionize!

Electrolytic when molten, not solid

Nonelectrolyte

High % ionic character Low % ionic character

Show resonance!

Total number of

sigma bonds and lone pairs on central

atom

Shape Atoms bonded to

central atom(# of

sigma bonds)

Lone pairs

on central atom

Molecular Shape Picture

And bond angles

Hybridization

Symmetrical in 3d space?

Example of a Lewis structure that is this shape. (Yes, draw the Lewis structure!)

2 Linear 2 0 sp yes

3 Triangular Planar

3 0 sp2 yes

3 Bent 2 1 sp2 no

4 Tetrahedral 4 0 sp3 yes

4 Triangular pyramidal

3 1 sp3 no

4 Bent 2 2 sp3 no

5 Trigonal bipyramidal

5 0 dsp3

orsp3d

yes

5 Unsymmetrical tetrahedron

(see saw)

4 1 dsp3

sp3dno

5 T-shaped 3 2 dsp3

sp3dno

5 Linear 2 3 dsp3

sp3dyes

6 Octahedral 6 0 d2sp3

sp3d2yes

6 Square pyramidal

5 1 d2sp3

sp3d2no

6 Square planar 4 2 d2sp3

sp3d2yes

TYPE USES TO DETERMINE/Produce HOWMass Spectroscopy Magnets Distinguishes isotopes (separates

by mass), can be used to determine % abundance of each isotope

The magnet bends heavier isotopes more than lighter ones

Photoelectron Spectroscopy

Beam of Light on Metal

Binding Energy of electrons in different sublevels within a given atom

“Light” beam (could be other form of ER) gives energy to electrons and removes them from the atom- this energy is

measured…takes more energy to remove innermost electrons than outermost electrons

Atomic absorption spectroscopy or Atomic Emission Spectroscopy

Electrical or Heat Energy added to excite electrons

Line Emission spectrum (or line absorption spectrum) which identifies the element

Energy excites electrons and they absorb a specific wavelength, then when they fall back down, they release a specific wavelength

NMR Spectroscopy (nuclear magnetic resonance)

Radio waves (LOW Energy right?) added to “excite” nucleus

Identity of a substance (types of atoms present)

Nucleus vibrational states altered by radio waves, gives “signature” of element

Infrared or Visible or UV spectroscopy (example: Using spectroscopes)

IR or Visible light or UV light to excite electrons in bonds

Identity of a substance/molecule, (types of bonds present) OR concentration of substance (when we used spectrophotometer)

Electrons in bonds also have quantized energies and these energies can be used to identify a molecule

Mass Spectrometry Data: …each bar represents the mass and abundance of a given isotope. What element? _______

PES data: (photoelectron spectroscopy)…each bar represents the abundance AND energy to remove each of the electrons from an atom (binding energy!). The outermost electrons take the least amount of energy to remove.

When bond is formed, 2 atoms rest at a distance of minimum potential energy. This gives the typical bond length of the molecule.

1s

2s2p3s 2s

----------------------------------------------------------∆E = q + w

∆E = internal energyq = heat added to the systemw = work done on the systemSo, for example if heat is added to the system AND work is done on the system, then the internal energy of the system increases and has a positive sign.

If a reaction is exothermic and it expands in volume, work is done on the surroundings and so work is negative, and heat is negative, and the internal energy decreases and change in energy has a negative sign.---------------------------------------------------------

During phase change, added or removed heat doesn’t change kinetic energy of particles and thus, does not change the temp. Energy is used to break IMFs or released when formed.

Boiling point – temp at which v.p equals atmospheric pressure.

Normal b.p. –temp at which v.p equals 1 atm.

Top of hump = activation complex/transition state/high-energy intermediate

Taller the “hump”, slower the reaction

Activation energy is used to break bonds, create an “effective” collision

Effective collision:1. enough energy to react/break bonds2. correct orientation

A 2 step mechanism might result in double hump.

Slow step only determines rate of reaction. (slow step has higher hump)

5 factors that affect rate of reaction: 1. nature of reactants2. concentration: increases # of collisions3. phase/ surface area: increases # of collisions4. temperature: increases # of collisions AND # of effective collisions5. catalyst: Increases # of effective collisions by lowering activation energy

Consider this reaction: A+ B D+ E Rate = k [A]x[B]y[C]z NOT to be confused with

equilibrium expression: K = 1. Exponents (in rate law) are NOT determined by

balanced equation.2. Exponents can match slow step of mechanism.3. Exponents can only be determined

experimentally.4. k depends on T.5. A catalyst can be included (C perhaps)6. Products are (almost) never included – no

denominator (in most cases)

How is rate law determined experimentally?1. Method of Initial Rates: Start with different

concentrations of reactants and measure initial rate.

a. If tripling A causes rate to increase 9-fold, exponent of A is 2.

2. Measure the concentration over time. Make a graph to see what order.

1 mol gas at STP = 22.4 L

2 ways a gas may not be ideal:1. Large size2. Polar

An unideal gas:1. Tends to have inelastic collisions, and/or2. Have a size that DOES matter 3. AND doesn’t obey the gas laws perfectly.

Any gas is more real (unideal) at 1. High pressure (size of particles does matter)2. Low temperature (particles more likely to have

inelastic collisions)

Any gas is more ideal at 1. Low pressure2. High temperature

2 balloons each filled with different gases at the same volume in the same room.

CO2 HeSame V, T, and P

Same KE.

Same number of gaseous particles

Higher mass, higher density Lower mass, lower density

Lower speed of particles Higher speed of particles

Zero Order 1st order 2nd orderRate law Rate = k Rate = k[A] Rate = k[A]2

Rate depends on [A] NO! Yes YesIntegrated Rate law [A] = -kt + [A]o ln [A] = -kt + ln [A]o 1 = kt + 1

[A] [A]o

Plot needed to give a straight line

[A] vs. t ln [A] vs. t 1/[A] vs. t

Relationship of rate constant to slope of

straight line

Slope = -k Slope = -k Slope = k

Half life t1/2 = [A]o/2k t1/2 = 0.693/k t1/2 = 1/k[A]o

Half life depends on [A] yes NO! yes

Activation energy with catalyst

What are IMFs?1. NOT bonds2. Attractions between neighboring species3. MUCH Weaker than bonds usually

The 5 types of intermolecular forces (IMFS) listed from strongest to weakest are:

1. Ion-ion (NOT really an IMF)Example: Salt bonds

2. Ion-dipoleExample: Salt dissolving in water

3. Dipole-dipole (Strongest type is hydrogen bonding)

Example: Water condensing4. Dipole-induced dipole

Example: CO2 dissolving in water5. Induced dipole-induced dipole

AKA London dispersion forcesAKA Van der Waals forces

Example: freezing nitrogen

The stronger the IMFS, the lower v.p. at given T.

The stronger the IMFs, the higher the normal b.p.

Types of crystalline solids:

Ionic –formed from ionic bonds, very strong, brittle, high MP, think NaCl

Molecular – formed from the regular arrangement of IMFS between covalent molecules, low MP, not strong, think snowflakes

Network – covalent bonds throughout, high MP, hard, nonconducting (usually), think SiO2 (sand/glass), C (diamond)

Graphite is a unique network solid made only of carbon atoms. They are bonded together in sheets. The sheets are attracted to each other by weak IMFS. (rubs off easily) Within the sheets, there is resonance throughout, allowing graphite (a nonmetal) to conduct electricity!

Ways to increase solubility of a

Solid gasIncrease temp Decrease temp

Increase surface area Increase pressure aboveAgitation, stirring Don’t agitate

ΔHsoln (heat of solution): SUM OF- Energy required to break the water-water IMFs- Energy required to break IMFS in solute- Energy released when IMFS form between solute

and solvent.

Hydrogen bond: the ATTRACTION between a H of one molecule (that is bonded to a very EN element such as F, O, N) and a very EN element (O, F, N) of another molecule. NOT A BOND!!!!!!!!!

Ways to measure the concentration of a solution:

Molarity (M) = mol solute/L solutionMolality (m) = mol solute/kg solventMole fraction (X) = mol solute/ mol solution% by mass = g solute/g solution% by volume = mL solute/mL solutionm/v % = g solute/mL solution

Colligative property: any property of a solution that changes based SOLELY on the NUMBER of solute particles dissolved. (not mass or size or IMF formed!)

-freezing point depression-Vapor pressure lowering-Boiling point elevation (as a result of v.p. lowering!)

WHY? Solute particles interfere with the process of evaporating, boiling, and freezing.

Incr

easin

g m

.p.,

b.p

, Hfu

s, H v

ap, c

ritica

l poi

nt, s

peci

fic

heat

Increasing vapor pressure, volatility, likelihood of being gas at room

temp

Alloys are metal solutions. (homogeneous mixtures)

Substitutional alloy: formed by 2 metals of similar radii; mp or ability to resist corrosion or conductivity might be affected. Hardness is NOT.

Interstitial alloy: formed by 2 metals of different radii; adding in smaller metals tend to make the alloy harder.

For both alloys, adding metals with more valence electrons increases conductivity.

Strong acid: Weak acid:

Strong base: Weak base:

Soluble ionic compound: Insoluble ionic compound:

Soluble covalent compound: Insoluble covalent compound:

DO NOT MIX CONCEPTS! Periodic trends do NOT explain boiling point for example.

WHY …? ONLY Reasons you will use

Does X have a larger radius than Y? a higher ionization energy than Y? a higher electronegativity than Y? a higher reactivity than Y?

Tell how many protons and electrons each species hasThen, useproton-electron attraction (Coulomb’s LAW)OR electron-electron repulsion (Coulomb’s LAW)OR shielding (electrons don’t feel protons pulling on them due to intervening energy levels)

Does Y make a given shape? Have a given bond angle? VSEPR THEORY!!! Bonding pairs of electrons experience

repulsion with each other so they spread out in 3d space as far as possible; Lone pair electrons induce greater repulsion and cause angles between bonding electrons to be smaller than expected.

Is Y polar? Electronegativity difference (polar or nonpolar bonds?)AND Symmetrical or Unsymmetrical shape

Does Y dissolve in Z? Have a high freezing pt, low volatility, low vp, high critical pt, high melting pt, high Hfus, high Hvap?

Identify intermolecular force involvedCompare strengths of intermolecular force

Does SOLUTION A Have a high b.p., high conductivity, low f.p., low v.p?

Compare [ ] of solutions

Ksp problem types: always AB A + + B - 1. Calculate solubility using Ksp: Do RICE solve for x

2. Calculate Ksp using solubility: Do RICE, solve for K3. Calculate solubility within a solution of common ion: Do RICE, solve for X (should not have all 0’s on I row)4. Calculate if precipitate forms when mix 2 solutions:

- MV = MV to determine new concentration- Plug into equil expression to get Q- Compare to K

5. Calculate how much of a given substance is required to form precipitate, (or compare which precip forms 1st)

- Solve equil expression for missing value!

Le Chatelier’s principle:If a reaction is stressed, the reaction responds in such a way as to remove that stress.

Equilibrium can be disturbed by:1. Adding products (shifts left)2. Adding reactants (shifts right)3. Removing reactants (shifts left)4. Removing products (shifts right)5. Adding volume or decreasing partial pressure (shifts

towards side with fewer gaseous or aqueous particles)6. Increasing temperature (shifts to get to new K)

a. Shifts right if endothermicb. Shifts left if exothermic

Equilibrium canNOT be disturbed by:1. Adding inert gas (no effect)2. Adding solid (no effect)

The ONLY way to change K is to change the temperatureK >1 products preferred at equilibrium

Buffer: combination of (weak) conj. acid/base pair that resists change in pH

A good buffer:1. desired pH = pKa of acid2. [conj. acid] = [conj. base]3. [conj.acid] and [conj. base] high (increases buffering capacity)

How to approach ANY titration calculation:1. Molarity x volume = moles for both substances.

Subtract to see what remains. Does a conjugate acid or base get produced? (it will in any weak titration.) Determine its moles.

2. Divide by volume to get molarity for substance remaining.

3. THINK….how do I get pH of this/these substances?

a. –log [strong]b. If weak, RICEc. If weak acid and weak base are present,

do H-Hasselbach.

At equivalence point of ANY titration: MaVa = MbVb, moles acid – moles base = 0Midpoint = half moles neutralized (AND [weak acid/base] = [conj. acid/base]); pH = pKa

TitrationsStrong acid/strong base

Strong acid/weak base

Strong base/ Strong acid

Strong base/weak acid

Starting point high not as high low not as lowMidpoint Not relevant pH= pKa Not relevant pH = pKaEquivalence point Exactly 7 Below 7

(conjugate acid present)

Exactly 7 Above 7(conjugate base present)

Vertical region Very long Not as long Very long Not as long Always pick an indicator whose pKa = pH at the equivalence point of the titration.

Various Ways to Describe Acid StrengthProperty Strong Acid Weak Acid

Ka Value Too large to measure << 1

Position of the dissociation equilibrium

right left

Equilibrium concentration of H+ compared to the original

concentration of HA

[H+ ] = [HA]o [H+ ] << [HA]o

Equilibrium concentration of HA compared to the original

concentration of HA

[HA] ~ 0 M[HA] << [HA]o

[HA] ~ [HA]o

Percent Dissociation 100% < 5%

Strength of conjugate base Conjugate base is nonexistent Conjugate base is also weak, but inversely proportional in strength to

the acid (Kb = Kw/Ka).How to represent in a net

ionic equationH+ + A- HA

Sign

Of Δ H0

System Entropy Change

Sign

ofΔ S0

Sign of Δ G0

Sign of E0

(if a redox)K Spontaneous?

(Thermodynamically favorable?)

1. Exothermic - Increasing + - + >1 always

2. Exothermic - Decreasing - + or - - or + Don’t know Low T

3. Endothermic + Increasing + + or - - or + Don’t know High T

4. Endothermic + decreasing - + - < 1 never

Galvanic Cell/ Voltaic Cell/ Electrochemical Cell

Spontaneous

Ex: battery!

Chemical energy electrical energy E = + ΔG = - anode= oxidation

Cathode = reduction

Electrolytic Cell NonspontaneousEx: electrolysis water

Electrical energy chemical energy E = - ΔG = + Anode= oxidation

Cathode= reduction

As reaction progresses, K doesn’t change BUT

When rxn is shifting right

When rxn is shifting left When a reaction is at equilibrium

Q approaches K Q < K Q > K Q = K

G approaches 0 G = - G = + G = 0

E approaches 0 E = + E = - E = 0

H+1/-1 Highest IE

Most reactive

nonmetal, highest

EN

Zn+2

Ag+1

Cd+2

Sn Sb

57-70

*

PbRoman

numeralMost

reactive metal,

lowest IE, lowest EN

89-102

**

*

**

Variable positive charge = roman numeral

roman numeral

+3

(+2) +4

(+3) +5 +6

d block (1 behind!)

Transition metals

p block

Alka

li m

etal

s+1

+2 -4 -3 -2 -1

No ions formed usually

+7

Variable positive charge = roman numeral

Rare earth metals

Lanthanide series, 4f

Actinide series, 5f

Increasing nuclear charge explains everything: increasing ability to attract electrons, decreasing AR, increasing EN, increasing IE, which both explain

decreasing reactivity across metals, increasing reactivity among nonmetals

Incr

easin

g #

of e

nerg

y le

vels

(incr

ease

d sh

ield

ing)

exp

lain

s eve

ryth

ing:

Decr

easin

g ab

ility

to a

ttra

ct e

lect

rons

, In

crea

sing

AR, D

ecre

asin

g IE

, Dec

reas

ing

EN Whi

ch e

xpla

in in

crea

sing

reac

tivity

am

ong

met

als,

dec

reas

ing

reac

tivity

am

ong

nonm

etal

s

No

mea

sura

ble

EN (t

oo lo

w),

high

IE, I

NER

T!

S bloc

kAl

kali

ne e

arth

met

als

57-70

*89-102

**

*

**

Page 16: Look on the internet for an interesting molecule that has many atoms. Draw its Lewis structure below and 1. Give its molar mass, 2. Give its mass per molecule. Then pick 2 central atoms that have different hybridizations and label their 3. Hybridization, 4. Bond angles, 5. Shape about that atom

A good website to try might be http://www.chm.bris.ac.uk/motm/motm.htm (Molecule of the Month).

Which atom?

1. 2. 3. 4. 5.

Which atom?

3. 4. 5.