Chem 105 final review

Chem 105 Final ReviewDR. SHIRTS-WINTER 2013

JOANNA WILLIAMS

Disclaimer

All the problems on here come from your textbook! If you need more, I suggest working through micro-exams and practice sheets/parallel example problems, seeing your Learning Community Mentor to run through previous exams, and looking in the book for more problems to work through over specific things your struggling with.

Measurements

Accurate: numbers close to the actual value

Precise: numbers close to each other

Significant Figures:

All non-zeros are sig. figs.

Zeros between two non-zeros are sig. figs.

Zeros left of first non-zero are NOT sig. figs.

If #>or=1, all zeros right of decimal are sig. figs.

If #<1, all zeros at end of # and between non-zeros are sig. figs.

Trailing zeros may or may not be sig. figs. (That’s why we use scientific notation)

Nomenclature

Metal + Nonmetal = Ionic compound

Charges designate formula, name the elements and add –ide to the end

Nonmetal + Nonmetal = Covalent molecule

Use prefixes and add –ide to the end

Polyatomic Ions

Organic Functional Groups

Organic Prefixes

Acids

If you “–ate” too much you feel “–ic”ky

“-ite”s like Nephites and Lamanites are people like “-ous”

Hypo-ous, ous, ic, per-ic increasing O

Hydro-ic

Dimensional Analysis

Use the Mole to Mole Ratio from stoichiometric coefficients in balanced chemical equation

Find limiting reactant

Practice Problems

3.51) Determine the empirical and molecular formulas of each of the following substances:

Styrene, a compound used to make Styrofoam cups and insulation, contains 92.3% C and 7.7% H by mass and has a molar mass of 104 g/mol

Caffeine, a stimulant found in coffee, contains 49.5% C, 5.15% H, 28.9% N, and 16.5% O by mass and has a molar mass of 195 g/mol

Monosodium glutamate (MSG), a flavor enhancer in certain foods, contains 35.51% C, 4.77% H, 37.85% O, 8,29% N, and 13.60% Na, and has a molar mass of 169 g/mol

Practice Problems

3.69) A piece of aluminum foil 1.00 cm square and 0.550 mm thick is allowed to react with bromine to form aluminum bromide.

How many moles of aluminum were used? (density of aluminum=2.699 g/mL)

How many grams of aluminum bromide form, assuming the aluminum reacts completely?

Practice Problems

3.76) Aluminum hydroxide reacts with sulfuric acid as follows:

2Al(OH)3(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 6H2O(l)

Which is the limiting reactant when 0.500 mol Al(OH)3 and 0.500 mol H2SO4 are allowed to react? How many moles of Al2(SO4)3 can form under these conditions? How many moles of the excess reactant remain after the completion of the reaction?

Reactions

Know the solubility rules and the exceptions for precipitation reactions

Oxidation-Reduction (RedOx)

123FHO7654

LEO goes GER / OIL RIG

Oxidizing agents/Reducing agents

Reactions

Acid/Base/Neutralization

Titrations and Dilutions: M1V1=M2V2 (molarity= mol/L)(volume=L)

Net Ionic Equations

Strong Acids (ionize completely)

H2SO4 , HNO3, HCl, HBr, HI, HClO4

Strong Bases (dissociate completely)

Group 1-OH, Group 2-OH from Ca down

pH

pH=-log[H+]

pOH=-log[OH-]

pH+pOH=14

[H+]+[OH-]=1x1014

pH<7 acidic

pH=7 neutral

ph>7 basic

Practice Problems

4.83) Some sulfuric acid is spilled on a lab bench. You can neutralize the acid by sprinkling sodium bicarbonate on it and then mopping up the resultant solution. The sodium bicarbonate reacts with sulfuric acid as follows:

2NaHCO3(s) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) + 2 CO2(g)

Sodium bicarbonate is added until the fizzing due to the formation of CO2(g) stops. If 27 mL of 6.0 M H2SO4 was spilled, what is the minimum mass of NaHCO3 that must be added to the spill to neutralize the acid?

Practice Problems

4.40) Write the balanced molecular and net ionic equations for each of the following neutralization reactions:

Aqueous acetic acid in neutralized by aqueous barium hydroxide

Solid chromium(III) hydroxide reacts with nitrous acid

Aqueous nitric acid and aqueous ammonia react

Practice Problems

4.51) Which element is oxidized and which is reduced in the following reactions?

N2(g) + 3H2(g) 2NH3(g)

3Fe(NO3)2(aq) + 2Al(s) 3Fe(s) + 2 Al(NO3)3(aq)

Cl2(aq) + 2NaI(aq) I2(aq) + 2NaCl(aq)

PbS(s) + 4H2O2(aq) PbSO4(s) + 4H2O(l)

Thermochemistry

Ek = ½(mv2)

1 Cal = 1000 cal

1 cal = 4.184 J (specific heat of water)

∆E=q+w

H=E+PV

Bond enthalpies: reactants – products aka bonds broken – bonds formed

∆Hfo: products – reactants (diatomics in natural state = 0)

q=mCs∆t

Hess’s Law

Practice Problems

5.43) Consider the following reaction:

2Mg(s) + O2(g) 2MgO(s) ∆H = -1204 kJ

Is this reaction exothermic or endothermic?

Calculate the amount of heat transferred when 3.55g of Mg(s) reacts at constant pressure

How many grams of MgO are produced during an enthalpy change of -234 kJ?

How many kilojoules of heat are absorbed when 40.3g of MgO(s) is decomposed into Mg(s) and O2(g) at constant pressure?

Practice Problems

5.56) When a 4.25g sample of solid ammonium nitrate dissolves in 60.0g of water in a coffee-cup calorimeter, the temperature drops from 22.0 C to 16.9 C. Calculate ∆H (in kJ/mol NH4NO3) for the solution process

NH4NO3(s) NH4+(aq) + NO3-(aq)

Assume that the specific heat of the solution is the same as that of pure water.

Is this process endothermic or exothermic?

Practice Problems

5.65) From the enthalpies of reaction

H2(g) + F2(g) 2HF(g) ∆H = -537 kJ

C(s) + 2F2(g) CF4(g) ∆H = -680 kJ

2C(s) + 2H2(g) C2H4(g) ∆H = +52.3 kJ

Calculate ∆H for the reaction of ethylene with F2:

C2H4(g) + 6 F2(g) 2CF4(g) + 4HF(g)

Electrochemistry

E=hv=hc/λ

1/λ=R(1/n12-1/n2

2)

E=-Rhc(1/n2)

λ=h/mv

Bohr’s Model

Practice Problem

6.37) Calculate the energy of an electron in the hydrogen atom when n=2 and when n=6. Calculate the wavelength of the radiation released when an electron moves from n=6 to n=2.

Is this line in the visible region of the electromagnetic spectrum? If so, what color is it?

Orbitals & Nodes

s=spherical

Radial nodes starting at 2s

p=peanut

1 planar node and radial nodes starting at 3p

d=dlover leaf?

2 planar nodes and radial nodes starting at 4d

f

Quantum Numbers

Pauli Exclusion Principle

n=shell (1, 2, 3….)

l=subshell (n-1 to 0) (0=s, 1=p, 2=d, 3=f…)

ml=orientation (-l to l)

ms=spin (-1/2 or +1/2)

Practice Problem

6.56) Which orbital goes with the following quantum numbers? Which are not allowed?

2, 1, -1

1, 0, 0

3, -3, 2

3, 2, -2

2, 0, -1

0, 0, 0

4, 2, 1

5, 3, 0

Electron Configuration

Expanded

Condensed using Noble Gas configuration

Cu and Cr exceptions

Why? Hund’s Rule

Practice Problem

6.61) For a given value of the principal quantum number, n, how do the energies of the s, p, d, and f subshells vary for

Hydrogen?

A many-electron atom?

Periodic Trends

Electronegativity

Size

Ionization Energy

Electron Affinity

Effective Nuclear Charge

Family Names

Cation and Anion Size

Lewis Dot Structures

Count total valence electrons

Least electronegative atom in the middle

Fill octet, create multiple bonds if too many electrons

Resonance structures: none actually what the molecule looks like, it’s a hybrid of all of them

Formal charges

Bond strengths

Molecular Orbitals

Bond Order

MO Diagrams

Paramagnetic vs. Diamagnetic

VSPER

Molecular shapes and Angles

Gases

PV=nRT

22.4 L = 1 mole gas @ STP

Ideal gas characteristics Small, high temp, low pressure

Partial pressures Pa=XaPt

Pt=P1+P2+P3…

Effusion and Diffusion Rates Urms=√(3RT/M)

r1/r2= √(M2/M1)

J=kgm2/s2

Practice Problems

10.54) Calculate the density of sulfur hexfluoride gas at 707 torr and 21 C

Calculate the molar mass of a vapor that has a density of 7.135 g/L at 12 C and 743 torr

Practice Problem

10.69) A piece of dry ice (solid carbon dioxide) with a mass of 5.50 g is placed in a 10.0 L vessel that already contains air at 705 torr and 24 C. After the carbon dioxide has totally vaporized, what is the partial pressure of carbon dioxide and the total pressure in the container at 24 C?

Intermolecular Forces

London Dispersion

Induced Dipole (Polarizability)

Dipole-Dipole

Hydrogen Bonding

Ion-Dipole

Liquids

Intermolecular force effect on

Viscosity

Surface Tension

Phase Changes

Phase Diagrams

Specific heats to heat the substance to melting or boiling point

Heats of vaporization or fusion to melt or evaporate substance

Practice Problems

11.43) For many years drinking water has been cooled in hot climates by evaporating it from the surfaces of canvas bags or porous clay pots. How many grams of water can be cooled from 35 C to 20 C by the evaporation of 60 g of water?

(The heat of vaporization of water in this temperature range is 2.4 kJ/g. The specific heat of water is 4.18 J/gK.)

Colligative Properties

Depends on number of solute particles present, not identity

Don’t forget ions dissociate!

Vapor Pressure ↓

Pa=XaP°

Boiling Point ↑

Freezing Point ↓

Osmotic Pressure ↑

π=iMRT

Solids

Unit Cells

Simple/Primitive Cubic

1 atom

Face-Centered Cubic

4 atoms

Body-Centered Cubic

2 atoms

Concentrations

[ ]=M=moles solute/L solution=molarity

m=moles solute/kg solvent=molality

X=moles substance/total moles=mole fraction

ppm=(mass substance/total mass)x106

mass %=(mass substance/total mass)x100

Equilibrium

kc=[products]/[reactants]

kp=Pproducts/Preactants

aA + bB cC + dD

kc=[C]c[D]d/[A]a[B]b

Le Chatelier’s Principle

Noble gases being pumped in to increase the pressure have no effect (don’t change individual partial pressures)

Catalysts have no effect

Solids and liquids have no effect