Chemistry: The Study Guide

7/27/2019 Chemistry: The Study Guide

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Topics by Chapter for Semester I F inal Exam: Uni ts A, B, C and D (Ch 1-11, 13 and 25)

Ch 1 &2: Measurement, Mixtures, Separation (emphasis on particulate diagrams) and Naming

Ch 3: Stoichiometry, Limiting Reactants, Percent Composition, Percent Yield, Mole Relationships, Empirical and

Molecular Formula

Ch 4: Molarity, Strong Acid- Strong Base Titrations, Types of Reactions, Net Ionic Equations

Ch 5: Thermochemistry, Specific Heat, Hesss Law, Coffee-cup Calorimetry, Enthalpies of Formation andCombustion

Ch 6: Wave Nature of Light, Wave length and frequency relationship, Bohrs Model of Hydrogen atom, Calculating

Energy of an Electron, Electron Configurations, Paramagnetism and Diamagnetism,

Ch 7: Periodic Trends- Atomic Radius, Ionization Energy, Electron Affinity, Metallic Character, Boiling/ Melting

Points and Electronegativity, Coulombs Law, Group Trends for common groups

Ch 8: Chemical Bonding- Ionic, Molecular, Metallic and Covalent Network Bonding, Lattice Energy, Bond Energy,

MO theory

Ch 9: Lewis Structures, Bond Polarity and dipole moments, Molecular Geometries, VSEPR model, Hybridization

Ch 10: Gases, Ideal Gas Law, Maxwells Distribution, Gas stoichiometry, Collecting gas over water, pa rtial

pressure, KMT, Real Gases

Ch 11: Intermolecular Forces- H bond, Dipole-dipole attraction, LDFs , Vapor Pressure, Properties of liquids-

viscosity and surface tension

Ch 12: Not included

Ch 13: Properties of solutions: Molarity and Solubility Curves

Ch 25: Organic Chemistry: Naming Hydrocarbons, alkanes, alkenes, alkynes (nomenclature), Isomerism- structural

and geometric, Addition and Substitution Reactions

AP Chemistry: Review for I Semester Final Exam 2013


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Unit: 1Chapter: 1

Elements Substances that cannot be decomposed further

Coumpounds Substances composed of two or more elements

Mixtures Combinations of two or more substances

Solid-SolidHand Seperation

Solid-Liquid

Solid-SolidFiltration

Shake it, let it sit

Used to seperate immiscible Liquids

Parts seperate out, removed one by one

Seperating Funnel

Heavier Particles on bottom, lighter particles on top

Solid liquidCentrifugation

Differences in boiling points to seperate homogenous mixture

Liquid-LiquidDistillation

Seperate Homogenous Mixtures on basis of solubility

Stationary, mobile phase, where mobile phase acts as a solvent

Extend of seperation depends on solubility of mobile phase.

Chromatography

Separation of Mixtures


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ExtensiveProperties

Depends on the amount ofmatter present

Examples: mass, heat, color,volume

IntensiveProperties

Does not depend on the amountof matter present

Examples: melting point, boilingpoint, density, conductivity


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- Laws

-

A: Mass number (protons + neutrons)

Z: Atomic number (number of protons, number of electrons)

- 1 AMU is 1/12 the mass of one C-12 atom

A pure compound always contains the same elements in the sameproportions by mass.

Law of DefiniteProportions

Mass is neither created nor destroyed during ordinary chemical orphysical reactions.

Law of Conservationof Mass

When elements combine to form more than a single compound, theratios of the masses of the combining elements can be expressed by aratio of small whole numbers.

Law of MultipleProportions

EX: Prove that SO2and SO3obey the law of multiple proportions":SO232:32 (1:1) ratioSO332:48 (1:1.5) ratioThe ratio of oxygen is 1:1.5, or 2:3, which is a whole number ratio.

ZEA

- Calculates theaverage atomic massof all the isotopes ofa particular element

- Tells us roughly theweight compared toother atoms.

- Weighted averageof the atomic massesof naturally occuringisotopes of anelementR

elativeAomic

Mass

AverageAtomicMass


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- Significant Figures

o All digits known with certainty, plus one digit, which is somewhat uncertain:

o

31.23: (31.2 is known for certain. The .03 is approximated)o IF DECIMAL POINT IS PRESENT (PACIFIC)

Count digits from left side, starting with the first nonzero digit.

EX: 4.5600 5 sigfigs

o IF NO DECIMAL POINT IS ABSENT (ATLANTIC)

Count digits from right side, starting from first nonzero digit.

EX: 1200 2 sigfigs


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o Operations

Addition, Subtraction:

Answer rounded to the smaller number of digits past the decimal point

from the two numbers

EX: 56.31g14.1g = 42.2

Multiplication, Division

Answer must have the same number of sigfigs as the least certain

number.

EX: 2.4/15.82 = 38


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Chapter 2:

Democritus, 400 BC

Everything is made up of a few simple particles called atomos(uncuttable)

envisioned atomos as small, solid particles of many different sizesand shapes.

Ideas rejected by Aristotle, who supported earth, wind, water, andfireapproach .

Dalton, 1808SEE DALTONS ATOMIC THEORYin vocab

Successfully explained the three laws below.

JJ Thomson(1897)

Discovery of electrons

Cathode ray tube experiment

Cations (+) and Anions (-) (ions)

Also deduced that Atoms must have positively charged particle

Charge to mass ratio (e/m)

Propsed Plum Pudding model, with electrons embedded in apositively charged spherical cloud .

Ernest Rutherford(1911)

Gold foil experiment, bombarding thin gold foil with alpha particles

Most of space in atom was empty

In center of atom, there was a small, positively charged nucleus.

Other Scientists

Millikan: Charge of electron

Goldstein: Discoverd the Proton, through canal waves.

Chadwick: Neutron


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- Common monatomic ions1+ 2+ 3+ 4+

Sodium Na+ Magnesium Li

2+ Aluminum Al

3+ Lead(IV) Pb

4+

Potassium K+ Calcium Mg

2+ Chromium(III) Cr

3+ Vanadium(IV) V

4+

Rubidium Rb+

Strontium Sr2+

Iron(III) Fe3+

Tin(IV) Sn4+

Cesium Cs

+ Barium Ba

2+ Lead(III) Pb

3+

Copper (I) Cu+

Cadmium Cd2+

Vanadium(III) V3+

Silver Ag+

Chromium(II) Cr2+

Lithium Li+

Cobalt(II) Co2+

Copper(II) Cu2+

Iron(II) Fe2+

Lead(II) Pb2+

Manganese(II) Mn2+

Mercury(II) Hg2+

Nickel(II) Ni2+Tin(II) Sn

2+

Vanadium(II) V2+

Zinc Zn2+

- Common Polyatomic ions:1- 2+ 3+

Acetate CH3COO-

Carbonate CO22-

Phosphate PO43-

Bromate BrO3-

Chromate CrO42-

Arsenate AsO43-

Chlorate CLO3-

Dichromate Cr2O72-

Chlorite CLO2-

Hydrogen

Phosphate

HPO42-

Cyanide CN-

Oxalate C2O42-

Dihydrogen

Phosphate

H2PO4- Peroxide O2

2-

Hydrogen

Carbonate

HCO3-

Sulfate SO42-

Hydrogen

Sulfate

HSO4- Sulfite SO3

2-

Hydroxide OH-

Hypochlorite CLO- 1+ 2+

Nitrate NO3- Ammonium NH4

-Dimercury Hg2

2+

Nitrite NO2-

Perchlorate CLO4-

Permanganate MnO4-


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BinaryIonicCompound

Compounds composed of two different ions,nonmetal and metalNaming: Name of cation, plus name of anion. Ifthe ion has multiple charges, name it with thecharge in Roman Numerals.

(Anions named withstem of name + ideadded to the end.)

ex: NACL:

Sodium Chloride

BinaryMolecularCompound

Compounds composed of two different nonmetals

Naming: Name of two elements with prefix

attached.

(For the first element,only add a prefix if > 1)

ex: P4O10:

TetraPhosphorousDecoxide

TernaryIonicCompound

Compounds composed of two different ions, with atleast one being a polyatomic ion.Naming: Name of first ion, plus name ofsecond ion.

Suffix of monotomicanions are changed to -ide

ex: (NH4)3PO4:

Ammonium Phosphate

Binary Acid Compounds with two nonmetallic ionic elements, thefirst being Hydrogen

Naming: Hydro + stem or root of nonmetal + ic+

___Acid(There can be multipleHydrogens) ex: HCL

Hydrochloric Acid

TernaryAcid

Compounds with two nonmetallic ionic compounds,the first being Hydrogen.

Naming: If polyatomic ion has suffix of -ate or -ide, + ic + acid

If the ion has suffixite, + ous+acid

ex: HNO2

Nitrous Acid

Hydrate Compound Hydrated with Water

Naming: Name the compound,then the Greek prefix, andthen add -hydrate.

ex: CuSO45 H2O

copper (II) sulfate pentahydrate


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PxQyYES NO

Does it start with Does it start with

H?

Does it start with

a metal?

YES NO

NY

N

Binary

Acid

Binary

IonicBinary

Molecular

TernaryIonic

TernaryAcid


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Chapter 3: Formula weight (FW): same as MW, except for ionic substances in which no molecule

exists. Then, FW is the simplest integer ratio of moles of each element (ions) present.

Ex: NaCl is a 3D array of ions

FW = 22.99 + 35.453 = 58.44 amu

Stoichiometry, Limiting Reactants, Percent Composition, Percent Yield, Mole Relationships, Empirical and

Molecular Formula problems

Chapter 4:

Molarity (M)= moles solute = mol

volume of solution LSince volume of a solution changes with temperature, M of a solution changes withtemp. too. M is a good unit in measuring the conc. of a solution under constanttemp conditions like during an experiment in the lab. Ex. Titration

Molality (m)= moles solute = molkg of solvent kg

Since mass of a solution is temp. independent, m does not change with change intemp, hence m is a good unit to use for solution conc. when temp is changing. Ex.

B.P. Elevation, F.P. Depression *Normality (N):associated with acid & base

strength. Normality = molarity x n (where n = thenumber of protons exchanged in a reaction).

= (moles solute)(# of acid/base equivalents) =volume of solution

mol/L Ex. N of 1M H2SO4will be 2N. N was used earlier to express concentration of acids/bases

but is obsolete now.


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Dilution: MV = MV

TITRATION

-It is a method to determine the molarity of unknown acid or base

- In titration, an acid or base of unknown molarity is titrated against a standard solution (whose M is known

of acid or base.

- The end point in a titration is indicated by a color change by the indicator. Indicators are weak acids or

bases and are added in small quantity (1-3 drops) to indicate the end point.

- At equivalence point (which should be close to end point),

moles of H+ = moles of OH-

M1V1= M2V2(sometimes used to get moles , M= moles/L , so moles= M XV)

-What other ways can you get the moles-for a solid acid or base? For a gas?

-color change by indicator indicates end point

-end point and equivalence points should be close. Equivalence point is defined when moles of acid and base

become equal in a titration and end point is where the indicator changes color.


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Unit: 2Chapter: 5

Thermochem Enthalpy change: Depends on physical state ofsubstance.

Calorimetry: Measuring enthalpy change.

Specific Heat: Amount of energy required to raise thetemperature of one substance by 1 degree kelvin

The heat evolved or absorbedin a chemical process is thesame whether the processtakes place in one or inseveral steps.

energy changes are statefunctions

Hess'sLaw

Breaking bonds is an endothermicprocess

Making bonds is an exothermicprocess


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System

Isolated portion ofstudy

Surroundings

Everything else

Force

A push or pull on anobject

Work

Energy transferred tomove an object

W = FD

Heat

Energy transferredfrom a hotter ovject toa colder one.

DE > 0 Increasein energy ofsystem (gained

fromsurroundings)

Endothermic

DE < 0 Decreasein energy ofsystem (lost to

surroundings) Exothermic

DE = q + w


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qrxn

+ qsolution

= 0

q > 0: Heat is added to system

q < 0 : Heat is removed from system(into surroundings)

w > 0: Work done to system

w < 0: System does work onsurroundings


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Chapter 6LIGHT:

- For waves traveling at the same velocity (longer

the wavelength, the smaller the frequency):

o Wavelength: Distance between two

consecutive peaks or troughs in a wave.

o Frequency: Shows how many waves pass

a given point per second.

o Speed: Indicates how fast a given peak is

moving through space.

o Speed of light ,c=,where l=wave lengthand n =frequency

- Two properties that exhibit the wave like

behavior of light are interference and diffraction.

- The wave nature of light does not explain how a

object can glow when its temperature increases.

o Max Planck explained it by assuming that

energy comes in packets called quanta.

Einstein Discovered E = h

E = energy of radiation

h= Plancks constant(6.626 x 1034 J-s.)

v= frequency of radiation

Wave Particle

Electromagnetic Radiation

Electromagnetic radiationis one of the ways in which energy travels through space. All forms of EMR

compose the electromagnetic radiation spectrum, which includes sun rays, microwaves, X- rays, visible

spectrum, UV rays and IR rays.All electromagneticradiation travels at thesame velocity: the speed

of light (c), 3.00 108

m/s.

Therefore, c=

Quantization ofEnergy:


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- Einstein and DeBroglie both talked about dualwaveparticle nature.

- Photoelectric effect:It refers to the emission of electrons from a metal, when the light shines on the metal. For each

metal the frequency of light needed to release the electrons is different. But the wave theory of

light could not explain it. The photoelectric effect led scientists to think about the dual nature of

light i.e. as a wave and a particle both.

Bohrs Model of the Hydrogen Atom An excited atom can release some or all of its excess energy by emitting a

photon, thus moving to a lower energy state.

The lowest possible energy state of an atom is called the ground state. Different wavelengths of light carry different amount of energy per photon. Ex.A beam of red light has a lower energy photons than beam of blue light.

Electrons in an atom can only occupy certain orbits (corresponding to certainenergies).

Electrons in permitted orbits have specific, allowed energies; these energieswill not be radiated from the atom.

Energy is only absorbed or emitted in such a way as to move an electron fromone allowed energy state to another; the energy is defined by

E= h

Black Body Radiation

Blackbody: object that absorbs all EM radiation that

strikes it; it can radiate all possible wavelengths of EM;

below 700 K, very little visible EM is produced; above

700 K visible E is produced starting at red, orange,

yellow, and white before ending up at blue as the

temperature increases

discovery that light intensity (energy

emitted per unit of time) is proportionalto T

4; hotter = shorter wavelengthS


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- Principles:

Impossible to determine simultaneously both the position and

velocity of an electron or other particle.

Heisenburg

UncertaintyPrinciple

Mathematically: The wave properties of electrons or other verysmall particles

Quantum Theory

Orbitals with the lowest energy fill up first.Aufbau Principle

No two electrons in the same orbital can have exactly the sameenergy, so no two electrons in the same atom have the same exactfour quantum #'s.

Pauli ExclusionPrinciple

When electrons are put into orbitals having the same energy,degenerate orbitals, one electron is put into each orbital beforeputting a second electron into an orbital.

Hunds Rule

ni= initial orbital of e-

nf= final orbital of e-in its transition

D2

f

2

i

H

n

1

n

1

h

R

h

E
http://chemmovies.unl.edu/ChemAnime/BOHRD/BOHRD.html


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Electron Configuration: Shows energy level of electron, plus subshell

- Electron Configuration:

SchrodingersQuauntum

MechanicalModel

The energy of electrons in atoms is quantized.

The number of possible energy levels for electrons in atoms of differentelements is a direct consequence of wave-like properties of electrons.

The position and momentum of an electron cannot both be determinedsimultaneously.

The region in which an electron with a specific energy will mostprobably be located is called an atomic orbital.

Shows energy of orbitals,in order of smallest to

greatest, increasing bythe arrow

4p5

Energy

level

(n)

P

orbital

5 electrons in the orbital


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- Orbital diagrams Each box in the diagram represents one

orbital. Half-arrows represent the electrons.

The direction of the arrow represents the

relative spin of the electron.

VOCAB

Ground State: Lowest energy state of an atom

Excited State: A state in which an atom has a

higher potential energy then it has in its ground

state

Orbital: A 3D region around the nucleus thatindicates the probable location of an electron

Notable Exceptions:

Cr & Mo: *Ar+ 4s1

3d5

not *Ar+ 4s2

3d4

Cu, Ag, & Au: *Ar+ 4s1

3d10

not *Ar+ 4s2

3d9


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Chapter 7

S block elements: Group 1 & 2

- Chemically reactive metals, group 1 more reactive than group 2. Group Configuration: ns1-2

o Alkali metals: silvery appearance, soft enough to cut with a knife, not found in nature as free

elements. Hydrogen shares electron configuration, but not properties.

o Alkaline-earth metals:harder, denser, stronger, and have a higher melting point than group 1.

Too reactive to be found uncombined in nature. Helium shares electron- configuration but not

properties.

P block elements: Group 13-18

- Includes all the three types of elements: metals, non-metals and metalloids. Group Configuration:ns

2np

1-6

o Includes Halogens, most reactive of the nonmetals. React vigorously with most metals to form

salts.

o P block metals are generally harder & denser than s block metals, but softer & less dense than

d block metals. They are found in nature solely as compounds, except for bismuth.

D block elements: Group 3-12

- Transition Elements: metals with typical properties; good conductors, high luster.

o Less reactive than s block, many existing in nature as free elements.

o Electrons added to the d sublevel of the preceding energy level (n-1).

o Group configuration: (n-1)d1-10ns 0-2

o Some deviations from orderly d sublevel filling occur in group 4-11(s electrons jumping to d

sublevel)

F-block elements

- F-block elements are wedged between groups 3 and 4 in the sixth and seventh period, consisting of

lanthanides and actinides

o Most elements are radioactive

o Trans Uranium elements are all synthetic

o Group Configuration: ns 0-2 (n-1) d 0-1 (n-2)f 1-14

o orderly d sublevel filling occur in group 4-11(s electrons jumping to d sublevel)


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- Atomic Properties

o Effective Nuclear Charge: approximate net nuclear charge felt by the highest energy

electrons.

Eeff= ZS

S = shielding effect (e present between nucleus and valence shell

electrons), Z = number of protons

o Ionic Radii

Cation = positive ion Anion = negative ion

Atomic Radius

Half the distancebetween the nuclei oftwo atoms of the sameelement

Decreases Across

The increase in positivecharge of eachsuccesive element pullsin the electrons closerand closer, as theyremain the samedistance from the

nucleus

Increases Down

Electrons are fartheraway from the nucleus,and there is a greaternumber of electronsbetween energy levels.

Ionization Energy

Amount of energyrequired to remove an efrom a neutral atom inits gaseous state.

Increases Across

Across the period, thereis a stronger Eeff, sothere is a strongerconnection between thevalence shell and thenucleus. Therefore, itbecomes harder to pull

electrons away.

Decreases Down

Since the chargedecreases down thegroup, it becomes easierto pull the electronaway, as there is less ofan attraction.

Electron Affinity

Amount of energyreleased when e isadded to a gaseousatom in its neutral state

Increases Across

Because the Eeffisstronger across aperiod, the chargebetween the nucleusand the valence shell isstronger. Therefore,when an electron istaken, more energy isreleased

Decreases Down

The atomic radiiincreases, indicatingthat there is not muchattraction down agroup. Therefore, lesscharge means lessenergy released whenan electron is added.

Electronegativity

The measure of theability of an atom in achemical compound toattract electrons.

Increases Across

As you go across aperiod, the Eeffisstronger. The nucleus ismore easily able to holdanother electron, andthe resulting diffusion ofcharge will still result ina high charge for thenucleus

Decreases Down

As the atomic radiiincreases, the totaldistance increasesbetween the valenceand the nucleus, so itbecomes harder toattract and retainelectrons.

Flourine is arbitrarilyassigned a value of 4.0


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- Explanation of Trends


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MO Theory

Antibonding, Higher in

Energy


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The more overlap between AOs the lower the energy of the bonding orbital they create and

the higher the energy of the antibonding orbital.

Coulombs Law

where Q1represents the quantity of charge on object 1 (in Coulombs), Q2represents the

quantity of charge on object 2 (in Coulombs), and drepresents the distance of separation

between the two objects (in meters). The symbol kis a proportionality constant known as the

Coulomb's law constant.

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Chemistry: The Study Guide