2011-2012MasterOutlineChem(1)

8/2/2019 2011-2012MasterOutlineChem(1)

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Chemistry 1 Regular 2011-2012

Ms. Ehrenwerth, Mr. Hagarman, Mr. Simpson, Mr. Yinger

Fall Term

Chapter 1: Measurements and CalculationsChapter 2: MatterChapter 3: Elements, Atoms, and IonsChapter 4: Modern Atomic TheoryChapter 5: Nuclear ChemistryChapter 6: Nomenclature

Winter Term

Chapter 7: Balancing Chemical EquationsChapter 8: Reactions in Aqueous SolutionsChapter 9: The Mole and StoichiometryChapter 10: Chemical BondingChapter 11: EnergyChapter 12: Gases

Spring Term

Chapter 13: Liquids and SolidsChapter 14: Solutions and Colligative PropertiesChapter 15: Acids and BasesChapter 16: Chemical EquilibriumChapter 17: Oxidation Reduction ReactionsChapter 18: Organic Nomenclature


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FALL TERM

CHEMISTRY 1: CHAPTER 1

Measurements and Calculations

Key words or terms from this chapter: measurement, unit, scientific notation,dimensional analysis, metric system, S.I., Kelvin, significant figures,density.

: Sometimes qualitative information is good: Sometimes quantitative information is good

: When both types of information are included, these are called

measurements: Must be as accurate as possible: Must contain a unit and a number

: Both are mandatory

: Scientific Notation: A simpler way to express really big or really small numbers.: Can we express the following in notation?

: 125, 1700, 0.000065, 540000000000

: Number as expressed must be in form Mx10

n

: M must be a number between 1-10, but not 10: We would like the n to not be a decimal if possible

: Calculations with Scientific Notation: Addition or subtraction: Must have same n numbers

: if not, adjust them to be the same: Add or subtract, then adjust to sig-figs and the correct notation orn number

: Multiplication: M factors are multiplied, n factors are added algebraically: Then adjust for sig-figs and the correct notation

: Division: M factors are divided, n factors are subtracted algebraically


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: Then adjust for sig-figs and the correct notation

: Units: Units can tell a great deal about the scale of a measurement: We use two major systems, Metric and English

: Why do the English, use the Metric?: In 1960 the world agreed to use a system called S.I., the units ofwhich would be based in metrics

: Examples = Kilogram, Meter, Second, Kelvin: We can then use prefixes to denote the size of the

measurement needed.: In science we tend to use kilo, base, centi, and milli, but micro

and nano are not uncommon.: Measuring cells and such.

: Prefixes for Metric units.: Kilo = 1000 x Base unit: Deca = 10 x Base unit: Base unit = 1 (meter, second, gram): Deci = 1/10 Base unit: Centi = 1/100 Base unit: Milli = 1/1000 Base unit

: Measurements of Length, Volume, and Mass: In S.I. the unit oflength is the meter = 39.37 in.: Fractions or multiples of meters can be expressed in powers of ten.

: Km, dm, cm, mm, etc.: 1 inch = 2.54 cm

: The S.I. unit ofvolume is the Liter: Volume is a measure of 3 dimensional space

: Start with 3 length measures: L x W x H = Volume ex. 1m3 box

: Cube could be broken into 1000 smaller ones which wouldeach be 1 dm3 which is the same as 1 liter.

: Break that cube down into 1000 smaller ones and you have1cm3, which is 1ml!

: The S.I. unit for mass is the Kilogram: Mass is = to the quantity of matter in a substance


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: grams are commonly used, but are a bit small in the grandscheme of things

: Uncertainty in Measures: Whenever a measurement is made, an estimate is required at the end: You must write down one more digit than you know for certain: In 2.85 cm, the 5 is an uncertain digit

: All numbers recorded are significant: Can use the sign: Example = 2.85 0.01 cm

: Significant Figure Rules: 5 rules total1. All non-zeros are significant2. Zero Rulesa. Leading Zeros --- Zeros that precede all other non-zeros

i. These are never significantb. Captive Zeros --- Zeros stuck between non-Zeros

ii. These are always significantc. Trailing Zeros --- Zeros after all non-Zeros

iii. Not significant by themselves ARE significant ifaccompanied by a decimal point

3. Exact Numbersa. All digits included are significantb. So, 10 cm has 2 sig-figs because the measurement is

assumed to be exact

: Rounding Off Numbers: Calculations using calculators can lead to crazy numbers: Rules for rounding off numbers

: Digits being removed: Less than 5, digit stays the same: Equal to or More than 5, digit increases by 1

: If doing a series of calculations, do ALL of them, then round off: NEVER look beyond the number used in rounding

: Example --- 4.448 rounded to 2 figures, the 8 does not matteror change the third 4.

: Calculations with Sig-Figs: For Multiplication or Division


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: Correct # of Sig-Figs = least known # of sig-figs: Examples --- 4.56 x 1.4 = 6.384 = ?

8.315/298 = 0.00279027 = ?: For Addition or Subtraction: Limited by smallest # of decimals

: 12.11 + 18.0 + 1.013 = 31.123 = ?: You are limited by the least known digit

: Problem Solving using Dimensional Analysis: Breaking a whole into its parts.: Converting from a known, to an unknown unit: Dollars to quarters, kilometers to miles

: Conversions will be based on definitions.: If 1 inch = 2.54 cm

: Then you can create 2 conversion factors from this information: 2.54 cm/1inch or 1 inch/2.54 cm: Can then use this to go form one unit to another in a problem

: Put the unit you want to get rid of, opposite itself in the fraction: All conversion factors are fractions that = 1

: You may have to use more than one step is you do not have a singleconversion that works.

: Days to seconds for example!

: Temperature Conversions: In the US we use the Fahrenheit scale

: Involves water freezing at 32 degrees andboiling at 212 degrees

: In most other countries they use the Celsius Scale of temperature: Involves water freezing at 0 degrees and

boiling at 100 degrees.: MUCH EASIER!

: There is also the Absolute or Kelvin Scale: Involves water freezing at 273 and boiling at 373.

These are NOT degrees.: Things to recognize:

: The increments in the Celsius and Kelvin scale are the same.100 between freezing and boiling

: the increments in the Fahrenheit Scale are much smaller, 180between the two points.

: The Zero Point on all three is different


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: Draw a line for each scale and label 0, freezing, and boiling!: Conversions between Celsius and Kelvin

: ToC = Tk--- 273: Tk= ToC + 273

: Conversions between Celsius and Fahrenheit: Need to adjust not only for size of units, but also the different zeropoints

: ToF = 1.80 (oC) + 32

: ToC = (ToF32)/1.80

: Density: Density = Mass/ Volume

: This is a complex unit as the mass has one unit and thevolume has another.


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Matter

: Key words or terms from this chapter: matter, elements, compounds, statesof matter, chemical and physical change, chemical and physical properties,pure substances, mixtures, alloys, homogeneous, heterogeneous, filtration,distillation, heat energy.

: The Nature of Matter: Matter is Stuff: Matter has mass, and occupies space

: Matter exists in all 3 states, and sub-atomics: We usually only go to atomic levels

: Use STM to see really small things: Most things look similar at that level, sort of likegumdrops

: Elements and Compounds: All atoms are obviously not alike: 118 different kinds, + all isotopes: Compoundsa combination of atoms bonded together in a specificway

: Contain 2 or more different types of atoms: Consider H2O, this is a molecule

: Elements --- Cannot be broken down: Sometimes 2 of the same type of atoms form a molecule likeO2, H2 etc. Carbon based compounds such as diamond andgraphite.

: The Three common states of Matter

Sublimation

Melting Evaporation

Solid Liquid Gas

Freezing Condensation

Deposition


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: Must always have same chemical and physical properties: Mixtures can be separated into pure substances

: Homogeneous vs Heterogeneous Mixtures: Homogeneous --- same throughout

: Salt dissolved in watersolution: can add salt, but not necessarily a heterogeneous

: Heterogeneous --- not same throughout: Add a lot of salt!: Lower concentrations at top: Higher as you go down: maybe pure at bottom?

: Separation of Mixtures

: Distillationprocess of trapping condensed steam of a heated,impure liquid

: Process of purifying alcohol: FiltrationPhysical process dependant on particle size and filterprecision

: Coffee

: Energy and Heat: Energy = the capacity to do work: Heat = the flow of energy due to a temperature difference


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Chemical Foundations; Elements, Atoms, Ions

: Key words or terms for this chapter: elements, Daltons Atomic Theory,Law of Definite Proportions, Law of Multiple Proportions, compounds,atoms, Cathode Ray Tube, Plum Pudding Model, Gold Foil Experiment,nucleus, proton, neutron, electron, isotope, ion, average atomic mass,periods, groups, families, metals, non-metals, metalloids, alkali metals,alkaline Earth metals, halogens, noble gases, transition metals, allotropes,ionic compounds.

: Chemicals have a profound effect on how we act: Lithium treats depression

: Nitroglycerine can treat heart disease: Greeks broke down things into four elements

: Earth, Water, Air, and Fire: Alchemy ruled the chemical world till ~1600

: Why is Alchemy not possible?: Remember Gold is an Element!!!!

:The Elements: There are roughly 118 elements, but millions of compounds

: These are set combinations or ratios: Compare to letters and words in English Language: There are 88 naturally occurring elements

: 9 of those make up 98% of Earths Crust, Oceans, andAtmosphere by mass

: Oxygen, Silicon, Aluminum, Iron, Calcium,Sodium, Potassium, Magnesium, Hydrogen

: Incidentally, ants make up 18% of the Earths animal biomass(random fact)

: The human body differs not too significantly

: Oxygen, Carbon, Hydrogen, Nitrogen,Calcium, Phosphorus make up 98.4% by Mass

: Symbols for Elements: Often named for Greek or Latin names

: Just like biological taxonomy: Explains why Gold = Au Aurum: Can sometimes be named for people or places


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: Rules:: If one letter only = Must be capitalized: If two letters = First must be capitalized, second must be lower cased

: Talk about Novs--- NO

: Daltons Atomic Theory: 1808 English School Teacher

: His Atomic Theory1. All matter is composed of small particles called atoms2. Atoms of a given element are identical in size, Mass, and

properties3. Atoms cant be subdivided, created, or destroyed. (sound like

Law of Cons. Of Mass?)4. Atoms of different elements combine to form Chemical

compounds in whole number ratios: Also came up with support laws

: Law of Definite Proportions = A chemical compound contains thesame ratios of atoms, regardless of source or size.

: NaCl is NaCl is NaCl: Law of Multiple Proportions = If you have 2 compounds made oftwo identical elements; of A is constant, ratios of B are in whole #integers.

: CO and CO2

: Formulas of Compounds: Compounds = distinct substance that is composed of atoms of two ormore elements and always contains the same relative masses ofatoms

: H2O is always water: Look at H2O and what that tells us as a formula

: Two atoms of Hydrogen: One atom of Oxygen

: Rules:

: Must use good symbols of elements: Place amounts in subscript to right

: C6H12O6, N2O5, SO3

: Structure of the Atom: Not until ~1900 did we think atoms had parts!: A series of experiments by a bunch of geniuses proved this to us


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: J.J. Thompson: Said that an atom of an element can be made to repel a negativecharge

: he felt that this must mean that they contain a negative part?: The Cathode Ray Tube

: Also then deduced some things: Since atoms are neutral, they must have a positive part!: Felt that since eare so small, there must be more things in

the atom that have mass (no): William Thompson (Lord Kelvin)

: The Plum Pudding Model: Decided that there were both (+) and (--) charged particles floatingaround in a confined area

: Like Pudding with raisins

: Ernest Rutherford (1911): Sent a (+) beam through a gold foil

: About 1/8000 were deflected back: He though that the (+) particles must have hit something that wasvery massive and also (+).

: He called this the nucleus (1919): By 1932 he had also defined the neutron as the other part of thenucleus that had mass.

: Modern Atomic Structure: Since the above experiments, much has been clarified about the atom: Sizes in an atom:

: Nucleus = 1013 cm in diameter: Electron orbitals = 108 cm from nucleus

: Masses in Atom:: Proton (p+) = 1.673 x 1027 Kg: Electron (e--) = 9.109 x 10---31 Kg: Neutron (no) = 1.675 x 10---27 Kg

: Isotopes: Consider Sodium and its 11 protons: We know that it must also have 11 electrons

: Atoms want to be neutral: Neutrons are not considered here, and they, unlike electrons have asubstantial mass


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: Group 18 are the Noble Gases: They are also generally unreactive

: Many things on the table are diatomic: Means they form two atom molecules: H, O, N, F, Cl, Br, I

: Metaloids are kind of crazy: Br, I, etc.

: Very few elements are liquid at 25oC: Br, Hg only two: Gallium and Cesium almost qualify (30oC)

: Allotropes = different forms of the same substanceI.E.= carbon molecules

: Graphite, Diamond, Bucky-Balls: Most common in solid, non-metalics

: Ions: We know that the number of protons = the number of electrons in anatom

: This means that atoms are neutral: We can create a charged entity by adding or subtracting one or moreelectrons from a species

: + Ions we call Cations: Sodium is a good example. Na+1

: - Ions we call Anions: Chlorine is a good example. Cl-1: Anions are named using the root of the ion + --ide.Chlorine becomes Chloride!

: We can never form an ion by adding or subtracting protons from theatom, why?

: Metals tend to lose one or more electrons: Non-Metals tend to gain one or more electrons: Trends in the Periodic Table

: Group 1 = Forms +1 Ions Na+1

: Group 2 = Forms +2 Ions Mg+2: Group 13 = Forms +3 Ions Al+3: Group 16 = Forms ---2 Ions S-2: Group 17 = Forms ---1 Ions Cl-1: Transition Metals = Form usually + ions of differing charges

: Compounds that Contain Ions


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: Many of the compounds that we encounter contain ions: NaCl is a good example

: Melting Point = 500 oC: Boiling Point = 1500 oC

: As a solid, it will not conduct electricity: As a solution, it will conduct electricity: Why does it exhibit these characteristics?

: The ions that make it up, are so attracted, that they really do not want toseparate

: This would make boiling really hard.: Current travels along a wire due to the fact that electrons of the wire

material are allowed to move through the wire freely: The ions in a compound can not move

: Do tons of examples of Ionic Compounds

: Must have a net charge of 0


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Modern Atomic Theory

: Key words or terms in this chapter: Rutherford, electromagnetic radiation,wavelength, frequency, speed, photon, ground state, excited state, Hydrogenline emission spectrum, Bohrs Model, Wave-Mechanical model, orbit,orbital, S,P,D,F-Blocks, Orbital notation, Noble gas notation, Hunds Rule,octet rule, Aufbau Principle, Pauli Exclusion Principle, Valence electrons,unpaired electrons, atomic radius, ionization energy.

: Once atoms were an accepted thing, we next wanted to know howthey worked

: + nucleus and a group of electrons

: We most of all wanted to know how elements in same groups actedsimilarly, and why those in different groups acted differently

: Rutherfords Atom: Used the metal foil and alpha particles: Showed that nucleus contained protons and neutrons, densely packed: Did NOT tell us what electrons were doing

: Why didnt atoms collapse?: Electrons are attracted to the protons and much smaller?

: Energy and Light: Why does fire warm your body if you are not on fire yourself?: Energy can be transferred by light

: Electromagnetic Radiation: Energy is ultimately transferred in waves, and each wave has threeparts

: Wavelength --- peak to peak, trough to trough: Frequency --- how many waves past a point in a given time

period

: Speed --- how fast is the wave moving: Photons --- smallest form of light energy (photon torpedo?): SO, we see light as both waves and particles

: Small waves carry higher energy photons: Larger waves carry smaller energy photons

: In general sort of like crowd surfing!


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: Emission of Light by Atoms: Different wavelengths of light fall into energy ranges: Different colors can be emitted as atoms fall from these excitedstates

--- Energy Levels of Hydrogen: Any atom with excess energy is in an excited state: This can be released in a photon: Now the atom is at a lower energy state

: All H atoms will release certain light: All light is = to a certain value

: NO in betweens, like steps or a ladder

4th___________________

3rd ________________________

2nd _____________________________

1st ___________________________________

Ground State _______________________________________________

: Atoms emit light based on where they start, and where they fall to: Remember, no in between, only on the steps!!!!!!!

: Bohrs Model of the Atom: Bohr knew the atom pretty well!: He said electrons orbit the nucleus on levels: Electrons may jump levels or fall levels based on the levels of thehydrogen line emission spectrum

: Turns out he was only right for Hydrogen: Turns out it didnt matter

: At age 37, for his theory, which we now know is mostlywrong, he received the Nobel Prize!

: The Wave Mechanical Model: DeBroglie and Schrodinger suggest that maybe electrons work asboth waves and particles

: Schrodinger, the mathematician, found that this model worked for


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many models of atoms, not just hydrogen: Orbitals, NOT Orbits: Orbital = a set of locations that the mass could be in (3

dimensional): Orbit = a set path, of movement (2 dimensional)

: The Hydrogen Orbitals: Orbitals do not have a defined exact size: There are no set boundaries

: The electron CAN travel outside of the orbital: Bohr was right on here, and we see this ability in the teaching oforbitals

: Orbitals and Notations

: S Orbitals --- groups 1--- 2 = (n): P Orbitals --- groups 13 --- 18 = (n): D Orbitals --- groups 3 ---12 = (n1): F Orbitals --- Lanthanide and = (n - 2)

Actinide Only

Draw a Map of the Periodic Table including Block name, number ofelectrons across each block, number of orbitals across each block, and n #.

Think about it this way, each house on a block must have a different address,so each element must have a different address so that we can find it. Alwaysstart at the beginning and then go until you reach the element that you areinterested in.

: Notations: Orbital, Long-Hand, and Noble Gas; do them all, about 1000 times.

: Hunds Rule = put one e in each orbital before putting asecond one. Like seats on the bus!

: Octet Rule = all elements want to have 8 valence electrons.

: Aufbau Principle = we add electrons in the lowest availableenergy level first.

: Pauli Exclusion Principle = in each orbital, you must have oneup and one down electron once full.

: Valence Electrons = the S and P electrons in the highest energy level onlyfor each element.


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: Unpaired Electrons= Go to the highest block being used and look for any

unpaired electrons left over.: Max in an S block = 1: Max in a D block = 5: Max in a P block = 3 Why?????

: Periodic Properties: Must look at both horizontal and vertical trends

: Metals vs Non-Metals: MetalsPhysicallylustrous, malleable, ductile, conducts heat and

electricityChemically --- tend to lose electrons, form + ions

: Non-Metals --- Physically --- Opposite, with exceptionChemically --- tend to gain electrons, form --- ions

: Metalloids --- tend to exhibit both qualities

: Numeric Trends ---: Atomic Radius--- how big is the atom from nucleus, to electronorbitals

: Decreases up groups, and decreases across periods: The larger energy levels are naturally larger than the smaller: As you add electrons across the period, you are also addingprotons, which are a significantly larger particle ( LikeGravity sort of.)

: Ionization Energy --- energy needed to remove an electron (1st, 2nd,and 3rd are all common)

: Increases up groups, and increases across a period: Francium really wants to lose an e---: Helium really doesnt want to!

: Chemical Reactivity, generally speaking.

: We would say that the most chemically active metals are inthe lower left hand corner, and the most chemically activenon-metals are in he upper right hand corner.


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Chemistry 1: Chapter 5

Nuclear Chemistry

: Key words or terms in this chapter: radioactive nuclide, Marie Curie,radioactive decay, Alpha, Beta, Positron, Gamma, Electron Capture, nuclearequation, life, decay series, radioactive dating, fusion, fission, somatic,genetic.

: Radioactive Decay: 1896Henri Becquerel = photographer.

: Looked at the emissions of light photons: Found that Uranium could expose light sensitive paper

without using light?: Some other energy source must have been at work here: He called this radioactive decaythe nucleus disintegrating into alighter nucleus which is accompanied by the release of energy,electromagnetic radiation, or both.

: radioactive Nuclideunstable nucleusthat undergoesradioactive decay

: Marie and Pierre Curie (1896)monumental work with radioactiveelements lead to the winning of 2 Noble Prizes. Too much radioactivity =not good!!!!

: Uranium and Thorium were thought to be the only radioactiveelements, but they proved Polonium and Thorium as well.

: We later learned that all nuclides over # 83 are radioactive: Types of Radioactive Decays: Types Symbol Charge Mass (amu)1. Alpha 4 2+ 4.00266

2 He2. Beta 0 1- 0.0005486

--1 B

3. Positron 0 1+ 0.0005486+1 B4. Gamma Ray Symbol 0 0.0


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: Nuclear Equations: Z# and A# must both be conserved, in short, the tops and bottoms

add up with the arrow being the = sign. Just dont forget to multiply thecoefficients of there are any.

14 14 06 C 7 N + + -1 e

: Specific Emissions

: Alpha Emissions: Contain 2 Neutrons and Protons bound together

: Really just a Helium nucleus (+2): This happens to only heavy nucleus

: Both Protons and Neutrons have to be released to maintainstability

210 206 484 Po 82 Pb + + 2 He

: Atomic Number decreases by 2, but Mass Number decreases by 4.Why?

: Beta Emissions: Happens to elements that have too many neutrons

: Above the band of stability: To decrease # of Neutrons, can convert into a Proton and an Electron

: Emitted as a Beta Particle

14 14 06 C 7 N + -1 B

: We lost a neutron, picked up a proton, and emitted an electron

: Positron Emissions: Happens to elements that have too many Protons

: Below the band of stability: To decrease # of Protons, can convert into a neutron with a +chargefrom the nucleus


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38 38 019 K 18 Ar + +1 B

: Electron Capture: An inner orbital electron is captured by its nucleus, combines with aproton, and forms a neutron

106 0 10647 Ag + -1 e 46 Pd

: Gamma Emissions: High energy electromagnetic waves emitted from the nucleus as itchanges its excited states

: The existence of this energy source supports nuclear shellmodels

: This is usually a secondary decay source

: Radioactive Half-Life: No two isotopes decay at the same rate: Life = t = m for x time

: x must be defined: each life, the remaining mass decays

: Nuclear Decay Series: A series of nuclear emissions that results in a more stable nucleus

: Radioactive Dating: Uses defined life information to date an artifact.: We often times use C-14 as the basis of this

: Discovered in the 1940s by Willard Libby: Nobel Prize

: Based on two opposite processes

: The breakdown of C-14 over time,

14 14 06 C 7 N +

-1 B

: And the atmospheric process of making C-14


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14 1 14 1

7 N + 0 n 6 C + 1 H

: This works well for plants as well due to atmospheric processes: Once cut, they can no longer take up the C-14 and then the decay begins

: Clothing from Biblical Times was made of Flax: Flax, like cotton or linen, was made from plants!

: Fusion and Fission: The nucleus is held together with some of the greatest forces in all ofchemistry

: FusionThe combining of two or more nucleus to form a more complexnuclide

: Problem is, the two nuclides repel each other: SO, you must have a great deal of energy to get them to bind (40

million K for Deuterium!): FissionThe splitting of a nuclide into two smaller, more stable nuclides

: Doesnt give as much energy as fusion, but it is a great deal easier toget going

1 235 141 92 1

o n + 92 U

56 Ba + 36 Kr + 3 0 n: Where is the problem here? Getting the reaction to slow down.

: Nuclear Reactors work on this process

: Effects of Radiation: Getting hit by a train is not good: Getting hit over and over again by a neutron is often-times just asbad

: The cumulative effect on the body is important

: Damage is broken into two types, somatic and genetic: Somaticdamage to the organism itself, causing sickness or death: Geneticdamage to the gametes or genetic material, often causingproblems with offspring

: The above depend on the following factors: Energy of the radiationhigher radiation, higher damages


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: Penetrating ability of the radiation --- vary according to values;gamma are high, beta can penetrate 1 cm into skin, alpha are stoppedby healthy skin

: Ionizing ability of the rays --- ability of the ray to form an ion from anormal atom or molecule.

: Inverse of the penetrating ability. Alpha are the Worst at this.: Chemical properties of the radiation source ---General reactivity must be analyzed.

: Calcium is very inert, yet Strontium can lodge in the bonesand cause cancers


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Nomenclature

: Key words or terms in this chapter: binary ionic compounds, polyatomicions, molecular compounds, acids, formula names.

: There used to be no system for naming compounds: We simply used their common names

: Problem, there are over 4 million recognized substances.Need names.

: To memorize all their names would be impossible: Need a set of rules, and then name things based on composition

: We will begin with compounds which can be broken down into 3 types: Metal + Non-Metal: Non-Metal + Non-Metal: More Complex Compounds

: Metal + Non-Metal Compounds: When as metal like Na and a non-metal like Cl form a compound,they use ions

: Metal must lose one or more electron: Non-Metal must gain one or more electron

: ResultBinary Ionic Compound: Type 1 and 2

: Type 1 --- Metals that only form one ion: Cation first, anion second: Simple cations take name of element: Simple anion takes root of element + -ide

: examples--- NaCl, KI, CaS, CsBr, MgO: Type 2 --- Metals that form more than one cation

: Gold Chloride is not enough information

: Gold forms+

1 or+

3 ions: Answer the ambiguity by using roman numerals: These will be placed between the cation and anion

: Gold (I) Chloride or Gold (III) Chloride: Sometimes you can work backwards from the anion and get

the charge of the cation: PbO2, must use Pb

+4 if it is going to cancel out 2 O2


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: Binary Compounds containing only Non-Metals

: Rules --- based on names of cations and anions, and also a prefixsystem to denote amounts

: First element in formula is named first: Full name is used no matter what

: Second element is named as if it is a normal anion: Prefixes are used to denote numbers

: Prefix monois never used for cation: Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Octa.: Sometimes we drop certain letters.

: Tetroxide, Monoxide.

: Naming Compounds that contain Polyatomic Ions: Polyatomic Ion --- a charged entity made up of more than one atomthat carries a defined charge.

: Their names MUST be memorized: Oxyanions--- contain oxygen combined with atoms

: Can remember by learning the chlorine series: ClO--- Hypo--- chlor ---ite: ClO2

--- Chlor ---ite: ClO3

--- Chlor ---ate: ClO

4

--- Per--- chlor ---ate: May need to use Roman Numerals if necessary

: Depends on the Cation type (1 or 2)

: Naming Acids: When dissolved in water certain compounds release H+ ions: We call these acids; first noted for sour taste: Acids can be viewed as molecules with one or more H attached to ananion

: Naming depends on Oxygen

: Anion which does not contain Oxygen: Use Prefix Hydro--- and Suffix ---ic attached to the root of the anionname

: HCl = Hydrochloric Acid: H2S = Hydrosulfuric Acid: HCN = Hydrocyanic Acid

: Anion which does contain Oxygen


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: uses root name of the central atom in the anion with the suffix ---icor ---ous.

: When anion name ends in ---ate, use ---ic for acid: When anion name ends in ---ite, use ---ous for acid

: Try these; H2SO

4, H

3PO

4, H

2SO

3, HNO

2, HC

2H

3O

2, HClO, HClO

2, HClO

3,

HClO4

: Formula Names: It is important to be able to go both ways.: Must be able to work backwards sometimes, but that is notnecessarily a bad thing.


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WINTER TERM


Balancing Chemical Equations

: Key words or terms in this chapter: chemical reaction, chemical equation,reactants, products, coefficients.

: How do you determine the correct amounts of things in a chemicalreaction

: there are millions of reactions around us: Things burn, plants grow, metabolism

: Evidence of a chemical reaction: How do you tell if a rxn has even occurred?: Sometimes visual --- a metal rusts: Sometimes thermal --- hot or cold produced

: Always, forms a new substance

: Chemical Equations: Shows the rearrangement of atoms: When one of more new substances has formed, a chemical equation

has occurred: When symbols are shown, this is the equation

: Contains reactants and products: All true chemical equations follow the law of conservation of mass: Also should show states of reactants and products. (s), (l), (g), (aq)

: Balancing Chemical Equations: Can only change coefficients: Rules ---

: Start with shopping list for reactants andproducts: Start with most out of balance: Domino Rule --- if that threw anything off,

go to it next: Water Rule --- Try even number of waters, if

you are having trouble


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: Treat polyatomics as a unit, not separate: Once you get an atom balanced, never

change it, but you could have to multiply


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Reactions in Aqueous Solutions

: Key words or terms in this chapter: precipitate, aqueous, solubility,molecular equation, complete ionic equation, net ionic equation, acid/basereaction, Redox, double replacement reaction, single replacement reaction,combustion reaction, synthesis, decomposition.

: Oftentimes, reactions of importance happen inwater

: Biochemical reactions are as such: The reactions of the body

: First we must decide if a reaction even occurs

: Predicting Reaction Occurrence: Must consider some ?s: Why does a reaction occur?

: What causes certain reactants to form certain products: We notice certain tendencies in reactions

: These tendencies push a reaction: The formation of a solid (most stable form): The formation of water

: Transfer of electrons (creates stability): Formation of a gas

: Rxns that form Solids: The formation of a solid is called precipitation: Solid = the precipitate, and the reaction is called a precipitation

reaction: K2CrO4(aq) + Ba(NO3)2(aq) a yellow solid: Which product is the yellow solid?

: We must first consider what happens to an ioniccompound in water: It will dissociate: This allows for the new water to hold a charge (Unlike pure

water): When a species completely dissociates, it is called a strong

electrolyte


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: Things we know:: There are four types of ions in solution: The precipitate must have a net charge of 0

: Look at all possible combinations of four ions: eliminate the reactants, then decide which is the solid and

which is aqueous: Using solubility rules

: Use the facts: Use concepts; memorize the rules, interpret

: Solubility Rules: NO3

- salts are soluble: Na+, K+, and NH4

+ salts are soluble: Cl-, Br-, and Isalts are soluble,

except Ag+

, Hg2+

, Pb2+

: SO4

2salts are soluble, except Ba2+, Pb2+, Ca2+

: S2-, CO32-, PO4

3- salts are insoluble: OHsalts are insoluble, except Na

+, K+, Ca2+

: Tips in predicting precipitate formation: Write reactants as the exist before the rxn

: Write them as ions: Consider the various solids that could form

: Simply exchange anions to get salts: Use the rules to predict solid formation

: --- Describing Reactions in Aqueous Solutions: Involves molecular, complete ionic, and net ionic equations

: Molecular equations --- Shows the complete formulas of all reactants andproducts

: Does not give a very accurate picture of what actually happens

: K2CrO4(aq) + Ba(NO3)2(aq)BaCrO4(s) + 2 KNO3 (aq)

: Complete Ionic Equations --- Shows the strong electrolytes as ions andprecipitate as solid

: Gives a better idea of what occurred

: 2 K+(aq) + CrO42-

(aq) + Ba2+

(aq) + 2 NO3-(aq)BaCrO4(s) + 2 K

+(aq) + 2 NO3

-(aq)


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: Notice the K+ and the NO3--- are on both sides of the reaction bythemselves

: They are watching the reaction and are thus called spectator ions

: Net Ionic Equations --- Only shows those components directly involved inthe reaction

: Ba2+(aq) + CrO42-

(aq) BaCrO4(s)

: Do all forms for the equation below;

: Pb(NO3)2 (aq) + Na2SO4 (aq)

: Reactions that form Water

: This involves adding Acids and Bases: Bases --- first associated with bitter taste and slippery feeling (think

soap): Acids --- first associated with sour taste and dry feelings (think

lemon juice on teeth): Arrhenius proposed theories on acids and bases inthe 1800s

: Acids --- produce H+ ions in solution: Strong acids completely dissociate

: Bases --- produce OH--- ions in solution: Strong bases completely dissociate

: Combining acids and bases forms water and a salt

: HCl (aq) + NaOH (aq)H2O(l) + NaCl (aq)

: Net Ionic says a great deal

: Metals + Nonmetals: Called an Oxidation Reduction Reaction: Involves the transfer of electrons

: 2 Na (aq) + Cl2 (aq) 2 NaCl (aq): What happened? Original species had no charge: Then electrons were transferred and ions were formed in the process: Two ions combine to form an uncharged species

: Metal gives away electrons --- oxidized: Non-Metal takes electrons --- reduced


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: Types of Reactions: Precipitation Reactions --- accompany the formation of a solid

: Also called Double Replacement Reactions: AB + CD AD + CB

: Acid Base Reactions --- forms water and a salt: Also called Neutralization Reaction

: Oxidation Reduction Reactions --- Involves electron transfer: Involves metal + non-metal

: Single Replacement Reactions --- format: A + BC B + AC

: Zn (s) + 2 HCl (aq) H2 (g) + ZnCl2 (aq)

: Combustion Reactions --- Usually form H2O and CO2: Actually a special type of Redox: Usually uses an organic solid with Oxygen

: C3H8(l)+ 5 O2 (g) 3 CO2 (g) + 4 H2O (g)

: 2 C8H18 (l) + 25 O2 (g) 16 CO2 (g) + 18 H2O (g)

: Synthesis Reactions --- simple materials into more complexmaterials

: 2 H2 (g) + O2 (g) 2 H2O (l)

: 2 Na (s) + Cl2 (g) 2 NaCl (s)

: Also technically redox!: A + B AB

: Decomposition Reactions --- complex material into simpler materials

: 2 H2O (l) 2 H2(g) + O2 (g)

: AB A + B


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: Scientific Notations and the above MassMass

: Do examples galore

: Comparing two Reactions: 1st you must have two balanced reactions: Consider the following two reactions

NaHCO3 (s) + HCl(aq) NaCl(aq) + H2O(l) + CO2 (g)

Mg(OH)2 (s) + 2 HCl(aq) 2 H2O(l) + MgCl2 (aq)

: If given 25 grams of each of the antacids, which would be capable ofkilling more HCl(aq)

: Limiting Reactants: Remember that one hockey team = 1 goalie, 2 defense, and 3

forwards: What if I have;

: 3 goal, 6 defense, and 8 forwards?, how many teams do I havenow?

: Not enough forwards?, then they are limiting you.: Not exactly like this, but close!!!!

: Calculations using limiting reactants: Consider; CH4(g) + H2O(g) 3 H2(g) + CO(g)

:How many grams of H2O in grams do I need to completely react with 249grams of CH4?

: What if I add 300 grams of H2O?: Who limits the reaction then?

: Now consider; N2(g) + 3 H2(g) 2 NH3(g): If I start with 2500 grams N2(g)

500 grams H2(g): How much NH3(g) can I form in grams?: What is in excess and how much?

: Steps in solving these problems:: First use the two scenario method

: Work out each problem separate, then decide


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: Choose the scenario that will actually happen and use the twoamounts from that to solve for the product amount.

: Then use the reactant amounts to solve excess: Then if given data, solve for a percent yield

: Actual Yield/ Theoretical Yield x 100 =

:Empirical and Molecular Formulas: Empirical Formulasleast common denominator: Molecular Formula --- a whole number multiple of the empirical

formula.

: How to calculate:: EmpiricalConvert all numbers to moles!: Remember!, moles are the only way to compare different

elements.: Now divide each number you have, by the smallestnumber you calculated.

: Now multiply if necessary to get all whole numbers.: MolecularThey have to give you the molar mass of the

compound.: Take the given molar mass and divide it by the molar mass of

the empirical formula you decided on.: That is your whole number multiple!


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Chemical Bonding

: Key words or terms in this chapter: chemical bond, electronegativity, ionicbond, polar-covalent bond, nonpolar-covalent bond, valence electrons,Lewis Structures, VSEPR Theory, linear, bent, tetrahedral, trigonal planar,trigonal pyramidal, octahedral, lone pairs.

: Most things in nature contain bound atoms: Exceptions to this are Argon and Helium

: The way that things are bound affects the way they work chemically andphysically

: Look at Carbon?, Diamond vs Graphite: Some bonds affect how certain reactions occur: Even small differences can be catastrophic: Heroin vs Endorphins vs Morphine

: Types of Chemical Bonds: Bond --- a force that hold together groups of two or more atoms and

makes them function as a unit.: We can look at the strength of the bond by how much energy itwould take to break it

: Bond Energy: Electronegativity --- The ability to attract an electron to an atom

: Highest --- Flourine: Lowest --- Francium

: Think about this trend, and why: Types of Bonds

: Ionic --- Bond forms due to charges of ions: Na and Cl: Very difficult to break apart

: High melting points (800o

C): Polar Covalent --- Forms due to small difference in electronegativity: CO is a good example.: Different, but not by much

: Like uneven sharing of electrons: Non-Polar Covalent --- Forms when two identical atoms are bound

together


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: O2 is a good example: Cant have a difference in electronegativity

: Lewis Structures: How to determine shapes of common molecules.: Formula help = (needshas) = # of bonds

2: each bond = 2 electrons: has is total # of electrons 2 for each bond = left overs.


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Energy

: Key words or terms in this chapter: calorie, joule, Q, specific heat capacity,endothermic, exothermic, Hesss Law.

: What exactly is a calorie anyway?: calorie = amount of energy needed to raise 1 gram of water 1oC

: food calorie is actually a Calorie = 1000 calories.: joule = 4.184 calories

: Q = s x m x T

: Q = heat energy gained in joules: s = specific heat capacity of sample material: m = mass: T = change in temp. in degrees C

: Solves for any of the above, or even starting or ending temp!

: Exothermic = reaction where heat energy is given off: Endothermicreaction where heat energy is taken in

: Hesss Law: a way of solving for thermal data indirectly using smallerequations.

: RulesIf a reaction is reversed, reverse the sign of the thermal dataIf a reaction is X or /, do the same to the thermal dataSum the reactions and cancel.


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Gases

: Key words or terms in this chapter: pressure, barometer, mmHg, atm, KMTof gases, directly proportional, inversely proportional, Boyles Law,

Charless Law, Avagadros Law, Ideal Gas Law, Daltons Law of PartialPressures.

: We actually live in a sea of gases: The atmosphere is mostly made up of N2, O2, andH2O gases

: Some others exist as trace gases

: Why is a balloon able to be filled up?: Why does it expand when it is heated up, and contract when it iscooled down?

: Pressure: Gas is able to fill a container completely: Gas is able to be compressed: Gas will uniformly mix

: Gases exert pressure on their surroundings: The pressure is coming from the inside and outside of the container

being viewed: The outside pressure we call atmospheric

: The Barometer measures atmospheric pressure: invented by Torricelli in 1643: At sea level, on average, a column of 760 mm Hgcan be supported by atmospheric pressure

: This comes from Gravity pulling on the molecules in the air; thehigher up you are, the less molecules to pull on, the less pressure!

: Units --- lots of them

: 1 mm Hg = 1 torr: SO, 760 mm Hg = 760 torr = 1 atm. = 101,325 pascals = 14.69 psi: You should, given this, be able to convert!

: Pressure and Volume Relationships: We call this Boyles Law: says that PV = K at constant T


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: P1V1 = P2V2: Can rearrange to solve for any of those: Make sure constants are same throughout

: Inversely or directly proportional?

: Volume and Temperature Relationships: We call this Charles Law: says that V/T = B (another Constant)

: at constant P: SO, V1 = V2

T1 T2: This was used to determine the value of AbsoluteZero, and so we use the Kelvin scale for these problems

: Inversely or directly proportional?

: Volume and Mole Relationships: We call this Avagadros Law: Says that V/N = F (another constant)

: at constant P and T: SO, V1 = V2

N1 N2: Basically, at a constant temperature and pressure, if you double thenumber of moles of gas, the volume doubles as well

: The Ideal Gas Law: How do we combine all the things we have learned about gases intoone big rule?

: Consider; Boyles LawCharles LawAvagadros Law

: Can combine into PV = nRT: R = Universal Gas Constant =

: 0.0821 (L x atm. / mol x K)

: I should be able to give you a group of numbers, andyou solve for the unknown

: Can have a few things change and still solve

: Daltons Law ofPartial Pressures: Many gases contain individual parts that we may be interested in =Air


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: Studies would show that these parts act independently of one another: In 1803, Dalton studied how parts of Gases act: Said that; For a mixture of gases in a sealed container, the totalpressures exerted is = to the sum of the pressures of the individualgases

: OR; Ptotal = P1 + P2 + P3 + Petc..: Solve the Ideal Gas Law for Ptotal of 3 gases: Find that ntotal = RT/V: Proves a few Things

: The Volume of the individual gas particles doesnt matter(they are all too small)

: The forces among the particles is negligible: SO, 1.75 mol Argon in 5.0 L at 20oC

1.75 mol Helium in 5.0 L at 20oC

: Same force!: A typical problem for this Law is the collection of a gas over water

: Must account for H2O pressure in flask: Given: O2 is collected over water at 22

oC resulting in O2 and H2O with avolumeof 0.650 L and a pressure of 754 torr. Calculate the pressure andnumber of moles of O2. The vapor pressure of water at 22

oC is 21 torr.: get Poxygen, convert to atm., solve for n using ideal gas law.

: Laws and Models: Ideal gas law can be used in all cases

: Sometimes gases do not obey the rules: Low temperatures, high pressures: SO, we avoid these conditions

: K.M.T. of Gases: Gases consist of tiny particles

: The particles are so small in relation to the distance betweenthe particles, that the particle size can be assumed to be zero

: The particles are in constant motion

: that motion, causes pressure: the particles neither attract, nor repel each other

: Average kinetic energy of a gas is proportional to the Kelvintemperature of the gas

: What did that just mean?: It tells us the true meaning of temperature


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: High temp. means high collisions, means high pressures, means theopposite is also true

: Pressure and temperature relationships: Pressure is a function of motion and collisions: SO, as temp. increases collisions, it must increase pressure: Volume and temperature relationships: Temp. increases so motion increases so pressure increases so if oneside of the container can collapse, then it will, and volume decreases.

: Real Gases: Ideal gases follow PV = nRT: But, Ideal gases are theoretical

: We should only try to do problems at normal conditions

: Gas Stoich: Like normal, but then use PV = nRT to get a volume: This means more than moles

: Generic Conversion = 1 mol of any gas at 1 atm. and 273 K =22.4 Liters!

: Can do regular, or limiting reactant.


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SPRING TERM


Liquids and Solids

: Key words or terms in this chapter: intermolecular forces, intramolecularforces, MHF, MHV, Hydrogen bonding, Dipole-Dipole, London Dispersionforces, phase change diagram, vapor pressure, ionic solids, molecular solids,atomic solids.

: Flying, ice skating, and swimming are all made possible throughinteractions with water

: We should now look into the other major states: Gaseslow densities and high compressibility

: Completely fill a container: Solidsmuch higher densities, slightly compressible

: Will maintain shape regardless of container: solids are rigid, and so pieces are close together

: LiquidsHave properties somewhere between solids and gases: NOT midway: H2O(s) H2O(l) = 6 Kj/mol: H2O (l) H2O (g) = 41 Kj/mol

: What does this mean?: Solids and Liquids are closer related than Gases

: We also see this in density values: Solid = 0oC and 1 atm. = 0.9168 g/cm3: Liquid = 25oC and 1 atm. = 0.9971 g/cm3: Gas = 100oC and 1 atm. = 0.000588 g/cm3

: Intermolecular Forces: We expect things with small molecules to be gases: Why isnt H2O a gas then at normal temperatures and pressures

: Intermolecular Forcesthe forces that exist between two molecules: Intramolecular Forces --- the forces within the molecule that hold ittogether: Dipole-Dipole Moment --- the forces between oppositely charged ends oftwo polar molecules in close proximity to one another: Hydrogen Bonding --- a particularly strong dipole moment betweenhydrogen, and another highly electronegative element


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: Can use water as an example and how Hs are attracted to the other Os: London Dispersion Forces --- forces between Noble Gases and non-polarmolecules

: Even Noble Gases can become liquid or solid at certain conditions: Like an instantaneous dipole

: Like if ALL electrons ACCIDENTALLY end up on the same side ofa molecule

: They wont be there for long, but when they are, they will beattractive

: Water and its Phase Changes: Making up 70 % of the Earths Surface and 97% of the oceans: Water both directly and indirectly effects us

: Cools atmosphere, cars, power plants

: Creates much of the Earths weather: At 1 atm. it is a liquid between 0oC and 100oC

: That is a HUGE range: Look at the Phase Change Diagram

: Water expands as it cools and therefore Ice floats on liquid

: Energy Requirements for Phase Changes: Remember, a state change is physical, not chemical: It takes energy to create steam or ice: Intermolecular Forces must be broken

: Molar Heat of Fusion --- amount of energy needed to melt one mole of asubstance

: For Ice, M.H.F. = 6.02 Kj/mol at 0oc: Molar Heat of Vaporization --- amount of energy needed to change one

mole of a liquid into a vapor: For Water, M.H.V. = 40.6 Kj/mol at 100oC

: Discrepancies in values show that liquids are closer related to solids, thangases

: Gases need 8 times the energy to form

: That is common in many substances: 2 types of problems we can do.

: Solve for energy amount needed or given off: Solve for MHF or MHV

: Evaporation and Vapor Pressure: for any substance, to create a gas takes an energy input


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: this causes evaporation, or vaporization: How does this happen?: Temperature = the average kinetic energy of the particles: Some particles are actually higher, some are lower: If the particle has enough energy to overcome the intermolecularforces, it leaves in the form of a gas

: This means that the average K.E. of the sample went DOWN: So, evaporation is a cooling process and is endothermic: the temperature has DECREASED: Vapor Pressureplace a liquid in a closed container and observe: Liquid will decrease in volume as it evaporates: When rate of evaporation = rate of condensation, Equilibrium has

been reached: This is a highly dynamic state

: The pressure that the newly formed gas places on the liquid, is calledthe Equilibrium Vapor Pressure

: The easier that something evaporates, the higher the E.V.P.: Pure Hg = 760 mmHg: Hg with H2O = 736 mmHg (24 mmHg for the water): Hg with Diethyl Ether = 215 mmHg (545 mmHg for the Ether)

: Boiling Point and Vapor Pressure: There is always gas in water: What causes the gases in a liquid to want top escape: When gases in a liquid gain enough energy to overcome atmospheric

pressure, they will escape, and boil out: At 1 atm, you need 100oC to allow for enough energy to have the gas

escape for a sample of water: At higher altitudes, there is less pressure, and so things boil at

lower temperatures: At lower altitudes, there is more pressure and so things boil at

higher temperatures: In a vacuum, room temperature water boils

: General Rules:: Molecules with higher electrical attractions boil at higher temps.: Molecules that are physically larger will boil at higher temps.: More polar molecules will boil at higher temps.

: Solid State and types of Solids: Solids are a part of our everyday lives


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: Most solids are mixtures: Some, such as diamonds and salt are pure substances

: Crystalline Solids --- contain regular arrangements of their parts: We call these geometric arrangements crystals: There are 2 major types and1minor type of these solids: Ionic Solids --- will form ions if crystal breaks

: Ions will conduct electrical current: Molecular Solids --- Forms only neutral solids

: Will not conduct electrical current: Atomic Solids --- Contains only one type of atom

: Covalently bonded to one another: Think of Diamond, Graphite, and pure metals

: Properties of Solids: Largely depends on forces holding them together

: Diamond, salt, and copper are all WAY different

: Bonding in Solids: Ionic Solids

: Very high melting points due to charged bonds: Best thought of as balls packed together tightly

: Molecular Solids: Tend to melt at low temps because the fundamental particle is the

molecule: This means that usually, not always, these are held together with some

form of the Dipole: Atomic Solids

: Vary greatly in tendencies: Usually form network solids involving insane amounts of

atoms for the sample: Metals --- sometimes for Alloys

: There are two major types of alloys: Substitutional Alloys --- where host molecule isreplaced in certain places by another type of molecule

: Brass = 66% Copper + 33% Zinc: Sterling Silver = 93% Silver + 7% Copper: Pewter = 85% Tin + 7% Copper + 6% Bismuth +2% Antimony

: All molecules are roughly same size: Interstitial Alloy --- Atoms much smaller then host gets stuck in

between the hosts


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: Steel = Iron + Carbon: Iron is huge, Carbon is tiny


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Solutions and Colligative Properties

: Key words or terms in this chapter: solution, solvent, solute, aqueoussolution, saturated, unsaturated, supersaturated, solubility curve, mass %,molarity, molality, dilutions, solution stoich, colligative properties.

: Most of life takes place in an aqueous Solution: Even water is usually a solution, and not a pure substance: Solution --- a homogeneous mixture

: samples from a solution are always identical: Solvent --- The thing being dissolved into

: Solute --- The thing being dissolved: Aqueous Solution --- a solution using water as the solvent

: Solubility: What does it mean when something dissolves?: Solubility of Ionic Solids

: The Ions are surrounded by water: And according to the charge of the ion

: NaCl(aq) Na+

(aq) + Cl---

(aq): The solid is broken into cations and anions

: Solubility of Polar Molecules: Water can dissolve non-ionic substances: If the substance has ANY polar bonds in its structure, it can be

broken up by water: Alcohols with Ice in them

: Insoluble substances: Petroleums and Oils: Both are relatively non-polar

: Contain mostly C---H bonds which are not pulled on by

water: How a substance dissolves: Fits into the holes between the waters: Like dissolves Like!

: Solution Composition: For every solvent, there is a given amount of solute that will fit


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into the sample at a given Temp. and Pressure: Saturated --- contains as much solute as possible: Supersaturated --- contains more solute than possible: Unsaturated --- contains less solute than possible

: all the above examples at X Temp. and Y Pressure!

: Mass Percent

Mass percent = Mass of Solute X 100Mass of Solution

: Keep in mind that Solution, by definition, has the solute in there!

: Be able to both calculate the Mass Percent, or the Mass of Solute used.

: Plug and Chug for the first version, multiply the %age by the gramssolution and be done.

1. A 135 gram sample of sea water is dissolved to leave only 4.73 gramsof solid residue. What was the mass percent of the sea water?

2. Cow milk is typically 4.5% lactose. Calculate the mass of lactose in175 grams of milk?

: Molarity

M = Molarity = Moles Solute__Liters of Solution

: Be able to solve for Molarity, Ion Concentration, or Moles in a MolarSolution.

: Could also go to grams then from Moles.

1. Calculate the Molarity of a solution prepared by dissolving 11.5grams of NaOH in enough water to make 1.50 L of Solution.

2. Calculate the concentration in Molarity of the Ions in a 1M FeCl3Solution.

3. How many moles of Ag+ ions are present in 25ml of a 0.75 M AgNO3Solution?

: Molalitym = molality = Moles solute

Kg Solvent


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: Be able to solve for molality, moles solute, grams solute.

: Dilutions

M1V1 = M2V2

: Simply plug and chug into this formula making sure you are in theappropriate units.

: When you are diluting, your number of moles will stay constant, but yourwater will not.

: The thing being diluted is called a Standard Solution because we knowits concentration.

1. Suppose I want to prepare 500.0 ml of 1.00 M acetic acid from a 17.5M standard solution. What volume of the standard solution and watermust be used?

If you are using an acid or a base that contains more than one H+ orOH- respectively, you can add a quality factor to the M1V1 = M2V2formula. Lets call that new formula Q1M1V1 = Q2M2V2 where Q is =to the number of ions we get. Example: H2SO4 has a Q of 2 since itproduces 2 H+ in solution.

: Solution Stoichiometry

Use the steps in the book.1. Write the balanced equation.2. Calculate moles of reactants.3. Determining Limiting if there is one.4. Calculate the moles of other reactants of products as asked for.5. Convert to units asked.

: Neutralization Reactions

: When we add Acids to Bases in identical strengths.


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: Colligative Properties

Why do we put salt on the roads in the winter? Its not because the saltmelts the ice in case you were thinking that.

: Freezing point depression and boiling point elevation.: Based off of the molality of the solution that you form, these pointswill change respectively.

Moles solute = m x f.p.d. = f.p. + normal f.p. = new f.p.Kg Solvent b.p.e. b.p. normal b.p. new b.p.

: There are many starting points here and a ton of information that

will be given depending on the starting point. Need to be able tosolve for change in f.p. or b.p., new f.p. or b.p., backwards for m,or for any part of m based on a solved m.

: Try them all!


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Acid and Base Chemistry

: Key words or terms in this chapter: Arrhenius, Bronsted-Lowry,hydronium, hydroxide, strong acid, weak acid, diprotic acid, amphoteric,pH, pOH, indicator.

: 80 Billion pounds of H2SO4 in the US each year: Used in fertilizers, detergents, plastics, batteries,pharmaceuticals, and the making of some metals

: Acids and bases

: Acids were first recognized as substances thattaste sour, and feel dry: Think lemon juice on your tongue

: Bases were first recognized as substances thattaste bitter, and feel slippery

: Think soap: Arrhenius Models

: First person to recognize essential nature of acids and bases: Used electrolyte theory as definition

: Acids = produce H+ ions in solution

: Bases = produce OH--- ions in solution: BronstedLowry Models

: Defined Acids = proton donorBases = proton acceptor

: According to this theory, when an acid like HA is added to H2O as abase, a proton is donated to theH2O creating a new acid and a conjugate base.

HA (aq) + H2O (l) H3O+

(aq) + A-(aq)

Acid Base C. Acid C. Base

: This model is very complete because it recognizes the importance of polarH2O pulling the proton from the acid in solution

: H3O+

(aq) is called the Hydronium Ion

: Acid Strength


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: We know that acids dissociate in water and that H+ is formed: This forms the Conjugate acid and base: This reaction is a reversible one

: Forward = HA(aq)

+ H2O

(l) H

3O+

(aq)+ A-

(aq)

: Reverse = H3O+

(aq) + A-(aq) HA (aq) + H2O (l)

: Can just as easily write with double arrows!!!

: There is an inherent competition for the H+ ion between A--- and H2O.Both are bases, and so both can accept protons.

: Strong Acid = If H2O is more attractive than A- for the H+, then the

forward reaction rules and the H3O+ and A- are present in greatest

concentration in solution: HA has been completely ionized

: Weak Acid = If A-

is more attractive than H2O for the H+

, the reversereaction dominates and there is more HA and H2O in solution

: HA has NOT been completely ionized: Can use different arrow sizes to show this!

: Diprotic Acids --- Release two protons: H2SO4 is a great example: Must dissociate in two steps

: 1. H2SO4 (aq) + H2O (l) H3O+(aq) + HSO4

-(aq)

2. HSO4-

(aq)+ H

2O

(l)H

3O+

(aq)+ SO

4

-2(aq)

: Notice that the second step involves a reversible reaction: Generic Types of acids

: Oxyacids = Involve oxygen as a major piece: Organic Acids = Involve a Carbon backbone

: Water as an Acid and Base: Amphoteric Substances --- can act as an acid and a base: Water is one of the most important amphoteric substances on earth: Lets see what happens when you dissociate water

: H2O(l) + H2O(l) H3O+(aq) + OH-(aq): One water is proton donating, the other is proton accepting since thereverse is the more powerful, only very small amounts ofH3O

+(aq) + OH

-(aq) exist in pure water

: At 25oC, [H3O+ ] = 1.0 x 107 and [OH- ] = 1.0 x 10-7

: SO, [H3O+ ] [OH- ] = 1.0 x 10 -14 = Kw

: If one goes down, the other must go up and vice versa


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: SO,: Neutral Solution = [H3O

+ ] = [OH- ]: Acidic Solution = [H3O

+ ] > [OH- ]: Basic Solution = [H3O

+ ] < [OH- ]: Remember, in all cases above, [H

3O+ ] [OH- ] = 1.0 x 10 -14

: The pH Scale: Based on the common Logarithm (base 10 log): We use this to describe the H+ ion concentration

: pH = -log [H+]: Calculate the pH for solutions

[H+] = 1.0 x 10-9[OH--] = 1.0 x 10-6

: What does all this mean?

: pH 4 is 10 times stronger for H+

than pH 5: pH 3 is 100 times stronger for H+ than pH 5

: pOH = - log [OH-]: Lets look at some numbers

: [H3O+ ] [OH--] = 1.0 x 10 -14

: -log [H3O+ ] [OH-] = -log (1.0 x 10 -14)

: -log [H3O+ ] -log [OH- ] = -log (1.0 x 10 -14) = 14

: pH + pOH = 14: Calculate pH, poH, [H3O

+ ], [OH- ] for a solution: Can include decimals

: pH = 7.41, Then [H3O+ ] = 1.0 x 10-7.41 = 3.9 x 10-8

:Testing for pH using indicators, litmus, and pH meters


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Equilibrium

: Key words or terms in this chapter: collision model, chemical equilibrium,homogeneous equilibrium, heterogeneous equilibrium, Q, Keq, activationenergy, dynamic condition, law of mass action, LeChateliers Principle, Ksp.

: Chemistry is about reactions and reorganization of atoms: We should be able to balance reactions, calculate reactants and

products

: How reactions occur

: Reactants on left, products on right, arrow in between: How do the reactants learn where to go?: Reactants change due to collisions

: Sometimes violent, sometimes casual: Examine - 2 BrNO(g) 2 NO(g) + Br2(g)

: 2 molecules collide, 3 fall out: We call this The Collision Model

: Explains why raising the concentration of the reactants raisesthe reaction rate

: More collisions means faster reaction

: Also works for temperature raising

: Conditions that effect reaction rate: Activation Energy = minimal amount of energy needed for a reactionto occur

: If there is < A.E., the reaction does not occur: Dont want to depend on raising the temperature or collisions, justadd a catalyst!

: Catalyst = a substance that speeds up the reaction without being

consumed: Enzymes = biological catalysts: How fast does a catalyst make things?

Cl + O3 ClO + O2+ O + ClO Cl + O2

Cl + O3 + O + ClO ClO + O2 + Cl + O2


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: Take out spectator species on both sides

O + O3 2 O2

: Notice Cl is both consumed and produced!: One Cl atom can destroy almost 1,000,000 Ozone molecules/second

: Heterogeneous Reactions: Reactions involving one type of phase = homogenous reaction: So, heterogeneous = involve more than one phase

Zn(s) + 2 HCl(aq) H2(g) + ZnCl2(aq): Large pieces of Zinc will slow the reaction rate

: Small grains will increase the rate

: Surface area here is huge!!!: SO, nature of reactants, concentration, pressure, temperature,

surface area

: The Equilibrium Condition: Balance or steadiness is imperative: Too many forces, but they must all balance out

: Chemical Equilibrium = exact balancing of two processes, one of which isthe opposite of the other

: Look at a closed container of water: Equilibrium is when rate of condensation =rate of evaporation: We often think of a reaction going to completion, or of at least onereactant going to completion

: Many reactions in closed containers will not go to this point: Instead, they will get to a point when the [ ] of all reactants andproducts is =, this is Equilibrium

Look at:NO2(g) + NO2(g) N2O4(g)Red-Brown Colorless

: Initial dark red-brown will decrease: Contents will not become colorless!: Contents will lighten and then stop: This does not mean that the reaction is stopped: Both reactants and products are being formed


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: 2 NO2(g) N2O4(g)

: Synthesis and Decomposition at the same time.: This is called a Dynamic Condition

: Chemical Equilibrium as a Dynamic Condition: The idea that a reaction at equilibrium has stopped is way off track: Like 2 islands with a bridge between them.

: Start cars driving across: Movement is in both directions, in even amounts

: No net movement: Consider :

H2O(g) + CO(g) H2(g) + CO2(g): Lets assume that you have = moles of each

: Immediately they should begin forming products: So the concentration of the reactants goes down: Now the reverse reaction is allowed to kick in

: So the concentration of the products goes up, then starts going downagain technically

: Eventually the rate of the forward reaction =the rate of the reverse reaction!: Draw the graph for this idea!!!!!!

: The Equilibrium Constant: The Law of Mass Action was proposed in 1864

: Would come to be called The Law of Equilibrium: Proposed that for a reaction;

a A + b B c C + d D

: Where A,B,C,D are species and a,b,c,d are their balancedcoefficients, you can write an expression;

K = [C]c[D]d

[A]a[B]b

K = Equilibrium Constant

: So Consider the reaction;

2 O3(g) 3 O2(g)


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: What is the Equilibrium Expression for this reaction?: Lets look at data for the reaction;

N2(g)

+ 3 H2(g)

2 NH3(g)

: No matter the original concentrations,there will eventually be a time when the K = expression

: And this will always = 0.0602 = K

: Heterogeneous Equilibrium: In the above examples, all species were gases: Now lets look at;

CaCO3(s) CaO(s) + CO2(g)

: A straight interpretation of this law would be.

K = [CO2] [ CaO][CaCO3]

: The problem is, that heterogeneous equilibrium does not depend onamounts of pure solids or liquids

: Reason = concentrations of pure solids or liquids can not change: These concentrations can be viewed as constant

: SO, K = [CO2] C1C

2

: SO, K = [CO2]

: What about the reaction; 2 H2O(l) 2 H2(g) + O2(g): K = ?: But what about if the water exists as a vapor?

: Then K = ?

: LeChateliers Principle: Chemists often want to maximize a species in a rxn

: Usually a $ product: Used extensively when creating ammonium

: We can therefore determine the effects of change in;: Change in temperature: Change in concentration ( pressure ): Changes in pressure


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: LeChateliers Principle = when a system in equilibrium is imposedupon by a change. That system tends to move to reduce the stress and re-establish equilibrium

: Effects of a change in concentration: Consider ammonia synthesis

: N2(g) + 3 H2(g) 2 NH3(g): At equilibrium we see [N2] = 0.399M

[H2] = 1.197M[NH3] = 0.203M

: What if we add 1.00 mol/L N2(g) into system?: There will suddenly be more collisions betweenN2 and H2 in the system

: You get a much larger forward reaction rate

: SO, equilibrium had shifted to the right: Original New[N2] = 0.399M [N2] = 1.348M[H2] = 1.197M [H2] = 1.044M[NH3] = 0.203M [NH3] = 0.304M

: The value of K has not changed!!!!!!!!!: Rule = When a reactant or a product is added, the

the system shifts AWAY from the added species

: What about:As4O6(s) + 6 C(s) As4(g) + 6 CO(g)

: CO is added; shift _______________: C or As4O6 is added; shift __________________: Removal of As4

: The effect of a change in volume: When the volume of a gas decreases, the pressure increases

according to Boyles Law: Molecules in a smaller space have more collisions

: Rule = When the volume of a gaseous reaction system in equilibrium issuddenly reduced, leading to a sudden increase in pressure, the system willshift in a manner that reduces the pressure

: Consider; CaCO3(s) CaO(s) + CO2(g)


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: As pressure is increased the system will shift left, reducing the number ofgas molecules, and therefore the pressure

: In this process, CO2 molecules are used up: Now consider; N

2(g)+ 3 H

2(g) 2 NH

3(g)

: Add pressure, shift right where less gas molecules exist: Look at the following;

PCl3(g) + Cl2(g) PCl5(g)

PCl3(g) + 3 NH3(g) P(NH2)3(g) + 3 HCl(g)

: The effects of change in temperature: So far, factors we have mentioned only effect the position, not the K

: Values of K will change with a temperature change: We will work to predict the change in K, and the shift

: We must decide between exo- and endothermic: Produce heat = exothermic: Consume heat = endothermic

: Consider the following: N2(g) + 3 H2(g) 2 NH3(g) + 92 Kj

: CaCO3(s) + 556 Kj CaO(s) + CO2(s)

: The system will move to consume the heat: Treat it like a reactant or product

: What if the reaction only says exo- or endothermic

:Applications of the Equilibrium Constant: NOW, we can analyze some reactions

: K can tell quite a bit about the reactions tendencies: Value of K larger than 1 means that at equilibrium, the system will have

mostly products!

: Think about it?: Consider the reaction;

A(g) B(g)K = [B]

[A]


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[B] = 10,000 or 10,000 SO, B is strongly favored[A] 1

: On the other hand, if K has a very small value, then A would be highlyfavored in the reaction

: Given; K = 4.34 for 2 SO2(g) + O2(g) 2 SO3(g)[ SO3] = 0.260M[ SO2] = 0.590M[ O2] = ?

Answer = 0.045M

: Ksp = Solubility Product Constant: Solubility is an important concept

: A great deal can be accomplished using ourknowledge of solubility

: We can flavor foods using salt or sugar: We can use Ba+2 to coat your insides for an X-Ray: We can consider the equilibrium for when a solid dissolves in water toform ions

: Look at reaction; CaF2(s) Ca+2

(aq) + 2 F---

(aq): When the reaction starts there are no ions present: Eventually there could be a reverse reaction: Ultimately equilibrium will be reached between the two processes

: No more solid dissolves, and the solution saturates: Ksp = [Ca

+2] [ F-]2: using concentrations in Moles/liter

: Remember, the concentration of the solid does not change!, but theseare aqueous ions!

: CuBr has a measured solubility of 2.0 x 10-4?: Ksp = [Cu+] [ Br-]

: SO, [2.0 x 10-4] [2.0 x 10-4] = 4.0 x 10-8: Solve for some of these problems

: Basic Formulas to think about for Ksp Problems.: Ksp can = X2, 4X3, 9X4, 108X5

: In all cases solve for X, then multiply by the coefficient of thation to find solubility for that ion.


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Oxidation-Reduction Reactions

: Key words or terms in this chapter: oxidation states, reaction method.

: We studied these very generically when we createdIonic Compounds and then defined Reaction Types

: Uses a metal + a non-metal

: Oxidation States: These numbers help to keep track of electrons: Looks at charges of individual atoms within a compound or

molecule: Sometimes these charges are evident, other times not: Binary Ionics are easy

: NaCl Na = +1 Cl = -1: MgO Mg = + 2 O = - 2

: Here, the charges of the ions = oxidation numbers: In uncombined elements, atoms are uncharged or 0: The atom in any pure element has oxidation state 0: In a covalent compound, the more electronegative element has itsnormal anion charge

: H2O H = +1 O = -2: NO2 N =

+4 O = -2: Na2CO3 Na =

+1 C = +4 O = -2: NH4Cl N =

---3 H = +1 Cl = -1

: What about if the species has a charge?: NO3

- N = +5 O = -2: MnO4

- Mn = +7 O = -2

: Oxidation Reduction Reactions: Reaction between metal and a non-metal: Look at how Sodium Chloride forms

2 Na(s) + Cl2(g) 2 NaCl(s)0 0 +1 -1

: Oxidation = loss of electrons


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: Reduction = gain of electrons: Some of the most important redox reactions involve only non-metals.

How is this possible????????

CH4(g)

+ 2 O2(g) CO

2(g)+ 2 H

2O

(g)+ Energy

-4 +1 0 +42 +1 -2: This is a sweet Combustion Reaction

: Redox with only Non-Metals: Take a look above!: What was oxidized and what was reduced?

: Balancing using the Reaction Method in Acid.Step 1 = Assign all Oxidation States and Separate Reactions

Step 2 = Balance all elements except H and O.Step 3 = Balance O using H2OStep 4 = Balance H using H+1Step 5 = Balance Charges using e-1Step 6 = Multiply reactions to make e-1 gained = e-1 lostStep 7 = Combine and cancel!


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Organic Chemistry

: Key words or terms in this chapter: alkanes, alkenes, alkynes, saturated,unsaturated, aromatic hydrocarbons, functional groups, isomers.

: Organic Chemistry: The study of carbon containing compounds and theirproperties.

: Carbon and Silicon form most natural substances: Silicon forms inorganics like rocks/minerals

Carbon Bonding

: Strong self bonding properties: CH4 is a solid example: Tetrahedral in nature

: Dont forget double and triple bonds: CO and CO2

: Also C2H4 = Ethylene: Also C2H2 = Acetylene

Alkanes: Hydrocarbons = compounds composed of C and H

: Those with all C-C single bonds = saturated: Those with C-C multiple bonds = unsaturated

: C2H4 + H2 C2H6Unsaturated Saturated

: Methane = CH4: Ethane = C2H6: Propane = C3H8: Butane = C4H10

: Propane and Butane have 109.5o CC bond

angles: Basic form is (CnH2n+2): These are all straight chained or unbranched

Structural Formulas and Isomerism: Isomerism = When two molecules have the sameatoms, but different bonds


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: n-butane versus isobutane: n-butane has a four carbon chain: isobutane has a three carbon chain with a

single carbon side chain: Draw five isomers for C

7H

16

Naming Alakanes: There must be a system: Otherwise you should just give up!

: Rule 1= Methane, Ethane, Propane, Butane; then usesGreek roots with (ane) added on till 10 carbons

: Rule 2 = You always use the longest continuous carbon chain to name analkane

: Rule 3 = You can substitute something else for an H

: These are named according to what they look like: ---CH3 is a methyl group, it looks like methane!

: ---CH2CH3 is an ethyl group

I: ---CH3CHCH3 is an isopropyl group

: ---CH2CH2CH2CH3 is a butyl group: also look at sec-butyl, isobutyl, and tert-butyl

: Rule 4 = We specify the substituent groups by numbering the carbonchains so that the group has the lowest possible number

: CH3I

CH3CH2CHCH2CH2CH31 2 3 4 5 6 Correct Numbering6 5 4 3 2 1 Incorrect Numbering

: This is called 3methylhexane: Note the dash between the number and name

: Rule 5 = When more than one group is involved, we use a prefix such asdior trito show this.

: CH3CH----CHCH2CH3I ICH3 CH3

: This is called 2,3dimethylpentane


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Petroleum: Petroleum = thick, black liquid, composed ofhydrocarbons from 5 to >25 carbons

: Natural Gas = mostly methane with some ethane,propane, and butane

: Forms according to chain length: C5---C12 = gasoline: C10---C18 = kerosene, jet fuel: C15---C25 = diesel, heating oil, lubricants: >C25 = Asphalt

Reactions of Alkanes: At lower temps., alkanes are not very reactive.

: CC and CH bonds are stable

: They do react with O2 very well in combustion!: 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(g)

: These also undergo substitution reactions: One or more of the Hs are replaced: This can be done many times over!

: CH4 + Cl2CH3Cl + HClChloromethane

: CH3Cl + Cl2CH2Cl2 + HClDichloromethane

: CH2Cl

2+ Cl

2CHCl

3+ HCl

Trichloromethane: CHCl3 + CL2CCl4 + HCl

TetrachloromethaneCarbon Tetrachloride

: Dehydrogenation reactions: When Hs are removed and the product is unsaturated: CH3CH3 CH2=CH2 + H2

: Uses Cr2O3 and 500oC to do this!

Alkenes and Alkynes: Alkenes = Hydrocarbons that contain doublebonds

: Alkynes = Hydrocarbons that contain triplebonds

: Rules for naming alkenes and alkynes: Rule 1= Select the longest continuous chain that


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contains the double or the triple bond: Rule 2 = For the alkene use naming like alkanes

but useene, for alkynes, use ---yne.: CH3CH3 = Ethane: CH

2CH

2= Ethene

: CHCH = Ethyne: Rule 3 = Number the parent chain so that the double or triple bond

gets the lowest number start point: CH2=CHCH2CH3 is called 1-butene: CH3CH=CHCH3 is called 2- butene

: Rule 4 = Sustituents on the parent chain are treated the same waythey were in alkanes

: ClI

CH=CHCH2CH3 is called 1-chloro-1-butene: Reactions using Alkenes

: Addition reactions = new atoms form single bonds to carbonspreviously unsaturated

: Example is Hydrogenation: CH2=CHCH3 + H2 CH3CH2CH3

1-propene propane

: Halogenation = addition of a halogen to an unsaturated hydrocarbon: CH

2=CHCH

2CH

2CH

3+ Br

2 CH

2BrCHBrCH

2CH

2CH

3

1Pentene 1,2dibromopentane

: Polymerization = small molecules are joined to form larger ones.

Aromatic Hydrocarbons: Some hydrocarbons smell good.: We refer to these as aromatic hydrocarbons.

: Cinnamon, wintergreen, vanillin: Most contain Benzene, C6H6 in a ring form.

Naming Aromatic Compounds.: Replaces one or more of the Hs with something

: Monosubstituted;: We simply add the name of the substituent group onto the

name benzene.: Some common name exceptions.


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: Toluene, Phenol, Styrene: If you would rather treat the benzene as the substituent, call it

a phenyl group.: Disubstituted:

: If more than one substitute, numbers are used to tell wherethey are on the ring

: Can use prefixes if need be: Ortho- = 2 adjacent substitutes: Para- = 2 directly across from oneanother

: Meta- = 2 substitutes, with one carbonbetween them

: SO, 1,2-dichlorobenzene = o-dichlorobenzene1,3-dichlorobenzene = m-dichlorobenzene

1,4-dichlorobenzene =p-dichlorobenzene: Benzenes with 2 methyl groups are called xylene

: use this when dealing with examples: 2 different substitutes

: One is assumed to be in the 1 position, the other is numbered.

Functional Groups: Common functional groups act to define certain types of organic

molecules: Name Functional Group Formula: Halohydrocarbons -- X RX: Alcohols --OH ROH: Ethers --O-- ROR

O OII II

: Aldehydes --CH RC--HO OII II

: Ketones --C RCR

O OII II

: Carboxylic Acids --COH RC--OHO OII II

: Esters --CO-- RCOR: Amines --NH2 RNH2


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: Oxidation of the primary alcohol gives a correspondingaldehyde

CH3CH2OH CH3C=OI

O: Oxidation of a secondary alcohol gives a corresponding

ketoneCH3CHCH3 CH3CCH3

I IIOH O

Naming Aldehydes and Ketones: We use the parent chain alkane

: For Aldehydes we remove thee and addal

: For Ketones we remove thee and addone: In the ketone, the C in the carbonyl gets the lowest number possible: In the aldehyde, the C in the carbonyl group is always at the one endand is therefore assumed to be the carbon 1

: Look out for Benzaldehyde

Carboxylic Acids and Esters: Carboxylic Acids are shown by the carboxyl group

: -COOH: These are typically weak acids: Named by dropping thee from the parent alkane and addingoic: Note that theCOOH is always given the 1spot on the chain

: A carboxylic acid reacts with an alcohol to form an ester and a watermolecule

: RCOOH + ROHRCOOR + H2O: Esters have the following generic formula

: RCORII

O

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