ࡱ> l@ >bjbj (&uuO# """6<6C2^Ȩ!!!BBBBBBB$/ERGB"$ !$$BȨcC111$Pz "ȨB1$B11?h "@Ȩ 왔E'F@IH"@P!"1K##&!!!BB66ڢ166ڢHONORS CHEMISTRY- IMPORTANT TERMS, RULES, AND FORMULAS 2006-2007 SCIENTIFIC MEASUREMENTS Significant Figures Addition/Subtraction Rule- When adding/subtracting, the answer should be rounded to the same number of decimal places as the measurement with the least precision (number of decimal places). Multiplication/Division Rule- Round the answer to the number of significant figures in the measurement with the fewest significant figures. A precaution for those planning to study from this: The summaries are only representative of my notebook, so some of them may not be complete study guides- for example, hybridization isn't in any of the summaries because there was a powerpoint on it that I printed. CHEMISTRY Atom- The smallest part of an element that retains its properties -smallest building block of matter -116 different kinds  Element- Only contains one kind of atom   Molecule- Two or more atoms chemically linked to behave as a unit Compound- A type of matter containing molecules composed of two or more different kinds of atoms Pure Substance- Matter containing only one type of particle (atoms or molecules) ( elements OR compounds -can be described using a single formula -unvarying composition Mixture- Contains two or more pure substances that are physically mixed together ~ components can be separated; composition varies Homogenous Mixture- Components are uniformly distributed throughout -ratio of components is constant -synonym = solution ~ solutions dont have to be liquid (ex: air) -synonym = alloy ~ solution of metals in solid phase (ex: brass, 14 K gold)    Heterogeneous Mixture- Components do not have uniform distribution ~ distinct phases or layers  Solute- The material being dissolved ~ smaller proportion of total Solvent- The dissolving material ~ greater proportion of total Organic Substance- Compounds containing the element carbon Ex: CH4 (methane) or C2H8 (ethane) Inorganic Substance- The compounds of all the other elements Physical Property- Observable without changing identity Ex: viscosity Extensive Physical Property- Depends on amount Ex: mass, volume, area Intensive Physical Property- Does not depend on amount; are unique to identity of substance Ex: boiling point, melting point, density, malleability, ductility Chemical Property- Result in a change in composition/identity Ex: flammability Specific Heat (C or Cp[ressure])- A measure of a substances ability to store heat; the amount of energy needed to raise the temperature of one gram of a substance by 1 degree Celsius -metals = low Cp values ( heat and cool quickly -water = high specific heat ( heats and cools slowly   ATOMS Cation- A positively charged ion (loss of e-) Anion- A negatively charged ion (gain of e-) Wavelength- The distance between two successive peaks ~ symbol =  Frequency- the number of wave crests that pass a certain point in a set amount of time -units = s-1, 1/sec, H2 (cycles per second) -symbol =  (lowercase nu) Wavelength and frequency are inversely related: - Long , small  - Short , large  Energy and wavelength are inversely related:   Energy and frequency are directly proportional:  Orbital- Region of space where electrons are likely to be found *review quantum number rules* Electron Configuration Rules 1) The Aufbau Principle- Fill in the lowest energy orbitals before higher energy orbitals Ex: H EMBED Equation.3  2) Pauli Exclusion Principle- No two electrons in the same atom can have all four quantum numbers be identical ~ any single orbital can hold a maximum of two electrons, but they must have opposite spins Ex:  He (2e-):  EMBED Equation.3   Li (3e-):  EMBED Equation.3  ~ spin doesnt matter on second s orbital since there is only one e- B (5e-):  EMBED Equation.3  ~ spin/position dont matter in 2p b/c they all have same energy (degenerate) 3) Hunds Rule- For p, d, and f orbitals: put one electron in each orbital, all spin aligned before pairing any electrons Ex: N (7e-):  EMBED Equation.3  Noble Gas Notation- Write the nearest noble gas with lower electron number, then whatever is left over Ex: Mo: [Kr] 5s2 4d4 CHEMICAL PERIODICITY Periodic Law For s and p block elements- Number of valence electrons = group number OR group number 10; varies from 1-8 For d block elements- Usually have two valence electrons Groups- Vertical columns on periodic table ~ elements in the same group have similar properties Periods- Horizontal rows 1-7 on the periodic table Metal(metalloid(nonmetal(gas Atomic Radius- One half the distance between the nuclei of two like atoms * Going down a group, atomic radius increases * Going left to right in a period, atomic radius decreases Zeff- Effective nuclear charge~ more protons as you go L(R Shielding- Core electrons effectively block nuclear charge from reaching outermost electrons -shielding is constant across a period -nuclear charge holds the outermost electrons more tightly as you go L(R -Zeff increases L(R -Z increases, constant shielding Ionization Energy- The amount of energy needed to remove an electron from an atom in its group state( the measure of how difficult it is to lose an electron~ endothermic in relation to atom -as you go L(R in a period, IE increases b/c nonmetals dont want to lose electrons(outermost electrons are held very tightly due to high Zeff -metals have a low IE -non-metals have high IE -second IE is usually much larger than 1st( atoms/ions dont want to disrupt noble-gas-like configuration Octet Rule- An atom with eight electrons in the outermost energy level is unreactive~ atoms will gain or lose electrons to achieve this stable octet (this is why they ionize) Electronegativity- The tendency of an atom (in a compound) to attract electrons to itself -as you go down a column, electronegativity tends to decrease -as you go from L(R, electronegativity tends to increase Electron Affinity- The attraction of an isolated (or gaseous) atom for an electron -metals typically have a low EA -nonmetals typically have a high EA -electron affinity decreases down group (decreasing attraction for e-) -electron affinity increases left to right across a period (greater attraction for e-) -(-) EA = electrons are repelled ( wants to push additional e- away -(+) EA = wants to gain e- NUCLEAR CHEMISTRY Nuclear Fission-  Nuclear Fusion-  Half Life- The amount of time it takes for 50% of the radioactive nuclides in a sample to decay *The combined mass of the protons and neutrons in the nucleus of real atoms is always less than the sum of the masses of the individual particles  COMPOUNDS Molecular Formula- Gives the actual numbers of each kind of atom in a molecule Empirical Formula- Gives the lowest whole number ration of each kind of atom in a molecule~ simplest formula Naming Rules Monatomic Ions- Cation = element name + ion; Anion = element name with ending changed to ide Transition Metals- can form multiple cations( ion name (element name) + Roman numeral in parentheses to indicate charge~ exception: silver only forms +1 ion, zinc only forms +2 ion( dont need to state charge Binary Ionic Formulas- (1) Cation then anion (2) no net charge (sum of charges must = 0) Formulas with Polyatomics- (1) Cation then anion (2) sum of the charges = 0 (3) use parentheses for multiple copies of a polyatomic ion Ionic Compounds- Cation name then anion name -include roman numerals as needed -no information on number of atoms Covalent Compounds- Use prefixes to indicate number of atoms CHEMICAL QUANTITIES    Second Definition of the Mole- The mass of a substance containing Avogadros number of particles Gram Formula Mass (gfm)- The mass, in grams, of one mole of a substance~ also called molecular weight, formula weight, molar mass Hydrates- Some compounds form complexes with water, such that the water molecules surround the formula unit Ex: CuSO4 5H2O name = copper (II) sulfate pentahydrate Ex: Zn(C2H 3O 2) 2 2H 2O name = zinc acetate dihydrate Percent Composition-   CHEMICAL REACTIONS Assigning Oxidation States- 1) Free (uncombined) elements have an oxidation number of zero 2) For monatomic ions, the charge of the ion is the oxidation number 3) Hydrogen in a compound has an oxidation state of one unless its combined with a metal (then it is -1) 4) Fluorine in a compound is always -1 5) Oxygen in a compound has an oxidation number of -2 unless it is combined with Fluorine (then it is +2) ( or if it is in a peroxide (then it is -1) 6) Sum of all oxidation numbers in a compound is zero 7) Sum of all oxidation numbers in a polyatomic ion is the same as the charge on the ion Six Classes of Reactions- 1) Combustion 2) Synthesis (aka combination) 3) Single Replacement 4) Double Replacement (aka double displacement) 5) Decomposition 6) Redox Key Classes of Organic Reactions- 1) Substitution 2) Addition 3) Esterification (dehydration synthesis) 4) Polymerization STOICHIOMETRY Theoretical Yield- The amount of product expected to be produced in a reaction based on a stoichiometry calculation Actual Yield- The amount of product experimentally formed in a reaction *The actual yield is often less than the theoretical yield Percent Yield- A measure of reaction efficiency  Limiting Reagent- Runs out( limits the amount of product that can be formed Excess Reagent- Left over at the end of the reaction CHEMICAL BONDS *Forming bonds gives off energy( always exothermic Three Main Classes of Bonds- 1) Ionic: EN > or = 1.7 2) Covalent: Non polar covalent: EN < or = 0.4; polar covalent: EN from .41-1.69 3) Metallic Resonance Structures- The multiple Lewis Structures that can be drawn for certain compounds that are not adequately represented by one Lewis Structure Bond Length- Average distance between two bonded atoms~ distance between nuclei at their minimum potential energy Bond Dissociation Energy (aka bond energy) - Energy required to break a bond and form neutral isolated atoms Formal Charge- Some molecules dont obey the octet rule -calculated for each atom -sum of FC must equal overall charge of species -the most appropriate Lewis Structures have the lowest FC possible( zero is best -FC is negative on most electronegative elements  VSPER Theory- Valence Shell Electron Pair Repulsion *review molecular geometries* Molecular Orbital- Can hold a maximum of two electrons, has a definite energy, and can be represented with an electron density cloud Linear Combination of Atomic Orbitals (LCAO)- Whenever two atomic orbitals overlap, two molecular orbitals form -# in = # out -energy is conserved( one orbital will be lower in energy, one will be higher in energy Bond Order- In MO Theory, bond stability of a covalent bond is related to its bond order  Intermolecular Attraction- Also called van derWaals forces or weak forces( generally weak -ion dipole attractions -dipole-dipole attractions -hydrogen bonding -dipole induced dipole attractions GASES Kinetic Molecular Theory of Gases -particles move non-stop in straight lines -particles have negligible volume (like points in geometry) -particles have no attraction to each other and no repulsion -particles exert pressure on the container by colliding with container walls -collisions between particles are elastic (no gain or loss of energy) Kelvin-Celsius Conversions- Kinetic Energy-  Ideal Gas- Fully obeys all statements of kinetic molecular theory~ most likely found at low pressure/high temperature Real Gas- Doesnt obey one or more parts of KMT~ most likely found at high pressure, low temperature, and where there are intermolecular attractions  Units of Pressure- 1 atmosphere (atm) = 101.3 kPa 760 mm mercury (Hg) 760 torr 14.7 psi Combined Gas Law- Temperature must be K Daltons Law of Partial Pressures- For mixtures of gases, the total pressure exerted by the mixture is equal to the sum of the pressures exerted by each individual gas Ideal Gas Law- Universal Gas Constant (R)- atm:  kPA: Diffusion- The gradual mixing of two gases due to random spontaneous motion Effusion- When molecules of a confined gas escape through a tiny opening in a container Grahams Law- At the same temperatures, molecules with a smaller gfm travel at a faster speed than molecules with a larger gfm~ as gfm goes up, velocity goes down( The relative rates of diffusion of two gases vary inversely with the square roots of the gram formula masses SOLUTIONS Miscible- Liquids that will dissolve in each other Ex: oil and gasoline for two stroke engines Immiscible- Two liquids that are insoluble in each other Ex: Oil and water Relative Humidity- A measure of how much water vapor is in the air, compared to the maximum amount of water the air can hold at that temperature Supersaturated Solution- A solution prepared with more dissolved solute than a saturated solution Molarity by Dilution-  Rule of Thumb for Solubility- Like dissolves like Molality- Used to calculate changes in physical properties of solutions  Vant Hoff Factor (i)- The theoretical maximum number of particles formed when a substance dissociates -for all covalently bonded substances: i = 1 -for ionic substances: i = number of ions present in formula Colligative Property- A property that depends on the number of solute particles but not their identity~ boiling point elevation; freezing point depression    Heterogeneous mixtures- Solutions: -homogeneous mixture -solute and solvent are evenly distributed throughout -typical particle size < 1 nm -particles are too small to filter -doesnt separate upon standing -Suspensions: -examples = clay in water, muddy water -particles will settle to bottom when undisturbed -particles can be recovered by filtration -particle diameters typically 100-1000 nm -Colloids: -may be milky or cloudy in appearance -dont separate on standing -cant recover particles through filtration -typical particle diameter = 1-100 nm Tyndall Effect- Suspensions and colloids scatter light beams, making them visible~ observed when high beams are used in a heavy fog THERMODYNAMICS   Energy Required for a Phase Change-  Exothermic Reaction- Releases energy to surroundings~ negative H ( products are lower in energy than reactants Endothermic Reaction- Absorbs energy from surroundings~ positive H ( products are higher in energy than reactants Enthalpy- Hrxn Hess Law of Heat Summation- If you add two or more thermochemical equations to give a final equation, then you can also add the heat changes to give the final enthalpy of reaction *The overall enthalpy of a reaction is the same whether the reaction occurs in one step or in several steps Standard Heat of Formation (Hf)- The change in enthalpy that accompanies the formation of one mole of a substance from its elements in their standard states~ the heat of formation of elements in their standard states is arbitrarily set to zero Thermodynamic Stability- A measure of the energy required to decompose a compound~ compounds with large, negative enthalpies of formation are thermodynamically stable  Entropy (S)- A quantitative measure of the degree of disorder in a system -solids have a higher degree of order (low entropy) -liquids have a low degree of order (high entropy) -more particles (moles) = higher entropy -systems tend to proceed to higher disorder (higher S) Gibbs Free Energy- The energy available from the system to do useful work  REACTION RATES For a general reaction (aA + bB ( cC + dD)-   Initial rate- Calculation of average rate for early part of data when plot is nearly linear Collision Theory- Molecules must collide to react( must have correct orientation and enough energy To Increase the Rate of Reaction- 1) Increase temperature 2) Increase concentration 3) Increase surface area Catalyst- Speeds up reaction without being consumed( effectively lowers the activation energy of the reaction Inhibitor- Causes reaction rate to slow down Reaction Mechanism- A step by step description of the steps that occur in a chemical reaction -includes all of the elementary steps that add up to the overall reaction -proposed mechanism must be consistent with experimental rate law Chemical Intermediates- Made in one step, consumed in another step Rate Limiting Step (RLS)- The slow step in a multi-step mechanism -has the greatest effect on the rate of a multi-step mechanism -molecularity of RLS should match the experimentally determined reaction order *The sum of the steps in the mechanism adds up to the overall balanced equation Rate Law- Rate is directly proportional to reactant concentration (in molarity) raised to some power, n   If more than one reactant is involved Reaction order- nOrderConcentration Dependence1FirstAs [A] is doubled, rate doubles2SecondAs [A] is doubled, rate quadruples Equilibrium constant- Constant ratio of product to reactant when reaction is at equilibrium  LeChateliers Principle- When a system at equilibrium is disturbed by the application of a stress, the system will attain a new equilibrium position that minimizes the stress ACIDS AND BASES Arrhenius Theory- Acids produce H+ ions in aqueous solution and bases produce OH- in aqueous solution Bronsted-Lowry Theory- Acids are H+ ion donors or proton donors and bases are H+ acceptors or proton acceptors Lewis Theory- Bases can donate a pair of electrons to form a covalent bond( atoms with LONE PAIRS; and acids can accept a pair of electrons to form a covalent bond Conjugate Acid-Base Pairs- Acids and bases occur in conjunction( differ only by a proton Ex: NH3/NH4+ H2O/OH- Amphiprotic Substances- Substances that can act as an acid and as a base~ also called amphoteric Acid anhydrides- An oxide that reacts with water to form an acid -nonmetal oxides -important component of acid rain Basic anhydrides- An oxide that reacts with water to form a base~ metal oxides   Autoionization of Water- Ionization constants (Ka or Kb)- Used to categorize acid/base strength - small K = low degree of dissociation, weak acid/base -large K = high degree of dissociation, strong acid/base -calculated the same way as Keq Buffer- a system that resists pH change when acids/bases are added( there is an acid to react with any base added and vice versa OXIDATION REDUCTION REACTIONS (REDOX REACTIONS) Reduction- gain of electrons( oxidation number becomes more negative Oxidation- loss of electrons( oxidation number becomes more positive *LoseElectronsOxidation the lion says GainElectronsReduction* Oxidizing agent- the species that gets reduced Reducing agent- the species that gets oxidized Conservation of Charge- number of electrons lost = number of electrons gained Electrode- a strip of metal( anode or cathode *Electrons flow from the anode to the cathode *Oxidation occurs at the anode *Reduction occurs at the cathode *An Ox( anode, oxidation* *Red Cat( reduction, cathode* ALL FORMULAS   Ex: granite; soil Individual atoms Diatomic elements: H2 N2 O2 F2 I2 Br2 Cl2 Q = mCt Q = amount of energy m = mass C = specific heat t = change in temperature C =   E  EMBED Equation.3   is proportional to E =  EMBED Equation.3  E E = h  Planck s constant = 6.6 * 10-34 JS ( Orbital notation Nucleus Nuclei + energy Nuclei + energy ( Nucleus Mass defect = Expected mass  actual mass 1 mole = 6.02*1023 particles Molar volume of a gas at STP = 22.4 L Mole = n =  EMBED Equation.3  % E =  EMBED Equation.3  100 % yield =  EMBED Equation.3  100 FC = # valence e- - (# lone pair e-) - # bonds Pairs of electrons will orient themselves in space to be as far away from each other as possible. 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TS Rate =  EMBED Equation.3  Average rate =  EMBED Equation.3  Rate = k[A]n n = the extent to which the rate depends on concentration [] = molarity Rate = k[A]n [B]m Keq =  EMBED Equation.3  pH = -log [H3O+] [H3O+] = 10-pH pOH = -log [OH-] [OH-] = 10-pOH Kw = [H3O+] [OH-] = 1.0 * 10-4 at 25C equilibrium expression -ignore solids and pure liquids -coefficients become exponents Keq =  EMBED Equation.3  Kw = [H3O+] [OH-] = 1.0 * 10-4 at 25C [OH-] = 10-pOH POH = -log [OH-] [H3O+] = 10-pH pH = -log [H3O+] Rate = k[A]n [B]m Average rate =  EMBED Equation.3  Rate =  EMBED Equation.3  $&(*,68^`bdfhv|~ג~vh]N]jh hmwCJUaJh hmwCJaJ jh|hmwmHsHhmwmHsH jhmwj"h|hmwEHUjhJ hmwCJUVaJh|hmwCJH*aJmHsHh|hmwCJaJmHsHh|hmwCJH*aJh|hmwCJaJhmwjhmwUja hl{hmwEHUjhJ hmwCJUVaJ,.F ` &(x 4gdj$a$gdjgdsgd<1gd|   ~vkh?lhmwCJaJhshmwH*hmwCJH*aJh<1hmwH* h<1hmw hmwH*h<1hmwCJaJmHsHh<1hmwCJH*aJh<1hmwCJaJhmwh hmwCJaJjh hmwCJUaJ!j$hSg3hmwCJEHUaJjhJ hmwCJUVaJ'&*6zrͲxlaVGVjhPhmwCJUaJhPhmwCJaJh;ZhmwCJaJhQQhmw6CJaJhSg3hmw6CJaJ!j/'hSg3hmwCJEHUaJ#j]J hSg3hmwCJUVaJhSg3hmwCJaJjhSg3hmwCJUaJhQQhmwCJH*aJhQQhmwCJaJh hmwCJaJ h?lhmw hmwH*hmwh?lhmwCJH*aJ4Z2pr "<>-.= pgdmwgd| Ugdj "8<ŲynbnSnjhE(KhmwCJUaJhE(KhmwCJH*aJhE(KhmwCJaJhmwCJaJhmwCJH*aJh| UhmwCJH*aJh| UhmwCJaJ!j*-hPhmwCJEHUaJjViJ hmwCJUVaJhmwhPhmwCJaJjhPhmwCJUaJ!j)hPhmwCJEHUaJjiJ hmwCJUVaJ'()*-.01239<=>LMPQTU[_`abdfhijklmpqr}¾ h hmwhmwCJH*aJhmwCJaJh hmwCJH*aJh hmwCJH*aJh hmwCJaJhmwhE(KhmwCJaJjhE(KhmwCJUaJ!jm/hE(KhmwCJEHUaJ#jfJ hE(KhmwCJUVaJ/=>PQ`a)*9:LM\]opgdmw pgdmw  ()*-.489:=HILMOPQRXǶ貧~wlh h@ CJaJ h hmwhmwCJH*aJhmwCJaJh hmwCJH*aJh hmwCJH*aJh hmwCJaJhmw!j1hE(KhmwCJEHUaJ#jfJ hE(KhmwCJUVaJjhE(KhmwCJUaJhE(KhmwCJaJhE(KhmwCJH*aJ*X[\]ijklop{|}䥖v䥖gV!j 6hPhmwCJEHUaJjiJ hmwCJUVaJ!j3hPhmwCJEHUaJjViJ hmwCJUVaJjhPhmwCJUaJhPhmwCJaJhmwCJaJhmwCJH*aJh| UhmwCJH*aJh| UhmwCJaJh hmwCJH*aJhmwh hmwCJaJh hmwCJH*aJ tv>@XZgd<$a$gd<gdmwG = H - 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