1. Pure Substance
    Made of One Substance....For example when you take water, purifiy...the result is a pure substance...not water from a creek...have compounds
  2. Compounds
    Can be broken down into there elements...and have constant composition
  3. Atom
    • basic building block of matter, representing the smallest unit of a chemical element;
    • In turn is composed of subatomic particles, called protons, neutrons, electrons
  4. Elements
    • Make up compounds
    • basic building blocks of matter
    • Cannot be broken down but can identify parts such as protons, neutrons, and electrons.
    • Defined by number of protons and electrons it has
    • Number of neutrons can vary from atom to atom
  5. Dalton's Theory
    • All elements are composed of very small particles called atoms
    • All compounds are composed of atoms of more than one element
    • A given chemical reactions involves only the separation, combination, or rearrangement of atoms; this does not result in the creation or destruction of atoms
  6. Neutron
    • Chargeless particle
    • In the nucleus
    • Isotopes have different numbers of neutrons but same number of protons
  7. Proton
    • In the nucleus
    • Smaller than neutron
    • Positive
    • Mass of one amu (atomic mass unit)
    • Carry same "quantity of charge" as electron but way heavier
  8. Electron
    • Negatively Charged
    • Outside Nucleus Circling it...
    • Almost No Mass
    • Can move around very quickly
    • Exist in regions of space called orbitals
  9. Valence Electrons
    • Electrons farthest from the nucleus;
    • Farther the valence electrons are from the nucleus, the weaker the attractive force of the positively charged nucleus and the more likely the valence electrons are to be influence by other atoms
    • Their activity determine the reactivity of the atom
  10. Ion
    Due to the result of gaining/losing an electron
  11. Nickel-60 2+ cation
    • Nickel is #28
    • E: 28 - 2: 26
    • P: 26
    • N: 60 - 28 = 32
  12. Nucleus
    • Does not move around
    • Contains Neutrons and Protons which both contribute tot he atomic mass of the element
  13. Isotope
    Each possible atom with varying neutrons; the mass given on the periodic table is a average of its isotopes; have different number of neutrons that change the weight of the element itself
  14. Quantum Theory
    • Energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta;
    • E = hf
    • where h is proportionality constant known as Planck's constant: equal to 6.626 X 10^-34 Js and f sometimes v is the frequency of the radiation
  15. Bohr Model
    • Determined the model of the Hydrogen Atom
    • Can only exist in certain fixed energy states
    • Smaller the radius, lower the energy state of the electron
    • The smallest orbit corresponds to n=1, which is the ground state of the Hydrogen electron

    • But did not prove for atoms containing more than one electrons;
    • Electrons are rather in a state of rapid motion within regions of space around the nucleus called orbitals
  16. Ground State
    Electron is its lowest electron state
  17. Atomic Emission Spectra
    • Electrons can be excited to higher energy levels by heat or other energy to yield the excited state of the atom
    • The lifetime of the excited state is minimum, the electron will return rapidly to the ground state emitting energy in the form of photons
    • E = hc/lambda
    • C= 3.00 X 10^8 m/s
    • Lambda: wavelength
    • Each element possesses a unique atomic emission spectrum which can be used as a fingerprint for the element

    • Lyman 1 2, 3, 4,... Ultraviolet
    • Balmer 2 3, 4, 5,... Visible and ultraviolet
    • Paschen 3 4,5,6, ... Infared
    • Brackett 4 5,6, 7... Infared
  18. Atomic Absorption Spectra
    • When an electron is excited to a higher energy level, it must absorb energy
    • Wavelengths of absorption correspond directly to the wavelengths of emission since the energy difference between levels remains unchanged
  19. Heisenberg Uncertainty Principle
    Impossible to determine, with perfect accuracy, the momentum and the position of the electron simultaneously
  20. Quantum Number
    Any electron in an atom can be described by four quantum numbers: n, l, ml, ms
  21. Pauli Exclusion Principle
    No two electrons in a given atom can possess the same set of four quantum numbers
  22. Energy State
    The position and energy of an electron described by its quantum numbers
  23. n
    • the size
    • Principle Quantum Number
    • Higher this value, higher the energy level and radius of electrons orbit
    • Maximum number of electrons in energy level n (electron shell n) 2n2
    • The difference in enegy between adjacent shells decreases as the distance from the nucleus increases
  24. l
    • about the shape
    • Azimuthal Quantum Number
    • refers to subshells or sublevels that occur within each principal energy level
    • The four subshells= 0, 1, 2, 3 and knowns as s, p, d, and f
    • The maximum number electrons that can exist is by the equation 4l + 2
    • Greater the value of L, the greater the energy of the subshell
    • Energies from different principle energy levels may overlap
  25. ml
    • about the orientation of the orbital
    • Magnetic Quantum Number
    • May contain no more than two electrons
    • from l to -l including 0
    • shape and energy of each is dependent on subshell in which orbital is found
  26. ms
    • Spin Quantum Number
    • intrinsic angular momentum
    • two spin orientations: 1/2 and -1/2
    • Two electrons in the same orbital have opposite spins
    • Electrons in different orbitals with the same ms values have parallel spins
    • Electrons with opposite spins in the same orbitals are often referred to as paired
  27. Electron Configuration
    • First number denotes the principal energy level, the letter designates the subshell, and the superscrip gives the number of electrons
    • Lower the values of the first and second quantum numbers, the lower the energy of the subshell
    • If the material has unpaired electrons, a magnetic field will align the spins of these electrons and weakly attract the atom: Paramagnetic
    • Material that have no unpaired electrons and are slightly repelled are diamagnetic
  28. Periodic Law
    states that the chemical properties of the elements are dependent, in a systematic way, upon their atomic numbers
  29. Periods
    Seven periods representing the principal quantum numbers n=1 to n=7 and each period is filled sequentially
  30. Groups
    Represent elements that have the same electronic configuration in their valence or outermost shell and share similar chemical properties
  31. Representative Elements
    • A elements
    • have either s or p sublevels as their outermost orbitals; in groups 1A to 7A who have incomplete filled s or p subshells with the exception of the noble gases
  32. Nonrepresentative Elements
    • B elements
    • including transtion elements which have partly filled d sublevels
  33. Lanthanide and Actinide Series
    Have partly filled f sublevels
  34. Trends
    • L to R, electrons added, electrons in outermost shell experience more nuclear attraction, more closer and tightly bound to the nucleus
    • U to D, outermost electrons become less tightly bound to the nucleus
  35. Atomic Radius
    Decreases across from L to R and increases from U to D
  36. Ionization Energy
    • energy required to completely remove an electron from a gaseous atom or ion
    • The closer and tightly bound the electron is to the nucleus, the more difficult it will be to remove the electron and the higher the ionization energy
    • L to R, increases
    • U to D, decreases
    • Group I has lower ionization energies because the loss of an electron results in the formation of a stable octet
  37. First Ionization energy
    energy required to remove one valence electron from the parent atom
  38. Second Ionization Energy
    energy needed to remove a second valence electron from the univalenet ion to form the divalent ion, second ionization energy is always larger than the first
  39. Electronegativity
    • L to R increases
    • D to U, increase
    • Measure of attraction an atom has for electrons in a chemical bond;
    • Greater the electronegativity, greater the attraction for bonding electrons
    • Directly related to Ionization Energies
  40. Electron Affinity
    • energy changed that occurs when an electron is added to a gaseous atom, and it represents the ease with which the atom can accept an electron
    • Stronger the attractive pull of the nucleus of the electrons, effective nuclear charge the greater the electron affinity will be
    • Positive Electron Affinity represents energy release when an electron is added to an atom;
    • Negative Electron Affinity represents release of energy
  41. Group IIA (Alkaline Earths)
    Have low electron Affinity values; stable because s subshell is filled
  42. Group VIIA
    • Halogens
    • have high electron affinities because the addition of an electron to the atom results in a completely filled shell which represents a stable electron configuration
    • achieving the stable octet involves a release of energy and the strong attraction of the nucleus for the electron leads to a high energy change
  43. Group VIII
    • Noble Gases
    • Have electron affinities on the order of zero because they already possess a stable octet and cannot readily accept an electron
  44. Metals
    • shiny solids except for mercuy
    • high melting points
    • High densities
    • Can be deformed without breaking
    • Malleable and Ductile
    • Good Conductors of heat and electricity
    • large atomic radius, low ionization, low electronegativity
  45. Nonmetals
    • brittle and show little or no metallic luster
    • High ionization energies and electronegativities
    • poor conductors of heat and electricity
    • Gain electrons easily
  46. Metalloids
    • Semimetals, properties fluctuate
    • Boron
    • Silicon
    • Germanium
    • arsenic
    • antimony
    • tellurium
  47. Hydrogen
    • Resemble halogens when H-
    • and alkali metals
  48. Alkali Metals
    • Group IA
    • densities lower
    • one loosely bound electron in outermost shell
    • easily lose valence electron to form univalent cations
    • react readily with nonmetals and halogens
  49. Alkaline Earths
    • Group IIA
    • two electrons in outer shell
    • smaller radii than Group IA
    • can be removed to form divalent cations
  50. Halogens
    seven valence electrons
  51. Noble Gases
    • group VIII
    • nonreactive
    • complete shell
    • low boiling points
  52. Transition Metals
    • hard, high melting points and boiling points, malleable and high conductivity has d orbitals fill up
    • have different oxidation states
    • can lose various amounts of electrons from s and d orbitals
  53. Chemical Bonds
    Strong attractive forces that hold together atoms in most molecules
  54. Octet Rule
    • which states that an atom tends to bond with other atoms until it has eight electrons in its outermost shell
    • forming a stable electron configuration similar to the noble gases
  55. Exceptions to the Octet Rule
    • Hydrogen: can have only two valence electrons
    • Lithium and Beryllium: which bond to attain two and four valence electrons
    • Boron: which bonds to attain six
    • Elements beyond second row: can expand octets
  56. Significant Figures
    • Any number that isn't a zero is always a sig. fig.
    • Zeros that don't have any non-zeros before them are not sig. figs. (Act as placeholders)
    • Zeros that are between two non-zeros (trapped) are always sig. figs
    • Zeros that are at the end of the number are sig. figs only if the number contains a decimal point
    • Numbers in equations and conversions are considered to be exact, that is have an infinite number of sig figs.
    • When you multiply or divide: Find number with least sig. figs and use that many sig figs in the answer
    • When you add or subtract: the last digit in the final answer must have the same place as the place of the least precise number, that is the leftmost ending
  57. Prefixes
    • Mega: 10^6
    • Kilo: 10^3
    • Hecto: 10^2
    • Deka: 10^1
    • Deci: 10^-1
    • Centi: 10^-2
    • Milli: 10^-3
    • Micro: 10^-6
    • Nano: 10^-9
  58. atomic number
    how many protons are in the nucleus of one atom of this element; also shows the number of the electrons in a stable atom of this element
  59. atomic mass
    How much one atom weighs; protons + neutrons
  60. Law of Multiple Proportions
    When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole number

  61. Symbol
    • Bottom number is always the number of protons
    • Bottom number is also usually the number of electrons but only in a neutral element, ions come later
    • Number of neutrons always the top number minus the bottom number...
  62. Intermolecular Forces
    Weaker forces between molecules, weaker than intramolecular chimical bonds are important in understanding physical properties
  63. Ionice Bonding
    Electrons from an atom with a smaller ionization energy is transferred to an atom with a greater electron affinity and the resulting ions are held together by electrostatc forces
  64. Covalent Bond
    • The compounds that share electrons and result in the formation of a molecule;
    • H20, shares electron with O, so O has two shared electrons completing its valence shell;
    • When two non-metal are put together
    • Contain discrete molecular units with weak intermolecular forces
    • low melting solids and do not conduct electricity in the liquid and aqueous states
    • Can have single, double or triple bonds
    • Can be characterized by two features: Bond Length and Bond energy
  65. Polar Covalent Bonds
    • Partially covalent and partially ionic
    • Between atoms with small difference in electronegativity
    • The bonding electron pair is not shared equally, but pulled towards the more electronegative element
  66. Bond Order
    The number of electron pairs between two atoms
  67. Bond Length
    • average distance between the two nuclei of the atoms involved in the bond
    • number of shared electron pairs increases, atoms are pulled closer together, decrease in bond length
    • Triple bond is shorter than a double bond which is shorter than a single bond
  68. Bond Energy
    • Energy required to separate two bonded atoms
    • Increases as the number of shared electron paires increases
  69. Polar Molecule
    Molecule that has such a separation of positive and negative charges is called a polar molecule
  70. Dipole Moment
    • Vector Quantity: mu
    • mu = qr
  71. Formal Charges
    • The number of electrons officially assigned to an atom in a lewis structure does not always equal the number of valence electrons of the free atom
    • The difference is the formal charge
    • Formal Charge = V - 1/2(Nbonding) - Nnonbonding
  72. Resonance
    • two or more nonidentical lewis structures
    • A lewis structure with small or no formal charge is preferred over a lewis structure with large formal charges
    • A lewis structure in which negative formal charges are placed on more electronegative atoms is more stable one in which the formal charges are placed on less electronegative atoms
  73. Nonpolar Covalent Bond
    • Between atoms that have the same electronegativities
    • Bonding electron pair is shared equally with no separation of charge across thebond
    • Occur in Diatomic Molecules
  74. Coordinate Covalent Bond
    • shared electron pair comes from the lone pair of one of the atoms in the molecule
    • found in Lewis acid-base compounds
  75. Lewis Acid
    compound that can accept an electron pair to form a covalent bond
  76. Lewis Base
    compound that can donate an electron to form a covalent bond
  77. Mass Number
    • = number of protons + number of neutrons
    • = atomic number + number of neutrons
  78. Molecule
    • Acts a unit by itself; moves together;
    • Held together by covalent bonds
    • Ionic Compounds do not form true molecules
  79. Anion
    • Negatively Charged Ion
    • Gains an electron
  80. Cation
    • Positively Charged Ion
    • When two atoms with large differences in electronegativity react; and complete transfer of electrons to the higher electronegative atom;
    • For transfer to occur, difference must be greater than 1.7
  81. Ionic Solid
    • Salt like NaCl
    • Na+ and Cl-
    • Salt refers to any ionic solid
    • Ions can be polyatomic
    • form crystal lattices consisting of infinite arrays of positive and negative ions in wihch the attractive forces between ions of opposite charge are maximized while the repulsive forces between ions of like charge are minimized
  82. Ionic
    Metal and Non-Metal except NH4Cl
  83. Binary Ionic Compounds
    • made up of two different ions, like sodium chloride
    • Put Name of Cation, followed by Anion
  84. KI
    • Potassium Iodide
    • K+ I-
  85. CaS
    • Calcium Sulfide
    • Ca2+ S2-
  86. Li3N
    Lithium Nitride; Li+ N-3
  87. Oxyanions
    An anion containing one atom and one or more oxygen atoms
  88. Covalent Compounds
    Are two non metals;
  89. Sodium Sulfate
  90. Vanadium(V) Fluoride
  91. Mercury (I) Chloride
    • Mercury (I) : Hg22+
    • Hg2Cl2
  92. Tin(II) Fluoride
  93. Mole (Avogadro's Number)
    • 6.022 x 1023
  94. Equivalent Weigths
    • Equivalents: Weight of Compound
    • ------------------------
    • Gram Equivalent Weight

    • GEW: Molar Mass
    • -------------
    • N
  95. Percent Composition
    • = Mass of X in formula X 100
    • Formula Weight of Compound

    • IF GIVEN: percentages of elements and told to find the formula, assume 100 gram and just divide the number by the gram/mol and then divide by the smallest digit;
    • then get the empiral formula; then find the grams per mole and divide the original by new one; and then times the empircal formula
  96. Percent Mass
    How much of something is in a compound
  97. Molecular Formula
    • How many atoms of each element are present in one molecule of the substance
    • Exact number ofatoms of each element in the compound
  98. Empirical Formula
    • Reduced form of this; ratioof atoms
    • simplest whole number ratio of elements in the compound
  99. Chemical Reaction
    When one or more molecules rearrange their atoms to form new molecules
  100. Reactants
    Things on the left side that react to form Products; formed into products until one of the reactants runs out or reaches equilibrium
  101. Products
    The things on the right side that are formed by reactants
  102. Combination Reactions
    • In which two more reactants form one product
    • Can also occur when two compounds react to form a new compound
  103. Single Displacement Reaction
    When an atom or ion of one compound is replaced by an atom of another element; further classified as redox reactions
  104. Decomposition
    • when a compound breaks down into two or more substances; result of heating or electrolysis
    • Oxygen is a usual product
  105. Net Ionic Equations
    • Written in ion form
    • Spectator Ions: do not take part in the overall reaction but simply remain in solution throughout
  106. Double Displacement Reactions
    • Metathesis Reactions
    • Elements from two different compounds displace each other to form two new compounds
  107. Neutralization Reactions
    • When an acid reacts with a base to produce a solution of a salt and water
    • Specific type of double displacements
  108. Limiting Reactants
    Look at number of particles not mass

    Tak + Asdf ---> TakAsdf

    • You are given 45.0 grams of Takalahium (411 g/mol) and 500 grams of Asdfur (5620.5g/mol).
    • Find the limiting reactant in the formation of Takalahium Asdfite.

    • moles:
    • 45.0 g Tak x 1 mol / 411 g = .109 moles.
    • 500.g Asdf x 1 mol / 5620.5 g = .0890 moles.
    • As both are being consumed, you'll see that Asdf is used up
    • first. Therefore it's the limiting reactant, even though there is over ten times as much
    • of it by mass.
  109. Percent Yield
    • Percent Yield= Actual Yield
    • ---------------- X 100%
    • Theoretical Yield
  110. theoretical yield
    amount of product that can be predicted from a balanced equation, assuming that all of the limiting reagent has been used; that no competing side reactions have occurred and all of the product has been collected
  111. actual yield
    amount of product that is isolated from the reaction experimentally
  112. VSEPR Theory
    • Valence Shell Electron Pair Repulsion Theory
    • The arrangement of atoms surrounding a central atom is determined by repulsions between the bonding and the nonbonding electron pairs in the valence shell of the central atom
    • electron pairs arrange themselves as far apart as possible, thereby minimizing repulsion

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  113. Polarity of Molecules
    • Polarity depends on the polarity of the constituent bonds and on the shape of the molecule;
    • Molecule with nonpolar bonds is always nonpolar
    • Molecule with polar bonds may be polar or nonpolar depending on the orientation of the bond dipoles
    • When orbitals overlap head to head the resulting bond is called a sigma bond
    • When orbitals are parallel, pi bond is formed
  114. Here's the combustion of ethane (not balanced): C2H6+ O2 ---> CO2 + H2O
    32.0 g of ethane was burned with 15.0 grams of oxygen gas,
    and 10.8 grams of carbondioxide was formed.
    Calculate the percent yield of carbon
    • Now to convert all given masses into moles:
    • 32.0 g ethane x 1 mol / 30.0 g = 1.07 mol
    • 15.0 g oxygen x 1 mol / 32.0 g = .469 mol

    • Which one is the LR? Let's look at ethane first.
    • If ethane is indeed the LR, then the moles oxygen reacted would be:
    • 1.07 mol ethane x 7 mol oxygen / 2 mol ethane = 3.74 moles oxygen needed.
    • Since the oxygen will have run out first, oxygen is the limiting reactant. Disregard
    • the 1.07 mol.

    • Now to find the theoretical yield. 4 moles of carbon dioxide will form from 7 moles of oxygen:
    • .469 mol oxygen x 4 mol CO2 / 7 mol oxygen = .268 moles.
    • Now to get the mass:
    • .268 moles x 44.0 g / mol = 11.8 grams.
    • That's how much you should get. But experiment shows that you get only 10.8 grams.
    • Finding theoretical yield:
    • 10.8 grams / 11.8 grams x 100 % = 91.5%
  115. Chemical Kinetics
    Study of Rates of Reactions
  116. Intermediate
    • A intermediate product between mechanisms which does not appear in the overall equation
    • Neither a product/reactant
  117. Rate Determining Step
    Slow step in a proposed mechanism; reaction cannot proceed faster than the step
  118. Rate
    • Decrease in concentration of reactants = Increase in concentation of products
    • Time Time
  119. Rate
    • Mol/L X s
    • Also, reactants have a - in front of there rate
  120. Rate Law
    k [A]x[B]y
  121. Rate Constant
    Defined as constant of proportionality between the chemical reaction rate and the concentration of the reactants
  122. Orders of reaction
    • x is the order with respect to A
    • y is the order with respect to B
    • Overall oder of a reactin or reaction order is defined by the sum of the exponents
  123. Rate Problems
    • 1. First look at data
    • 2. Where A is same but B is different in one trial
    • 3. write the trials
    • r1 = k[A]x[By]
    • r2 = k[A]x[B]y
    • Now the big rate is divided by the lower one as well as the equations
    • then solve for the exponent
    • Same thing for where B is same but A is different
  124. Zero Order Reactions
    • has a constant rate
    • Independent of the reactants concentrations
    • K = Msec-1
    • A = Ao - (Ko)(t)

    half life = 0.5Ao/Ko
  125. First Order Reactions
    • sec-1
    • radioactive decay
    • [At] = [Ao]e-kt

    • t1/2 = ln2/k
    • percentage of drug administered remains constant;
    • Lnx = (logx)2.3
  126. Second Order Reactions
    • M-1sec-1
    • rate proportional to the product of the concentration of two reactants or to the square of the concentration of a single reactant
  127. Higher order reactions
    order greater than 2
  128. Mixed Order Reactions
    has a fractional order
  129. Collision Theory of Chemical Kinetics
    • In order for a reaction to occur, molecules must collide with each other
    • Rate of a reaction is proportional to the number of collisions per second between the reacting molecules
    • Reaction rates almost alwyas increase with increasing temperatures
    • Reaction rates decrease with decreasing temperatures
    • Not all reactions result in a chemical reaction
  130. Effective Collision
    One that leads to the formation of products occurs only if the molecules collide with correct orientation and sufficient force to break the existing bonds and form new ones
  131. Activation Energy
    • Minimum energy of collision necessary for a reaction to take place
    • Only a fraction of colliding particles have enough kinetic enegy to exceed the activation energy
    • rate = fZ
    • Z = total number of collisions per second
    • f = fraction of collisions that are effective
  132. Transition State Theory
    • When molecules collide with sufficient energy, a state is formed in which the old bonds are weakened and the new bonds are beginning to form
    • then dissociates into products and new bonds are fully formed
    • Greater energy than reactants or products
    • Do not have a finite lifetime like intermediates
  133. Enthalpy
    • the difference between the potential energy of the products and the potential energy of the reactants
    • A negative enthalpy change: exothermic (heat given off)
    • A positive enthalpy change: endothermic (heat absorbed)
  134. Factors Affecting Reaction Rate
    • Rate of reaction with Increase: increase in number of effective collisions; stabilization of the activated complex compared to the reactants
    • Reactant Concentrations: greater the concentration of reactions, greater the number of effective collisions per time and reaction rate with increase for all but zero order reactions
    • Temperature: reaction rate will increase with temperature; molecules have more energy than activation energy
    • Medium: certain reactions proceed more rapidly in aqueous solution, whereas other reaction may proceed more rapidly in benzene
    • Catalysts: lower the activation energy; increase frequency of collision between the reactants; change orientation; donate electron density to the reactants and reduce intramolecular bonding
  135. Reversible Reaction
    Does not proceed to completion, because the products can react to reform the reactants
  136. Law of Mass Action
    • Equilibrium Constant Kc
    • = [C]c[D]d
    • ----------------
    • [A]a[B]b
    • Coefficients of the equation are used as exponents;
    • that is equal to Q
    • Kc = [product]coefficient/ [reactants]coefficient
    • Q is only constant at equilibrium when it is equal to Kc
  137. Equilibrium Constant
    • Pure solids and liquids do not appear in the equilibrium constant expression
    • Keq is characteristicof a given system at a given temperature
    • If the value of Keq is very large compared to 1, an equilibrium mixture of reactants and prducts will contain very little of the reactants compared to the products
    • If the value of Keq is very small compared to 1, an equilibrium mixture of reactants and products will contain very little of the products compared to the reactants
    • if the value of Keq is close to 1, an equilibrium mixture of products and reactants will contain approximately equal amount of reactants and products
  138. Le Chatelier's Principle
    • If an external stress is applied to a system currently at equilibrium the system will attempt to adjust itself in a way to partially offset the stress
    • Determine the direction in which a reaction at equilibrium will proceed when subjected to a stress such as change in concentration, pressure, temperature, or volume
  139. Change in pressure or volume
    • Change in pressure, causes a change in volume
    • Reactions involving gases greatly impacted
    • pressure and volume are inversely related
  140. Change in temperature
    • Affect equilibrium; heat may be considered at a product in an exothermic reaction and reactant in an endothermic reaction
    • alters equilibrium constant
  141. Law of Conservation of Energy
    • Thermal, chemical, potential and kinetic energies are interconvertible
    • All chemical changes accompanied by energy changes
  142. Spontaneous
    • If under a given set of conditions it can occur by itself without outside assistance
    • May or may not proceed to completion depending on the rate of the reaction
  143. System
    • Particular part of the universe being studied
    • System can be: Isolated
    • closed
    • open
  144. Surroundings or environment
    Everything outside the system
  145. Isolated System
    when it cannot exchange energy or matter with the surroundings as with an insulated bomb reactor
  146. Closed
    when it can exchange energy but not matter with the surroundings as with a steam radiator
  147. Open
    can exchange energy and matter with the surroundings as with a pot of boiling water
  148. Process
    • The system undergoes a process when one or more of its properties changes
    • Associated with a change of state
  149. Isothermal
    A process when the temperature of the system remains constant
  150. Adiabatic
    Process when no heat exchange can occur
  151. Isobaric
    Process where the pressure of the system remains constant
  152. Heat
    • form of energy which can easily transfer to or from a system
    • Result of the temperature difference between the system and its surroundings
    • Transfer will occur spontaneously from a warmer system to a cooler system
    • Heat Absorbed by a system: POSITIVE
    • Heat Lost by a system: NEGATIVE
  153. Endothermic
    Reactions that absorb heat energy
  154. Exothermic
    Reactions that release heat energy
  155. Calorimetry
    • Measures heat changes
    • Constant Volume Calorimetry and Constant-Pressure Calorimetry
    • q = mc(delta)T
    • c = specific heat
  156. Constant Volume Calorimetry
    the volume of the container holding the reacting mixture does not change during the course of the reaction
  157. Macroscopic Properties
    • Temperature
    • Pressure
    • Volume
  158. State Functions
    Properties whose magnitude depends only on the inital and final states of the system and not on the path of the change (how the change was accomplished)
  159. Standard Conditions
    25 degrees C and 1 atm
  160. Standard State
    A substance in its most stable form under standard condtions is said to be in its standard state
  161. Enthalpy
    • express heat changes in constant pressure
    • The change in enthalpy is equal to the heat absorbed or evolved by the system at constant pressure
    • A positive change in enthalpy: Endothermic
    • A negative change in enthalpy: Exothermic

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  162. Exothermic Potential Diagram
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  163. Endothermic Potential Diagram
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  164. Standard Heat of Formation
    • The enthalpy of formation of a compound is the enthalpy change that would occur if one mole of a compound were formed directly from its elements in their standard states
    • It is usually zero for an element in its standard state
  165. Standard Heat of Reaction
    the hypothetical enthalpy change that would occur if the reacter were carried out under standard conditions when reactants in their standard states are converted to products in their standard states at 298 K
  166. Hess's Law
    • states that enthalpies of reactions are additive
    • when reactants are transformed into products, the net change in enthalpy is the same irrespective if the reaction takes place in a series of steps or in a single step
    • Enthalpy depends on the initial and final states not on the path taken (state functions)
  167. Problem of Enthalpies
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  168. Bond Dissociation Energy
    • Averages of the energy required to break a particular type of bond in one more of gaseous molecules
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    • = total energy input - total energy released
  169. Bond Dissociation Energy Problem
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    • Bond------Bond Energy
    • -----------(kJ/mol)
    • H-H-------436
    • O=O------499
    • O-H ------463
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  170. Heat of Combustion
    Fast and Spontaneous
  171. Entropy
    • Measure of disorder
    • J/K or cal/K
    • The greater the order of system, the lower the entropy
    • The greater the disorder of the system, the greater the entropy
    • Solids will have lower entropy than a gas
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    • Also, qrev/T
    • qrev is the heat added tot he system undergoing reversible process
  172. Second Law of Thermodynamics
    • All spontaneous Reactions proceed such that the entropy of the system plus its surroundings
    • Suniverse = Ssystem + Ssurroundings > 0
    • when equal to zero, it is a reversible process
    • A system will tend toward an equilibrium state if left alone
  173. Gibbs Free Energy
    • combines the two factors which affect the spontaneity of a reaction--changes in enthalpy, change in entropy
    • change in free energy of a system represents the maximum amount of energy released by a process

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    • T = absolute temperature
    • The TdeltaS portion = total amount of heat absorbed by a system when its entropy increases reversibly
    • In an equilibrium state, free energy is at a minimum
    • A process will occur spontaneously if the Gibbs function decreases, when less than 0

    • 1. If (Delta)G = (-), reaction spontaneous
    • 2. If (Delta)G = (+), reaction is not sponataneous
    • 3. If (Delta)G = 0, state of equilibrium thus, then (Delta)G = 0, (Delta)H = T(Delta)S
  174. Spontaneous Reactions Conditions
    • (Delta)H (Delta)S Outcome
    • - + Spontaneous at all temperatures
    • + - Non-spontaneous at all temperatures
    • + + Spontaneous only at high temperatures
    • - - Spontaenous only at low temperatures

    Rate of a reaction depends on the activation energy not the delta G
  175. Standard Free Energy
    defined as a process occurring at 25 degrees C and 1 atm pressure and for which concentrations of any solutions involved are 1 M
  176. Standard Free Energy of Formation
    • is thre free energy change that occurs when 1 mole of a compound in its standard state is formed from its elements in their standard states under standard conditions
    • IS zero of any element in there standard states at standard conditions
    • Standard free energy of a reaction: free energy change that occurs when that reaction is carried out under standard state condtions
  177. Reaction Quotient
    DeltaG = -RTlnKeq

    DeltaG = DeltaG(standard) +RTlnQ
  178. States of Matter
    • Gas
    • Liquid
    • Solid
    • Plasma
  179. The forces in matter:
    From Weakest to Strongest
    • Intermolecular Forces
    • Covalent Bonding
    • Ionic Bonding
    • Network Covalent Bonding
  180. Intermolecular Forces
    • London Dispersion Forces
    • Polarizability
    • Dipole-Dipole Attractions
    • Hydrogen Bonding
  181. London Dispersion Forces
    • Forces between non-polar and noble gas molecules
    • Have zero dipole moment
    • More Electron means greater dispersion forces
    • Short lived dipole moments
    • most important when molecules closest to another
    • strength depends on how easily the electrons can move (be polarized)
    • Larger molecules easier to polarize and greater dispersion forces
  182. Polarizability
    • As the atomic number of the atom increase, the # of electrons and protons also increases, this increase changes the momentary dipole action
    • -Based on this principle, as # of molecules or atom increases, the london dispersion forces also increase
  183. Dipole-Dipole Attractions:
    • Forces between polar molecules
    • The positive pole of the molecule attracts the negative pole of another molecule
    • This force is only 1% as a strong as a covalent or ionic bond
    • The strength of the force is inversely proportional to the distance between the two molecules
    • The positive region of molecule is near the negative region of another molecule and this is energetically favorable
  184. Hydrogen Bonding
    • The strongest intermolecular force
    • Its strenght is largely due to the small size of the hydrogen atom and the large polarity in the resulting molecule
    • The atoms that form hydrogen bonding: Fluorine, Oxygen and Nitrogen
    • Can be intra- or inter-
  185. Liquid State
    • Low Compressibility
    • High Density compared to Gas
    • Lack of rigidity
    • Viscosity
    • Surface Tension
    • Capillary Actions
    • Are able to diffuse and evaporate
    • Mix easily to form solutions
  186. Miscibility
    • The degree to which two liquids can mix
    • Oil and water are immiscible
  187. Emulsion
    Violent shaking that can make immiscible liquids form a fairly homogeneous mixture
  188. Viscosity
    • Measure of resistance to flow
    • Stronger the intermolecular forces = greater viscosity
  189. Surface Tension
    • Resistance of liquid to increase its surface area
    • Stronger the intermolecular forces = Greater surface tension
  190. Capillary Action
    the sudden rise in level of liquid in tight space
  191. Cohesive Forces
    Intermolecular forces between the molecules in the liquid
  192. Adhesive Forces
    Intermolecular forces between the liquid and the container
  193. Water creeps up glass because...
    • Water and glass are polar
    • Molecules in glass orient themselves to attract the water molecules (Dipole Dipole Attractions)
    • Both water and glass have strong intermolecular forces so adhesive forces equal cohesive forces
    • Once they reach a balance, the result is a concave meniscus

    For Mercury: Cohesive Forces > Adhesive forces since no attraction between mercury, nonpolar with glass, polar; result convex meniscus
  194. Solid
    • Crystalline Solid: Organized Structures of its Compounds
    • Amorphous Solid: Disorganized Arrangement of its Compounds
    • Lattice: A 3D representation of the components of the solid
    • Unit Cell: The smallest repeating unit of the lattice
    • The two most commong forms of crystal are metallic and ionic crystals
  195. Types of Crystalline Solids
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  196. Ionic Solid
    • aggregates of positively and negatively charged ions
    • no discrete molecules
    • high melting points
    • high boiling points
    • poor electrical conductivity in solid phase
  197. Metallic Solids
    • metal atoms packed together as closely as possible
    • high melting points and boiling points as a result of strong covalent attractions
  198. Unit Cells
    Three types: Simple Cubic, Body centered Cubic, and face centered cubic
  199. Transitions between the 3 States of Matter
    • Fusion
    • Sublimation
    • Evaporation
    • Condensation
  200. Melting Point
    • temperature at which fusion or melting occurs
    • Pure Crystals: have distinct, very sharp melting points
    • Amorphous Solids: such as glass tend to melt over a larger range of temperatures due to their less ordered distribution
  201. Fusion
    • Melting: Solid Liquid
    • Heat of Fusion, D, Hfus is enthalpy change at the melting point of the solid
  202. Sublimation
    • Solid Gas
    • Solid to GAS directly
    • Reverse transition is deposition
    • Change in G = 0
    • Change in G = G (g) - G(s)
    • so G(g) = G(s) at equilibrium
  203. Evaporation
    • Liquid Gas
    • Vaporization, the energy to vaporize 1 mole of a liquid at 1 atm is the heat of vaporization, D Hvap
    • A cooling process
  204. Condensation
    • Gas liquid
    • When a liquid is sealed in a closed container, the liquid evaporates but at the same time the gas condenses back to the liquid state.
    • Equilibrium is reached whent he rate of condensation equals the rate of evaportation.
    • The equilibirum vapor pressure is also the vapor pressure of the liquid
  205. Vapor Pressure
    • The pressure the gas exerts over the liquid
    • Increases as temperature increases
  206. Boiling Point
    temperature at which the vapor pressure of the liquid equals the external pressure
  207. Graph 1
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  208. States of Matter
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    • Solid Liquid Line: Freezing/Melting
    • Liquid Gas Line: Evaporation/Condensation
    • Triple Point Intersection of the lines;
    • Solid Gas Line: Sublimation/Deposition
    • Critical Point: Beyond this point there is no distinction between liquid and gas
    • SOLID: Low temp, and high pressure
    • LIQUID: High temp, and High Pressure
    • GAS: High temp, and low pressure
  209. Triple Point
    The temperature and pressure in which all 3 states of matter co-exist in equilibrium
  210. Critical Point
    the endpoint of the liquid gas line where no matter how much the pressure and temperature are varied, the gas will not liquify
  211. Solid Liquid Line
    The slope of the solid and liquid line is negative if the liquid state is denser than the solid state
  212. Colligative Properties
    Physical Properties derived solely from the number of particles present, not the nature of the particles;
  213. Freezing-Point Depression
    • Solute particles lower the temperature at which the molecules can align themselves into a crystalline structure
    • Change in T = Km
    • m is molality
    • Salt depresses the freezing point of the water when sprinkled on ice
  214. Boiling Point Elevation
    • If Vapor Pressure of a solution is lower than that of the pure solvent, more energy and at a higher temperature will be required before its vapor pressure equals atmospheric pressure
    • Change in T = Km
  215. Osmotic Pressure
    • Substances tenc to flow or diffuse from ligher to lower concentrations which increases entropy, water will diffuse from the compartment containing pure water to the compartment containing the water-solute mixture
    • Water level will rise only to the point at which it exerts a sufficient pressure to counterbalance the tendency of water to flow across the membrane
    • Osmotic Pressure = MRT
    • M = molarity
    • R = Ideal Gas Constant
    • T = Temperature
    • As concentration of the solution increases, osmotic pressure increases
    • Depends on amount of solute, not its identity
  216. Raoult's law
    • When solute B is added to pure solvent A, the vapor pressure of A above the solvent decreases
    • Change in P = Pa - Pb
    • Pa = XaPa(standard)
    • only when the attraction between molecules of the different components of the mixture is equal to the attraction between the molecules of any one component in its pure state
    • Only Ideal Solutions
  217. Gases
    • Move rapidly
    • Are far aparf from each other
    • only very weak intermolecular forces exist between gas particles
    • Expand to fill any volume
    • Compressible
    • Four variables define: Pressure, Volume, Temperature and Number of Molecules
    • 1 atm = 760 mm Hg = 760 torr
    • Temperature usually in Kelvin
    • STP: 273.15 K (0 degrees Celsius) and 1 atm
    • Standard Conditions: 25 degrees Celsius and 298K with enthalpy, entropy, Gibbs and voltage
  218. Ideal Gas
    • A gas whose molecules have no intermolecular forces and occupy no volume
    • At low pressures, High Temperatures behavior like ideal gases
  219. Boyles Law
    • given gasious sample held at constant temperature , the volume of the gas is inversely proportional to its pressure
    • PV = k
    • P1P2 = V1V2
  220. Charles and Gay-Lussac Law
    • At constant pressure, volume of a gas is directly proportional to its absolute temperature
    • Absolute temperate = Kelvins
    • Tk = Tc + 273.15

    V1/T1 = V2/T2

    • Absolute Zero 0K -273.15 degrees celsius
    • Water Freeze 273.15K 0 degrees celsius
    • Water Boils 373.15K 100 degrees celsius
  221. Avogadro's Principle
    n1/V1 = n2/V2
  222. Ideal Gas law
    • PV = nRT
    • R: 8.21 X 10-2 Latm/molK

    Another R is 8.314 J/(K mol)
  223. Density
    • PV = nRT
    • n = m/MM m- mass in g; MM - molar mass
    • PV = mRT/MM
    • D = m/v = P(MM)/RT

    • P1V1/T1 = P2V2/T2
    • stp: 22.4 L for volume

    • Increase pressure, decrease volume
    • Increase temperature, increase volume
  224. Molar Mass
    • first use the P1V1/T1 = P2V2/T2 equation to solve for Volume
    • then do mass/volume
    • then times by 22.4 to get molecular weight
  225. Deviations due to Pressure
    • As pressure of a gas increases, the particles are pushed closed and closer together,
    • As condensation pressure for a given temperatue is approached , intermolecular attraction become more and more significant until the gas condenses into the liquid state
  226. Deviations Due to temperature
    • Temperature is decreased, velocity of molecules decreases, intermolecular forces become significant
    • Closer the temperature to its boiling point, less ideal behavior
  227. Dalton's lawof partial pressures
    • Pressure exerted by each individual gas
    • Total pressure of the gaseous mixture is equal tot he sum of the partial pressure of the individual components
    • PA= PtXa
    • Xa = na/nt

    • first calculate Xa, mole fraction for each gas,
    • then calculate Pa for each gas
  228. Kinetic Molecular Theory
    • Gases are made up of particles whose volumes are negligible compare tot he container volume
    • Gas atoms or molecules exhibit no intermolecular attractions or replusions
    • Gas particles are in continuous, random motion, undergoing collisions with other particles and the container walls
    • Collisions between any two gas particles are elastic, meaning that there is no overall gain or loss of energy
    • The average kinetic energy of gas particles is proportional to the absolute temperature of the gas and is the same for all gases at a given temperature
    • Average Molecular Speed: KE = 1/2mv2 = 3/2KT
  229. Diffusion
    • gas molecules diffuse through a mixture
    • r1/r2= square root of (MM2/MM1)
  230. Effusion
    flow of gas particles under pressure from one compartment to another through a small opening

    R1/r2 = (MM2/MM1)1/2
  231. Solvent
    • component of the solution whose phase remains the same after mixing;
    • If in the same phase, then solvent is is usually in greater quantity
    • Solute molecules move about freely in the solvent
  232. Solvation
    • Interaction between solute and solvent molecules
    • when water is the solvent, its hydration and the resulting solution is known as an aqueous solution
    • When attractive forces between solute and solvent are stronger than those between the solute particles
    • For Non-ionic: Solvation involves van der Waals forces between the solute and solvent molecules
    • Like Dissolves Like
  233. Solubility
    Maximum amount of that substance that can be dissolved in a particular solvent at a particular temperature
  234. Saturation
    • When this maximum amount of solute has been added, solution is at equilibrium
    • It will not dissolve, if more solute is added
  235. Dilute
    A solution in which the proportion of solute to solvent is small
  236. Concentration
    A solution in which the proportion of solute to solvent is large
  237. Crystallization
    When a dissolved solute comes out of solution and forms crystals
  238. Supersaturated Solutions
    • solutions that contain more solute than found in a saturated solution
    • the addition of additional solute will cause the excess solute in the solution to separate, and a saturated solution will form
  239. Aqueous Solutions
    • Solvent is water
    • All Salts of alkali metals are water soluble
    • All salts of the ammonium ion (NH4+) are water soluble
    • All chlorides, bromides, and iodides are water soluble, with the exceptions of Ag+, Pb2+, and Hg22+
    • All salts of the sulfate ion (SO42-) are water soluble, with the exception of Ca2+, Sr2+, Ba2+, and Pb2+
    • Al metal oxides are insoluble, with the exception of the alkali metals and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides
    • All carbonates (CO32-), phosphates (PO43-), sulfides (S2-) and sulfites (SO32-) are insoluble with the exception of the alkali metals and ammonium
  240. Ion Charges
    • Metals, generally form positive ions
    • Nonmetals, form negative ions
  241. Electrolytes
    • Solutes whose solutions are conductive
    • Strong: if it completely dissociates in its constituent ions
  242. Strong Electrolytes
    • Ionic compounds
    • Molecular Compounds with highly polar covalentbonds that dissociate into ions when dissolved
  243. Weak electrolytes
    Ionizes or hydrolyzes incompletely in aqueous solution and only some of the solute is present in ionic form
  244. Nonelectrolytes
    Do not ionize at all in aqueous solution, and retain their molecular structure which limits their solubility
  245. Molarity
    • number of moles of solute per liter of solution
    • Depends on the volume of the solution, not on the volume of solvent used to prepare the solution
  246. Molality
    • solution is the number of moles of solute per kilogram of solvent
    • For dilute, molality = molarity
  247. Normality
    • number of gram equivalent weights of solute per liter of solution
    • Must know the purpose the solution is being used because it is the concentration of the reactive species with which we are concerned
    • reaction dependent
  248. Dilution
    • when solvent is added to a solution of high concentration to produce a solution of lower concentration
    • MiVi = MfVf
  249. Solution Equilibria
    • Solvation tends towards an equilibrium
    • As soon as solute introduced, most of the change taking place is dissociation
    • Rate of solute dissociatio equal to the rate of precipitation and the net concentration of the dissociate solute remain unchanged regardless of the amount of solute added
  250. Solubility Product Constant
    • Ion Product: [An+]m[Bm-]n
    • For saturated solution: [An+]m[Bm-]n
    • Ion product is defined with respect to initial concentrations and does not necessarily represent either an equilibrium or a saturated solution which Ksp does
    • If IP is equal to Ksp, solution is saturated, and the rate at which the salt dissolves equals the rate at which it precipitates out of solution
    • if IP > Ksp, solution is supersaturated, holding more salt than it should be able to at a given temperature and unstable; if the supersaturated solution is disturbed then the salt will precipitate until IP equals Ksp
    • If IP < Ksp, solution is unsaturated and no precipitate will form
  251. Common Ion Effect
    reduction of solubility
  252. Indicators
    • Lithmus paper will turn red: Acidic Solution
    • Lithmus paper will turn blue: Basic Solution
  253. Acids
    • Have sour tast
    • Aqueous solutions can conduct electricity
    • React with bases to form water and a "salt"
    • Nonoxidizing acids react with metals to produce hydrogen gas
    • Causes color changes in plant dyes
  254. Bases
    • Have a bitter taste
    • Feel slippery to the touch
    • Causes color changes in plant dyes
    • React with acids to form water and a "salt"
    • Aqueous solutions can conduct electricity
  255. Arrhenius Acids
    Species that produces H+ (a proton) in an aqueous solution
  256. Arrhenius Bases
    Species that produces OH- (a hydroxide ion) in an aqueous solution
  257. Bronsted Acid
    Species that donates a proton
  258. Bronsted Base
    Species that accepts protons
  259. Conjugate Acid-Base Pairs
    • Bronsted Lowry acids and bases always occurs in paris
    • Are related by the transfer of a proton
  260. Lewis Acid
    Electron Pair Acceptor
  261. Lewis Base
    Electron pair donor
  262. Oxyacids
    Acids formed from oxyanions
  263. pH
    pH = -log[H+] = log(1/[H+])
  264. pOH
    pOH = -log[OH-] = log (1/[OH-])
  265. Kw
    • Kw = [H+][OH-] = 10-14
    • Kw = Ka x Kb
    • If Ka is larger, kb is small
    • Inversely related, Ka and Kb
  266. Percent Ionization
    • Ionized acid concentration at equilibrium
    • ------------------------------------------ X 100%
    • initial concentration of acid

    • strength of an acid to completely ionize
    • The stronger the acid, weaker the conjugate base
  267. Strong Acids
    • HClO4
    • HNO3
    • H2SO4
    • HCl
  268. Strong Bases
    • NaOH
    • KOH
  269. Ka
    • Ka = [H3O][A-]
    • ----------
    • [HA]
    • weaker the acid, lower the Ka
  270. Kb
    • Kb = [B+][OH-]
    • -----------
    • [BOH]
    • Weaker the base, loewr the Kb
  271. Conjugate Acid
    Acid formed when a base gains a proton
  272. Conjugate Base
    when an acid loses a proton
  273. Neutralization reactions
    • Forming a salt and water sometimes
    • salt may precipitate out or remin ionized in solution depending on its solubility and the amount produced
    • go to completion
    • reverse reaction in which the salt ions react with water to give back the acid or bases - hydrolysis
  274. Strong Acid and Strong Base
    • produce a salt and water
    • acid and base neutralized each other
    • ions formed in the reaction do not react with water
  275. Strong Acid and a weak base
    • produces a salt but usually no water is formed since weak bases are usually not hydroxides
    • cation of the salt will react with the water solvent, reforming weak base
    • hydrolysis
  276. Weak acid reacts with strong base
    • solution is basic
    • due to hydrolysis
  277. Weak acid and weak base
    depends on the strengths of the reactants
  278. Acid Equivalence
    An acid equivalent is equal to one mole of H+
  279. Base Equivalence
    A base equivalent is equal to one mole of OH-
  280. Polyvalent
    each mole of the acid or base liberates more than one acid or base equivalent
  281. Amphoteric
    • can act either as an acid or a base depending on the environment
    • can be oxidizing or reducing agents
  282. Titration
    • procedure used to determine the molarity of an acid or base
    • reacting a known volume of a solution of unknown concentration with a known volume of a solution of known concentration
  283. Equivalence Point
    • When the number of acid equivalents equals the number of base equivalents added
    • Not always at pH 7
    • When titrating polyprotic acids/bases there are several equivalence points as each different acidic or basic species is titrated separately
    • estimated by:
    • plotting the pH as a pH meter or an indicator
  284. End Point
    point at which the indicator actually changes color
  285. Strong Acid and Strong base titration
    • The equivalence point : at pH 7
    • solution: Neutral
  286. Weak acid and Strong Base titration
    • pH significantly changes early
    • and equivalence point in the basic ranges
  287. Buffers
    • consists of a mixture of a weak acid and its salt(consists of its conjugate base and a cation) or a mixture of a weak base and its salt(consists of its conjugate acid and an anion)
    • resists changes in pH
  288. Henderson-Hasselbalch Equation
    Estimamte pH of a solution in the buffer region where the concentrations of the species and its conjugate are present in approximately equal concentrations

    For Weak Acid Buffer:

    • pH = pKa + log([conjugate base]/[weak acid])
    • pOH = pKb + log([conjugate acid]/[weak base])

    • [conjugate base] = [weak acid], in a titration, half way to the equivalence point
    • pH = pKa because the log 1 = 0

    • for the base: [conjugate acid] = [weak base]
    • pOH = pKb
  289. Electrochemical Reactions
    • include spontaneous reactions that produce electrical energy, and nonspontaneous reactions that use electrical energy to produce a chemical change
    • reactions alwyas involve a transfer of electrons with conservation of charge and mass
    • Electrons flow from the ANODE to the CATHODE
  290. Oxidation
    Loss of electrons
  291. Reduction
    Gain of Electron
  292. Oxidizing Agent
    causes another atom in a redox reaction to undergo oxidation
  293. Reducing Agent
    cause the other atom to be reduced
  294. Oxidation Numbers
    • element is said to be oxidized if its oxidation number is increased
    • element is said to be reduced if its oxidation number is decreased
    • oxidation number of a free element is zero
    • oxidation number for a monoatomic ion is equal to the charge of the ion
    • oxidation number of Hydrogen is +1, can be negative with less electronegative elements
    • oxidation number of Oxygen is -2, but can be positive with more electronegative elements; also as in a peroxided ion, O is -1
    • sume of the oxidation number is zero; and for polyatomic ions, equal to the charge
    • Metallic have positive oxidation numbers, and nonmetals have positive and negative
  295. Three types of redox reactions
    • Combination: oxxur with one or more free elements
    • Decomposition: lead to the production of one or more free elements
    • Displacement: an atom or an ion of one element is displaced from a given compound by an atom from a totally different element
  296. Galvanic Cell
    • Spontaneous Reactions occur
    • Voltaic cells
    • Negative G
    • supply energy and are used to do work
    • energy can be harnessed by separating into half cells
    • anode: negative because the spontaneous oxidation reaction that takes place at the anode is the original source of the cell's negative charge
    • charge is spontaneously creates as electrons are released by the oxidizing species or the anode, this is the source of electrons, the anode is considered to be negative
  297. Electrolytic Cells
    • Non spontaneous reactions
    • Positive G
    • Electrical energy required to induce reaction
    • placed in one container
    • uses faraday's constant
    • anode: Positive, since attached to the positive pole of the battery and so attracts anion from the solution
    • electrons are forced through the cathode where they encounter the species which is to be reduced; cathode is providing electrons and cathode is negative
  298. Electrodes
    At which oxidation and reduction occur
  299. Anode
    Electrode at which oxidation occurs
  300. Cathode
    electrode where reduction occurs
  301. Salt Bridge
    • dissipates the charge gradient
    • permits the exchange of cations and anions
  302. Cell Diagram
    • Shorthand notation representing the reactions
    • reactants and products are always listed from left to right in the form:
    • anode I anode solution II cathode solution I cathode
    • single vertical line: phase boundary
    • double vertical line: indicates the presence of a salt bridge or some other type of barrier
  303. Reduction Potential
    • Tendecy of a species to acquire electrons and be reduced
    • more positive, greater the species' tendency to be reduced
  304. Standard Hydrogen Electrode
    0.00 volts
  305. standard reduction potention
    • measure under standard conditions 25 degrees celsius, a 1 M concentration for each ion participating in the reaction, a partial pressure of 1 atom for each gas that is part of the reaction, and metals in their pure state
    • a higher E, means greater tendency for reduction to occur
    • lower E, means greater tendency for oxidation to occur
  306. Standard Electromotive Force
    • Difference between two half cells
    • When adding, do not multiply by the number of moles oxidized or reduced
    • EMF for Galvanic: Positive
    • EMF for Electrolytic: Negative
    • Change in G = -nFEcell
    • Faradays in J/V (coulombs); G in Joules
    • concentration does effect EMF
  307. Nernst Equation
    E cell = Ecell - (RT/nF)(Ln Q)

    Change in G = -RTlnKeq

    nFEcell = RTlnKeq
  308. Binding Energy
    • Energy required to break up a given nucleus into its constituent protons and neutrons
    • energy converted using E=mc2 resulting in a larger mass for the constituent protons and neutrons than that of the original nucleus, difference being called the mass defect
    • holds the nucleons together in the nucleus
    • energy per nucleon peaks at iron, which implies that iron is the most stable atom
    • in general, intermediate-sized nuclei are more stable than large and small nuclei
  309. Nucleons
    • Protons or neutrons
    • Make up the nucleus are held together with considerably more energy than the energy needed to hold electrons in orbit around the nucleus
  310. Number of Nucleons
    A is an integer equal to it which is neutrons and protons
  311. Radionuclide
    for any radioactive isotope
  312. Nuclear Reactions
    • Elements or isotopes are changed from one to another
    • Reactions result in the release or absorption of large amounts of energy
    • Reaction rates are generally not affected by catalysts, temperature, or pressure
    • Protons, neutrons, or electrons can be involved
  313. Chemical Reactions
    • Atoms can be rearranged by the formation or breaking of chemical bonds
    • Reactions generally result in the release or absorption of small amounts of energy
    • Reaction rates are generally affected by catalysts, temperature, or pressure
    • Only electrons in the affected orbital of the atom are involved in the formation and breaking of bonds
  314. Mass Defect
    Every nucleus has a smaller mass than the combined mass of its constituent protons and neutrons
  315. Fusion
    • when small nuclei combine with a larger nucleus
    • can only take place at extremely high temperatures, referred to as thermonuclear reactions
  316. Fission
    • process in which a large, heavy atom splits to form smaller nuclei and one or more neutrons
    • large nucleus is more unstable, than its products there is the release of a large amount of energy
    • Spontaneous fission rarely occurs
    • chain reaction can occur when fission reactions release more neutrons and in turn these neutrons cause other atoms to undergo fission
  317. Radioactive Decay
    • naturally occuring spontaneous decay of certain nuclei accompanied by the emission of specific particles
    • certain type of fission
    • Integer arithmetic of particle and isotope species
    • radioactive half-life problems
    • use of exponential decay curves and decay constants
  318. Alpha Decay
    • emission an a-particle, which is a 4He nucleus that consists of two protons and two neutrons
    • alpha particle is very massive (compared to a beta particle) and doubly charged
    • alpha particles interact with matter very easily
    • do not penetrate shielding very far
    • 2 less the Z and 4 less the mass number
    • Zdaughter = Zparent - 2
    • Adaughter = Aparent - 4
  319. Beta Decay
    • emission of a B-particle
    • electron given the symbol e- or B-
    • more penetrating than alpha radiation
    • in some cases a positron is emitted
    • positron: e+ or B+
    • neutron disappears and a proton takes its place
    • mass number unchanged, atomic number increased by 1

    • in positron decay,
    • proton splits into a positron and a neutron
    • mass number is unchanged, and parent's atomic number is decreased by 1
  320. Gamma decay
    • emission of y-particles which are high energy photons
    • carry no charge and simply lower the energy of the emitting (parent) nucleus without changing the mass number or the atomic number,
    • A is the same and z is the same
  321. Electron Capture
    • Certain unstable radionuclides are capable of capturing an inner electron that combines with a proton to form a neutron
    • atomic number one less but mass same
    • rare process
    • inverse B- decay
Card Set
Chemistry Section of PCAT