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84 Cards in this Set
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Saturation |
Maximum amount of solute that can be dissolved in a solution at equalibrium. This amount is called solubility |
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Unsaturated |
Contains less than the maximum amount which can be dissolved. |
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Supersaturated |
Contains more than the maximum amount which can be dissolved at equalibrium. |
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Electrolytes |
Are substances whose water solutions will conduct an electric current because they dissociate into ions in solution. Can produce electricity and conduct electricity. |
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Strong electrolytes |
Substances which are essentially completely dissociated in aqueous solution. Completely separate in solution |
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Weak electrolytes |
Dissociated to a relativley small extent in solution. And produce a few ions the solution is thus much smaller conductor than a strong electrolyte. |
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Nonelectrolytes |
are substances which produce nonconducting solutions because they do not dissociate into ions in water. Examples: Sugars, alcohols |
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Arrhenius definition of acid and base |
Acid- substance produces hydrogen ions in solution (Acid anhydride Generally non Metallic oxides which reacts with water to give acidic solutions) Base- substance that produces hydroxide ions in solution (Base anhydrides generally metal oxides which react with water gives basic solutions) |
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Amphoteric substances |
Can function as either an acid or a base. In water they do not disolve and behaves as neither acids nor bases. This indicates that not all metal oxides or hydroxide are basic in water. |
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Bronsted Lowry definition |
Acid: proton donor Base: proton acceptor Proton = H+ ion Conjugate acid base theory Acid That gives an H+ will yield conjugate base. Base that will accept an H+ will yield conjugate acid. |
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Lewis definition |
Acid: electron pair acceptor Base: electron pair donor |
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Electro negativity of O |
Each lone O will pull an electron from the center element allowing the H to leave more readily Acidity increases |
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Pg 216 |
A |
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Bases in water OH NH |
1 bases are strong (100% ionized) 2 bases are weak (few % ionized) |
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Salts |
Ionic compounds containing cations and anions Strong acid and base = neutral _____ Strong acid weak base = acidic _____ Weak acid strong base= basic _____ |
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Oxidation |
Loss of electron Increased oxidation state |
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Reduction |
Gains electron Decreasing in oxidation state. |
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Oxidation states |
0 Free elements (Cu Na O2 Cl2) 1+ Hydrogen with nonmetals (1- if with metals hydrides) 2- Oxygen in most compounds (1- in peroxides H2O2) 1- Halides (simple salts NaCl cacl2) 1+ Group IA alkali metals 2+ Group IIA alkaline earth metals |
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Oxidation states of metals |
A moving down a representative group of metal the lower oxidation state is more stable E.G. Al prefers 3+ while TI prefers 1+ |
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Transition (B) groups |
(1) Lose s electrons first (2) lose any number of d electrons next |
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Balancing oxidation and reduction reaction step 1 |
Assign oxidation states to all reactants and products |
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Balancing oxidation and reduction reaction step 2 |
Decide what is oxidized and what is reduced |
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Balancing oxidation and reduction reaction step 3 |
Mass balance Pull out the species involved in each half reaction and balance each half reaction with respect to atoms (mass) First balance the atoms oxidized and reduced then: If acidic Add water to the side needing an O for balancing Ad H+ to the opposite side to balance the Hs If basic Add OH- to the side needing O (normally double) Add water to the opposite side to balance Hs |
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Balancing oxidation and reduction reaction step 4 |
Charge balance Balance with respect to charge using electrons. Check see that the oxidation loses the appropriate number of electrons and the reduction gains the appropriate number of electrons. |
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Equivalents |
Number of H+,OH-, or e- gained or lost, per mole. |
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Eq wt For wt |
g/eq g/mole |
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Normality |
EQ solute/L Soln N=EQ/V EQ=NV |
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Molarity |
Moles solute/L soln M=MOLES/V MOLES=MV |
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Titration equations |
Eq acids=Eq base Eq oxidized = EQ reduced Eq=NV (eq/L ×L= Eq) EQ= wt/eq wt (g/(g/eq) = Eq) |
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Balancing oxidation and reduction reaction step 5 |
Balance so that when we add half reaction the electrons will cancel from opposite sides of the equation |
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Ionic (electrovalent)bond
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Occurs when elements on the left electro positive metals meet those on the right electronegative nonmetallic if electronegative difference on the Pauling scale is greater than 1.7 this will form
Face centered cubic |
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Covalent bond
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Sharing of electrons formed when both atoms are on the right (electro negative nonmetals) and has a small difference in electronegativities
Unequal sharing |
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Ionic bond
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C
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Metal bonds to nometals
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Forms ionic salt
Metals conduct electricity High melting point |
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Electrostatic attraction
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Example NaCl
Na atom Losses of an electron (+) Cl gains the electron (–) |
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Water and ionic compounds
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Water breaks bonds and electrolyte are free and can now conduct electricity.
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Vesper theory
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like charges repell
Valance shell like charges 180 degrees apart. |
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Lewis structures
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A
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Poly atomic ions cation(s)
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Ammonium
NH4+ only positive |
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Polyatomic ions anion(s)
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Nitrate NO3^–
Sulfate SO4^(2–) Phosphate PO4^(3–) |
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Allotropes
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Different forms of the same element as a result of the differences in molecular structure such as O2 and O3 or different packing of atoms in a solid.
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Sp Hybridization
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2 paired electrons
Orbital geometry Linear |
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Sp^2 hybridization
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3 pairs of electronsOrbital geometry Triangular planar
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Sp^3 hybridization |
4 sets of paired electrons Orbital geometry Tetrahedral |
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Dsp^3 hybridization
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5 paired electronsOrbital geometry Triangular bipyramidal
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D^2sp^3 hybridization |
6 electron pairs Orbital geometry |
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Factors influencing formation of bonds
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Energy must be put into system (endothermic) more energy in than out
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Coordinate covalent bond
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When both electrons shared come from the same atom.
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Polar covalent bond
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Any difference in electronegativity will result in an unequal sharing of electrons
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Polar and nonpolar shape
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Dipole cancel in p usually linear
Dipoles do not cancel in np will not be symmetrical in all positions |
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Hydrogen bonding
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An extreme polarization when the smallest most electronegative element (N,O,F) attach to H.
Strong bond responsible for water being liquid at room temp. And for water to be polar enough to disolve ionic salts. |
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unpolarized
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Ionic potential equals q/r (low)
More ionic in nature Downward trend of periodic table More soluble I'm H2O More basic Less color Higher melting point Less cation hydrolysis (S orbitals) |
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Polarized
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Ionic potential equals q/r (high)
More covalent in nature Trends down periodic table Less soluble in water More acidic More color Lower melting point More cation hydrolysis |
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Pg 182
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Examples
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Effect of anion on the character of compounds
Polarization (distortion) increases with |
1 increase in size
2 increase in charge Rows right to left Groups top to bottom |
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Valence bond theory
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Explains how a bond forms between 2 neighboring atoms using ovelapping one atomic orbital from each yielding a new wave form which holds two electrons of opposite spin.
The degree of the overlap is due to the strength of the covalent bond. |
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Molecular Orbital (MO) |
Theory considers all of the atomic orbitals from both atoms and shows how these interact to form orbitals on the resulting molecule called molecular orbitals Even orbitals which are not involved in the bond are described. Two types of covalent bonds are formed: |
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Hybridization of d orbitals
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From the third period down on down we get hybridization ro include an aditional available (d) orbital. These structures will exceed the octet rule.
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Multiple bonds n2
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Triple bond
1 sigma 009999652627351861618 |
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Resonance
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Is the mathematical combination of two or more valence bond structures that are equivalent
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Electron pair repulsion theory
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Principle
Pairs of electrons in the valence shell(whether bonded or nonbonded) are always arranged to allow the greatest spacial separation. The arrangement depends only on number of electron pairs. |
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Molecules with resonance
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NO2 SO2 .........
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Triple bonds
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N2 .........
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Oxidation reduction balance electrons
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Multiply both half reactions by common multiple.
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Find oxidation state
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Use knowns E.G. O^2– H^+ and note charge of molecule and find unknown.
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More acidic
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Bigger molecule weaker bonds
Less electronegative weaker bonds More Os that are separate from Hs |
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Oxidation trends
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Top down and right to left
Larger size smaller charge |
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Find percentage by weight
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Use 100 as the value of solvent in grams
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Mole fractions
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Moles of solute over moles of solution(solvent and solute)
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Base and conjugate acid and vice versa.
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Base recieves proton to yield acid
Acid gives proton yield base |
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Lewis acid
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Accepts electron pair
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Aqueous solution check
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The charge of ions in solution
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Acid anhydrides
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Nonmetal oxides are acidic
Metal oxides are basic ...... |
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More covalent in cations
Less ionic |
Ionic potential equals charge(q)/radius(r)
Smaller radius Larger positive charge |
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More covalent and less ionic in anions
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Larger radius
Larger negative charge |
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Hybridization
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A
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More covalent and less ionic
radius |
Cation smaller
Anion larger |
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More covalent and less ionic charge
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Cation larger Charge(+)
Anion larger charge (–) |
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Ionic properties (covalent vise versa)
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Greater
Melting point Solubility in water Brittleness Hardness Less color |
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Exceptions to octet rule
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BA Only needs 4 valence electrons
B and all 3a rule only contains 6 electrons Elements that have an expanded valence has d orbitals transitions. |
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Molecular geometry
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Lone pairs bent
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Normality
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Find molarity and convert with conversion factor for
equivalents per mole |
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Acid strength
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Central molecule more electro negative (a proximity to F) will not pull as hard and h will disassociate.
Ratio of Os to H |
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Conjugate acid
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Base takes on h
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