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102 Cards in this Set
- Front
- Back
- 3rd side (hint)
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Common Strong Acids - 7
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HCl
HVr HI HClO3 HClO4 HNO3 H2SO4 |
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Common Strong Bases
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Group 1A metal hydroxides
(LiOH, NaOH, KOH, RbOH, CsOH) Heavy group 2A metal hydroxides [Ca(OH)2, Sr(OH)2, Ba(OH)2] Common soluble metal hydroxides |
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Monoprotic acid
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when molecules of different acids can ionize to form different #s of H+ ions
ex. HCL, HNO3 > 1 H+ ion to 1 ion- |
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diprotic acid
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yields 2 H+ per molecule of acid
ex. H2SO4 > 2 H+ to 1 ion- |
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strong acids and bases
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strong electrolytes (completely ionized in solution), more reactive than weak when the reactivity depends only on the concentrations of H+, most acids are weak
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most common weak base
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NH3
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Acids
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1. substances that ionize in aq solns. to form hydrogen ions
2. increases concentration of H+ (aq) ions 3. called proton donors - H+ is simply a proton (solvated by water molecules. 4. in a chem eq = H+ (aq) |
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Bases
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1. substances hat accept (react with) H+ ions
2. produce hydroxide ions (OH-) when the dissolve in water 3. some bases do not contain (OH-) ex, NH3 |
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Electrolytes - soluble ionic compounds
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all ionic are strong electrolytes; no weak, no non
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Electrolytes - soluble molecular compounds
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Strong acids are strong electrolytes; weak acids and bases are weak electrolytes; all other compounds are non
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Identifying strong and weak electrolytes
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1. determine ionic or molecular
2. If ionic - its strong. 3. If molecular, determine acid (H) or base (OH) 4. determine strong from memory 5. non - not a acid or NH3 |
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oxidation-reduction reactions
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or redox
a reaction in which electrons are transferred between reactants |
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precipitation reactions
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cations and anions come together to form an insoluble ionic compound
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neutralization reactions
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H+ ions and OH- ions come together to form H2O molecules
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oxidation
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loss of electrons by a substance; when an atom, ion or molecule has become more positively charged
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ex. Ca, which has no net charge, is oxidized - forming Ca2+.
aka rusting - when metal corrodes, it loses electrons and forms cations (+) |
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Oxidation # Rules
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1. Sum of ox # of each element in a compound = 0
2. Ox # of 0 in a compound = -2 (except in peroxides (O2 2-) = -1 3. Ox # of H in a compound = +1 to NM (except in metallic hydride = -1 to M) 4. Ox # of free element = 0 (M 5. Ox # of diatomic molecule = 0 (NM) 6. Group 1A and 2A elements in combined state = +1, +2 |
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reduction
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gain of electrons by a substance; when an atom, ion, or molecule has become more negatively charged
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when Ca is oxidized, oxygen is transformed fom neutral O2 to two O2- ions
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relationship b/w oxidation and reduction
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When one reactant loses electrons (oxidation), another must gain them. The oxidation of one substance is always accompanied by the reduction of another as electrons are transferred among them
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oxidation numbers
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a way of keeping track of electrons gained by the substance reduced and those lost by the substance oxidized
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when each atom in a neutral molecule or charged species is assigned an ox number, which is actual charge for monatomic ion
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When does oxidation occur?
When does reduction occur? |
when oxidation numbers increase
when oxidation numbers decrease |
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elements most readily oxidized?
elements least readily oxidized? |
metals in Group 1A and 2A; then transition metals. Finally, nonmetals
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Free element example - ox#
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Mg =0
K =0 Cu =0 Fe =0 |
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Diatomic molecule example - ox#
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N2 =0
Cl2 =0 |
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Combined examples - ox#
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MgSO4
+2, +6, -8 =0 KNO3 +1, +5, -6 =0 CH3OH -2, +3, -2, +1 =0 |
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Activity Series of metals in aqueous soln -
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a list of metals arranged in order of decreasing ease with which they are oxidized
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Ease of Oxidation
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any metal on the list can be oxidized by the ions of elements below it
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Zn is oxidized by aq Cu 2+, but Ag is not.
aka Zn loses electrons more readily than Ag Zn is easier to oxidize than Ag |
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Active metals
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1A, alkali and 2A, alkaline earth metals at top of the list are easily oxidized (aka easily make into compounds)
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Noble metals
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at bottom of series, ex transition elements from 8B and 1B, are stable and form compounds less readily
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Activity Series rules
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1. all meal below H will show no reaction w/acid (NR)
2. all metals above H will react w/acid (HCl) and release H2 gas |
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Electrolyte
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ionic compounds - NaCl
acids - H+Cl base - KOH |
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Nonelectrolyte
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compounds w/C and H
ex. CH3HO, C6H12O6 |
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Concentration of ions example
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Given 2.5 M KCl
K+ = 2.5M Cl- = 2.5 M |
Given 1.28 Mg3(PO4)2
Mg 2+ = 3 x 1.28 = 3.84M PO4 3- = 2 x 1.28 = 2.56M |
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Concentration of Ions problems
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1. separate elements into ions
2. multiply subscript by M given |
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Calculate based on equation
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when only given wt of one reactant, find mass of product.
Could be in gram, mols, or molecules |
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Calculate Limiting Reactant
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if mass of both reactants are given, find initial, change and final; find LR and excess
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LR steps
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1. Initial - convert weight to moles and / by lowest coefficient
2. Change - use LR mol from Initial, x by coefficient, add - to reactants, and + to product 3. Final - add/subtract; LR = 0, excess is coefficient |
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Calculate Molarity
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M = moles soluble/liter of soln
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Precipitation reactions - question sample
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write balanced net ionic equation that occurs in each case. Identify the spectator ions in each reactions
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Precipitation reactions eq steps
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1. Write out reaction equation and cross 1sts and 2nds
2. Match + and - ions w/subscripts 3. balance eq 4. Split: but NOT (s, l, g) 5. cancel out spectator ions 6. write net-ionic eq In answer, keep ionic charges and carry (aq, s, l, g) |
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Precipitation Reaction: Solubility Rules
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1. NH4 compounds - soluble (aq)
2. Group 1A metals > compounds > soluble (aq) 3. NO3- and CH3COO- compounds > soluble (aq) 4. Cl-, Br-, I- compounds (aq) - except compounds with Ag+, Hg2+ and Pb2+ 5. SO4 2- (aq) - except compounds with Sr 2+, Ba 2+, Hg 2+, and Pb 2+ 6. do not split (s, l, g) |
Metals are (s)
Gases are (s) |
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Soluble
Insoluble |
aq
s, l, g |
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spectator ions
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similar ions on both side of the reaction; coefficient must be the same or it was not balanced correctly
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Acid-Base reactions > neutralizations
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1. net-ionic answer always H+ + OH- > H2O
2. any element w/H in front is an acid 3. H2O is always (l) and H+ is always (aq) |
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calculate Molar Mass
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use mass of each element
answer in grams/moles |
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calculate formula weight
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use mass of element
answer in amu |
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how to balance equations
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1. start w/subscript. if not balanced, than make even # by multiplying
3. if polyatomic ions on both sides, balance individually 4. if polyatomic ions on one side, ignore polyatomic ion and follow subscript |
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g to mols
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(mm)
=mol/g |
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mol to g
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(mm)
=g/mol |
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mol to # molecules
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=molecule/mol
molecule is 6.02 x 10,23 |
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g to # molecules
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= mol/g x molecule/mol
molecule is 6.02 x 10,23 |
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combination reactions
(simple) |
two or more substance react to form one product
M + NM > ionic compound |
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decomposition reactions
(simple) |
one substance undergoe a reaction to produce two or more other substances; heat is often involved; triangle above arrow
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combustion reactions
(complex) not on exam |
rapid reactions that product a flame; involve O2 in air as reactant
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formula weight
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(fw) of a substance is the sum of atomic (aw) weights of each atom in its chemical formula
= mass of a substance |
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% composition formula
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% element = (# of atoms of that element)(atomic weight of element)(100%) / fw of compound
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composition by mass
= empirical formula |
1. assume 100% = 100 grams
2. convert grams > mols 3. divide each # mol by lowest # mol 4. write whole # as subscript |
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molecular formula
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= empirical formulat
1. subscripts x atomic weight of each element = mm ef 2. mm / mass of ef |
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phosphate ion
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PO4 3-
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carbonate ion
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CO3 2-
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sulfate ion
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SO4 2-
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acetate ion
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C2H3O2 -
or CH3COO - |
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ammonium ion
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NH4 +
only ion w/+ charge |
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nitrate ion
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NO3 -
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nitrite ion
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NO2 -
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hydroxide ion
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OH -
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cyanide ion
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CN -
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-ide
-hypo/ite -ite -ate -per/ate |
for Cl, Br, I only
Br - > bromide ion BrO - > hypobromite BrO2 - > bromite BrO3 - > bromate BrO4 - > perbromate |
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element to ion
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+ ion, subtract # of electron
- ion, add # of electron |
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exponents
/ x + - |
move decimal <, +
move decimal >, - / - exp are subtract top from bottom x - exp are added + and - - powers of 10 must be same |
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significant figures
x and / + - |
x and / - round answer to lowest digit of all #s
+ and - - round answer to lowest decimal place |
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Ions
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charged atoms
M > lose electrons, become + ion NM > gain electrons, become - ion |
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Density formula
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d = Mass / volume
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Element to Ion math examples
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Metal: Li 3p 3e
Electron loss: 3p 2e- Ion gains: Li+ NonMetal: N 7p 7e electron gain: 7p 10e+ ion loss: N3- |
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naming acids
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1. H in front is always H+
2. -ide to -ic and add "hydro" (when no oxygen is present) and "acid" -ate to -ic and add "acid" -ite to -ous and add "acid" |
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naming ionic compound
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1st name: metal, 2nd name: NM or polyatomic
Group A: no roman numeral needed except Ag +, Zn 2+, Cd 2+ Group B: all metals need roman numerals including a metals |
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naming ionic compounds: Fe, Cu, Co, Cr, Pb
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Fe 2+, Iron (II)
Fe 3+, Iron (III) Cu 1+, Copper (I) Group B Cu 2+, Copper (II) Cr 2+, Chromium (II) Co 2+, Cobalt (II) Pb 2+, Lead (II) |
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naming molecular compounds
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NM to NM
use 5A, 6A, 7A 1st word: never modify ending 2nd word: modify to -ide and change beginning -mono, di, tri, tetra, penta, hexa |
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Periodic table
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1A = alkali metals; +
2A = alkaline earth metals; 2+ 3A = 3+ 4A = 0 5A = 3- 6A = Chalcogens 2- 7A = Halogens - 8A = Noble; stable |
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polyatomic rule
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less is -ite
more is -ate |
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volume conversion
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1 mL = 1 cm^3 = 760 mm Hg
1 gal = 3.8754 L |
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meter conversions
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1 m = 100 cm = 1000 mm = 39.37 cm
1 in = 2.54 cm 1 L = 1.06 qt |
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mass conversion
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1 kg = 1000 g = 2.205 lbs
1 lb = 453.6 g |
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-1 polyatomic ions
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Acetate C2H3O2
Nitrate NO3 Nitrite NO2 |
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-2 polyatomic ions
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carbonate CO3
sulfate SO4 sulfite SO3 |
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-3 polyatomic ions
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phosphate PO4
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only +1 polyatomic ion
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ammonium NH4
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Giga
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G
10^9 |
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Mega
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M
10^6 |
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Kilo
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k
10^3 |
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Deci
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d
10^-1 0.1 m |
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Centi
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c
10^-2 0.01 m |
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Milli
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m
10^-3 0.001 m |
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Micro
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u
10^-6 |
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Nano
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n
10^-9 |
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Pico
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p
10^-12 |
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Femto
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f
10^-15 |
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Temperature conversion:
Celsius= |
5 / 9 (F - 32)
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Temperature conversion:
Fahrenheit = |
9 / 5 (C + 32)
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Temperature conversion:
Kelvins = |
C + 273.15
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Millikan
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oil drop experiment; discovered charge of an electron
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extensive properties
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a property that depends on the among of material considered; for ex. mass or volume
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intensive property
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a property that is independent of the amount of material considered; for ex. density
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