Download Review: key points in General Chem I

REVIEW
1
Copyright (c) 2011 by Michael A. Janusa, PhD. All rights reserved.
Measurement and Significant Figures
• To indicate the precision of a measured
number (or result of calculations on
measured numbers), we often use the
concept of significant figures.
2 mL
1.5 mL
– Significant figures are those digits in a
1 mL
measured number (or result of the calculation
with a measured number) that include all
certain digits plus a final one having some
uncertainty (first digit basically guessing).
30.2246743 cm
cm
30.22 cm
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• Rules for Significant Figures:
– All nonzero digits are significant.
i.e. 111
1286
111 3 SF
1286 4 SF
– Zeros between significant figures are significant.
i.e. 1001
20,006
1001 4 SF
20,006 5 SF
– Zeros preceding the first nonzero digit are not significant.
i.e. 0.0002
0.00206
0.0002 1 SF 0.00206 3 SF
– Zeros to the right of the decimal after a nonzero digit are
significant.
i.e. 0.00300
9.00
9.10
90.0
0.00300 3 SF 9.00 3 SF 9.10 3 SF
90.0 3 SF
– Zeros at the end of a nondecimal number may or may not be
significant. (Use scientific notation.)
i.e. 900
900.
1, 2, or 3 SF
3 SF
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• Scientific notation – is the representation of a number in the form
A. x 10n, where A is a number (sign digits only) with a single
nonzero digit to the left of the decimal point and n is an integer or
whole number.
900
9 x 102 1 SF
9.0 x 102 2 SF
9.00 x 102 3 SF
300,000,000 (write with 3 SF)
0.0000301
843.4
3.00 x 108 3 SF
3.01x 10-5 3 SF
8.434 x 102 4 SF
0.00421
4.21 x 10-3 3 SF
6.39 x 10-4
0.000639
3.275 x 102
327.5
Note: exp or EE represents “x 10”
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• Number of significant figures refers to the number of digits
reported for the value of a measured or calculated quantity,
indicating the precision of the value. [Basically means if all
quantities have X sign fig can’t report final answer with
more than X sign figs: measurement or calculation dictates
sign figs.]
– When multiplying and dividing measured quantities, give as many
significant figures as the least found in the measurements used.
• 2.1 x 3.52 = 7.392 = 7.4
– Which gets us to rounding: left most digit to be dropped – 5 or
greater add 1 to last digit to be retained, less than five leave alone –
1.2143 -- 1.21
– Multiple step calculation - Guard digit: 1.214
– When adding or subtracting measured quantities, give the same
number of decimals as the least found in the measurements used.
• 84.2
(3 sign)
• +22.321
(5 sign)
• 106.521
• 106.5 (4 sign)
arithmetic rules if combined ( ), x / , + -
5
3.38 – 3.012 = 0.368 = 0.37
2.4 x 10-3 + 3.56 x 10-1 =
0.0024
+0.356
0.3584 = 3.58 x 10-1
2.568 x 5.8 = 14.8944 = 3.55814 = 3.6
4.186
4.186
or 14.9 gives 3.55948
4.18 – 58.16 x (3.38 – 3.01) =
4.18 – 58.16 x (0.37) = 4.18 – 21.5192 = -17.3392 = -17
6.3 + 7.2 = 13.5 = 25.685 = 25.7
0.5256
0.5256
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Measurement and Significant Figures
(cont’d)
• An exact number is a number that arises
when you count items or when you define
a unit (conversion 12 in = 1 ft).
– For example, when you say you have nine
coins in a bottle, you mean exactly nine
(9.00000…. - infinite).
– When you say there are twelve inches in a
foot, you mean exactly twelve.
– Note that exact numbers have no effect on
significant figures in a calculation.
HW 1
code: dog
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The Periodic Table
• Metals, Nonmetals, and Metalloids –
generally, left of staircase metals,
staircase metalloids, right of staircase
nonmetals. This is important for
determining bond type, using proper
terminology, and making decisions.
nonmetals
metals
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Chemical Formulas; Molecular and
involves covalent bond – share electrons
Ionic Substances
between atoms – typically nonmetal/nonmetal
involves ionic bond – transfer electrons between
atoms – attraction between charged particles –
+ Na Cl
typically metal/nonmetal or polyatomic ions
C:C
• Molecular substances
– A molecule is a definite group of atoms that are
chemically bonded together through sharing of
electrons (covalent bonding, generally nonmetalnonmetal including H).
– A molecular substance is a substance that is
composed of molecules, all of which are alike.
– A molecular formula gives the exact number of
atoms of elements in a molecule (i.e. C2H6O).
– Structural formulas show how the atoms are bonded
to one another in a molecule.
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i.e. ethanol, CH3CH2OH
• Ionic substances
– Although many substances are molecular, others are
composed of ions (charged particles, transfer of
electrons, ionic bonding, generally metal-nonmetal).
– An ion is an electrically charged particle obtained from
an atom or chemically bonded group of atoms by
adding or removing electrons.
+
Na  1e  Cl
– Sodium chloride is a substance made up of ions.
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Chemical Formulas; Molecular
and Ionic Substances
• Ionic substances
– The formula of an ionic compound is written by giving
the smallest possible whole-number ratio of different
ions in the substance.
– The formula unit of the substance is the group of
atoms or ions explicitly symbolized by its formula.
C:O
Covalent bond (share e-)
Ionic bond (transfer e-/
attraction charged particles
nm –nm
m – nm and charged ions
Molecules
Formula unit
Molecular substance
Ionic substance
Molecular formula
formula
+ -
Na Cl
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• Ionic substances
– When an atom gains extra electrons, it becomes a
negatively charged ion, called an anion (more
electrons than protons).
– An atom that loses electrons becomes a positively
charged ion, called a cation (more protons than
electrons).
– An ionic compound is a compound composed of
cations and anions.
ionic or molecular; formula unit or molecule; ionic or covalent bonds involved?
NaCl
CaBr2
Na2SO4
CO2
ionic substance; formula unit; ionic bond
ionic substance; formula unit; ionic bonds
ionic substance; formula unit; ionic and covalent bonds
molecular substance; molecule; covalent bonds
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Ions in Aqueous Solution
• Many ionic compounds (ionic bond/m-nm)
dissociate into independent ions when
dissolved in water
NaCl (s)  Na+(aq) + Cl-(aq)
Soluble salt
Soluble ionic compounds dissociate 100%
- referred to as strong electrolytes –
breaks into charged particles
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Ions in Aqueous Solution
• Most molecular (covalent bond/nm-nm)
compounds dissolve but do not dissociate
into ions, exception acids.
C6H12O6 (s)  C6H12O6 (aq)
These compounds are referred to as
nonelectrolytes; no charged particles;
soluble but no ions formed.
How would the sodium sulfate dissolve?
Na2SO4 (s)  2Na+(aq) + SO42-(aq)
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Chemical Substances;
Formulas and Names
• Ionic compounds
– Most ionic compounds contain metal and nonmetal
atoms; for example, NaCl.
– You name an ionic compound by giving the name of
the cation followed by the name of the anion.
Sodium chloride, NaCl
Calcium Iodide, CaI2
Potassium Bromide, KBr
– A monatomic ion is an ion formed from a single atom.
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How get charge for ions?
Rules for predicting charges on monatomic ions
– Most of the main group metals form cations with the
charge equal to their group number.
– The charge on a monatomic anion for a nonmetal
equals the group number minus 8.
– Most transition elements form more than one ion, each
with a different charge (exceptions Cd2+, Zn2+, Ag+).
– Other important elements with variable charge
Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+
0
43+ 4+ 3- 2- 1-
1+
2+
varies
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• Rules for naming monatomic ions
– Monatomic cations are named after the element. For
example, Al3+ is called the aluminum ion.
– If there is more than one cation of an element (charge),
a Roman numeral in parentheses denoting the charge
on the ion is used. This often occurs with transition
elements.
Na+ sodium ion
Ca2+ calcium ion
Fe2+ iron (II) ion
Fe3+ iron (III) ion
Older name: higher ox state (charge) – ic, / lower, -ous
Fe3+ ferric ion
Cu+ cuprous ion
Fe2+ ferrous ion
Cu2+ cupric ion
Hg2+ mercuric ion Hg22+ mercurous ion
The names of the monatomic anions use the stem
name of the element followed by the suffix – ide. For
example, Br- is called the bromide ion. Br bromine 17
• The formula of an ionic compound is written by giving
the smallest possible whole-number ratio of different
ions in the substance.
Sodium chloride
Na+ ClNaCl
Iron (III) sulfate
Fe3+ SO42-
Chromium (III) oxide
Cr3+ O2-
Calcium nitrate
Ca2+ NO3-
Ca(NO3)2
Sodium phosphate
Na+ PO43-
Na3PO4
Strontium oxide
Sr2+ O2-
Fe2(SO4)3
Cr2O3
SrO
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Naming Binary Compounds
• NaF
-
sodium fluoride
LiCl
-
lithium chloride
• MgO
-
magnesium oxide
• MnBr2
-
manganese (II) bromide
Co2O3
-
cobalt (III) oxide
CuCl2
-
copper (II) chloride or cupric chloride19
Chemical Substances;
Formulas and Names
• Polyatomic ions
– A polyatomic ion is an ion consisting of two or more
atoms chemically bonded together and carrying a net
electric charge.
– Table in book lists some common polyatomic ions.
Most are oxo anions – consists of oxygen with another
element (central element).
NO3- nitrate
SO42- sulfate
NO2- nitrite
SO32- sulfite
Most groups –ate, -ite
Mn, Br, Cl, I
differ by O
per- -ate, -ate, -ite, hypo- -ite
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Ions You Should Know
Polyatomic ions
• NH4+ - Ammonium
• OH- - Hydroxide
• CN- - Cyanide
• SO42- - Sulfate
• SO32- - Sulfite
• ClO4- - perchlorate
• ClO3- - chlorate
• ClO2- - chlorite
• ClO- - hypochlorite
• Hg22+ - mercury (I) or
mecurous
• S2O32- - thiosulfate
• SCN- - thiocyanate
• CNO- - cyanate
• MnO4- - permanganate
•
•
•
•
•
•
•
•
•
•
•
•
•
•
O22- - Peroxide
PO43- - Phosphate
PO33- - Phosphite
CO32- - Carbonate
HCO3- - Bicarbonate or
Hydrogen Carbonate
N3- - azide
NO3- - nitrate
NO2- - nitrite
C2H3O2- - acetate
Cr2O72- - dichromate
CrO42- - chromate
C2O42- - oxalate
HSO4- - bisulfate or
hydrogen sulfate
H2PO4- - dihydrogen
phosphate
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SnSO4
Na2SO3
tin (II) sulfate or stannous sulfate
sodium sulfite
Ca(ClO)2
calcium hypochlorite
Ba(OH)2
barium hydroxide
KClO4
potassium perchlorate
Cr2(SO4)3
chromium (III) sulfate
Mg3N2
magnesium nitride
Fe3(PO4)2
iron (II) phosphate or ferrous phosphate
Ti(NO3)4
titanium (IV) nitrate
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Chemical Substances;
Formulas and Names
• molecular compounds
– Binary compounds composed of two nonmetals are
usually molecular and are named using a prefix
system (name same as ionic except must indicate
how many atoms are present using mono, di, tri,
etc.). No charges involved with molecular compounds
but we typically put more metallic compound first.
NF3
F3N
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Chemical Substances;
Formulas and Names
• Binary molecular compounds
– The name of the compound has the elements in the
order given in the formula.
– You name the first element using the exact element
name.
– Name the second element by writing the stem name of
the element with the suffix “–ide.”
– If there is more than one atom of any given element,
you add a prefix (di, tri, tetra, penta, hexa, hepta, octa,
etc.)
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• Binary molecular compounds
– N2O3
dinitrogen trioxide
– SF4
sulfur tetrafluoride
– ClO2
chlorine dioxide
– SF6
sulfur hexafluoride
– Cl2O7
dichlorine heptoxide
– HCl (g)
hydrogen chloride
Name this compound but think about bonding:
magnesium chloride; ionic bond no prefix
MgCl2
Older names: water - H2O, ammonia – NH3,
hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4
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Chemical Substances;
Formulas and Names
• Acids
– Acids are traditionally defined as compounds with a
potential H+ as the cation.
– Binary acids consist of a hydrogen ion and any single
anion. For example, HCl (aq) is hydrochloric acid.
– An oxoacid is an acid containing hydrogen, oxygen,
and another element. An example is HNO3, nitric acid.
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oxoacids
Anion prefix/suffix
per- -ate ion
-ate ion
-ite ion
hypo- -ite ion
acid prefix/suffic
per- -ic acid
-ic acid
-ous acid
hypo- -ous acid
NO3- nitrate ion
NO2- nitrite ion
ClO4- perchlorate ion
HNO3 nitric acid
HNO2 nitrous acid
HClO4 perchloric acid
SO42PO43-
sulfate ion
phosphate ion
HW 2
H2SO4
H3PO4
code: cat
sulfuric acid
phosphoric acid
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Molecular Weight and Formula
Weight, Molar Mass
• The molecular weight of a substance is the sum
of the atomic weights of all the atoms in a
molecule of the substance.
– For, example, a molecule of H2O contains 2 hydrogen
atoms (at 1.01 amu each) and 1 oxygen atom (16.00
amu), giving a molecular weight of 18.02 amu.
– Molecular wt – mass one molecule
or
– do for 1 mole of substance called molar mass:
18.02 g H2O/mol H2O
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Working with Solutions
Molar Concentration
• When we dissolve a substance in a
liquid, we call the substance the solute
(being dissolved) and the liquid the
solvent (doing the dissolving).
– The general term concentration refers to the
quantity of solute in a standard quantity of
solution. There are many concentration terms but
we will concentrate on one.
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Working with Solutions
Molar Concentration
• Molar concentration, or molarity (M),
is defined as the moles of solute
dissolved in one liter (cubic decimeter)
of solution.
moles of solute
Molarity (M) 
liters ofsolute
solution
+ solvent volume
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Working with Solutions
Molar Concentration
• The molarity of a solution and its volume are
inversely proportional. Therefore, adding water
makes the solution less concentrated. Most of time
will be using a stock solution and diluting to new
concentration. Basically using
Cc  Vc  Cd  Vd
– So, as water is added, increasing the final volume, Vf,
the final molarity, Mf, decreases. Thing to realize
here is that M x V = mols: want new concentration of
substance take mols and divide by total volume
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• Mixture example
– A solution is prepared by mixing 12.9 mL of 0.245 M HCl and 56.7 mL of
Na2SO4
0.847 M HCl,
then add 630.4 mL of water. Assuming the liquid volumes
are additive, calculate the molarity of HCl in the resulting solution.
mmol = mol
mol x mL = ____
L
tot mL
L
mol
mol
mmols of HCl  (12.9 mL) x (0.245
HCl)  (56.7 mL ) x (0.847
HCl)
L
L
 3.161 mmol HCl  48.025 mmol HCl  51.19 mmol HCl
3.161 mmol
51.19 mmol HCl
M HCl 
 0.00452 M HCl
12.9 mL  56.7 mL  630.4 mL
51.19 mmol HCl
HW 3

 0.0731 M HCl
700.0 mL
code: bird
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Solubility Rules for Ionic Compounds (Dissociates 100%)
1.) All compounds containing alkali metal cations and the ammonium ion
are soluble.
2.) All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are
soluble.
3.) All chlorides, bromides, and iodides are soluble except those containing
Ag+, Pb2+, or Hg22+.
4.) All sulfates are soluble except those containing Hg22+, Pb2+, Ba2+, Sr2+,
or Ca2+. Ag2SO4 is slightly soluble.
5.) All hydroxides are insoluble except compounds of the alkali metals and
Ca2+, Sr2+, and Ba2+ are slightly soluble.
6.) All other compounds containing PO43-, S2-, CO32-, CrO42-, SO32- and
most other anions are insoluble except those that also contain alkali
metals or NH4+.
Generally, compound dissolves
Hg2Cl2 (s)
KI (aq)
Pb(NO3)2 (aq)
insoluble
soluble
soluble
> 0.10 M - soluble (aq)
< 0.01 M - insoluble (s)
in between - slightly soluble
(this class we will assume slightly soluble as soluble)
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Strong Acids (Ionizes 100%)
HCl, HBr, HI, HClO4, HNO3, H2SO4
Strong Bases (Dissociates 100%)
NaOH, KOH, LiOH, Ba(OH)2, Ca(OH)2,
Sr(OH)2
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Ions in Aqueous Solution
Molecular and Ionic Equations
• A molecular/formula unit equation is one in which the
reactants and products are written as if they were
molecules/formula units, even though they may actually
exist in solution as ions.
Calcium hydroxide + sodium carbonate
M.E.
Ca(OH)2 (aq) +
Na2CO3 (aq)  CaCO3 (s) + 2 NaOH (aq)
strong base
soluble salt
insoluble salt
strong base
s solid
l liquid
aq aqueous (acid/bases and soluble salts dissolve in water)
g gases
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Ions in Aqueous Solution
Molecular and Ionic Equations
• An total ionic equation, however, represents strong electrolytes as
separate independent ions. This is a more accurate representation of the
way electrolytes behave in solution.
– A complete ionic equation is a chemical equation in which strong
electrolytes (such as soluble ionic compounds, strong acids/bases) are
written as separate ions in solution. (note: g, l, insoluble salts (s), weak
acid/bases do not break up into ions)
M.E.
Ca(OH)2 (aq) +
strong base
Na2CO3 (aq)  CaCO3 (s) + 2 NaOH (aq)
soluble salt
insoluble salt
strong base
Total ionic
Ca2+ (aq) + 2OH- (aq) + 2Na+ (aq) + CO32- (aq) 
CaCO3 (s)
+ 2Na+ (aq) + 2OH- (aq)
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Net ionic equations.
– A net ionic equation is a chemical equation from
which the spectator ions have been removed.
– A spectator ion is an ion in an ionic equation that
M.E. does not take part in the reaction.
Ca(OH)2 (aq) +
Na2CO3 (aq)  CaCO3 (s) + 2 NaOH (aq)
Total Ionic
Ca2+ (aq) + 2OH- (aq) + 2Na+ (aq) + CO32- (aq)  CaCO3 (s) + 2Na+ (aq) + 2OH- (aq)
Net
Ca2+ (aq) + CO32- (aq)
 CaCO3 (s)
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Types of Chemical Reactions
• Oxidation-Reduction Reactions (Redox rxn)
– Oxidation-reduction reactions involve the
transfer of electrons from one species to another.
– Oxidation is defined as the loss of electrons.
– Reduction is defined as the gain of electrons.
– Oxidation and reduction always occur
simultaneously.
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27.1 Reduction and Oxidation
Redox reactions – transfer of ereduction – oxidation reactions
Reduction – gain of e- / gain of H / lost of O
Fe3+ + 1e-  Fe2+
(lower ox state)
note: must balance atoms and charges
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Oxidation - loss of e- / loss of H / gain of O
Fe2+
 Fe3+ + 1e-
(higher ox state)
Br + 4(-2) = -1
Br = -1 +8 = +7
H2O + BrO3-  BrO4- + 2H+ + 2eBr + 3(-2) = -1
(Br oxidized: charge 5+  7+)
Br = -1 +6 = +5
2H+ + 2e-  H2
(H reduced: charge 1+  0)
Oxidizing agent is species that undergoes reduction.
Reducing agent is species that undergoes oxidation.
Note: need both for reaction to happen; can’t have
something being reduced unless something else is being
oxidized.
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27.3 Balancing Redox Reactions
- Must know charges (oxidation numbers) of species
including polyatomic ions
- Must know strong/weak acids and bases
- Must know the solubility rules
Oxidation Numbers – hypothetical charge assigned to the
atom in order to track electrons; determined by rules.
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Rules to balance redox
1.) Convert to net ionic form if equation is originally in molecular form
(eliminate spectator ions).
2.) Write half reactions.
3.) Balance atoms using H+ / OH- / H2O as needed:
– acidic: H+ / H2O put water on side that needs O or H (comes from
solvent)
– basic: OH- / H2O put water on side that needs H but if there is no H
involved then put OH- on the side that needs the O in a 2:1 ratio
2OH- / H2O balance O with OH, double OH, add 1/2 water to
other side.
4.) Balance charges for half rxn using e-.
5.) Balance transfer/accept number of electron in whole reaction.
6.) Convert equation back to molecular form if necessary (re-apply
spectator ions).
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Zn(s) + AgNO3(aq)  Zn(NO3)2(aq) + Ag(s)
Total ionic:
Zn(s) + Ag+(aq) + NO3-(aq)  Zn2+(aq) + 2NO3-(aq) + Ag(s)
Net ionic:
Zn(s) + Ag+(aq)  Zn2+(aq) + Ag(s)
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Zn(s) + Ag+(aq)  Zn2+(aq) + Ag(s)
Net:
Oxidation:
Zn(s)  Zn2+(aq) + 2e-
Reduction:
[ 1e- + Ag+(aq)  Ag(s) ] 2
Balanced net:
Zn(s) + 2 Ag+(aq)  Zn2+(aq) + 2 Ag(s)
Balanced eq:
Zn(s) + 2 AgNO3(aq)
 Zn(NO3)2(aq) + 2 Ag(s)
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MnO4-(aq) + Fe2+(aq)
Net:
[ Fe2+(aq)  Fe3+(aq) + 1e- ] 5
Ox:
Red:
H+
 Mn2+(aq) + Fe3+(aq)
5e- + 8 H+(aq) + MnO4-(aq)  Mn2+(aq) + 4 H2O(l)
Balanced net:
8 H+(aq) + MnO4-(aq) + 5 Fe2+(aq)  Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l)
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KMnO4(aq) + NaNO2(aq) + HCl(aq)  NaNO3(aq) + MnCl2(aq) + KCl(aq) + H2O(l)
Net:
MnO4-(aq) + NO2-(aq) + H+(aq)  NO3-(aq) +
Mn2+(aq)
+ H2O(l)
H2O(l) + NO2-(aq)  NO3-(aq) + 2 H+(aq) + 2 e- ] 5
Ox:
[
Red:
[ 5 e- + 8 H+(aq) + MnO4-(aq)  Mn2+(aq) + 4 H2O(l) ] 2
Balanced net:
2 MnO4-(aq) + 5 NO2-(aq) + 16 H+(aq) + 5 H2O(l)  2Mn2+(aq) + 8 H2O(l) + 5 NO3-(aq) +10 H+(aq)
2 MnO4-(aq) + 5 NO2-(aq) + 6 H+(aq)  2Mn2+(aq) + 3 H2O(l) + 5 NO3-(aq)
Balanced eq:
2 KMnO4(aq) + 5 NaNO2(aq) + 6 HCl(aq)  2MnCl2(aq)+ 3 H2O(l)+ 5 NaNO3(aq) + 2 KCl
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OH-
Net:
CrI3 (s) + Cl2 (g)  CrO42-(aq) + IO4-(aq) + Cl-(aq)
Ox:
[ 32 OH-(aq) + CrI3(s)  CrO42-(aq) + 3 IO4-(aq) + 16 H2O(l) + 27 e- ] 2
Red:
[ 2 e- + Cl2(g)  2 Cl-(aq) ] 27
Balanced net:
64 OH-(aq) + 2 CrI3(s) + 27 Cl2(g)  2 CrO42-(aq) + 6 IO4-(aq) + 54 Cl-(aq) + 32 H2O(l)
HW 4
code: cow
47