Download Metals

EIT Chemistry Review
F2009
Dr. J.A. Mack
Chemistry is really…
Chem
is
try
To do chemistry, you must:
Download available at:
Practice
www.csus.edu/indiv/m/mackj/
Practice
Or look for Jeffrey Mack under CSUS faculty web pages
Practice
Part 1
2
Where do we begin…
4
The modern periodic table is defined by:
The Periodic Table
Groups (families)
(columns down)
Periods (rows across)
Dmitri Mendeleev (1834 ‐
Dmitri Mendeleev (1834 ‐ 1907)
5
6
1
Ch. 2
Periodic Table: Metallic arrangement
1
IA
1
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
3
IIIB
4
IVB
4
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
Nonmetals
Metals
5
6
Tin is in group 4A (14) in the 5th period. 7
Ch. 2
7
Periodic Table: Metallic arrangement
1
IA
1
Chemical Symbols and Formula:
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
8
17
VIIA
Elements:
O = oxygen
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
4
5
6
7
H = hydrogen
somewhere in between metals and non‐metals
C = carbon
12
IIB
Metalloids Molecules:
H2 = hydrogen
O2 = oxygen
H2O = water
Uh‐
Uh‐Oh! this is confusing…
this is confusing…
Yes it is…
Yes it is…
Get over it and get used to it!
CO2 = carbon dioxide
9
10
2
Which elements do we need to know?
Al
B
Cr
Co
Cu
F
Fe
Au
Pb
Ag
Hg
P
K
Na
S
Start here as a minimum…
As many as possible!
Aluminum
Boron
Chromium
Cobalt
Copper
Fluorine
Iron
Gold
Lead
Silver
Mercury
Phosphorus
Potassium
Sodium
Sulfur
11
12
inert or “noble” gasses:
Atom: The smallest divisible unit of an element
Compound: A substance made of two or more atoms
Ion: A charged atom or molecule
Cation: Positive ion
Anion: Negative ion
NOMENCLATURE
Format for naming chemical compounds using prefixes, suffixes, and other modifications of the names of elements which constitute compounds. Metals form Cations
Cations
13
non‐metals form anions
anions Metalloids can do both
14
3
Ion Charges:
Compounds fall into one of two classes:
+1
+3
‐3 ‐2
‐1
Inorganic Salts
Molecules
metal cation
non‐metal
+
+
non‐metal or
polyatomic anion
non‐metal
(no cations or anions)
+2
Variable
The two use different formalisms for naming…
15
Chemical Nomenclature:
Go to the lab page on my website and download the worksheet!
17
16
Binary Compounds: Metal & non‐Metal
Metal of fixed oxidation (charge) state combined with a non‐metal.
non‐metal takes on “ide” suffix
Examples:
Cation
Anion
Formula
Name
K+
Cl−
KCl
Ca2+
O2‐
CaO
Calcium Oxide
Na+
S2‐
Na2S
Sodium sulfide
Al3+
S2‐
Al2S3
Aluminum sulfide
Potassium chloride
18
4
Metals of variable charge (transition) with a non‐metal
Examples:
Some common polyatomic ions:
modify transition metal name with roman numeral
Cation
Anion
Pb2+
Cl−
Formula
Name
PbCl2
lead (II) chloride
NH4+
H3O+
CO32–
HCO32–
NO2–
NO3–
SO42–
SO32–
PO43–
C2H3O2–
pronounced: lead ‐
lead ‐ two ‐
two ‐ chloride
Pb4+
Cl−
PbCl4
lead (IV) chloride
Fe3+
O2−
Fe2O3
Iron (III) oxide
ammonium
hydronium
carbonate
hydrogen carbonate or bicarbonate
nitrite
nitrate
sulfate
sulfite
phosphate
acetate
19
Ternary Compounds: Those with three different elements
metal of fixed charge with a complex ion
Cation
Anion
Formula
21
Metal of variable charge transition with a complex ion
Cation
Anion
Formula
Name
Name
K+
OH−
KOH
Potassium hydroxide
Ca2+
OH−
Ca(OH)2
Calcium hydroxide
Na+
SO 24−
Na2SO4
Sodium sulfate
Al3+
SO 24−
Al2(SO4)2
Aluminum sulfate
22
Fe3+
NO3−
Fe(NO3)3
Iron (III) nitrate
Fe2+
NO −2
Fe(NO2)2
Iron (II) nitrite
23
5
Examples:
Formula Name:
Non‐
Non‐metal with a non‐
metal with a non‐metal
When non‐metals combine, they form molecules.
They may do so in multiple forms: CO
CO2
Because of this we need to specify the number of each atom by way of a prefix.
BCl3
boron trichloride
SO3
sulfur trioxide
NO
nitrogen monoxide
we don’t write:
1 = mono
5 = penta
2 = di
3 = tri
6 = hexa
4 = tetra
N2O4
7 = hepta
nitrogen monooxide
or mononitrogen monoxide
dinitrogen tetraoxide
24
Writing formulas for acids and Bases
25
Type I Acids: Acids derived from –
–ide anions.
The names for these acids follows the formula:
•An acid
acid is a substance that produces H+ when dissolved in water.
•Certain gaseous molecules become acids when dissolved in water.
•A base
base produces OH− when dissolved in water.
•Bases often are Group I and Group II hydroxide salts.
26
“hydro”
+
the root of the ide anion +
ic “acid”
Anion:
Acid:
Name:
chloride
HCl
hydrochloric acid
fluoride
HF
hydrofluoric acid
27
6
Examples:
H+ and S2‐
Anion:
it takes 2 H+ to cancel one S2‐
Acid:
Name:
(nitrate)
NO3−
HNO3
nitric acid
(sulfate)
SO24 −
H2SO4
sulfuric acid
(acetate)
C2 H 3O2−
HC2H3O2
acetic acid
vinegar
H2S
hydro sulfuric acid
28
29
Common Names:
CO2
carbon dioxide
carbon monoxide
CO
P2O5
diphosphorous pentaoxide
nitrogen trihydride
NH3
(ammonia)
30
H2O
water
ammonia
NH3
CH4
methane
NO
nitric oxide
N2O
nitrous oxide
31
7
It’s time to play…
Practice:
N2SO4
sodium sulfate
barium carbonate
BaCO3
FeO
Iron (II) oxide sodium chloride
NaCl
sodium nitrate
NaNO3
Ca3N2
calcium nitride
zinc phosphide
Zn3P2
Fe(NO3)3
iron (III) nitrate
or ferric nitrate*
NiBr2
nickel (II) bromide
PbS
channel 6
lead (II) sulfide
32
33
The Composition of an Atom: Modern Atomic Theory:
•Atoms are made of protons, neutrons and electrons.
•The nucleus of the atom carries most of the mass.
The atom is mostly
empty space
• It consists of the protons and neutrons surrounded by a cloud of electrons.
The charge on the electron is –1
The charge on the proton is +1
•protons and neutrons in the nucleus.
There is no charge on the neutron
•the number of electrons is equal to the number of protons.
The Atomic Number
Atomic Number or number of protons in the nucleus defines an element.
•electrons in space around the nucleus.
34
35
8
Ch. 2
Atomic Number, Z
Isotopes, Atomic Numbers, and Mass Numbers
Elemental Isotopic Symbols: For a given element “XX”, an An element’s identity is defined by the number of protons in the nucleus: Z
Atomic number (Z) = number of protons in the nucleus.
Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
Atomic number
13
Al
isotope is written by:
One nucleon has a mass of 1 amu
(Atomic Mass Unit
Atomic Mass Unit) a.k.a “Dalton
Dalton” or u
Atom symbol
26.981
Atomic weight
A
Z
X
Isotopes have the same Z but different total number of nucleons (A).
36
Example: Selenium‐75
37
Ch. 2
The average weighted atomic mass is determined by the following mathematical expression:
Average mass of a Cl atom
75 protons + neutrons
mass of a Cl‐35 atom
m Cl (u) = m Cl‐35 ×
75
34 Se
fraction that are Cl‐35
mass of a Cl‐37 atom
abundance + m Cl‐37 ×
of Cl‐35
fraction that are Cl‐37
abundance of Cl‐37
35.45 u = 34.96885u × 0.7553 + 36.96590 × 0.2447
34 protons
(4 sig. fig)
protons = 34
This is the value that is reported on the periodic table.
Note that: 0.7553 + 0.2447 = 1.0000 (100%)
electrons = 34
neutrons = 75‐34 = 41
38
39
9
Avogadro’s Number and the Mole
Ch. 2
The concept of a mole is defined so that we may equate the amount of matter (mass) to the number of particles (mole).
Ch. 2
Avagagro’s Number
Since one mole of 12C has a mass of 12g (exactly), 12g of 12C contains 6.022142 × 1023 C‐12 atoms.
The Standard is based upon the C‐12 isotope.
But carbon exists as 3 isotopes: C‐12, C‐13 &C‐14
The mass of one 12C‐atom is 1.99265 × 10‐23 g.
The average atomic mass of carbon is 12.011 u. The atomic mass of 12C is defined as exactly 12 u.
From this we conclude that 12.011g of carbon contains Therefore: 1u = (the mass of one 12C atom ÷12) = 1.66054 × 10‐24g = 1.66054 × 10‐27 kg
6.022142 × 1023 C‐atoms
Is this a valid assumption?
Yes, since NA is so large, the statistics hold.
40
Ch. 2
Molar Masses
Since we can equate mass (how much matter) with moles (how many particles) we now have a conversion factor
that relates the two.
mols
×
41
So if you have 12.011g of carbon…
you have 6.022×1023 carbon atoms!
Ch. 2
So if you have 39.95g of argon…
you have 6.022×1023 argon atoms!
molar mass (g/mol) = grams
The Molar Mass of a substance is the amount of matter that contains one‐mole or 6.022 × 1023 particles.
if you have a mole of dollar bills…
aka: Avogadro's number (N
Avogadro's number (NA)
you are Bill Gates…
you have 6.022×1023 bucks!
and if you have 6.022×1023 avocados…
The atomic masses on the Periodic Table also represent the molar masses of each element in grams per mole (g/mol)
you have…
42
a “guacamole”
44
10
Question: How many magnesium atoms are there in 150.0 g of Question: magnesium?
Grams, Moles and Molar Mass:
The molar mass of an atom is a conversion factor that relates mass (grams) to moles
moles and vice versa.
(how much matter)
Solution: Use the molar mass of Mg from the periodic table.
(number of atoms)
150.0g Mg ×
grams ÷ molar mass = moles
g ×
1 mol Mg
6.022 x 1023 Mg atoms
×
24.305g Mg
1 mol Mg
mol
= mol
g
= 3.717 × 1024 Mg atoms
Avogadro's number (NA) relates moles numbers of individual particles:
1 mol ×
(4 sig figs) big number as expected!
6.022 × 1023 particles
= 6.022 × 1023 particles
mol
45
Molar Masses (Molecular Weights) of Compounds:
Ch. 2
46
Ch. 2
The molar mass of a molecular compound is the sum of the molar masses of its atoms.
Example:
Ch. 2
The molar mass of an ionic compound is the molar mass of one formula unit of the compound.
For the compound Al2(SO4)3 the molar mass would be:
2 × (26.98 g/mol) The molar mass of CO2 is:
1 × (12.01 g/mol) + 2 × (16.00 g/mol) = 44.01 g/mol
2 Al’s
+ 3 × (32.07 g/mol) 3 S’s
+ 3 × 4 × (16.00g/mol) 3 × 4 O’s
342.17 g/mol
47
48
11
How many oxygen atoms are there in 25.1g of chromium (III) acetate?
The relative amounts of each atom in a molecule or compound can be represented fraction of the whole.
step 1: write the correct chemical formula…
Cr3+ & C2H3O2−
Cr(C2H3O2)3
step 2: calculate the molar mass…
Question: What is the weight % of each element in C2H6?
First determine the molar mass of C2H6:
229.13 g/mol
1 mol of C 2 H 6 has a mass of 30.07 g
step 3: use dimensional analysis to solve the problem…
25.1g Cr(C 2 H3O 2 )3 × 1mol Cr(C 2 H3O 2 )3
229.13g Cr(C2 H 3O 2 )3
6.022 × 1023 O − atoms
×
=
1mol O
3×2 = 6
6mol O
×
1mol Cr(C 2 H 3O 2 )3
3.96 ×
1023
Percent Composition:
Next determine the mass of hydrogen in 1 mol of Next the compound:
1 mol C2 H 6 ×
O‐atoms 3 sf
3 sf
(2×12.01 + 6 ×1.008) g/mol
6 mol H
1.0079 g H
×
= 6.047 g H
1 mol C2 H 6
1 mol H
Ch. 2
49
Chemical Reactions
Now relate the mass of H in one mol of the compound to the molar mass of the compound
6.047 g H ×
Reactants
Since there is only C as the remaining element:
− % H
Ch. 3
A chemical reaction
chemical reaction can be described by a chemical equation
chemical equation.
All chemical equations have three parts:
1
× 100 = 20.11% H
30.07g C2 H 6
% C = 100
50
Æ
Products
When a chemical changes occurs:
Atoms, Compounds or Molecules
= 79.89 % C
The compound C2H6 is 20.11% H & 79.89% C
51
Æ
New Compounds, Molecules or atoms
52
12
How does one know when a reaction has occurred?
Generally, what you end up with looks nothing like what you started with.
Before
Ch. 3
Quantitative Information About Chemical Reactions Mass must be conserved in a chemical reaction.
After
Total mass of reactants
Total mass of products
=
Chemical equations must therefore be balanced for mass.
Numbers of atoms on the reactant side
Numbers of atoms on the product side
=
53
The chemical equation for the formation of water can be visualized as two hydrogen molecules reacting with one oxygen molecule to form two water molecules:
55
Balancing Chemical Reactions: Example
C2H6
+
2 C’s & 6 H’s
→
O2
2 O’s
balance last
CO2
+
H2O
1 C & 2 O’s
2 H’s & 1 O
balance H first
2H2 + O2 → 2H2O
2 2H6
___C
+
→
O2
CO2
3 2O
6
___ H
+
balance C next
2C2H6 +
O2
→
4 2
___ CO
+
6H2O
balance O
2C2H6 +
11.2g
+ 88.8g → 100.0g
mass is conserved!
56
4 C’s
7 2
____ O
12 H’s 14 O’s
→
4CO2 +
4 C’s
6H2O
12 H’s 14 O’s
57
13
When writing chemical reactions one starts with:
Reactants
Ch. 3
Reactions in Aqueous Solutions
Ch. 3
products
Aqueous Solutions: Water as the solvent
N2(g) + 3H2(g)
2NH3(g)
Solution =
solute
+ solvent
Some reactions can also run in reverse:
2NH3(g)
That which is dissolved (lesser amount)
N2(g) + 3H2(g)
That which is dissolves (greater amount)
There are three types of aqueous solutions:
Under these conditions, the reaction can be written:
Those with Strong Electrolytes
Strong Electrolytes Those with Weak Electrolytes
Weak Electrolytes
3H2 (g) + N2 (g) R 2NH3 (g)
& those with non
non‐‐Electrolytes
Dynamic Equilibrium!
58
Species in Solution, Electrolytes:
Strong electrolytes:
Ch. 3
59
It’s time to play…
What’s in the beaker !!!
Characterized by ions only (cations & anions) in solution (water).
Conduct electricity well
In da
In da beaker…
beaker…
Weak electrolytes:
Characterized by ions (cations & anions) & molecules in solution.
ions only in solution!
Conduct electricity poorly
Na+
Non‐electrolytes:
PO43‐
Characterized by molecules in solution.
Na+
Na+
Na3PO4(aq) → 3Na+(aq) + PO43‐(aq)
Strong electrolyte
Do not conduct electricity
60
61
14
How do we know if a compound will be soluble in water?
It’s time to play…
For molecular compounds
molecular compounds, the molecule must be polar
polar.
What’s in the beaker !!!
We will discuss polarity later, for now I will tell you whether or not a molecular compound is polar…
In da beaker…
molecules only in solution!
H2S(aq)
H2S
weak
electrolyte
→ 2H+(aq) + S2‐(aq)
Two bears, one from Yosemite and one from the Arctic jump in a pond. Because he’
Because he’s…
Which one dissolves?
The one from the arctic of course…
arctic of course…
right?
Nope!
joke
a “
a “polar”
polar” bear!
Ch. 3
62
How do we know if a compound will be soluble in water?
For molecular compounds
molecular compounds, the molecule must be polar
polar.
63
The solubility of salts (metal / non‐
metal compounds) is determined by the solubility rules
solubility rules
We will discuss polarity later, for now I will tell you whether or not a molecular compound is polar…
For ionic compounds
ionic compounds, the compound solubility is governed by a set of SOLUBILITY RULES!
•You must learn the basic rules on your own!!!
•Please download the table on my website for reference.
•I will provide a limited solubility guideline for exams and quizzes.
I will provide a limited solubility guideline for exams and quizzes.
Ch. 3
64
Ch. 3
65
15
Solubility of Inorganic Compounds: The easy way!
Ch. 3
You can memorize all of the rules and the exceptions therein, or
you can remember two easy rules! (Your choice...)
Ch. 3
Types of Reactions in Solution:
Precipitation Reactions: A reaction where an insoluble solid (precipitate) forms and drops out of the solution
Rule 1: Compounds containing one of the following cations are likely soluble:
Group 1A cations (Li+, Na+, K+, Rb+, Cs+)
Rule 2:
Acid–
Acid–base Neutralization: A reaction in which an acid reacts with a base to yield water plus a salt.
Ammonium ion (NH4+)
Compounds containing one of the following anions are likely soluble:
Nitrate (NO3–), chlorate (ClO3–), perchlorate (ClO4–) or acetate (C2H3O2–)
Between these two rules, one can identify 90–95% of all soluble salts.
Gas forming Reactions: A reaction where an insoluble gas is formed.
Reduction and Oxidation Reactions (
RedOx): A reaction Reduction and Oxidation Reactions (RedOx)
where electrons are transferred from one reactant to another.
66
Ch. 3
Precipitation Reactions:
Exchange (Metathesis) Reactions
67
Ch. 3
Consider the following reaction:
Aqueous solutions
2KI(aq)
& Pb(NO3)2(aq) are mixed
Metathesis reactions involve swapping ions in solution:
AX + BY
→ AY + BX.
Metathesis reactions will lead to a precipitate if one of the products is an insoluble solid.
A reaction occurs:
KBr(aq) + AgNO
AgBr(s))
KBr(aq) + AgNO3(aq) →
(aq) → KNO3 (aq) + aq) + AgBr(s
A yellow precipitate forms
68
69
16
Strong acids are almost completely ionized in water. (strong electrolytes)
Reactions of Acids & Bases: Acid‐
Reactions of Acids & Bases: Acid‐Base Neutralization
Examples:
Acid + Base →
HX (aq) (X = Cl, Br & I)
hydro ___ ic acid
HNO3 (aq)
nitric acid
HClO4 (aq)
perchloric acid
H2SO4 (aq)*
sulfuric acid
* Only the 1st
Ch. 3
Salt + Water (usually)
HA (aq) + MOH(aq) →
MA(aq) + HOH(l)
Strong acid ‐ Strong base neutralization:
HBr(aq)/KOH(aq)
Molecular Equation:
HBr(aq) + KOH(aq) →
KBr (aq) + H2O(l)
Total Ionic Equation:
H is strong, sulfuric acid dissociates via:
H2SO4 (aq) →
H+
(aq) + HSO4–
/
/
H+ (aq) + Br– (aq) + K+(aq) + OH– (aq) →
(aq)
/
/
K+(aq) + Br– (aq) + H2O(l)
Net Ionic equation:
H+ (aq) + OH– (aq) → H2O (l)
Ch. 3
70
Gas Forming Reactions:
Ch. 3
Metal carbonate salts react with acids to the corresponding metal salt, water and carbon dioxide gas.
2HCl(aq) + CaCO3(s)
→ CaCl2(aq) + H2CO3(aq)
71
Gas Forming Reactions:
Ch. 3
Metals & water:
Gr I metals: Na, K, Cs etc.. react vigorously with water
2K(s) + 2H2O(l) →
2KOH(aq)
+ H2(g)
decomposes
H2O(l) + CO2(g) Similarly:
HCl(aq) + NaHCO3(s) →
acid
base
Metals & acid:
Some metals react vigorously with acid solutions:
NaCl(aq) + H2O(l) + CO2(g) salt
water
Zn(s) + 2H+(aq) →
Zn2+(aq)
+ H2(g)
Neutralization!!!
74
75
17
Gas Forming Reactions: More examples
Ch. 3
Oxidation & Reduction Reactions:
Ch. 3
Oxidation involves an atom or compound losing electrons
Reduction involves an atom or substance gaining electrons
CO2(g)
Niether process can occur alone… that is, there must be an exchange of electrons in the process.
The substance that is oxidized
oxidized is the reducing agent
reducing agent
H2S(g)
The substance that is reduced
reduced is the oxidizing agent
oxidizing agent
SO2(g)
oxidized reduced
Mg(s) + 2H+(aq) →
NH3(g)
9/18/08
Chem. 1A F2008 J.Mack
Oxidation & Reduction Reactions:
76 76
Ch. 3
2Ag + (aq) + Cu(s) → 2Ag(s) + Cu2+ (aq)
reducing agent
Mg2+(aq) + H2(g)
oxidizing agent
9/19/08
9/18/08
Chem. 1A F2008 J.Mack
77 77
Oxidation & Reduction Reactions:
Ch. 3
Electron Transfer in a RedOx Reaction:
e−
2Ag+(aq)
e−
+ Cu(s) →
Cu2+(aq)
+ 2Ag(s) Two electrons leave copper. The silver ions accept them.
The copper metal is oxidized to copper (II) ion.
The silver ion is reduced to solid silver metal.
78
80
18
Oxidation & Reduction Reactions: Oxidation Numbers
Oxidation Numbers
Chemists use oxidation numbers
oxidation numbers to account for the transfer of electrons in a RedOx reaction.
Oxidation numbers are the actual on an atom or the apparent charge on atom when combined in a compound.
Chemists use oxidation numbers
oxidation numbers to account for the transfer of electrons in a RedOx reaction.
Oxidation numbers are the actual on an atom or the apparent charge on atom when combined in a compound.
1.
2.
3.
4.
The atoms of pure elements always have an oxidation number of zero.
If an atom is charged, then the charge is the oxidation number.
In a compound, fluorine always has an oxidation number of −1.
Oxygen most often has an oxidation number of −2.
• When combined with fluorine, oxygen has a positive O.N.
• In peroxide, the O.N. is −1.
5. In compounds, Cl, Br & I are generally −1 so long as F and O are not 1. The atoms of pure elements always have an oxidation number of
zero.
present.
6. In compounds, H is +1, except as a hydride (H −: −1)
7. For neutral compounds, the sum of the O.N.’s equals zero. For a Ch. 3
poly atomic ion, the sum equals the charge.
I2(s)
Examples:
Ch. 3
Mg(s)
Hg(l)
All have an oxidation number of zero (0)
O2(g)
81
Oxidation Numbers
Ch. 3
82
Most common oxidation numbers
oxidation numbers
+1
+2
Ch. 3
+3
‐3 ‐2 ‐1
2. If an atom is charged, then the charge is the oxidation numbers
oxidation numbers .
Examples:
Ion
Oxidation Number
Mg2+(aq)
+2
Cl−(aq)
−1
Sn4+(s)
+4
Variable (to be deduced)
Metals take on their formal charge:
Non‐metals do the same
9/19/08
9/18/08
Chem. 1A F2008 J.Mack
83 83
84
19
Oxidation Numbers
Ch. 3
magnesium chloride +2
? +
+ 2 × (−1) = 0
3×(−1) =
0
Fe(OH)3
MgCl2
−3
Ch. 3
Determine the oxidation numbers
oxidation numbers of iron in the following compound:
7. For neutral compounds, the sum of the O.N.’s equals zero.
For a poly atomic ion, the sum equals the charge.
Example:
Oxidation Numbers
+ 4 × (+1) = +1
Iron must have an oxidation number of +3!
+
4
NH
86
Oxidation Numbers
Ch. 3
Determine the oxidation numbers
oxidation numbers of manganese in the following compound:
+1 + ? + 4×(−2)
9/19/08
9/18/08
Chem. 1A F2008 J.Mack
Oxidation & Reduction Reactions:
87 87
Ch. 3
Recognizing a RedOx Reaction:
In a RedOx
RedOx reaction, the species oxidized and the species reduced are identified by the changes in oxidation numbers
oxidation numbers :
= 0
KMnO4
Oxidation numbers:
+1
0
2Ag + (aq) + Cu(s) → 2Ag(s) + Cu2+ (aq)
0
+2
Manganese must have an oxidation number of +7!
Oxidation numbers:
Since silver goes from +1 to zero, it is reduced
reduced.
Since copper goes from zero to +2, it is oxidized
oxidized.
88
89
20