Chapter 4 Aqueous Reactions/Solution Chemistry
Chapter 4 Aqueous Reactions/Solution Chemistry
There are many different ways to express concentration quantitatively. Common way to express concentration:
• What is a solution ? – Homogeneous mixture of two or more substances. – Composed of a solvent with solutes dissolved. • Solvent: Substance that solutes are dissolved in. • Solute: Substance dissolved in the solvent to give a solution. • Solutes: Can be a gas (HCl, CO2), liquid, or solid.
Molarity ( M ) =
moles solute solution volume (liters)
(molar concentration) Example: Prepare a HNO3 solution by dissolving 0.0375 mol HNO3 in water to yield a 0.250 L solution. What is the molarity of this solution?
Solution Preparation There can be numerous solutes dissolved in the solvent to give a complex solution. We will consider water as the solvent Î
aqueous solutions. l ti n
Concentration: Amount of solute dissolved in a given quantity of solvent or solution.
Qualitative Description: – Concentrated solution: large amount of solute. – Dilute solution: Only a small amount of solute.
Quantitative Description: – Exact measurement of the amount of solute in a given quantity of solvent or solution.
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QUANTITATIVE DESCRIPTION OF AQUEOUS SOLUTIONS
mL solution Î L solution:
(200 mL ) ⎛⎜
1L ⎞ ⎟ = 0.200 L ⎝ 1000 mL ⎠
Our concentration is calculated as:
0.0375 mol HNO 3 = 0.150 M HNO 3 0.250 L solution Conversion factor Very useful as a ratio in calculations:
0.150 mol HNO 3 1.00 L solution or 1.00 L solution 0.150 mol HNO 3
Example Calculations 1) How many mL of a 0.150 M HNO3 solution is required to give 0.0450 mol HNO3? Set up the problem: given i
X
conversion i factor f t
=
M=
0.0476 mol HNO 3 = 0.238 0 238 M HNO 3 0.200 L
We can combine these conversions into one series ⎛ 3.00 g HNO 3 ⎞⎛ 1 mol HNO 3 ⎞⎛ 1000 mL ⎞ ⎟⎜⎜ ⎜ ⎟ = 0.238 M HNO 3 ⎟⎟⎜ ⎝ 200 mL ⎠⎝ 63.02 g HNO 3 ⎠⎝ 1 L ⎠
3) How many moles of HNO3 are in 50.0 mL of 0.238 M HNO3?
Conversion: Volume Î moles 1 L ⎞⎛ 0.238 mol HNO 3 ⎞ ⎟ = 0.0119 mol HNO 3 ⎟⎜ 1L ⎝ 1000 mL ⎠⎝ ⎠
(50 mL )⎛⎜
desired d i d
⎛
⎞ 1L ⎟⎟ = 0.300 L 0.150 mol HNO ⎝ 3⎠
(0.0450 mol HNO 3 ) ⎜⎜
300 mL of 0.150 M HNO3 contains 0.0450 mol HNO3
2) What is the concentration of a HNO3 solution prepared by dissolving 3.00 g HNO3 to prepare a 200 mL solution? g HNO3 Î mol HNO3
Conversions:
Calculate Molarity (moles solute/liter solution):
mL solution Î L solution
Conversion factor Alternative calculation ⎛ 0.238 mol HNO 3 ⎞⎛ 1 L ⎞ ⎟⎜ ⎜ ⎟(50 mL ) = 0.0119 mol HNO 3 1L ⎝ ⎠⎝ 1000 mL ⎠
4) How many grams of HNO3 is needed to
prepare a 250 mL solution that has a concentration of 0.100 M HNO3 ? The concentration required is:
0.100 M =
Solve: g HNO3 Î mol HNO3
⎛ 1 mol HNO 3 ⎞ ⎟⎟ = 0.0476 mol HNO 3 ⎝ 63.02 g HNO 3 ⎠
(3.00 g HNO 3 ) ⎜⎜
0.100 mol HNO 3 1 L solution
Two steps in calculation: First: Calculate mol HNO3 in 250 mL Second: Convert mol HNO3 to g HNO3
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Calculate mol HNO3 in 250 mL: ⎛ 1 L ⎞⎛ 0.100 mol HNO 3 ⎞ (250 mL)⎜ ⎟ = 0.0250 mol HNO 3 ⎟⎜ 1L ⎝ 1000 mL ⎠⎝ ⎠
Convert mol HNO3 to g HNO3:
⎛ 63.02 g HNO 3 ⎞ ⎟⎟ = 1.58 g HNO 3 ⎝ 1 mol HNO 3 ⎠
(0.0250 mol HNO 3 )⎜⎜
Figure 4.19 Only added more solvent upon dilution, same amount (moles) of solute.
DILUTION OF SOLUTIONS • Often, a solution of known concentration must be diluted to prepare a solution with a lower desired concentration.
• For any solution:
DILUTION OF SOLUTIONS
M i Vi = M f V f What volume of a 0.150 M HNO3 solution is required to prepare 250 mL of a 0.0100 M HNO3 solution? M i Vi = M f V f
Vi =
MfVf Mi
molarity X volume (L) = moles solute (0.0100 M )(250 mL) = 16.7 mL 0.150 M Take 16.7 mL of 0.150 M HNO3 solution and dilute to a final volume of 250 mL with the solvent. Vi =
mol solute X L solution = moles solute L solution
DILUTION OF SOLUTIONS • Take a given amount of solute (volume of solution) and add solvent to give a new concentration. • Upon dilution, moles solute is constant, the volume has simply increased. M i X Vi = moles solute
M f X V f = moles solute
Pure water: A very poor conductor of electricity (does not conduct electricity). m solutes, u , when w n dissolved in n w water,, y yield an n Some
electrically conducting solution.
Initial
Final (diluted)
Electrical Conductivity of Aqueous Solutions
equal moles solute
These solutes are called electrolytes.
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Other solutes, when dissolved in water, yield a solution that does not conduct electricity (non electrically conducting solution).
Electrolytes Î STRONG OR WEAK Strong electrolytes: Dissolve in the solvent exclusively as dissociated ions.
These solutes are called nonelectrolytes.
Examples of strong electrolytes:
Ionic compounds Î electrolytes Molecular Î generally nonelectrolytes compounds
All ionic i i compounds d CaCl2 NaNO3 KClO3 Pb(NO3)2 Each of these compounds dissociates in water to yield ions exclusively.
Dissolving of an Ionic Compound
Pure water
nonelectrolyte solution
electrolyte solution
Electrical Conductivity of Aqueous Solutions Svante Arrhenius: 1884; postulated that electrolytic solutions are composed of dissolved ions (Nobel prize (Chemistry) - 1903)
NaCl (s)
water
Na+(aq) + Cl-(aq)
Dissociation of Strong Electrolytes
CaCl2
water
NaNO3 KClO3 Pb(NO3)2
Ca2+ + 2Cl-
water
Na+ + NO3-
water
K+ + ClO3-
water
Pb2+ + 2NO3-
Ions that allow for the conduction of electricity in aqueous solutions.
Note: Polyatomic ions stay intact when dissolved (do not break apart further).
Consequently, electrolytes dissociate to ions upon dissolving.
Note: There are no intact ionic compounds (M+X-) dissolved in the solution; only dissociated ions.
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Weak Electrolytes • Only partially dissociate to ions for the dissolved species. • Solutions poor conductors of electricity. • Some molecular compounds are weak electrolytes; e.g., ammonia (NH3) and acetic acid (HC2H3O2). Most , however, are nonelectrolytes. • Molecular compounds that are acids or bases are electrolytes (dissociate to ions in solution).
Complete dissociation to ions, hence, HCl is a molecular compound that is a strong electrolyte. Contrast this with the weak acid, hydrofluoric acid, which is a weak electrolyte:
water
HF
H+ + F -
There is only partial dissociation to ions. Both molecular HF and the dissociated ions, H+ and F-, exist in solution.
ACIDS Arrhenius Definition (1884): An acid is a substance that produces H+ ions (H3O+) when dissolved in water. Very y limited definition,, restricted to aqueous q solutions. Brønsted Definition (1932): An acid is a hydrogen ion, H+ (proton), donor.
HF is, therefore, a weak electrolyte. In an HF solution, HF exists mostly as undissociated HF molecules. O l a small Only ll amountt dissociates di i t to t ions i (ionized). (i i d) Acids can be either strong or weak electrolytes.
Broader definition of acids; does not require the acid to be dissolved in water.
ACIDS
CHARACTERISTICS OF ACIDS
Example: HCl H3O+ + Cl-
HCl + H2O
C Common I Inorganic i A Acids: id
Short-hand representation:
HCl
water
1) Sour (tart) taste in solution (citrus drinks have citric acid). 2) Turn litmus paper red. 3) React with many metals to release H2.
+
-
H + Cl
Ionization reaction: A reaction that forms ions from the dissociation of molecules.
HCl HNO3 HF
water
H+ + Cl-
water
H+ + NO3-
water
H+ + F-
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Triprotic Acid Common Organic Acid:
Three acidic Hydrogen atoms H3PO4 (Phosphoric Acid)
Acetic Acid – HC2H3O2 water
HC2H3O2
H+ + C2H3O2acetate ion
All previous acids are monoprotic; i.e., one acidic hydrogen.
H3PO4
H+ + H2PO4-
H2PO4-
H+ + HPO42-
HPO42-
H+ + PO43-
BASES
Diprotic Acids Acids with two acidic hydrogen atoms. Examples: H2SO4 H2SO3 H2CO3 H2C2O4
sulfuric acid sulfurous acid carbonic acid oxalic acid (organic acid)
What happens in solution with a diprotic acid ?
Arrhenius Definition (1884): A base is a substance that produces OH- ions when dissolved in water. Very limited definition, restricted to aqueous solutions. Brønsted Definition (1932): A base is a hydrogen ion, H+ (proton), acceptor. Broader definition of bases; does not require the base to be dissolved in water.
Diprotic Acids in Solution
BASES
Sulfurous Acid (H2SO3)
H2SO3
H+ + HSO3- (hydrogen sulfite ion)
HSO3-
H+ + SO32- (sulfite ion)
H+ + HCO3- (hydrogen carbonate ion)
HCO3-
1) Bitter taste in solution 2) Solutions are often slippery (soapy feel) 3) Turn litmus paper blue Two Categories:
Carbonic Acid (H2CO3) H2CO3
Characteristics of bases:
H+ + CO32- (carbonate ion)
1) Ionic compounds that contain hydroxide ions (OH-). 2) Molecular compounds that react with water to yield hydroxide ions (OH-).
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Ionic Hydroxides:
NaOH
Seven strong acids: Know these!!!!
water
Na+ + OH-
water
Ca2+ + 2OH-
Ca(OH)2
The hydroxide ion is a hydrogen ion (H+, proton) acceptor, therefore, it is a base by our definition. The two ions combine to form the weak electrolyte, water.
H+ + OH-
H2O
Strong Bases: Know these!!!! Group 1A metal hydroxides:
Bases that do not contain OH- ions.
NH3 + H2O
NH4+ + OH-
off of water to yield
– LiOH, NaOH, KOH, RbOH, CsOH
Heavier Group 2A metal hydroxides: – Ca(OH)2, Sr(OH)2, Ba(OH)2 – Exceptions: (Be(OH)2 and Mg(OH)2 are not strong bases)
Hydrogen ion (proton) acceptor. Ammonia pulls
HCl - hydrochloric acid HBr - hydrobromic acid HI - hydroiodic acid HNO3 - nitric acid HClO3 - chloric acid H HClO4 - perchloric acid H2SO4 - sulfuric acid
All other acids are weak, e.g., HC2H3O2 - acetic acid HF - hydrofluoric acid
Molecular Bases:
H+
• • • • • • •
OH-.
ACIDS AND BASES
Strong acids and bases Î Strong Electrolytes Weak acids and bases Î Weak Electrolytes
Acid/Base Definitions:
Classified as either strong or weak.
Arrhenius Definition:
Strong acids and Strong bases: Completely dissociate to ions in solution. Hence they are also strong electrolytes.
Acid: A substance that increases the H3O+ concentration in water.
Weak acids and Weak bases: Only undergo partial dissociation to ions in solution. Hence, they are also weak electrolytes.
Base: A substance that increases the OH- concentration is water.
Most acids and bases are weak.
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Brønsted definition: Acid: Hydrogen ion (H+, proton) donor. Base: Hydrogen ion (H+, proton) acceptor. HCl + H2O
H3O+ + Cl-
acid
NH3 + H2O
NH4+ + OH-
base
What happens when an acid and a base react?
acid + base
salt + H2O
Ionic compounds Î dissolve in water dissociating to ions How much of the ionic compound can dissolve in a given quantity of water? Î solubility Solubility: The maximum amount of solute that can dissolve l in a given quantity of solvent l at a specific temperature Î saturated solution. Solubility expressed as:
mol solute grams solute or liter solution 100 g solvent
Solubility of Ionic Compounds Compound NaCl
Acid and base neutralize (destroy) each other. This is called a neutralization reaction.
• Salt: Ionic compound whose cation comes from a base and whose anion comes from an acid. – cations other than H+ – anions other than OH- or O2-
HCl + NaOH acid
base
NaCl + H2O salt
Solubility (g/100 g water) 35.7 g @ 0 ºC 39.1 g @ 100 ºC
CuS (copper(II) sulfide)
3.3 X 10-5 g @ 18 ºC
CaCl2
74.5 g @ 20 ºC
CaCO3
1.5 X 10-3 g @ 25 ºC
All ionic compounds are soluble to some extent in water. However, many have very low solubilities ( 0.1 0 1 g/100 g water) or insoluble (solubility < 0.1 g/100 g water).
NaCl and CaCl2 are soluble CuS and CaCO3 are insoluble
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General rules for predicting the solubility of ionic compounds. Summarized in Table 4.2, page 123; KNOW THIS TABLE!!!
Solubility Guidelines for Ionic Compounds Soluble Compounds All alkali metal ions ions, NH4+
Important Exceptions None
NO3-, C2H3O2-, HCO3-, ClO3-, ClO4Cl-, Br-, I-
None
SO42-
Strong Acid
H+,
Cl-
H2O
In solution solution:
Salts of Ag+, Hg22+, Pb2+
water
H+(aq) + Cl-(aq)
H+,
Cl-,
H2O
Salts of Ca2+, Sr2+, Ba2+, Hg22+, Pb2+, Ag+
Insoluble Compounds CO32-, PO43-, SO32-, S2-, CrO42-
Important Exceptions Salts of NH4+, all alkali metal ions
OH-
NH4+, all alkali metal ions ions, Ca2+, Sr2+, Ba2+
(acid oxides)
HCl solution:
HCl
Solubility Guidelines for Ionic Compounds
O2-
What are the main species in solution??
All alkali metal ions, Sr2+, Ba2+
Solubility Examples: Na2CO3
Soluble (Na+)
CaCO3
Insoluble (CO32-)
AgNO3
Soluble (NO3-)
AgCll
Insoluble l bl ((Ag+)
PbSO4
Insoluble (Pb2+)
CaSO4
Insoluble (Ca2+)
Ca2+,
What are the main species in solution?? HF solution:
Weak Acid
H+, F-
In solution solution:
HF, H2O
HF
water
HF,
H+(aq) + F-(aq)
H+,
F-,
H2O
What are the main species in solution?? NaOH solution:
Strong Base
Na+,
H2O
In solution solution:
NaOH
water
Na+,
OH-
Na+(aq) + OH-(aq)
OH-,
H2O
Apply the solubility rules to determine if an ionic compound is soluble or insoluble.
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What are the main species in solution??
Molecular Equation:
HCl(aq) + NaOH(aq)
NH3 solution:
NaCl(aq) + H2O(l)
Weak Base
NH4+, OH-
Ionic Equation:
In solution solution:
NH3, H2O
Strong g electrolytes y written as dissociated ions.
water
NH3 + H2O NH3,
NH4+(aq) + OH-(aq)
NH4+,
OH-,
H2O
What are the main species in solution??
Soluble Ionic Compound
Mg2+,
Cl-
H2O
Mg2+,
Net Ionic Equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l) Na+ and Cl- do not react, spectator ions.
In solution: water
Na+(aq) + Cl-(aq) + H2O(l)
Eliminate spectator ions form the ionic equation.
MgCl2 solution:
MgCl2
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)
Mg2+(aq) + 2Cl-(aq)
Cl-,
H2O
Reactions of Ions in Solution 3 ways to represent reactions in solution 1) Molecular equation: Each substance written as intact units (molecules). 2) Ionic equation: Strong electrolytes written as dissociated ions (how they exist in solution). 3) Net ionic equation: Only the species that actually take part in the reaction are shown.
Net Ionic Equation: H+(aq) + OH-(aq)
H2O(l)
A neutralization reaction
Balanced Ionic Reactions: 1) Need atom balance 2) Net charge balance reactant charge = product charge
Molecular Equation:
HC2H3O2(aq) + NaOH(aq) NaC2H3O2(aq) + H2O(l)
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Ionic Equation:
Molecular Equation:
HC2H3O2(aq) + Na+(aq) + OH-(aq) weak electrolyte
CaCl2(aq) + Na2CO3(aq) CaCO3(s) + 2 NaCl(aq)
Na+(aq) + C2H3O2-(aq) + H2O(l)
AB + CD Net Ionic Equation (eliminate spectator ions):
HC2H3O2(aq) + OH-(aq) weak electrolyte
C2H3O2-(aq) + H2O(l)
CaCl2
+
Soln 1
Reactions yielding a precipitate are a common type of ionic reaction. Precipitate formation drives the reaction; ions are removed from the solution.
Precipitation Reaction
Reactions between ionic substancessoluble ionic compounds.
CaCl2(aq) + Na2CO3(aq)
AD + CB
???
Na2CO3
Î
Soln 2
mix the two solutions
Ca2+(aq) + 2Cl -(aq) + 2Na+(aq) + CO32- (aq) What are the products ?
2KI(aq) + Pb(NO3)2(aq)
PbI2(s) + 2K+(aq) + 2NO3-(aq)
Ca2+(aq) + 2Cl -(aq) + 2Na+(aq) + CO32- (aq) CaCO3(s) + 2Na+(aq) + 2Cl -(aq) precipitate The reaction is called a metathesis reaction
(double displacement reaction).
AB + CD
AD + CB
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2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2K+(aq) + 2NO3-(aq) Net Ionic Equation:
Reactions of insoluble salts:
MgCO3(s) + 2 HCl(aq) MgCl2(aq) + H2CO3(aq)
2I-(aq) + Pb2+(aq)
PbI2(s)
There are many examples Th l of f precipitation i it ti reactions-need to know the solubility rules and how to apply them.
Remember, if you are not part of the solution…… then you are part of the precipitate.
Some ionic reactions proceed to yield a weak electrolyte. Molecular: HCl(aq) + NaC2H3O2(aq)
NaCl(aq) + HC2H3O2(aq)
Ionic: H+(aq) + Cl-(aq) + Na+(aq) + C2H3O2-(aq) +
H2CO3(aq)
CO2(g) + H2O(l)
Baking: Na2CO3 baking powder CO2(g) NaHCO3 baking soda during baking (raises cake)
TITRATIONS • Use chemical reactions to determine concentration of a solution. • Use a solution of known concentration (standard solution) to determine the concentration of another solution of unknown concentration. • Reactions between the two solutions. • A common titration is an acid/base titration
-
Na (aq) + Cl (aq) + HC2H3O2(aq) weak electrolyte
Simple Titration
Net Ionic: H+(aq) + C2H3O2-(aq)
HC2H3O2(aq) weak electrolyte
Weak Electrolytes: Consist of weak acids and bases: HC2H3O2, NH3, HF, HNO2
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HCl(aq) + NaOH(aq)
Titration
NaCl(aq) + H2O(l)
neutralization reaction 1 mol HCl (H+) reacts with 1 mol NaOH (OH-)
Example titration: A 20.0 mL HCl solution
required 37 37.8 8 mL of 0 0.100 100 M NaOH to reach the endpoint in the titration. What is the concentration of the HCl solution ? At the endpoint (equivalence point):
mol NaOH (OH-) = mol HCl (H+)
First: Calculate mol NaOH used in the titration. Second: Calculate concentration of HCl. Note that mol HCl = mol NaOH Concentration X Volume = number of moles
Oxidation-Reduction Reactions The following is a common teaching laboratory reaction:
Zn(s) (s) + 2 HCl(aq) C (aq)
(Molarity)
Mol NaOH:
⎛ 0.100 mol NaOH ⎞⎛ 1 L ⎞( ⎜ ⎟⎜ ⎟ 37.8 mL ) 1L ⎝ ⎠⎝ 1000 mL ⎠ = 0.00378 mol NaOH
Second, calculate concentration of HCl At the endpoint, mol HCl (H+) = mol NaOH (OH-): mol NaOH = mol HCl = 0.00378 mol
The concentration of HCl (Molarity) is: MOLARITY =
moles solute liter of solution
⎛ 0.00378 mol HCl ⎞⎛ 1000 mL ⎞ ⎜ ⎟⎜ ⎟ = 0.189 M HCl 20 mL ⎝ ⎠⎝ 1 L ⎠
Use acid/base titrations to determine concentration
ZnCl C 2(aq) + H2(g)
Ionic Equation:
Zn(s) + 2 H+(aq) + 2 Cl-(aq) Zn2+(aq) + 2 Cl-(aq) + H2(g)
Net Ionic Equation:
Zn(s) + 2 H+(aq) + 2 Cl-(aq) Zn2+(aq) + 2 Cl-(aq) + H2(g) Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
Zn has lost 2 electrons !!!! 2 H+ have gained 2 electrons !!!! This reaction involves the transfer of electrons from Zn metal to hydrogen ions (H+).
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CORROSION OF METALS BY ACIDS Corrosion: The conversion of the elemental form of a metal (native metal) to a metal compound by reaction with a substance in its environment. Metals tend to lose electrons to form cations:
Fe(s)
Fe3+ + 3 e-
4 Fe(s) + 3 O2(g)
2 Fe2O3(s)
4 Fe3+ + 12 e- oxidation
4 Fe(s)
3 O2 + 12 e-
6 O2-
reduction
Redox reactions are a very important class of reactions. Many reactions proceed with the transfer of electrons from one species to another. Electrons are central to chemistry!!!!
This often involves reaction with oxygen:
4 Fe(s) + 3 O2(g)
2 Fe2O3(s)
Since most metals lose electrons by reaction with O2, the metal is said to be oxidized.
Loss of electrons is called oxidation. oxidation Since we must maintain electrical neutrality, some other species gains the electrons that are lost by the metal. Oxygen gains the electrons.
O2 + 4 e-
Oxidation of metals by acids: Zn loses 2 electrons (oxidized)
Zn(s) (s) + 2 HCl(aq) C (aq)
ZnCl C 2(aq) + H2(g)
2 H+ gain 2 electrons (reduced)
2 O2-
• When a species gains electrons (becomes more negatively charged) it is said to be reduced. • A gain in e-’s by a species is called reduction. • If f one species gains i electrons, l t another h must
lose electrons.
• Oxidation and reduction reactions are coupled together. • These types of reactions are called
Oxidation of metals by acids: Metals vary in their tendency for oxidation. Some metals are more easily oxidized than others. Metals are arranged in a table by their tendency to be oxidized. This is called the activity series (Figure 4.16, page 139.) Know how to use this table.
oxidation-reduction (redox) reactions.
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ACTIVITY SERIES Metals at the top are more easily oxidized than metals towards the bottom.
Metals at the top of the table are called active metals; those at the bottom are called noble metals. Any neutral metal in the table can be oxidized by ions of elements below it.
OXIDATION OF METALS BY ACIDS Only metals above hydrogen can be oxidized by reaction with acids (H+); e.g., reaction
of Zn(s) with HCl.
Metals below hydrogen do not react with acids. Consider the reaction of Fe(s) with HCl:
Fe(s) + 2 HCl(aq)
FeCl2(aq) + H2(g)
Activity Series for Metals
Increasing easse of oxidation
Reaction occurs, but slowly.
Figure 4.15 a
Figure 4.16
Fe(s) + 2 H+(aq)
Fe2+(aq) + H2(g)
Oxidation of metals by acids: Zn loses 2 electrons Figure 4.15 b
+
Zn(s) + 2H (aq)
2+
Zn (aq) + H2(g)
2H+ gain 2 electrons
More active the metal the faster the reaction with acids.
Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
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ACTIVITY SERIES TABLE Consider the reaction: Figure 4.15 c
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Zinc is oxidized since it is above Copper; Cu2+ is reduced since it acquires the electrons that Zn looses. Mg(s) + 2 H+(aq)
Mg2+(aq) + H2(g)
In the laboratory Metals at the top of to the table are called active metals.
Zn(s) + CuSO4(aq)
ZnSO4(aq) + Cu(s)
Net Ionic Equation: Metals at the bottom of the table are called noble metals.
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Most Active (active)
Activity Series for M t l Metals
Least Active (noble)
Increasing ea ase of oxidation
Figure 4.10a illustrates this reaction with a Zn bar placed in a solution containing Cu2+ ions.
Cu(s) + ZnSO4(aq)
CuSO4(aq) + Zn(s)
Cu(s) + Zn2+(aq)
Cu2+(aq) + Zn(s)
This reaction does not occur because Zn is a more active metal (tends to loose its electrons) than Cu (activity series).
Figure 4.16
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ACTIVITY SERIES TABLE Consider the reaction:
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Copper is oxidized since it is above Silver; Ag+ is reduced since it acquires the electrons that Cu looses. Refer to Figure 4.16
Activity Series for M t l Metals
Least Active (noble)
Zn(s) + Cu2+(aq)
Anytime an element (H2, O2, Cl2, metal, etc.) reacts to yield a compound, the reaction is a redox reaction (involves a transfer of electrons). Examples:
4Fe(s) + 3O2(g)
2Fe2O3(s)
2H2(g) + O2(g)
2H2O(g)
CH4(g) + 2O2(g)
Increasing ease of oxidation
Most Active (active)
There are many types of redox reactions.
CO2(g) + 2H2O(g)
All reactions with oxygen (O2) to yield compounds are redox reactions (involve the transfer of electrons).
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
Reactions R ti of f molecular l l substances b t d do nott involve i l the complete transfer of electrons. However, a partial transfer of electrons does occur. Figure 4.16
Zn2+(aq) + Cu(s)
Zinc is oxidized, it looses electrons; Cu2+ is reduced, it acquires the electrons that Zn looses. Reducing agent: Acts to reduce another species, it is therefore oxidized. Zn is the reducing agent. Oxidizing agent: Acts to oxidize another species, it is therefore reduced. Cu2+ is the oxidizing agent.
Oxidation number (oxidation state): The charge that an atom would have in a molecule (or ionic compound) if there was complete electron transfer. A bookkeeping system to follow electron transfer in reactions. Need to assign OXIDATION NUMBERS (STATES) to elements in compounds. Set of rules (guidelines) to follow for assigning oxidation numbers.
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RULES ( page 134): 1) The oxidation number of any free element is zero: Examples: Fe(s) Î 0 O2(g) Î 0 O3(g) Î 0 Cl2(g) Î 0 Cl2(l) Î 0 Cl(g) Î 0
2) The oxidation number of any monatomic ion is equal to the charge on the ion.
There can be conflicts in these rules. In those cases, the rule with the lower number applies. Do not confuse oxidation number with charge on an ion: Ion charge Î 2+ (number followed by the sign) Oxidation number Î +2 (sign followed by the number). H2O: H Î zero charge; Oxidation number Î +1 O Î zero charge; Oxidation number Î -2
Examples:
Metals always have positive oxidation numbers in compounds. In binary ionic compounds, nonmetals have oxidation numbers equal to the charges on their stable monatomic anions.
H2O:
What is the oxidation number for O ?? 2 (# H) + 1 (# O) = 0 2 (+1) + 1 (# O) = 0 # O = 0 – 2 = -2
Examples: Na+ Î +1 Ca2+ Î +2 O2- Î -2 Fe3+ Î +3
3) The sum of all the oxidation numbers of the atoms in a molecule or polyatomic ion must equal the charge on the unit. 4) In compounds, fluorine (F) has an oxidation number of –1. 5) In compounds, hydrogen (H) has an oxidation number of +1. 6) In compounds, oxygen (O) has an oxidation number of –2.
H Î +1
O Î -2
Examples: H2S:
H Î +1
What is the oxidation number for S ?? 2 (# H) + 1 (# S) = 0 2 (+1) + 1 (# S) = 0 # S = 0 – 2 = -2 S Î -2
18
Examples: NH3:
4 H lose 4 electrons
H Î +1
What is the oxidation number for N ?? 3 (# H) + 1 (# N) = 0 3 (+1) + 1 (# N) = 0 # N = 0 – 3 = -3
2H2(g) + O2(g)
2 O gain 4 electrons
N Î -3
-4 +1
Examples: NO3-:
2H2O(g)
O Î -2
What is the oxidation number for N ?? 3 (# O) + 1 (# N) = -1 1 3 (-2) + 1 (# N) = -1 # N = -1 + 6 = +5 N Î +5
0
+4 -2
CH4(g) + 2O2(g)
+1 -2
CO2(g) + 2H2O(g)
Assign oxidation numbers to atoms of each element in this reaction: O Oxygen h been has b reduced d d (gain ( i of f electrons). l ) Carbon has been oxidized (loss of electrons). No change in the oxidation number of hydrogen.
Look at Example 4.4 (pages 135-136).
0
0
2H2(g) + O2(g)
+1 -2
2H2O(g)
Assign oxidation numbers to atoms of each element in this reaction:
1 C loses 8 electrons
Oxygen has been reduced (gain of electrons). Hydrogen has been oxidized (loss of electrons). Reduction Î The oxidation number has been reduced!!
CH4(g) + 2O C O2(g)
CO2(g) + 2H2O(g)
4 O gain 8 electrons
Oxidation Î The oxidation number has been increased!!
19
Types of Redox Reactions
Iron production in a blast furnace: Involves the reduction of Fe(III) in the form of Fe2O3 ore to metallic iron, Fe. 2 Fe gain 6 electrons
Combination Reactions: Two or more substances combine to form a single product. 0
0
+1 -2
2H2(g) + O2(g) Fe2O3(s) + 3 C(s)
2 Fe(l) + 3 CO(g)
100.0 g
69.9 g
0
2H2O(g)
0
+1 -1
2 Na(s) + Cl2(g)
2 NaCl(s)
3 C lose 6 electrons
Thermite reaction
Types of Redox Reactions
2 Fe gain 6 electrons Decomposition Reactions: A compound is broken down to two or more substances.
2 Al(s) + Fe2O3(s)
Al2O3(s) + 2 Fe(l)
2 Al lose 6 electrons This reaction produces enough heat to yield molten iron!!
+1 -2
2 H2O(g) +1 -1
2 NaCl(s)
0
0
2 H2(g) + O2(g) 0
0
2 Na(s) + Cl2(g)
It can be used to fuse iron bars together (a chemical weld!!!).
Types of Redox Reactions
Thermite reaction used to weld railroad rails together. together
Displacement Reactions: An ion (or atom) in a compound is replaced by an ion (or atom) of another element. General Reaction:
A + BC
AC + B
Several types of displacement reactions (hydrogen, metal, halogen).
20
Types of Redox Reactions
Types of Redox Reactions Disproportionation Reactions: One element is both oxidized and reduced in the reaction.
Hydrogen Displacement Reactions: 0
+1 -2
2 Na(s) + 2 H2O(l) 0
+1 -1
Zn(s) + 2 HCl(aq)
+1, -2, +1
2 NaOH(aq) + H2(g) +2 -1
+2 -2
0
0
ZnCl2(aq) + H2(g)
3 NO(g)
+1 -2
+4 -2
N2O(g) + NO2(g)
Nitrogen is reduced to +1 oxidation state and oxidized to a +4 oxidation state. There is no change in the oxidation state of oxygen. +4
3 NO2(g) + H2O
+5
+2
2 HNO3(aq) + NO(g)
Types of Redox Reactions Metal Displacement Reactions: Thermite reaction 0
+3 -2
2 Al(s) + Fe2O3(s)
+3 -2
0
Al2O3(s) + 2 Fe(l)
Types of Redox Reactions Halogen Displacement Reactions: 0
+1 -1
Cl2(g) + 2 KBr(aq) 0
+1 -1 1
Br2(g) + 2 KI(aq)
+1 -1
0
2 KCl(aq) + Br2(l) +1 -1 1
0
2 KBr(aq) + I2(s)
Order of reactivity of halogens: F2 > Cl2 > Br2 > I2 F2 is so reactive that it reacts directly with water.
21
There are many different ways to express concentration quantitatively. Common way to express concentration:
• What is a solution ? – Homogeneous mixture of two or more substances. – Composed of a solvent with solutes dissolved. • Solvent: Substance that solutes are dissolved in. • Solute: Substance dissolved in the solvent to give a solution. • Solutes: Can be a gas (HCl, CO2), liquid, or solid.
Molarity ( M ) =
moles solute solution volume (liters)
(molar concentration) Example: Prepare a HNO3 solution by dissolving 0.0375 mol HNO3 in water to yield a 0.250 L solution. What is the molarity of this solution?
Solution Preparation There can be numerous solutes dissolved in the solvent to give a complex solution. We will consider water as the solvent Î
aqueous solutions. l ti n
Concentration: Amount of solute dissolved in a given quantity of solvent or solution.
Qualitative Description: – Concentrated solution: large amount of solute. – Dilute solution: Only a small amount of solute.
Quantitative Description: – Exact measurement of the amount of solute in a given quantity of solvent or solution.
1
QUANTITATIVE DESCRIPTION OF AQUEOUS SOLUTIONS
mL solution Î L solution:
(200 mL ) ⎛⎜
1L ⎞ ⎟ = 0.200 L ⎝ 1000 mL ⎠
Our concentration is calculated as:
0.0375 mol HNO 3 = 0.150 M HNO 3 0.250 L solution Conversion factor Very useful as a ratio in calculations:
0.150 mol HNO 3 1.00 L solution or 1.00 L solution 0.150 mol HNO 3
Example Calculations 1) How many mL of a 0.150 M HNO3 solution is required to give 0.0450 mol HNO3? Set up the problem: given i
X
conversion i factor f t
=
M=
0.0476 mol HNO 3 = 0.238 0 238 M HNO 3 0.200 L
We can combine these conversions into one series ⎛ 3.00 g HNO 3 ⎞⎛ 1 mol HNO 3 ⎞⎛ 1000 mL ⎞ ⎟⎜⎜ ⎜ ⎟ = 0.238 M HNO 3 ⎟⎟⎜ ⎝ 200 mL ⎠⎝ 63.02 g HNO 3 ⎠⎝ 1 L ⎠
3) How many moles of HNO3 are in 50.0 mL of 0.238 M HNO3?
Conversion: Volume Î moles 1 L ⎞⎛ 0.238 mol HNO 3 ⎞ ⎟ = 0.0119 mol HNO 3 ⎟⎜ 1L ⎝ 1000 mL ⎠⎝ ⎠
(50 mL )⎛⎜
desired d i d
⎛
⎞ 1L ⎟⎟ = 0.300 L 0.150 mol HNO ⎝ 3⎠
(0.0450 mol HNO 3 ) ⎜⎜
300 mL of 0.150 M HNO3 contains 0.0450 mol HNO3
2) What is the concentration of a HNO3 solution prepared by dissolving 3.00 g HNO3 to prepare a 200 mL solution? g HNO3 Î mol HNO3
Conversions:
Calculate Molarity (moles solute/liter solution):
mL solution Î L solution
Conversion factor Alternative calculation ⎛ 0.238 mol HNO 3 ⎞⎛ 1 L ⎞ ⎟⎜ ⎜ ⎟(50 mL ) = 0.0119 mol HNO 3 1L ⎝ ⎠⎝ 1000 mL ⎠
4) How many grams of HNO3 is needed to
prepare a 250 mL solution that has a concentration of 0.100 M HNO3 ? The concentration required is:
0.100 M =
Solve: g HNO3 Î mol HNO3
⎛ 1 mol HNO 3 ⎞ ⎟⎟ = 0.0476 mol HNO 3 ⎝ 63.02 g HNO 3 ⎠
(3.00 g HNO 3 ) ⎜⎜
0.100 mol HNO 3 1 L solution
Two steps in calculation: First: Calculate mol HNO3 in 250 mL Second: Convert mol HNO3 to g HNO3
2
Calculate mol HNO3 in 250 mL: ⎛ 1 L ⎞⎛ 0.100 mol HNO 3 ⎞ (250 mL)⎜ ⎟ = 0.0250 mol HNO 3 ⎟⎜ 1L ⎝ 1000 mL ⎠⎝ ⎠
Convert mol HNO3 to g HNO3:
⎛ 63.02 g HNO 3 ⎞ ⎟⎟ = 1.58 g HNO 3 ⎝ 1 mol HNO 3 ⎠
(0.0250 mol HNO 3 )⎜⎜
Figure 4.19 Only added more solvent upon dilution, same amount (moles) of solute.
DILUTION OF SOLUTIONS • Often, a solution of known concentration must be diluted to prepare a solution with a lower desired concentration.
• For any solution:
DILUTION OF SOLUTIONS
M i Vi = M f V f What volume of a 0.150 M HNO3 solution is required to prepare 250 mL of a 0.0100 M HNO3 solution? M i Vi = M f V f
Vi =
MfVf Mi
molarity X volume (L) = moles solute (0.0100 M )(250 mL) = 16.7 mL 0.150 M Take 16.7 mL of 0.150 M HNO3 solution and dilute to a final volume of 250 mL with the solvent. Vi =
mol solute X L solution = moles solute L solution
DILUTION OF SOLUTIONS • Take a given amount of solute (volume of solution) and add solvent to give a new concentration. • Upon dilution, moles solute is constant, the volume has simply increased. M i X Vi = moles solute
M f X V f = moles solute
Pure water: A very poor conductor of electricity (does not conduct electricity). m solutes, u , when w n dissolved in n w water,, y yield an n Some
electrically conducting solution.
Initial
Final (diluted)
Electrical Conductivity of Aqueous Solutions
equal moles solute
These solutes are called electrolytes.
3
Other solutes, when dissolved in water, yield a solution that does not conduct electricity (non electrically conducting solution).
Electrolytes Î STRONG OR WEAK Strong electrolytes: Dissolve in the solvent exclusively as dissociated ions.
These solutes are called nonelectrolytes.
Examples of strong electrolytes:
Ionic compounds Î electrolytes Molecular Î generally nonelectrolytes compounds
All ionic i i compounds d CaCl2 NaNO3 KClO3 Pb(NO3)2 Each of these compounds dissociates in water to yield ions exclusively.
Dissolving of an Ionic Compound
Pure water
nonelectrolyte solution
electrolyte solution
Electrical Conductivity of Aqueous Solutions Svante Arrhenius: 1884; postulated that electrolytic solutions are composed of dissolved ions (Nobel prize (Chemistry) - 1903)
NaCl (s)
water
Na+(aq) + Cl-(aq)
Dissociation of Strong Electrolytes
CaCl2
water
NaNO3 KClO3 Pb(NO3)2
Ca2+ + 2Cl-
water
Na+ + NO3-
water
K+ + ClO3-
water
Pb2+ + 2NO3-
Ions that allow for the conduction of electricity in aqueous solutions.
Note: Polyatomic ions stay intact when dissolved (do not break apart further).
Consequently, electrolytes dissociate to ions upon dissolving.
Note: There are no intact ionic compounds (M+X-) dissolved in the solution; only dissociated ions.
4
Weak Electrolytes • Only partially dissociate to ions for the dissolved species. • Solutions poor conductors of electricity. • Some molecular compounds are weak electrolytes; e.g., ammonia (NH3) and acetic acid (HC2H3O2). Most , however, are nonelectrolytes. • Molecular compounds that are acids or bases are electrolytes (dissociate to ions in solution).
Complete dissociation to ions, hence, HCl is a molecular compound that is a strong electrolyte. Contrast this with the weak acid, hydrofluoric acid, which is a weak electrolyte:
water
HF
H+ + F -
There is only partial dissociation to ions. Both molecular HF and the dissociated ions, H+ and F-, exist in solution.
ACIDS Arrhenius Definition (1884): An acid is a substance that produces H+ ions (H3O+) when dissolved in water. Very y limited definition,, restricted to aqueous q solutions. Brønsted Definition (1932): An acid is a hydrogen ion, H+ (proton), donor.
HF is, therefore, a weak electrolyte. In an HF solution, HF exists mostly as undissociated HF molecules. O l a small Only ll amountt dissociates di i t to t ions i (ionized). (i i d) Acids can be either strong or weak electrolytes.
Broader definition of acids; does not require the acid to be dissolved in water.
ACIDS
CHARACTERISTICS OF ACIDS
Example: HCl H3O+ + Cl-
HCl + H2O
C Common I Inorganic i A Acids: id
Short-hand representation:
HCl
water
1) Sour (tart) taste in solution (citrus drinks have citric acid). 2) Turn litmus paper red. 3) React with many metals to release H2.
+
-
H + Cl
Ionization reaction: A reaction that forms ions from the dissociation of molecules.
HCl HNO3 HF
water
H+ + Cl-
water
H+ + NO3-
water
H+ + F-
5
Triprotic Acid Common Organic Acid:
Three acidic Hydrogen atoms H3PO4 (Phosphoric Acid)
Acetic Acid – HC2H3O2 water
HC2H3O2
H+ + C2H3O2acetate ion
All previous acids are monoprotic; i.e., one acidic hydrogen.
H3PO4
H+ + H2PO4-
H2PO4-
H+ + HPO42-
HPO42-
H+ + PO43-
BASES
Diprotic Acids Acids with two acidic hydrogen atoms. Examples: H2SO4 H2SO3 H2CO3 H2C2O4
sulfuric acid sulfurous acid carbonic acid oxalic acid (organic acid)
What happens in solution with a diprotic acid ?
Arrhenius Definition (1884): A base is a substance that produces OH- ions when dissolved in water. Very limited definition, restricted to aqueous solutions. Brønsted Definition (1932): A base is a hydrogen ion, H+ (proton), acceptor. Broader definition of bases; does not require the base to be dissolved in water.
Diprotic Acids in Solution
BASES
Sulfurous Acid (H2SO3)
H2SO3
H+ + HSO3- (hydrogen sulfite ion)
HSO3-
H+ + SO32- (sulfite ion)
H+ + HCO3- (hydrogen carbonate ion)
HCO3-
1) Bitter taste in solution 2) Solutions are often slippery (soapy feel) 3) Turn litmus paper blue Two Categories:
Carbonic Acid (H2CO3) H2CO3
Characteristics of bases:
H+ + CO32- (carbonate ion)
1) Ionic compounds that contain hydroxide ions (OH-). 2) Molecular compounds that react with water to yield hydroxide ions (OH-).
6
Ionic Hydroxides:
NaOH
Seven strong acids: Know these!!!!
water
Na+ + OH-
water
Ca2+ + 2OH-
Ca(OH)2
The hydroxide ion is a hydrogen ion (H+, proton) acceptor, therefore, it is a base by our definition. The two ions combine to form the weak electrolyte, water.
H+ + OH-
H2O
Strong Bases: Know these!!!! Group 1A metal hydroxides:
Bases that do not contain OH- ions.
NH3 + H2O
NH4+ + OH-
off of water to yield
– LiOH, NaOH, KOH, RbOH, CsOH
Heavier Group 2A metal hydroxides: – Ca(OH)2, Sr(OH)2, Ba(OH)2 – Exceptions: (Be(OH)2 and Mg(OH)2 are not strong bases)
Hydrogen ion (proton) acceptor. Ammonia pulls
HCl - hydrochloric acid HBr - hydrobromic acid HI - hydroiodic acid HNO3 - nitric acid HClO3 - chloric acid H HClO4 - perchloric acid H2SO4 - sulfuric acid
All other acids are weak, e.g., HC2H3O2 - acetic acid HF - hydrofluoric acid
Molecular Bases:
H+
• • • • • • •
OH-.
ACIDS AND BASES
Strong acids and bases Î Strong Electrolytes Weak acids and bases Î Weak Electrolytes
Acid/Base Definitions:
Classified as either strong or weak.
Arrhenius Definition:
Strong acids and Strong bases: Completely dissociate to ions in solution. Hence they are also strong electrolytes.
Acid: A substance that increases the H3O+ concentration in water.
Weak acids and Weak bases: Only undergo partial dissociation to ions in solution. Hence, they are also weak electrolytes.
Base: A substance that increases the OH- concentration is water.
Most acids and bases are weak.
7
Brønsted definition: Acid: Hydrogen ion (H+, proton) donor. Base: Hydrogen ion (H+, proton) acceptor. HCl + H2O
H3O+ + Cl-
acid
NH3 + H2O
NH4+ + OH-
base
What happens when an acid and a base react?
acid + base
salt + H2O
Ionic compounds Î dissolve in water dissociating to ions How much of the ionic compound can dissolve in a given quantity of water? Î solubility Solubility: The maximum amount of solute that can dissolve l in a given quantity of solvent l at a specific temperature Î saturated solution. Solubility expressed as:
mol solute grams solute or liter solution 100 g solvent
Solubility of Ionic Compounds Compound NaCl
Acid and base neutralize (destroy) each other. This is called a neutralization reaction.
• Salt: Ionic compound whose cation comes from a base and whose anion comes from an acid. – cations other than H+ – anions other than OH- or O2-
HCl + NaOH acid
base
NaCl + H2O salt
Solubility (g/100 g water) 35.7 g @ 0 ºC 39.1 g @ 100 ºC
CuS (copper(II) sulfide)
3.3 X 10-5 g @ 18 ºC
CaCl2
74.5 g @ 20 ºC
CaCO3
1.5 X 10-3 g @ 25 ºC
All ionic compounds are soluble to some extent in water. However, many have very low solubilities ( 0.1 0 1 g/100 g water) or insoluble (solubility < 0.1 g/100 g water).
NaCl and CaCl2 are soluble CuS and CaCO3 are insoluble
8
General rules for predicting the solubility of ionic compounds. Summarized in Table 4.2, page 123; KNOW THIS TABLE!!!
Solubility Guidelines for Ionic Compounds Soluble Compounds All alkali metal ions ions, NH4+
Important Exceptions None
NO3-, C2H3O2-, HCO3-, ClO3-, ClO4Cl-, Br-, I-
None
SO42-
Strong Acid
H+,
Cl-
H2O
In solution solution:
Salts of Ag+, Hg22+, Pb2+
water
H+(aq) + Cl-(aq)
H+,
Cl-,
H2O
Salts of Ca2+, Sr2+, Ba2+, Hg22+, Pb2+, Ag+
Insoluble Compounds CO32-, PO43-, SO32-, S2-, CrO42-
Important Exceptions Salts of NH4+, all alkali metal ions
OH-
NH4+, all alkali metal ions ions, Ca2+, Sr2+, Ba2+
(acid oxides)
HCl solution:
HCl
Solubility Guidelines for Ionic Compounds
O2-
What are the main species in solution??
All alkali metal ions, Sr2+, Ba2+
Solubility Examples: Na2CO3
Soluble (Na+)
CaCO3
Insoluble (CO32-)
AgNO3
Soluble (NO3-)
AgCll
Insoluble l bl ((Ag+)
PbSO4
Insoluble (Pb2+)
CaSO4
Insoluble (Ca2+)
Ca2+,
What are the main species in solution?? HF solution:
Weak Acid
H+, F-
In solution solution:
HF, H2O
HF
water
HF,
H+(aq) + F-(aq)
H+,
F-,
H2O
What are the main species in solution?? NaOH solution:
Strong Base
Na+,
H2O
In solution solution:
NaOH
water
Na+,
OH-
Na+(aq) + OH-(aq)
OH-,
H2O
Apply the solubility rules to determine if an ionic compound is soluble or insoluble.
9
What are the main species in solution??
Molecular Equation:
HCl(aq) + NaOH(aq)
NH3 solution:
NaCl(aq) + H2O(l)
Weak Base
NH4+, OH-
Ionic Equation:
In solution solution:
NH3, H2O
Strong g electrolytes y written as dissociated ions.
water
NH3 + H2O NH3,
NH4+(aq) + OH-(aq)
NH4+,
OH-,
H2O
What are the main species in solution??
Soluble Ionic Compound
Mg2+,
Cl-
H2O
Mg2+,
Net Ionic Equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l) Na+ and Cl- do not react, spectator ions.
In solution: water
Na+(aq) + Cl-(aq) + H2O(l)
Eliminate spectator ions form the ionic equation.
MgCl2 solution:
MgCl2
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)
Mg2+(aq) + 2Cl-(aq)
Cl-,
H2O
Reactions of Ions in Solution 3 ways to represent reactions in solution 1) Molecular equation: Each substance written as intact units (molecules). 2) Ionic equation: Strong electrolytes written as dissociated ions (how they exist in solution). 3) Net ionic equation: Only the species that actually take part in the reaction are shown.
Net Ionic Equation: H+(aq) + OH-(aq)
H2O(l)
A neutralization reaction
Balanced Ionic Reactions: 1) Need atom balance 2) Net charge balance reactant charge = product charge
Molecular Equation:
HC2H3O2(aq) + NaOH(aq) NaC2H3O2(aq) + H2O(l)
10
Ionic Equation:
Molecular Equation:
HC2H3O2(aq) + Na+(aq) + OH-(aq) weak electrolyte
CaCl2(aq) + Na2CO3(aq) CaCO3(s) + 2 NaCl(aq)
Na+(aq) + C2H3O2-(aq) + H2O(l)
AB + CD Net Ionic Equation (eliminate spectator ions):
HC2H3O2(aq) + OH-(aq) weak electrolyte
C2H3O2-(aq) + H2O(l)
CaCl2
+
Soln 1
Reactions yielding a precipitate are a common type of ionic reaction. Precipitate formation drives the reaction; ions are removed from the solution.
Precipitation Reaction
Reactions between ionic substancessoluble ionic compounds.
CaCl2(aq) + Na2CO3(aq)
AD + CB
???
Na2CO3
Î
Soln 2
mix the two solutions
Ca2+(aq) + 2Cl -(aq) + 2Na+(aq) + CO32- (aq) What are the products ?
2KI(aq) + Pb(NO3)2(aq)
PbI2(s) + 2K+(aq) + 2NO3-(aq)
Ca2+(aq) + 2Cl -(aq) + 2Na+(aq) + CO32- (aq) CaCO3(s) + 2Na+(aq) + 2Cl -(aq) precipitate The reaction is called a metathesis reaction
(double displacement reaction).
AB + CD
AD + CB
11
2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2K+(aq) + 2NO3-(aq) Net Ionic Equation:
Reactions of insoluble salts:
MgCO3(s) + 2 HCl(aq) MgCl2(aq) + H2CO3(aq)
2I-(aq) + Pb2+(aq)
PbI2(s)
There are many examples Th l of f precipitation i it ti reactions-need to know the solubility rules and how to apply them.
Remember, if you are not part of the solution…… then you are part of the precipitate.
Some ionic reactions proceed to yield a weak electrolyte. Molecular: HCl(aq) + NaC2H3O2(aq)
NaCl(aq) + HC2H3O2(aq)
Ionic: H+(aq) + Cl-(aq) + Na+(aq) + C2H3O2-(aq) +
H2CO3(aq)
CO2(g) + H2O(l)
Baking: Na2CO3 baking powder CO2(g) NaHCO3 baking soda during baking (raises cake)
TITRATIONS • Use chemical reactions to determine concentration of a solution. • Use a solution of known concentration (standard solution) to determine the concentration of another solution of unknown concentration. • Reactions between the two solutions. • A common titration is an acid/base titration
-
Na (aq) + Cl (aq) + HC2H3O2(aq) weak electrolyte
Simple Titration
Net Ionic: H+(aq) + C2H3O2-(aq)
HC2H3O2(aq) weak electrolyte
Weak Electrolytes: Consist of weak acids and bases: HC2H3O2, NH3, HF, HNO2
12
HCl(aq) + NaOH(aq)
Titration
NaCl(aq) + H2O(l)
neutralization reaction 1 mol HCl (H+) reacts with 1 mol NaOH (OH-)
Example titration: A 20.0 mL HCl solution
required 37 37.8 8 mL of 0 0.100 100 M NaOH to reach the endpoint in the titration. What is the concentration of the HCl solution ? At the endpoint (equivalence point):
mol NaOH (OH-) = mol HCl (H+)
First: Calculate mol NaOH used in the titration. Second: Calculate concentration of HCl. Note that mol HCl = mol NaOH Concentration X Volume = number of moles
Oxidation-Reduction Reactions The following is a common teaching laboratory reaction:
Zn(s) (s) + 2 HCl(aq) C (aq)
(Molarity)
Mol NaOH:
⎛ 0.100 mol NaOH ⎞⎛ 1 L ⎞( ⎜ ⎟⎜ ⎟ 37.8 mL ) 1L ⎝ ⎠⎝ 1000 mL ⎠ = 0.00378 mol NaOH
Second, calculate concentration of HCl At the endpoint, mol HCl (H+) = mol NaOH (OH-): mol NaOH = mol HCl = 0.00378 mol
The concentration of HCl (Molarity) is: MOLARITY =
moles solute liter of solution
⎛ 0.00378 mol HCl ⎞⎛ 1000 mL ⎞ ⎜ ⎟⎜ ⎟ = 0.189 M HCl 20 mL ⎝ ⎠⎝ 1 L ⎠
Use acid/base titrations to determine concentration
ZnCl C 2(aq) + H2(g)
Ionic Equation:
Zn(s) + 2 H+(aq) + 2 Cl-(aq) Zn2+(aq) + 2 Cl-(aq) + H2(g)
Net Ionic Equation:
Zn(s) + 2 H+(aq) + 2 Cl-(aq) Zn2+(aq) + 2 Cl-(aq) + H2(g) Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
Zn has lost 2 electrons !!!! 2 H+ have gained 2 electrons !!!! This reaction involves the transfer of electrons from Zn metal to hydrogen ions (H+).
13
CORROSION OF METALS BY ACIDS Corrosion: The conversion of the elemental form of a metal (native metal) to a metal compound by reaction with a substance in its environment. Metals tend to lose electrons to form cations:
Fe(s)
Fe3+ + 3 e-
4 Fe(s) + 3 O2(g)
2 Fe2O3(s)
4 Fe3+ + 12 e- oxidation
4 Fe(s)
3 O2 + 12 e-
6 O2-
reduction
Redox reactions are a very important class of reactions. Many reactions proceed with the transfer of electrons from one species to another. Electrons are central to chemistry!!!!
This often involves reaction with oxygen:
4 Fe(s) + 3 O2(g)
2 Fe2O3(s)
Since most metals lose electrons by reaction with O2, the metal is said to be oxidized.
Loss of electrons is called oxidation. oxidation Since we must maintain electrical neutrality, some other species gains the electrons that are lost by the metal. Oxygen gains the electrons.
O2 + 4 e-
Oxidation of metals by acids: Zn loses 2 electrons (oxidized)
Zn(s) (s) + 2 HCl(aq) C (aq)
ZnCl C 2(aq) + H2(g)
2 H+ gain 2 electrons (reduced)
2 O2-
• When a species gains electrons (becomes more negatively charged) it is said to be reduced. • A gain in e-’s by a species is called reduction. • If f one species gains i electrons, l t another h must
lose electrons.
• Oxidation and reduction reactions are coupled together. • These types of reactions are called
Oxidation of metals by acids: Metals vary in their tendency for oxidation. Some metals are more easily oxidized than others. Metals are arranged in a table by their tendency to be oxidized. This is called the activity series (Figure 4.16, page 139.) Know how to use this table.
oxidation-reduction (redox) reactions.
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ACTIVITY SERIES Metals at the top are more easily oxidized than metals towards the bottom.
Metals at the top of the table are called active metals; those at the bottom are called noble metals. Any neutral metal in the table can be oxidized by ions of elements below it.
OXIDATION OF METALS BY ACIDS Only metals above hydrogen can be oxidized by reaction with acids (H+); e.g., reaction
of Zn(s) with HCl.
Metals below hydrogen do not react with acids. Consider the reaction of Fe(s) with HCl:
Fe(s) + 2 HCl(aq)
FeCl2(aq) + H2(g)
Activity Series for Metals
Increasing easse of oxidation
Reaction occurs, but slowly.
Figure 4.15 a
Figure 4.16
Fe(s) + 2 H+(aq)
Fe2+(aq) + H2(g)
Oxidation of metals by acids: Zn loses 2 electrons Figure 4.15 b
+
Zn(s) + 2H (aq)
2+
Zn (aq) + H2(g)
2H+ gain 2 electrons
More active the metal the faster the reaction with acids.
Zn(s) + 2 H+(aq)
Zn2+(aq) + H2(g)
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ACTIVITY SERIES TABLE Consider the reaction: Figure 4.15 c
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Zinc is oxidized since it is above Copper; Cu2+ is reduced since it acquires the electrons that Zn looses. Mg(s) + 2 H+(aq)
Mg2+(aq) + H2(g)
In the laboratory Metals at the top of to the table are called active metals.
Zn(s) + CuSO4(aq)
ZnSO4(aq) + Cu(s)
Net Ionic Equation: Metals at the bottom of the table are called noble metals.
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Most Active (active)
Activity Series for M t l Metals
Least Active (noble)
Increasing ea ase of oxidation
Figure 4.10a illustrates this reaction with a Zn bar placed in a solution containing Cu2+ ions.
Cu(s) + ZnSO4(aq)
CuSO4(aq) + Zn(s)
Cu(s) + Zn2+(aq)
Cu2+(aq) + Zn(s)
This reaction does not occur because Zn is a more active metal (tends to loose its electrons) than Cu (activity series).
Figure 4.16
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ACTIVITY SERIES TABLE Consider the reaction:
Cu(s) + 2 Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Copper is oxidized since it is above Silver; Ag+ is reduced since it acquires the electrons that Cu looses. Refer to Figure 4.16
Activity Series for M t l Metals
Least Active (noble)
Zn(s) + Cu2+(aq)
Anytime an element (H2, O2, Cl2, metal, etc.) reacts to yield a compound, the reaction is a redox reaction (involves a transfer of electrons). Examples:
4Fe(s) + 3O2(g)
2Fe2O3(s)
2H2(g) + O2(g)
2H2O(g)
CH4(g) + 2O2(g)
Increasing ease of oxidation
Most Active (active)
There are many types of redox reactions.
CO2(g) + 2H2O(g)
All reactions with oxygen (O2) to yield compounds are redox reactions (involve the transfer of electrons).
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
Reactions R ti of f molecular l l substances b t d do nott involve i l the complete transfer of electrons. However, a partial transfer of electrons does occur. Figure 4.16
Zn2+(aq) + Cu(s)
Zinc is oxidized, it looses electrons; Cu2+ is reduced, it acquires the electrons that Zn looses. Reducing agent: Acts to reduce another species, it is therefore oxidized. Zn is the reducing agent. Oxidizing agent: Acts to oxidize another species, it is therefore reduced. Cu2+ is the oxidizing agent.
Oxidation number (oxidation state): The charge that an atom would have in a molecule (or ionic compound) if there was complete electron transfer. A bookkeeping system to follow electron transfer in reactions. Need to assign OXIDATION NUMBERS (STATES) to elements in compounds. Set of rules (guidelines) to follow for assigning oxidation numbers.
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RULES ( page 134): 1) The oxidation number of any free element is zero: Examples: Fe(s) Î 0 O2(g) Î 0 O3(g) Î 0 Cl2(g) Î 0 Cl2(l) Î 0 Cl(g) Î 0
2) The oxidation number of any monatomic ion is equal to the charge on the ion.
There can be conflicts in these rules. In those cases, the rule with the lower number applies. Do not confuse oxidation number with charge on an ion: Ion charge Î 2+ (number followed by the sign) Oxidation number Î +2 (sign followed by the number). H2O: H Î zero charge; Oxidation number Î +1 O Î zero charge; Oxidation number Î -2
Examples:
Metals always have positive oxidation numbers in compounds. In binary ionic compounds, nonmetals have oxidation numbers equal to the charges on their stable monatomic anions.
H2O:
What is the oxidation number for O ?? 2 (# H) + 1 (# O) = 0 2 (+1) + 1 (# O) = 0 # O = 0 – 2 = -2
Examples: Na+ Î +1 Ca2+ Î +2 O2- Î -2 Fe3+ Î +3
3) The sum of all the oxidation numbers of the atoms in a molecule or polyatomic ion must equal the charge on the unit. 4) In compounds, fluorine (F) has an oxidation number of –1. 5) In compounds, hydrogen (H) has an oxidation number of +1. 6) In compounds, oxygen (O) has an oxidation number of –2.
H Î +1
O Î -2
Examples: H2S:
H Î +1
What is the oxidation number for S ?? 2 (# H) + 1 (# S) = 0 2 (+1) + 1 (# S) = 0 # S = 0 – 2 = -2 S Î -2
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Examples: NH3:
4 H lose 4 electrons
H Î +1
What is the oxidation number for N ?? 3 (# H) + 1 (# N) = 0 3 (+1) + 1 (# N) = 0 # N = 0 – 3 = -3
2H2(g) + O2(g)
2 O gain 4 electrons
N Î -3
-4 +1
Examples: NO3-:
2H2O(g)
O Î -2
What is the oxidation number for N ?? 3 (# O) + 1 (# N) = -1 1 3 (-2) + 1 (# N) = -1 # N = -1 + 6 = +5 N Î +5
0
+4 -2
CH4(g) + 2O2(g)
+1 -2
CO2(g) + 2H2O(g)
Assign oxidation numbers to atoms of each element in this reaction: O Oxygen h been has b reduced d d (gain ( i of f electrons). l ) Carbon has been oxidized (loss of electrons). No change in the oxidation number of hydrogen.
Look at Example 4.4 (pages 135-136).
0
0
2H2(g) + O2(g)
+1 -2
2H2O(g)
Assign oxidation numbers to atoms of each element in this reaction:
1 C loses 8 electrons
Oxygen has been reduced (gain of electrons). Hydrogen has been oxidized (loss of electrons). Reduction Î The oxidation number has been reduced!!
CH4(g) + 2O C O2(g)
CO2(g) + 2H2O(g)
4 O gain 8 electrons
Oxidation Î The oxidation number has been increased!!
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Types of Redox Reactions
Iron production in a blast furnace: Involves the reduction of Fe(III) in the form of Fe2O3 ore to metallic iron, Fe. 2 Fe gain 6 electrons
Combination Reactions: Two or more substances combine to form a single product. 0
0
+1 -2
2H2(g) + O2(g) Fe2O3(s) + 3 C(s)
2 Fe(l) + 3 CO(g)
100.0 g
69.9 g
0
2H2O(g)
0
+1 -1
2 Na(s) + Cl2(g)
2 NaCl(s)
3 C lose 6 electrons
Thermite reaction
Types of Redox Reactions
2 Fe gain 6 electrons Decomposition Reactions: A compound is broken down to two or more substances.
2 Al(s) + Fe2O3(s)
Al2O3(s) + 2 Fe(l)
2 Al lose 6 electrons This reaction produces enough heat to yield molten iron!!
+1 -2
2 H2O(g) +1 -1
2 NaCl(s)
0
0
2 H2(g) + O2(g) 0
0
2 Na(s) + Cl2(g)
It can be used to fuse iron bars together (a chemical weld!!!).
Types of Redox Reactions
Thermite reaction used to weld railroad rails together. together
Displacement Reactions: An ion (or atom) in a compound is replaced by an ion (or atom) of another element. General Reaction:
A + BC
AC + B
Several types of displacement reactions (hydrogen, metal, halogen).
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Types of Redox Reactions
Types of Redox Reactions Disproportionation Reactions: One element is both oxidized and reduced in the reaction.
Hydrogen Displacement Reactions: 0
+1 -2
2 Na(s) + 2 H2O(l) 0
+1 -1
Zn(s) + 2 HCl(aq)
+1, -2, +1
2 NaOH(aq) + H2(g) +2 -1
+2 -2
0
0
ZnCl2(aq) + H2(g)
3 NO(g)
+1 -2
+4 -2
N2O(g) + NO2(g)
Nitrogen is reduced to +1 oxidation state and oxidized to a +4 oxidation state. There is no change in the oxidation state of oxygen. +4
3 NO2(g) + H2O
+5
+2
2 HNO3(aq) + NO(g)
Types of Redox Reactions Metal Displacement Reactions: Thermite reaction 0
+3 -2
2 Al(s) + Fe2O3(s)
+3 -2
0
Al2O3(s) + 2 Fe(l)
Types of Redox Reactions Halogen Displacement Reactions: 0
+1 -1
Cl2(g) + 2 KBr(aq) 0
+1 -1 1
Br2(g) + 2 KI(aq)
+1 -1
0
2 KCl(aq) + Br2(l) +1 -1 1
0
2 KBr(aq) + I2(s)
Order of reactivity of halogens: F2 > Cl2 > Br2 > I2 F2 is so reactive that it reacts directly with water.
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