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Evidence of a chemical reaction:
• Gas Evolution
• Temperature Change
• Color Change
• Precipitation (insoluble species forms)
In general, a reaction involves a rearrangement
or change in oxidation state of atoms from
reactants to products.
Chemical Reactions
Chemical Equations show:
• the reactants and products in a reaction.
• the relative amounts in a reaction.
Example:
4 Al(s) + 3 O2(g) 2 Al2O3(s)
• The numbers in the front are called
stoichiometric coefficients
• The letters (s), (g), (l) and (aq) are the
physical states of compounds.
Chemical Equations
Notice the stoichiometric coefficients and the physical states of
the reactants and products.
Reaction of Phosphorus with Cl2
Notice the stoichiometric coefficients and the physical states of
the reactants and products.
Reaction of Iron with Cl2
4 Al(s) + 3 O2(g) 2 Al2O3(s)
This equation states that:
4 Al atoms + 3 O2 molecules
react to form 2 formula units
of Al2O3
or...
4 moles of Al + 3 moles of
O2 react to form 2 moles of
Al2O3
Chemical Equations
Law of the
Conservation of Matter
• Because the same
number of atoms are
present in a reaction at
the beginning and at
the end, the amount of
matter in a system
does not change.
2HgO(s) 2 Hg(l) + O2(g)
Chemical Equations
• Since matter is
conserved in a chemical
reaction, chemical
equations must be
balanced for mass!
• In other words, there
must be same number of
atoms of the each kind
on both sides of the
equatoin. Lavoisier, 1788
Chemical Equations
Steps in balancing a chemical reaction using coefficients:
1. Write the equation using the formulas of the reactants
and products. Include the physical states (s, l, g, aq
etc…)
2. Balance the compound with the most elements in the
formula first using integers as coefficients.
3. Balance elements on their own last.
4. Check to see that the sum of each individual elements
are equal on each side of the equation.
5. If the coefficients can be simplified by dividing though
with a whole number, do so.
Balancing Chemical Reactions
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
Balancing Chemical Equations: Example
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
3
balance H first
___C2H6 + O2 CO2 + ___ H2O
This side has an odd # of
O-atoms
This side will always have
an even # of O-atoms
Balancing Chemical Equations: Example
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
3
balance H first
___C2H6 + O2 CO2 + ___ H2O 2
Balancing Chemical Equations: Example
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
3
balance H first
___C2H6 + O2 CO2 + ___ H2O 2
balance C next
2C2H6 + O2 ___ CO2 + 6H2O 4
Balancing Chemical Equations: Example
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
3
balance H first
___C2H6 + O2 CO2 + ___ H2O 2
balance C next
2C2H6 + O2 ___ CO2 + 6H2O 4
balance O
2C2H6 + ____ O2 4CO2 + 6H2O 7
Balancing Chemical Equations: Example
C2H6 + O2 CO2 + H2O
2 C’s & 6 H’s 2 O’s 1 C & 2 O’s 2 H’s & 1 O
balance last
3
balance H first
___C2H6 + O2 CO2 + ___ H2O 2
balance C next
2C2H6 + O2 ___ CO2 + 6H2O 4
balance O
2C2H6 + ____ O2 4CO2 + 6H2O 7
4 C’s 12 H’s 14 O’s 4 C’s 12 H’s 14 O’s
Balancing Chemical Equations: Example
___C3H8(g) + ___ O2(g)
___ CO2(g) + _____ H2O(g)
___B4H10(g) + ___ O2(g)
___ B2O3(g) + ___ H2O(g)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
Balancing Equations: Practice
_ Mg(OH)2(s) + _ HCl(aq)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq) _ MgCl2(aq) + _ H2O(l)
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq) _ MgCl2(aq) + _ H2O(l)
Balance with a coefficient of ―2‖ in front of both HCl
and water.
Balancing Equations: Practice
• Solid magnesium hydroxide reacts with hydrochloric
acid to form aqueous magnesium chloride and
water.
• Write the balanced chemical equation for this
reaction.
_ Mg(OH)2(s) + _ HCl(aq) _ MgCl2(aq) + _ H2O(l)
Balance with a coefficient of ―2‖ in front of both HCl
and water.
Mg, Cl, O and H are now balanced.
• What Scientific Principles are used in the process of
balancing chemical equations?
• What symbols are used in chemical equations:
gasses: _____
liquids: _____
solids: _____
aqueous species in solution: _____
• What is the difference between P4 and 4P in an
eq.?
• In balancing a chemical equation, why are the
reactant and product subscripts not changed?
Chemical Equations: Review
When writing chemical reactions one starts with:
Reactants products
N2(g) + 3H2(g)
2NH3(g)
Some reactions can also run in reverse:
2NH3(g)
N2(g) + 3H2(g)
Under these conditions, the reaction can be written:
2 2 33H (g) N (g) 2NH (g)
Double arrows indicate ―Equilibrium‖.
Chemical Equilibrium
Chemical Equilibrium
Once equilibrium is achieved, reaction continues, but there
is no net change in amounts of products or reactants.
• Salts (ionic compounds): Composed of a
metal and non metal element(s).
• Acids: Arrhenius definition
Produce H+(aq) in water
Examples: HCl, HNO3, HC2H3O2
• Bases: Arrhenius definition
Produce OH (aq) in water
Examples: NaOH, Ba(OH)2, NH3
Classifying Compounds
• Molecular Compounds:
• Covalently bonded atoms, not acids, bases or
salts.
• Compounds like alcohols (C2H5OH) or table
sugar (C6H12O6)
• These never break up into ions.
Classifying Compounds
• Classify the following as ionic, molecular,
acid or base.
Compound Type
Na2SO4
Ba(OH)2
H3PO4
CH4
P2O5
NH3
HCN
Classifying Compounds
Classifying Compounds
• Classify the following as ionic, molecular,
acid or base.
Compound Type
Na2SO4 ionic
Ba(OH)2 base
H3PO4 acid
CH4 molecular
P2O5 molecular
NH3 base
HCN acid
Aqueous Solutions:
There are three types of aqueous solutions:
Those with Strong Electrolytes
Water as the solvent
Solution = solute + solvent
That which is dissolved
(lesser amount)
That which is dissolves
(greater amount)
Those with Weak Electrolytes
& those with non-Electrolytes
Reactions in Aqueous Solutions
Many reactions involve ionic compounds, especially reactions in water — aqueous solutions.
KMnO4 in water K+(aq) + MnO4-(aq)
Reactions in Aqueous Solutions
When ions are present in water,
the solutions conduct
electricity!
Ions in solution are called
ELECTROLYTES
Examples of Strong Electrolytes:
HCl (aq), CuCl2(aq) and NaCl
(aq) are strong electrolytes.
These dissociate completely (or
nearly so) into ions.
Strong Electrolytes conduct
electricity well.
Strong Electrolyte
HCl(aq), CuCl2(aq) and NaCl(aq) are strong electrolytes.
These dissociate completely (or nearly so) into ions.
Strong Electrolytes
Acetic acid ionizes only to a small
extent, it is a weak electrolyte.
Weak electrolytes exist in solution
under equilibrium conditions.
The small concentration of ions
conducts electricity poorly.
Weak electrolytes exit primarily in
their molecular form in water.
3 2 3 2CH CO H(aq) CH CO (aq) H (aq)
Weak Electrolytes
Weak electrolytic solutions are characterized by
equilibrium conditions in solution:
When acetic acid dissociates, it only partially
ionizes. +
2 3 2 2 3 2HC H O (aq) H (aq) + C H O (aq)
The majority species in solution is acetic acid in its
molecular form.
When writing a weak electrolyte in solution, one
NEVER breaks it up into the corresponding ions!
95% 5%
+2 3 2 2 3 2HC H O (aq) H (aq) + C H O (aq)×
Weak Electrolytes
Acetic acid ionizes only to a small extent, so it
is a weak electrolyte.
CH3CO2H(aq) CH3CO2-(aq) + H+(aq)
Weak Electrolytes
Some compounds dissolve in
water but do not conduct
electricity.
They are non-electrolytes.
Examples include:
• sugar
• ethanol
• ethylene glycol
Non-electrolytes do not
dissociate into ions!
Non-Electrolytes
Strong electrolytes: Characterized by ions only (cations &
anions) in solution (water).
Weak electrolytes: Characterized by ions (cations & anions)
& molecules in solution.
Non-electrolytes: Characterized by molecules in solution.
Conduct electricity well
Conduct electricity poorly
Do not conduct electricity
Species in Solution: Electrolytes
How do we know if a compound will be soluble in
water?
For molecular compounds, the molecule must be
polar.
We will discuss polarity later, for now I will tell you
whether or not a molecular compound is polar…
For ionic compounds, the compound solubility is
governed by a set of SOLUBILITY RULES!
You must learn the basic rules on your own!!!
Solubility Rules
Water Solubility of Ionic Compounds
If one ion from the ―Soluble Compound‖ list is present in a compound, then the compound is water soluble.
Precipitation Reactions: A reaction where an
insoluble solid (precipitate) forms and drops out
of the solution.
Acid–base Neutralization: A reaction in which an
acid reacts with a base to yield water plus a salt.
Gas forming Reactions: A reaction where an
insoluble gas is formed.
Reduction and Oxidation Reactions (RedOx): A
reaction where electrons are transferred from
one reactant to another.
Types of Reactions in a Solution
REDOX REACTIONS
EXCHANGE
Acid-Base
Reactions
EXCHANGE
Gas-Forming
Reactions
EXCHANGE: Precipitation Reactions
REACTIONS
EXCHANGE REACTIONS
The anions exchange
places between cations.
A precipitate forms if one of
the products in insoluble.
Pb(NO3) 2(aq) + 2 KI(aq)
PbI2(s) + 2 KNO3 (aq)
Chemical Reactions in Water
The ―driving force‖ is the formation of
an insoluble solid called a precipitate.
Pb(NO3)2(aq) + 2 KI(aq)
2 KNO3(aq) + PbI2(s)
BaCl2(aq) + Na2SO4(aq)
BaSO4(s) + 2 NaCl(aq)
Precipitates are determined from the
solubility rules.
Precipitation Reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?) 2KNO3(aq) + PbI2(?)
Precipitation reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?) 2KNO3(aq) + PbI2(?)
All potassium salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(?)
Precipitation Reactions
Which species is the precipitate?
Pb(NO3)2(?) + 2KI(?) 2KNO3(?) + PbI2(?)
From the solubility rules:
All nitrate salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(?) 2KNO3(aq) + PbI2(?)
All potassium salts are soluble, therefore:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(?)
By the solubility rules:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s)
PbI2 is the ppt.
Precipitation Reactions
Molecular Equation: all species listed as formula units or in
molecular form. reactants products
• Note all states of each reactant or product by: (s), (l), (g) or
(aq)
Ionic Equation: All soluble (aq) species present are listed as
ions.
• Leave all (s), (l) or (g) species as is. They do not dissociate
into ions
Net Ionic Equation:
• From the ionic equation, cancel out any species that appear
on either side of the equation.
• These are known as the ―spectator ions‖ and they are
never part of a net ionic equation!
Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
2K+(aq) + 2NO3– (aq) + PbI2(s)
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
2K+(aq) + 2NO3– (aq) + PbI2(s)
Never break up
any (s), (l) or (g)
or molecular
(aq) species!
Writing Net Ionic Equations
Molecular Equation:
Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s)
Total Ionic Equation:
Pb2+ (aq) + 2NO3– (aq) + 2K+(aq) + 2I–(aq)
2K+(aq) + 2NO3– (aq) + PbI2(s)
Never break up
any (s), (l) or (g)
or molecular
(aq) species!
Cancel out the spectator ions to yield the net ionic equation:
Pb2+ (aq) + 2I–(aq) PbI2(s)
Writing Net Ionic Equations
Arrhenius Definition:
• An acid is any substance that increases the
H+(aq) concentration in an aqueous solution.
HX(aq) H+(aq) + X–(aq)
• A base is any substance that increases the
OH–(aq) concentration in an aqueous
solution.
MOH(aq) M+(aq) + OH–(aq)
Acids & Bases
Brönsted-Lowry:
• An acid is any substance that donates H+(aq)
to another species in an aqueous solution.
HX(aq) + H2O(l) H3O+(aq) + X–(aq)
• A base is any substance that accepts an
H+(aq) in an aqueous solution.
H+(aq) + NH3(aq) NH4+(aq)
H3O+(aq) = H+(aq)
Acids and Bases
Strong acids are almost completely ionized in
water. (strong electrolytes)
Examples:
HX (aq) (X = Cl, Br & I) hydro ___ ic acid
HNO3 (aq) nitric acid
HClO4 (aq) perchloric acid
H2SO4 (aq)* sulfuric acid
* Only the 1st H is strong, sulfuric acid dissociates via:
H2SO4 (aq) H+ (aq) + HSO4– (aq)
Strong Acids
Weak Acids are incompletely ionized in water.
(weak electrolytes) Weak acids are governed by
dynamic equilibrium.
Examples:
HC2H3O2 (aq)
nitrous acid
hydrosulfuric acid
hydrogen sulfate ion
See you text and home work for more examples.
Weak acids are always written in their molecular form.
acetic acid (vinegar)
HNO2 (aq)
H2S (aq)
HSO4–(aq)
Weak Acids
2H O( )NaOH(s) Na (aq) ΟΗ aq
Bases: A base is a substance that produces OH– (aq) ions in
water by dissociation in water:
Strong bases are almost completely ionized in aqueous
solution. (Strong electrolytes)
Examples: Hydroxides of Group 1 (MOH(aq) where M = Li,
Na, K ect…) and Ca, Sr, Ba.*
*Ca(OH)2, Sr(OH)2 & Ba(OH)2 are slightly soluble, but that
which dissolves is present as ions only.
Strong Bases
NH3 acts as a base by reacting with water:
NH3(aq) + H2O(l)
Weak Bases:
Ammonia can also accept H+ from an acid:
NH3(aq) + H+(aq)
NH4+(aq) + OH –(aq)
NH4+(aq)
Weak Bases
K+(aq) + Br– (aq)
Salt + Water (usually)
HA (aq) + MOH(aq)
Acid + Base
Strong acid - Strong base neutralization: HBr(aq)/KOH(aq)
Molecular Equation:
Total Ionic Equation:
HBr(aq) + KOH(aq) KBr (aq) + H2O(l)
H+ (aq) + Br– (aq) + K+(aq) + OH– (aq) + H2O(l)
Net Ionic equation:
H+ (aq) + OH– (aq) H2O (l)
/ / / /
MA(aq) + HOH(l)
Reactions of Acids & Bases: Acid-Base Neutralization
• The ―driving force‖ is the formation of water.
NaOH(aq) + HCl(aq) NaCl(aq) + H2O(liq)
• Net ionic equation
OH-(aq) + H3O+(aq) 2 H2O(l)
• This applies to ALL reactions
of STRONG acids and bases.
Acid-Base Reactions
Total Ionic Equation:
Reactions of weak acids and strong bases:
Molecular Equation:
HC2H3O2(aq) + NaOH(aq) NaC2H3O2(aq) + H2O(l)
HC2H3O2(aq) + Na+(aq) + OH–(aq) Na+(aq) + C2H3O2–(aq) + H2O(l)
Leave in
molecular
form
Net Ionic: HC2H3O2(aq) + OH–(aq) C2H3O2–(aq) + H2O(l)
/ /
Reactions of Acids & Bases: Acid-Base Neutralization
Nonmetal oxides can form acids in
aqueous solutions:
Examples:
CO2(aq) + H2O(s) H2CO3(aq)
SO3(aq) + H2O(s) H2SO4(aq)
Both gases come from the burning
of fossil fuels.
Non-Metal Acids
Metal oxides form bases in aqueous solution
CaO(s) + H2O(l) Ca(OH)2(aq)
CaO in water. Indicator
shows solution is basic.
Bases
Metal carbonate salts react with acids to the corresponding
metal salt, water and carbon dioxide gas.
2HCl(aq) + CaCO3(s) CaCl2(aq) + H2CO3(aq)
H2O(l) + CO2(g)
decomposes
Similarly:
NaCl(aq) + H2O(l) + CO2(g) HCl(aq) + NaHCO3(s)
acid base salt water
Neutralization!!!
Gas-Forming Reactions
Group I metals: Na, K, Cs etc.. react vigorously
with water
2K(s) + 2H2O(l) 2KOH(aq) + H2(g)
Metals & acid:
Some metals react vigorously with acid solutions:
Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)
Gas-Forming Reactions
CaCO3(s) + H2SO4(aq) 2 CaSO4(s) + H2CO3(aq)
Carbonic acid is unstable and forms CO2 & H2O
H2CO3(aq) CO2 + water
(The antacid tablet contains citric acid + NaHCO3)
Gas-Forming Reactions
• Oxidation involves a reactant atom or compound losing
electrons.
• Reduction involves a reactant atom or substance gaining
electrons.
• Neither process can occur alone… that is, there must be
an exchange of electrons in the process.
• The substance that is oxidized is the reducing agent
• The substance that is reduced is the oxidizing agent
Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
oxidized reduced
reducing
agent
oxidizing
agent
Oxidation-Reduction Reactions
• Chemists use oxidation numbers to account for the transfer of electrons in a RedOx reaction.
• Oxidation numbers are the actual or apparent charge on atom when alone or combined in a compound.
1. The atoms of pure elements always have an
oxidation number of zero.
Examples: Mg(s)
Hg(l)
I2(s)
O2(g)
All have an
oxidation number
of zero (0)
Oxidation Numbers
2. If an atom is charged, then the charge is the
oxidation numbers .
Examples: Ion Oxidation Number
Mg2+(aq)
Cl (aq)
Sn4+(s)
22Hg (aq)
+2
1
+4
+2/2 = +1 for each Hg atom
Oxidation Numbers
3. In a compound, fluorine always has an oxidation numbers of 1.
4. 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 1 (Except with F and O present)
6. In compounds, H is +1, except as a hydride
(H : 1)
Oxidation Numbers
Examples:
compound Oxidation Numbers
HF(g) H = +1 F = 1
H2O(l) H = +1 O = 2
OF2(g) O = +2 F = 1
Na2O2(s) Na = +1 O = 1
HCl(g) H = +1 Cl = 1
NaH(l) Na = +1 H = 1
Oxidation Numbers
7. For neutral compounds, the sum of the oxidation numbers
equals zero.
For a poly atomic ion, the sum equals the charge.
Examples:
MgCl2
+2 + 2 × (−1) = 0
3 + 4 × (+1) = +1
4NH
Oxidation Numbers
Determine the oxidation number of iron in the
following compound:
Fe(OH)3
0 = 3 ( 1) ? +
Iron must have an oxidation number of +3!
Oxidation Numbers
In a RedOx reaction, the species oxidized and the
species reduced are identified by the changes in
oxidation numbers :
+ ++ ® + 22Ag (aq) Cu(s) 2Ag(s) Cu (aq)
Oxidation numbers:
+1 0
Oxidation numbers:
0 +2
Since silver goes from +1 to zero, it is reduced.
Since copper goes from zero to +2, it is oxidized.
The reaction is balanced for both mass and charge.
Recognizing a Redox Reaction
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
2
4 2 7
3
3 2
3CH (g) Cr O (aq) 8H (aq)
3CH OH(l) 2Cr (aq) 4H O(l)
Answer:
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
2
4 2 7
3
3 2
3CH (g) Cr O (aq) 8H (aq)
3CH OH(l) 2Cr (aq) 4H O(l)
Answer:
• The carbon in methane (CH4) is oxidized ( 4 to 2)
2
4 2 7
3
3 2
3CH (g) Cr O (aq) 8H (aq)
3CH OH(l) 2Cr (aq) 4H O(l)
Answer:
• Chromium in dichromate is reduced (+6 to +3)
Practice:
Identify the species that is Oxidized and
Reduced by assigning oxidation numbers in the
following reaction.
• The carbon in methane (CH4) is oxidized ( 4 to 2)
• Iron gains 3 electrons
(+3 to 0) oxidation
number change. It is
Reduced.
• Carbon loses 2
electrons (+2 to +4) it
is Oxidized.
Oxidation-Reduction Reactions
REDOX = reduction & oxidation
Corrosion of aluminum
2 Al(s) + 3 Cu2+(aq) 2 Al3+(aq) + 3 Cu(s)
Redox Reactions
In all reactions if a
species is oxidized then
another species must
also been reduced
Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)
Redox Reactions
e e
• 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.
2Ag+(aq) + Cu(s) Cu2+(aq) + 2Ag(s)
Electron Transfer in a Redox Reaction
Metal + acid
Mg + HCl
Mg = reducing agent
H+ = oxidizing agent
Metal + acid
Cu + HNO3
Cu = reducing agent
HNO3 = oxidizing agent
Examples of Redox Reactions
• You have the following items available to you:
Deionized water, pH paper, test tubes various
metal nitrate salts, common acid and base
solutions.
• Suggest a simple test or set of tests for
identifying the unknown substances. Use
proper terminology and write balanced
chemical equations where applicable.
• Justify your answers thoroughly.
Reviewing What You’ve Learned
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