Chapter 4 Reactions in Aqueous Solutions
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Transcript of Chapter 4 Reactions in Aqueous Solutions
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Chapter 4Reactions in Aqueous Solutions
Many chemical and almost all biological reactions occur in the aqueous medium
Substances (solutes) that dissolve in water (solvent) can be divided into two categories: Electrolytes Non-Electrolytes
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Three Major Types of Reactions Precipitation Reaction – the product an
insoluble substance separates from the solution
Acid/Base Reactions – A proton transfer from an acid to a base
Oxidation/Reduction (Redox) “the bane of the AP Test” – Electrons are transferred from a reducing agent to an oxidizing agent
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Solution Stoichiometry Quantitative studies with known
concentrations (Molarity) of solutions Gravimetric Analysis Titrations
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General Properties of a Solution Solution – a homogenous mixture of two or
more substances Solution may be gaseous (air), solid (alloy) or
Liquid (salt water)
In this chapter we will deal only with aqueous solutions Most common Solvent - Water
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Electrolytes versus Nonelectrolytes Electrolytes – Ionic compounds that
completely or partially dissociate in solution with the ability to pass electric current in solution Acids/Bases will ionize in solution, therefore
electricity can be conducted Non-Electrolytes – Molecular compounds
that do not dissociate in solution, therefore no electric current can be pass
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Ionic Compounds in Solution Water is a great solvent for ionic
compounds because it is polar, the positive end attracts the Negative Ion and vice versa
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Acids and Bases as Electrolytes Some acid/bases competely dissociate in solution
HCl HNO3 H2SO4 Ba(OH)2 NaOH
While others only partially dissociate CH3COOH HF HNO2 HN3
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Writing Partial Dissociation Equations
Partial dissociation equations are written with a double arrow, indicating a reversible reaction
Write partial dissociation for CH3COOH
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Precipitation Reactions A double replacement reaction (metathesis)
in which a product is insoluble
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Solubility RulesIn water at 25 Degrees
All common compounds of Group I and ammonium ions are soluble.
All nitrates, acetates, and chlorates are soluble. All binary compounds of the halogens (other than F) with
metals are soluble, except those of Ag, Hg(I), and Pb. All sulfates are soluble, except those of barium, strontium,
calcium, lead, silver, and mercury (I). The latter three are slightly soluble.
Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble.
Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.
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Soluble or Insoluble at 25 Degrees Celsius in Water
PbSO4
BaCO3
Li3PO4
FeS Ca(OH)2
Co(NO3)3
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Net Ionic Equations Write the correctly balanced equation and
decide on state of each product Write free state of all ions and insoluble
product Cancel out spectator ions – anyone not part
of the reaction Check charges and balancing in net ionic
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Practice Net IonicPredict, Balance and write net ionic
1) Lead Nitrate and Potassium Iodide
2) Barium Chloride and Sodium Sulfate
3) Potassium Phosphate and Calcium Nitrate
4) Aluminum Nitrate and Sodium Hydroxide
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Acid – Base Reactions Acids react with metal such as Zn, Mg and
Fe to produce hydrogen gas Acids react with carbonates and
bicarbonates to produce carbon dioxide gas, water and the salt
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Bronsted Acid and Bases Bronsted Acid is a proton donor Bronsted Base is a proton acceptor
HCl (aq) H+ (aq) + Cl-(aq)
In water the H+ attracts to the water molecule producing the hydronium ion
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Monoprotic Acids Each unit of acid yields one hydrogen ion
upon ionization
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Diprotic Acids Each unit of the acid gives up two hydrogen
ions in two separate steps (they strip)
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Triprotic Acids Yield three hydrogen ions in three separate
steps (they strip)
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Bronsted Acid is a proton donorBronsted Base is a proton acceptor
Classify each of the following as an Bronsted acid or Bronsted base, explain your reasoning based on the definition HBr
SO-2
4
HI
HCO-
3
NO2
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Neutralization ReactionAcid and Base will form Salt and
Water
Write the net ionic for the following Hydrochloric acid and Sodium Hydroxide Sulfuric acid and Aluminum Hydroxide
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Acid – Base Reactions Leading to Formation of a Gas
Certain Salts – Carbonates, bicarbonates, sulfites and sulfides react with acids to form gaseous products
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Oxidation Numbers Oxidation Reaction – refers to half reaction that
involves loss of electrons Reduction reaction – refers to a half reaction that
involves the gain of electrons The extent of oxidation in a redox reaction must
be equal to the extent of reduction; that is the number of electrons lost by a reducing agent must be equal to the number of electrons gained by an oxidizing agent
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The half-reactions of a redox reaction or oxidation-reduction reaction
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Oxidation Number The number of charges the atom would
have in a molecule if electrons are transfer completely
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The convention is that the cation is written first in a formula, followed by the anion.
For example, in NaH, the H is H-; in HCl, the H is H+. The oxidation number of a free element is always 0. The atoms in He and N2, for example, have oxidation numbers of 0. The oxidation number of a monatomic ion equals the charge of the ion. For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3. The usual oxidation number of hydrogen is +1. The oxidation number of hydrogen is -1 in compounds containing elements that are
less electronegative than hydrogen, as in CaH2. The oxidation number of oxygen in compounds is usually -2. Exceptions include OF2, since F is more electronegative than O, and BaO2, due to
the structure of the peroxide ion, which is [O-O]2-. The oxidation number of a Group IA element in a compound is +1. The oxidation number of a Group IIA element in a compound is +2. The oxidation number of a Group VIIA element in a compound is -1, except
when that element is combined with one having a higher electronegativity. The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl. The sum of the oxidation numbers of all of the atoms in a neutral compound is
0. The sum of the oxidation numbers in a polyatomic ion is equal to the charge of
the ion. For example, the sum of the oxidation numbers for SO42- is -2.
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Assign oxidation numbers to all the elements in the following compounds
Na2O
HNO2
Cr2O7-2
PF3
MnO4-
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Arrange the following species in order of increasing oxidation number of the sulfur atoms
H2S
SO2
SO3
S8
H2SO4
S-2
HS-
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Concentration Molarity = moles of solute/liters of solution
Example: How many grams of potassium dichromate are required to prepare a 125ml solution whose concentration is 1.83M
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Concentration In a biochemical assay a chemist needs a
to add 4.07g of glucose to a reaction mixture. Calculate the volume in milliliters the volume of a 3.16M glucose she should use
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Dilution of Solutions The procedure of making a less
concentrated solution from a high concentration solution
MiVi = MfVf
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Dilution Problem Describe hou you would prepare 2.50 * 102
ml of a 2.25M H2SO4 solution, starting with a 7.41 M stock solution of H2SO4
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Dilution Problem #2 How would you prepare a 200ml of a .866M
KOH solution, starting with 5.07M stock solution
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Acid – Base Titrations In a titration a solution of an accurately
known concentration, called the standard is added gradually to another solution of unknown until reaction is neutralized (equivalence point)
Indicators are used to color the reaction when it is complete
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Titration Problem In a titration experiment, a student finds
that 25.46ml of a NaOH solution is needed to neutralize .6092g of KHP. What is the concentration of the NaOH solution?
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Titration Problem #2 How many milliliters of a .836M NaOH
solution is needed to neutralized 25ml of a .335M of H2SO4?
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Solution Stoichiometry
When sodium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 150ml 3M of silver nitrate?
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1. When Magnesium chloride reacts with silver nitrate, silver chloride precipitates. What mass of silver chloride is produced from 4.5M in 250ml of silver nitrate? What is the name of the other product of the reaction?
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