Chapter 4
Alex
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Describe the general properties of a solute, solvent, and solution
a. Solute-the substance that gets dissolved
b. Solvent- the substance that does the dissolving
 Present in the greatest amount
c. Solution- homogeneous mixture
 Has 2 or more substances
Classify common compounds as strong electrolytes, weak electrolytes, or nonelectrolytes
a. Electrolyte- a substance that, when dissolved in water, results in a solution that can conduct electricity
b. Strong Electrolytes- Strongest conductivity when the solute completely falls apart into ions (dissociates)
 All soluble ionic compounds
 Strong acids (H2SO4, HNO3, HBr, HI, HCl, HCO3, HClO4)and strong bases (LiOH, NaOH, KOH,
RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2)
c. Weak Electrolytes- weak conductivity (still conductive) when the solute only partially dissociates
 Weak acids and weak bases
d. Nonelectrolytes- No conductivity, solute doesn’t dissociate
 Covalent (molecular) compounds EX) sugar, ethanol, sulfur dioxide, etc
Describe precipitation reactions using terms solubility and precipitate
a. Precipitation reaction- a reaction that results in the formation of an insoluble (cannot be dissolved) product, or
precipitate
b. Precipitate- an insoluble solid that separates from the solution
c. Solubility- Can be dissolved
d. Stoichiometry for a Precipitation Reaction Identify the species and write a molecular or net ionic equation, calculate the moles of reactants,
determine which is limiting and calculate the moles of product, convert to desired unit
Classify common ionic compounds as soluble or insoluble
a. Soluble Compounds- Compounds containing alkali metal ions (Li+, Na+, K+, Rb+, Cs+) and the ammonium ions
(NH4+), Nitrates (NO3-), acetates (CH3COO-), bicarbonates (HCO3-), Chlorates (ClO3-), and perchlorates (ClO4), Halides (Cl-, Br-, I-), Sulfates (SO42-) exceptions include halides of AG+, Hg22+, and Pb2+ and Sulfates of Ag+,
Ca2+, Sr2+, Ba2+, Hg22+, and Pb2+
b. Insoluble Compounds- Carbonates (CO32-), phosphates (PO43-), chromates (CrO42-), sulfides (S2-), Hydroxides
(OH-), exceptions include compounds containing alkali metal ions and the ammonium ion and compounds
containing alkali metal ions and the Ba2+ ion
Predict the resulting products and write the molecular equation, ionic equation, and net ionic equation; and identify
spectator ions given the reactants of a chemical reaction
a. Molecular Equation- lists the reactant and products in their molecular form
b. Ionic Equation- shows all strong electrolytes dissociated into their ions. It more accurately reflects the species
that are found in the reaction mixture
c. Net ionic Equation- cross out anything that doesn’t change from the left side of the equation to the right. What
is left is the net ionic equation
d. Spectator ions- are anything that didn’t change and deleted from the net ionic
e. Metathesis Reaction- double replacement reactions where positive ions swap places with each other
Explain the difference between Arrhenius acids and bases and Brønsted acids and bases
a. Arrhenius’ definition of acids and bases is limited to aqueous solutions- acids: substances that ionize in water to
produce H+ ions, bases: substance that ionize in water to produce OHb. Brønsted’s definition of acids and bases is more general, it doesn’t require them to be in aqueous solutionsAcids: proton donor, bases: proton acceptor
Compare and contrast the properties of acids and bases
a. Acids-sour taste, cause color changes in paint dyes, react with certain metals to produce hydrogen gas, react
with carbonates and bicarbonates to produce carbon dioxide gas, and aqueous acid solutions conduct electricity
b. Bases- bitter taste, feel slippery, cause color changes in plant dyes, and aqueous base solutions conduct
electricity
By using a chemical reaction, explain why ammonia (NH3) is classified as a Brønsted base
a. NH3(aq)+H2O(l)
NH4+(aq)+OH- (aq)
The ionization of ammonia in water forms the ammonium ion and the hydroxide ion. NH3 is considered a
Brønsted base because it can accept an H+ ion.
9.
Predict the products formed by acid-base neutralization reactions
a. The acid donates a proton (H+) to the base so generally the products of an acid-base neutralization reaction are
a salt and water
10. Describe what factor results in an oxidation-reduction reaction
a. In order to have an oxidation-reduction reaction there has to be a transfer in electrons
 Oxidation- an atom or ion loses electrons
 Reduction- an atom or ion gains an electron
 Can’t have an oxidation-reduction reaction without both an oxidation and reduction
11. Identify oxidation half-reactions and reduction half reactions
a. All oxidation-reduction equations are thought to have 2 halves one oxidation (loss of electrons) and the other
reduction (gain in electrons)
b. Example: 2Mg(s) + O2(g)
2MgO
2Mg2+ + 2O22MgO2Mg
2Mg2+ + 4e
Oxidation Half
O2+4e
2O2Reduction Half
12. Assign oxidation numbers to elements in compounds and ions
a. Oxidation Number- the number of charges the atom would have in a molecule (or an ionic compound) if
electrons were transferred completely
b. Example:
0
0
+1 -1
H2(g) + Cl2(g)
2HCl(g)
 There is no charge on the atoms in the reactants thus their oxidation number is 0
Page Reference: Chapter 4 pages 118-157, practice questions and key items page 157- 169
Multiple Choice Questions:
1.
Which of the following is correctly characterized?
a) HClO4- Strong electrolyte
b) Ca(OH)2- weak electrolyte
c) H2O- strong electrolyte
d) Li3PO4- nonelectrolyte
Answer- A because HClO4 is a strong acid which causes it to be a strong electrolyte. B is incorrect because Ca(OH)2 is a
strong bases which means it is a strong electrolyte. Water would not dissociate to a great degree causing it not to be a strong
electrolyte. Li3PO4 is a soluble salt so it would dissociate completely causing it to be a strong electrolyte.
Caroline
12. Assign oxidation numbers to elements in compounds and ions.
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Elements in their elemental form have an oxidation number of
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The oxidation number of a monatomic ion is the same as its charge
Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation
number of −1
Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal
Fluorine always has an oxidation number of −1
The other halogens have an oxidation number of −1 when they are negative; they can have
positive oxidation numbers, however, most notably in oxyanions
The sum of the oxidation numbers in a neutral compound is 0
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The sum of the oxidation numbers in a polyatomic ion is the charge on the ion
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13. Categorize redox reactions in terms of combination reactions, decomposition reactions,
displacement reactions, and disproportion reactions.
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Combination: a reaction that combines elements to form a chemical compound. The general
equation is: A + B→ AB
Decomposition: a reaction involving the breakdown of a chemical compound into the individual
elements. The general equation is: AB → A + B
Displacement reactions are broken down into two categories: single replacement and double
replacement.
Single replacement: a reaction that involves the replacing of an element in the reactants with
another element in the products. The general equation is: A + BC → AB + C
Double replacement: a reaction between two compounds where the positive ion of one
compound is exchanged with the positive ion of another compound. The general equation is: AB
+ CD → AD + CB
Disproportion reactions: a reaction that begins with an element in one oxidation state but produce
products in which that element appears in two different oxidation states. The general equation is:
2A → A' + A"
14. Predict the results of a chemical reaction involving metals given the activity series
(electrochemical series).
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An activity series of elements is often used to determine if A will displace B in a single
replacement reaction
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A less reactive metal cannot replace a more reactive metal. Therefore, no reaction would occur
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Reactivity series of metals: potassium > sodium > calcium > magnesium > aluminium > zinc >
iron > tin > lead > copper > silver > gold > platinum
15. Predict the results of a chemical reaction involving halogens given the halogen activity series.
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A less reactive halogen cannot replace a more reactive halogen
Reactivity series of halogens: fluorine > chlorine > bromine > iodine
16. Calculate the molarity of a solution given the mass (number of moles) of a solute and the volume
of solution.
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Molarity: the concentration of a solution measured in moles of solute per liter of solution.
Formula: Molarity = moles of solute/ liters of solution
To change the mass of a solute to moles, divide the mass of the solute by the mass of one mole of
the solute. Then, plug this number into the equation along with the given liters of solution
17. Describe the method for preparing a specific molar solution given the volume of solution
required and the solute to be used.
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To make a solution: add solute (with number of moles known) to volumetric flask and add
enough water to meet needed volume
18. Describe in detail how to prepare a specific dilute solution given a known stock solution using
dilution techniques.
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Stock solution: a solution of known concentration (molarity) used to make a more dilute solution
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The known molarity and volume of the stock solution and the desired molarity of the final
solution can be plugged into the below formula. The volume (V2) obtained from this calculation
is the required volume for the final solution. Add water to the stock solution until this volume is
met. Formula: M1 V1 = M2 V2
19. Deduce the mass percent of specific ions present in an original solution given the results of a
gravimetric analysis.
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Gravimetric analysis: a process involving the isolation of an ion in solution by a precipitation
reaction, filtering, washing the precipitate free of contaminants, conversion of the precipitate to a
product of known composition, and finally weighing the precipitate and determining its mass by
difference. From the mass and known composition of the precipitate, the amount of the original
ion can be determined.
Use stoichiometry to determine the mass of the ion being analyzed. (mass of precipitate into
moles by dividing by molar mass of percipitate, moles of percipitate in proportion with moles of
analyte by using a balanced chemical equation, moles of analyte into grams by multiplying
moles by the mass of one mole of analyte)
Find mass percent of analyte by dividing the mass of the anayte by the mass of the unknown
20. Use the terms titration, standard solution, equivalence point, and indicator to describe
quantitative studies of acid-base neutralization reactions.
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An acid-base titration is the determination of the concentration of an acid or base by exactly
neutralizing the acid or base with an acid or base of known concentration, which is the standard
solution
The titrant is the standard, known solution and the analyte is the solution with the unknown
concentration. The analyte is put into a flask with an indicator, like phenolphthalein, that will
show a color change when end point is met. Then, the titrant is titrated into the analyte. When the
end point is met, it is possible to subtract the initial volume from the final volume to yield the
amount of titrant used to reach the end point.
The equivalence point of the reaction is the point at which equivalent amounts of the reactants
have reacted and neutralized each other. If you know the volume and concentration of the titrant,
the standard solution, at the equivalence point, you can calculate the concentration of a base or
acid in the unknown solution.
21. Determine the concentration of an unknown acid (base) given the results of an acid-base
titration.
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Take the amount of known base (acid) in grams used and convert it into moles by dividing my
the molar mass of that substance. Next, find out how many moles it takes for the base (acid) to
neutralize the acid (base) by using a chemical formula. Then, convert the moles of base (acid)
into moles, and plug this into the molarity equation to find out the concentration. The liters
would be given to you in the problem, and you just solved for the moles
22. Predict the amount (mass, moles, or volume of solution) of an acid (base) required to neutralize a
base (acid).
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Identify the acid and base. Then, write a balanced chemical equation, convert the moles of titrant
to analyte using the balanced chemical equation, and use the molarity equation to calculate the
molarity or volume of the analyte.
23. Predict the volume of an oxidizing (reducing) reagent solution provided that the net ionic
equation is given.
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Find the number of moles of the known solution by using the molarity equation (M=mol/L).
Then, find the stoichiometric equivalence between the reagents using the net ionic equation.
Now that the moles of the oxidizing reagent is known, it can be plugged into the molarity
equation to solve for liters since the concentration is given
Important Equations Needed:
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Molarity = Moles of Solute/ Liters of Solution
M1 V1 = M2 V2
Refer to Page 141, figure 4.16 for the activity series chart; Page 146, figure 4.18 for making solutions;
Page 150, figure 4.20 for gravimetric analysis; Page 152, figure 4.21 for titrations.
Problems:
1. What is the oxidation number of B in NaBH4?
a. 3 b.2 c.5 d.1 (The answer is a because Na is +1, H4 is -4, and the compound has to have a total
of 0)
2. How many grams of K2Cr2O7 are needed to prepare a 250-mL solution whose concentration is
2.16 M?
a. 294.2g b. 159g c. 209g d. 540g (The answer is b because first the number of moles of
K2Cr2O7 is determined by M=mol/L→2.16=mol/.250 → mol=.540 then multiply by the molar
mass (294.2g) to get the answer in grams)