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Unformatted text preview: Chapter 4: Objectives
4.1. Discuss the idea of a chemical solution and the components of a
solution. Emphasize that water is the most common solvent in a solution.
4.2. Introduce molarity (M) as a calculation of a way describing the
atomic/molecular composition of a solution.
4.3. Emphasize acid-base, precipitation, and redox reactions as common
acidsolution chemistry reactions.
4.4. Include examples of solution stoichiometry calculations that will be
emphasized throughout the semester in both lecture and laboratory.
laboratory. 8/31/2009 Zumdahl Chapter 4 1 Water, the Most Common Solvent Hydration or
Dissolution: Two (or more) substances spread out, or disperse, into each
other at the level of individual atoms, molecules, or ions. Solution:
In general, the solute and solvent can be any combination of solid (s), liquid (l), and
gaseous (g) phases. However, we will be primarily interested in liquid solutions.
8/31/2009 Zumdahl Chapter 4 2 1 Ionic Compounds in Water
All ionic compounds are solids (salts) in the range of temperature in which water
is a liquid. They have rigid lattices in which strong (Coulomb) forces hold the
constituent ions in place.
The high melting points of ionic compounds indicate that a good deal of energy
must be supplied to destroy the lattice and produce a liquid (molten) form in which
the ions move more freely. Water is a polar molecule. When water interacts with an ionic
compound such as NH4NO3, the ionic compound dissociates into a
positive ion (NH4+, a cation) and a negative ion (NO3−, an anion).
anion). 8/31/2009 Zumdahl Chapter 4 3 Ionic Compounds in Water
NH4NO3(s) + H2O(l) → NH4+(aq) + NO3–(aq) 8/31/2009 Solvation of the ions by water molecules provides the
Zumdahl Chapter 4
energy gain needed to break apart the crystal lattice. 4 2 Dissolution of One Liquid in Another Liquid
Pairs of liquids that mix in any proportion are termed miscible.
Liquids that do not mix in any proportion are termed immiscible.
Liquids that are partially miscible mix in some proportions but not in others.
"Like Dissolves Like"
Substances with similar intermolecular attractive forces tend to be soluble in one
▬► Polar liquids tend to be miscible in other polar liquids & non-polar
nonliquids tend to be miscible in other non-polar liquids.
Ethanol: CH3CH2OH 8/31/2009 n-octane: CH3CH2CH2CH2CH2CH2CH2CH3 Zumdahl Chapter 4 5 Electrolytes and Non-Electrolytes
Mainly ionic compounds such as potassium sulfate (K2SO4)
and sodium chloride (NaCl) which, as a solute, increase the
electrical conductance of water when they dissolve.
Do not increase the conductance of water when
~100% dissociated in aqueous
solution. Increase conductivity
much more than weak electrolytes.
electrolytes. 8/31/2009 Zumdahl Chapter 4 6 3 The Composition of Solutions
The actual amount of solute in any given volume of solution depends
on how concentrated or dilute the solution happens to be.
This is expressed by a ratio:
The concentration (c) of a solute in a solution equals the chemical
amount (i.e., no. of moles) of the solute (n) divided by the volume (V)
of the entire solution. csolute = molarity
8/31/2009 ≡ ≡ M
Zumdahl Chapter 4 7 Calculate the molarity of a solution made by dissolving 10.0
g of Al(NO3)3 in enough water to make 250.0 mL of solution.
= 0.0469 moles of Al(NO3)3
= 0.250 L = 0.188 mol L-1 cAl(NO3)3 =
8/31/2009 Zumdahl Chapter 4 8 4 Types of
Reactions 8/31/2009 1. Precipitation Reactions (Chapter 13)
– Ionic Equations and Net Ionic Equations
– Predicting Precipitation Reactions
2. Acids and Bases (Chapters 12 & 13)
– Arrhenius Acids and Bases Theory
– Strong and Weak acid
– Naming Acids
– Weak Bases
– Modifying the Arrhenius model
– Acid-Base Titrations
Acid3. Oxidation−Reduction (Redox) Reactions
– Oxidizing and Reducing Agents
– Oxidation Number
– Types of Redox Reactions
» Combination and Decomposition
» Displacement Reactions
Zumdahl Chapter 4
» Disproportionation Approximate Solubility Designations
• More than of solute per 100 g water • Slightly Soluble
• of solute per 100 g water • Insoluble
• Less than 8/31/2009 of solute per 100 g water Zumdahl Chapter 4 10 5 Predicting Dissolution and Precipitation Reactions CHEM1310 McKelvy Lecture 8/31/2009 Zumdahl Chapter 4 11 Ionic Equations and Net Ionic Equations
BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq)
Ba2+(aq) + 2 Cl−(aq) + 2 K+(aq) + SO42−(aq) →
BaSO (s) + 2 K+(aq) + 2 Cl−(aq)
4 Ba2+(aq) + 2 Cl−(aq) + 2 K+(aq) + SO42−(aq) →
Spectator ions BaSO4(s) + 2 K+(aq) + 2 Cl−(aq)
(aq) Net ionic equation 8/31/2009 Zumdahl Chapter 4 12 6 8/31/2009 Zumdahl Chapter 4 13 Write a net ionic equation to represent the formation of the
precipitate observed when aqueous solutions of CaCl2 and
NaF are mixed. Identify the spectator ions in this process. CaCl2(aq) ≈ Ca2+(aq) + 2 Cl−(aq)
aq) NaF(aq) ≈ Na+(aq) + F−(aq)
aq) CaCl2(aq) + NaF(aq) →
Ca2+(aq) + 2 Cl−(aq) + Na+(aq) + F−(aq) →
aq) Spectator ions CaF2(s) + Na+(aq) + 2 Cl−(aq)
aq) Net ionic equation
8/31/2009 Zumdahl Chapter 4 14 7 Acids and Bases and Their Reactions (more in Chapter 12)
1. Arrhenius Acids and Bases
• Acids are H+ donors
• Bases are OH− donors
2. Arrhenius-Broadened Definition
Arrhenius• Acids increase H+ concentration
• Bases increase OH− concentration
3. Brønsted-Lowry Acids and Bases
Brø nsted• Acids donate H+
• Bases accept H+
8/31/2009 Zumdahl Chapter 4 Arrhenius
Nobel Prize 15 The Person Behind the Science Svante Arrhenius (1859-1927)
– Specialized in the conductivities of electrolytes.
Postulated that in forming a solution, a salt
dissociates into charged particles.
– Formulated the concept of activation energy. The
Arrhenius equation gives the quantitative basis of
the relationship between the activation energy and
the rate at which a reaction proceeds (Chapter 17).
– In 1884 Arrhenius proposed definitions for acids
and bases. Acids produce hydrogen ions in solution
and bases produce hydroxide ions in solution.
Moments in a Life
– Awarded 1903 Nobel Prize in chemistry
– The lunar crater Arrhenius is named after him
8/31/2009 Zumdahl Chapter 4 16 8 Acid/Base Neutralization Reactions H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) →
H2O(l) + Na+(aq) + Cl−(aq)
aq) (net ionic equation) Neutralization Concept (at the equivalence point)
= 8/31/2009 Zumdahl Chapter 4 17 Titration of an Acid with a Base
Standard Solution 8/31/2009 Indicator Equivalence point
Stoichiometric point Zumdahl Chapter 4 18 9 Strong Acids and Bases
Strong acids and bases are strong electrolytes which
dissociate essentially completely in water.
Example: n HCl(g) → n H+(aq) + n Cl−(aq)
7 Common Strong acids: Hydrochloric acid
Perchloric acid Chloric acid Hydrobromic acid Sulfuric acid Hydroiodic acid Nitric acid Example: n NaOH(s) → n Na+(aq) + n OH−(aq)
hydroxide salts are strong bases
8/31/2009 Zumdahl Chapter 4 19 Hydrofluoric acid (HF) Many Weak Acids
Oxalic acid (H2C2O4) Acetic acid (CH3COOH) Formic acid (HCOOH) Phosphoric acid (H3PO4) CH3COOH(aq) → CH3COO−(aq) + H+(aq) (dissociation of acetic acid)
CH3COOH(aq) + NaOH(aq) → NaCH3COO(aq) + H2O(l) (overall equation) CH3COOH(aq) + Na+(aq) + OH−(aq) →
Na+(aq) + CH3COO−(aq) + H2O(l)
O(l (ionic equation) (net ionic equation)
8/31/2009 Zumdahl Chapter 4 20 10 Compute the molarity of a solution of sodium hydroxide if 25.64 mL of
solution must be added to a solution containing 0.5333 g of KHC8H4O4
(potassium hydrogen phthalate) to reach the equivalence point . nKHP =
= 2.611 x 10-3 mol KHP
nNaOH = (at the equivalence point)
point) [NaOH] =
Concept = 0.1018 mol L-1
nbase = nacid
nNaOH = nKHP 8/31/2009 (at the equivalence point)
point) Zumdahl Chapter 4 21 Oxidation − Reduction Reactions
Redox: an extensive and important class of reactions that is
characterized by the transfer of electrons.
2 Mg(s) + O2(g) → 2 MgO(s)
2x 2e2e ↑gain
= 1 x 2 x 2e- Magnesium is
: it gives up electrons as the charge
on its atoms increases from zero to +2.
: it gains electrons as the charge on its
atoms decreases from zero to -2 (i.e., becomes more negative).
8/31/2009 Zumdahl Chapter 4 22 11 8/31/2009 Zumdahl Chapter 4 23 Oxidation Numbers (also called oxidation states) are determined
for the atoms in covalently bonded compounds by applying the
following set of simple rules:
1. In the free elements, each atom has an oxidation number of Zero
no matter how complicated the element.
i.e., Na (s), Fe(s), K(s), Mg(s) H2, O2, N2, Cl2 all have zero oxidation states
Fe(s), K(s), Mg(s) 2. The oxidation number of the atoms in a neutral molecule must add
up to zero; those in an ion must add up to the charge on the ion.
3. Alkali metal (Group I) atoms have oxidation number +1, and alkaline
earth (Group II) atoms have oxidation number +2 in their compounds;
atoms of Group III elements usually have oxidation number +3 in their
8/31/2009 Zumdahl Chapter 4 24 12 Oxidation Numbers (con’t)
4. Fluorine always has an oxidation number of -1 in its compounds. The
other halogens have oxidation number -1 in their compounds, except
in compounds with oxygen and with other halogens, in which they can
have positive oxidation numbers.
5. Hydrogen is assigned an oxidation number of +1 in its compounds,
except in metal hydrides such as LiH, in which rule 3 take precedence
and hydrogen has an oxidation number of -1.
6. Oxygen is assigned an oxidation number of -2 in compounds. There
are two exceptions: in compounds with fluorine, rule 4 takes
precedence, and in compounds that contain O—O bonds, rules 3 − 5
take precedence. Thus, the oxidation number of oxygen in OF2 is +2; in
peroxides (e.g., HO-OH, FO-OF and NaO-ONa), its oxidation number is
±1, and in superoxides (e.g., KO-O, FO-O), its oxidation number is ±½.
KO- FO8/31/2009 Zumdahl Chapter 4 25 Oxidation and Reduction
An atom is oxidized (_____ electrons) if its
oxidation number increases in a chemical reaction
Na · → Na+ + e- Loss of a valence electron An atom is reduced (
electrons) if its
oxidation number decreases.
8/31/2009 - e - + Cl → Cl− Chapter 4 Gain of a valence electron
26 13 Oxidation State (Oxidation Number)
Can assign for Ionic and Covalent bonded compounds
Not formal electric charges, rather what the charge would be if
the compound were ionic
Range from -3 to +7
Examples 4. VCl4 1. CrO3
0 = (Cr * 1) + (O * 3) 5. Mn2O7
Net Charge = ∑ (Oxid State x N)
where " Oxid State" is the oxidation number 3. Mn3N2 of an element and N is the number of atoms
of that element
8/31/2009 Zumdahl Chapter 4 27 Examples (con’t)
-1 = (Mn * 1) + (O * 4)
-1 = (Mn *1 ) + (-2 * 4)
(- Net Charge = ∑ (Oxid State x N)
where " Oxid State" is the oxidation number
of an element and N is the number of atoms
of that element
8/31/2009 Zumdahl Chapter 4 28 14 Oxidation
Number Change Change Term
Oxidation Increase Loss of
Electrons Reduction Decrease Gain of
Electrons Oxidizing Agent, does
the oxidizing Decrease Picks Up
electrons Reducing Agent, does
the reducing Increase Supplies
Electrons Substance Oxidized Increase Loses
Electrons Substance Reduced Decrease Gains
Electrons 8/31/2009 Zumdahl Chapter 4 29 Oxidation-Reduction reactions
Change Electron Change Oxidizing Agent, does the oxidizing Decrease Picks Up electrons Reducing Agent, does the reducing Increase Supplies Electrons Substance Oxidized Increase Loses Electrons Substance Reduced Decrease Gains Electrons Term ½ O2 H2 → H2O ½ Cl2 + Na → NaCl 2 H+ 8/31/2009 + + Mg → Mg++ + H2 Zumdahl Chapter 4 30 15 A summary of an
oxidationprocess, in which M
is oxidized and X is
reduced. Li F Li + 8/31/2009 e- + e- F Zumdahl Chapter 4 31 Chapter 4
Types of Chemical Reactions and Solution Stoichiometry
4.2 Water, the Common Solvent
The Nature of Aqueous Solutions: Strong and Weak
4.3 The Composition of Solutions
4.4 Types of Chemical Reactions
4.5 Precipitation Reactions
4.6 Describing Reactions in Solution
4.7 Selective Precipitation
4.8 Stoichiometry of Precipitation Reactions
4.9 Acid-Base Reactions
Acid4.10 Oxidation-Reduction Reactions
Oxidation4.11 Balancing Oxidation-Reduction Equations (skip)
Oxidation4.12 Simple Oxidation-Reduction Titrations (skip)
Oxidation8/31/2009 Zumdahl Chapter 4 32 16 ...
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This note was uploaded on 10/06/2009 for the course CHEM 1101 taught by Professor Bottomley during the Fall '08 term at Georgia Institute of Technology.
- Fall '08