Chemistry Exam 3 Study Guide PDF

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Chemistry Exam 3 Study Guide (Chapters 8,11,13-14) Chapter 8: Chemical Bonding 8.1 Types of Bonds:  Chemical Bond- the force that holds atoms together in a molecule or compound. General Properties of Ionic and Covalent Substances

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Ionic Bonding- the bonding between cations and anions resulting from electrostatic attractions of opposite charges. Covalent Bonding- the bonding between two atoms resulting from the sharing of electrons. Polarity- the separation of electronic charge within a bond (or molecule).

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Nonpolar Covalent Bond- a bond resulting from the equal sharing of electron pairs. Polar Covalent Bond- a bond resulting from the unequal sharing of electron pairs.

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Electronegativity- a measure of the ability of an atom to attract electrons within a bond to itself.

8.2 Ionic Bonding:  Lewis Symbol- a representation of an atom consisting of the symbol for the element surrounded by a number of dots equal to the number of valence electrons. o Transition metals are not usually represented with Lewis symbols.  Crystal Lattice- the repeating pattern of atoms or ions in a crystal.  Ionic Crystal- a solid structure in which the ions are arranged in a regular repeating pattern.

8.3 Covalent Bonding:  Octet Rule- a rule stating that atoms tend to gain, lose, or share electrons to achieve an electronic configuration with eight electrons in the valence shell.  Single Bond- a covalent bond that involves the sharing of one pair of electrons.    

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Lewis Formula- a representation of a molecule or polyatomic ion consisting of the symbols of the component elements, each surrounded by dots representing shared and unshared electrons. Double Bond- a covalent bond that involves the sharing of two pairs of electrons. Triple Bond- a covalent bond that involves the sharing of three pairs of electrons. Resonance Hybrid- an average or composite Lewis formula derive from two or more valid Lewis formulas that closely represents the bonding in a molecule.

Exceptions to the Octet Rule o In a molecule that contains an odd number of valence electrons, one electron must remain unpaired, so one of the atoms cannot have an octet. Generally, the one with an incomplete octet is the one with the lower electronegativity. o Some compounds have more than the eight electrons around the central atom.

o Some atoms participate in covalent bonding but do not have enough valence electrons to form an octet.

8.4 Bonding in Carbon Compounds:  Hydrocarbon- a compound that contains only carbon and hydrogen.

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Alkane- a hydrocarbon that has a straight-chain or branched-chain structure with only single bonds. Alkene- a hydrocarbon that contains one or more localized carbon-carbon double bonds. Alkyne- a hydrocarbon that contains one or more triple bonds. Aliphatic Hydrocarbon- a hydrocarbon that contains localized single, double, or triple bonds. Aromatic Hydrocarbon- a hydrocarbon that contains delocalized bonds made up of alternating single and double bonds that impart special stability. Functional Group- a group of atoms substituted for hydrogen in the formula of a hydrocarbon that gives the compound its characteristic properties; a reactive part of a molecule that undergoes characteristic reactions. Alcohol- an organic compound containing the functional group –OH.

8.5 Shapes of Molecules:  Valence-Shell Electron-Pair Repulsion(VSEPR) Theory- the shapes of molecules result from the tendency for electron pairs to maximize the distance between them to minimize repulsions.  Bond Angle- the angle between the two lines defined by a central atom attached to two surrounding atoms.

Geometric Structures Arising from Different Numbers of Atoms

Arrangement of Electron Domains and Molecular Shapes

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If the bonds are polar and the molecular shape is not symmetrical, then the molecule is polar. If all the bonds are nonpolar or if the molecular shape is symmetrical, the molecule is nonpolar.

Chapter 11: Solutions 11.1 The Composition of Solutions:  Solution- a homogenous mixture of two or more substance uniformly dispersed at a molecular or ionic level.  Solute- the substance being dissolved; usually the component of a solution that is present in the lesser amount.  Solvent- the substance doing the dissolving; usually the component of a solution that is present in the larger amount.  Electrolyte solutions contain a solute that dissociates or ionizes in a solvent, producing ions. Rules Used to Predict the Solubility of Ionic Salts

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Weak electrolytes are usually weak acids and bases. Miscible- the ability of liquids to mix in all proportions.

11.2 The Solution Process:  Ion-Dipole Force- an intermolecular force between an ion and a polar molecule.  When an ionic compound dissolves in water, the process can be summarized as follows: o Ionic bonds in the solute break. o Hydrogen bonds between water molecules break. o Ion-dipole forces form between ions and water molecules.  Entropy- a measure of the tendency for the energy of matter to become more dispersed. 11.3 Factors That Affect Solubility:  Structure

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Temperature o The solubility of ionic solids in water increase with increasing temperature. Gas solubility, in contrast, decreases.

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Pressure o An equilibrium exists between the solute molecules in the gas phase and in the solution. An increase in pressure changes the equilibrium, resulting in increases gas solubility. o Henry’s law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid.

11.4 Measuring Concentrations of Solutions:  Concentration- the relative amounts of solute and solvent in a solution.

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Solubility- a ratio that describes the maximum amount of a solute that dissolves in a particular solvent to form an equilibrium solution under specified conditions. o Saturated Solution- a solution that is in equilibrium with excess solute. o Unsaturated Solution- a solution that contains less than the maximum amount of solute possible in a stable system. o Supersaturated Solution- an unstable solution that contains more dissolved solute than the maximum dictated by the solubility. Percent by Mass- the mass of solute divided by mass of solution and multiplied by 100%. Percent by Volume- the volume of solute divided by volume of solution and multiplied by 100%.

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Mass/Volume Percent- mass of solute divided by volume of solution and multiplied by 100%. Parts per Million- the mass of solute divided by mass of solution and multiplied by 106; in aqueous solutions, it is essentially the milligrams of solute per liter of solution. Parts per Billion- the mass of solute divided by mass of solution and multiplied by 109. Molarity(M)- the number of moles of solute per liter of solution. Molality(m)- moles of solute per kilogram of solvent.

Concentration Units

11.5 Quantities for Reactions That Occur in Aqueous Solution:  Precipitation Reactions

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Example: Titration- a process of determining the concentration of one substance in solution by reacting it with a solution of another substance of known concentration. o To determine the concentration of a measured volume of acid, a known volume of a base of a known concentration is added until some sign of a reaction can be observed, such as a color change of an indicator.

Chapter 13: Acids and Bases 13.1 What are Acids and Bases?  Arrhenius Model of Acids and Bases- a simple model that describes acids as substances that generate H+ ions in solution and bases as substance that generate OH- ions in solution.  Hydronium Ion- an aqueous hydrogen ion; H3O+(aq).  Bronsted-Lowry Theory- a theory that defines an acid as a substance that donates an H+ to another substance in solution and defines a base as a substance that accepts an H+ in solution.  Conjugate Acid- a substance that forms after a base gains an H+ ion.  Conjugate Base- a substance that forms after an acid loses an H+ ion.

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Amphoteric Substance- a substance that can act as either an acid or a base.

13.2 Strong and Weak Acids and Bases:  Strong Acid- an acid that ionizes completely when dissolved in water. Common Strong Acids

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Strong Base- a base that ionizes or dissociates completely when dissolved in water. Common Strong Bases

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Weak Acid- an acid in which only a fraction of the molecules ionizes when dissolved in water; an acid that forms an equilibrium with its conjugate base. Common Weak Acids

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Weak Base- a base in which only a fraction of the molecules ionizes when dissolved in water; a base that forms an equilibrium with its conjugate acid. Common Weak Bases

13.3 Relative Strengths of Weak Acids:  Acid Ionization Constant- an equilibrium constant for the ionization of an acid in water; a value that expresses the strength of a weak acid.  Polyprotic Acid- an acid containing more than one acidic hydrogen.

Chapter 14: Oxidation-Reduction Reactions 14.1 What is an Oxidation-Reduction Reaction?  Oxidation-Reduction Reaction- a reaction in which electrons are transferred. o Oxidation- a chemical process, coupled with reduction, in which an atom increases in oxidation number. o Reduction- a chemical process, coupled with oxidation, in which an atom decreases in oxidation number. o “LEO says GER” Loss of Electrons is Oxidation. Gain of Electrons is Reduction.

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Reducing Agent- a reactant that causes the reduction of an element in another reactant and contains an element that is oxidized during a chemical reaction. Oxidizing Agent- a reactant that causes the oxidation of an element in another reactant and contains an element that is reduced during a chemical reaction.

14.2 Oxidation Numbers:  Oxidation Number- a charge assigned to each atom in a compound according to a set of rules used in an electron accounting system. Rules for Assigning Oxidation Numbers

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The rules are a hierarchy. The first rule that applies takes precedence over any subsequent rules that may apply. For an isolated atom or a molecule that contains only one element, only rule 1 need be applied. An uncombined element, whether occurring as an atom, such as He, or a molecule, such as H2, has an oxidation number of 0. Thus, the oxidation number of sulfur in S, S2, and S8 is 0. A monatomic ion has an oxidation number equal to its ionic charge, according to rule 2. For example, the oxidation number of lithium in Li+ is 1+. Similarly, the oxidation number is 1− for chlorine in Cl−. Rule 2 can be applied whenever oxidation numbers have been assigned to all but one element. Rule 5 uses position in the periodic table to cover situations that are not handled by one of the other rules. In general, the more electronegative element in a compound is assigned the oxidation number that would be its charge if it were a monatomic ion. To assign oxidation numbers to binary compounds, handle one of the elements with an appropriate rule and the other with rule 2.

14.4 Balancing Simple Oxidation-Reduction Equations: 

Balancing oxidation-reduction reactions begins with the identification of which element is oxidized and which is reduced. o Half-reactions account for changes in oxidation number and describe the oxidation and reduction processes that occur at each electrode. o To balance half-reactions, changes in oxidation number, charge, and numbers of atoms must be balanced. o An oxidation-reduction reaction equation is balanced when the number of atoms and charges are the same on both sides of the equation and the number of electrons gained equals the number lost....