Ch4 outline - Summary Adv Freshm Chem I PDF

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notes for chapter four with robinson...

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Chapter 4: Reactions in Aqueous Solution 4.1 General Properties of Aqueous Solutions ● Solution- a homogeneous mixture of two or more substances ● Solvent- the substance present in the greatest quantity ● Solute- substances said to be dissolved in the solvent ● When NaCl is dissolved in water ○ Water is the solvent ○ NaCl is the solute Electrolytes and Nonelectrolytes ● Electrolyte- a substance whose aqueous solutions contain ions ● Nonelectrolyte- a substance that does not form ions in solution How Compounds Dissolve in Water ● Dissociation- separates from structure and disperses throughout the solution ● Water is a very effective solvent for ionic compounds ○ Although it is an electrically neutral molecule, the O atom is rich in electrons and has a partial negative charge, while each H atom has a partial positive charge ○ The lowercase Greek letter delta (ẟ ) is used to denote partial charge ○ A partial negative charge→ delta minus ẟ○ A partial positive charge→ delta plus ẟ+ ○ Cations are attracted by the negative end of water ○ Anions are attracted by the positive end ● Solvated- solute is completely surrounded by solvent molecules ● Solvation- the clustering of solvent molecules around a solute molecule ○ Helps stabilize the ions in solution ○ Prevents cations and anions from recombining ○ Ions become dispersed uniformly throughout the solution ● Most molecular compounds are nonelectrolytes ● A few molecular substances do have aqueous solutions that contain ions ○ Acids are the most important ○ When HCl(a) dissolves in water to form HCl(aq) ○ The molecules ionizes; dissociate into H+(aq) and Cl-(aq) ions Strong and Weak Electrolytes ● Strong electrolytes- those solutes that exist in solution completely or nearly completely as ions ○ All water soluble ionic compounds ○ A few molecular compounds ● Weak electrolytes- those solutes that exist in solution mostly in the form of neutral molecules with only a small fraction in the form of ions ● Just because an electrolyte dissolves easily does not mean it is a strong electrolyte ● Just because an electrolyte doesn’t dissolve easily does not mean it is a weak electrolyte ● Half-arrows in chemical reactions→ the reaction is significant in both directions ○ Represent reactions that go both forward and backward to achieve equilibrium

○ Such as the ionization of weak electrolytes Chemical equilibrium- a state in which the relative numbers of each type of ion or molecule in the reaction are constant over time ● Water soluble ionic compounds are strong electrolytes**** ○ Ionic compounds containing the ammonium ion are exceptions 4.2 Precipitation Reactions ● Precipitation reactions- reactions that result in the formation of an insoluble product ● Precipitate- an insoluble solid formed by a reaction in solution ● Precipitation reactions occur when pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic solid Solubility Guidelines for Ionic Compounds ● Solubility- the amount of the substance that can be dissolved in a given quantity of solvent at a given temperature ● Any substance with a solubility of less than .01 mol/L will be considered insoluble ○ Substance remains largely undissolved ● No rules based on simple physical properties to guide us in predicting solubility ● ALL COMMON IONIC COMPOUNDS OF THE ALKALI METAL IONS (GROUP 1A) AND OF THE AMMONIUM IONS (NH4+) ARE SOLUBLE IN WATER Solubility Table Soluble Ionic Compounds ●

Compounds containing

NO3-

None

CH3COO-

None

Cl-

Compounds of Ag+, Hg22+, and Pb2+

Br-

Compounds of Ag+, Hg22+, and Pb2+

I-

Compounds of Ag+, Hg22+, and Pb2+

SO42-

Compounds of Sr2+, Ba2+, Hg22+, and Pb2+

F- salts

Mg2+, Ca+2, Sr+2, Ba+2, Pb+2

Insoluble Ionic Compounds Compounds containing

S2-

Compounds of NH4+, alkali metal cations, Ca2+, Sr2+, and Ba2+

CO32-

Compounds of NH4+ and the alkali metal cations

PO43-

Compounds of NH4+ and the alkali metal cations

OH-

Compounds of NH4+, alkali metal cations, Ca2+, Sr2+, and Ba2+

How to Predict Whether a Precipitate Forms When Strong Electrolytes Mix 1. Note the ions present in the reactants 2. Consider the possible cation-anion combinations 3. Use this table^^^ to determine if any of the combinations is insoluble Exchange Metathesis Reactions ● THE EQUATION CAN ONLY BE BALANCED AFTER THE CHEMICAL FORMULAS OF THE PRODUCTS HAVE BEEN DETERMINED ● Reactions in which cations and anions appear to exchange partners conform to the general equation ○ AX+BY→ AY+BX ● Exchange reactions/metathesis reactions- such equations^^ ○ Precipitation reactions conform to this pattern, as do many neutralization reactions between acids and bases Ionic Equations and Spectator Ions ● Molecular equation- a chemical equation in which the formula for each substance is written without regard for whether it is an electrolyte or a nonelectrolyte ○ Shows chemical formula without indicating ionic character ● Complete ionic equation- an equation written with all soluble strong electrolytes shown as ions ● Spectator ions- ions that appear in identical forms on both sides of a complete ionic equation ○ No direct role in the reaction ○ Can be cancelled on either side of the reaction arrow since they aren’t reacting with anything ○ Once we cancel the spectator ions we are left with: ● Net ionic equation- an equation that includes only the ions and molecules directly involved with the reaction ● Sum of ionic charges must be the same on both sides of a balanced net ionic equation

●

IF EVERY ION IN A COMPLETE IONIC EQUATION IS A SPECTATOR, NO REACTION OCCURS ● More than one set of reactants can lead to the same net reaction 4.3 Acids, Bases, and Neutralization Reactions ● Acids and bases can be ○ Industrial and household substances ○ Components of biological fluids ○ Common electrolytes Acids ● Acids- substances that ionize in aqueous solution to form hydrogen ions H+ (aq) ○ A hydrogen atom consists of a proton and an electron→ H+ is simply a pton ○ Acids are often called proton donors ● Protons in aqueous solution are solvated by water molecules→ we write H+ (aq) ● Molecules of different acids ionize to form different numbers of H+ ions ● Monoprotic acids→ yielding one H+ per molecule of acid ○ HCl and HNO3 ● Diprotic acids→ yielding two H+ per molecule of acid ○ H2SO4 Bases ● Bases- substances at accept (react with) H+ ions ○ Produce hydroxide ions (OH-) when they dissolve in water ● Ionic hydroxide compounds are the most common bases ○ NaOH, KOH, Ca(OH)2 ○ When they dissolve in water, they dissociate into ions, introducing OH- to the solution ● Compounds that do not contain OH- can also be bases ○ NH3; ammonia ■ When added to water, it accepts an H+ ion from a water molecule and thereby produces an OH- ion ○ Ammonia is a weak electrolyte because about only 1% of the NH3 forms NH4+, and OH- ions Strong and Weak Acids and Bases ● Strong acids and strong bases- acids and bases that are strong electrolytes (completely ionized in solution) ● Weak acids and weak bases- acids and bases that are weak electrolytes (partly ionized) ● When reactivity depends only on H+ (aq) concentration, strong acids are more reactive than weak acids ● The reactivity of an acid can depend on the anion as well as on H+ (aq) concentration ○ HF (hydrofluoric acid) → weak acid (only partly ionized in aqueous solution) but very reactive and vigorously attacks many substances ■ This is due to the combined action of H+ (aq) and F- (aq)

Strong Acids Hydrochloric acid, HCl Hydrobromic acid, HBr Hydroiodic acid, HI Chloric acid, HClO3 Perchloric acid, HClO4 Nitric acid, HNO3 Sulfuric acid (first proton), H2SO4 ● This table shows us that most acids are weak ● The only common strong bases are the common soluble metal hydroxides ● The most common weak base is NH3→ reacts with water to form OH- ions Identifying Strong and Weak Electrolytes Strong Electrolyte

Weak Electrolyte

Nonelectrolyte

Ionic

All

None

None

Molecular

Strong acids

Weak acids, weak bases

All other compounds

Neutralization Reactions and Salts ● The properties of acidic solutions are quite different from those of basic solutions ○ Acids have a sour taste; bases have a bitter taste ○ Acids change the colors of certain dyes in a different way that bases affect the same dyes ■ Litmus paper principle ● Neutralization reaction- occurs when a solution of an acid and a solution of a base are mixed ○ Products of the reaction have none of the same characteristic properties of the acidic solution or the basic solution ● Water and table salt are the products of the reaction below ○ Salt- any ionic compound whose cation comes from a base and whose anion comes from an acid ○ A NEUTRALIZATION REACTION BETWEEN AN ACID AND A METAL HYDROXIDE PRODUCES WATER AND A SALT ●

The main feature of the neutralization reaction between any strong acid and any strong

base: H+ (aq) and OH- (aq) ions combine to form H2O (l) Example: Chemical equation: HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq) (Acid)

(base)

(water)

(salt)

Complete ionic equation: H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) → H2O (l) + Na+ (aq) + Cl- (aq) Net ionic equation: H+ (aq) + OH- (aq) → H2O (l) ● Notice that the OH- ions (this time in a solid reactant) and H+ ions combine to form H2O ● Neutralization reactions between acids and metal hydroxides are metathesis reactions Neutralization Reactions with Gas Formation ● Many bases besides OH- react with H+ to form molecular compounds ○ Ex: sulfide ion and carbonate ion ○ Both of these anions react with acids to form gases that have low solubilities in water 4.4 Oxidation-Reduction Reactions ● Oxidation reduction reactions/redox reactions- reactions in which electrons are transferred from one reactant to another Oxidation and Reduction ● One of the most familiar redox reactions is corrosion of a metal ○ Sometimes limited to the surface of the metal ○ Sometimes the corrosion goes deeper ● Corrosion is the conversion of a metal into a metal compound by a reaction between the metal and some substance in its environment ● When a metal corrodes, each metal atom loses one or more electrons to form a cation ○ Which can combine with an anion to form an ionic compound ● Oxidized- when an atom, ion, or molecule becomes more positively charged ○ Loses electrons ○ Oxidation- loss of electrons by a substance ● Reduced- when an atom, ion, or molecule becomes more negatively charged ○ Gains electrons ○ Reduction- the gain of electrons by a substance ● When one reactant loses electrons, another must gain them ○ Oxidation of one substance must be accompanied by reduction of some other substance Oxidation Numbers ● Oxidation number/oxidation state- a positive or negative whole number assigned to an element in a molecule or ion on the basis of a set of formal rules; to some degree it reflects the positive or negative character of that ion ● For monatomic ions (single atom)→ oxidation number is the same as the charge ● Neutral molecules and polyatomic ions→ oxidation number of a given atom is a hypothetical charge

● Rules: 1. For an atom in its ELEMENTAL FORM→ oxidation number is always 0 a. H2 = 0 2. For any MONATOMIC ION→ oxidation number is ionic charge a. K+ = +1; S2- = -2 3. NONMETALS usually have negative oxidation numbers, although can sometimes be positive a. Oxidation number of OXYGEN is usually -2 in both ionic and molecular compounds. Exception is in compounds called peroxides which contain the O22ion→ each oxygen has an oxidation number of 1 b. Oxidation number of HYDROGEN is usually +1 when bonded to nonmetals and -1 when bonded to metals c. Oxidation number of FLUORINE is -1 in all compounds. The other halogens have an oxidation number of -1 in most binary compounds. When combined, with oxygen, they have positive oxidation states 4. THE SUM OF THE OXIDATION NUMBERS of all atoms in a neutral compound is 0. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion Oxidation of Metals by Acids and Salts ● The reaction between a metal and either an acid or a metal salt→ A + BX → AX + B Ex: Zn (s) + 2HBr ( aq) → ZnBr2 (aq) + H2 (g) ● These reactions are called displacement reactions- ions in the solution are displaced (replaced) through oxidation of an element ● REMEMBER THAT WHENEVER ONE SUBSTANCE IS OXIDIZED, ANOTHER SUBSTANCE MUST BE REDUCED ● The Activity Series ● Activity series- a list of metals arranged in order of decreasing ease of oxidation ○ Active metals- metals at the top of the activity series, such as alkali and alkaline earth metals→ MOST EASILY OXIDIZED; they react the most readily to form compounds ○ Noble metals- metals at the bottom of the activity series, such as transition elements from groups 8B and 1B→ VERY STABLE; form compounds less readily low reactivity ● The activity series can be used to predict the outcome of reactions between metals and either metal salts or acids ● ANY METAL ON THE LIST CAN BE OXIDIZED BY THE IONS OF ELEMENTS BELOW IT ● Only metals above hydrogen in the activity series are able to react with acids to form H2 4.5 Concentrations of Solutions ● Concentrations- the term used to designate the amount of solute dissolved in a given quantity of solvent or quantity to solution ○ The greater the amount of solute dissolved in a certain amount of solvent, the more concentrated the resulting solution Molarity

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Molarity (symbol M)- expresses the concentration of a solution as the number of moles of solute in a liter of solution Molarity = (moles solute-mol) / (volume of solution in liters-L) Expressing the Concentration of an Electrolyte ● When an ionic compound dissolves, the relative concentrations of the ions in the solution depend on the chemical formula of the compound ○ 1.0 M of NaCl is 1.0 M Na+ ions and 1.0 M Cl- ions ○ 1.0 M of Na2SO4 is 2.0 M Na+ ions and 1.0 M SO2- ions ● Concentration of electrolyte solutions can be specified in either terms of the compound used to make the solution (1.0 M NaCl) or in terms of the ions in the solution (1.0 M Na+ and 1.0 M Cl-) Interconverting Molarity, Moles, and Volume ● If we know any two of the three quantities in the definition of molarity, we can calculate the third ● Molarity is a conversion factor between volume of a solution and moles of solute ● If one of the solutes is a liquid, we can use its density to convert its mass to volume and vice versa Dilution ● Solutions used routinely in the laboratory are often purchased or prepared in concentrated form (called stock solution) ● Dilution- the process of adding water to obtain aqueous solutions of lower concentrations ● When solvent is added to a solution, the number of moles of solute remains unchanged ● Moles solute in concentration solution = moles of solute in dilute solution ● MV = MV 4.6 Solution Stoichiometry and Chemical Analyses ● Moles of solute = (M*V) Titrations ● Titration- involves combining a solution where the solute concentration is not known with a reagent solution of known concentration ○ Used to determine the concentration of a particular solute in a solution ● Standard solution- the reagent solution of known concentration ● Just enough standard solution is added to completely react with the solute in the solution of unknown concentration ● Equivalence point- the point at which stoichiometrically equivalent quantities are brought together ● Titrations can be conducted using neutralization, precipitation, or oxidation reduction reactions ● Indicator- a dye that changes color on passing the equivalence point

Chapter 16 Sections 1, 2, 11 16.1 Arrhenius acids and Bases ● Arrhenius defined

○ acids-substances that produce H+ ions in water ○ Bases-substances that produce OH- ions in water ● An acid is a substance that, when dissolved in water, increases the concentration of H+ ions ● A base is a substance that, when dissolved in water, increases the concentration of OHions ● HCl hydrogen chloride gas is an example of an Arrhenius acid ○ When dissolved in water, it produces hydrated H+ and Cl- ions ○ The aqueous solution of HCl is hydrochloric acid ● NaOH sodium hydroxide is an Arrhenius base ○ When dissolved in water, it dissociates into Na+ and OH- ions ● This Arrhenius concept is useful but rather limited ○ Restricted to aqueous solutions 16.2 Bronsted-Lowry Acids and Bases ● Bronsted Lowry concept is based on the fact that acid base reactions involve the transfer of H+ ions from one substance to another + The H ion in Water ● Hydronium ion: H3O+ ● Chemists use the notations H+ (aq) and H3O+ (aq) interchangeably to represent the hydrated proton responsible for the characteristic properties of aqueous solutions of acids ○ H+ is more convenient and simplistic ○ H3O+ is more realistic Proton-transfer Reactions ● Proton donor (gives proton) ● Proton acceptor (accepts proton) ● Bronsted lowry acid- an acid is a substance (molecule or ion) that DONATES a proton to another substance ● Bronsted lowry base- a base is a substance that ACCEPTS a proton ● Emphasis is on proton transfer in this theory→ concept also applies to reactions that do not occur in aqueous solutions ● THE TRANSFER OF A PROTON ALWAYS INVOLVES BOTH AN ACID(DONOR) AND A BASE(ACCEPTOR) ● A substance can function as an acid only if another substance simultaneously behaves as a base ● To be a Bronsted Lowry acid, a molecule or ion must have a hydrogen atom it can lose as an H+ ion ● To be a Bronsted Lowry base, a molecule or ion must have a nonbonding pair of electrons it can use to bind the H+ ion ● Some substances can act as an acid in one reaction and a base in another ● Amphiprotic- a substance capable of acting as either an acid or a base ○ Acts as a base when combined with something more strongly acidic than itself ○ Acts as an acid when combined with something more strongly basic than itself Conjugate Acid-Base Pairs

●

In any acid base equilibrium, both the forward reaction (to the right) and the reverse reaction (to the left) involve proton transfer ● Conjugate acid base pair- an acid and a base that differ only in the presence or absence of a proton ● Conjugate base- formed by removing a proton from the acid ○ Every acid has one ● Conjugate acid- formed by adding a proton to the base ○ Every base has one Relative strengths of Acids and Bases ● Some acids are better proton donors than others ● Some bases are better proton acceptors than others ● The more easily a substance gives up a proton, the lese easily its conjugate base accepts a proton ● The more easily a base accepts a proton, the lses easily it conjugate acid gives up a proton ● THE STRONGER AN ACID, THE WEAKER ITS CONJUGATE BASE ● THE STRONGER A BASE, THE WEAKER ITS CONJUGATE ACID ● Groupings of acids and bases into three broad categories based on their behavior in water: 1. A STRONG ACID completely transfers its protons to water, leaving essentially no undissociated molecule in solution. Its conjugate base has a negligible tendency to accept protons in aqueous solution. The conjugate base of a strong acid shows negligible basicity 2. A WEAK ACID only partially dissociates in aqueous solution and therefore exists in the solution as a mixture of the undissociated acid and its conjugate base. The conjugate base of a weak acid shows a slight ability to remove protons from water. The conjugate base of a weak acid is a weak base 3. A substance with NEGLIGIBLE ACIDITY contains hydrogen but does not demonstrate any acidic behavior in water. Its conjugate base is a strong base, reacting completely with water, to form OH- ions. The conjugate base of a substance with negligible acidity is a strong base ● The ions H3O+ and OH- are the strongest possible acid and strongest possible bse that can exist at equilibrium in aqueous solution ● The leveling effect ○ Stronger acids react with water to produce hydro...