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Chapter 4 Notes

by: Corey Stewart

Chapter 4 Notes CHEM 1110 - 01

Corey Stewart

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These notes are really long and cover acids and bases mostly.
Gen Chem
Class Notes
General Chemistry, acids, acids and bases, bases, chapter4
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This 12 page Class Notes was uploaded by Corey Stewart on Saturday October 8, 2016. The Class Notes belongs to CHEM 1110 - 01 at East Tennessee State University taught by Mohseni in Fall 2016. Since its upload, it has received 119 views. For similar materials see Gen Chem in Chemistry at East Tennessee State University.

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Date Created: 10/08/16
Ch. 4 Water  The universal solvent  One of the most common compounds on Earth (The Blue Planet)  60% of the human body  Dissolves ionic compounds (Dissociation Reaction)  Acid-base reactions occur in water (Neutralizations) Reactions in Solution  For a reaction to occur reactants need to touch which is why reactions happen best in liquid or gas phase.  Solutions - Homogeneous mixture (Means there is no distinction between the chemicals used to make the mixture) - At least two substances that are completely intermingled (fancy word for mixed) freely - Made of molecules or ions  Solvent - Medium that dissolves solutes (Present in largest amounts) - Can be solid, liquid, or gas - In an aqueous solution, water is the solvent (aqua means water)  Solute - Substance that is dissolved (The Kool-Aid powder) - Solution is named by solute - Can be solid (Kool-Aid powder), liquid (ethylene glycol in antifreeze), gas (CO in2Coke)  Solutions are Characterized by: - Concentration ((g solute) / (g solvent)) or ((g solute) / (g solution)) - Percent concentration = (g solute) / (100 g solution) Relative Concentration  Dilute Solution - Small solute to solvent ratio (Sour Kool-Aid)  Concentrated Solution - Large solute to solvent ratio (Sweet Kool-Aid) Concentration  Solubility = (g solute to make a saturated solution) / (100 g solvent) - Temperature dependent (Hot tea holds more sugar than cold tea)  Saturated Solution - Solution that can hold no more solute at that temperature  Unsaturated Solution - Solution that can still hold some more solute at that temperature  Supersaturated Solutions - When you heat up a solution, add a solute, then cool the solution down so that it will hold more solute at that temperature (super sweet tea) - These are unstable solutions - They will crystallize when a crystal seed is added; forms precipitates Precipitates  Solid formed when a reaction occurs in water and one product is in soluble; separates out of solution. Electrolytes in Aqueous Solution  Ionic compounds conduct electricity (Ions can hold a charge and dissociate in water)  Molecular compounds don’t conduct electricity (Molecules are neutral and can’t hold a charge) Ionic Compounds (Salts) in Water  Water molecules arrange themselves around the salt and pulls the ions from the lattice (Water attacks a salt and pulls the ions away)  Dissociation - Salts break apart into ions when entering solution  Separated Ions - Hydrated - Conduct electricity Molecular Compounds in Water  Solute surrounded by water  Molecules don’t dissociate Electrical Conductivity  Electrolyte - Solutes which separate into ions and yield electrically conducting solutions (Sodium in Gatorade)  Strong Electrolyte - Electrolyte that dissociates completely in water. This causes it to conduct electricity really well. - Strong electrolytes include ionic compounds (NaCl) and strong acids/bases (HCl, NaOH)  Non-Electrolyte - Aqueous solution that doesn’t conduct electricity. Composed of molecules. (Sugar water)  Weak Electrolyte - As their name suggests, these solutions don’t conduct electricity very well. That means the solute only very slightly dissociates. Things that do this are weak acids and bases (Acetic acid (CH 3OOH), ammonia (NH )). 3 Dissociation Reactions  Ionic compounds, when dissolved in water, dissociate, or separate, into its ions. The ions then become hydrated. This means they have been surrounded by water molecules.  A hydrated ion is indicated by an (aq) subscript (H (aq). + -  When written in an equation, the ions are written separate (NaCl →Na (s) (aq) Cl (aq) Equations of Ionic Reactions  Molecular - A molecular equation is a basic equation. Nothing is written separate. No charges are written because molecules are neutral.  Ionic - Break soluble things down into the respective ions. Everything insoluble is left as a molecule. - You can find precipitates easily if you see something insoluble on the products side.  Net Ionic - The only things you write are things that participated in the reaction everything else is a spectator. (You only write the things that formed a precipitate.) Arrhenius Acid +  A substance that, when combined with water, forms the hydronium (H O ) ion. 3  Ionization Reaction Definition - Ions form where none have been before Arrhenius Base  An Arrhenius base is the exact opposite of an Arrhenius acid. Where an Arrhenius acid releases an ionized hydrogen ion, an Arrhenius base releases a hydroxide (OH) ion. (If it contains OH, more than likely [like a 99.9 % chance] that it is basic.)  Ionic Compound Containing OH - - Metal Hydroxides  These dissociate into a metal and the hydroxide (OH) ion. These are the simplest to tell if it is a base. Strong Acids  Strong acids completely dissociate in water. (See list to the right).  These acids are great electrical conductors because they break down into strong electrolytes. Strong Bases  Strong bases also completely dissociate in water.  These are also great electrical conductors for the same reason as the strong acids.  Behave as aqueous ionic compounds. (They behave as if they hadn’t been dissociated in water.)  Common strong bases are hydroxides combined with group 1 and 2 metals. - LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) , Sr(OH) 2 Ba(OH) 2 2 Weak Acids  Any acid other than the 7 in the list.  These acids do not completely dissociate in water and, therefore, are not good conductors of electricity.  These include mostly organic acids (acids containing both Carbon and Hydrogen i.e., acetate [C 2 3 ]2 Dynamic Equilibrium  Happens when two opposing reactions are occurring at the same rate (2H O → 2H + O ; 2H 2 2 2 2 O → 2H O) 2 2  Also called chemical equilibrium. + -  Written using a double arrow (HCl ↔ H + Cl)  Equilibrium means the concentration of the substances.  Dynamic means happening at the same time. Molecular Bases  Undergo hydrolysis (splitting using water) to become ions - (NH 3 H O 2 NH + OH)4+ -  These bases are harder to tell because you have to add water to them before the hydroxide ion is released. (NH + H O → NH + OH [To make it easier to see the hydroxide you can write 3 2 4 water as HOH (NH + HO3 → NH + OH)]). 4 - Weak Bases  These bases are mostly molecular so they don’t dissociate in water. Due to this, they are also weak electrolytes. +  They also accept the hydrogen ion (H ) from water inefficiently. Equilibrium for Weak Bases  These equilibrium reactions are just like a dynamic equilibrium. + -  NH 3 (aq) H2O ↔ NH 4 (aq) OH (aq)  The particles collide and react together both ways until they are both happening at the same time. General Ionization Equations This graphic does a really good job at explaining ionizations of acids and bases Brief Summary  Strong acids and bases both completely dissociate in water making them strong electrolytes.  Weak acids and bases do not completely dissociate in water making them weak electrolytes.  Strong electrolytes - Completely ionizes - Forward reaction dominates - Mostly products in solution - Equation written with →  Weak electrolytes - Only portion ionizes - Reverse reaction dominates - Mostly reactants - Equation written using ↔  IN GENERAL - Strong acids and bases have the characteristics of strong electrolytes - Weak acids and bases have characteristics of weak electrolytes - STRONG ELECTROLYTE CHARACTERISTICS ARE THE EXACT OPPOSTIE OF WEAK ELECTROLYTES Polyprotic Acids  Monoprotic Acids + - Mono- means one so these acids only release one H ion - HNO 3 (aq) H 2 → H O 3 +(aq)+ NO 3 (aq) - The hydrogen from HNO breaks3off and combines with the water to make a hydronium ion (H3O ) and leaves behind a nitrate ion (NO ) 3-  Diprotic Acids - Di- means two so these furnish two H ions +  Polyprotic Acids - Poly- means multiple so these furnish multiple H ions + - Diprotic can be considered polyprotic. Polyprotic is just a general term for acids that furnish more than one hydrogen ion. - Ex. Diprotic and triprotic HOW MANY H IONS DOES A TRIPROTIC ACID FURNISH? Acidic Anhydrides  Nonmetal Oxides - Act as acids - React with water to form molecular acids that contain hydrogen - This is what happened in the lab when we added baking soda (NaHCO ) to vinegar (3C H O ) 2 3 2 then trapped the gas (CO ) in2a test tube with water combined with phenolphthalein and it turned it a yellowish orange color. - CO 2 H O 2 HCO 3 Ionic Oxides  Basic Anhydrides - Soluble metal oxides - Reacts with water in ionization to form hydroxide ions - Forms metal hydroxide - CaO +(s)H → Ca(OH) 2 (aq) - The metal hydroxide then dissociates in water. 2+ - - Ca(OH) 2 (aq) Ca (aq) 2OH (aq) Acid – Base Nomenclature  Acids - Binary acids  Take the molecular name (HCl – hydrogen chloride)  Drop the -gen from hydrogen and merge it with the nonmetal (HCl – hydrochloride)  Drop the -ide from the nonmetal and add -ic (HCl – hydrochloric) - Oxoacid  Acids with hydrogen, oxygen and another nonmetal; polyatomic ions including oxygen (i.e., ClO3, CO3, SO 4 etc.)  Naming is based on parent oxoanion name; based on the negative polyatomic ion name  If the polyatomic ion ends in -ate change it to – ic  If it ends in -ite change it to -ous  Then add acid  H 2O –4Sulfuric acid, H SO2– S3lfurous acid - Acid salt  To name an acid salt, you use the name of the metal and combine it with the acid.  However, to find the formula, you have to make the charge of the acid be neutralized by the metal.  NaHSO – 4odium sulfurous acid, H SO – Su2fur4us acid  To make sulfurous acid have a (-1) charge to match sodium, take off a hydrogen.  Bases - Metal hydroxides and oxides  These are ionic compounds so they are named like ionic compounds  Li2O – Lithium oxide - Molecular Bases  These are molecules so they are named like molecules  ([CH 3 = methyl group, [NH ] =2amine group (This is only necessary for the naming I’m about to do.))  CH 3H – 2ethylamine  (CH 3 2H –2dimethylamine  (CH 3 3H –2trimethylamine Predicting Precipitation Reactions  Metathesis (Double Replacement) Reactions - Cations (positive ions) and anions (negative ions) exchange partners  Pb(NO ) 3 2 (aq)2KI (aq) PbI 2 (s)2KNO 3 (aq)  AB + CD → AD + CB - Charges on each ion don’t change (Pb 2+on the reactants side is still Pb on the products side). - ONLY OCCURS IF A SOLID, GAS, WEAK ELECTROLYTE, OR NON-ELECTROLYTE FORMS. - Metathesis reactions can produce a precipitate if the reactants yield an insoluble product - Knowing the solubility rules can help you know if a precipitate is going to form.  Solubility Rules 1. All salts with an alkali metal (Group 1A) are soluble. 2. All salts containing NH , N4 , ClO3, ClO 4 and C3H O are 2ol3bl2. – – – 3. All chlorides, bromides, and iodides (salts containing Cl , Br , or I ) are soluble except when combined with Ag , Pb , and Hg 22+ (note the subscript 2). 4. All salts containing SO 4– are soluble except those of Pb , Ca , Sr , Ba , and Hg . 22+  5. All metal hydroxides (ionic compounds containing OH ) and all metal oxides (ionic compounds containing O are insoluble except those of Group 1A and those of Ca , Sr , 2+ 2+ and Ba .+  When metal oxides do dissolve, they react with water to form hydroxides. The oxide ion, 2 O , does not exist in water. For example: Na 2(s) + H O 2 2NaOH(aq) 3– 2– 2– 2– 6. All salts containing PO ,4CO , S3 3 and S are insoluble except those of Group 1A and NH .4 Predicting Products of Double Replacement Reactions 1. Identify ions making sure to distinguish between subscripts telling how many ions are present and subscripts characteristic of a polyatomic ion. 2. Swap the partners making sure to make neutral compounds by adding necessary subscripts 3. Assign states based on solubility rules 4. Balance Predicting If Ionic Reaction Occurs 1. Write balanced molecular equation for double replacement reaction. 2. Determine if the ions on the product side form insoluble salts, water, weak electrolytes, non-electrolytes, or gas. 3. Translate the molecular equation to an ionic equation. 4. Cancel the spectator ions and write the net ionic equation. 5. Look for the driving force of the reaction which is listed above: the formation of form insoluble salts, water, weak electrolytes, non-electrolytes, or gas. Predicting Acid-Base Reactions  Neutralization Reaction - Combination of an acid and a base to form a salt and water. - Can be viewed as a double replacement reaction. - HCl (aq) NaOH (aq) NaCl (aq) H 2  This basic structure of acid-base neutralization forming a salt and water is true for any neutralization other than a strong acid-weak base neutralization.  NH 3 (aq) HCl(aq) NH Cl4 (aq) Acid Salts  These formulas will contain a cation, a hydrogen, and an anion (NaHCO ) 3  The acid salt can react with a base - H 2O 3 (aq)NaOH (aq) NaHCO 3 (aq) H2O  Polyprotic acids can be neutralized stepwise  This means they may stop neutralization part of the way through.  Can be neutralized by an additional base.  As you name the acid throughout the neutralization, the name must indicate how may more hydrogen are left.  NaH PO 2 So4ium dihydrogen phosphate  Na HP2 – So4ium hydrogen phosphate  Some acid salts have common names (NaHCO – Sodium hydr3gen carbonate or sodium bicarbonate) Metathesis and Gas Formation  Metathesis reactions involving certain ions lead to formation of gas  The low solubility of gas leads to it escaping the solvent (most cases water)  Gas cannot redissolve once it escapes. This will drive the reaction to completion.  Most compounds that contain an anion that causes a gas to form are insoluble.  Adding an acid to these will cause the gas to form  Some gases formed are CO , SO , an2 NH 2 3  Unstable compounds decompose and form gas - H 2O →3H O + 2O 2 (g) - H 2O →3H O +2SO 2 (g) - NH O4 → H O + 2H 3 (g) Reactions That Release CO 2 -  Acid with Bicarbonate (HCO ) 3 - NaHCO 3 (aq) HCl → NaCl (aq)+ H 2 + CO 2 (g) 2-  Acid with Carbonate (CO ) 3 - Na 2O + 3HCl → 2NaCl (aq)+ H 2 + CO 2 (g)  Acids with Sulfites (SO ) or3Bisulfites (HSO ) 32- - K 2O 3(aq) 2HClO 4(aq) SO 2(g) 2KClO 4(aq)+ H 2 - LiHSO 3(aq) HClO 3(aq) SO 2(g)+ H 2 + LiClO 3(aq)  Acids with Sulfides - 2HCl (aq)+ Na 2 (aq)  2NaCl (aq) H 2 (g)  Acid with Cyanides - HNO + CsCN  HCN + CsNO 3(aq) (aq) (g) 3(aq)  Bases with Ammonium salts-- - NaOH (aq)+ NH 4l (aq) NH 3(g) H 2 + NaCl (aq) Metathesis Overview  Precipitation - A solid forms from two solutions  Neutralization - Acid + Base → Water + Salt  Gas forms - Metathesis reaction in which one of these products form: HCN, H S, H CO 2 2 3(aq) H2SO 3(aq) NH 3(aq)  Formation of a Weak Electrolyte - Weak acid reacts with strong base to form molecule Predicting Reaction and Writing Their Equations  This is just an example. Some equations you predict will have those listed at the end of the example. Synthesize Salts via Metathesis Reactions  This is a practical use for metathesis reaction  The desired compound should easily separated from the solution. This follows three principal approaches: - The desired compound is insoluble in water.  The desired compound can be separated by filtration. - The desired compound is soluble in water  Acid-base reactions  Desired compound can be separated by evaporation - The desired compound is soluble in water  Acid is added to supply the needed anion  A metal carbonate, metal sulfide, or metal sulfite are added to produce a gas, a salt, and water (the latter two are generally avoided are the gases produced are poisonous)  The desired compound is then isolated by evaporation Molarity (M)  (Moles of solute)/(Liters of solution)  Expression of relationship between moles of solute and liters of solution  .1 M solution of NaCl contains .1 moles of NaCl in 1 liter of solution Molarity as Conversion Factor  M = (Moles of solute)/(Liters of solution)  Moles of solute = Molarity * Liters of solution  Liters of solution = (Moles of solute)/(Molarity) Preparing Solution of Known Molarity  Weigh out the amount of solid needed.  Add enough water needed to dissolve the solid.  Then add water to the etched line in a volumetric flask. Diluting Solutions  Moles do not change when a solution is diluted only the volume does. Solution Stoichiometry  Often these problems will work with solutions when conducting reactions (i.e, determining the amount of reactant needed to react with another reactant completely)  Worked like any other stoichiometry problem Stoichiometry of Ionic Equations  Used to determine amount of ions in a solution  Molarity is used  Molar concentration of ions = Molarity of compound * number of ions of that element in that compound  Molar concentration of a .1 M solution of NaCl - Na = .1 M * 1 ion of Na in NaCl = .1 - Cl = .1  Total Molar concentration is the sum of the molar concentration of all ions in the compound - Total molar concentration in .1 M solution of NaCl = .2 - .2 = .1 molar concentration of Na + .1 molar concentration Cl - Chemical Analysis  Two types - Qualitative analysis  What substances are present in a sample - Quantitative analysis  Measure the amounts of various substances in a sample  Convert all of an element present in a sample into a substance of known formula  Use the amount of this known to determine amount of element present in the original sample (unknown or analyte) Titrations  Controlled addition of one reactant to another until reaction is complete.  Analytical technique to determine the amount of solute in a solution.  Acid-base titrations  It’s used daily to determine the purity of water (Don’t know if this is gonna be on the test, but is an interesting tidbit if you find this fun).  How it works: - Need to know that quick, complete reaction takes place - Need to know exact quantity of one reactant - Use stoichiometry to find exact amount of any other substance in solution Titration: Definitions  Buret – Volumetric measuring device with 0.10 mL markings holding titrant  Stopcock – Permits flow of titrant to stop when reaction is complete  Titrant – Solution in the buret; known concentration; can be either an acid or base depending on the analyte  Analyte – Solution being analyzed in the flask; unknown concentration  Equivalence point- Volume of titrant where moles of titrant and moles of analyte are stoichiometriclly equal  Indicator – Dye that is one color in an acid and another in a base  Endpoint – Volume of titrant at which acid or base is completely neutralized; color of indicator changes  Millimoles – moles divided by 1000; easier to do stoichiometric equations in mL Summary of Stoichiometry Calculations


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