Aqueous Solutions – a solution which water is the solvent Solution = Solute + Solvent Solute = smaller part being dissolved Solute = larger fraction doing the dissolving Molarity is used to express conc.
Molarity • M = moles of solute/ liters of solution • A 0.38g sample of sodium nitrate is placed in a 50.0 ml volumetric flask and is filled with water to the mark. What is the molarity of the solution?
How many grams of potassium permanganate are needed to prepare 2.0 liters of a 0.2 molar solution of the salt? Molarity
Dilution Problems • M1V1= M2V2 • How many liters of 12 molar sulfuric acid is necessary to make 2 liters of a 6 molar conc. of the acid?
Specific Gravity or Density Problems • Commercially available conc. HCl is in a aq sol cont 38% HCl by mass. D = 1.19g/ml a) what is the molarity of the solution? b) how many ml of conc HCl are required to make 1.00 l of a 0.10 molar solution?
Solutes that separate into tow or more ions when dissolved in water are called electrolytes. Solutes that remain uncharged molecules are called non-electrolytes. Either may be soluble in water. Electrolytes
Electrolytes • NaCl(s) Na+ + Cl- ionic solids dissociate into ions. • Electrical Conductivity Strong electrolytes conduct electricity very well because ions totally dissociate Weak electrolytes only weakly conduct because ions only partially dissociate Nearly all ionic cmps are strong and molecular cmps are usually non-electrolytes
Acids – are solutes that can ionize to produce hydrogen ions H+(aq) Strong acids dissociate and exist almost totally as ions Ex HCl(aq) H+ + Cl- Acids, Bases, And Salts
HC2H3O2 H+ + C2H3O2- Weak acids only partially dissociate the double arrows indicate that the reaction occurs in significant amounts in both directions Acedic Acid or Vinegar is the classic example. Weak Acids
Acids Cont. • H+ is simply a proton – acids are often called proton donors • Monoprotic acids – contribute one H+ per molecule of acid ex. HCl or HNO3. • Diprotic acids contribute two H+ per moleclue ex. H2SO4 • Polyprotic acids contribute greater than one H+ ex. H2SO4 or H3PO4.
Bases • Are substances that accept H+ ions or a proton. • Ex H+(aq) + OH-(aq) H2O(l) • Ex H+(aq) + NH3-((aq) NH4(aq) • OH- & NH3- are the bases. OH- is the most common base in aqueous solution. Sometimes bases are defined as substances with increased OH- ions in solution.
Note • Acids and bases are electrolytes – if a strong acid or base they are strong electrolytes if a weak acid or base a weak electrolyte. • Strong Acids HCl (aq) Hydrochloric H2SO4 Sulfuric HNO3 (aq) Nitric HBr Hydrobromic HClO4 (aq) Prechloric HI Hyrooiodic HClO3 (aq) chloric
Weak Acids HF Hydrofluoric H2C2H3O2 Acedic • Strong bases Grp I Hydroxides LiOH, NaOH, KOH, RbOH, and CsOH Grp II Hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2 Weak Base NH3 Ammonia
Are ionic compounds formed from acids by the replacement of one or more H+ ions with some other cation. For example replacing the H+ ion in HCl(aq) with Na+ ion yields NaCl. Salts
most salts are strong electrolytes most acids are weak electrolytes except for the strong acids most strong bases hydroxides of metals are strong electrolytes except NH3 which is a weak base and electrolyte Identifying strong and weak electrolytes
Neutralization reactions are those between an acid and a base they yield salt and water • Ionic Equations a. varying amounts of detail can be included in chem eq depending on what information is relevant to the problem at hand b. molecular equations - have all the species written as associated molecules even though they may be disassociated c. complete ionic eq – strong electrolytes are written as dissociated ions; molecules, solids, and weak electrolytes are written as associated ions
d. net ionic equations – an ionic eq in which the spectator ions have been cancelled. • Rules for Converting Molecular eq to Ionic 1. Ionic Sub indicated as dissolved in solution such as NaCl(aq) are normally written as ions 2. Ionic substances that are insoluble (do not dissolve readily) either as reactants or products (precipitates) are rep as formulas of the compounds
d. Molecular substances that are strong electrolytes such as strong acids are written as ions. e. Molecular substances that are weak electrolytes are represented as their molecular formulas f. Spectator ions – ions in an equation that do not take part in the equation are cancelled from both sides of the equation
Examples • Ba(C2H3O2)2 + K2CrO4 • Ba(C2H3O2)2 + K2CrO4 BaCrO4 + 2KC2H3O2 Ba + 2C2H3O2 + 2K + CrO4 BaCrO4 + 2K +2 C2H3O2 Ba + CrO4 BaCrO4
Examples • Ca(C2H3O2)2 + H2C2O4 • Ca(C2H3O2)2 + H2C2O4 CaC2O4 + HC2H3O2 • Ca+ 2C2H3O2 + H2C2O4 CaC2O4 + HC2H3O2 • Weak acids do not split
Examples • NH4NO3 + NaOH • NH4NO3+ NaOH NH4OH + Na + NO3 • NH4+NO3 + Na + OH NH3 + H2O+ Na + NO3 • NH4 + OH NH3 + H2O
Metathesis Reactions • Reactions in which two ionic reactants exchange ion partners • AX + BY AY + BX • Precipitation reactions are those metathesis reactions in which an insoluble solid product or precipitate forms A. Precipitate –insoluble solid product B. Solubility – the amount of a substance that can dissolve in a given amount of solvent
(expressed in g/100 ml or in liters/mol) A substance is considered insoluble if its solubility is less that 0.01 mol/l. • Predicting solubility is a matter of know the solubility rules (Table 4.3 or what I should know for AP list) • Some metathesis reactions from water or another weak electrolyte or non-electrolyte in ionic and net ionic equations (appear as associated species) • Some methesis reactions form gases these are usually binary compounds of non-metals exam ples H2S, CO2, HCN.
I. Oxidization Numbers A. Oxidization – loss of electrons. LEO B. Reduction – gain of electrons – GER C. When one reactant loses e’s another reactant must gain them Redox It’s electro-magic
D. reactions in which electrons are trans are called oxidization – reduction or redox reactions E. oxidization number (ON) rules 1. atoms in the elemental form – oxidization # = 0 2. monatomic ions – the oxidization # = the charge of the ion ex Na+ = 1 S-2 = -2 3. non-metals usually have neg ON although they can sometimes be positive a> Oxygen usually = -2 except peroxides O2-2 = each O = -1 Intro to electrochemistry
b) hygrogen usually = +1 when bonded to non-metals = -1 when bonded to metals c) Fluorine = -1 in all cmps – halogens = -1 in most binary cmps but when combine with oxygen as in oxyanions they have a + number 4. The sum of oxidizations #’s a) in a neutral cmp = 0 b) in a polyatomic ion = charge of ion Work sample exercise 4.6