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Types of Chemical Reactions and Solution Stoichiometry.

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Presentation on theme: "Types of Chemical Reactions and Solution Stoichiometry."— Presentation transcript:


2 Types of Chemical Reactions and Solution Stoichiometry

3 Classification of Matter Solutions are homogeneous mixtures

4 Solute A solute is the dissolved substance in a solution. A solvent is the dissolving medium in a solution. Solvent Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Water in salt waterWater in soda

5 Saturation of Solutions  A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated.  A solution that contains less solute than a saturated solution under existing conditions is unsaturated.  A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.

6 The ammeter measures the flow of electrons (current) through the circuit.  If the ammeter measures a current, and the bulb glows, then the solution conducts.  If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting. Electrolytes vs. Nonelectrolytes

7 An electrolyte is:  A substance whose aqueous solution conducts an electric current. A nonelectrolyte is:  A substance whose aqueous solution does not conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes… Definition of Electrolytes and Nonelectrolytes

8 1.Pure water 2.Tap water 3.Sugar solution 4.Sodium chloride solution 5.Hydrochloric acid solution 6.Lactic acid solution 7.Ethyl alcohol solution 8.Pure, solid sodium chloride Electrolytes?

9 ELECTROLYTES:NONELECTROLYTES: Tap water (weak) NaCl solution HCl solution Lactate solution (weak) Pure water Sugar solution Ethanol solution Pure, solid NaCl But why do some compounds conduct electricity in solution while others do not…? We must understand the nature of water Answers…

10 The Nature of Water Water molecules are BENT (105⁰<) H & O share electrons UNEVENLY – O pulls harder on e-s, so the O end of the molecule is slightly negative – H pulls less hard on e-s, so the H end of the molecule is slightly positive This makes H2O a POLAR molecule (oppositely charged ends), like a little magnet AKA “a dipole”

11 The Nature of Water The polarity of water gives it great ability to dissolve compounds

12 1. Ionic Compounds Ionize in Solution + ions associate with the - (oxygen) end of the water dipole. - ions associate with the + (hydrogen) end of the water dipole. Ions tend to stay in solution where they can conduct a current rather than reforming a solid. IONIC CPDS ARE ELECTROLYTES! Dissociation of sodium chloride Dissociation of sodium chloride

13 Many Ionic Compounds Dissociate (break apart) NaCl(s)  AgNO 3 (s)  MgCl 2 (s)  Na 2 SO 4 (s)  AlCl 3 (s)  Na + (aq) + Cl - (aq) Ag + (aq) + NO 3 - (aq) Mg 2+ (aq) + 2 Cl - (aq) 2 Na + (aq) + SO 4 2- (aq) Al 3+ (aq) + 3 Cl - (aq)

14 Solubility Rules – Mostly Soluble (Strong Electrolytes) IonSolubilityExceptions NO 3 - SolubleNone ClO 4 - SolubleNone Na + SolubleNone K+K+K+K+SolubleNone NH 4 + SolubleNone Cl -, I - Soluble Pb 2+, Ag +, Hg 2 2+ SO 4 2- Soluble Ca 2+, Ba 2+, Sr 2+, Pb 2+, Ag +, Hg 2+

15 Some Ionic Compounds Dissociate only Slightly

16 Solubility Rules – Mostly Insoluble (Weak Electrolytes) IonSolubilityExceptions CO 3 2- Insoluble Group 1 (IA) and NH 4 + PO 4 3- Insoluble Group 1 (IA) and NH 4 + OH - Insoluble Group 1 (IA) and Ca 2+, Ba 2+, Sr 2+ S 2- Insoluble Groups 1 (IA), 2 (IIA), and NH 4 +

17 Again, b/c water is polar…and also b/c covalent acids are also polar! For instance, hydrogen chloride molecules, which are polar, give up their hydrogens to water, 2. Covalent (molecular) acids IONIZE in solution forming chloride ions (Cl - ) and hydronium ions (H 3 O + ).

18 Examples of strong acids include:  Hydrochloric acid, HCl  Sulfuric acid, H 2 SO 4  Nitric acid, HNO 3  Hydroiodic acid, HI  Perchloric acid, HClO 4 Strong acids are completely ionized in solution (Strong Electrolytes) In general, we can assume nearly all other acids are weak

19 Many of these weaker acids are “organic” acids that contain a “carboxyl” group. The carboxyl group does not easily give up its hydrogen. Weak acids ionize only slightly (Weak Electrolytes) HC 2 H 3 O 2 (acetic acid or vinegar)is a weak acid

20 Other organic acids and their sources include: o Citric acid – citrus fruit o Malic acid – apples o Butyric acid – rancid butter o Amino acids – protein o Nucleic acids – DNA and RNA o Ascorbic acid – Vitamin C This is an enormous group of compounds; these are only a few examples. FYI: Because of the carboxyl group, organic acids are sometimes called “carboxylic acids”.

21 4.Some Other Polar Covalent Compounds are Electrolytes They ionize very slightly in water Ex: Ammonia, NH 3 –Only about 1% of the molecules dissociate!

22 Sugar (sucrose – C 12 H 22 O 11 ), and ethanol (ethyl alcohol – C 2 H 5 OH) 5. MOST covalent compounds do not ionize at all in solution. (Nonelectrolyes) NOTE: These molecular compounds DISSOLVE (the water pulls the molecules away from each other, but do not IONIZE (the molecules are not charged.) DISSOLVE

23 Molarity The concentration of a solution measured in moles of solute per liter of solution. mol = M L

24 Preparation of Molar Solutions Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of M NaCl solution?  Step #1: Ask “How Much?” (What volume to prepare?) L  Step #2: Ask “How Strong?” (What molarity?) mol 1 L  Step #3: Ask “What does it weigh?” (Molar mass is?) g 1 mol = 43.8 g

25 Serial Dilution It’s not practical to keep solutions of many different concentrations on hand, so chemists prepare more dilute solutions from a more concentrated “stock” solution. Problem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution? (Note: must convert volumes to L) M stock V stock = M dilute V dilute OR M 1 V 1 =M 2 V 2 (11.6 M)(x Liters) = (3.0 M)(0.250 Liters) x Liters = (3.0 M)(0.250 Liters) 11.6 M = L

26 Single Replacement Reactions Replacement of:  Metals by another metal  Hydrogen in water by a metal  Hydrogen in an acid by a metal  Halogens by more active halogens A + BX  AX + B BX + Y  BY + X

27 The Activity Series of the Metals Lithium Potassium Calcium Sodium* Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Metals above hydrogen can replace hydrogen in acids. *Metals from sodium upward can replace hydrogen in water

28 The Activity Series of the Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F 2 (g)  2NaF(s) + Cl 2 (g) MgCl 2 (s) + Br 2 (g)  ???No Reaction ???

29 Double Replacement Reactions The ions of two compounds exchange places in an aqueous solution to form two new compounds. AX + BY  AY + BX One of the compounds formed is usually a precipitate (an insoluble solid), an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

30 Double replacement forming a precipitate… Pb(NO 3 ) 2 (aq) + 2KI(aq)  PbI 2 (s) + 2KNO 3 (aq) Pb 2+ (aq) + 2 NO 3 - (aq) + 2 K + (aq) +2 I - (aq)  PbI 2 (s) + 2K + (aq) + 2 NO 3 - (aq) Pb 2+ (aq) + 2 I - (aq)  PbI 2 (s) Double replacement (ionic) equation Complete ionic equation shows compounds as aqueous ions Net ionic equation eliminates the “spectator ions”

31 Oxidation and Reduction (Redox) Electrons are transferred Spontaneous redox rxns can transfer energy Electrons (electricity) Heat Non-spontaneous redox rxns can be made to happen with electricity

32 Oxidation and Reduction G ain E lectrons = R eduction An old memory device for oxidation and reduction goes like this… LEO says GER L ose E lectrons = O xidation

33 Oxidation Reduction Reactions (Redox) Each sodium atom loses one electron: Each chlorine atom gains one electron:

34 LEO says GER : LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

35 Rules for Assigning Oxidation Numbers Rules 1 & 2 1.The oxidation number of any uncombined element is zero 2. The oxidation number of a monatomic ion equals its charge

36 Rules for Assigning Oxidation Numbers Rules 3 & 4 3. The oxidation number of oxygen in compounds is The oxidation number of hydrogen in compounds is +1

37 Rules for Assigning Oxidation Number Rule 5 5. The sum of the oxidation numbers in the formula of a compound is 0 2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H

38 Rules for Assigning Oxidation Numbers Rule 6 6. The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its charge X + 3(-2) = -1 N O  X = +5  X = +6 X + 4(-2) = -2 S O

39 Reducing Agents and Oxidizing Agents The substance reduced is the oxidizing agent The substance oxidized is the reducing agent Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

40 Trends in Oxidation and Reduction Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents

41 Redox Reaction Prediction #1 Important OxidizersFormed in reaction MnO 4 - (acid solution) MnO 4 - (basic solution) MnO 2 (acid solution) Cr 2 O 7 2- (acid) CrO 4 2- HNO 3, concentrated HNO 3, dilute H 2 SO 4, hot conc Metallic Ions Free Halogens HClO 4 Na 2 O 2 H 2 O 2 Mn(II) MnO 2 Mn(II) Cr(III) NO 2 NO SO 2 Metallous Ions Halide ions Cl - OH - O 2

42 Redox Reaction Prediction #2 Important ReducersFormed in reaction Halide Ions Free Metals Metalous Ions Nitrite Ions Sulfite Ions Free Halogens (dil, basic sol) Free Halogens (conc, basic sol) C 2 O 4 2- Halogens Metal Ions Metallic ions Nitrate Ions SO42- Hypohalite ions Halate ions CO 2

43 Not All Reactions are Redox Reactions Reactions in which there has been no change in oxidation number are not redox rxns. Examples:

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