1: Water, the Common solvent

• Water

  • POLAR molecule

    • unequal distribution of charge

    • polarity facilitates water’s ability to dissolve compounds

  • Ability to DISSOLVE many substances

    • homogenous mixture, soltuion, aqueous

  • Shape

    • collection of H₂O molecules

    • H-O-H 105 degree

    • O-H covalent in nature

    • unequal sharing between the atoms

    • oxygen has a greater attraction for electrons → helps it gain slight excess of negative charge

      • Nacl(s) → Na(s) + Cl₂(g) : CHEMICAL

      • Nacl(s) → Na^{+2_{(aq) + Cl}-1 _{(aq) : PHYSICAL DISSOLUTION}}

    • δ - delta (partical charge symbol

      

  • Hydration

    • process by which positive ends of H₂O molecules are attached to negatively charged ions and vice versa

      

  • Solubility

    • solubility of ionic substances in water varies depending on

      • the attraction among ions

      • the attraction of ions for water molecules

    • polar and ionic subtances are more soluble in water than nonpolar substances (like likes like) -

      • ionic conducts electricity

    • nonionic substances are soluble because they are polar - cannot conduct electricity

2: Electrical conductivity

• ability of a solution to conduct electric currents

  • solute: being dissolved

  • solvent: dissolving medium

  • nonpolar covalent = insoluble

  • ionic compounds = soluble

• electrolyte

  • solute that dissolves in water to produce a solution that can conduct electricity

    • strong electrolytes

      • highly efficient conductors of current in aqueous solutions = many ions

      • completely pulled apart in water

      • NaCl - salts

        • soluble salts (ionic compounds)

        • strong acids

        • strong bases

    • weak electrolytes

      • conduct small current in aqueous solutions = few ions

        • partially dissolved

      • acetic acid - weak acid

    • nonelectrolytes

      • do not conduct current in aqueous solutions= no ions

      • sugar! alcohol and molecules

• Svante Arrhenius

  • definition of acid and bases

  • indentified the basis for the conductivity properties of solutions

  • extenet to which a solution can conduct an electric current directly depends on the number of ions present

• Soluble salts

  • contrian array of cations and anions

    • disintegrate and undergo hydration when the salt dissolves

    • the attraction of ions for water molecules is greater

• Nature of acids

• when dissolved in water, acids act strong electrolytes - produce ions

  

• Acids: substance that produces H⁺¹ ions when it is dissolved in water

• polarity of water helps produce H⁺¹ ions

  

• strong acids

  • every molecule fully dissociates (completely ionized)

  • ex: HCl, HNO3 H2SO4

  • arrow to the right

  • produces max moles of ions possible

• Strong bases

  • soluble ionic compounds that contain the OH ion

  • fully dissociate when dissolved in water, cations and OH ions separate and move independently

    

• Weak electrolytes

• Weak bases: resulting solution will be weak electrolyte ( partical dissociation)

  

• Weak acids: dissociate only a slight extent

• double arrow indicates that the reaction can occur in either direction

  

  • small degree of ionization in water

  • weak acids and weak bases

  • formula of acids

  • atoms that produce H ion is acid

  • OH is base

  • Non electrolytes

    • substances that dissolve in water but do not produce any ions

  

3: Composition of solutions (molarity M)

• commonly used expression for concentration

• expressed as moles of solute per volume of solution in liter

  

• Standard solution

  • solution whose concentration M is accurately known

  • process of preparation

    • 1. place a weighed amount of the solute into a volumetric flask and add a small amount of water

    • 2. dissolve the solid swirling the flask

    • 3. add more water until the level of the solution reaches the mark etched on the flask

      • mix by invention

      • label with name

  • Dilution

    • adding water to a concentrated solution to achieve molarity desired for a particular solution

    • MOLES NOT CHANGED

      

    • pipet

      • device used for the accurate measurement and transfer of a given volume of solution

4: Types of chemical reaction

• Types of solution reactions

  • precipitation reactions

    • when two solutions are mixed, a precipitate separates from the solution

      • precipitate

        • insoluble solid that is formed in a precipitation reaction

  • acid base reactions

  • oxidation-reduction reactions (redox)

6: Equations used to represent reactions in solution

• formula equation

  • overall reaction - stoich

• complete ionic equation

  • all reactants and products (that are strong electrolytes) are represented as ions

• net ionic equation

  • includes only those solution components that undergo change

    • no spectator ions

    • spectator ions

      • ions that do not directly participate in a reaction

7: Problem solving stragety

• solving stoich probelms for reactions in a solution

  

8: Acid-base reactions

• definitions

  • Bronsted-Lowry

    • acid = proton donor

    • base = acceptor

  • Aerenous

    • OH - base

    • H - Acid

• Neutralization reaction 1:1

  • general name given to acid-base reactions

    • acid is “neutralized” when enough base reacts exactly with it in a solution

  • Strong acid + strong base → ionic compound and water

    

• Alkali + Alkaline earth = bases (STRONG)

• Problem solving strategy

  1. list the species present

     1. decide what reaction occurs

  2. write balanced net ionic equation

  3. calculate moles of reactant

     1. for reaction in a solution use the volumes of the original solutions and their molarities

  4. determine the limiting reactant, sometimes

  5. calculate the mole of the required reactant product

  6. covert to grams

     

• strong

  

• Volumetric analysis

  • titration!

    • delivery of a tritrant into analyte

      • titrant

        • solution of known concentration (in buret)

      • analyte

        • solution containing the substance being analyzed (unknown)

  • Equivalence (stoichiometric) point

    • marks the point in titration where enough titrant has been added to react exactly with the analyte

  • indicator

    • methyl red that changes the color at the equivalence point

    • Phenolphtalein

      • pink in basic

  • endpoint

    • point where the indicator actually changes color

  • NO LIMITING REACTANT

9:Oxidation Reduction

• transfer of one or more electrons

• everything else that isnt acid-base or precipitant

• Oxidation states (identifying oxidation numbers)

  • the oxidation state of an uncombined elemnt is zero

  • Na, Ca, H₂ and HOFBrINCl - NO redox

  • the sum of the oxidation states of all the atoms or ions in a neutral compound is zero

    • Na⁺¹Cl^-1

    • # of electrons lost = # of electrons gained

  • The more electronegative element = negative oxidation state

  • less electronegative one is given a positive oxidatioon state

    • F = most electronegative

    • O = second most electronegative

  • SUM of the oxidation states in an ion is = the charge on the ion

• IDENTIFYING

  • some elements almost always have the same oxidation states in their compounds

    • group 1 = +1

    • group 2 = +2

    • Oxygen = -2

    • Flourine = -1

    • Hydrogen = +1 (when covalent)

    • When hyride = -1

  • Compound order= least electronegative most

    • Na Cl

• termanoligy

  • oxidation

    • increase in oxidation state = lose electrons (+ ion)

  • reduction

    • decrease in oxidation state = electron gain (- ion)

  • reducing agent

    • electron donor (oxidized) = lost electrons

  • oxidizing agent

    • electrong acceptor ( reduced) = gain electrons

• LEO the lion says GER

  • Losing Elections Oxidation

  • Gain Electrons Reduction

10: Balancing Oxidation-Reduction

• 2 methods

  1. inspection

     1. always try

  2. is the reaction redox?

     1. assign and compare oxidation #’s

     2. determine net charge in oxidation # for each

     3. determine a ratio of oxidized to reduced atoms that would = a net increase in oxidation number equal to net decrease