Chem Lab 2
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Chem Lab 2

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Aqueous Solutions.

Aqueous Solutions.

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Chem Lab 2 Chem Lab 2 Presentation Transcript

  • Chem Lab 2 Solutions
  • Lab 2.1 Adding Liquids to Water
  • Lab Procedures- Behavior of Drops
    • Water
    • Drops are high and round
    • Drops stay together when dragged
    • Wet toothpick pulls drop upward
    • Many more drops can be added to a penny before falling off.
    • Propanol
    • Drops are flat
    • Drops leave a trail when dragged
    • Less drops can be added to a penny before falling off.
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  • Attractions between Molecules
    • Water molecules are more attracted to each other than alcohol molecules
    • Dipole-dipole attractions keep the molecules together as the drop is dragged or piled upon a penny.
    • Hydrogen bonds are a stronger version of dipole-dipole attractions when hydrogen atoms are one of the atoms.
  • Water vs Alcohol
    • Water is a polar molecule with two polar O-H bonds (oxygen pulls harder on negative e- resulting a partial negative charge while hydrogen has a partial positive charge)
    • Alcohol has both a polar part and nonpolar part- therefore alcohol molecules are less attracted to each other.
  • Attractions between Molecules
    • Water
    • Alcohol
  • Lab Procedures- solubility
    • Nonpolar substances do not noticeably dissolve in water
      • Oil
      • Hexane
      • Kerosene
    • A thin stream of liquid did not bend towards a charged rod
    • Polar substances dissolve in water
      • Propanol
      • Ethanoic acid
      • Hydrochloric acid
    • A thin stream of water net towards a charged rod. Other polar liquids bent slightly towards charged rod.
  • Polar solutes dissolve in polar water. Dipole-dipole attractions form between the polar molecules. The covalent bonds between atoms remain in tact while the attractions between molecules are broken and formed. The ethylene glycol molecules move between the water molecules due to the attractions these molecules.
  • Nonpolar molecules are not attracted to polar water molecules. Like dissolves Like. Nonpolar solutes dissolve in nonpolar solvents. Since the induced dipole attractions between the molecules is relatively weak, random mixing is responsible for the dissolving process.
  • Like dissolves Like
  • The nonpolar iodine dissolved to a small extent in the water. Weak induced dipole-dipole attractions form between the molecules. Upon shaking the iodine leaves the water and becomes dissolved in the nonpolar carbon tetrachloride. Weak induced dipole-induced dipole attractions form between the molecules.
  • Predicting the polarity of a bond Use electronegativity
  • Electronegativity
    • A measure of atom’s pull on electrons within a bond
    • The most electronegative atom is fluorine with an electronegativity of 4.0
    • The least electronegative atom is francium with an electronegativity of 0.8.
  • Electronegativity
  • Using Electronegativity to determine the nature of the bond between atoms
    • Calculating the difference between the electronegativity (EN) indicates how much harder one atom is pulling compared to the other atom.
    • Difference in EN
    • 0.0------0.5------------------------1.9----------
    • Nonpolar Polar Ionic
  • Little or no difference in the electronegativity
    • If both atoms exert the same pull, they share electrons equally. A nonpolar covalent bond forms between the atoms. NO OPPOSITELY CHARGED ENDS are formed.
    • If the difference in EN is between 0 – 0.5, NONPOLAR COVALENT BONDS FORM.
  • POLAR COVALENT BONDS
    • If one atom pulls harder than the other,
      • a difference between 0.5 – 1.9.
      • an unequal sharing of electrons may occur
    • The pair of electrons will spend more time at the more electronegative end.
    • This end will become partially negative and the other end partially positive.
  • IONIC BONDS
    • One atom pulls so much harder that the atom takes the electron, rather than sharing the electron(s)- difference in EN greater than 1.9
    • The more electronegative atom gains an electron(s) and becomes a negative ion.
    • The less electronegative atom loses an electron(s) and becomes a positive ion.
  • Attractions between molecules
    • Between nonpolar molecules– weak induced dipole-induced dipole attractions act between the molecules. There are no oppositely charged ends, so temporary dipoles may be induced.
    • Between polar molecules– dipole-dipole attractions act between the molecules. Oppositely charged ends or dipoles are attracted.
  • Dipole-dipole attractions form between oppositely charged ends of different polar molecules polar molecule
  • Dipole-dipole attractions form between polar water molecules and polar alcohol molecules. The molecules have moved between each other. The individual molecules remain intact. METHANOL DISSOLVES IN WATER. Notice: The covalent bonds between the atoms remain intact the molecules have simply moved between each other.
  • Dipole-dipole attractions between ethanoic acid molecules. Also known as glacial acetic acid.
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  • Polarity and Solubility
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  • Oil in water
  • Induced dipole – Induced dipole Attractions Temporary shifts in electron density lead to temporary dipoles
  • Packing peanuts will dissolve in acetone Nonpolar Nonpolar Due to random motion and and mixing (entropy driven- increasing disorder) the molecules mix between each other. Since there is little attraction between any of the molecules, they simple mix Weak induced dipole – induced dipole attractions
  • Styrofoam in acetone Styrofoam in water Is styrofoam polar or nonpolar?
  • Dipole-induced dipole Attractions Polar molecule Nonpolar molecule
  • Nonpolar oxygen dissolves in water to a small extent
    • Fish breath oxygen molecules that are dissolved in the water
    • Weak attractions form between oxygen molecules that allow some oxygen to dissolve in water
    • With increasing temperature of the water however, less oxygen will remain dissolved in the water
  • Listed from strongest IMF to weakest IMF IMF s
    • A solute:
    • Dissolves in water or other solvent
    • Changes phase if different from the solvent
    • Present in lesser amount
    • (if the same phase as the solvent)
    • A solvent:
    • Retains its phase (if different from the solute)
    • Present in greater amount
      • (if the same phase as the solute)
    • Rule for predicting solubility:
    • Like dissolves Like
  • Naming carbon compounds Organic Chemistry Nomenclature
    • Carbon with four valence electrons forms four covalent bonds to achieve a stable octet (eight valence)
    • Organic chemistry is the branch of chemistry dealing with carbon compounds.
  • Simple alkanes- carbon chains with single bonds
    • Methane- one carbon
    • Ethane- two carbons
    • Propane- three carbons
    • Butane- four carbons
    • Pentane- five carbons
    • Hexane- six carbons
    • Heptane- seven carbons
    • Octane- eight carbons
  • Memorize these.
  • Hexane
  • Branching chains in alkanes
  • Same attractions (induced dipole-induced dipole attractions) Different states of matter- What difference explains this?
  • Alkenes- a double between two carbon atoms
    • Ethene (also known as acetylene)
    • C = C
    • Use the alkane name.
    • Add a double bond
    • Adjust the number of hydrogens
    H H H H
  •  
  • Alkynes- a triple bond between two carbon atoms
    • Ethyne (also known as acetylene)
    • H C C H
    • Use the alkane name.
    • Add a triple bond
    • Adjust the number of hydrogens
  • ethene ethyne
  • Acteylene torches are used in welding. The combustion of ethyne.
  • Again, notice the difference in boiling points
  • Alcohols
    • Alcohol functional group -OH
    • Add the alcohol functional group to an alkane or other hydrocarbon
    • Add the –ol ending to the name
    • Example
      • Methanol
      • Ethanol
  • Methanol methanol
  • methanol
  • Alcohol- part polar and part nonpolar
  • Antifreeze- soluble in water due to the polar part Note: two alcohol groups attached
  •  
  • Organic Acids
    • Carboxyl group - COOH
    • The hydrogen sometimes breaks off the carboxyl group making this an acid in solution
    • Add the carboxyl group
    • Adjust the number of hydrogens accordingly
    • Add –oic acid to alkane name
  • Ethanoic Acid Also known as acetic acid or vinegar Ethan- therefore 2 carbons -oic acid add –COOH
  • Polarity and Alkanes
    • C-H bond
      • EN of carbon 2.5
      • EN of hydrogen 2.1
      • Therefore the bond is nonpolar
    • Alkanes are nonpolar
    • Alkanes will NOT dissolve in water
  • Polarity and Alcohol
    • C-H bond is nonpolar
    • O-H
      • EN of oxygen 3.5
      • EN of hydrogen 2.1
      • Difference in EN is 1.4
      • Therefore the bond is polar
    • Alcohol molecules are polar
    • Alcohol molecules will dissolve in water
  • The longer the nonpolar hydrocarbon chain the less soluble the alcohol more soluble in water less soluble in water
  •  
  •  
  • Attractions and Size determine the state of compound
    • Differences in Attractions
    • Compare ethane and ethanol
      • Ethane is nonpolar, so attractions between the molecules are weak induced dipole-induced dipole attractions.
      • Ethanol is polar, so the attractions between the molecules are dipole-dipole attractions
      • With weaker attractions between the molecules ethane is a gas at room temperature, while ethanol is a liquid
  • Given the same attractions between molecules, size can be used as a predictor of state Differences in Size
  • Attractions and Vapor Pressure
    • Ethane will have a higher vapor pressure than ethanol at the same temperature
    • With weaker attractions between the molecules, more molecules will escape the liquid and move into the gas phase causing more collisions and a greater VP
    • With a higher VP a lower temperature is required to get the liquid to boil
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  •  
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  • Which liquid has a higher boiling point? What is the boiling point of each liquid at standard pressure? Normal pressure (1 atm or 760 mm Hg
  • Boiling and Vapor Pressure
    • Boiling occurs when the vapor pressure (collisions of the molecules that have escaped the liquid) equals the external or outside pressure
    • On the previous graph
      • Find the “normal” boiling point of ethanol and heptane
      • Find the boiling point of three liquids at 500 mm Hg
  • Which liquid has greater attractions between the molecules? If the outside pressure is 200 mmHg what is the BP temp? If the outside pressure is 400 mmHg and the temperature is 10  C, will CS 2 be a liquid or gas?
  • Notice the difference in the axis- pressure compared temperature This is external pressure to show the state given different combinations of pressure and temperature S  ℓ ℓ  g S  g
  • Critical point triple point
  •  
  • Lab 2.2 Adding Different Solids to Water
  • Conductivity
    • Electrical current is a flow of charge
    • Conductivity indicator
      • Light glows if substance conducts
      • to complete circuit
    • In order to conduct electricity:
      • Must have charged particles
      • Must be able move
  • Conductivity of Solids
    • Solids that conduct
    • Metal elements
    • Composed of atoms
      • Atoms held in fixed positions by metallic bonds
    • Metallic bond
      • Sea of free moving electrons shared by all the atoms
      • Charged particles (e-) are free to move
    • Solids that do NOT conduct
    • Nonmetal elements
    • Composed of molecules
    • Molecules are held together by covalent bonds between the atoms within the molecules
    • Molecules have no charge and cannot move, so do not conduct
  •  
  •  
  • Ionic Solids
    • While the ions have a charge
    • They are not free to move in the solid
    • The solid does NOT conduct electricity
    - + - + - - + + IONIC BONDS hold the ions in fixed positions
  • Ionic Crystal Lattice
  • The ionic compound has different properties than its elements.
  • Conductivity of Solutions
    • Solutions that conduct
    • Ionic compounds
    • Combination of
    • Metals and Nonmetals
    • Composed of ions
    • Ions are free to move in solution, so the solution conducts electricity
    • Solutions that do NOT conduct
    • Covalent Compounds
    • Combination of
    • Nonmetals and Nonmetals
    • Composed of molecules
    • Molecules are free to move, but have no charge
  • Electrolytes conduct electricity in solution.
    • Strong electrolytes:
    • Conduct current very efficiently.
    • completely ionized when they are dissolved in water
    • (soluble salts, strong acids, strong bases).
    • Weak electrolytes:
    • Conduct only a small current.
    • exhibit a small degree of ionization in water
    • (weak acids, weak bases).
    • Non electrolytes:
    • Permit no current to flow.
    • that dissolve in water but do not produce any ions.
    • (pure water, sugar solution).
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  • Acid is a substance that produces H + ion (protons) when dissolve in water. HCl, H 2 SO 4 Base is a substance that produces OH - ion (hydroxide) when dissolve in water. NaOH
  • Ionic Compounds dissociate in solution
    • Ions split up
      • Dissociate
      • Ionic bonds break between ions
    • Separated ions are attracted to water molecules
      • Ion-dipole attractions form between ions and water molecules
      • Ions are surrounded by water molecules
  • Ions dissociate in solution. Ions are attracted to the oppositely charged ends of water molecules by ion-dipole attractions dipole ion
  • Water- a dipole With two oppositely charged ends Ions attracted to oppositely-charged ends of water molecule ION-DIPOLE ATTRACTION
  •  
  •  
  • Energy Change Associated with Dissolving Ionic bond break Ion-dipole attractions form Net change +2 kJ/mol Note: the net change in energy +9 kJ/mol Energy absorbed to break ionic bonds Energy released as ion-dipole attractions form
  • Calcium chloride dissolves in water
    • The temperature of the solution increases
    • Describe the direction of energy transfer.
      • Energy must be transferred out of the system to the surroundings increasing the kinetic energy (must have been PE before the reaction)
    • Is the reaction endothermic or exothermic?
      • An exothermic reaction
    • What bonds must be broken?
      • Energy must be transferred in to break ionic bonds
    • What attractions are formed?
      • Energy is transferred out as the ion-dipole attractions break
    • Is more energy transferred in or out?
      • Since the net energy change is negative, more energy was transferred out than in for an overall net transfer out
  • Calcium chloride dissolves in water Energy in to break ionic bonds Energy out as ion-dipole attractions form Net change in enthalpy Enthalpy CaCl 2 (s) Ca 2+ (aq) + 2Cl - (aq)
  • Ionic solid with ionic bonds between the ions Water with dipole-dipole attractions between the water molecules Ionic bonds break Ions dissociate Ions move between water molecules Ion-dipole attractions form Solution conducts electricity since charged particles (the ions) are free to move
  • Visualize the solution with ions and water molecules in motion
  • Ammonium nitrate dissolves in water
    • The temperature of the solution decreases
    • Describe the direction of energy transfer.
      • Energy must be transferred into the system from the surroundings, since the temperature of the surroundings decreases that KE must have been transferred into the system
    • Is the reaction endothermic or exothermic?
      • An endothermic reaction
    • What bonds must be broken?
      • Energy must be transferred in to break ionic bonds
    • What attractions are formed?
      • Energy is transferred out as the ion-dipole attractions break
    • Is more energy transferred in or out?
      • Since the net energy change is positive, more energy was transferred in than out for an overall net transfer in
  • Ammonium nitrate dissolves in water Energy in to break ionic bonds Energy out as ion-dipole attractions form Net change in enthalpy Enthalpy NH 4 NO 3 (s) NH 4 + (aq) + NO 3 - (aq)
  • Electrical Force and ion-dipole Attractions
    • Ion-dipole attraction is the attraction between charges therefore an electrical force
    • Electrical force depends upon both
      • Charge
        • (increase charge – increase force)
      • Distance
        • (decrease distance- increase force)
  • With a greater positive charge, the Mg 2+ ion is more attracted to the water molecule than Li +1 . Remember: electrical force depends on magnitude of charge- greater charge greater electrical force With a greater size, the K +1 ion cannot get as close to the water molecule as the smaller Li +1 ion. With a lesser distance, the electrical force is greater Remember: electrical force depends on distance- greater distance lesser electrical force
  • Inorganic Nomenclature- Covalent compounds
    • Covalent – Nonmetal/Nonmetal Combinations
      • Use prefixes to indicate number of atoms
        • Mono (1)
        • Di (2)
        • Tri (3)
        • Tetra (4)
      • Example
        • Carbon dioxide (CO 2 )
        • Carbon tetrachloride (CCl 4 )
        • Disulfur trioxide (S 2 O 3 )
    • Ionic - Metal/Nonmetal Combinations
  • Inorganic Nomenclature- Ionic compounds
    • Ionic - Metal/Nonmetal Combinations
      • Use charges on the ions to find ratio of ions in neutral compound
      • Metal listed first, nonmetal listed second with ending changed to –ide
      • Example
        • Calcium (+2) chloride (-1) CaCl 2
        • Sodium (+1) sulfide (-2) Na 2 S
      • Transition metals- form ions with different charges
        • Iron (III) Fe 3+ and Iron (II) Fe 2+
        • Copper (II) Cu 2+ and Copper (I) Cu +
      • Polyatomic ions
        • Nitrate (NO 3 - )
        • Carbonate (CO 3 2- )
        • Sulfate (SO 4 2- )
        • Ammonium (NH 4 + )
        • Hydroxide (OH-)
  • Inorganic Nomenclature
    • Acids
      • Positive hydrogen ions
      • And a negative ionWith monoatomic ions
        • Hydro—ic acid
        • Ex. Hydrochloric acid H w/ Cl
        • Ex. Hydrobromic acid H w/ Br
        • Ex. Hydrosulfuric acid H w/ S
      • With polyatomic –ate ions
        • — ic acid
        • Ex. Sulfuric acid H with sulfate (H 2 SO 4 )
        • Ex. Nitric acid H with nitrate (HNO 3 )
      • With polyatomic –ite ions
        • — ous acid
        • Ex. Sulfurous acid H with sulfite(H 2 SO 3 )
        • Ex. Nitrous acid H with nitrite(HNO 2 )
  • Polyatomic Ions Polyatomic ions are charged groups of covalently bonded atoms
  • Formulas and Names for Acids EOS Ternary acids simply use the polyatomic anion name with “ate” changing to “ic” plus the word acid
  • Oxidation Numbers
    • Assigned values that allow us to track the electrons
    • Represent a relative share of electrons
    • Oxidation allow us to identify the species oxidized and reduced
  • Rules for assigning Oxidation Numbers #1 elements and monoatomic ions
    • Uncombined elements have oxidation numbers of zero
    • ex. Oxygen O 2 oxid # is 0
    • ex. Iron Fe oxid # is 0
    • Monoatomic ions have oxid # is the charge on the ion
    • ex. Iron (III) ion Fe 3+ oxid # is +3
    • ex. Chloride ion Cl - oxid # is -1
  • Rules for assigning Oxidation Numbers #2 Hydrogen and Oxygen
    • Hydrogen oxid # is usually +1
    • Unless
    • as uncombined element H 2 then 0
    • or in metal hydride then -1
    • Oxygen oxid # is usually -2
    • Unless
    • as uncombined element O 2 then 0
    • or in peroxide ion then -1
  • Rules for assigning Oxidation Numbers #3 Calculating Other Oxidation numbers
    • Sum of oxidation numbers for neutral compounds will be zero
      • Ex. SO 2 Since the oxygen must be -2 and there are two of them, the sulfur must be +4 so that the sum will be zero
      • Ex. NH 3 Since the hydrogen must be +1 and there are three of them, the nitrogen must be -3 so that the sum will be zero
    • Sum of oxidation numbers for polyatomic ions will be the charge on the ion
      • Ex. NO 3 - Since the oxygen must be -2 and there are three of them, the nitrogen must be +5 so that the sum will be -1
      • Ex. NH 4 + Since the hydrogen must be +1 and there are four of them, the nitrogen must be -3 so that the sum will be +1