Aqueous Solutions

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Aqueous Solutions

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Essential Questions

• Why does something dissolve?
• Can an unlimited quantity of a substance dissolve
  in H2O?
• What quantities (units) are used to express
  solubility concentration?
• What are acids and bases?

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Solutions

• A solution is

A homogeneous mixture of two or more components.
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• A solution consists of two component types.
• solvent - component doing the dissolving
• solute - component being dissolved
• (You may have more than one.)

• In a solution
• The solute cant be filtered out.
• The solute always stays mixed.
• Particles are always in motion.
• Volumes may not be additive.
• A solution will have different
• properties than the solvent

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Physical states of solutions

• Solutions can be made that exist in any of the
  three states (or phases).
• Solid solutions
• dental fillings, 14K gold, sterling silver
• Liquid solutions
• saline, vodka, vinegar, sugar water
• Gas solutions
• the atmosphere, anesthesia gases

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Physical states of solutions

• The solution is always in the same phase as the
  solvent.
• THIS IS NOT ALWAYS THE LARGER
• VOLUME OR MASS COMPONENT.
• If the solvent and the solution are
• the same phase, then the solvent is
• the larger quantity.

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Physical states of solutions

• Solid solutions
• the solvent is solid like the resulting
  solution.
• dental amalgams liquid mercury in solid silver
• Liquid solutions
• the solvent is liquid like the resulting
  solution.
• sugar water solid sugar in liquid water
• rubbing alcohol liquid water in liquid
  isopropanol
• soda under pressure CO2 gas in liquid water
• tinctures solid medicines in liquid ethanol
• Gas solutions
• the solvent is gas like the resulting solution.
• humid air liquid water in gas phase

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Solutions

• Solutions are always in a single phase.
• Carbonated soda, once opened and you can see the
  gas bubbles in the liquid is NOT a solution. When
  the soda is sealed and the carbon dioxide is
  dissolved in the liquid (usually due to
  pressure), it is a solution.
• Aqueous solutions
• Solutions where water is the solvent.
• Other common terms
• alloy solutions of two or more metals steel
• tincture ethanol is solvent used in
  medicine/pharmacy
• amalgam mercury is solute with another metal
  dentistry

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HOMEWORK Define following terms

• Heterogeneous
• Homogeneous
• Solution
• Aqueous solution
• Solute
• Solvent
• Tincture
• Miscible
• Immiscible
• Alloy
• Amalgam

• Electrolyte
• Nonelectrolyte
• Hydrated ion
• Dissociation
• Solubility limit
• Solubility curve
• Solution equilibrium
• Unsaturated solution
• Saturated solution
• Supersaturated solution
• Colligative property

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Properties of Aqueous Solutions

• There are two general classes of solutes.
• Electrolytic
• ionic compounds in polar solvents
• dissociate in solution to make ions
• conduct electricity
• may be strong (100 dissociation) or weak (less
  than 100)
• Nonelectrolytic
• do not conduct electricity
• solute is dispersed but does not dissociate

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Dissolving Ionic Compounds in H2O

• When an ionic solid dissolves in water, the
  solvent removes individual ions from the crystal.
  The positive ions separate from the negative
  ions.

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Dissolving Ionic Compounds in H2O

• This process is called DISSOCIATION. The
  individual ions are hydrated with many water
  molecules.

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Dissolving Ionic Compounds in H2O
http//www.northland.cc.mn.us/biology/Biology1111/
animations/dissolve.html
Click link for alternate movie/video
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Electrolytes in H2O

• The dissociated ions are charged particles!
• These charged particles act as a conductor and
  can transfer electricity.
• NaCl(s) ? Na(aq) Cl-(aq)
• Strong electrolyte dissociate completely.
• The dissociate ions are free to move around.
• Many salts(ionic compounds), some acids and some
  bases are strong electrolytes.

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Weak Electrolytes in H2O

• Weak electrolytes only dissociate slightly.
• These charged particles act as a conductor and
  can transfer electricity, but there are not as
  many of them.
• HC2H3O2(l) D H(aq) C2H3O2-(aq)
• For weak electrolytes, the dissociation is
  reversible and in equilibrium.
• Some acids and some bases are
• weak electrolytes.

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Dissolving covalent compounds

• Covalent compounds do not dissociate.
• Acids are an exception (see weak electrolytes)

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Dissolving Covalent Compounds in H2O

• Ability of a solvent to dissolve a solute depends
  on attraction of polar dipole ends of water to
  the solute.

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Dissolving Covalent Compounds in H2O

• Non polar solutes do not dissolve in polar H2O.
• No attraction!

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Typical substances

• Common POLAR substances
• Water H2O
• Ammonia NH3
• Acetone C3H6O
• Alcohols, -OH, are slightly polar, bridge the
  two categories POLAR and NONPOLAR
• Common NON POLAR Substances
• Benzene C6H6
• Toluene C7H8
• Carbon tetrachloride CCl4
• Oils, gasoline, waxes, tar, diatomic molecules

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Solubility

• Factors affecting solubility

LIKES DISSOLVE LIKES!
General rule only!
Substances, like alcohols, which are only
slightly polar, can dissolve nonpolar substances,
like oils/waxes/etc. and polar substances, like
water/sugar/etc. However, its not polar enough
to dissolve ionic compounds.
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Solubility

• A measure of how much of a solute can be
  dissolved in a solvent.
• Factors affecting solubility

POLARITY of Solute AND Solvent!
Ionic Compounds (so extremely polar, there
ionic) typically dissolve in polar solvents
like water but not in nonpolar solvents like oil.
Covalent Compounds Nonpolar dissolves in
nonpolar solvents Polar dissolves in polar
solvents
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Solubility

• A measure of how much of a solute can be
  dissolved in a solvent.
• Common unit
• - grams solute / 100 g of solvent
• Factors affecting solubility
• Polarity
• Temperature

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Solubility of some substances

• Temperature Solubility
• Substance oC g/100 ml water
• NaCl (s) 100 39.12
• PbCl2 (s) 100 3.34
• AgCl (s) 100 0.0021
• CH3CH2OH (l) 0 - 100 infinity
• CH3CH2OCH2CH3 (l) 15 8.43
• O2 (g) 60 0.0023
• CO2 (g) 40 0.097
• SO2 (g) 40 5.41

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Saturation

• When a solution contains as much solute as it can
  at a given temperature.
• Unsaturated Can still dissolve more.
• Saturated Have dissolved all you can.
• Supersaturated Temporarily have dissolved
• too much.
• Precipitate Excess solute that falls out
• of solution.

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Concentration

• Molarity
• M - special symbol which means molar ( mol/L )
• Recognizes that compounds have different
• formula weights (or molar masses). Easy for
• calculating number of moles in volume used.
• A 1 M solution of magnesium atoms contains the
• same number of molecules as 1 M Hydrochloric
  acid.

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Molarity

• Molarity
• M - special symbol which means molar ( mol/L )
• Recognizes that compounds have different
• formula weights.
• A 1 M solution of hydrochloric acid contains
• 1 mole of HCl in each liter of solution

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Molarity

• Calculate the molarity of 2.0 L solution that
  contains 10 moles of NaOH.
• Use definition of Molarity MNaOH moles
  /liter
• Plug in data from problem
• MNaOH 10 molNaOH / 2.0 L
• Calculate answer
• MNaOH 5.0 M

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Molarity

• Whats the molarity of a solution that has 18.23
  g HCl in 2.0 liters?
• First, you need the molar mass of HCl.
• Molar MassHCl 1.008 x 1 H 35.45 x 1 Cl
• 36.46 g 1 mol
• Next, find the number of moles.
• molesHCl 18.23 gHCl x ( 1mol )
• 36.46 g
• 0.50 mol
• Finally, divide by the volume.
• MHCl 0.50 mol / 2.0 L
• 0.25 M

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Solution stoichiometry

• Extension of earlier stoichiometry problems.
• First step is to determine the number of moles
  based on solution concentration and volume.
• Final step is to convert back to volume or
  concentration as required by the problem.

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Solution stoichiometry example

• Determine the volume of 0.100 M HCl that contains
  5.60 grams of HCl.
• The first step is to determine how many moles of
  NaOH we have.

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Solution stoichiometry example

• We have 5.60 grams of a HCl. This is the starting
  point. (The 0.100 M is actually a conversion
  factor!)
• Start your set up.this should look familiar
• 5.60 grams HCl x ( ____________) x (___________)
• What unit do we want to cancel?
• What fact might we need?

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Solution stoichiometry example

• 5.60 grams HCl x ( ____________) x (___________)
• 
  g HCl
• Molar Mass
• 1 mol sum masses in grams
• HCl H 1 x 1.008 1.008
• Cl 1x 35.45 35.45
• 36.458g HCl 1mol HCl
• 5.60 grams HCl x ( __1mol HCl__) x (___________)
• 36.458g HCl

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Solution stoichiometry example

• 5.60 grams HCl x ( __1mol HCl__) x (___________)
• 36.458g HCl
• Whats next? Reread problem
• Determine the volume of 0.100 M HCl
• that contains 5.60 grams of HCl.
• Cancel out moles and go to liters, right?
• 5.60 grams HCl x ( __1mol HCl__) x (___L HCl___)
• 36.458g HCl mol HCl

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Solution stoichiometry example

• Determine the volume of 0.100 M HCl
• that contains 5.60 grams of HCl.
• 5.60 grams HCl x ( __1mol HCl__) x (___L HCl___)
• 36.458g HCl mol HCl
• We cant use 1mol 22.4L because its not a gas!
• However, we CAN use the MOLARITY as a
• conversion factor because M (mol)
• L

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Solution stoichiometry example

• Determine the volume of 0.100 M HCl
• that contains 5.60 grams of HCl.
• 0.100 M HCl means 0.100 mol HCl are in every1 L.
• This is a conversion factor 0.100 mol HCl 1L
  HCl for this 0.100 M solution of HCL
• 5.60 grams HCl x ( __1mol HCl__) x (_0.100_L
  HCl__)
• 36.458g HCl 1 mol HCl

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Solution stoichiometry problems

• How many grams of CaCl2 are needed to make 0.250
  L of a 4 M CaCl2 solution?
• How many molecules are in 25 L of 1.25 M solution
  of ethanol?
• What volume of 6 M H2SO4 is needed to get 75 g of
  H2SO4?

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Acids, bases and salts

• Three types of compounds are electrolytes
• Acid - a compound that increases the
  concentration of hydrogen ions in water.
• HCl H Cl-
• Base - a compound that increases the
  concentration of hydroxide ion in water.
• NaOH Na OH-
• Salt - the ions that remain after an acid and
  base react with each other - neutralization.
• HCl(aq) NaOH(aq) NaCl(aq) H2O(l)

water
water
water
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Uses
ACIDS

• H3PO4 - soft drinks, fertilizer, detergents
• H2SO4 - fertilizer, car batteries
• HCl - gastric juice
• HC2H3O2 - vinegar

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Uses
BASES

• NaOH - lye, drain and oven cleaner
• Mg(OH)2 - laxative, antacid
• NH3 - cleaners, fertilizer

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Properties
ACIDS
BASES

• sour taste
• corrosive
• electrolytes
• turn litmus red
• react with metals to form H2 gas

• bitter taste
• corrosive
• electrolytes
• turn litmus blue
• slippery feel

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Acids

• Produce H (as H3O) ions in water
• Produce a negative ion (-) too
• Taste sour
• Corrode metals
• React with bases to form salts and water

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ACID

• Acids
• Ionize to form hydronium ions (H3O) in water

HCl H2O ? H3O Cl
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Bases

• Produce OH- ions in water
• Taste bitter, chalky
• Are electrolytes
• Feel soapy, slippery
• React with acids to form salts and water

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Bases

• Bases
• Dissociate or ionize to form hydroxide ions (OH-)
  in water

NH3 H2O ? NH4 OH-
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Arrhenius Acids and Bases

• Acids produce H in aqueous solutions
• water
• HCl H(aq) Cl- (aq)
• Bases produce OH- in aqueous solutions
• water
• NaOH Na(aq) OH- (aq)

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Arrhenius Acids and Bases

• Acids ionize H in aqueous solutions
• water
• HCl H(aq) Cl- (aq)
• Bases dissociate OH- in aqueous solutions
• water
• NaOH Na(aq) OH- (aq)

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Bronsted-Lowry Acids

• Acids are hydrogen ion (H) donors
• Bases are hydrogen ion (H) acceptors
• HCl H2O H3O
  Cl-
• donor acceptor
  -

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Some acids, bases and their salts

• Acid Sodium salt
• Name Formula Name Formula
• Acetic acid HC2H3O2 Sodium acetate NaC2H3O2
• Hydrogen chloride HCl Sodium chloride NaCl
• Nitric acid HNO3 Sodium nitrate NaNO3
• Phosphoric acid H3PO4 Sodium phosphate Na3PO4
• Sulfuric acid H2SO4 Sodium sulfate Na2SO4
• Base Chloride salt
• Name Formula Name Formula
• Sodium hydroxide NaOH Sodium chloride NaCl
• Barium oxide BaO Barium chloride BaCl2
• Sodium oxide Na2O Sodium chloride NaCl
• Ammonia NH3 Ammonium chloride NH4Cl

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Some acids, bases and their saltsDO NOT
dissociate completely.

• Their strength is expressed
• Weak Acid
• Name Formula
• Acetic acid HC2H3O2
• Weak Base
• Name Formula
• Ammonia NH3

Ka is the ionization constant for acids It
tells the strength
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Weak Electrolytes in H2O

• Weak electrolytes only dissociate or ionize
  slightly.
• These charged particles act as a conductor and
  can transfer electricity, but there are not as
  many of them.
• HC2H3O2(l) D H(aq) C2H3O2-(aq)
• For weak electrolytes, this is reversible and in
  equilibrium.
• Some acids and some bases are weak electrolytes
  their strength is represented by Ka or by Kb

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• Arrhenius Concept

• Acids H donors
• Bases OH- donors
• Bronsted-Lowry concept
• Acids H donors
• Bases H acceptors
• Lewis concept
• Acids electron pair acceptors
• Bases electron pair donors

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Indicators

• Indicator
• substance that changes color in an acid or base
• Examples
• litmus - red/blue
• phenolphthalein - colorless/pink
• goldenrod - yellow/red
• red cabbage juice - pink/green

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Indicators and pH Ranges
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Indicator examples

• Acid-base indicators are weak acids that undergo
  a color change at a known pH.

pH
phenolphthalein
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Indicator examples
bromthymol blue
methyl red
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Defining Indicators

• Most solutions of acids or bases are clear and
  colorless. Therefore they cannot be distinguished
  from ordinary water by appearance alone.
• The simplest way to distinguish them from water
  is to use an indicator. A pH meter can also be
  used.
• An indicator is a chemical that changes colour as
  the concentration of H (aq) or OH (aq) changes.

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Defining Indicators (cont)

• Two of the most common indicators are
  phenolphthalein and litmus.
• Litmus is a compound that is extracted from
  lichens, a plant-like member of the fungi
  kingdom. Litmus paper is made by dipping paper in
  litmus solution.

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Neutralization

• The reaction of an acid with a base to produce a
  salt and water.
• HBr (aq) LiOH (aq) LiBr (aq) H2O (l)
• If we prepare a standard solution of LiOH, we can
  then use it to determine the concentration of HBr
  in a sample.
• This is an example of Analytical Chemistry.

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Titration

• Method based on measurement of volume.
• You must have a solution of known concentration -
  standard solution.
• It is added to an unknown solution while the
  volume is measured.
• The process is continued until the end point is
  reached - a change that we can measure.
• Acids and bases are commonly measured using
  titrations.

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Titrations
Buret - volumetric glassware used for
titrations. It allows you to add a known
amount of your titrant to the solution you are
testing. An indicator will give you the
endpoint.
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Titrations
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Titrations

• Note the color change which indicates that the
  endpoint has been reached.

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Essential Questions Enduring Understandings Knowledge and Skills
What quantities do scientist use to express solubility? Molarity is a measure of concentration commonly used. Parts per million (ppm) also is used. Express concentrations of solutions quantitatively.
Why does something dissolve? The structure of matter determines whether solutions form. Define solubility and the factors that affect it.
Can an unlimited amount of a substance dissolve in H2O? Solubility limits vary based on the substance and temperature. Use a solubility curve to make predictions about a solution.