Download Outline for Unit 1 Solutions, Acid/Base, and Gases

Unit 3: Solutions/ Acids and Bases
Matter as Solutions, Acids, and Bases
Examples:
Pure:
Heterogeneous:
Homogeneous:
Solutions: made by the homogenous mixture of 2 substances (can only see 1 substance)
a) Solute: substance that dissolves (normally a solid, but can be a liquid or gas), or the
substance in lesser quantity.
b) Solvent: substance that does dissolving (normally a liquid, but can be a gas), the
substance in greater quantity.
When substances are in the same state, the one with greater volume is considered to be the solvent.
Solutions:
1. Allow reactions to happen faster due to the Brownian motion of fluids
-more movement of particles = more likely particles will react in a favorable way
2. Allow reactions to happen that would otherwise not happen – surface area increases,
possible entry shape is changed.
Examples of solutions:
Types of solutions:
DATE:
NAME:
CHAPTER 5
ASSESSMENT
CLASS:
Classification of Matter
Review
BLM 5.0.1
1. Use the terms provided on p. 160 of the text to complete this flow chart on the classification of
matter.
H2O(ℓ), CH4(g),
NaNO3(s)
AgCl(s), Na3PO4(s),
KOH(s)
Has a definite
composition
H2(g), Ne(g),
Cu(s)
Anything
with mass
and volume
CO2(g), CH4(g),
O2(g), P2O5(s)
Air, brass,
stainless steel
Separated by
physical means
Orange juice,
chocolate chip
cookies
2. Classify the following substances using the terms provided on p. 160 of the text.
(a) tap water
(d) gasoline
(b) gold
(e) ocean water
(c) sweetened tea
(f) seasoning salt
3. Many ionic compounds are soluble in water. Using the solubility table on p. 161 of the text,
determine which of the following ionic compounds would dissolve to form an aqueous solution.
(a) NaF (s)
(b) AgCH3COO(s)
(c) (NH4)2S(s)
(d) Ba(OH)2(s)
(e) Ca(OH)2(s)
(f) PbSO4(s)
(g) CuCl2(s)
(h) CuCl(s)
4. List two properties that are different in molecular and ionic compounds.
Dissolving (forming a solution): A physical change (not chemical change-rxn)
‘like’ dissolves ‘like’
ie. Polar molecules will dissolve in a polar substance. Non-polar molecules will not
dissolve in a polar substance. (vice-versa)
Water is the most common solvent used to form a solution. A solute is the substance
dissolved into the solvent to form a solution. A solution is a mixture of 2 or more
things that appears the same throughout (homogeneous mixture). Solutes and
solvents may be gases, liquids, or solids, but we are going to focus on solutions that
use water as the solvent ie. Aqueous solutions. They are clear and transparent.
Ionic compounds:
positive end of water attracts the
negative ions and draws them out of the crystal and
surrounds them. The same process occurs for the negative
end of water and the positive ions.
Ionic crystals dissociate in water
molecular compounds:
the intermolecular forces break
and the water is attracted to the polar parts of the molecule
and forms H-bonds if possible.
Molecular compounds break into individual
molecules in water.(if they dissolve)
Any liquid substance that is cloudy (opaque or translucent) has undissolved particles
in it and is NOT aqueous. They are heterogeneous!
See sheet 5.1.2-practice with solubility table
Conductivity
Electrolyte: An aqueous solution that conducts electricity.
Non-electrolyte: An aqueous solution that does NOT conduct electricity.
If a substance ionizes (separates into ions when dissolved) it will be an electrolyte.
This happens for any aqueous solutions of ionic compounds and for strong acids
and bases. Some substances only partially ionize and therefore are WEAK
electrolytes.
Conductivity can be tested with a conductivity apparatus or an ohmmeter.
Place the following terms in the correct column of the chart below:
aqueous ionic compounds, aqueous molecular compounds, low solubility ionic
compounds, insoluble molecular compounds, strong acids, weak acids, C8H18(l)
C6H12O6(aq), NaOH(aq), H2SO4(aq), CH3COOH(aq), CaPO4(s), C2H5OH(aq), NaCl(aq),
Electrolytes
Non-electrolytes (or weak)
Text questions and 5.1 review pg 175
Lab 5A pg 174
Section 5.2 – Solubility
Solubility: the amount of a solute that dissolves in a given quantity of solvent, at a given
temperature. (g/mL)
Things that affect the rate or degree of Solubility:
1. Nature of solute and solvent – like dissolves like
2. Agitation – solute dissolves more rapidly since fresh solvent is brought into
contact with the solute. Only affects the rate not amount of solute that
dissolves.
3. Temperature – at higher temp kinetic energy of the solvent is higher so more
collisions of solvent molecules with solute
4. Particle size – smaller particles dissolve faster since there is more surface area
available to solvent
5. Pressure (partial pressure) – only affects gas in liquids – solubility of gas
increases as partial pressure of the gas above the solution is increased.
Miscible: liquids will dissolve in each other in all proportions (can never reach saturation
point) ex. alcohol and water
Immiscible: liquids that are insoluble in one another. Ex. oil and water
Remember: Ionic compounds break apart into positive and negative ions when they are
dissolved in solution.
KBr(s) in water  K+ (aq) and Br- (aq)
See your new solubility chart.
Slightly soluble ionic compounds are solid crystals in solution e.g. LiF(s)
General rules:
 all ionic compounds are at least partially soluble
 All compounds containing group 1 ions or ammonium ions are soluble.
(except LiF)
 All compounds containing nitrates are soluble.
 Solubility for solutions is dependent on temperature
Temperature:
The solubility of solids increase with temperature
The solubility of liquids does not change appreciably with temperature
The solubility of gases decreases with temperature
Can you explain why??
What are the problems with the temperature increasing (thermal pollution) in a river or
lake?
Pressure:
The solubility of solids does not change appreciably with pressure
The solubility of liquids does not change appreciably with pressure
The solubility of gases increases with pressure.
Explain why SCUBA divers need to rise slowly after diving.
Unsaturated solution: a solution that is not saturated and, therefore, can dissolve more
solute at that particular temperature.
Saturated solution: a solution that contains the max.amount of dissolved solute at a given
temperature in the presence of undissolved solute. It is said to be in a state of equilibrium.
In equilibrium, the forward and the reverse reaction are occurring at the same rate.
In a saturated solution, the dissolved solute will crystallize at the same rate as the
undissolved solute dissolves
Supersaturated solution: contains more dissolved solute than its solubility at a given
temperature.
How can you make a supersaturated solution???
http://gizmodo.com/magically-turn-liquid-soda-into-a-delicious-slushie-wit-525655870
Investigation 5B (blm 5.2.2)pg 180 or solution saturation lab
5.2 review pg 183
5.3 Concentration of solutions
Concentration: conc. of a solution is a measure of the amount of solute that is dissolved
in a given quantity of solvent.


Dilute solution: contains small amount of solute
Concentrated solution: contains a large amount of solute
Ways to express concentration:
1.
Percent by mass =
Common use: alloys
Example:
Page 186 practice problems
 100
2.
Concentration in ppm =
 106ppm
Common use: small solute amounts (toxins; CO detectors, HS gas)
3.
Concentration in ppb =
 109 ppb
Eg. Symptoms of mercury poisoning become apparent after a person has accumulated
20mg of mercury in his or her body.
a) Express this concentration as parts per million for a 60 kg person.
b) Express this concentration as parts per billion for the same person.
c) Express this amount as a percent by mass for the same person.
Page 188 practice problems
Sheet 5.3.2
4. Molarity (Molar concentration): the number of moles of solute dissolved in one
liter of solution. Square brackets
(mol/ L)
(M) = n/V
= number of moles of solute (mol)/ number of liters of solution (L)
Page 191 practice problems
Example:
Sheet 5.3.3
Molarity is the most common method used for expressing concentration of aqueous
solutions. We will use this method for the rest of the unit.
If asked for ion molarity, watch your ratio of compound/ion.
Concentration of solution vs. Concentration of ions (dissociation/ionization)
With strong electrolytes, we said that the solute ionizes in order to dissolve in water.
This means that the compound has separated into ions.
Different compounds have different ratios of ions within them. A substance such as
NaOH has one Na ion and one hydroxide ion. However, Ba(OH)2 would separate to form
one barium ion and 2 hydroxide ions. In other words the same conc. of the solutes would
result in twice the amount of hydroxide for the barium compound.
(write dissociation equations and show mole ratio)
eg. Page 193 practice problems
Sheet 5.3.5
5.3 review page 196
5.4 Solution
Preparation:
Volumetric Techniques must be used when preparing STANDARD solutions
(SOLUTIONS OF KNOWN CONCENTRATIONS).
This involves the use of precision equipment like electronic balances and containers
called VOLUMETRIC glassware. They come in a variety of shapes and sizes.
Volumetric flask;
volumetric pipette:
A: to prepare from a pure solid:
use the formula C=n/V to find the moles required,
and then the formula n=m/M to find the mass of solid needed to prepare the
solution.
Eg. Calculate the moles of SALT required to make 900mL of 0.225 mol/L sodium
phosphate solution. (can be used to remove grease from ovens)
What mass is required to prepare the above solution?
What steps would need to be taken?
Step 1 – Using an electronic balance, measure the mass of solute onto weigh paper.
 Place the paper on the balance. Press
“tare” or “zero” to re-zero the
balance (so that you will be massing
only the solute).
 Add the solute using a scoopula until
you reach the desired mass. Transfer
the mass to a beaker.
 The weigh paper must be rinsed into
the beaker to reduce the amount of
solute lost during the transfer.
Step 2 – Add distilled water to the solute in the beaker. Stir with a stirring rod to
dissolve.
 Add approximately ½ of the total
volume of water desired for the final
solution (ie: if you are making 100
mL of solution, add about 50 mL –
this volume does not have to be
precise)
 At this point, all of the solute must
be dissolved. Add more water if
necessary or heat the solution to
dissolve all of the solute (do not add
more than the total volume of the
solution and leave some room for the
water from rinsing.)
Step 3 – Transfer your solution to a volumetric flask of appropriate size.
 A volumetric flask is a pear-shaped
glass container with a flat bottom
and a long neck. They are used to
make up standard solutions because
they are carefully calibrated to the
precise volume listed.
 Ensure that the volumetric flask is
clean before using. You may rinse
the flask with distilled water. The
flask does not need to be dry.
 Use a funnel to transfer the solution
from the beaker to the volumetric
flask.
Step 4 – Rinse the beaker and stirring rod into the funnel with distilled water. Rinse the
funnel into the volumetric flask.
 Rinse all glassware that came into
contact with the solution into the
funnel. This will reduce the amount
of solute lost during the transfer.
 Thoroughly rinse the funnel into the
flask.
 Be careful not to add too much water
– the level MUST remain below the
etched line on the neck of the
volumetric flask.
Step 5 – Add distilled water until the bottom of the meniscus reaches the etched line on
the volumetric flask. Stopper the flask, keep your thumb on the stopper and invert the
flask several times to mix the solution.
 Add the rest of the water slowly. The
narrow neck of the volumetric flask
will fill up quickly.
 When the flask is almost full, add the
water drop by drop until the bottom
of the meniscus rests at the etched
line. The flask must be on a level
surface and the line must be at eye
level.
 Depending on the size of the
volumetric flask used, the volume
will be precise to either one or two
digits after the decimal point. Check
with your teacher as to the precision
of your laboratory glassware.
Page 194,195
Sheet 5.3.6
Mixed conc review 5.3.7
B: To prepare a dilute solution from a concentrated solution (Dilutions):
When a more dilute solution is being made from a STANDARD solution then how do
you know how much of each should be used?
The amount of solute (moles) in the solution drawn into a pipette, is the same amount of
solute (moles) in the final diluted solution.
Therefore:
n=C1V1
before dilution
moles is the same so:
n=C2V2
after dilution
C1V1(before)= C2V2(after)
eg. A 18.0 mol/L concentrated sulfuric acid is used to make 1.00L of 0.100 mol/L
solution. How much of the concentrated acid is needed?
**In dilutions, the number of moles of solute stay the same because you are using
ALL of the original solute in solution and just adding more solvent (water).
Step 1 – Rinse the pipette with the standard solution using the technique in Step 2.
Discard the solution.
 A pipette is a thin glass tube that is
used to measure a precise volume of
solution. There are two types of
pipettes: graduated and volumetric.
 Pipette A is a graduated pipette. It is
calibrated to measure a number of
volumes, like a graduated cylinder.
 Pipette B is a volumetric pipette. It is
calibrated to measure one volume
only, like a volumetric flask.

Step 2 – Draw the desired volume of solution into the pipette.
 Squeeze some air from the pipette
bulb before lightly placing it over
the pipette to draw up the solution
(draw up more solution than desired
without overfilling the pipette and
flooding the bulb).
 Remove the bulb and quickly place
your index finger over the top of the
pipette.
 By slightly rolling your index finger
to the side, release the solution from
the pipette until the bottom of the
meniscus is on the line etched into
the glass. (This must be done while
looking at the line at eye level)
 You may need to practise to get the
hang of this technique.
Step 3 – Release the standard solution into a clean volumetric flask.
 DO NOT force the last drop from
the pipette using the bulb. The
pipette should be vertical and should
touch the side of the flask as shown
below.
 After the solution has visibly stopped
flowing, keep the pipette touching
the side of the flask for three seconds
to ensure the full volume has had
time to exit the pipette.
 A small volume of standard solution
will remain in the pipette. Pipettes
are calibrated in the factory based on
how much solution they deliver, not
on how much solution they contain.
Step 4 – Add distilled water until the bottom of the meniscus reaches the etched line on
the volumetric flask. Stopper the flask, keep your thumb on the stopper, and invert
several times to mix the solution.
Page 198 practice problems
Sheet 5.4.1
Lab page 200
** When taking a small sample of solution from a larger volume of solution, the
concentration stays the same. You have NOT used ALL of the solute (total moles
has changed), but you do have the same amount of solute per volume (conc is the
same)
Eg. A 20mL sample is taken from a 0.25mol/L solution of volume 500mL. What is
the concentration of the sample?
What is the amount of solute (moles) present in the sample?
5.3 review page 202
chapter 5 review/test
Solution Stoichiometry:
Remember that in our gas unit we learned the four steps to stoich:
1. Write a balanced chemical equation/reaction.
2. Find the moles of a substance in the reaction.
3. Convert the moles of the substance from step 2 to moles of the unknown
substance.
4. Find the unknown value.
In steps 2 and 4 we are using equations that we have learned to find the moles, or the
unknown value. These equations, so far, were;
Gravimetric(solids);
n=m/M for solids
ideal gas law(gases); PV=nRT for gases
NOW with solutions we also have the equation C=n/V that can be used to find moles or
the unknown value for a solution (liquids).
Eg. What is the minimum volume of 0.250mol/L MgCl2(aq) needed to precipitate all the
silver ions in 60mL of 0.30 mol/L AgNO3(aq)?
Pg 282 practice
lab
ACIDS AND BASES:
Name them IONIC first!
Classical:
Eg.
IUPAC naming for acids:
Pg 209 naming acids chart and questions
DATE:
NAME:
CHAPTER 6
CLASS:
Naming Acids and Bases
Review
ASSESSMENT
1. Fill in the following chart:
Name of pure
substance
IUPAC name for
acid
Classical name
for acid
Chemical
formula
hydrogen cyanide
HBr(aq)
aqueous hydrogen
sulfate
nitrous acid
HSCN(aq)
hydrogen
phosphate
aqueous hydrogen
chlorite
hydroiodic acid
H3BO3(aq)
2. Using the trends you discovered in Question 1, fill in the following naming grid.
Ending of name of pure substance
Classical name of acid
-ate
hydro-----ic acid
-ous acid
3. Name or give the formula of the following bases:
(a) NaOH
(c) strontium hydroxide
(b) Mg(OH)2
(d) potassium hydroxide
(c) NH3
(e) cesium hydroxide
BLM 6.1.2
Empirical Properties of acids and bases (using senses, not theory)
ACIDS
Sour taste
Conduct electricity
React with base to neutralize
Reacts with active metals to produce gas
Ph 1-6.9
Sheet 6.1.4
BASES
Bitter taste
Slippery feel
Conduct electricity
Reacts with acid to neutralize
Doesn’t react with active metals
Ph 7.1-14
Arrhenius Theory of acids: is a substance that ionizes to form hydrogen ions, H+ (aq),
when placed in water
Ex. HCl  H+ (aq) + Cl – (aq)
Arrhenius Theory of bases: is a substance that ionizes to form hydroxide ions, OH(aq),when placed in water
Ex. NaOH  Na + (aq) + OH – (aq)
Ionization/Dissociation: when acidic molecular substances or soluble ionic substances
dissolve(dissociate) in water, they ionize or form ions.
Page 213 questions
Limitations of Arrhenius Definition of acids and bases:
His definition did not include substances such as water or hydrogen carbonate ions, who
can play both an acid or a base depending on what it is mixed with. Also doesn’t explain
the behavior of substances that do not have OH in their formula like ammonia.
Modified Theory of Acids and Bases: demo lab 6b page 214
Acid: reacts with water to produce H3O+ (aq) in aqueous solution.
Base: dissociates or reacts with water to produce OH- (aq) in aqueous solution.
Examples:
Limitations; some substances can behave as both! How do you tell whether it is really
an acid or base?
Sheet on oxides
6.1 review
Strong vs. Weak Acids: according to the theory of acids and bases , all acids produce
hydronium ions in water. However not all acids ionize to the same degree.
Strong acids ionize 100% in water, whereas weak acids ionize less than 50% in water.
Dissociation equation for strong acid (single arrow):
Dissociation equation for weak acid (equilibrium arrow):
This means different amounts of hydronium ion are present for strong vs weak acids.
Strong acids:
 Have high concentrations of hydronium ions, making their pH closer to 1
 Have a high conductivity
 React quicker
 React with active metals vigorously producing hydrogen gas.
Weak Acids:
 Have a low concentration of hydronium ions, making their pH closer to 7
 Have a low conductivity
 React more slowly
 Have a little reaction with active metals.
Compare amount of ionization/strength.
Strong vs. Weak Bases:
Strong base: completely (100%) dissociates into hydroxide (OH-) ions when dissolved
in water. All oxides and hydroxides of the alkali metals and alkaline earths – group 1 and
2 are strong bases. (some are not very soluble in water but what does dissolve will
dissociate completely.
Examples: (arrow!)
What would you expect with their:
 Conductivity?
 pH?
 Reactivity?
Weak bases: < 50% of the substance dissociates into ions (resulting in hydroxide ions)
when combined with water.
Examples: (arrow!)
Pg 222 #5-7
Following sheet
CHAPTER 6
ASSESSMENT
Properties of Strong and Weak
Acids and Bases
BLM 6.2.4
1. Fill in the chart of expected results for acids that all have a concentration of 0.1
mol/L.
Acid
Chemical
pH
Conductivity
Reactivity with
formula
(slightly less (high or low)
magnesium
than 7 or
metal (high or
much less
low)
than 7)
hydrochloric
acid
ethanoic
acid
boric acid
hydrofluoric
acid
sulfuric acid
HClO4(aq)
H3PO4(aq)
HBr(aq)
H2SO3(aq)
2. Two different acidic solutions have a concentration of 0.1 mol/L. Solution A
conducts electricity extremely well, while solution B conducts very poorly.
Which of the solutions will have a lower pH? Explain using a description of what
is happening on a molecular level.
3. You have two basic solutions. One has a concentration of 1.0 mol/L and one has
a concentration of 0.1 mol/L. You know one is a strong base and one is a weak
base. Can you determine which solution is which based on pH? If so, explain
how. If not, explain why not.
Monoprotic acid – will only give one H + ion (PROTON) when in water.
Ex. HCl, CH3COOH, HBr, HNO3
Polyprotic acid contains two or more hydrogen ions. Each hydrogen ionizes less
completely than the previous hydrogen ion.
Ex. H2SO4(aq)
H3PO4(aq)
Sulfuric acid, a strong diprotic acid, is a very important industrial acid. It is so important
that the mass of sulfuric acid used by a nation is a measure of the strength of a country’s
economy. Sulfuric acid is used in all of the products and processes in this diagram.
Polyprotic Bases: reacts with water in two or more steps to generate hydroxide ions:
Examples: a) sodium carbonate
First dissociate:
Then write base reactions:
Try writing the base reactions for all steps of potassium phosphate:
Dealing with Spills of acids and bases
If acid spills you can either?
a) Dilute it with water
b) Neutralize it
When would you use each method?
Neutralization reactions:
When an acid and base react to form water and an ionic salt (ionic compound that does
not produce hydrogen or hydroxide ions in water).
Eg:
All neutralization reactions come down to hydronium (or hydrogen ion) and hydroxide
ion reacting to form water. Neutralization reactions are used to maintain how acidic or
basic a solution should be (e.g. pool water).
Example:
Do flipchart review: acids and bases
HYDROGEN IONS FROM WATER
Hydroxide ion: a negatively charged ion formed when a water molecule loses a hydrogen
ion. (OH-)
Hydronium ion: a positively charged ion formed when a water molecule gains a hydrogen
ion or when H+ binds to a water molecule . (H3O+)
In pure water or a neutral substance: H3O+= OH- = 1x10-7 mol/L
The concentrations of H3O+and OH- must be equal in pure water.(neutral, pH= 7)
Kw= [H3O+] x [OH-] = 1.0 x 10-14 (mol/L)2 Ion product constant for water.
This is true for every aqueous solution regardless of the presence of other ions.
Acidic solutions [H3O+] greater then [OH-]
[H3O+] greater than 1 x 10-7
Basic/alkaline solutions [H3O+] less than [OH-]
[H3O+] less than 1 x 10-7 mol/L
BOTH Hydrogen(hydronium) and hydroxide ions will always be present in aqueous
solutions, even in acids or bases. If we know the conc. of one of the ions we can always
find the conc. of the other using Kw.
Example
If H3O+= 1 x 10-5 mol/L, is the solution acidic, basic, or neutral? What is the
concentration of OH-?
Sheet 6.3.2
The pH concept (power of hydrogen)
The pH of a solution is the negative logarithm of the hydrogen ion concentration.
pH = -log[H+] or
–log [H3O+]
Ex 1: [H+] = 1 x 10-7 mol/L of a neutral solution.
pH= -log( 1x 10-7)
= -(log 10-7)
pH = 7.0
Neutral!
Ex 2: H+ = 1 x 10-10 mol/l
pH: = -log (H+)
= -log (1 x 10-10)
= -(0.0 + (-10) )
= 10.0
pOH = -log[OH-]
pH + pOH=14
**SI digit rule
Fig 18-6 (p 435)
Most solutions fall between 0 and 14 on the pH scale.
Hydronium ion
conc
Hydroxide ion
conc
pH
pOH
Acids
Neutral
Bases
A change in pH of one unit is a 10X difference in strength(conc)
Dilutions of acid or base with water decrease the strength of the acid or base
but CANNOT change an acidic solution into a basic one! (or vice versa)
Can you find the concentration when give the pH or pOH?
Ex 1. If the pH of a solution is 11.638, what is the
a) pOH?
b) [H3O+]?
c) [OH-]?
Ex 2. What mass of potassium hydroxide is required to make 250mL of solution with a
pH of 11.70?
Read pH and acid deposition on page 238
‘pH calculations’ worksheet
‘more practice’ worsheet
Sheet 6.3.4
Sheet 6.3.8 after dilution
Indicators are substances that change colour at different pH values
Indicators are used to differentiate amoung acidic, basic, and neutral solutions.
Sheet 6.3.6
Lab 6D with indicators-6.3.7
6.3.13 putting it all together
6.3.14 review
Titrations 8.3???