Transcript Document

Chapter 6: Sports Drink
Introductory Activity
• What do you think are the benefits of drinking
a sports drink while exercising rather than
plain water?
• How are your ideas influenced by the
marketing strategies of the companies that
sell these drinks?
Sports Drinks
• This chapter will introduce the chemistry
needed to understand how Sports Drinks work
– Section 6.1: Solutions & electrolytes
– Section 6.2: Concentrations of solutions
– Section 6.3: Acidity & pH
– Section 6.4: Solubility & precipitates
– Section 6.5: Stoichiometry
– Section 6.7: Limiting Reactants
– Section 6.6: Properties of solutions
Sports Drinks
Is a
Differ from pure
liquids in
Properties
Solution
With
How much
solute is in it?
Concentrations
Electrolytes
Some affect
that need to all
dissolve when mixed
together
Solubility
pH
Can be
determined by
Titrations
Section 6.1—Solutions &
Electrolytes
What are those “electrolytes” they say you’re replacing by drinking sports drinks?
Dissolving substances
• Substances are dissolved by a process called
hydration
– The solvent and solute need to break
intermolecular forces within themselves
– New intermolecular forces are formed between
the solvent and solute
– The solvent “carries off” the solute particles
Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
Water molecules are polar and
they are attracted to the charges
of the ions in an ionic
compound.
-
+
-
+
-
+
-
+
-
When the intermolecular forces
are stronger between the water
and the ion than the
intramolecular between the ions,
the water carries away the ion.
Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
As more ions are “exposed” to
the water after the outer ions
were “carried off”, more ions can
be “carried off” as well.
+
-
+
-
+
-
+
-
Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
+
-
These free-floating ions in the
solution allow electricity to be
conducted
-
+
-
-
+
+
-
Electrolytes
• When there are free-floating charges in a
solution then it can conduct electricity.
• Things that produce free-floating charges
when dissolved in water are called
electrolytes.
Dissolving Covalent Compounds
+
-
Solvent, water (polar)
+
Solute, sugar (polar)
-
-
+
+
-
+
Polar covalent molecules are
formed in the same way—water
forms intermolecular forces with
the solute and “carries” the
solute particles away.
+
-
+
Dissolving Covalent Compounds
+
-
Solvent, water (polar)
+
Solute, sugar (polar)
-
-
+
+
+
-
+
-
+
However, the polar covalent
molecules themselves do not
split into charged ions—the
solute molecule stays together
and just separates from other
solute molecules.
Non-electrolytes
• When molecules separate from other
molecules (breaking intermolecular forces),
but free-floating charges are not produced
from breaking intramolecular forces, the
solution cannot conduct electricity.
• These are called non-electrolytes
Types of Electrolytes
Strong Electrolytes
Weak Electrolytes
Non-Electrolytes
Ionic compounds
Ionic Compounds
Covalent Compounds
Almost all ions are
separated when
dissolved in water.
Only a few ions are
separated when
dissolved in water
No molecules
separate—ions are not
formed
Easily conducts
electricity when
dissolved in water
Conducts electricity
slightly when dissolved
in water
Does not conduct
electricity at all when
dissolved in water
Breaking up Electrolytes
• Leave polyatomic ions in-tact (including the
subscript within the polyatomic ion)
• All subscripts not within a polyatomic ion
become coefficients
• Be sure to include charges on the dissociated
ions!
Example:
Break up the
following ionic
compounds into
their ions
KNO3
Ca(NO3)2
Na2CO3
Misconceptions about dissolving
• People often describe something that
dissolves as having “disappeared”
• Before the solute dissolves, it’s in such a large
group of particles that we can see it.
• After dissolving, the solute particles are still
there—they’re just spread out throughout the
solution and are in groupings so small that our
eyes can’t see them
Types of Solutions
Unsaturated
More solute can be
dissolved
Saturated
Super-Saturated
No more solute can be
dissolved—it’s “full”
Has more solute than
would make a
saturated solution
dissolved
In general, the higher the temperature of a solution, more solid can be dissolved.
Section 6.2—Concentration
How do we indicate how much of the electrolytes are in the drink?
Concentrated versus Dilute
solute
Lower
concentration
Not as many solute (what’s
being dissolved) particles
solvent
Higher concentration
More solute (what’s being
dissolved) particles
Concentration
• Concentration gives the ratio of amount
dissolved to total amount
• There are several ways to show concentration
Percent Weight/Volume
• This is a method of showing concentration that
is not used as often in chemistry
• However, it’s used often in the food and drink
industry
– For example, your diet drink can might say you have
less than 0.035 g of salt in 240 mL.
– That would give you a concentration of
0.035 g /
240 mL, which is 0.015% solution
%( W / V ) 
grams solute
mL solvent
 100
%(W/V) Example
%( W / V ) 
grams solute
mL solvent
Example:
If you dissolve 12 g of
sugar in 150 mL of
water, what percent
weight/volume is the
solution?
 100
Concentration using # of molecules
• When working with chemistry and molecules, it’s
more convenient to have a concentration that
represents the number of molecules of solute
rather than the mass (since they all have different
masses)
• Remember, we use moles as a way of counting
molecules in large numbers
Quick Mole Review
• 1 mole = 6.02 × 1023 molecules
• The molecular mass of a molecule is found by
adding up all the atomic masses in the atom
• Molecular mass in grams = 1 mole of that
molecule
Quick Mole Example
Example:
How many
moles are
in 25.5 g
NaCl?
Molarity
• Molarity (M) is a concentration unit that uses
moles of the solute instead of the mass of the
solute
M 
moles solute
L solvent
Molarity Example
Example:
If you dissolve 12 g of
NaCl in 150 mL of
water, what is the
molarity?
Molarity Example
Na 1  22.99 g/mole
Cl 1  35.45 g/mole
Example:
If you dissolve 12 g of
NaCl in 150 mL of
water, what is the
molarity?
12 g NaCl
= 22.99 g/mole
= + 35.45 g/mole
58.44 g/mole
1 mole NaCl molecules = 58.44 g
1
mole NaCl
0.21 mole NaCl
= _______
58.44 g NaCl
Remember to change mL to L! 150 mL of water = 0.150 L
M 
moles solute
L solvent
M 
0 . 21 moles
0 . 150 L
1.4 M NaCl
Converting between the two
• If you know the %(W/V), you know the mass
of the solute
• You can convert that mass into moles using
molecular mass
• You can then use the moles solute to find
molarity
Converting from % to M Example
Example:
What molarity is a
250 mL sample of 7.0
%(W/V) NaCl?
Concentration of Electrolytes
• An electrolyte breaks up into ions when
dissolved in water
• You have to take into account how the
compound breaks up to determine the
concentration of the ions
CaCl2  Ca+2 + 2 Cl-1
For every 1 CaCl2 unit that dissolves, you will produce 1 Ca+2 ion and
ions
2 Cl-1
If the concentration of CaCl2 is 0.25 M, the concentration of Ca+2 is 0.25 M and
Cl-1 is 0.50 M
Let’s Practice #1
%( W / V ) 
grams solute
mL solvent
Example:
You want to make
200 mL of a 15%
(W/V) solution of
sugar. What mass of
sugar do you need to
add to the water?
 100
Let’s Practice #2
Example:
What is the %(W/V)
of a 500. mL sample
of a 0.25 M CaCl2
solution?
Let’s Practice #3
Example:
What are the
molarities of the ions
made in a 0.75 M
solution of Ca(NO3)2
Section 6.3—Acidity, pH
How does concentration of acid affect the pH of a sports drink?
A Review of Acids & Bases
Acids – Arrhenius Definition
• Produce Hydronium ion (H3O+1) in water
• Hydronium ion is water + a hydrogen cation
+1
+1
water
By this definition, if an acid is to give a H+1 to water, then all acids
will have hydrogen as the cation (first element written).
How do Acids produce Hydronium?
water
acid
Hydrogen cation with some anion
How do Acids produce Hydronium?
+1
-
How do Acids produce Hydronium?
+1
Hydronium ion
Anion
Bases – Arrhenius Definition
• Bases produce the hydroxide ion in water
-1
Hydroxide Ion
Characteristics of Acids & Bases
Acids
Bases
Produce H3O+1 (hydronium ion)
in water
Produce OH-1 (hydroxide ion) in
water
Tastes sour
Tastes Bitter
React with active metals to
form hydrogen gas
Feels slippery
Strength versus Concentration
Strong versus Weak Acids
-
-
How many hydronium ion – anion pairs
can you find?
3
+
-
How many intact acid molecules can
you find?
1
Strong acid
Most of the acid molecules have
donated the H+1 to water
Strong versus Weak Acids
How many hydronium ion – anion pairs
can you find?
+
1
How many intact acid molecules can
you find?
3
-
Weak acid
Only a few of the acid molecules
have donated the H+1 to water
Concentrated versus Dilute
solute
Lower
concentration
Not as many solute (what’s
being dissolved) particles
solvent
Higher concentration
More solute (what’s being
dissolved) particles
Combinations of Concentration & Strength
Concentrated
Dilute
Strong
A lot of acid
added & most
dissociates
Not much acid
added, but most of
what’s there
dissociates
Weak
A lot of acid
added, but most
stays together
Not much acid
added and most of
what is there stays
together
Acids and Bases as Electrolytes
Acids and bases
dissociate into
ions in water
Free-floating
ions in water
conduct
electricity
Strong acids and bases are strong electrolytes
Weak acids and bases are weak electrolytes
Acids & Bases are
electrolytes
pH
pH Scale
Is a scale to measure the acidity of a sample
1
7
Highly acidic
neutral
Chapter 6 will give more detail about how pH is calculated!
14
Very basic
(not acidic)
pH is a Logarithmic Scale
Logarithm –The number of times a base
must be multiplied by itself to reach a given
number
x  log
# of multiples
b
Base
# you’re trying to reach
y
Calculating pH
pH scale – Logarithmic scale of the acidity
of a solution
The pH scale uses base “10”
pH   log[ H 3 O
[ H 3O
1
]  10
pH has not units
1
 pH
]
[ ] = concentration in
Molarity
The “-” in the pH equation
Because pH is the negative log of concentration of hydronium, as concentration
increases, the pH goes down.
The lowest pH is the highest concentration of hydronium
Concentration of Hydronium ion versus pH
[H3O+]
1
0.5
0
0
0.5
1
pH
1.5
2
What does a “log” scale really mean?
Every change of 1 in pH shows a change of 10x in concentration of hydronium
Level of acidity increases
pH
4
3
2
1
100x
more
10x more acidic
acidic
1000x
more
acidic
An example of calculating pH
Example:
Find the pH if the
concentration of
[H3O+1] is 0.25 M
An example of calculating hydronium
Example:
Find the [H3O+1] if
the pH is 2.7
Auto-ionization of Water
• Water will split into ions
– 2 H2O  H3O+1 + OH-1
• Water will do this to make sure that at 25°C
the following is true:
– [H3O+1] × [OH-1] = 1 × 10-14
• So if you know the hydronium concentration
at 25°C (which can be found from pH), then
you can also find the hydroxide concentration
An example of calculating hydroxide
Example:
Find the [OH-1] if
the pH is 10.7
Let’s Practice #1
Example:
Find the pH if the
concentration of [H3O+1]
is 2.5 × 10-5 M
Let’s Practice #2
Example:
Find the [OH-1] if
the pH is 3.6
Let’s Practice #3
Example:
Find the [H3O+1] if
the pH is 11.2
Section 6.4—Solubility &
Precipitation
How can we make sure everything that’s added to the sports drink will dissolve?
A Review of Double-Replacement
Reactions
Double Replacement Reactions
The cations from two compounds replace each
other.
NaCl + AgNO3  AgCl + NaNO3
Two ionic compounds switch ions
Double Replacement Reactions
General format of a double replacement reaction:
A
X
B
Z
A
Z
B
X
Products of a Double Replacement
1
Ca Cl2
Combine the cation of the first reactant with the anion of the second
reactant
+ Ag NO3
Products of a Double Replacement
2
Ca Cl2
Combine the cation of the second reactant with the anion of the first
reactant
+ Ag NO3
Products of a Double Replacement
3
Remember to write cations first …
& balance charges with subscripts when writing formulas
Only leave subscripts that are in the original compound there if
they are a part of a polyatomic ion!
Ca Cl2
Ca Cl2
+ Ag NO3
+
Ag NO3
Ca(NO3)2
+
AgCl
Precipitation Reactions
Precipitation Reactions
• A precipitation reaction is when 2 soluble
substances are mixed together and they form
an insoluble substance
2 soluble chemicals: NaOH and Cu(NO3)2
Precipitation Reactions
• A precipitation reaction is when 2 soluble
substances are mixed together and they form
an insoluble substance
2 soluble chemicals: NaOH and Cu(NO3)2
2 soluble chemical: NaNO3
1 insoluble chemical (the precipitate): Cu(OH)2
Why do some things dissolve and
others don’t?
Remember the dissolving process?
• Substances are dissolved by a process called
hydration
– The solvent and solute need to break
intermolecular forces within themselves
– New intermolecular forces are formed between
the solvent and solute
– The solvent “carries off” the solute particles
Review--Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
Water molecules are polar and
they are attracted to the charges
of the ions in an ionic
compound.
-
+
-
+
-
+
-
+
-
When the intermolecular forces
are stronger between the water
and the ion than the
intramolecular between the ions,
the water carries away the ion.
Review--Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
As more ions are “exposed” to
the water after the outer ions
were “carried off”, more ions can
be “carried off” as well.
+
-
+
-
+
-
+
-
Review--Dissolving Ionic Compounds
O
H
H
-
water
+
Ionic compound
+
+
-
These free-floating ions in the
solution allow electricity to be
conducted
-
+
-
-
+
+
-
What about with stronger ionic bonds?
O
H
H
2- 2+
water
Ionic compound
+
Ion charge can affect strength
of ionic bond—the higher the
charges, the stronger the
bond.
2-
2+
2-
2+
2-
2+
2-
2+
2-
(How closely the two ions
can pack together also affects
ionic bond strength)
What about with stronger ionic bonds?
O
H
H
2- 2+
water
Ionic compound
+
2-
2+
2-
2+
2-
2+
2-
2+
2-
If the connection between the
water and the ions is not similar
in strength or stronger than the
ion-ion and water-water
connections that are being
broken…
What about with stronger ionic bonds?
O
H
H
2- 2+
water
Ionic compound
+
2-
2+
2-
2+
2-
2+
2-
2+
2-
If the connection between the
water and the ions is not similar
in strength or stronger than the
ion-ion and water-water
connections that are being
broken…
The water won’t be able to carry
the ions away…it won’t dissolve
the solid.
Solubility Rules
Solubility Rules Table
THESE
ANIONS
FORMS SOLUBLE COMPOUNDS
WITH THESE CATIONS
FORM INSOLUBLE COMPOUNDS
WITH THESE CATIONS
NO3- nitrate
Most cations
No common cations
CH3COO–
acetate
Most cations
Ag+
Cl- chloride
Most cations
Ag+, Pb2+, Hg22+, Tl+
Br- bromide
Most cations
Ag+, Pb2+, Hg22+, Tl+
I- iodide
Most cations
Ag+, Pb2+, Hg22+, Tl+
SO42- sulfate
Most cations
Ba2+, Sr2+, Pb2+, Ag+, Ca2+
CrO42chromate
Most cations
Ba2+, Sr2+, Pb2+, Ag+
S2- sulfide
NH4+, cations of column 1,
cations of column 2
Most other cations
OHhydroxide
NH4+, cations of column 1,
and Ba2+ and Sr2+
Most other cations
CO32carbonate
NH4+, cations of column 1
except Li+
Most other cations
PO43phosphate
NH4+, cations of column 1
except Li+
Most other cations
This table, found at the
end of Chpt 6 and in the
Appendix, can help you
figure out which
compounds dissolve (those
that are soluble) and which
form precipitate (insoluble)
Let’s Practice #1
NaNO3
Example:
Decide whether
each is soluble or
not
AgCH3COO
CaBr2
Ba(OH)2
Cu(OH)2
Let’s Practice #2
Example:
Write the
products for this
reaction
Remember to indicate compounds that
dissolve with “aq” for “aqueous” and
compounds that don’t dissolve with “s”
for “solid”
Na2CrO4 (aq) + BaCl2 (aq) 
Let’s Practice #3
Example:
Write the
products for this
reaction
Remember to indicate compounds that
dissolve with “aq” for “aqueous” and
compounds that don’t dissolve with “s”
for “solid”
NaCH3COO (aq) + KCl (aq) 
Section 6.5—Stoichiometry
How can we determine in a lab the concentration of electrolytes?
What do those coefficients really mean?
For every 2 moles of
H2…
and 2 moles of H2O are
produced
2
2
2 H2 + O2  2 H2O
No coefficient = 1
1 mole of O2 is
need to react…
What is stoichiometry?
Stoichiometry – Using the mole ratio from
the balanced equation and information
about one compound in the reaction to
determine information about another
compound in the equation.
Stoichiometry with Moles
Example:
If 4.2 mole of H2 reacts
completely with O2, how
many moles of O2 are
needed?
2 H2 + O2  2 H2O
Stoichiometry with Moles
Example:
If 4.2 mole of H2 reacts
completely with O2, how
many moles of O2 are
needed?
2 H2 + O2  2 H2O
4.2 mole H2
1
mole O2
2
mole H2
From balanced equation:
2 mole H2  1 mole O2
= ________
2.1 mole O2
But we can’t measure moles in lab!
We can’t go to the lab and count or
measure moles…so we need a way to work
in measurable units, such as grams and
liters!
Molecular mass gives the grams = 1 mole of a compound!
Stoichiometry with Moles & Mass
Example:
How many grams of AgCl will be precipitated
if
0.45 mole AgNO3 is reacted as follows:
2 AgNO3 + CaCl2  2 AgCl + Ca(NO3)2
Stoichiometry with Mass
Example:
How many grams Ba(OH)2 are precipitated
from 14.5 g of NaOH in the following
reaction:
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
But what about for solutions?
Molarity

moles solute
L solution
Molarity gives the number of moles of the solute
that are in 1 L of a solution
Stoichiometry with Solutions
Example:
If you need 15.7 g Ba(OH)2 to precipitate,
how many liters of 2.5 M NaOH solution is
needed?
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
Stoichiometry with Solutions
Example:
If you need 15.7 g Ba(OH)2 to precipitate,
how many liters of 2.5 M NaOH solution is
needed?
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
Concentration of NaOH:
2.5 mole NaOH = 1 L
From balanced equation:
2 mole NaOH  1 mole Ba(OH)2
15.7 g Ba(OH)2
1
mole Ba(OH)2
171.35 g Ba(OH)2
Molar Mass of Ba(OH)2:
1 mole Ba(OH)2 = 171.35 g
2
mole NaOH
1
mole Ba(OH)2
1
L NaOH
2.5 mole NaOH
0.0733 L NaOH
= ________
What about gases?
Standard Temperature and Pressure (STP)
– 1 atm (or 101.3 kPa) and 273 K (0°C)
Molar Volume of a Gas – at STP, 1 mole of
any gas = 22.4 liters
Stoichiometry with Gases
Example:
If you need react 1.5 g of zinc completely,
what volume of gas will be produced at STP?
2 HCl (aq) + Zn (s)  ZnCl2 (aq) + H2 (g)
Keeping all these equalities straight!
TO GO BETWEEN
USE THE EQUALITY
Grams & moles
Molecular Mass in grams = 1
mole
moles & liters of a solution
Molarity in moles = 1 L
Moles & liters of a gas at
STP
1 mole = 22.4 L at STP
2 different chemicals in a
reaction
Coefficient ratio from balanced
equation
Titrations—Using Stoichiometry
Titration – Addition of a known volume of a
known concentration solution to a known
volume of unknown concentration solution
to determine the concentration.
End Point
End Point (or Stoichiometric Point) – When
there is no reactant left over—they have all
be reacted and the solution contains only
products
The end point must be reached in order to use stoichiometry to
calculate the unknown solution concentration
Indicators – Paper or liquid that change
color based on pH level.
If the pH of the products is known, the indicator can be chosen to
indicate the end point
Gravemetrics—Using Stoichiometry
Gravemetric Analysis – Using a reaction to
precipitate out an insoluble compound.
The solid is dried and massed.
Stoichiometry can then be used to
determine the original substance’s
concentration from the mass of the
precipitate
Let’s Practice #1
Example:
If you are making 0.57
moles H2O, how many
moles of O2 are needed?
2 H2 + O2  2 H2O
Let’s Practice #2
Example:
If you need to precipitate 10.7 g of
Ba(OH)2, how many grams NaOH are
needed?
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
Let’s Practice #3
Example:
How many moles AgNO3 are needed to react
with 10.7 g CaCl2?
2 AgNO3 + CaCl2  2 AgCl + 2 Ca(NO3)2
Let’s Practice #4
Example:
How many liters of 0.10 M NaOH is
needed to react with 0.125 L of 0.25 M
BaCl2?
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
Let’s Practice #5
Example:
If you produce 15.4 L of H2 at STP, how many
grams of ZnCl2 is also produced?
2 HCl (aq) + Zn (s)  ZnCl2 (aq) + H2 (g)
Section 6.6—Limiting Reactants
What happens if you don’t add reactants in a molar ratio?
Planning a Meal
You go to the grocery store and you buy 1 package of Brats (5 Brats), 1 package
of cheese (16 slices) and 1 package of hot dog buns (8 buns).
If you use all of these…
You can make this many…
5 Brats
5 meals
16 slices of cheese
16 meals
8 hot dog buns
8 meals
So you have the possibility of making 5, 16 or 8 meals…which is it?
You’ll never get the chance to make 8 or 16 meals…you’ll run out of Brats after 5.
Once you run out of one component, you have to stop making meals.
What’s a limiting reactant?
Limiting Reactant – The reactant that runs
out and causes the reaction to stop.
In the previous example, the Brats were the limiting
reactant—once they were gone, you had to stop!
Once even one of the reactants runs out, the reaction stops…it
can’t make any more product.
Limiting Reactant Example
Example:
How many moles of H2O is
produced when 2.3 moles
O2 and 2.3 moles H2 react?
2 H2 + O2  2 H2O
Let’s Practice
Example:
If you react 10.5 g of NaOH and 7.5 g of
BaCl2, how many grams NaCl is produced?
2 NaOH + BaCl2  Ba(OH)2 + 2 NaCl
Section 6.7—Properties of
Solutions
How do all those dissolved things affect the properties of the drink?
What’s Vapor Pressure?
Vapor Pressure – Pressure created above a
sample by particles evaporating from the
sample and becoming gas particles.
Vapor Pressure & Temperature
Temperature is proportional to the average kinetic energy of the molecules.
Average means some will have more and some will have less!
To evaporate,
molecules must
break
intermolecular
forces—this
requires a
minimum
amount of
energy
As temperature
increases, the
average energy
of the
molecules
increase
More molecules
will have the
minimum
needed to
evaporate from
the liquid
As temperature increases, the vapor pressure increases.
Vapor Pressure of Solutions
Only solvent particles on the very top layer of the sample can
evaporate
Looking down on the top of beaker:
If a solvent particle on the top layer has
enough energy, it can break the IMF’s
and evaporate
Solvent particles
Beaker with solvent only
Once evaporated, they cause vapor pressure
Vapor Pressure of Solutions
Only solvent particles on the very top layer of the sample can
evaporate
Looking down on the top of a solution in a beaker:
Solute
particles
Beaker with solvent only
Solvent particles
The solvent and solute form intermolecular forces (connections) with each other.
The solvent must now break those connections in order to evaporate.
The connections are holding the solvent particles back.
Vapor Pressure of Solutions
Only solvent particles on the very top layer of the sample can
evaporate
Looking down on the top of a solution in a beaker:
Beaker with solvent only
The fewer particles that evaporate, the lower the vapor pressure.
The vapor pressure of a solution is always less than the pure solvent.
When does something boil?
Atmospheric pressure pushes down on the
top of the liquid
Molecules are gaining the
energy to break intermolecular
forces and become a gas
Heat source (usually underneath) heats the
molecules closest to it the fastest
When does something boil?
When enough water molecules
turn to gas and create as much
pressure as the atmosphere is
pushing down with, a bubble
can form (counter-act the
atmospheric pressure)
Boiling and Atmospheric Pressure
Boiling occurs when vapor pressure of liquid = atmospheric pressure
Higher altitude
means lower
atmospheric
pressure
The vapor
pressure of the
liquid doesn’t
need to be as
high to boil
with lower
atmospheric
pressure
The bubbles can
form at a lower
temperature
The boiling point of a liquid is always lower at higher altitude
Boiling Points of Solutions
Boiling occurs when vapor pressure of liquid = atmospheric pressure
Solutions have
lower vapor
pressure than
the pure
solvent.
The solution does
not have high
enough vapor
pressure to boil
at the solvent’s
boiling point
The temperature
needs to be
raised until the
vapor pressure
of the solution
= atmospheric
pressure
The boiling point of a solution is always higher than the pure solvent
When do things freeze?
When you’re above the freezing point, solid will melt to liquid
When you’re below the freezing point, liquid will freeze to solid
Freezing point is when there is equilibrium between solid & liquid—the amount of
solid and liquid stay the same
This occurs when the rate of evaporation from the solid is the same as the rate of
evaporation from the liquid
Every time a molecule evaporates from the solid,
one also evaporates from the liquid.
Every time a molecule re-forms into the solid, one
also reforms into the liquid.
Neither one can “get ahead”—it’s at equilibrium
Freezing Points of Solutions
In order for a liquid to freeze, the solid’s vapor pressure and the liquid’s
vapor pressure must be equal
This is the point where the speed of molecules joining to form a solid equals the
speed molecules leave the solid to be liquid
The solid is the
pure solvent.
The liquid is the
solution.
The vapor
pressure of the
liquid (solution)
is lower than
the solid’s
(solvent)
The temperature
is lowered until
the solid’s
vapor pressure
= the liquid
solution’s
vapor pressure
The freezing point of a solution is always lower than the pure solvent
What effects do electrolytes
cause?
When solutes are electrolytes, the impact is greater
Electrolytes break up into more than one particle when added to water.
Therefore, there are even more particles when considering colligative
properties.
For every 1 mole of ___
added
__ moles of particles are in
solution
Sugar (non-electrolyte)
1 (C6H12O6 stays together)
NaCl (electrolyte)
2 (Na+ + Cl-)
CaCl2
3 (Ca2+ + 2 Cl-)
What did you learn about sports
drinks?
Sports Drink
Is a
Differ from pure
liquids in
Properties
Solution
With
How much
solute is in it?
Concentrations
Electrolytes
Some affect
that need to all
dissolve when mixed
together
Solubility
pH
Can be
determined by
Titrations