Chapter 13
Water and Properties of Liquids
Liquids have intermediate properties between solids
and gases. Liquids are almost incompressible, have
definite volume and assume the shape of the container.
Densities of
liquids are
usually lower
than that of their
solids. Water is
an exception.
Evaporation or vaporization is the escape of
molecules from liquid into gaseous state.
During evaporation, liquid that stays behind is
cooler. The opposite process is condensation.
Sublimation is the escape of molecules directly
from solid into gas, bypassing liquid state.
Vapor pressure is the pressure exerted by a gas at
evaporation
equilibrium with its liquid, so that:
liquid
gas
Vapor pressure depends only on
condensation
temperature, not on the amount of
liquid.
Open container
completely evaporates.
Closed container
reaches equilibrium
between liquid and gas.
Vapor Pressure Measurement
1 atm = 760 torr
20 oC
20
oC
a.
b.
a. The system is evacuated.
Manometer attached to the
flask shows equal pressure
in both legs.
b. Water is added.
Liquid evaporates.
Manometer shows
increase in pressure.
20 oC
30 oC
c.
d.
c. Equilibrium established.
Manometer shows constant
pressure difference, 17.5 torr.
d. Temperature raised to 30 oC.
Equilibrium reestablished.
Manometer shows constant
pressure difference of 31.8 torr.
Vapor pressure
and temperature
1 atm = 760 torr
Vapor pressure of any
gas at the boiling point
is equal to the
atmospheric pressure.
Vapor pressure of
ethyl ether is the
highest at any temp.
TBP
TBP
Vapor pressure:
Ether > Alc. > Water.
Rate of evaporation:
Ether > Alc. > Water.
proportional to vapor
pressure.
TBP
Volatility
Boiling point:
Ether < Alc. < Water
Substances that readily evaporate are volatile.
Vapor pressure of ethyl ether at 20 oC: 442.2 torr
Volatile
Vapor pressure of water at 20 oC: 17.5 torr
Vapor pressure of mercury at 20 oC: 0.0012 torr
Moderately volatile
Nonvolatile
Boiling Point Curves
Normal Boiling Point
Boiling point at standard pressure
(1 atm, or 760 torr).
Each point on the curve represents a
vapor-liquid equilibrium at a
particular temperature and pressure.
At 500 torr, ethyl ether boils at
~22 oC, alcohol at ~68 oC, and
water at 89 oC.
Freezing or Melting Point
The temperature at which the solid
and liquid are in equilibrium.
Changes of State
Majority of substances change phases
upon heating: solid liquid gas.
1 atmosphere
pressure
TBP ethyl ether
TBP alcohol
TBP water
34.6oC
78.4oC
100.0oC
Heating curve for a pure
substance
CO2 is an exception (dry ice sublimes).
A – B: solid state
B – C: melting
C – D: liquid state D – E: evaporation
E – F: vapor state
Temperature is constant during melting
and boiling – all heat used to break
solid (at boiling point) or liquid forces.
liquid
solid
evaporation
condensation
melting
freezing
gas
liquid
Heat of Fusion and Heat of Vaporization
We learned before that amount of heat
Qheating = (mass) (spec.heat) (temp.change)
depends on mass and temp. change.
Energy (heat) needed to change 1 g
Energy (heat) needed to change 1 g
of a liquid at its boiling point into
of a solid at its melting point into Constant
liquid is heat of fusion.
temperature! vapor is heat of vaporization.
Qfusion = (mass) (spec.heat of fusion)
Qvaporization = (mass) (spec.heat of vaporization)
Example 1: How many joules is
needed to change 20.0 g of ice at
0 oC to steam at 100. oC?
Qheating = (mass) (spec.heat) (temp.change)
Qtot = Qfusion + Qheating + Qvaporization
Qfusion = (20.0 g) x (335 J/g)
Qheating = (20.0 g) x (4.184 J/goC) x (100. oC)
Qvaporization = (20.0 g) x (2260 J/g)
Hydrogen Bond
}
Qtot = 60.3 kJ
produces unusually high melting & boiling point
Hydrogen Bonding (cont.)
H bonding exists between H directly bonded
to one of the three most electronegative
elements (Fluorine, Oxygen, and Nitrogen),
and F, O or N of another molecule.
H bond
..
..
.
.
.
H–O:
H–O:
|
|
H
H
H
bonded
to O
No H bond
H
H
| .. |
H–C–O–C-H
| .. |
H
H
Ethyl ether
Surface Tension and Capillary Action
A droplet of liquid
falling forms a
sphere due to
attractions to other
liquid molecules –
surface tension.
Cohesive forces within
Spontaneous rise mercury liquid (left) are
of liquid in a
stronger than adhesive
forces between Hg and
narrow tube –
walls of the container.
capillary action.
Opposite is true for H O.
2
No H bonded
to F, O, or N
H bonds are
intermolecular
forces.
Hydrates
Some ionic solutions retain water upon evaporation. It becomes
the part of the crystalline compound – water of crystallization.
The formula is written as: ionic compound, dot , # water moleculesL
CuSO4 5 H2O and name them by adding # (Latin) hydrate.
.
Copper(II) sulfate pentahydrate.
Hydrates are true compounds and the water is an integral part of it.
Formula mass CuSO4 5 H2O: 63.55+32.07+64.00+5x18.02 = 249.7
Percent composition of water is (5x18.02 / 249.7) x 100 = 36.08%
dry CuSO4 – white
Hydrate = blue
.
Water can be removed by intense heat: CuSO4
The reaction is reversed when water is added.
Water, a Unique Liquid
Water indicator
. 5 H O(s) CuSO (s) + 5 H O(g)
2
4
2
δO
H
H
Water covers ~75% of Earth. 97% of water is in the oceans. Only
3% is fresh water, of which 2/3 is locked up in ice polar caps.
δ+
Solid form (ice) has lower density than liquid water.
Water is very stable molecule, can stand temperatures up to 2000 oC. It does not
conduct electricity when pure, but decomposes into H2 and O2 in solutions of ions.
2 H2 + O2 --> 2 H2O + 484 kJ
Water can be formed by
2 C2H2(g) + 5 O2 4 CO2 + 2 H2O(l) + 1212 kJ
Combustion,
Neutralization,
HCl(aq) + NaOH(aq) --> NaCl(aq) + 2 H2O
Metabolic reaction
C6H12O6(aq) + 6 O2 6 CO2(g) + 6 H2O(l) + 2519 kJ
Water reactions with metals:
Cold water reacts with Na, K, Ca:
Steam reacts with Zn, Al and Fe:
Reactions of water
Na + H2O H2 + NaOH
Fe + H2O(g) --> H2 + Fe3O4
Remind yourself of the activity series: the above six metals are the most active. Another
three metals are more active than H: Pb, Sn, and Ni and react with acids only; Cu, Ag,
Hg and Au are below H in the series and do not react with acids or H2O.
Water also reacts with certain nonmetals.
Anhydride means:
without water.
Most reactive:
2 F2 + 2 H2O(l) --> 4 HF(aq) + O2
Less reactive:
Cl2(g) + H2O(l) HCl(aq) + HOCl(aq) To test whether a metal or
nonmetal is an anhydride, try
Least reactive: C(s) + H2O(g) CO(g) + H2(g)
Water reacts with oxides of metals and non-metals:
Basic anhydride: CaO + H2O(l) Ca(OH)2(aq)
Acid anhydride:
CO2(g) + H2O(l) H2CO3(l)
Water Purification
Screening, flocculation and sedimentation,
sand filtration, aeration, disinfection.
Hard water contains Mg2+ and Ca2+ ions
Additional water purification is done by
distillation, Ca2+, Mg2+ precipitation, ion
exchange and demineralization.
to remove H2O until all
hydrogen is removed.
Ca
OH ∆
OH
CaO + H2O
H2SO4 ∆ SO3 + H2O
Solutions
Chapter 14
Homogeneous mixture
Solute
separation
A solution is a homogeneous mixture of
Heterogeneous mixture
two or more substances. Its composition is
the same throughout, i.e. concentration
of the substances is the same.
A solution consists of solute (or
solutes) present in lesser amount, and
solvent, present in greatest amount.
Solvent / solute separation
A solution can be made of substances
steps require energy.
Solvent separation
in different phases.
Solvation step
Solvent solute
example
must
Gas
gas
air
overcome
gas
solid
dusty air
separation
gas
liquid
humid air
steps for the
Liquid
solid
sugar solution
substance to
Solvation
liquid
liquid
vinegar
be
soluble.
liquid
gas
soda
Solid
solid
all alloys
Grinding and mixing two or more
solids will never yield a true solution. A
liquid (or gaseous) phase is needed.
Attractive forces release energy.
Solubility of ions in water
insoluble
soluble
all Na+, K+, NH4+ most CO32-, PO43-,
all NO3-, C2H3O2- OH-, S2-.
(except Na+,K+,NH4+)
most SO42-, Cl-,
Br-, I-
Solubility and effects of P,T
Solubility is the maximum mass of
solute that dissolves in 100 g of solvent.
Solubility of most solids increases with
an increase in temperature.
Solution containing the maximum mass
of solute in 100. g H2O is called
saturated solution; this mass can be
estimated from the solubility graph.
Example 1: prepare saturated solution of
NaNO3 in 250. g of water at 20 oC.
250. g H2O x 90. g NaNO3 = 230 g NaNO3
100. g H2O
Soaps and detergents wash (non-polar)
grease by the action of their molecules,
consisting of polar head and non-polar
tail. “Like dissolves like”
Solubility of gases decreases with
temperature, but increases with pressure.
“Tadpole’s”
tail is
attracted to
non-polar
grease.
Water
makes
hydrogen
bonds with
polar head.
Saturated solution of NaNO3 is 90. g NaNO3 in 100. g H2O at 20 oC. If one
adds 100. g NaNO3 in 100. g H2O, 90. g dissolves, 10. g remain undissolved.
At a given temperature, the mass of solid that goes into solution
equals to the mass that precipitates from the saturated solution.
solute (undissolved)
solute (dissolved)
Unsaturated solution contains less
solute per 100 g H2O than
Supersaturated solution has more
the saturated solution.
Supersaturated solutions are unstable, a small disturbance will cause rapid precipitation of the excess of solute.
Cooling of a saturated solution makes the excess of solute precipitate.
Rate of dissolving depends on temperature, particle size, concentration and stirring.
In solid state no reaction will occur:
Solution as a Reaction Medium
NaCl(s) + AgNO3(s) no reaction
Water breaks the crystal lattice, reaction
Ions are locked within the crystal lattice.
proceeds. NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq)
Na+(aq) + Cl-(aq) + Aq+(aq) + NO3-(aq) AgCl(s) + Na+(aq) + NO3-(aq)
Percent Composition
% composition by Grams of solute x 100
Volume of solution
mass / volume
Concentration can be expressed as
% composition by
Volume of solute x 100
percent composition. There are 3 types.
volume / volume
Volume of solution
Amount of solute
Example 2: What is the mass % composition
x 100
Concentration =
Amount of solution
of 273 g solution that contains 35.0 g NaCl?
Grams of solute
% composition by mass Grams of solute x 100
Mass % = Grams of solution x 100
Grams of solution
35.0 g NaCl
=
x 100 = 12.8 mass %
Units must be specified to prevent ambiguity!
273 g solution
Chemists usually work with moles.
Molarity
Molarity is defined as the number of moles
of solute dissolved in 1 L of solutionL
Molarity (M) = Moles of solute
1 L of solution
Not # of moles of solute plus 1 L solvent!
To prepare 1 L of NaCl, measure molar
mass of NaCl, pour it to a 1L volumetric
flask with enough water to dissolve it,
and then add water to the mark.
Another way to do it is to use more
concentrated solution (called stock
solution), by dilution equation:
Mstock soln. x Vstock soln. = Mdiluted soln. x Vdiluted soln.
Moles before dilution = Moles after dilution
Find the number of milliliters of the stock
solution and dilute it with water.
Moles solute x Liters of solution = Moles solute
M
x Vdiluted soln.
Vstock soln. = diluted soln.
Liters solution
Mstock soln.
Dilution does not affect the
number of moles (M x V)!
M
V
Example 3a: Prepare 500. mL of 0.15 M
NaCl using solid NaCl.
1 L solution
0.15 mol NaCl
500. mL x
x
1000 mL
1 L solution
58.44 g NaCl
x
= 4.4 g NaCl
1.00 mol NaCl
Example 3b: Prepare 500. mL of 0.15 M
NaCl using stock solution of 0.50 M NaCl.
0.15 M NaCl x 500. mL soln..
Vstock soln. =
0.50 M NaCl
= 150 mL stock solution + 350 mL water
Molar mass of solute
dilute conc.
M1V1=M2V2
V2 =
Example 4: How many grams of NaCl is in
200. mL of 1.00 M NaCl?
1 L solution
1.00 mol NaCl
200. mL x
x
1000 mL
1 L solution
58.443 g NaCl
x
= 11.7 g NaCl
1.00 mol NaCl
M1V1
M2
Example 5: Prepare 400. mL of 2.00 M NaCl
using 3.00 M NaCl stock solution.
2.00 M NaCl x 400. mL soln.
Vstock soln. =
3.00 M NaCl
= 267 mL stock solution + 133 mL water.
Molality (m) is the number of moles of solute per kilogram of solvent. Note the
difference in writing molarity (M) and molality (m) and the difference in their
definitions. m = mol solute / kg solvent, therefore, molality is independent of volume.
Colligative Properties of Solutions
NaCl is often used to “melt” ice from streets; ethylene glycol/water mixture in car
radiators works both as antifreeze and boiling point elevation agent.
Freezing Point Depression, Boiling Point Elevation and vapor-pressure lowering
are known as colligative properties of solutions. They depend on the number of
solute particles in a solution and not on the nature of the particles.
Reasoning: Vapor pressure is the tendency of water molecules to escape from the liquid
surface into gas. If, say, 10 % of molecules on the surface molecules are nonvolatile, the
resulting effect is that the vapor pressure is lowered.
If the vapor pressure is lowered, the boiling point is elevated because liquid boils
once its vapor pressure equals that of the atmospheric pressure.
Also, if the vapor pressure is lowered,
so is the freezing point because liquid
vapor pressure curve of the solution
does not intercept the solid vaporpressure at the freezing point of the
pure solution, but rather below it.
∆tf = m x Kf
∆tb = m x Kb
oC
mol solute oC kg solvent
=
x
kg solvent
mol solute
Example 6: Find molality of solution
prepared by dissolving 150.0 g C6H12O6
(molar mass 180.156) in 600.0 g H2O.
150.0 g x 1 mol
= 0.8326 mol
180.156 g
molality = 0.8326 mol / 0.6 kg = 1.388 m
Example 7: what is the freezing point of
solution made by dissolving 100.0 g
ethylene glycol in 200. g H2O?
molar mass C2H6O2 = 62.05 g
100.0 g (1mol / 62.05 g) = 1.61 mol
molality = 1.61 mol / 0.2 kg = 8.05 m
Kf = 1.86 oC kg solvent/mol solute
∆tf = m x Kf = 8.05 (mol solute/kg solvent) x
(1.86 oC kg solvent/mol solute) = 15.0 oC
The freezing point of solution is lowered
by 15 oC from that of solvent (water), so
The actual freezing point is:
0 oC – 15.0 oC = -15.0 oC (5.00 oF).
Osmosis
L is diffusion of a liquid through a semipermeable membrane. The membrane
allows solvent (usually water) to diffuse, but prevents diffusion of larger molecules.
Solvent diffuses from the place with more solvent (water) to the place with less solvent.
Blood cells: in blood plasma (0.15 M NaCl, 0.9% saline), they are normal.
in hypertonic solution (1.6% saline) the blood cells shrink because water leaves the
cell plasma. In hypotonic solution (0.2%
saline) blood cells swell because water
diffuses into them.
Osmosis is also the reason bacteria
cannot survive in sugar solutions.
Plants take H2O from ground by osmosis.
normal
hypertonic
hypertonicsoln.
soln. hypotonic
hypotonicsoln.
soln.
HW, Chapter 13: 7, 11 17, 25
7. Name these hydrates: BaBr2 2H2O;
AlCl3 6H2O; FePO4 4H2O.
11. In which of the following substances
would you expect to find hydrogen
bonding? C3H7OH; H2O2; CHCl3; PH3;
HF.
17. How many moles of compound are in
25.0 g of Na2CO3 10 H2O
25. How many joules of energy are
needed to change 275 g of H2O from
15 oC to steam at 100 oC?
.
.
.
.
1.
3
21.
27.
33.
HW Chapter 14: 1, 3, 21, 27(a-c), 33.
Oft the following substances, which ones are
generally soluble in water? (See fig. 14.2):
AgCl; K2SO4; Na3PO4; NaOH; PbI2; SnCO3.
Calculate the mass percent of the following
solutions: 15.0 g KCl + 100.0 g H2O; 2.50 g
Na3PO4 + 10.0 g H2O; 0.20 mol NH4C2H3O2 +
125 g H2O; 1.50 mol NaOH in 33.0 mol H2O.
What will be the molarity of the resulting
solutions made by mixing of the following?
Assume volumes are additive: 125 mL of 5.0
M H3PO4 with 775 mL H2O; 250 mL of 0.25 M
Na2SO4 with 750 mL of H2O; 75 mL of 0.50 M
HNO3 with 75 mL of 1.5 M HNO3.
Use the eq. to calculate the following:
Ca(NO3)2(aq) + 2 Na3PO4(aq) Ca3(PO4)2(s) + 6 NaNO3(aq)
the moles of Ca3(PO4)2 produced from 2.7
mol Na3PO4; the moles NaNO3 produced
from 0.75 mol Ca(NO3)2; the moles Na3PO4
required to react with 1.45 L of 0.225 M
Ca(NO3)2.
What is the molality, freezing point, and
boiling point of the solution containing 2.68 g
of naphtalene (C10H8) in 38.4 g of benzene
(C6H6)?