Water
Focus 1: Water
is distributed
on Earth as a
solid, liquid
and gas
Definitions
• Solution: A
homogeneous mixture
of one substance
dissolved in another.
• Solute: The substance
that is dissolved
• Solvent: The
substance that does
the dissolving (usually
a liquid)
• Examples: salt in
water, iodine in
alcohol
Outcome 1
Density
• All these cubes have a volume of 1 cm3
11.3g
lead
7.9g
0.3g
iron
Wood
8.9g
copper
• Can you arrange them in order of mass?
• The density of a substance is its mass per unit volume.
Outcome 5
Calculating Density
Density = Mass
Volume
Units: g/cm3 or g/mL
Outcome 5
Calculating the density of water and
ice
• Discuss: How would you determine the
density of:
– Liquid water
– Ice
• Write down the steps you would take for
each.
• Carry out the investigation!! (Use worksheet)
Outcome 5
Homework:
• Use your textbook to answer outcomes 2-4
• Exercises 1 – 2 on p.185-186
Explaining the density of water vs ice –
a model
• Complete practical investigation “Using
models to explain the densities of water and
ice”
Investigating some properties of water
• First we will compare these properties of
water with other substances:
– Melting and boiling points
– Surface tension
– Viscosity
• Then we will study intermolecular forces in
order to explain these unusual properties!
Comparing the melting and boiling
points of the hydrides
Use Excel and the worksheet “Boiling and
Melting point data” to graph the melting and
boiling points of different element hydrides.
Group 4
Group 5
Group 6
Group 7
Period 2
CH4
NH3
H2O
HF
Period 3
SiH4
PH3
H2S
HCl
Period 4
GeH4
AsH3
H2Se
HBr
Period 5
SnH4
SbH3
H2Te
HI
Outcome 15
Boiling points of hydrides
Outcome 8-9
Boiling points of hydrides - conclusions
Tin Hydride – Melting point -146C
Outcome 15
• As you go down a
group, the boiling
points of the hydrides
increase.
• However, H2O, HF,
and NH3 do not fit the
pattern – we would
expect them to have
the lowest boiling
points in their group,
but instead they have
anomalously high
boiling points!
Stations activity
• Complete the investigations of two other
important properties of water – surface
tension and viscosity!
Outcome 14
Surface Tension
• the tendency of a liquid to resist increase in its
surface area.
• A high surface tension means the surface
behaves like a taut skin and “holds in” the water
inside.
• The liquid forms spherical droplets (as this
minimises the surface area) rather than
spreading out as a thin film.
• Measured in J/m2(it is the energy
required to increase the surface
area
by 1m2)
Outcome 14
Can you walk on water?
Outcome 14
What shape has the lowest surface
area for a given volume?
• Answer: a sphere
Outcome 14
Viscosity
– The resistance of a liquid to flow
– The higher the viscosity, the less easily the liquid
flows.
– Measured in N s /m2) (it is the force required in 1
second for the liquid to move 1m2)
Outcome 14
Ballpoint pens
Outcome 14
Motor oils contain “Polymeric viscosity
index improvers”
Outcome 14
Capillary action (non-syllabus)
Outcome 14
Capillary action (non-syllabus)
• The tendency of a liquid to rise up a tube
against the pull of gravity.
Outcome 14
Capillary action – why?
• Cohesive forces = forces between
the molecules of the liquid
• Adhesive forces = forces between
the liquid and the tube walls.
• These two forces are enough to
overcome gravity
• In water, the adhesive forces
between water and glass are
stronger than the cohesive forces
within the water, resulting in a
convex meniscus.
• In mercury, it is the opposite.
Outcome 14
Adhesion
Cohesion
Capillary action in
plants
Outcome 14
Spillproof tablecloth?
Outcome 14
Comparing the surface tension of water
and other liquids
Liquid
Formula
Surface tension at 20°C (J m2)
Water
H2O
7.3 x 10-2
Methanol
CH3OH
2.3 x10-2
Ethanol
CH3CH2OH
2.3 x 10-2
Propanol
CH3CH2CH2OH
2.4 x 10-2
Butanol
CH3CH2CH2CH2OH
2.5 x 10-2
Ethylene glycol CH2(OH)CH2(OH)
4.8 x 10-2
Acetone
CH3COCH3
2.4 x 10-2
Chloroform
CHCl3
2.7 x 10-2
Hexane
CH3CH2CH2CH2CH2CH3 1.8 x 10-2
Mercury
Hg
48 x 10-2
Comparing the viscosity of water and
other liquids
Liquid
Formula
Viscosity at 20°C
( Ns/m2)
Water
H2O
1.00 x 10-3
Ethanol
CH3CH2OH
1.20 x 10-3
Ethylene glycol
CH2(OH)CH2(OH)
19.9 x 10-3
Glycerol
CH2(OH)CH(OH)CH2(OH)
1490 x 10-3
Acetone
CH3COCH3
0.33 x 10-3
Chloroform
CHCl3
0.58 x 10-3
Hexane
CH3CH2CH2CH2CH2CH3
0.33 x 10-3
Mercury
Hg
1.55 x 10-3
Explaining these properties of water
• Summary so far: Water has:
– An unusually high melting and boiling point for its
molecular weight
– The highest surface tension of any molecular liquid
– A high viscosity for its molecular weight
–WHY???????????????????
Comparing water, ammonia and
hydrogen sulfide
1. Construct Lewis Dot diagrams of:
a. Methane
b. Ammonia (NH3)
c. Water
d. Hydrogen sulfide (H2S)
2. Construct molecular models of compounds
a-d
The next section covers outcomes 9-14
Methane - tetrahedral
H
H
C
H
4 Bonding pairs
H
Ammonia – Trigonal Pyramidal
1 non-bonding pair
H
3 Bonding pairs
N
H
H
Water - Bent
2 non-bonding pairs
H
O
2 Bonding pairs
H
Hydrogen sulfide - Bent
2 non-bonding pairs
H
S
2 Bonding pairs
H
Polar Covalent Bonds
• When the two elements in a chemical bond have
different electronegativities, the electrons will be
shared unevenly between them; i.e the electrons will
spend more time near one nucleus than the other.
Unequal sharing
Equal sharing
H
Electronegativity:
2.1
H
H
Electronegativity:
2.1
O
δ+
δ-
Electronegativity:
2.1
Outcome 8-9
Dipole
Electronegativity:
3.5
Polar vs Nonpolar Molecules
• HF
δ+
δ-
H
F
• CO2
δ-
δ+
O
C
Shape = linear
Net dipole present
– molecule is polar
δO
Outcome 8-9
Shape = Linear
No net dipole
– molecule is nonpolar
Polar vs Nonpolar Molecules
δ-
• H2O
O
δ+
H
• NH3
H
δ+
δN
δ+
Shape – Bent
Net dipole present
– molecule is
polar
H
H
δ+
δ+
H
Outcome 8-9
Shape: Pyramidal
Net dipole present
– molecule is polar
Polar vs Nonpolar Molecules
δ+
• CH4
H
Shape – tetrahedral
No Net dipole – molecule
is nonpolar
δ-
δ+
C
δ+
H
H
H
δ+
• BF3
δ-
δ+
F
δ-
F
B
F
δOutcome 8-9
Shape – trigonal planar
No Net dipole – molecule
is nonpolar
Polarity - summary
• A bond is polar if one end is slightly positive and
one end is slightly negative, thanks to different
electronegativities of the atoms.
• A molecule is polar if one end of the molecule is
slightly positive and the other is slightly negative
due to the additive effect of the polar bonds.
• If the molecule’s shape means that the dipoles of
each bond cancel each other out, the molecule is
nonpolar overall.
Polarity - summary
• To decide whether a molecule is polar or not:
1. Use electronegativities to decide the polarity of
the bonds
2. Use the shape of the molecule to decide
whether the polar bonds cancel out or combine
to produce a net dipole.
Intermolecular forces – take notes on w/s
1. Dispersion Forces – occur in all
substances.
2. Dipole - Dipole Forces – occur
in polar substances only
3. Hydrogen Bonding – occur in
polar substances that have an H
bonded to an F, O or N
Outcome 14
Strength
Dipole - Dipole Forces
•
•
Occurs between polar molecules only
Arise from the transient attraction of the positive
pole of one molecule to the negative pole of the
other.
•
Click here to view an animation of dipole-dipole forces
Outcome 14
Hydrogen Bonding
•
•
•
•
•
•
•
A special type of dipole-dipole
force
Occurs between molecules that
have an H atom bonded to an N, O
or F atom.
The H-N, H-O and H-F bonds are
extremely polar, so the electron
density is withdrawn strongly from
H.
As a result, the partially positive H
of one molecule is attracted
strongly to the partially negative
lone pair on the N, O or F of
another molecule.
This attraction is called a hydrogen
bond.
Click here to view an animation of H bonding
Another one!
Outcome 14
Dispersion forces
•
•
•
Electrons are constantly
moving, so at any instant
they can be unevenly
distributed across a
molecule.
This can leave one end of a
molecule slightly negative
and the other end slightly
positive – i.e there is a
temporary (instantaneous)
dipole.
This induces a dipole in a
neighbour molecule, causing
a short-lived attraction
between them.
Outcome 14
F
F
Comparing the four molecules
Molecule
Shape
Polar or
nonpolar?
Intermolecular Melting
forces present Point
Boiling
Point
H2O
0
100
H2S
-86
-60
NH3
-78
-33
CH4
-183
-162
Outcome 8-9
Comparing the four moleculesanswers
Molecule
Shape
Polar or
nonpolar?
Intermolecular Melting
forces present Point
Boiling
Point
H2O
Bent
Polar
-Dispersion
-Hydrogen
0
100
H2S
Bent
Polar
-Dispersion
-Dipole-Dipole
-86
-60
NH3
Pyramidal
Polar
-Dispersion
- Hydrogen
-78
-33
CH4
Tetrahedral
Nonpolar
- Dispersion
-183
-162
Outcome 8-9
Questions
What causes surface tension?
• Molecules in the interior experience
intermolecular attractions in all directions
• Molecules on the surface experience
intermolecular attractions only from below
and the sides i.e a net attraction downwards
and want to move into the interior.
• The stronger the IMFs in a liquid, the greater
its surface tension
What causes viscosity?
• When a liquid flows, the molecules slide past one
another.
• Intermolecular attractions hinder this movement,
resulting in viscosity (resistance to flow)
• Large, long molecules have higher viscosity than
small, spherical ones.
• Due to strong H bonding, water has a much
higher viscosity than its small molecular size
might suggest.
Writing explanations
1. Explain what causes liquids to have surface
tension.
2. Explain what causes liquids to be viscous
3. Melting and boiling point
Game – matching properties to explanations
Polarity - Summary
• Both ICl and Br2 have the same number of
atoms and approximately the same molecular
weight, but ICl is a solid whereas Br2 is a liquid
Polar compounds like H S, NH and H O have
o
at 0 C. Why?dipole-dipole forces present; nonpolar
2
3
compounds like CH4 do not.
2
Homework: Outcomes 2-4
Polar Covalent Bonds
• When the two elements in a chemical bond have
different electronegativities, the electrons will be
shared unevenly between them; i.e the electrons will
spend more time near one nucleus than the other.
Equal sharing
H
Electronegativity:
2.1
Unequal sharing
H
H
Electronegativity:
2.1
O
δ+
δ-
Electronegativity:
2.1
Dipole
Electronegativity:
3.5
Polar vs Nonpolar Molecules
• HF
H
F
• CO2
O
C
O
Polar vs Nonpolar Molecules
• H2O
O
H
H
• NH3
N
H
H
H
Polar vs Nonpolar Molecules
• CH4
H
C
H
H
H
• BF3
F
F
B
F
Polarity - summary
• A bond is polar if one end is slightly ................
and one end is slightly ..................., thanks to
different .................................of the atoms.
• A molecule is polar if one end of the molecule is
slightly .............. and the other is slightly .............
due to the additive effect of the polar ..............
• If the molecule’s ............... means that the
dipoles of each bond ............... each other out,
the molecule is ..................overall.
Polarity - summary
• To decide whether a molecule is polar or not:
1. Use ............................ to decide the polarity of
the bonds
2. Use the .................... of the molecule to decide
whether the polar bonds cancel out or combine
to produce a net dipole.
Water
Focus 4: Water
is distributed
on Earth as a
solid, liquid
and gas
Heat vs temperature
• Temperature:
– how “hot” or “cold” something feels.
– Measured in degrees celsius (°C) or in kelvin (K)
• Heat:
– A form of energy
– Heat flows from a hotter object to a colder object
until their temperatures are equal.
– Measured in joules (J)
Heat vs temperature
• Two objects at the same temperature can
contain different quantities of heat
100g water
Temperature: 100°C
1 tonne water
Temperature: 100°C
Heat vs temperature
• Two substances at the same temperature can
contain different quantities of heat
20g water at 100°C
added
100g water at 25°C
Temperature increase: 12.5°C
20g copper at 100°C
added
100g water at 25°C
Temperature increase: 1.5°C
Heat vs temperature
• When given the same amount of heat, two
substances can have different final
temperatures
100g water
Heat on
bunsen burner
for 1 min
Temperature
increase: 20°C
100g ethanol
Heat on bunsen
burner for 1 min
Temperature
increase: 11°C
Specific heat capacity
• The amount of heat required to raise the
temperature of 1g of a substance by 1°C (1K)
Specific heat capacities of various
liquids (textbook. p.223)
Substance
Water
Ethanol
Ethylene glycol
Octane
Acetone
Chloroform
Specific heat
capacity (J K-1 g-1)
4.18
2.44
2.39
2.22
2.17
0.96
• Water has a very high specific heat capacity –
this has implications for living things and the
environment
Worked example 12 in textbook
• Calculate the quantity of heat needed to raise
the temperature of 155g water from 17.0°C to
35.5°C. The heat capacity of water is 4.18 JK1g-1
• Note: A temperature change of 1°C is the
same as a temperature change of 1K
Homework
• Textbook questions 30-32 on p.224
• Use textbook p.226-227 to write “Explain”
answers to each of outcomes 40 and 41.
Heat changes when substances
dissolve
• Exothermic process: - releases heat into
surroundings.
• Eg when some substances dissolve, heat is
released into the surroundings (the water), which
become hotter.
• Endothermic process: - absorbs heat from
surroundings
• Eg when some substances dissolve, heat is
absorbed from the surroundings (the water)
which become colder.
Molar heat of solution (ΔHsoln)
• The heat absorbed when one mole of a
substance dissolves in water
• If ΔHsoln is positive, the process is endothermic
(heat is absorbed: the temperature of the
solution falls)
• If ΔHsoln is negative, the process is exothermic
(heat is released; the temperature of the
solution rises)
Calculating ΔHsoln – worked example 13
in textbook
a. When 11.2g sodium hydroxide at 19.2°C was
dissolved in 200mL water also at 19.2°C, the
temperature rose to 31.4°C. Calculate the
molar heat of solution of sodium hydroxide.
Take the specific heat capacity of the final
solution as 4.2 J K-1 g-1
b. Use this heat of solution to calculate the
expected temperature rise when 23.6g
sodium hydroxide is dissolved in 1.00L
(=1000g) of water
Homework
• Complete exercises 33-36 on p.226