Water Focus 1: Water is distributed on Earth as a solid, liquid and gas.

WaterFocus 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

lead iron Wood copper

• Can you arrange them in order of mass?• The density of a substance is its mass per unit volume.

Outcome 5

11.3g 7.9g 0.3g 8.9g

Calculating Density

Density = MassVolume

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.

Outcome 15

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

Boiling points of hydrides

Outcome 8-9

Boiling points of hydrides - conclusions

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!

Tin Hydride – Melting point -146C- Boiling point -52C

Stations activity

Outcome 14

• Complete the investigations of two other important properties of water – surface tension and viscosity!

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. Methaneb. Ammonia (NH3)

c. Waterd. Hydrogen sulfide (H2S)

2. Construct molecular models of compounds a - d

The next section covers outcomes 9-14

Methane - tetrahedral

4 Bonding pairs

H C H

H

H

Ammonia – Trigonal Pyramidal

3 Bonding pairs

H N H

H

1 non-bonding pair

Water - Bent

H O H

2 Bonding pairs

2 non-bonding pairs

Hydrogen sulfide - Bent

H S H

2 Bonding pairs

2 non-bonding pairs

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.

Outcome 8-9

H H H O

Electronegativity: 3.5

Electronegativity: 2.1

Unequal sharingEqual sharing

Electronegativity: 2.1

Electronegativity: 2.1

δ+ δ-

Dipole

Polar vs Nonpolar Molecules• HF

• CO2

Outcome 8-9

Shape = linearNet dipole present – molecule is polar

Shape = LinearNo net dipole – molecule is nonpolar

δ+ δ- H F

CO O

δ+ δ-

δ-

Polar vs Nonpolar Molecules• H2O

• NH3

Outcome 8-9

Shape – BentNet dipole present – molecule is polar

δ+

δ-

δ+

δ-

δ+

O

H H δ+

N

H

H

H

δ+

Shape: PyramidalNet dipole present– molecule is polar

Polar vs Nonpolar Molecules

• CH4

• BF3

Outcome 8-9

Shape – tetrahedralNo Net dipole – molecule is nonpolarδ+

δ- δ+

F

B

FF δ+

δ-

δ-

δ- Shape – trigonal planarNo Net dipole – molecule is nonpolar

H

H

H

C

Hδ+

δ+

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 bonds2. 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 moleculesMolecule Shape Polar or

nonpolar?Intermolecular forces present

Melting Point

Boiling Point

H2O 0 100

H2S -86 -60

NH3 -78 -33

CH4 -183 -162

Outcome 8-9

Comparing the four molecules- answersMolecule Shape Polar or

nonpolar?Intermolecular forces present

Melting 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 viscous3. 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 at 0oC. Why?

Polar compounds like H2S, NH3 and H2O have dipole-dipole forces present; nonpolar compounds like CH4 do not.

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.

H H H O

Electronegativity: 3.5

Electronegativity: 2.1

Unequal sharingEqual sharing

Electronegativity: 2.1

Electronegativity: 2.1

δ+ δ-

Dipole

Polar vs Nonpolar Molecules• HF

• CO2

H F

CO O

Polar vs Nonpolar Molecules• H2O

• NH3

O

H H

N

H

H

H

Polar vs Nonpolar Molecules

• CH4

• BF3

F

B

FF

H

H

H

C

H

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 bonds2. Use the .................... of the molecule to decide

whether the polar bonds cancel out or combine to produce a net dipole.

WaterFocus 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 waterTemperature: 100°C 1 tonne water

Temperature: 100°C

Heat vs temperature

• Two substances at the same temperature can contain different quantities of heat

100g water at 25°CTemperature increase: 12.5°C

20g water at 100°C added

20g copper at 100°C added

100g water at 25°CTemperature increase: 1.5°C

Heat vs temperature

• When given the same amount of heat, two substances can have different final temperatures

100g waterHeat on bunsen burner for 1 minTemperature increase: 20°C

100g ethanolHeat on bunsen burner for 1 minTemperature 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)

• Water has a very high specific heat capacity – this has implications for living things and the environment

Substance Specific heat capacity (J K-1 g-1)

Water 4.18Ethanol 2.44Ethylene glycol 2.39Octane 2.22Acetone 2.17Chloroform 0.96

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 JK-

1g-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