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# Water - PowerPoint PPT Presentation

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)

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### Water

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

• 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

• 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.

11.3g 7.9g 0.3g 8.9g

Outcome 5

Density = Mass

Volume

Units: g/cm3 or g/mL

Outcome 5

• 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

• 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”

• 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!

Use Excel and the worksheet “Boiling and Melting point data” to graph the melting and boiling points of different element hydrides.

Outcome 15

Outcome 8-9

• 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!

Outcome 15

Tin Hydride – Melting point -146C

- Boiling point -52C

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

Outcome 14

• 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

Outcome 14

• Answer: a sphere

Outcome 14

• 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

Outcome 14

Outcome 14

• The tendency of a liquid to rise up a tube against the pull of gravity.

Outcome 14

• 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.

Cohesion

Outcome 14

Outcome 14

Spillproof tablecloth?

Outcome 14

• 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???????????????????

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

4 Bonding pairs

H

H

C

H

H

Ammonia – Trigonal Pyramidal

1 non-bonding pair

3 Bonding pairs

H

N

H

H

2 non-bonding pairs

H

O

H

2 Bonding pairs

Hydrogen sulfide - Bent

2 non-bonding pairs

H

S

H

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

O

H

δ+

δ-

Electronegativity: 2.1

Electronegativity: 2.1

Electronegativity: 2.1

Electronegativity:

3.5

Dipole

Outcome 8-9

Polar vsNonpolar Molecules

• HF

• CO2

δ+

δ-

F

H

Shape = linear

Net dipole present – molecule is polar

δ-

δ+

δ-

Shape = Linear

No net dipole

– molecule is nonpolar

O

C

O

Outcome 8-9

Polar vsNonpolar Molecules

δ-

• H2O

• NH3

O

Shape – Bent

Net dipole present – molecule is polar

δ+

δ+

H

H

δ-

N

Shape: Pyramidal

Net dipole present

– molecule is polar

H

H

δ+

δ+

H

δ+

Outcome 8-9

Polar vsNonpolar Molecules

δ+

H

• CH4

• BF3

Shape – tetrahedral

No Net dipole – molecule is nonpolar

δ-

C

δ+

δ+

H

H

H

δ+

δ-

Shape – trigonal planar

No Net dipole – molecule is nonpolar

F

δ+

F

δ-

B

δ-

F

Outcome 8-9

• 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.

• To decide whether a molecule is polar or not:

• Use electronegativities to decide the polarity of the bonds

• 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

• Dispersion Forces – occur in all substances.

• Dipole - Dipole Forces – occur in polar substances only

• Hydrogen Bonding – occur in polar substances that have an H bonded to an F, O or N

Strength

Outcome 14

• 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

• 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

• 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.

F F

Outcome 14

Outcome 8-9

Outcome 8-9

• 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

• 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.

• Explain what causes liquids to have surface tension.

• Explain what causes liquids to be viscous

• Melting and boiling point

Game – matching properties to explanations

• 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.

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

O

H

δ+

δ-

Electronegativity: 2.1

Electronegativity: 2.1

Electronegativity: 2.1

Electronegativity:

3.5

Dipole

Polar vsNonpolar Molecules

• HF

• CO2

F

H

O

C

O

Polar vsNonpolar Molecules

• H2O

• NH3

O

H

H

N

H

H

H

Polar vsNonpolar Molecules

H

• CH4

• BF3

C

H

H

H

F

F

B

F

• 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.

• To decide whether a molecule is polar or not:

• Use ............................ to decide the polarity of the bonds

• 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

20g copper at 100°C added

100g water at 25°C

Temperature increase: 1.5°C

100g water at 25°C

Temperature increase: 12.5°C

Heat vs temperature

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

100g ethanol

Heat on bunsen burner for 1 min

Temperature increase: 11°C

100g water

Heat on bunsen burner for 1 min

Temperature increase: 20°C

• The amount of heat required to raise the temperature of 1g of a substance by 1°C (1K)

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

• 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 p.223)

• Textbook questions 30-32 on p.224

• Use textbook p.226-227 to write “Explain” answers to each of outcomes 40 and 41.

• Exothermic process: - releases heat into surroundings.

• Eg when some substances dissolve, heat is releasedinto the surroundings (the water), which become hotter.

• Endothermic process: - absorbs heat from surroundings

• Eg when some substances dissolve, heat is absorbedfrom the surroundings (the water) which become colder.

Molar heat of solution ( p.223)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 p.223)ΔHsoln – worked example 13 in textbook

• 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

• 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 p.223)

• Complete exercises 33-36 on p.226