Why do some solids dissolve in water but others do not?



Download 4.1 Mb.
Date11.09.2018
Size4.1 Mb.
#65335

Questions

  • Why do some solids dissolve in water but others do not?
  • Why are some substances gases at room temperature, but others are liquid or solid?
  • What gives metals the ability to conduct electricity, what makes non-metals brittle?
  • The answers have to do with …
  • Intermolecular forces

Overview

  • There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces
  • We have already looked at intramolecular bonds (ionic, polar, non-polar)
  • Intermolecular forces (IMF) have to do with the attraction between molecules (vs. the attraction between atoms in a molecule)
  • Intermolecular forces

Ionic, Dipole - Dipole attractions

  • We have seen that molecules can have a separation of charge
  • This happens in both ionic and polar bonds (the greater the EN, the greater the dipoles)
  • H
  • Cl
  • +
  • –
  • Molecules are attracted to each other in a compound by these +σ and - σ forces
  • +
  • –
  • +
  • –
  • +
  • –
  • +
  • –

H - bonding

  • H-bonding is a special type of dipole - dipole attraction that is very strong
  • It occurs when N, O, or F are bonded to H
  • Q- Calculate the EN for HCl and H2O
  • A- HCl = 2.9-2.1 = 0.8, H2O = 3.5-2.1 = 1.4
  • The high EN of NH, OH, and HF bonds cause these to be strong forces (about 5x stronger than normal dipole-dipole forces)
  • They are given a special name (H-bonding) because compounds containing these bonds are important in biological systems

Hydrogen Bond

  • Strongest of all “weak” forces
  • Is caused when H is bonded to F, O, or N
  • These are so electronegative that the H is a “naked nucleus” or bare proton
  • Very attractive!
  • Will bond to nearby electron pairs

Importance of Hydrogen Bonding

  • Biological systems – DNA, proteins
  • Water chemistry (MP, BP, specific heat, surface tension)
  • Density of ice

Density

  • Most solids are more dense than liquid
  • Water is less dense because of hydrogen bonding
  • At 4°C, water becomes less dense
  • Important for life in winter
  • Causes lake turnover
  • Alum example

London forces

  • Non-polar molecules do not have dipoles like polar molecules. How, then, can non-polar compounds form solids or liquids?
  • London forces are named after Fritz London (also called van der Waal forces)
  • London forces are due to small dipoles that exist in non-polar molecules
  • Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical
  • The resulting tiny dipoles cause attractions between atoms/molecules

London forces

  • Instantaneous dipole:
  • Induced dipole:
  • A dipole forms in one atom or molecule, inducing a dipole in the other

London Dispersion Forces

  • All molecules have this
  • Only attraction in nonpolar molecules
  • How can Iodine be a solid?
  • Temporary lopsided charge builds up from random motion of electrons - 1930
  • Increases with mass – we say it has greater polarizability
  • Straight molecule is more polarizable than a curled up molecule – why?
  • Halogen Family is a great essay

Ionic, H-bonding, Dipole, or London?

  • Details
  • Bond
  • Molecule
  • IMF
  • EN = 0 - 0.5
  • nonpolar
  • nonpolar
  • London
  • EN = 0.5 - 1.7
  • polar
  • polar
  • dipole-dipole*
  • EN = 1.7 - 3.2
  • ionic
  • Ionic
  • ionic*
  • H + N,O,F
  • polar
  • polar
  • H-bonding*
  • Symmetrical molecule (any EN)
  • --
  • nonpolar
  • London
  • *Since all compounds have London forces. London forces are also present. However, their affect is minor and overshadowed by the stronger forces present. Note: the term “polar” is used interchangeable with “polar covalent”. Likewise, “nonpolar” and “nonpolar covalent” mean the same thing.

Mixing oil and water

  • Lets take a look at why oil and water don’t mix (oil is non-polar, water is polar)
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +
  • +
  • –
  • +

Properties of Liquids

  • Viscosity
  • “Slower than…..
  • Resistance of a liquid to flow
  • Time it as it goes through a small tube with gravity acting upon it.
  • Poise – 1g/cm-s
  • Trends – same substance – decreases with increasing temperature
  • series (same structure) – increases with increasing mass

Surface Tension

  • How many drops on a penny?
  • Uneven forces at surface
  • Acts like pond scum
  • Definition – energy needed to increase the surface area of a liquid by a certain amount
          • Water is high – why?
  • Called “cohesive” force – together
  • Water moving up a stem – adhesive force
  • Capillary acion – rise up a thin tube
  • Meniscus!

Phase Changes

  • Solid to Liquid is called Heat of Fusion Hfus
  • For water, 6 kJ/mol
  • Liquid to Gas is called Heat of Vaporization
  • Hvap
  • For water, 40.7 kJ/mol
  • Hsub is sum of each

Heating Curve

  • Try a problem
  • Remember - flat during phase change, temperature change when heating a single phase
  • Cooling is opposite

Supercooling

  • Happens with some liquids - remove heat and it doesn’t freeze when it should
  • Very unstable
  • May happen during hibernation

Critical Temperature

  • Highest temperature at which a liquid can form from a gas when pressure is applied.
  • Above this, the substance is called a supercritical fluid.
  • Gas just becomes more compressed.
  • Critical pressure - pressure at the critical temperature

Vapor pressure

  • Vapor pressure forms above any liquid if container is closed – why?
  • Equilibrium is reached
  • This is vapor pressure
  • Higher if forces holding liquid together are weak - called a volatile (fleeing) liquid

Boiling Point

  • Temperature at which the VP equals atmospheric pressure
  • Normal BP - boiling point at 1 atm
  • Everest? Autoclave?

Phase Diagram

  • Handout
  • Look at lines
  • Look at slope of AB
  • Freeze-drying - library book example

Water vs. CO2

Structure of Solids

  • Amorphous (rubber, plastics) - large or mixtures - no true structure
  • Crystalline - highly ordered structure
  • Crystalline solids have true melting points

Unit Cell

  • Repeating unit of a solid
  • 7 types – (6-sided parallelograms)
  • Ni, Na, NaCl
  • Array of points in the crystal lattice

3 cubic unit cells

Total Atoms for each unit cell

Packing

Bonding

  • Shown by x-ray diffraction
  • Molecular - low MP
  • If unit packs well, mp can be high
  • Covalent Network Solid - very strong
  • Many covalent bonds in 3-D
  • Diamond, graphite, SiO2, SiC, BN
  • Ionic - greater charge, greater MP
  • Metallic solids - hexagonal close packed, mp varies

Network solids (covalent crystals)

  • There are some compounds that do not have molecules, but instead are long chains of covalent bonds (E.g. diamond)
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • C
  • This happens in 3 dimensions, creating a crystal
  • Because there are only covalent bonds, network solids are extraordinarily strong

Metallic crystals

  • Metals normally occur as solids (high melting points).
  • Thus, there must be strong bonds between the atoms of metals causing them to bond
  • Bonding in metals and alloys is different from in other compounds: positive nuclei exist in a sea of electrons (this explains why metals conduct electricity)
  • +
  • +
  • +
  • +
  • +
  • +
  • +
  • +
  • +

Crystal types

  • There are 6 types of intermolecular forces
  • These forces are associated with certain crystal types. By comparing solids we have a common frame of reference.
  • The crystal types and their basic units are
  • 1) Network (covalently bonded atoms)
  • 2) ionic (electrostatic attraction of ions),
  • 3) metallic (positive nuclei in electron sea),
  • 4) Molecular (electrostatic attraction of dipoles in molecules)
  • a) Polar (dipole-dipole and H-bonding)
  • b) Non-polar (London forces)

Properties of crystals

  • Boiling and melting occur when the forces between molecules are overcome and a change of state occurs
  • The higher the force of attraction between molecules (IMF) the higher the melting/-boiling point (see previous slide for order)
  • Only metallic crystals conduct electricity in solid state (they also conduct in liquid state)
  • Ionic crystals will conduct electricity in molten state or dissolved because ions are free to move to positive and negative poles

Solubility of crystal types

  • Solute = what is dissolving (e.g. salt)
  • Solvent = what it is dissolving in (e.g. water)
  • Strong attractions between the basic units of covalent crystals cause them to be insoluble.
  • Metallic crystals are likewise insoluble
  • The solubility of other crystals depends on solute and solvent characteristics
  • We will see that polar/ionic solutes dissolve in polar/ionic solvents and non-polar solutes dissolve in non-polar solvents
  • This is known as the like-dissolves-like rule

Crystals worksheet

  • Place numbers in boxes according to descriptions
  • Use pg. 431-439 and class notes, and internet as references
  • At end rows should add up to 126 each

Crystals worksheet answers

  • P
  • A
  • M
  • E
  • T
  • C
  • E
  • Metallic
  • 2
  • 7
  • 13
  • 16
  • 23
  • 30
  • 35
  • Network
  • 3
  • 10
  • 11
  • 20
  • 21
  • 29
  • 32
  • Ionic
  • 5
  • 8
  • 12
  • 17
  • 24
  • 27
  • 33
  • Molecular polar
  • 4
  • 9
  • 14
  • 18
  • 22
  • 28
  • 31
  • Molecular non-polar
  • 1
  • 6
  • 15
  • 19
  • 25
  • 26
  • 34
  • For more lessons, visit www.chalkbored.com

Testing concepts

  • Which attractions are stronger: intermolecular or intramolecular?
  • How many times stronger is a covalent bond compared to a dipole-dipole attraction?
  • What evidence is there that nonpolar molecules attract each other?
  • Suggest some ways that the dipoles in London forces are different from the dipoles in dipole-dipole attractions.
  • A) Which would have a lower boiling point: O2 or F2? Explain. B) Which would have a lower boiling point: NO or O2? Explain.

Which would you expect to have the higher melting point (or boiling point): C8H18 or C4H10? Explain.

  • Which would you expect to have the higher melting point (or boiling point): C8H18 or C4H10? Explain.
  • What two factors causes hydrogen bonds to be so much stronger than typical dipole-dipole bonds?
  • So far we have discussed 4 kinds of intermolecular forces: ionic, dipole-dipole, hydrogen bonding, and London forces. What kind(s) of intermolecular forces are present in the following substances:
  • a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2
  • (hint: consider EN and molecular shape/polarity)
  • Challenge: Ethanol (CH3CH2OH) and dimethyl ether (CH3OCH3) have the same formula (C2H6O). Ethanol boils at 78 C, whereas dimethyl ether boils at -24 C. Explain why the boiling point of the ether is so much lower than the boiling point of ethanol.
  • Challenge: try answering the question on the next slide.

Testing concepts

  • Intramolecular are stronger.
  • A covalent bond is 100x stronger.
  • The molecules gather together as liquids or solids at low temperatures.
  • London forces
    • Are present in all compounds
    • Can occur between atoms or molecules
    • Are due to electron movement not to EN
    • Are transient in nature (dipole-dipole are more permanent).
    • London forces are weaker

Testing concepts

  • A) F2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger London attractions and higher melting/boiling points.
  • B) O2 because it has only London forces. NO has a small EN, giving it small dipoles.
  • C8H18 would have the higher melting/boiling point. This is a result of the many more sites available for London forces to form.
  • 1) a large EN, 2) the small sizes of atoms.

Testing concepts

  • a) NH3: Hydrogen bonding (H + N), London.
  • b) SF6: London only (it is symmetrical).
  • c) PCl3: EN=2.9-2.1. Dipole-dipole, London.
  • d) LiCl: EN=2.9-1.0. Ionic, (London).
  • e) HBr: EN=2.8-2.1. Dipole-dipole, London.
  • f) CO2: London only (it is symmetrical)
  • Challenge: In ethanol, H and O are bonded (the large EN results in H-bonding). In dimethyl ether the O is bonded to C (a smaller EN results in a dipole-dipole attraction rather than hydrogen bonding).

H – bonding and boiling point

  • why does BP as period , why are some BP high at period 2?

Testing concepts

  • Boiling points increase down a group (as period increases) for two reasons: 1) EN tends to increase and 2) size increases. A larger size means greater London forces.
  • Boiling points are very high for H2O, HF, and NH3 because these are hydrogen bonds (high EN), creating large intermolecular forces
  • For more lessons, visit www.chalkbored.com


Download 4.1 Mb.

Share with your friends:




The database is protected by copyright ©www.sckool.org 2022
send message

    Main page