Properties of Aggregates

1
Properties of Aggregates

• Review Chapters Ch. 3 Ch. 4 (p. 222 251)
• Rest of Chapter 4
• Handouts

2
Review of Grade 11 Chem

• Quantum Mechanical Model of the Atom
• Lewis Dot Structures
• Nature of Chemical Bonds
• VSEPR Theory
• Electronegativity and Polarity

3
Quantum Mechanics (p. 185 203)
4
Quantum Mechanics (p. 185 203)
5
Quantum Mechanics (p. 185 203)

• E level Diagrams
• Electron Arrangements
• Orbital Filling
• Lewis Dot / Valence Levels (orbital filling)
• Electron Configurations
• Core Notation

6
Lewis Dot Structures (p. 224 - 230)

• Based upon the attraction of electrons
• Electronegativtiy
• The attraction for electrons in a chemical bond
• Ionic bonds (Metal and Non-Metal)
• The electrons will go to the more electronegative
  atom, creating an anion and a cation. The
  positive charge attracts the negative charge

7
Lewis Dot Structures (p. 224 - 230)

• Covalent Bonds
• Electrons are shared between an overlap of
  orbitals
• The electrons will be shared and form a bond
• The electrons will be held closer to the more
  electronegative atom and form a slightly positive
  or negative end (polar ends)

8
Lewis Dot Structures (p. 224 - 230)

• Summary for Lewis Dot structures (p. 229)
• When we look at bonds and molecules, there are
  some exceptions
• Electron Deficient Structures (Trigonal Planar)
• Valence Expansion
• Coordinate Covalent Bonds
• Still follow the process of creating a Lewis Dot
  Structure

9
VSEPR Theory (p. 243 250)

• The shape of a molecule is determined by the
  central atoms electrons and the ligands (atoms
  bonded)
• The shape will also help with determining
  polarity of the molecule
• Summary of Shapes of Molecules (p. 245)

10
Polar Molecules (p. 251 255)

• Table p. 251
• Based upon the Electronegativity Difference (EN
  Diff) between the bonded atoms
• Very polar bond Ionic
• Non-polar bond pure covalent bond
• Polar covalent bond there is a difference
  between the bonded atoms that create a dipole
  moment
• Summary p. 255

11
Review Assignment

• For each of the following molecules and
  compounds
• Give the formula
• Draw the Lewis Dot Structure
• Designate polarity, if possible
• State the shape

12

• Magnesium oxide
• Water
• Calcium chloride
• Hydrogen flouride
• Carbon dioxide
• Nitrogen trihydride
• Carbonate ion

• Carbon tetrachloride
• Beryllium chloride
• Boron trihydride
• Phosphorus pentachloride
• Sulfur dichloride
• Chlorotrifluoro carbon

Do question p. 229 11
13
The Structure and Properties of Solids

• Structure and Properties are based upon
• The combination of atoms
• Electronegativity
• Shapes
• Bond Energies
• Radii of Atoms
• Intermolecular Forces

14
Ionic Solids (p. 268)

• Ionic Bonds
• This type of bonding occurs when an element of
  high electronegativity reacts with an element of
  low electronegativity.
• A transfer of electrons produces oppositely
  charged ions, which then exert a non-directional
  electrostatic force of attraction on each other.
• There is no tendency to share electrons to form a
  noble gas electron configuration.

15
Ionic Solids (p. 268)

• The Crystalline Arrangement
• When two atoms are bound via ionic interactions,
  no diatomic molecule is formed, but an infinite
  packing of anions and cations produce the final
  stable ionic compound.
• Both attractive (anion-cation) and repulsive
  (cation-cation, anion-anion) forces result and
  the global arrangement satisfies the rule of the
  minimal electrostatic energy.
• The ions are thus regularly packed in three
  dimensions in order to put neighbours closer with
  opposite charges.

16
Ionic Solids (p. 268)

• A crystalline network or crystal lattice is
  formed in which the global arrangement is
  neutral.
• The stability of the crystal can be measured in
  terms of the crystal lattice energy, which is
  defined as the amount of energy released when one
  mole of an ionic compound is formed from its
  gaseous ions.
• Example Na (g) Cl- (g) ? NaCl (s)
  761 kJ
• The presence of energy on the product side of the
  thermochemical equations indicates that this
  reaction is exothermic (i.e. energy is being
  released).

17
Ionic Solids (p. 268)
18
Ionic Solids (p. 268)

• Properties of Ionic Compounds
• In solid phase at room temperature, they do not
  conduct an appreciable electric current.
• In liquid phase, they are relatively good
  conductors of electricity (but not as good as
  metals)
• They have relatively high melting points and
  boiling points.

Compound Melting Point Boiling Point
KI 686oC 1330oC
MgO 2800oC 3600oC
19
Ionic Solids (p. 268)

• They have low volatilities and low vapour
  pressures.
• They are hard but very brittle.
• Those which are soluble in water form
  electrolytic solutions.

Name Formula Solubility in water at 25oC
Sodium nitrate NaNO3 92 g/100 ml
Barium sulfate BaSO4 0.0002 g/100 ml
20
Ionic Solids (p. 268)

• Factors Affecting Bond Strength and Crystal
  Stability
• Once we have determined that the bond between two
  elements will be ionic rather than covalent, the
  electronegativity of the elements involved no
  longer is a consideration in terms of the
  strength of the ionic bond.
• The two factors that do influence the bond
  strength are related to the force of attraction
  between the ions, which is given by Coulombs Law
• F (k?q1?q2)/r2

21
Ionic Solids (p. 268)

• Where q1 and q2 represent the charges on the ions
  and r is the sum of the ionic radii.
• A stronger force of attraction between ions means
  that the ionic bond will be stronger the
  stronger bond will result in a more stable
  crystal.
• In terms of physical properties, we can expect an
  ionic compound with stronger internal bonding to
  have a higher melting point and a lower
  solubility because more energy will be required
  to separate the ions from each other.

22
Ionic Solids (p. 268)

• Effect of radius
• Cesium ions are larger than sodium ions, but have
  the same charge. Therefore the value of F will be
  small for CsCl than NaCl due to the larger value
  of r.

Compound Mp (oC) Solubility g / 100 ml water at 25oC
CsCl 646 161
NaCl 800 35.7
23
Ionic Solids (p. 268)
Compound Mp (oC) Solubility g / 100 ml water at 25oC
MgO 2800 0.0006
NaF 993 4.22

• Effect of charge
• The sum of the ionic radii in MgO is similar to
  the sum of the ionic radii in NaF, however the
  charges on the ions are larger. (Mg2 compared to
  Na and O2- compared to F-). The larger values
  of q1 and q2 result in a value for F.

24
Ionic Solids (p. 268)
Ionic Solids (p. 268)

• Summary
• Form a crystal lattice
• Held together by electrostatic attraction
• Stronger the intramolecular forces, the higher
  the mp, but lower solubility
• Charge and ionic radius determine the strength of
  the bond

25
Molecules (Section 4.5 pages 257-267)

• Formed from covalent bonds, which incorporates
  the overlap of valence energy levels and the
  electrons are shared to create a pair of
  electrons in a bond. (Intramolecular Force)
  (Bonding Theory p. 231- 240)
• This type of bonding occurs when an element of
  similar electronegativity reacts with an element
  of similar or same electronegativity.
• The overlap and sharing of electrons creates a
  bond composed of two electrons. The sharing of
  electrons creates a noble gas electron
  configuration.

26
Molecules (Section 4.5 pages 257-267)

• If the two bonded elements are the same, a
  non-polar covalent, or pure covalent bond forms.
• If the two bonded elements are of differing
  electronegativities, a polar covalent bond forms,
  creating a dipole moment.
• Each molecule is held close to each other by
  intermolecular forces. The intermolecular forces
  determine the physical properties of the
  molecular substance.

27
Molecules (Section 4.5 pages 257-267)

• Properties of Molecular Crystals
• Neither liquids nor solids conduct an electrical
  current.
• Many exist as gases at Standard Temperature and
  Pressure (STP, 101.3 kPa and 0C) and many solids
  and liquids are relatively volatile.
• The melting points of solid crystals are
  relatively low

28
Molecules (Section 4.5 pages 257-267)

• The boiling points of liquids are also relatively
  low.
• Solids tend to be soft and waxy.
• Although the crystals melt and boil at low
  temperatures, large amounts of energy are
  required to decompose the compounds back into
  their constituent elements.

29
Molecules (Section 4.5 pages 257-267)

• This last property indicates that the
  intramolecular forces (those within the molecule,
  i.e. covalent bonds) are stronger than the
  intermolecular forces (those holding the
  molecules to each other to form the crystal).
• The properties of molecular crystals can vary
  considerably due to both the shape and the size
  of the molecules.

30
Molecules (Section 4.5 pages 257-267)

• Intermolecular Forces
• There are three types of interactions between
  covalent molecules to hold them together in a
  molecular crystal.
• More than one of these interactions can be at
  work in any specific crystal although generally
  one type will predominate.

31
Molecules (Section 4.5 pages 257-267)

• Dipole-dipole forces
• This interaction exists between all polar
  molecules (those that have a permanent dipole or
  separation of charge within the molecule).
• The molecules tend to align themselves so that
  the positive end of one molecule is attracted to
  the negative end of an adjacent molecule.
• The strength of this attraction will vary with
  the degree of polarity of the molecule (its
  dipole moment) with more polar molecules having
  stronger attractions than those which are less
  polar.

32
Molecules (Section 4.5 pages 257-267)

• Dipole Moment

33
Molecules (Section 4.5 pages 257-267)

• London dispersion forces
• While this interaction is present between all
  molecules, it is only of importance with
  non-polar molecules (those with no permanent
  dipole) because it is the only interaction.
• It is a very weak force that results from a
  temporary dipole produced by a fluctuation in the
  electron cloud density.
• The strength of the London force interaction is
  influenced by atomic and molecular size and shape.

34
Molecules (Section 4.5 pages 257-267)

• This dipole, however fleeting, can induce a
  dipole in a neighbouring atom, causing a force.
• This force is always attractive but even shorter
  ranged (and weaker) than permanent dipole-induced
  dipole forces.

35
Molecules (Section 4.5 pages 257-267)

• Factors Affecting the Strength of London
  Dispersion Forces
• Effect of Atomic Size
• Atomic size increases as we go down a column of
  the periodic table, so each successive element
  has more electrons and a larger radius.
• With more electrons, the temporary dipole can be
  larger which will result in a stronger attraction
  between molecules

36
Molecules (Section 4.5 pages 257-267)

• Effect of Molecular Size
• While the atomic size remains the same, the
  surface area increases with the number of atoms
  in the molecule.
• Although the attractions at any one point will
  not change, there will be more attractions
  between adjacent molecules if the molecules are
  larger.

37
Molecules (Section 4.5 pages 257-267)

• Effect of Molecular Shape
• Similarly, a less compact molecular shape will
  produce a larger surface area on the molecule
  which will result in a greater number of
  attractions between adjacent molecules

38
Molecules (Section 4.5 pages 257-267)

• Hydrogen bonding
• This interaction occurs between molecules that
  have hydrogen covalently bonded to a highly
  electronegative element that has a small radius
  (N, O or F).
• The interaction exists between the hydrogen atom
  on one molecule and the highly electronegative
  element on an adjacent molecule.

39
Molecules (Section 4.5 pages 257-267)

• The presence of hydrogen bonds between molecules
  results in boiling points that are higher than we
  might expect based on the trend produced by the
  other hydrides in the same family of elements.

40
Molecules (Section 4.5 pages 257-267)
41
Metallic Crystals (Section 4.6 pages 269 in text)

• What is a metallic bond ?
• Occurs when both atoms have low ionization
  energies and low electronegativities and will
  lose electrons easily (i.e. 2 metals). In
  metallic bonding, positive metal ions are
  arranged with valence electrons delocalized
  around them.
• Since the electrons are delocalized, they are
  mobile and able to move throughout the metal
  structure.

42
Metallic Crystals (Section 4.6 pages 269 in text)

• The structure of a metallic crystal and the
  nature of the metallic bond can be describes as
  follows
• Metallic crystals consist of a three-dimensional,
  closely packed latticework of atomic kernels
  surrounded by a sea of delocalized, mobile
  electrons.
• The electrons can move throughout the crystal
  rather like gas molecules confined in a
  container. The electrons are held within the
  metal by the attraction of the positive atomic
  kernels.

43
Metallic Crystals (Section 4.6 pages 269 in text)

• The atomic kernels are held together by the
  electrostatic attraction of the electrons which
  move between them. It is this force of
  attraction which results in a metallic bond
  between the atoms of the metal.

44
Metallic Crystals (Section 4.6 pages 269 in text)

• The strength of a metallic bond depends on the
  nuclear charge and the number of electrons in the
  outer energy levels.
• As the number of outer electrons increases, the
  strength of the metallic bond increases. This
  trend would be expected, since a greater number
  of electrons would give each atomic kernel a
  higher charge and a greater share of the bonding
  electrons. In general, as the strength of the
  metallic bond increases, the melting and boiling
  points and the hardness of the metal increase.

45
Metallic Crystals (Section 4.6 pages 269 in text)

• In metallic bonding, metal ions are tightly
  packed with their outer shell electrons overlap,
  so each electron becomes detached from its parent
  atom (delocalized). The metal is held together by
  the strong forces of attraction between the
  positive nuclei and the surrounding sea of
  delocalized electrons.
• Examination of the electron configurations and
  ionization energies of atoms of substances with
  metallic properties reveals two common features
  (1) a number of vacant outer shell orbitals and
  (2) low ionization energies.

46
Metallic Crystals (Section 4.6 pages 269 in text)

• Properties of Metals
• Luster or reflectivity.
• High electric conductivity.
• High heat conductivity.
• Workability (malleable and ductile).
• Electron emission caused by heat or light.

47
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• Covalent crystals, also called network solids or
  macromolecules, consist of atoms of the same or
  different elements joined together by a network
  of single covalent bonds.
• This network of bonds can exist in one, two or
  three dimensions to produce a variety of
  properties

48
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• Allotropes elements that exist in different
  physical forms with different physical properties
  but the same chemical properties.
• For example graphite and diamond are allotropes
  of carbon (as is Buckeyball).
• Both form carbon dioxide and water when
  undergoing combustion but only graphite conducts
  electricity and diamond is one of the hardest
  substances known.

49
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• 3-D network examples carbon as diamond,
  quartz, silicon carbide
• These network solids consist of covalently bonded
  atoms which form regular 3-D arrays or crystals.
• Much like an ionic crystal, the intermolecular
  bonds are the same as the intramolecular bonds
  (covalent bonds).

50
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• Due to the very strong bonds, these solids will
  have
• very high melting and boiling points
• will be solids at room temperature
• be extremely hard
• not soluble in polar or nonpolar solvents
• do not conduct electricity.

51
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)
52
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• 2-D network example carbon as graphite, mica
• Networks which form 2-D arrays or sheets.
• The properties of graphite are quite different
  from those of diamond.
• While the melting point and boiling point are
  still very high, graphite is soft and a good
  conductor of electricity.
• The layers will slide over each other allowing
  them to be used as a lubricant.

53
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• Each carbon atom is covalently bonded to only 3
  other carbon atoms to form flat sheets or layers
  with weak London force attractions existing
  between the layers.
• The electrons making up the double bondsare
  actually delocalized throughout the structure
  (i.e. not true double bonds). The layers of
  hexagons are held in place.
• The delocalized electrons are able to move freely
  therefore allowing graphite to conduct
  electricity.
• When you write with a pencil you are breaking of
  layers of graphite

54
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)
55
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)

• 1-D network example asbestos
• These are solids that form networks in a one
  dimensional array or fibre.
• They consist of long chains held together by
  covalent bonds.
• The forces between adjacent chains are very weak
  therefore the solids will form threads.
• They have very high melting and boiling points
  due to the strong covalent bonds.
• They are solids at room temperature and are not
  soluble in water
• Asbestos is composed of silicon, oxygen and
  certain metallic elements such as magnesium and
  calcium.

56
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)
57
Covalent Network Crystals (Section 4.6 pages 270
- 273 in text)