PROPERTIES OF SOLIDSSCH4U1

Intra vs. Intermolecular Bonds• The properties of a substance are influenced by the force

of attraction within and between the molecules.

Intra vs. Intermolecular Bonds

Intramolecular Bonds: Bonds within a molecule (covalent or polar covalent)

Intermolecular Bonds: Bonds between molecules.

Intermolecular Forces• The physical properties of a molecule (e.g. melting point)

are mainly due to the strength of intermolecular bonding:

H2O (s) H2O (l) H2O (g) 2H + O MP = 0oC BP =1000C Decomposes: >2000oC

intermolecular bonds breaking intramolecular bonds breaking

1) Atomic Solids• Noble gases form liquids and solids at very low

temperature due to the very weak bonds between the atoms.

• Since the attraction is so weak, they are weakened and broken at very low temperature.

• E.g. argon (Ar): mp = -189oC; bp = -186oC

van der Waals (London) Forces• Since they do form solids, a very weak attraction must

exist between the atoms.• These are explained by weak attractions between

molecules called van der Waals forces.

• London forces are the weakest type of van der Waals attraction.

London Forces• London forces form due to the attraction between

instantaneous dipoles (charge imbalances) that form in the atoms.

• At low temperature can induce dipoles in other atoms, causing solidification of helium:

London Forces and Electrons• As the number of electrons in an atomic solid increases,

the mp/bp also increases.

Group 18 Electrons Boiling Point ('C)

He 2 -268.6

Ne 10 -245.9

Ar 18 -185.7

Kr 36 -152.3

Xe 54 -107.1

Rn 86 -61.8

Summary: Properties of Atomic Solids

• very low melting points/ boiling point• do not conduct electricity• mp/bp increase down the group (increasing London

forces)

• Liquid helium is a very strange substance.

Molecular Solids• Substances with covalent bonds or polar covalent bonds

(N2, CH4, H2O, C6H12O6 etc.).

• Can exist in all states at room temperature.

Check out liquid nitrogen!

2) Non-Polar Molecular Compounds

• Compounds without bond dipoles only have London Forces between molecules.

• This results in LOW bp/mp

Group 17 Boiling Point ('C)

F2 -188.1Cl2 -34.6Br2 58.8I2 184.4

BOILING POINTS OF RELATED MOLECULAR COMPOUNDS

Formula Number of

Electrons

Boiling Point (oC)

CH4 10 -161

SiH4 18 -112

GeH4 36 -90

SnH4 54 -52

Comparing Larger Compounds• When comparing non-polar compounds, the forces of

attraction are greater between molecules with the greatest number of atoms.

• There are more locations for London (van der Waals) forces to occur between adjacent molecules.

Boiling Points of Hydrocarbons

MolecularFormula

Boiling Point (oC)

State at STP

CH4 -161.5 gas

C2H6 -88.6 gas

C3H8 -42.1  gas

C4H10 -0.5  gas

C5H12 36.1 liquid

C6H14 68.7 liquid

     liquid

C10H22 174.1  liquid

     liquid

C22H46 327 solid

3) Polar Molecular Compounds• Compounds with bond dipoles AND molecular dipoles

(e.g. HCl, H2S, CF2H2) have higher boiling points.

• This is due to intermolecular forces between permanent dipoles.

• These are called dipole-dipole forces or bonds.

Dipole-Dipole Force (Bond)

Boiling Points of Some Polar and Nonpolar Substances

Substance

  Boiling Point (oC)

Molar Mass (g/mol)

Number of Electrons

HClpolar

(molecular dipole)

-84.9 36 18

H2S -60.7 34 18

F2 nonpolar(NO molecular

dipole)

-188.1 38 18

Ar -185.7 40 18

4) Polar Molecules: Hydrogen Bonding

• If hydrogen is bonded to a VERY electronegative atom (F, O or N), a very strong dipole forms.

• These atoms are also very small, concentrating this positive and negative charge.

• Dipole-dipole bonds between molecules containing O-H, N-H or H-F bonds form “hydrogen bonds”.

Water “bends” near a charged object.

Properties of Hydrogen-bonded Molecules

• A hydrogen bond is about 10x weaker than a covalent bond BUT 10x stronger than a normal dipole-dipole bond

• Thus H-bonded molecules have the highest mp/bp of the molecular compounds:

mp (oC) bp (oC)

• Propane (C3H8) -188 -42

• Propanol (C3H7OH) -126 97

• Glycerol (C3H6(OH)3) 18 290

General Properties of Molecular Compounds

• Molecular compounds do not conduct electricity sine their electrons cannot move between molecules.

• They have relatively low bp/mp due to the existence of weaker intermolecular forces

• As the strength of these intermolecular forces increase, so does the mp and bp.

Unusual Properties of Water• Water is called the “universal solvent” since it dissolves

both polar molecules (e.g. sugar) and ionic compounds (e.g. NaCl).

• Water expands when it freezes due to the organization of the many hydrogen bonds in the solid.

5) Metallic Solids

General Properties:• Few valence electrons• Low ionization energies• Malleable, ductile and shiny• Moderate mp/bp• Good conductors of heat and electricity in the solid and

liquid states.

Metallic Bonding• Metal properties can be explained by considering them as

postivie ions in an “electron sea” or “electron cloud”• Delocalized or conduction electrons are shared among

multiple cations are free to move throughout a crustal of positive ions.

The Electron Sea Model of a Metallic Crystal

Positive Metal Ion

Delocalized Electron “Cloud”

Explaining Metallic PropertiesProperty Explanation

Conductivity (Electricity / Heat)

Delocalized electrons can move between ions.

Ductility and Malleability

The plane of ions can move by distorting the electron cloud.

Lustre Reflection is caused by loosely bonded electrons absorbing and remitting all wavelengths of light.

• e.g. 1 Lithium is far more malleable than aluminum. Propose an explanation for this observation using the model of metallic bonding.

Metallic bonding occurs since the loosely held (delocalized) electrons are mutually shared by a crystal of positive ions. Since Li has only 1 delocalized valence electron compared with aluminum which has 3 and aluminum has a greater nuclear charge, we can deduce that the additional protons & electrons strengthen the metallic bonding and make it more difficult to displace the network of atoms in the crystal.

• e.g. 2 Which element would require the most energy to undergo vapourization, K or Sc? Explain.

• Scandium. • The stronger the metallic bonding, the more energy

required to change state. Similar explanation as above…..scandium has more delocalized electrons.

Kl (l) + 77 kJ K (g)

Sc (l) + 333 kJ Sc (g)

6) Ionic Solids• Solids formed by ionic bonds between metal cations (+)

and non-metallic anions (-).• Bonded together by a 3D array or crystal lattice without

distinct molecules.

Properties of Ionic Solids• High melting points and boiling points (many ionic bonds

that must be broken to change states).• Hard but brittle.• Many are soluble in water.• DO NOT conduct electricity in the solid state since ions

cannot move.• DO conduct electricity in the liquid or aqueous states

since charged ions are mobile.

Crystal Packing• Properties of ionic compounds are related to the packing

of the crystals:

Factors Affecting the Strength of Ionic Bonding

1. Ionic Radius of the Cation and Anion: As the radius of the ions increases, the attraction between oppositely charged ions decreases.

Factors Affecting the Strength of Ionic Bonding

2. Ionic Charge: As the charge of the cation and anion increases, the attraction increases.

Melting Point (oC)

Solubility(g/100g H2O

@ 0oC)  Melting Point

(oC)

Solubility(g/100g H2O

@ 0oC)

CsCl   MgO

NaCl   NaCl

646 161

800 35.7

2800 0.0006

800 35.7

7. Covalent Network Solids• Form a lattice of continuous covalent / polar-covalent

bonds.• Do not contain molecules.• Very hard, brittle substances.• Most do not conduct since electrons are either in sigma

bonds or lone pairs (filled orbitals). • Some exist as different allotropes (forms with different

properties)

Quartz: A Common Network Solid

• Quartz (SiO2) and Feldspars (KAlSi3O8 , NaAlSi3O8 & CaAl2Si2O8) make up most of the Earth`s crust.

• Quartz is a continuous framework of tetrahedral SiO4

Comparing CO2 and SiO2

Property Carbon dioxide (CO2) Quartz (SiO2)

Type of Solid Non-polar Molecular Covalent Network

Melting Point (oC) -78(sublimates at 1 atm) 1650

Boiling Point oC) N/A 2230

Bond angle (o) 180o 109o

Geometry Linear (sp) Tetrahedral (sp3)

Intramolecular bondType(s)

Polar covalent

Polar covalentIntermolecular bond Types

London forces

Allotropes of Carbon: Diamond• Covalent network of sp3 hybridized carbon (tetrahedral).• Very hard; very high sublimation point (3642 oC)• Does not conduct electricity.

Allotropes of Carbon: Graphite

• Network of sp2 hybridized carbon (trigonal planar)

• Half-filled p-orbitals form pi bonds• Graphite conducts electricity along the plane of the layers

due to the network of delocalized p-orbital electrons/• Graphite is a good lubricant since the planes can slip over

each other.

Summary: Types of Bonds

Intramolecular Metallic Intermolecular

1. Ionic bonds2. Covalent (Polar and Non-polar) 

Metallic bonds

1. London forces2. Dipole-dipole forces3. Hydrogen bonds

strong bonds weak bonds

increasing bond strength

Types of Solids formed by Elements

Atomic Solids

Non-polar Molecular Solids

H2 He

Li Be B C N2 O2 F2 Ne

Na Mg Al Si P8 S4 Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br2 Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I2 Xe

Cs Ba La-Lu

Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At2 Rn

Fr Ra Ac-Lw

Metallic Solids

Network Covalent Solids

Summary: Types of SolidsType Examples Intramolecular

BondsIntermolecular

BondsRelative Melting

Point

Atomic 

He, Ar van der Waals (London forces) very low

Molecular 

Cl2, HCl, H2O(non-metals)

covalent bonds (polar or non-polar)

van der Waals,dipole-dipole and hydrogen bonds 

low

Metallic Cu, Mg, Fe(metals) 

 metallic bonds

moderate-high

Ionic NaClNaNO3

(metal + non-metal)

 ionic bonds

high

Network quartz (SiO2)diamond (C)Silicon (Si)(non-metals) 

 covalent bonds

very high