Solids
CHEM HONORS
Solids: Crystalline vs. Amorphous
a) Amorphous - Amorphous
solids have irregular or
curved surfaces, do not give
well-resolved x-ray
diffraction patterns, and
melt over a wide range of
temperatures.
b) Crystalline – Crystalline
solids have well-defined
edges and faces, diffract xrays, and tend to have sharp
melting points.
Types of Solids
Ionic Solids (Ionic Bonds)
• Molecular Solids (Covalent Bonds)
• Atomic Solids
• Metallic Solids (Metallic Bonds)
•
Ionic Bonds
• Ionic bonds commonly contain a metal and a non-metal
• The metal is oxidized and loses an electron – becomes a CATION (positive charge)
• The non-metal is reduced and gains an electron – becomes an ANION (negative charge)
• An electrostatic attraction holds the two ions together
Ionic Bonds – Types of Ions
Charge of ion predicted using the electronegativity and electron configuration of the atom
- Atoms with loosely held electrons form positive ions (Cations)
- Atoms able to hold additional electrons form negative ions (Anions)
Ex) Group I atoms loose electrons, become positive (by losing an electron they fulfill the octet
rule)
Ionic Solids
• Contains ions at the point of the lattice that
describes the structure of the solid (ex NaCl)
• VERY high MP’s (500-3000C⁰)
• Very Hard
• Ion-ion Coulombic forces are the strongest of all
attractive forces
• IMF usually implies covalently bonded substances,
but can apply to both types
• Very STABLE because they are held together by
strong electrostatic forces.
NaCl crystal lattice: Na –
purple, Cl - green
Ionic Solids – Recap
Strong attraction between the 2 ions due to the Coulombic force
Solids at room temp
High melting points
Huge, 3D networks of ions of opposite charge held together by ionic bonds
Questions
•Which has a stronger ionic bond, NaCl or NaF?
◦ NaF because F- is smaller; the particles/ions will pack tighter in the crystal lattice
•Melting point measures what?
◦ How strongly ions are held together in a bond; the more strongly the ions are held together, the higher
the MP
•Which has a higher MP? NaCl vs. NaBr; MgO vs. CaO
• NaCl and MgO because they’re smaller
• Smaller ions = greater bond, because the in Coulombic force equation distance and force are inversely
proportional:
Lattice Energy
•Lattice Energy: amount of energy needed to completely break apart the ions in one mole of an
ionic compound
•Which has a larger lattice energy: NaCl v. NaBr; NaCl v. MgS:
• NaCl because it has a higher melting point, and thus stronger ionic bonds, so it will take more energy to
break the bonds
• MgS because it has a higher charge (Mg2+ and S2- v. Na+ and Cl-), and thus stronger bonds (because in
Coulombic force equation force and charge are directly proportional), so it will take more energy to
break the bonds
Summary Ionic Bonds + Ionic Solids
•Ionic Bonds
• NM (gains electrons)+ Metal (loses electrons)
• Smaller ion = greater force = higher MP
• Greater charge = greater force = higher MP
•Ionic Solids
• Solids at room temp
• High melting points
• Huge, 3D networks of ions of opposite charge held together by ionic bonds
Molecular Solids
• discrete covalently bonded molecules at each of its lattice
points (water or sugar)
Molecular Solids
• Characterized by strong COVALENT bonding within the molecule yet weak forces between the
molecule
• It takes 6.0 kJ of energy to melt one mole of solid water since you have to overcome H-bonding
while it takes 470 kJ of energy to break one mole of O-H bonds
• Molecules such as CO2, I2, P4, and S8 have no dipole moment (London Dispersion Forces)
• As the size of the molecule increases the London Dispersion Forces increase because the larger
the molecule the more electrons, the more polarizable its electron cloud.
• If it is polarizable, temporary dipoles can easily form. This causes the melting point and boiling
point to increase due to the molecules becoming more attracted to one another
Molecular Solids
Carbon Dioxide
Iodine (I2)
Sulfur (S8)
Network Atomic Solids
• AKA Network Covalent
• Composed of strong directional covalent bonds that are best
viewed as a “giant molecule”.
• Examples include diamond and graphite (both are composed
of strictly carbon atoms)
• Allotropes are different structural modifications of an
element; the atoms of the element are bonded together in a
different manner.
Diamond
• Diamond - bonds with neighbors in a tetrahedral 3-D fashion
• Graphite – only has WEAK bonding in the 3rd dimension
Graphite
Network Atomic Solids
Network Solids are often:
• Brittle – diamond is the hardest substance on the planet, but when
a diamond is “cut” it is actually fractured to make the facets
• They DO NOT conduct heat or electricity
• They are carbon or silicon based
Diamond
Network atomic solid
• Diamond
is hard, colorless and
an insulator.
• It consists of carbon atoms ALL
bonded tetrahedrally.
• Therefore they are sp3
hybridization and 109.5o bond
angles
Graphite
Network Atomic Solid
• Graphite is slippery, black and a conductor.
• Graphite is bonded so that it forms layers of carbon atoms
arranged in fused six-membered ring.
• This indicates sp2 hybridization and 120o bond angles within
the fused rings.
• The unhybridized p orbitals are perpendicular to the layers
and form π bonds.
• The delocalized electrons in the π bonds account for the
electrical conductivity while also contributing to the
mechanical stability of the layers.
• It is often used as a lubricant in locks – grease or oil collects
dirt, graphite does not.
Silica
• Empirical Formula – SiO2
• Nothing like CO2 due to its bonding. SiO2 –
sand or glass; CO2 – clear colorless gas
• Silicon cannot use its valence 3p orbitals to
form strong π bonds with oxygen, mainly due
to the larger size of the silicon atom and its
orbitals – you get inefficient overlap
• INSTEAD of forming π bonds , the silicon atom
satisfies the octet rule by forming single σ
bonds with FOUR OXYGEN atoms.
Silica
Network Atomic Solid
• Each silicon is in the center of a tetrahedral
arrangement of oxygen atoms.
• The structure is based on a network of SiO4
tetrahedral with shared oxygen atoms.
• When silica is heated above its melting point
of about 1600oC and cooled rapidly, an
amorphous solid forms. We call it glass – it’s
really a super-cooled, ultra viscous liquid with
a great deal of disorder.
Structure and Bonding in Metals
• Metals are characterized by high thermal and electrical
conductivity, malleability, and ductility.
• Closest Packing – a model that uses hard spheres to
represent the atoms of a metal. These atoms are packed
together and bonded to each other equally in all
directions.
• It will be easiest for you to understand if you can imagine
taking a cubic box and pouring in golf balls.
• The balls will layer, perhaps directly on top of one
another, but perhaps one layer slides into the “dimple”
made by the first layer so that the two
Metals – Properties
•Have shine or luster
•Malleable – can be hammered or pressed into different shapes without breaking
•Ductile – can be drawn into thin sheets or wires without breaking
•Conduct heat and electricity
Nonmetals
•NOT malleable
•NOT ductile
•Poor conductors of heat and electricity
•High Electronegativity
•Generally, a gas
Bonding Models for Metals
Metals:
◦
◦
◦
◦
◦
Conduct Heat
Conduct Electricity
Are Malleable
Are Ductile
Have High Melting Points
Indicates that bonding in metals is both strong and non-directional
Difficult to separate atoms, but easy to move them provided they stay in contact with each other
Bonding Models in Metals
Electron Sea Model: A regular array of metals in a “sea” of electrons.
Electronic Structure of Metals
•Valence electrons on a metal atom are shared with many neighboring atoms, not just one
• Electrons not tightly bonded to individual atoms, so they move freely throughout the metal
• Force of attraction btwn (+) metal ions and the sea of mobile (–)electrons forms a metallic bond that
holds these particles together
• Metals hold electrons loosely due to low electronegativity
•Valence electrons are delocalized over a number of metal atoms
•Metals exist as extended arrays of packed spherical atoms so each atom can touch as many
neighboring atoms as possible
Bonding Models of Metals
Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs
formed from valence atomic orbitals of metal atoms
Metal Alloys
Metal Alloys – a substance that has a mixture of elements and has metallic properties
Substitution Alloys –
◦
◦
◦
◦
in brass 1/3 of the atoms in the host copper metal have been replaced by zinc atoms.
Sterling Silver – 93% silver and 7% copper
Pewter – 85% Tin, 7% copper, 6% bismuth, and 2% antimony
Plumber’s solder – 95% tin and 5% antimony
Review
• Molecular Solids
• Crystallizes
• Low MP
• Metallic Solids
• High MP
• Insoluble in H2O
• Used in Structural Material
• Ductile
• Malleable
• Not Crystalline
• Conducts electricity and heat
• Ionic Solids
• High MP
• Crystallizes
• Hard
• Brittle
• Soluble in H2O
• Atomic Network Solids
• Insoluble in H2O
• Crystallizes
• Doesn’t conduct
electricity or heat
• High MP
• Brittle
Bond-Type Triangle
•Bond-type triangle: chart used to predict properties of a compound based on the
electronegativities of the elements that compromise the compound
• Can be divided into regions which indicate the predominant type of bonding present in compounds
•Semiconductor: compound with properties intermediate btwn metallic and covalent
• Ex) Si
ΔEN =
(absolute)
difference in
electronegati
vity btwn 2
elements
͞ = avg.
EN
electronegati
vity of 2
elements
Checkpoint
Identify each solid as molecular, ionic, or atomic.
A.) Ar (s)
B.) H2O (s)
C.) K2O (s)
D.) Fe (s)
Checkpoint
Identify each solid as molecular, ionic, or atomic.
A.) CaCl2 (s)
B.) CO2 (s)
C.) Ni (s)
D.) I2 (s)
Checkpoint
Rank the solid by increasing melting points. Why?
Ar (s), CCl4 (s), LiCl (s), CH3OH (s)
Checkpoint
Rank the solids by increasing melting points. Why?
C (s, diamond), Kr (s), NaCl (s), H2O (s)
Checkpoint
Which solid in each pair has the higher melting point and why?
A.) TiO2 (s) or HOOH (s)
B.) CCl4 (s) or SiCl4 (s)
C.) Kr (s) or Xe (s)
D.) NaCl (s) or CaO (s)
Checkpoint
Which solid in each pair has the higher melting point and why?
A.) Fe (s) or CCl4 (s)
B.) KCl (s) or HCl (s)
C.) Ti (s) or Ne (s)
D.) H2O (s) or H2S (s)