CDS CHAPTER 10 SUMMARY 1 10: Bonding in Metals and Metal-­‐Non-­‐Metal Salts Concept Development Studies in Chemistry, 2013 •
10.3 – Observation 1: Properties of Metals
o Important properties of metals:
§ Malleability – metals can be shaped, bent, pressed, flattened without breaking
§ Conductivity – when a piece of metal is bridged across an electric potential, electrons
flow from the negative electrode to the positive electrode, creating a current
§ Ductility – metals can be drawn into thin wires while maintaining strength
§ Appearance – metals have shiny finishes
o Ionization energies
§ The ionization energies of metals are significantly lower than those of non-metals
§ These ionization energies do not vary much from metal to metal
§ The relatively low ionization energies of metals is due to their collective position on the
“left” side of the periodic table, since core charge increases as one moves right across a
row of the periodic table
• However, this does not explain the lack of variation in the ionization energies of
metals
§ Electron configurations (table 10.2)
• For the transition metals (V to Cu), both 4s and 3d electrons are present and the
number of electrons in each orbital is very sensitive to the number of valence
electrons and the core charge of the specific element
• The outermost electron in each of these atoms is always a 4s electron, but 3d
electrons are added as the atomic number increases
o 3d electrons shield 4s electrons from the increasing nuclear charge
o Therefore, there is not much increase in the core charge or ionization
energy
o How do electron configurations affect the bonding of metal atoms to each other? Look at each
property
§ Electrical conductivity
• A current is the movement of electrons through metal from the negative end of an
electric field to the positive end
o This requires that the electrons be “loose” – not strongly attracted to the
nuclei
o Remember that the ionization energy of metal atoms is low – this is the
“looseness” of electrons that creates current
• Conclusion: valence electrons in metal atoms are not localized to individual
nuclei – they can move about many nuclei
§ Malleability and ductility
2 CDS CHAPTER 10 SUMMARY These properties imply that the bonding of the metal atoms to each other is not
affected when the atoms are rearranged – the bonds do not break
• This is consistent with the idea of “loose” or “free” valence electrons
§ Metal atoms are arranged in an “array”
• Non-valence electrons remain strongly attracted to each nucleus
• Valence electrons are free to move about the nuclei and core electrons
• This is known as the electron sea model
§ Shininess
• Remember that light is a form of electromagnetic energy that can be absorbed or
emitted by electrons when they move to higher or lower energy levels
• Since there are so many electrons in the electron sea model, there are an
enormous number of energy levels and differences between these levels
• This makes it easy for metal to absorb and reemit light that hits it
10.4 – Observation 2: Properties of Salts
o One of the types of compounds formed by combining metal and non-metal atoms is a salt
o Important properties:
§ Brittleness – salts cannot be reshaped like metals; they will break into pieces
§ Lack of ductility – cannot be made into a wire or thread
§ Nonconductivity – insulates against the movement of current
§ Dissolves in water – when dissolved, the resulting solution conducts electricity
o Salt solutions and electricity
§ A current can pass through a salt solution – this means there must be charged particles
dissolved in the solution that carry the charge
§ These particles are ions (Na+, Cl-)
§ This does not reveal whether NaCl itself has ions
o Melting NaCl
§ NaCl changes its phase from solid to liquid at 808°C – this shows there are strong forces
keeping the solid bonds together
§ The resulting liquid conducts electricity – therefore, the ions Na+ and Cl– should exist in
solid NaCl
§ Why doesn’t solid NaCl conduct electricity if it is made up of ions?
• Current is charge in motion – the ions must be able to move to carry electric
current, but they are fixed in place in NaCl
• Since opposite charges strongly attract each other (Coulomb’s Law), the bonds in
NaCl are due to the attraction of Na+ and Cl– ions – this is why they cannot move
to form an electric current
o Why does a solid like NaCl contain ions?
§ Na has a low ionization energy – Cl has a high ionization energy and a high electron
affinity – therefore, a lot of energy is released when an electron is added to Cl
§ The energy released when an electron is attached to Cl is not enough to ionize the Na
atom – so this does not explain the presence of ions
§ Analyzing ionization energy and electron affinity isn’t enough
•
•
CDS CHAPTER 10 SUMMARY 3 If we used the energy comparison above, the positive and negative ions that
would be formed from Na and Cl wouldn’t interact with each other
• In reality, these ions are strongly attracted to each other – by Coulomb’s Law, this
lowers the potential energy
§ An NaCl crystal consists of an array of positive ions surrounded by multiple negative
ions
• All of these opposite charges so close to one another lowers the energy – this is
known as the “lattice energy”
o Since the bonding in NaCl is different from covalent and metallic bonding, it is called “ionic
bonding”
o Not malleable and brittle
§ NaCl crystals cannot be deformed – it would shatter
§ This shows that the bonding in NaCl depends strongly on the arrangement of the atoms
• For ionic bonding to work, the positive and negative ions must be surrounded
by each other (remember the lattice)
• Rearranging the atoms would result in like charges being next to each other – they
would repel each other, by Coulomb’s Law, and the lattice would collapse
o Differences in lattice energies
§ Look to table 10.3 – trends:
• The largest lattice energy corresponds to the combination of the two smallest
ions
• Lattice energy decreases when either or both of the ions are larger in size
o Remember Coulomb’s Law – energy is inversely proportional to the
distance between two nuclei – smaller ions can be found closer to each
other than larger ones due to their smaller atomic radii
• The lattice energies for ions with multiple charges is much larger
o Again, Coulomb’s Law – energy is directly proportional to charge
10.5 – Observation 3: Properties and Bonding in Solid Carbon
o Not all solids are formed from ionic and metallic bonding – there is also the bonding of nonmetals (covalent bonding)
§ Example: diamond, formed from pure solid carbon
o Properties of diamond
§ Very hard – the hardest solid available in bulk – not malleable or brittle
§ Very high melting point
§ Most thermally conducting material
§ Does not conduct electricity
o Inferences about diamond
§ Since it isn’t brittle, it can’t be ionic
§ Since it doesn’t conduct electricity, the electrons aren’t delocalized – doesn’t follow the
electron sea model, so it isn’t metallic
§ Since it is hard, the bonding depends on the specific arrangement of the atoms
o Modeling the structure of diamond
•
•
4 CDS CHAPTER 10 SUMMARY §
•
Carbon atoms have a valence of 4 and 4 valence electrons which they share with other
atoms to form covalent bonds
• We can assume that all of the electrons in diamond are localized, since it isn’t
malleable – therefore, all of the bonding in diamond is covalent
§ If we bond each carbon atom to four other carbon atoms, we can build a “network” of
carbon atoms that could theoretically go on forever
§ In this model, all of the electrons are localized and precisely arranged – this accounts for
the hardness of diamond
10.6 – A Model for Predicting the type of Bonding: Electronegativity
o Three types of bonding:
§ Metallic – bonding valence electrons are delocalized in an “electron sea” which allow
for malleability, ductility and conductivity
§ Ionic – adjacted positive metal ions and negative non-metal ions are strongly attracted to
each other in an array, creating a hard, brittle solid that doesn’t conduct electricity
§ Covalent – each atom in a “covalent network” forms covalent bonds (electron pairs) with
other atoms in order to satisfy its valence (the octet rule) – forms a hard solid that isn’t
brittle, malleable or conductive
o Metal atoms bond with metal atoms in metallic bonding; metal atoms bond to non-metal atoms
in ionic bonding; non-metal atoms bond to non-metal atoms in covalent bonding
§ One major difference between metal and non-metal atoms: the electronegativity of nonmetals is high while the electronegativity of metals is low
o Electronegativity by type of bond:
§ Ionic – an atom with low electronegativity (metal) is combined with an atom with high
electronegativity (non-metal)
• Since their electronegativities are so different, the atoms are not likely to share
electrons
• Instead, the very electronegative atoms form negative ions and the weakly
electronegative atoms form positive ions – these oppositely charged ions attract
each other, forming ionic bonds
§ Metallic – all the atoms are metals, so they all have low electronegativity
§ Covalent – all the atoms are non-metals, so they all have high electronegativity
o Must consider both the differences in electronegativity between atoms in a compound as
well as the actual magnitude of the electronegativities
o Know the bond type triangle (figure 10.2)