Types of Bonding
The names Bond, James Bond
Electronic Nature of Chemical Bonding
• Electrostatic forces between charged particles
– Protons are positive
– Electrons are negative
– Neutrons are neutral
• All bonding is based on
– Opposite charges attract
– Like charges repel
– The greater the distance between 2 particles, the
weaker their bond
– The greater the charge between 2 particles, the
stronger their bond
Shells
• Open shell = a shell that contains less than the
maximum number of electrons
• Closed shell = A shell with the maximum
number of electrons (noble gases)
• Valence electrons = all electrons in the
outermost shell, not including d or f orbitals
– Open shell!
– S and P orbitals only!
Three rules
• Four rules about valence electrons in orbital
theory
– Each orbital holds up to 2 electrons
– Electrons repel so fill each orbital of the same
subshell first before pairing
– After all orbitbals are filled with 1, pair them up
– Fill from lowest energy and up
Valence of an Atom / Combining
Capacity
• It is different from valence electrons
• It is the number of UNPAIRED electrons
around an atom
• Going from left to right in the periodic table,
ignoring the transition metals
Ionization Energy
• The amount of energy required to remove an
electron from an atom to turn it into an ion
• Less energy required to remove electrons
from atoms with a few valence electrons
– Such as Na or Li
• Decreases as we go from top to bottom
• Increases as we go from left to right
• Helium has the highest, Francium the lowest
Types of chemical bonds
Ionic Bonds
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Metal and non-metal
Metal atom donates an e- to the non-metal atom
Metal atom becomes a cation
Non-metal atom becomes an anion
Forms what we call a crystal lattice (forms a
crystal)
• As ions are the exchanging of electrons, ionic
compounds are very conductive
Ionic Bonds
• Ionic compounds are usually high solubility
and dissolve into aqueous solutions
– More on this next unit
• Examples
– NaCl, KCl, NaF, MgCl2
Covalent Bonds
• Non-metal and non-metal
• When 2 atoms have an opened valence shell
• Non-metal atoms share 2 electrons to form a
bond and are paired
• Still must follow the octet rules
• Example
– H2, CO2, O2, CO
Covalent Bonds
• Can be split into 2 types depending on how
electrons are shared amongst the atoms
– Polar
– Non-polar
• Remember the term electronegativity?
• Some atoms will want more electrons than
others
• When this happens, a dipole is formed which
means an unequaled sharing of electrons
Metallic Bonds
• Metal and Metal
• Metals will form bonds too, think about big
pieces of iron and gold. Those are pure
substances!
Forming more than 1 bond
• Remember, molecules aren’t just always 2
atoms put together
• We can have more than 2 atoms
• If this is the case, we must remember
– Octet rule still applies for all the atoms present
– All valence electrons are involved
Electronegativity and Ionization Energy
• Remember, EN is how much an atom wants an
electron
• IE is how much energy we need to put in to
remove an electron
• Atoms with high EN will also have high IE!
– They want the electron so it takes a lot to remove
it!
• Like Mr. Chio with his McDonald’s McDouble.
Melting Points and Bonds
• Melting point is the point in which bonds are
breaking!
• Ionic bonds are strong! Hard to break bonds
– High melting point
• Examples
– NaF = 993°C, KCl = 770°C, LiCl = 605°C and LiF =
845°C
Melting Points and Bonds
• Covalent bonds are also strong, but have
varied melting point depending on their
intermolecular interactions
• Some such as pure carbon have HIGH melting
points, over 4000°C due to strong bonds.
• Some such as CH4 with weak intramolecular
forces have low melting points, -182°C due to
weak bonds
Predicting the formula for covalent
compounds
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Same as ionic, criss cross method
No charges needed
This is due to combining capacity
Example
– Boron and carbon
– Arsenic and sulphur
• You try
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Nitrogen and oxygen
Carbon and oxygen
Nitrogen and chlorine
Silicon and oxygen
London Forces
• Intramolecular forces = bonds holding atoms
together to form molecules = strong
• Intermolecular forces = bonds holding
molecules together = weak
– Van der waals forces
• We will only be dealing with one type of
intermolecular forces/van der waals force
called London Forces
London Forces
• Remember dipole?
• In a covalent molecule, some atoms want the electron
more than others and shifts it slightly towards them
– The atom that wants the electron becomes –ve
– The atom that loses the electron becomes +ve
• But this is temporary
• This creates a positive and negative end to the
molecule! But temporary
• With these positive and negative ends, we can have
molecules bonding +ve end to –ve end! But still
temporary as electrons always move!
Example
• H2O
London Forces
• These are what determines melting points of
covalent molecules due to low or no dipole
and thus no London Forces so low melting
points!
• Main Rule = the more electrons an atom has,
the higher their atomic number, the stronger
the London Forces
• London Forces are the weakest of all types of
bonding!
Homework
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Page 166 #42 and 44
Page 172 #57
Page 177 #68
Page 179 #72
Worksheet on bond type!