Bonding. Year 11 DP Chemistry. What is a bond?. A chemical bond is a force that holds atoms together making a new substance. Ionic Bonds result from electrostatic attraction between oppositely charged ions
Year 11 DP Chemistry
A chemical bond is a force that holds atoms together making a new substance
Ionic Bond: As a rule of thumb, we say that the difference between the electronegativity values needs to be high (i.e. greater than 1.7) to be ionic. They form between cations on the left and anions on the right of the Periodic Table.
Covalent Bond: If the difference between the electronegativity values of two highly electronegative atoms is low, a covalent bond is formed. They tend to form between non-metals, but sometimes metals are involved (eg Al2Cl6)
Metallic Bond: If the difference between the electronegativity values of two highly electropositive atoms is low, a metallic bond is formed. These form between metals of the same or different type of atom
The relative tendency of an atom to attract bonding electrons to itself on the Pauling Scale
Group 5 – gains 3 e- to gain a full valence shell
Group 6 – gains 2 e- to gain a full outer shell
Group 7 - ??
For example – Fe forms two ions
Can you deduce which two ions and why?
Fe(II) – losing two 4s electrons
F(III) – losing two 4s e- and one 3d e- to give a half-filled d
Oppositely charged ions are formed by electron transfer due to a large electronegativity difference(> 1.7 difference)
Na has a low electronegativity relative to Cl, so ions are formed by a transfer of an electron to achieve a full valence shell for both atoms. These oppositely charged ions then form a bond.
This shows a model of a NaCl lattice with alternating positive and negative ions
Li+ F- LiF Mg2+Cl- MgCl2
Some ions contain more than one element and the charge on the ion is spread (delocalised) over the entire ion. They have specific names and act as a single unit.
Important ones to know:
HCO3- (bicarbonate or hydrogen carbonate)
Ionic compounds form in the same way with polyatomic ions. Here we see Na2SO4. Notice the sulfate did not change formula.
Electrons are shared between two atoms. These atoms are most commonly non-metals.
The more shared pairs the stronger the bond and the shorter the bond. Note the trend in C-C bonds.
Note: bond energy is the amount of energy required to break a bond
In this type of covalent bond, the difference is that one of the atoms in the pair donates both of the electrons in the bond.
Examples: CO, NH4+, H3O+
The dots represent the valence electrons for each element
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
Here, the octet rule is satisfied for Cl, but is irrelevant for H which can only hold 2 e-.
Two pairs of shared electrons
Draw the Lewis structure for Acetylene (ethyne) – C2H2
Notice the double bond in O2. No other configuration will satisfy the octet rule. Why is H4O not formed?
Notice that the bond lengths would be different in benzene unless there is a resonance structure like the one represented above
All octets (ignoring H) have been satisfied.
N2 has a triple bond. Using dot diagrams, show why a single or double bond is incorrect.
Illustrate how CO and H3O+ contain dative bonds.
Draw a Lewis diagram for HCN
Draw the two resonance structure of ozone O3
Use a drawing to show how many resonance structures are possible for the nitrate ion (NO3-)
Compare the bond lengths and strengths of the two C,O bonds in the carboxyl group below.
There are 3 ways the octet rule breaks down:
NO (nitrous oxide)
This mostly occurs with H, B and Be
BF3 (boron trifluoride)
Octets to outer atoms
Extra e- (24-24=0) to B
Therefore no extra electrons to add. B has only 6 e- (< octet). What about double bonds??? see next…
This would give 3 resonance structures. What would they look like?
Add a double bond to BF3 for a possible octet…
The above structure would lead to a δ+ on F and a δ- on B. Is this likely considering the electronegativites?
Because B has only 6 valence electrons, BF3 reacts strongly with compounds that have unshared pairs of electrons
Starting in period 3, expanded valence shells are possible.
This is the most likely exception to the octet rule.
The octet rule is based upon valence orbitals containing an s and p orbital. This gives 2 + 6 = 8 e- (an octet). In the third shell (n=3) d orbitals become available. P is below. A 3s can be exited to the 3d, which allows for 5 valence shell bonding electrons.
Promote one e-
Sulfur (also in group III) can expand it’s octet to have more than 8 electrons as well. Sulfur can form SF2, SF4, SF6
Go to this website to see an animation on expanded octets in sulfur:
Other notable expanded octets…
PF6- (12 e-)
Some bonds are not purely ionic, but they still have a significant difference in electronegativity that leads to one atom pulling the electrons more strongly than the other.
This electronegativity difference leads to a partially positive end δ+ of a molecule and a partially negative end δ- (note: δ is the Greek letter delta and means partial)
F is more electronegative, so the electrons spend more time around the F nucleus
Small or no difference in electronegativity values leads to non-polar substances.
Cl is a highly electronegative element, but there is no difference when one Cl atom bonds to another Cl atom
C and H have very little difference in electronegativity, so methane is non-polar
Some compounds contain polar bonds, but the polarity is cancelled out due to the structure.
O is more electronegative than C meaning each bond is polar towards the O atom, but due to its linear shape, these polarities cancel each other resulting in a non-polar molecule.
To determine the shape of covalent molecules, we use the Valence Shell Electron Pair Repulsion Theory (VSEPR) which states:
“The geometric arrangement of atoms around a central atom is determined by the repulsion between electron pairs in the valence shell of the central atom.”
Stay away! I am repulsed by you
In other words…
So, VSEPR theory says that molecular geometry is determined by the shape that keeps e- pairs as far apart as possible
We know C forms four bonds and O forms 2 bonds
What arrangement will allow the valence e- around the central atom (C) to be as far apart as possible?
The Lewis structure looks like this:
Linear molecules have two areas of high electron density around the central atom.
Other examples of linear molecules :
Ethyne (acetylene) C2H2
Molecular chlorine (Cl2)
3 e- pairs around the central atom leads to a trigonal planar shape as in BF3
Trigonal planar - Angles are1200
If one of those pairs is a non-bonding or lone pair of electrons, the shape is described as bent or v-shaped as in SO2
Bent – angles are less than 1200 due to lone pairs taking up more space than bonding pairs. These angles are 1170
However, in 3-D space, it is possible to allow the electrons to be further apart using a tetrahedral shape with bond angles of 109.50
If we look at the Lewis structure for CCl4, we might assume a flat structure with 900 bond angles
(0 lone pairs)
(1 lone pair)
Angles = 900 & 1200
2 lone pairs
3 lone pairs
Octahedral (0 lone pairs)
(1 lone pair)
(2 lone pairs)
Justify the shape of XeF4. Why are the lone pairs at 1800?What other two shapes are possible with 6 pairs?
Each carbon atom in graphite is bonded to 3 other carbon atoms forming flat sheets of carbon rings. These layers are loosely bonded to each other making graphite soft
In diamond, each carbon is bonded to 4 other carbons in a giant repeating lattice. This lattice is non-polar and very strong, making diamond the hardest mineral on Earth. It’s m.p. is over 35000C!
The fullerene contains 60 carbons arranged like a soccer ball with alternating 5 and 6-member rings.
Carbon bonded to three other carbon atoms leaves one valence electron per carbon atom. These electrons are delocalised allowing graphite to conduct electricity
The individual layers contain strong covalent bonds, but are only loosely bonded to other layers. This allows them to easily slide over one another making graphite useful in pencils and as a solid lubricant
The third allotrope of carbon was discovered in 1985 and includes some weird and wonderful shapes. The first was known as a “Bucky Ball” which has a structure like a soccer ball and contains 60 carbons- C60
There are other fullerenes that contain more or less than 60 carbons. These have been detected in space leading scientists to suggest that they could be the origin of life in the universe
Nanotubes are another area of current research by materials scientists. These tubes have high tensile strength and are able to conduct electricity. These have potential medical applications by attaching to resistant bacteria or cancer cells. They are also being researched for the proposed “space elevator” to be used as cables.
Silicon dioxide (SiO2), which is the chemical component of sand. has a structure similar to diamond with a repeating giant lattice of tetrahedral shapes. This makes it very hard. However, it has polar bonds allowing it to be dissolved slowly by some alkaline solutions and it also has a much lower m.p. than diamond (17700C)
Silicon atoms contain 4 valence electrons and form regular repeating covalent bonds. It s chemically similar to C. The picture below show it’s crystalline form
This shows a computer generated model of the complicated giant lattice of silicon dioxide.
Metallic bonding can be described as
A repeating lattice of positive metal ions in a sea of delocalised electrons
Delocalised electrons are not attached to any particular metallic nucleus and are free to move about the lattice.
This is due to the electrons being able to move freely through the lattice.
This means the electrons can act as charge carriers for conducting electricity and energy carriers for conducting heat
Electrical Conductivity of Metals
Malleability and Ductility
The delocalised electrons in the 'sea' of electrons in the metallic bond, enable the metal atoms to roll over each other when a stress is applied.
Source of animations and diagrams:
www.google.com (images search)
Sundin, C.; (n.d.); found at http://www.uwplatt.edu/ ; accessed 04/2011
Dynamic Science; (n.d.); found at http://www.dynamicscience.com.au/tester/solutions/chemistry/chemistry%20index.htm ; accessed 04/2011
Courtland, R.; New Scientist (online); 28/10/2010; found at http://www.newscientist.com/blogs/shortsharpscience/2010/10/buckyballs-abound-in-space.html ; accessed 04/2011