New type of hydrogen bond discovered

An entirely new class of hydrogen bond that forms between a boron–hydrogen group and the aromatic, π-electron system of a benzene ring has been discovered. The non-classical B–Hπ bond can be seen in the gas phase locking together diborane and benzene with a strength comparable to the hydrogen bonds that hold water dimers together.

Dieter Cremer and Wenli Zou of the computational and theoretical chemistry group at Southern Methodist University, Dallas in the US, worked with coordination chemists from Nanjing University in China to investigate the theory of non-classical hydrogen bonds that might form between a B–H group and organic structures and to demonstrate one such system experimentally.

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Covalent Bonding!

Before I even begin talking about what Covalent bonds are I want to clear up some possible confusion. If you’re learning general chemistry now you’ll thank me later.

The last two graphics above are typically what are used to illustrate what a covalent bond is between atoms. However this picture is not quite accurate. In reality atoms don’t align themselves in a Bohr model form of atoms and electron orbiting around like planets and moons.

Reality it turns out is just a bit weirder than that. Because quantum mechanics is really weird. The reality is that what exists are orbitals, which are a probability distribution of where an electron may be around an atom. These take on particular shapes depending on the element (heavier elements have more and different orbitals). This is because according to Heisenberg’s Uncertainty Principle, it is not possible to know both where an electron is and its exact momentum. Because of this we can only rely on probabilities on where electrons may be.

The first image is a closer representation of the way electrons are distributed than the circular models used in Basic/General Chemistry.

But this isn’t about orbitals. This is about Covalent bonds, so lets begin.

Covalent bonds are essentially what takes place when the valence electrons (the outermost electrons in an atom) are shared between atoms forming a bond. These occur between non-metals. They can form single, double or triple bonds depending on the amount of pairs of electrons that are shared. In more detail, these bonds consist of sigma and pi bonds, but I won’t go into those today.

These covalent bonds lead to stable molecules if they share electrons in such a way to create a noble gas configuration (where the electrons amounts are the same as those for the noble gas elements) for each atom. Essentially, when the ns2 and np6 orbitals are filled, this leads to stable orientation for the molecules. This is what is known as the Octet Rule, as it involves filling 8 electrons up in the mentioned ns2 and np6 orbitals.

There are different types of Covalent Bonds too.

Nonpolar covalent bonds occur between identical, non-metal atoms. These are the simplest types of covalent bonds. Because the atoms are the same, the electrons are shared equally between the atoms.

Examples of Nonpolar covalent bonds are: Oxygen (O2), Iodine (I2) and Fluorine (F2).

Polar Covalent Bonds are covalent bonds in which the shared electrons are not shared equally. Different atoms pull differently on the shared electrons, which will lead to an increased electron density to one side of the molecule rather than the other. The atom that pulls more strongly on the electrons becomes slightly more negative, and the other slightly more positive. This positive and negative end of a molecule is known as a dipole (like a magnet)  How much this occurs is related in the numerical value of Electronegativity. Electronegativity in simpler terms is an atoms affinity for attracting electrons to itself.

Examples of Polar Covalent bonds are: Water (H2O), Ammonia (NH3) and Hydrogen Fluoride (HF).

Whether a molecule is polar or nonpolar covalent can be determined by the electronegativity difference of the atoms that compose it. From 0.0-0.2 is nonpolar covalent. From 0.3-1.4 is Polar Covalent. From >1.5 is Ionic (but I won’t talk about that now). Covalent bonds are stronger than Hydrogen bonds, but are in general weaker than Ionic bonds.

If I’ve made any errors please let me know and I’ll correct them ASAP!

Sources: [1][2][3]

Other articles by me:

Hydrogen Bonding

Whales are Really, Really Fascinating


Note: As written, this activity requires that students hold hands. Younger students may not have any problems with this, however, the self-consciousness of adolescents may hinder the spontaneous movement and physical contact required for this activity. If you think this will be problematic in your classroom, cut 12-inch lengths of string for the students to hold to make the ‘bonds.’

*is miraculously born on one of the few planets out of 100 billion in our galaxy that are capable of supporting life through a delicate balance of heat, size, and chemical bonding*
ahh this is nice , quite nice
*sees an insect*
why hath god condemned me to the Hellplanet of hither Universe


Ionic Bonds

Ionic Bonds are a type of chemical bond that typically forms between metal and non-metal atoms. I say typical because it is also possible between certain non-metal molecular ions such as Ammonia (NH3). Unlike Covalent bonds, the electrons in Ionic bonds are not shared equally. They are instead transferred between species. Ionic bonds are formed by the electrostatic attraction between oppositely charged ions.

The images above show beautifully, that when a metallic species gives up its electron(s) it becomes more positively charged (what we call a cation), and the atom receiving the extra electron(s) becomes negative (what we call an anion). It is these positive and negative charges on these ions that form ionic bonds. For the most part these compounds follow the octet rule of filling electron shells to became stable, but this is a general rule, not absolute. There are many exceptions (particularly with metals).

Properties of Ionic bonded compounds are that they conduct electricity while molten or in a solution, they generally have a high melting point and tend to be soluble in water. In the real world, ionic compounds exist not as pairs of bonded atoms, but in arrangements of ions set in very specific shapes and forms. What we call a crystal lattice. This specific form repeated millions of times gives ionic compounds their 3 dimensional shape (salt crystals for example). Ionic bonds are the strongest of all bonds (Ionic > Covalent > Hydrogen).

Pure ionic bonds do not exist in nature. They are all covalent to some degree. The greater the electronegativity difference between atoms, the more ionic the bond. For reference again:

>1.7 EN difference = Ionic bond

0.4-1.7 EN difference = Polar Covalent

< 0.4 EN difference = Covalent

The example of Sodium Chloride (NaCl), or table salt is a good one to explain ionic bonds. When Na and Cl are combined, the sodium atoms lose an electron becoming cations (Na+), while the Chlorine atoms each gain an electron to form anions (Cl-). These oppositely charged ions are then attracted to each other in a 1:1 ratio to form Sodium Chloride (NaCl). Depending on the charges on the atoms, this ratio can vary. The compound Potassium Oxide (K2O) for example, the potassium exists in a 2:1 relationship with Oxygen because Oxygen’s charge is -2; so two K atoms (charge of +1) are needed per O atom to balance the charges. The greater the charge, the greater the attraction, and the stronger the ionic bond.

To remove electrons from atoms in this process requires the input of energy or heat, and is therefore endothermic. However the formation of ionic compounds is exothermic and creates heat or energy as a byproduct. Ionic bonding however will only occur if the reaction is what we call favourable. Essentially the bonded atoms must have a lower energy requirement than the free atoms. The larger the energy change, the stronger the ionic bond. Because metals have low electronegativity, and non-metals have high electronegativity, the resulting energy change of the reaction is more favourable when metals lose electrons and non-metals gain them.

This can be explained easier by reviewing the 3 steps in ionic bonding, in this case again with NaCl:

Please note that this is a simplification that does not take into effect a solid product forming into a crystal lattice (as occurs in reality) or beginning with gaseous Cl (Cl2). See here for a more detailed picture.

Ionic bonding can be broken down into 3 steps: ex. sodium chloride

Formation of the sodium ions (ionization)

Na(g) + energy —–> Na+(g) (DH = +496 kJ/mol)

Requires 496kj/mol input to remove electrons.

Formation of the chloride ions (electron affinity)

Cl(g) + e- —–> Cl-(g) + energy (DH = -348 kJ/mol)

Releases 348kJ/mol of heat/energy in this reaction.

Formation of the ion pair (bond energy) 

Na+(g) + Cl-(g) —–> NaCl(g) + energy (DH = -504 kJ/mol)

Releases 504kJ/mol of heat/energy in the final formation of NaCl.

DH = ΔH = Change in Enthalpy (heat energy).

We use Hess’s Law, which simply put means that the total energy change in a reaction is equal to the sum of the energy changes that occur.

So adding up each reaction we get:

DHrxn = DH1+ DH2 + DH3 = -356 kJ/mol

Sources: [1][2][3][4][5]

Read my other science writing posts here.


In astronomy and physical cosmology, the metallicity of an object is the proportion of its matter made up of chemical elements other than hydrogen and helium. Because stars, which comprise most of the visible matter in the universe, are composed mostly of hydrogen and helium, astronomers use for convenience the blanket term “metal” to describe all other elements collectively. Thus, a nebula rich in carbon, nitrogen, oxygen, and neon would be “metal-rich” in astrophysical terms even though those elements are non-metals in chemistry. This term should not be confused with the usual definition of “metal”; metallic bonds are impossible within stars, and the very strongest chemical bonds are only possible in the outer layers of cool K and M stars. Earth-like chemistry therefore has little or no relevance in stellar interiors.

The metallicity of an astronomical object may provide an indication of its age. When the universe first formed, according to the Big Bang theory, it consisted almost entirely of hydrogen which, through primordial nucleosynthesis, created a sizeable proportion of helium and only trace amounts of lithium and beryllium and no heavier elements. Therefore, older stars have lower metallicities than younger stars such as our Sun.

Image credit: NASA, ESA, and H. Richer (University of British Columbia)

Dash is dead. Need new blogs to follow.

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Doctor Who


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James Bond

Now You See Me

All Time Low

We The Kings

Fall Out Boy

Green Day

My Chemical Romance

Neck Deep

Mayday Parade

Arctic Monkeys

Panic! At the Disco

Let me tell you about jasmines.
The Persian meaning of a jasmine is “a gift from God”;
a pure, white gift from God
whose scent slithers in the wind
and curls up in your heart.
The jasmine is the cousin you envy,
the one who got away from the olive tree
and grew beautiful, exotic, different.
I once asked my chemistry teacher about jasmines.
He looked at me bluntly and threw words
like carbonyl, acid, compound.
I sat perplexed,
wondering how one could crush beauty
and reduce it to numbers and chemical bonds.
I wondered if someone somewhere knew
they were bottling the scent of my home.

Let me tell you
about the jasmines of my home as I remember them.
I was born in Homs.
My mother was rushed past a jasmine tree
before I entered this world.
She stopped and reached out
to pluck that orb of white heaven
and gripped it while I was led here, crying and bare.
She held my hand and the jasmine she had plucked
pushed into the fine lines of my palms –
engrossing its scent into my skin.

My grandmother grew jasmines.
They wound up the walls of her home
and although those walls became cracked and grey,
the jasmines stitched them up like fine lace.
She would place one in my hair.
I felt radiant and until this day,
I can see her frail smile as she admired her artwork.
The stems of those jasmines have never left me
and I feel them, like angels
whispering on the tips of my ears.

(Day 679 of the besiegement of Homs)

Now, let me tell you
about the jasmines of my home as they are.
I have not seen my jasmines for a long time.
I have not felt their soothing glances
from behind the trees;
They used to grow stretched in the sun,
across the walls of Homs.
My jasmines are rotting, they are bled.
My jasmines are brown.
My jasmines cannot run away from murder,
tyranny, missiles,
the trod of fleeing feet.
They are there to starve as my people starve.
There to be crushed under the vultures
who used the jasmine trees
to pick my people out of their teeth.
We are not people
but statistics to the world;
The jasmines have fallen with us
and we have caused as much sound as they.

But my jasmines are what comforted me out of the womb.
And they are what keep me fighting,
keep me digging,
until all I see under the rubble of Homs
is the sunshine petals of my jasmines.
The jasmines of Homs are bleeding;
yet I continue to smell them…
I hold them with me wherever I go.

—  “The Jasmines of Homs”, Tala E.