162 Zero to Genetic Engineering Hero - Chapter 6 - Processing Enzymes
Figure 6-10. Classic out-dated illustration of an atom. Elec-
trons (yellow) orbit the positively charged nucleus (red) like
moons orbiting a planet.
The classic illustration of an atom can be found in
Figure 6-10. In this illustration, negatively charged
electrons were thought to orbit the positively charged
nucleus of an atom in a single ring-like path like
the ring around the planet Saturn. Throughout the
20th century, however, our understanding of atoms
changed, and we now know that electrons do orbit the
nucleus, but rather than doing so in a single ring, they
make three-dimensional structures that look sort of
like different balloon shapes. Scientists call these
orbitals. Every electron orbiting an atom has a slightly
different but predictable orbital “balloon shape” due
to the size of the nucleus (number of protons and
neutrons) and the number of electrons. (Figure 6-11).
Because this section is only meant to provide a basic
introduction to atoms and enable you to start thinking
more deeply about chemical bonding and interac-
tions, Figure 6-11 shows you the simplest orbital
paths that electrons take around a nucleus. The most
simple orbital is called an “s-orbital” where up to
two electrons travel around the nucleus very quickly
creating a spherical balloon-like pattern (Figure
6-11 (left)). It is important to note that just like in a
balloon where the rubber material makes up only a
thin layer, the electron orbit is the same. The electron
path does not ll the entire sphere but instead orbits
it in a thin layer. In the second example in Figure 6-11
(right), called the “p-orbitals”, there are three differ-
ent p-orbital paths that up to three electron pairs, or
six electrons, can take. Each of the paths look simi-
lar, they are simply in different orientations around
the nucleus. They have what look like a “dumbbell”
Further electron orbitals called “d-orbitals” and
“f-orbitals” have yet further different and interesting
shapes, but all orbit around the nucleus of the atom.
Every one of the electrons orbiting an atom has a
particular orbit and energy.
In the later sections when we will talk about chemi-
cal reactions, know that the different atoms involved
in the chemical reactions give away, accept, or share
electrons from these clouds. Imagine in your minds-
eye these balloon-like electron orbitals forming,
disappearing, and overlapping around the nuclei of
atoms in molecules as chemical reactions happen.
Figure 6-11. 3D Balloon-like electron clouds called orbitals. Electrons orbit the nucleus very fast and create three-dimensional pat-
terns that look like balloons. S-orbitals are spherical, and p-orbitals look like dumbbells.
Nucleus (+)
Nucleus (+)
Electron Cloud (-)
Electron Cloud (-)
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163Zero to Genetic Engineering Hero - Chapter 6 - Processing Enzymes
Now that you have a basic idea of what an atom is, we
can discuss bonding. Bonds are the “joints” between
atoms. Strong bonds can hold atoms together to form
stable molecules, and weaker bonds can temporarily
hold atoms and molecules together. Molecules are two
or more atoms. There are many kinds of bonds. Here
is a high-level overview of what they are:
Covalent Bonding: These bonds are the strongest
bonds, generally having a bonding energy of ~200-
400 kJ/mol. Covalent bonds occur when an electron
from an atom is shared with another. More specif-
ically, when an electron in an orbital of one atom
overlaps with another atom’s electron orbital, the
electron can become shared. This sharing of elec-
trons creates an attraction between the atoms. For
example, in a carbon-carbon bond (a common bond
in living organisms - like in the sugar molecules in
Figure 3-18) you’ll see that the electron orbitals of one
carbon can overlap with another, which results in one
or more bonds (Figure 6-12).
Covalent bonds are abundant in biochemistry. These
bonds bring the hydrogen (H), carbon (C), nitrogen
(N), oxygen (O), phosphorous (P), and sulfur (S) atoms
together to form all of the biomolecules in cells.
Proteins, DNA, RNA, lipids, and sugars molecules are
made of CHOPNS and are all held together by covalent
bonds, which you can learn more about in Table 6-1.
What does “a bond energy of ~200-400 kJ/mol” mean?
The basic unit of energy is a joule (J), named after the
famous scientist James Prescott Joule who dened
the joule. The joule is used to measure the amount of
energy in the world around us, including in chemical
bonds. For example, as a general rule, the thermal
heat energy available at room temperature is 3,000
joules per mole or 3 kilojoules per mole (~3 kJ/mol),
which is much lower energy than the typical covalent
bond energy of 200-400 kJ/mol. In other words, all
of the matter around you that make up chairs, desk,
table, and even the gas that you’re breathing have a
heat energy of about 3 kJ/mol. What is a mole? Going
Deeper 6-5 will help you dive into this topic.
Figure 6-12. Orbitals in two carbon atoms can overlap to form
bonds. On the top there are two carbon atoms bound because
of the overlapping yellow dumbbell shaped orbitals - this re-
sults in a single bond between the carbon atoms. Under the
right conditions, the green “vertical” orbital can also overlap
to create a second bond (bottom).
Blades of a Fan and Electron Orbitals Going Deeper 6-4
It can be difcult to understand in your mind how one or two electrons can create a cloud. An analogy
that can help is to think about the blades of a fan.
You may have noticed that when looking at a fan when it is in the “off” position you very clearly see the
blades of the fan. Perhaps three blades that are very clear and distinct. In this analogy, electrons are like
the blades. The electron(s) are distinct particles. When you turn on the fan, the blades spin quickly, and
they appear to be a disk. It is not because the blades morph into a new form, but rather because the blades
are moving fast in a circle and they are perceived to be a disk. This is similar to an electron and an electron
cloud orbital. The electron moves very fast around the nucleus a spherical path (in s-orbitals) that appears
to be a spherical electron cloud!
To learn more about orbitals and orbital theory you can look to Khan Academy. If you’re interested in seeing
how two atomic orbitals can come together to form a molecular orbital, have a look at a Youtube video where
the instructor discusses what happens when two oxygen atoms come together to form a molecule of oxygen
gas that youre breathing right now! https://amino.bio/molecularorbital
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164 Zero to Genetic Engineering Hero - Chapter 6 - Processing Enzymes
Table 6-1. Enthalpies and lengths of different bonds
Bond Length (pm) Energy (kJ/mol) Bond Length (pm) Energy (kJ/mol)
H-H 74 436 C–N 142.1 305
H-C 106.8 413 C=N (double) 130 615
H-N 101.5 391 CN (triple) 116.1 891
H-O 97.5 467 C–O 140.1 358
C-C 150.6 347 C=O (double) 119.7 745
C=C (double) 133.5 614 CO (triple) 113.7 1072
CC (triple) 120.8 839 O–O 148 146
C-S 182 272 O=O (double) 120.8 498
Source: http://philschatz.com/chemistry-book/contents/m51056.html
What is a mole (mol)? Going Deeper 6-5
The next time you eat a chocolate bar, have a look at the nutritional information to see how much sugar
there is in it - you might be surprised! Lets assume for this discussion that it has 10 grams of sugar in it. Of
what use to scientists is “10 grams of sugar”? It turns out that simply knowing the mass of something does
not tell you very much. What is much more helpful is knowing how many sugar molecules there are. Look
back at Figures 6-8 and 6-9 and consider the enzyme reaction. Is it more useful to know that there are 10
substrate molecules per one protein enzyme or 0.00000001 g of substrate and 0.0000001 g of enzyme? In
chemistry, the number of molecules and ratios of molecules provides you with more information than just
the mass. This is in part because molecules are different “sizes” and so one gram of sugar does not have
the same number of molecules as one gram of salt. Inventing the mole was a way to calculate the number
of molecules of any substance from the mass so they could be compared and more accurate calculations
could be completed.
In the 1800s, scientists wanted a way to find the equivalent number of atoms/molecules in
different substances using their mass. Using chemistry, scientists determined a very close approx-
imation of the number of carbon atoms in 12 grams of pure carbon to be 6.022 x 10
atoms. That is
602,200,000,000,000,000,000,000 carbon atoms in 12 grams of pure carbon. They then decided that a
mole would be 6.022 x 10
atoms or molecules of any pure substance. With this knowledge, chemists then
determined how many grams of each chemical element is needed to get 6.022 x 10
atoms (one mole). This
became known as the atomic weights. If you look at the periodic table at the end of this book, you’ll see the
atomic weights of the elements. Hydrogen, for example, has an atomic weight of 1.008 g/mol. This means
that in 1.008 grams of hydrogen, there are 6.022 x 10
atoms, 1 mole, of hydrogen atoms.
You can also combine atomic weights to nd the molecular weights. Water (H
0) has one oxygen (15.999 g/
mol) and two hydrogens (1.008 g/mol x 2 = 2.016 g/mol). Add the atomic weights up, and you get a molecu-
lar weight of 18.015 g/mol. This means that 18.015 grams of water, 1 mole, has 6.022 x 10
water molecules!
Next time when you drink a 500 mL bottle of water, you can boast that you just drank:
500 mL water is 500 g water, divide by, 18.015 g/mol = 27.75 mol
27.75 mol water x 6.022 x 10
molecules per mol = 1.67 x 10
water molecules!
(That is 16,700,000,000,000,000,000,000,000 molecules of water in a 500 mL bottle!)
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165Zero to Genetic Engineering Hero - Chapter 6 - Processing Enzymes
Now that you have a better idea of what a mole is (Going
Deeper 6-5) lets get back to the kJ/mol of heat energy.
This means that a mole of substance around you, such
as one mole of oxygen gas in the room around you,
or one mole of molecules in the chair you’re sitting
on, has heat energy of around 3 kilojoules. However,
in one mole of atoms bound together via a covalent
bond, the bond energy holding the atoms together is
200-400 kilojoules (kJ/mol).
To break a bond, you need to add at least the energy
of that bond. For example, ~200-400 kJ/mol of energy
needs to be added to a molecule in order to over-
come and break the covalent bonds. Because normal
room temperature energy is ~3 kJ/mol, all bonds
with higher than 3 kJ/mol energy are stable at room
temperature because there is not enough thermal
energy (heat) to break them. This is, for example
why the graphite in your pencil stays stable - the
carbon-carbon bonds that make up the graphite have
a bond energy of around 347 kJ/mol (single bond) and
614 kJ/mol (double bond). In the coming sections,
you’ll soon learn about some other types of bonds
that are affected by normal room temperature energy
and are a reason why butter may melt if left out of
the fridge on a warm day and why the heatshock step
during a transformation makes the membranes of
cells more uid.
Because most chemical reactions that happen in cells
involve the creation or breaking of covalent bonds,
which have much higher bond energy than the 3 kJ/
mol energy available at room temperature, a catalyst
is needed to replace the need for such energy. As you
saw in the Going Deeper 3-7 on chemical reactions
of Chapter 3, a catalyst is a substance that lowers the
activation energy needed to cause chemical reactions
to happen. In the case of biology, the catalysts that
lower the activation energy to break covalent bonds,
are protein enzymes, such as the beta-galactosidase
and ATF1 protein enzymes you genetically engineered
the E. coli cells to produce. Rather than needing the
“height off the ground” 327 kJ/mol energy to break
a carbon-carbon bond, the enzyme helps create a
“tunnel” to make the reaction happen without the
excess energy.
Protein enzymes bind to substrate molecules, twist
them, bend them, and even bring substrate molecules
into close proximity to “force” chemical reactions to
happen. In some instances, the amino acids that make
up the protein enzyme have extra electrons that will
kick-start the chemical reaction by forming bonds
with the substrate. When a protein enzyme does this,
it changes the rules of the game and it lowers the
energy required to break and form bonds and make
the chemical reaction happen.
This is the magic of living systems. Without protein
enzymes, very few chemical reactions would happen
because covalent bonds are quite stable. Life would
not exist without enzymes. It is the thousands of
protein enzymes in the cell that catalyze specific
chemical reactions to happen and sustain life.
Have a look back at the various chemical diagrams
of the different macromolecules - nucleic acids,
lipids, sugars, and proteins. You’ll see that all of these
important molecules, made up of CHOPNS, are joined
together by covalent bonds. Thousands of molecules
are integral to life and are made up of an assortment
of covalently bonded CHOPNS atoms. All of the atoms
are joined together by the sharing of electrons in their
balloon-like orbits (orbitals).
Electron orbitals and valence shells Going Deeper 6-6
The outer most electron orbital in an atom is called a ‘valence shell’. Some atoms have their outer most
valence shells almost full or almost empty. To become more stable, atoms have natural propensity to be
full or empty and to do this, atoms can gain or lose electrons. Alkali metals such as lithium (Li), sodium
(Na), and potassium (K) all have one valence electron in their outer s-orbital. To become more stable, they
prefer to lose this to another atom. When they lose an electron, they get a positive charge. This is why you
typically see Li
, Na
, and K
, these are the most stable forms of those atoms.
Alkali metals are well known to lose their single valence electron to halides such as uorine (F), chlorine
(Cl), bromine (Br), and iodine (I). This is because the halides have an almost full valence p-orbital shell and
they would like to ll it up with one more electron to become more stable.
As an example of this, when sodium (Na) and chlorine (Cl) are combined, the sodium will spontaneously
transfer an electron to chlorine to become sodium (Na
) and chloride (Cl
). Now that the atoms have become
charged ions, they participate in ionic bonding. Search “valence shell” online to learn more.
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166 Zero to Genetic Engineering Hero - Chapter 6 - Processing Enzymes
Ionic Bonding: Electromagnetism, a well-understood
eld, describes how positive charges repel, negative
charges repel, but a positive and a negative charge
attract. In cells, there are many cases where positively
charged atoms called cations (‘cat-ions’), and nega-
tively charged atoms called anions (‘an-ions’) form
an interaction due to the rules of electromagnetism
(Figure 6-13). This is called an ionic bond. Ionic bonds
typically have an energy of ~30 kJ/mol - 10 times less
strong than covalent bonds. Ionic interactions are
long-range interactions, as compared to covalent
interactions that require two atoms to be so close
that their electron orbitals overlap and electrons
are shared. In other words, ionic bonds do not bring
atoms together to form molecules, rather they can be
the temporary bonds between atoms and molecules.
You learned about calcium in Chapter 4 when you
completed your transformation experiment. The
positively charged cationic metal calcium (Ca
is able to interact through ionic bonding with the
DNA plasmids and the outer cell membrane. This is
because the DNA has a strongly negatively charged
backbone partially comprised of phosphate (PO
You also learned about the 20 standard amino acids,
and how each of them has a unique side chain. If you
look back at Figure 3-30, you’ll see that some of the
amino acids have a positive charge, negative charge,
and some are uncharged. Ionic bonding is key in driv-
ing some of the interactions that cause the proteins to
fold up into larger 3D structures (Figure 3-28).
Ionic interactions are also important in causing
molecules to begin interacting with one another
from a distance. Recall that during the Four B’s, the
rst being bump, molecules will bump around until
they bind. Ionic interactions can pull two bumping
molecules together during the bumping phase. For
example, ionic bonding can help the substrate “key
nd the enzyme “lockfrom a distance. It is important
to note that unlike how electrons are actually shared
between atoms in covalent bonding, ionic bonds do
not result from electron sharing. Ionic bonding is
caused by the electromagnetic forces of positive or
negative charges of atoms and molecules attracting
or repulsing.
Hydrogen Bonding: This is the third strongest class
of bonds generally with 5 - 30 kJ/mol bond energy.
Hydrogen bonds are similar in nature to the ionic
bonds you just learned about. Hydrogen bonds are
weak ionic bonds created when a hydrogen atom (H)
that is covalently bound to an electronegative atom
such as a nitrogen (N), oxygen (O), or uorine (F), is
attracted to another negatively charged atom.
Electronegative atoms are those that have a strong
desire for electrons, and they have the ability to tug
on the electron s-orbital in a hydrogen atom that is
covalently attached to them. This results in the elec-
tron orbital cloud around the hydrogen atom being
slightly positively charged because the positively
charged nucleus is no longer fully shielded by its
s-orbital (Figure 6-14).
A good example is in a water molecule (Figure 6-14),
where there is one electronegative oxygen atom
covalently attached to two hydrogen atoms. Normally
the negatively charged s-orbital of the hydrogen fully
surrounds positively charged nucleus resulting in
hydrogen having no charge (Figure 6-14 A). However,
in the case of when hydrogen atoms are covalently
bonded to an electronegative oxygen (water), the
oxygen atom pulls the hydrogen’s electron toward
it (Figure 6-14 B). Thats right, oxygen is
reedy for
Figure 6-14. A. Hydrogen atom with s-orbital electron fully
orbiting the nucleus. B. Water molecule demonstrating elec-
tronegative oxygen tugging on hydrogen atoms’ s-orbitals
leading to a hydrogen bond opportunity because of the (+)
charge from the nucleus.
- charge electron orbital
+ charge nucleus
Figure 6-13. Positively charged and negatively charged ions
are attracted to one another by electromagnetic force. The
dashed line indicates an attraction interaction.
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