Loudon Ch. 1 Review: Chemical Structure & Bonds
Jacquie Richardson, CU Boulder – Last updated 9/2/2020
Organic chemistry is focused on carbon, and is used in making pharmaceuticals and
materials including polymers and liquid crystals. It’s also the basis for biochemistry. It uses
primarily the top three rows (a.k.a. periods) of the periodic table, plus a few elements from
lower periods. The rightmost column (a.k.a. group) is the noble gases (He, Ne, etc.), and the
one next to it (F, Cl, etc.) is the halogens.
Group 1 Group 2 … Group 13 Group 14 Group 15 Group 16 Group 17 Group 18
Period H (1) He (2)
1
Period Li (3) Be (4) B (5) C (6) N (7) O (8) F (9) Ne (10)
2
Period Na (11) Mg (12) Al (13) Si (14) P (15) S (16) Cl (17) Ar (18)
3
Period Br (35)
4
Period I (53)
5
The number next to each element is its atomic number, which describes the number of
protons in the atom. Each atom consists of a small, heavy nucleus made up of protons (with
positive charge) and neutrons (neutral, with no charge), surrounded by much lighter
electrons (negative charge). Protons determine which element an atom is, electrons
determine the overall charge on that atom, and neutrons don’t have much effect on chemical
behavior except in a few specific cases. For an atom to be electrically neutral, it must have
the same number of protons and electrons. For example, a neutral fluorine atom (F) has 9
protons and 9 electrons, plus some number of neutrons.
The electrons in an atom exist in shells around the nucleus. Each shell has a maximum
amount of electrons it can hold (2 for the first, 8 for subsequent shells…as far as organic
chemists are concerned most of the time). The shells fill up starting closer to the nucleus and
moving outwards. The electrons in filled shells are core electrons, and those in the unfilled
outermost shell are valence electrons. The core electrons of each atom have the same
configuration as the noble gas immediately before that atom in the table. For example,
lithium (Li) has the same core electron configuration as helium (He), and chlorine (Cl) has
the same core electron configuration as neon (Ne). The remaining electrons above the noble
gas configuration are the valence electrons – Li has 1 valence electron, and Cl has 7.
Chemical behavior is based primarily on the behavior of valence electrons, since these are
furthest from the nucleus and can interact with the “outside world” most easily. When an
atom becomes a charged ion by losing or gaining electrons, it tends to do this by going to
the same electron configuration as the nearest noble gas. For example, sodium (Na) has one
valence electron and is closest to Ne, so it will most likely drop a single electron to become a
sodium cation, Na+. Chlorine has 7 valence electrons and will most likely pick up one more
to become a chloride anion, Cl-. (Note that because electrons have negative charge, extra
electrons give something a – charge, not +. This is a very common mistake.)
Chemical Bonds
Atoms can bond together in several different ways. The two biggest subdivisions are
between ionic and chemical bonds.
1
, Loudon Ch. 1 Review: Chemical Structure & Bonds
Jacquie Richardson, CU Boulder – Last updated 9/2/2020
Ionic bonds involve electrostatic attraction between ions, each with a complete positive or
negative charge(s). The most common example is table salt, NaCl. It consists of alternating
Na+ and Cl- ions. Each ion is surrounded by the opposite type of ion, but it has no
preference for one neighboring ion over the others; there is no directionality to ionic bonds.
If this ionic solid were placed in something that dissolved it (for example, water in this case),
each of the ions would dissociate entirely from the others and become an isolated ion.
Covalent bonds involve atoms sharing electrons, instead of losing or gaining them entirely.
This is a way for multiple atoms to use the same set of electrons to get a completed shell of
valence electrons – to “fill their octets” (or “duets” in the case of hydrogen, since it can only
fit two electrons into its shell). The simplest example is hydrogen gas, H2. Each hydrogen
atom has a single valence electron, but pairing up with another H atom gives each one access
to two shared electrons, or a filled duet. This is shown with a Lewis dot diagram or Lewis
structure, where each electron is a dot. Shared electron pairs can either be shown as two
dots between those atoms, or as a line linking those atoms.
Another example is methane, or CH4, where a central carbon atom is surrounded by four
hydrogen atoms. After sharing electrons, each H atom has a filled duet and the carbon has a
filled octet.
Another example is water, or H2O. Here, oxygen already has 6 valence electrons, so it only
needs to share with two hydrogen atoms for a full octet. In this case, oxygen ends up with
two lone pairs – electrons that are not shared – as well as two shared pairs. (Later on, these
lone pairs will sometimes be omitted to simplify things, although they’re still implied to be
there.)
Two atoms can share more than just one pair of electrons, if that’s the only way to complete
their octet in a given compound. In ethene (a.k.a. ethylene), or C2H4, there are not enough H
atoms to share electrons with each carbon. In this case, the carbon atoms share two pairs of
electrons between themselves, shown as two pairs of dots between C atoms. This can also be
shown as a double bond linking the carbon atoms.
In ethyne (a.k.a. acetylene), or C2H2, the carbon atoms share three pairs of electrons. This is
shown as a triple bond between carbon atoms.
Another way to describe this is with bond order, the number of shared electron pairs
between two atoms. The carbon atoms in ethene have a bond order of two (a double bond),
while the carbon atoms in ethyne have a bond order of three (a triple bond).
2
Jacquie Richardson, CU Boulder – Last updated 9/2/2020
Organic chemistry is focused on carbon, and is used in making pharmaceuticals and
materials including polymers and liquid crystals. It’s also the basis for biochemistry. It uses
primarily the top three rows (a.k.a. periods) of the periodic table, plus a few elements from
lower periods. The rightmost column (a.k.a. group) is the noble gases (He, Ne, etc.), and the
one next to it (F, Cl, etc.) is the halogens.
Group 1 Group 2 … Group 13 Group 14 Group 15 Group 16 Group 17 Group 18
Period H (1) He (2)
1
Period Li (3) Be (4) B (5) C (6) N (7) O (8) F (9) Ne (10)
2
Period Na (11) Mg (12) Al (13) Si (14) P (15) S (16) Cl (17) Ar (18)
3
Period Br (35)
4
Period I (53)
5
The number next to each element is its atomic number, which describes the number of
protons in the atom. Each atom consists of a small, heavy nucleus made up of protons (with
positive charge) and neutrons (neutral, with no charge), surrounded by much lighter
electrons (negative charge). Protons determine which element an atom is, electrons
determine the overall charge on that atom, and neutrons don’t have much effect on chemical
behavior except in a few specific cases. For an atom to be electrically neutral, it must have
the same number of protons and electrons. For example, a neutral fluorine atom (F) has 9
protons and 9 electrons, plus some number of neutrons.
The electrons in an atom exist in shells around the nucleus. Each shell has a maximum
amount of electrons it can hold (2 for the first, 8 for subsequent shells…as far as organic
chemists are concerned most of the time). The shells fill up starting closer to the nucleus and
moving outwards. The electrons in filled shells are core electrons, and those in the unfilled
outermost shell are valence electrons. The core electrons of each atom have the same
configuration as the noble gas immediately before that atom in the table. For example,
lithium (Li) has the same core electron configuration as helium (He), and chlorine (Cl) has
the same core electron configuration as neon (Ne). The remaining electrons above the noble
gas configuration are the valence electrons – Li has 1 valence electron, and Cl has 7.
Chemical behavior is based primarily on the behavior of valence electrons, since these are
furthest from the nucleus and can interact with the “outside world” most easily. When an
atom becomes a charged ion by losing or gaining electrons, it tends to do this by going to
the same electron configuration as the nearest noble gas. For example, sodium (Na) has one
valence electron and is closest to Ne, so it will most likely drop a single electron to become a
sodium cation, Na+. Chlorine has 7 valence electrons and will most likely pick up one more
to become a chloride anion, Cl-. (Note that because electrons have negative charge, extra
electrons give something a – charge, not +. This is a very common mistake.)
Chemical Bonds
Atoms can bond together in several different ways. The two biggest subdivisions are
between ionic and chemical bonds.
1
, Loudon Ch. 1 Review: Chemical Structure & Bonds
Jacquie Richardson, CU Boulder – Last updated 9/2/2020
Ionic bonds involve electrostatic attraction between ions, each with a complete positive or
negative charge(s). The most common example is table salt, NaCl. It consists of alternating
Na+ and Cl- ions. Each ion is surrounded by the opposite type of ion, but it has no
preference for one neighboring ion over the others; there is no directionality to ionic bonds.
If this ionic solid were placed in something that dissolved it (for example, water in this case),
each of the ions would dissociate entirely from the others and become an isolated ion.
Covalent bonds involve atoms sharing electrons, instead of losing or gaining them entirely.
This is a way for multiple atoms to use the same set of electrons to get a completed shell of
valence electrons – to “fill their octets” (or “duets” in the case of hydrogen, since it can only
fit two electrons into its shell). The simplest example is hydrogen gas, H2. Each hydrogen
atom has a single valence electron, but pairing up with another H atom gives each one access
to two shared electrons, or a filled duet. This is shown with a Lewis dot diagram or Lewis
structure, where each electron is a dot. Shared electron pairs can either be shown as two
dots between those atoms, or as a line linking those atoms.
Another example is methane, or CH4, where a central carbon atom is surrounded by four
hydrogen atoms. After sharing electrons, each H atom has a filled duet and the carbon has a
filled octet.
Another example is water, or H2O. Here, oxygen already has 6 valence electrons, so it only
needs to share with two hydrogen atoms for a full octet. In this case, oxygen ends up with
two lone pairs – electrons that are not shared – as well as two shared pairs. (Later on, these
lone pairs will sometimes be omitted to simplify things, although they’re still implied to be
there.)
Two atoms can share more than just one pair of electrons, if that’s the only way to complete
their octet in a given compound. In ethene (a.k.a. ethylene), or C2H4, there are not enough H
atoms to share electrons with each carbon. In this case, the carbon atoms share two pairs of
electrons between themselves, shown as two pairs of dots between C atoms. This can also be
shown as a double bond linking the carbon atoms.
In ethyne (a.k.a. acetylene), or C2H2, the carbon atoms share three pairs of electrons. This is
shown as a triple bond between carbon atoms.
Another way to describe this is with bond order, the number of shared electron pairs
between two atoms. The carbon atoms in ethene have a bond order of two (a double bond),
while the carbon atoms in ethyne have a bond order of three (a triple bond).
2
