- Atoms consist of electrons surrounding a positively charged nucleus made up of protons and neutrons. Different types of atoms have different numbers of protons, electrons, and neutrons. - There are two main types of chemical bonding: ionic bonding and covalent bonding. Ionic bonding involves the transfer of electrons between atoms to form oppositely charged ions that are attracted to each other. Covalent bonding involves the sharing of electrons between atoms. - Covalent bonding allows for many different molecular structures and isomers. Molecules can be nonpolar or polar depending on electron distribution. Metals exhibit metallic bonding where electrons are delocalized and shared among many neighboring atoms. Hydrogen bonding is a weaker interaction between polar molecules that

1. Structure and bonding of
chemistry
Kelompok 1 :
Nurul Safitry
Arum estu tami
Novita Kumala Sari
Yolanda Aprita

2. Structure Atom
What are molecules made from? They are made from atoms, which are
themselves made from nuclei and electrons. These building blocks carry an
electrical charge: nuclei are positively charged, and electrons are negatively
charged. The nuclei themselves are made up of (positively charged) protons and
(neutral) neutrons. This is all summarised on the following picture:
Different types of atom have different numbers of protons, neutrons, and
electrons. For example, carbon atoms have 6 protons, 6 neutrons, and 6
electrons.

3. • Charged species interact with each other: like
charges (+ and + or - and -) repel each other,
opposite charges (- and +) attract each other.
This well-known principle from physics is
summarised by Coulomb's law:

4. • Atoms consist of electrons surrounding a nucleus that
contains protons and neutrons. Neutrons are neutral,
but protons and electrons are electrically charged.
Protons have a relative charge of +1, while electrons
have a relative charge of -1.
• The number of protons in an atom is called its atomic
number. In the periodic table atoms are arranged in
atomic number orderElectrons are arranged in energy
levels or shells, and different energy levels can hold
different numbers of electrons. The electronic
structure of an atom is a description of how the
electrons are arranged, which can be shown in a
diagram or by numbers.

5. LEWIS STRUCTURES OF THE ELEMENTS
• Examine the Sideways Periodic Chart With Electron
Shell Numbers again. All of the Group I elements and
hydrogen (the top row of the chart) have one and only
one electron in the outside shell. That single electron is
what gives these elements the distinctive character of
the group. The Lewis structures are just an attempt to
show these valence electrons in a graphic manner as
they are used to combine with other elements. The
element symbol is in the center and as many as four
groups of two electrons are shown as dots above,
below, to the right and left of the element symbol to
show the valence electrons.

6. Chemical bonding
• Elements in the first few columns of the periodic table have a few
more electrons than predicted by the octet rule: they therefore lose
the electrons in the outermost shells fairly easily. For example, the
alkali metals (group I), such as sodium (Na) or potassium (K), which
have, respectively, 11 and 19 electrons, easily lose one electron to
form monopositive ions, Na+ and K+. These ions have 10 and 18
electrons, respectively, so they are quite stable according to the
octet rule.
• Elements in the last few columns of the periodic table have one,
two or three fewer electrons than predicted by the octet rule: they
therefore gain electrons fairly easily. For example, the halogens
(group VII), such as fluorine (F) or chlorine (Cl), which have,
respectively, 9 and 17 electrons, easily gain one electron to form
mononegative ions, F- or Cl-. These ions have 10 and 18 electrons,
respectively.

7. IONIC BONDING
Likewise, elements in group II form doubly positive ions such as
Mg++ or Ca++, and elements in group VI form doubly negative
ions such as O-- or S--. All these ions obey the octet rule and so
are fairly stable.
Now, imagine what will happen when one sodium atom meets one
chlorine atom: the sodium atom will lose one electron to give
Na+, and the chlorine atom will gain that electron to give Cl-. This
can be represented schematically in the following way:

8. • The resulting ions, which have Opposite charges, will be
attracted to one another, and will draw closer, until they
"touch". This happens when the inner shell of electrons on
the sodium ion (shown in blue) starts to overlap with the
outer shell of electrons on the chloride anion (shown in
green). This pair of ions looks something like this:

9. • The stable form of sodium chloride involves a very large
number of NaCl units arranged in a lattice (or regular
arrangement) millions of atoms across. Because the lattice is
rigid, this means that one gets a solid: the ions do not move
much one with respect to another.

10. • Ionic bonds form between elements which
readily lose electrons and others which readily
gain electrons. Because the interaction between
charges as given by Coulomb's law is the same in
all directions, ionic compounds do not form
molecules. Instead, periodic lattices with billions
of ions form, in which each ion is surrounded by
many ions of opposite charge. Therefore, ionic
compounds are almost always solids at room
temperature. By careful consideration of the
properties of each ion, it is possible to design
ionic solids with certain well-defined and
desirable properties, like superconductors.

11. COVALENT BONDING
In the previous page, we saw
how atoms could achieve a
complete shell of electrons by
losing or gaining one or more
electrons, to form ions. There is
another way atoms can satisfy
the octet rule: they can share
electrons. For example, two
hydrogen atoms can share their
electrons, as shown below.
Because each of the shared
electrons then "belongs" to both
atoms, both atoms then a fulled
shell, with two electrons. The
pair of shared electrons is
symbolised by the heavy line
between the atoms.

12. • The methane (CH4) molecule illustrates a more complex
example. Each of the 4 electrons in the outermost ("valence")
shell of carbon is shared with one hydrogen. In turn, each of
the hydrogens also shares one electron with carbon. Overall,
carbon "owns" 10 electrons - satisfying the octet rule - and
each hydrogen has 2. This is shown here:

13. • When a molecule of methane is studied
experimentally, it is found that the four
hydrogens spread out evenly around the
carbon atom, leading to the three-
dimensional structure shown here:

14. • When we add two more carbon atoms and 4 more hydrogens,
to make butane (C4H10), an interesting situation arises: There
are two different ways of bonding the carbons together, to
form two different molecules, or isomers!!! These are shown
below. For one of the isomers, the first carbon is bonded to
three hydrogens, and to the second carbon, which is itself
bonded to another two hydrogens and to the third carbon,
which is itself bonded to the fourth carbon. In the other
isomer, one of the carbons forms a bond to all three carbon
atom s:

15. • Covalent bonds involve sharing electrons between atoms.
The shared electrons "belong" to both atoms in the bond.
Each atom forms the right number of bonds, such that they
have filled shells. There is lots of flexibility in terms of
which atom bonds to which other ones. This means that
many isomeric molecules can be formed, and Nature as
well as chemists are skilled at designing and making new
molecules with desirable properties. In most cases, only a
small number of atoms are bonded together to make a
molecule, and there is no bonding between atoms in one
molecule and other atoms in other molecules. This means
that molecules are only very slightly "sticky" between
themselves, and covalent compounds are either gases, or
liquids, or sometimes solids. In some cases, bonding occurs
to form large molecules with thousands or millions of
atoms, and these can be solids.

16. OTHER BONDING
• There are other principles which can lead to atoms bonding
to each other, and we will examine here two important
cases: metallic bonding, and hydrogen bonding.
• Metallic Bonding
• Metals are well known to be solids (except for Mercury!).
The bonds between metals can loosely be described as
covalent bonds (due to sharing electrons), except that the
metal atoms do not just share electrons with 1, 2, 3 or 4
neighbours, as in covalent bonding, but with many atoms.
The structure of the metal is determined by the fact that
each atom tries to be as close to as many other atoms as
possible. This is shown here for one typical metal structure
(adopted, for instance, by iron at some temperatures):

17. Can you count how many neighbours each iron atom is bonded to? Contrast
this with the structure of diamond seen previously.
Because the electrons are shared with all the neighbours, it is quite easy for
the electrons in metals to move around. If each "shared" electron shifts one
atom to the right or left, this leads to a net shift of charge. This occurs quite
easily in metals, but much less so in ionic solids, or covalent ones, where the
electrons are rigidly associated with either a particular atom or ion, or a
particular pair of atoms. It is because electrons can move around so easily
inside metals that the latter conduct electricity.

18. • HYDROGEN BONDING
• In covalent bonds, the electrons are shared, so that each atom
gets a filled shell. When the distribution of electrons in
molecules is considered in detail, it becomes apparent that
the "sharing" is not always perfectly "fair": often, one of the
atoms gets "more" of the shared electrons than the other
does.
• This occurs, in particular, when atoms such as nitrogen,
fluorine, or oxygen bond to hydrogen. For example, in HF
(hydrogen fluoride), the structure can be described by the
following "sharing" picture:

19. • However, this structure does not tell the whole truth about
the distribution of electrons in HF. Indeed, the following,
"ionic" structure also respects the filled (or empty) shell rule:
• In reality, HF is described by both these structures, so that the
H-F bond is polar, with each atom bearing a small positive
(δ+) or negative (δ-) charge. When two hydrogen fluoride
molecules come close to each other, the like charges attract
each other, and one gets a "molecule" of di-hydrogen fluoride
as shown:

20. Hydrogen Bonding
• The weak "bond" between the F atom and the H is called a
Hydrogen Bond, and is shown here as the dotted green
line.
• Hydrogen bonds can also occur between oxygen atoms and
hydrogen. One of the most important types of hydrogen
bonds is of this type, and is the one occurring in water. As
discussed for HF, the electrons in H2O molecules are not
evenly "shared": the oxygen atom has more of them than
the hydrogen atoms. As a result, oxygen has a (partial)
negative charge, and the hydrogens have a positive charge.
When you have two water molecules close to another, a
hydrogen atom on one of the molecules is attracted to the
oxygen of the other molecule, to give a dimer

21. Chemical bonding

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