Structure and bonding of chemistry

Kelompok 1 :Nurul Safitry

Arum estu tami Novita Kumala Sari

Yolanda Aprita

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.

• 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:

• 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.

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.

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.

IONIC BONDINGLikewise, 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:

• 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:

• 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.

• 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.

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.

• 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:

• 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:

• 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:

• 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.

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):

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.

• 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:

• 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:

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

Chemical bonding

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