3. KEY POINTS
• A chemical bond is any force of attraction that holds two atoms or
ions together.
• In most cases, that force of attraction is between one or more
negatively charged electrons held by one of the atoms and the
positively charged nucleus of the second atom.
• Chemical bonds vary widely in their strength, ranging from
relatively strong covalent bonds (in which electrons are shared
between atoms) to very weak hydrogen bonds.
• The term chemical bond also refers to the symbolism used to
represent the force of attraction between two atoms or ions. For
example, in the chemical formula H—O—H, the short dashed lines
are known as chemical bonds.
4. HISTORY
• Theories of chemical bonds go back a long time.
• One of the first was developed by Roman poet Lucretius (c. 95–c.
55 B.C. ), author of De Rerum Natura (title means "on the nature of
things").
• In this poem, Lucretius described atoms as tiny spheres with
fishhook-like arms.
• Atoms combined with each other, according to Lucretius, when the
hooked arms of two atoms became entangled with each other.
5.
6. IONIC BONDING
Ionic bonding
Consider now a more difficult situation, one in which two different
atoms compete for electrons.
• One example would be the case involving a sodium atom and a
chlorine atom. If these two atoms come close enough to each
both nuclei pull on all electrons of both atoms. In this case,
a very different result occurs. The chlorine nucleus has a much
charge than does the sodium nucleus. It can pull on sodium's
electrons much more efficiently than the sodium nucleus can pull
the chlorine electrons.
• In this case, there is a winner in the battle: chlorine is able to pull
of sodium's electrons away. It adds that electron to its own
of electrons. In a situation in which one atom is able to completely
remove an electron from a second atom, the force of attraction
between the two particles is known as an ionic bond.
9. To determine the chemical formulas of ionic
compounds, the following two conditions
must be satisfied:
1.Each ion must obey the octet rule for
maximum stability.
2.Ions will combine in a way that the overall
ionic compound will be neutral. In other
words, the charges of the ions must balance
out.
10.
11. ionic bond: sodium chloride,
or table salt Ionic bonding in
sodium chloride. An atom of
sodium (Na) donates one of
its electrons to an atom of
chlorine (Cl) in a chemical
reaction, and the resulting
positive ion (Na+) and
negative ion (Cl−) form a
stable ionic compound
(sodium chloride; common
table salt) based on this ionic
bond.
12. EXAMPLE 1: CHLORIDE SALTS
In this example, the sodium atom is donating its 1 valence electron to the chlorine
atom. This creates a sodium cation and a chlorine anion. Notice that the net charge of
the resulting compound is 0.
In this example, the magnesium atom is donating both of its valence electrons to chlorine
atoms. Each chlorine atom can only accept 1 electron before it can achieve its noble gas
configuration; therefore, 2 atoms of chlorine are required to accept the 2 electrons
donated by the magnesium. Notice that the net charge of the compound is 0.
13. EXAMPLE 2: MAGNESIUM FLUORIDE
Mg most commonly forms a 2+ ion. This is because Mg has two valence electrons and it would like
to get rid of those two ions to obey the octet rule. Fluorine has seven valence electrons and as such,
usually forms the F– ion because it gains one electron to satisfy the octet rule. When Mg2+ and F–
combine to form an ionic compound, their charges must cancel out. Therefore, one Mg2+ needs two
F– ions to balance. The 2+ of the Mg is balanced by having two -1 charged ions. Therefore, the
formula of the compound is MgF2. The subscript two indicates that there are two fluorine's that are
ionically bonded to magnesium.
14. KEY POINTS
•Ionic bonds are formed through the exchange of
valence electrons between atoms, typically a
metal and a nonmetal.
•The loss or gain of valence electrons allows ions
to obey the octet rule and become more stable.
•Ionic compounds are typically neutral. Therefore,
ions combine in ways that neutralize their
charges.
15. COVALENT BONDING
• Covalent bonding is the sharing of electrons
between atoms. This type of bonding occurs
between two atoms of the same element or of
elements close to each other in the periodic table.
• This bonding occurs primarily between nonmetals;
however, it can also be observed between
nonmetals and metals.
16. Covalent bonding
• Today, it is widely accepted that most examples of chemical
represent a kind of battle between two atoms for one or more
electrons. Imagine an instance, for example, in which two hydrogen
atoms are placed next to each other.
• Each atom has a positively charged nucleus and one electron
spinning around its nucleus. If the atoms are close enough to each
other, then the electrons of both atoms will be attracted by both
nuclei. Which one wins this battle?
• The answer may be obvious. Both atoms are exactly identical. Their
nuclei will pull with equal strength on both electrons. The only
possible result, overall, is that the two atoms will share the two
electrons with each other equally. A chemical bond in which two
electrons are shared between two atoms is known as a covalent
bond.
17.
18. EXAMPLE 1: PHOSPHORUS TRICHLORIDE
(PHOSPHORUS CHLORIDE, PCL3)
In this example, a phosphorous atom is sharing its three unpaired
electrons with three chlorine atoms. In the end product, all four of
these molecules have 8 valence electrons and satisfy the octet rule.
19. KEY POINTS
•Covalent bonds involve two atoms, typically nonmetals,
that share electron density to form strong bonding
interactions.
•Covalent bonds include single, double, and triple
bonds and are composed of sigma and pi bonding
interactions where 2, 4, or 6 electrons are shared
respectively.
•Covalent compounds typically have lower melting and
boiling points than ionic compounds.
20. WHICH TYPE OF BOND IS STRONGER?
Ionic bonds take more energy to break than
covalent bonds, so ionic bonds are
stronger. The amount of energy needed to
break a bond is known as bond dissociation
energy, which is basically the force it takes
to break bonds of any type.
