Transcript Chapter 18: Chemical Bonds

Chapter 18: Chemical Bonds
• Section 1: Stability in Bonding
• Section 2: Types of Bonds
• Section 3: Writing Formulas and
Naming Compounds
Section 1: Stability in Bonding
Why do atoms combine to form compounds?
• Because the atoms are “looking” for a stable structure
• Stability is achieved when the outermost energy level of the
atom is full
 With the exception of hydrogen and helium atoms, an
atom is structurally and chemically stable when there are
8 electrons in the outer energy level
• The group 8A elements have a full outer energy level and are
stable and do not need to combine with atoms
Atoms achieve stability through chemical bonding
• Chemical bond – the force that holds the atoms in a
compound together
• There are two types of chemical bonds
 Ionic – the giving or accepting of electrons to achieve a
full energy level
 Covalent – the sharing of one or more pair of electrons to
achieve a full energy level
ONLY THE ELECTRONS IN AN ATOM’S OUTERMOST ENERGY
LEVEL ARE AVAILABLE TO FORM CHEMICAL BONDS
Section 2: Types of Bonds
Ionic bonding
• When atoms gain or lose electrons they acquire a charge
 Ion – an atom that carries a charge
 Typically, metal will give up electron(s) and become ions
with a positive
 Nonmetals will accept electron(s) and become ions with a
negative charge
 On your Periodic Table, the main group elements (the A’s),
the number before the letter “A” indicates the number of
electrons in the outermost energy level.
Exs.: Calcium – Group = 2A – 2 electrons
Oxygen – Group = 6A – 6 electrons
 The B group, or transition, elements can have a varying
number of valence electrons
• The formation of ionic bonds
 Example: the formation of salt (NaCl)
 According to the Periodic Table, sodium (Na) is a 1A
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


metal and chlorine (Cl) is a 7Anonmetal
Na will give up 1 electron to gain a noble gas
configuration while chlorine will accept 1 electron to
achieve the full outer energy level
By giving up a electron, sodium becomes an ion with a
+ charge, Na+, while chorine, by accepting an electron,
becomes an ion with a negative charge, ClBecause the two ions have opposing charges, they
attract and form an ionic compound
The compound NaCl is an ionic compound
Section 2: Types of Bonds
Ionic bonding
• Conclusions in the formation of ionic compounds
 metals bond with nonmetals
 metals give up electrons, nonmetals accept electrons
 the compound formed is electrically neutral – the (+)
charges and the (–) charges add up to 0
Covalent bonding
• A covalent bond forms when two atoms share one or more
pair of valance electrons
 Only nonmetals form compounds through covalent bonds
 The reason for this is that nonmetals have nearly
complete outermost energy levels so it is easier for
them to share electrons than it is to give up or add
electrons when bonding with other nonmetals
 Example: the formation of the gas fluorine (F2)
 According to the Periodic Table, fluorine is a 7A
nonmetal, and needs one electron to achieve stability
 If the two fluorine atoms bond ionically, one fluorine
atom would give up an electron and the other accept
it. As a result, one atom would be stable but the other
would be two electrons short of having a full outer
level. This is not good. Remember, the whole goal of
chemical bonding is that all the atoms involved in
bonding achieve stability
 Solution: each fluorine atom shares one electron with
the other
 This sharing of electrons to form stable octets is called
covalent bonding, and the smallest particle of a
covalent compound is called a molecule
Section 3: Writing Formulas and
Naming Compounds
Naming compounds
1. Ionic compounds
 Use the name of the metal
 Use the root of the nonmetal with the “ide” suffix
 Example: MgBr2 - magnesium `bromide
 Use the name of the polyatomic (table 4, page 619) if
present
 Example: CaCO3 – calcium carbonate
2. Binary covalent compounds
 Use the name of the first nonmetal followed by the root
of the second with the “ide” suffix
 Use a prefix to indicate how many of each element is in a
molecule
 Always use prefix with 2nd nonmetal
 Only use prefix with 1st nonmetal when there is more
than one atom in the formula
 Example: CO2 = carbon dioxide (refer to table 5, page 621
for common prefixes)
Section 3: Writing Formulas and
Naming Compounds
The chemical formula contains the symbol for each atom in the
compound, and tells something about number of atoms of each
element in the compound
• The formula for an ionic compound gives the simplest ratio of
the different ions in the pound.
 Example: Calcium chloride
Formula: CaCl2
The subscript 2 tells you that for every Ca2+ ion
there are two Cl- ions
• Covalent compounds
 Formulas for covalent compound tell the exact number of
each kind of atom in the molecule
 Example: glucose and sucrose
1. both are sugars – glucose is plant sugar used by plants
in photosynthesis and sucrose is ordinary table sugar
2. both are composed of atoms of carbon, oxygen, and
hydrogen covalently bonded together
3. the formulas: glucose: 𝑪𝟔 𝑯𝟏𝟐 𝑶𝟔
sucrose: 𝑪𝟏𝟐 𝑯𝟐𝟒 𝑶𝟏𝟐
Section 3: Writing Formulas and
Naming Compounds
Formulas of Ionic Compounds
In any ionic compound the net charge must equal zero
• Use the criss-cross method to determine the formula for
ionic compounds
 Ex 1: Magnesium fluoride
Mg → 2A metal – will give up 2
electrons
F → 7A nonmetal – will accept 1
electron
1. write symbols of the ions showing the
charge of each
2. criss-cross the superscript numerals
3. the numerals now become the subscript
designating how many of which ion is in the compound
4. disregard the + and - signs
5. when writing the formula never show the subscript “1”
Section 3: Writing Formulas and
Naming Compounds
Formulas of Ionic Compounds
Ex. 2: calcium hydroxide
Ca → 2A metal, will lose 2 electrons
OH → a polyatomic ion with a net
charge of –1
1. criss-cross numeral superscripts
2. ignore the signs, numbers become subscripts
3. when writing the formula enclose the
OH ion in parentheses to indicate that it
is a single unit
Oxidation number – A positive or negative number that
indicates how many electrons an atom has gained, lost, or
shared to become stable
• The oxidation number for the “A” group elements is easy to
calculate
 Metals – the oxidation # is number with the letter
 Ex: the 1A metals have an oxidation number of +1
 Nonmetals – the oxidation # is the number that must be
added to reach 8
 Ex: Oxygen, 6A, so: 6 + x = 8, x = 8-6, x = 2, the 6A
nonmetals have an oxidation # of -2
Section 3: Writing Formulas and
Naming Compounds
• For the “B” group elements, known as the
transition metals, determining the oxidation
number can be more difficult
 The atoms of the transition elements can have
their valence electrons distributed between
two sublevels
 During the formation of an ionic compound, the
transfer of electrons may involve one or both of
the sublevels
 A varying number of electrons can be lost, any
number between 2 and 7
 It is impossible to predict the oxidation number
of the element
• Analysis of the ionic compound can give the
simplest whole number ratio of the metal and
nonmetal ions
 Ex: what is the oxidation # of iron in the compound
FeCl3
1. we know that chlorine always forms an ion with a
-1 charge
2. we know the net charge of the compound must =
0
so: total (-) = -1x3 = -3, and: total (+) = total (-)
so: the oxidation # for Fe = +3
Section 3: Writing Formulas and
Naming Compounds
• The oxidation number of the transition metal is given as a
roman numeral in the compound name
 Ex: Iron (II) oxide and Iron (III) oxide
• We can now determine the formula for those compounds
using the criss-cross method:
Iron (II) Oxide
Iron (III) Oxide