Chapter 4:
Forces Between Particles
1
NOBLE GAS CONFIGURATIONS
• An electronic configuration that is characterized by two
electrons in the valence shell of helium and eight electrons
in the valence shell of all other group VIIIA noble gases.
2
LEWIS STRUCTURES
• A representation of an atom or ion in which the elemental
symbol represents the atomic nucleus and all but the
valence-shell electrons. The valence electrons are
represented by dots arranged around the elemental
symbol.
3
DETERMINATION OF THE NUMBER OF VALENCE
ELECTRONS IN AN ATOM
• Write an electronic configuration for the atom and
identify the valence electrons as those having the largest
n value in the configuration.
• A simpler alternative for representative elements is to
refer to the periodic table and note the group to which
the element belongs. The number of valence electrons
is the same as the Roman numeral group number.
• Examples: Calcium, Ca, is in group IIA. The number of
valence electrons is 2. Phosphorus, P, is in group VA.
The number of valence electrons is 5.
4
EXAMPLE OF A LEWIS STRUCTURE
• Potassium, K, is in group IA and so it has one valence
electron. The Lewis structure is K.
• Aluminum, Al, is in group IIIA and so it has three valence
electrons. The Lewis structure is Al : .

5
THE OCTET RULE
• According to the octet rule, atoms will gain or lose
sufficient electrons to achieve an outer electron
arrangement identical to that of a noble gas. This
arrangement usually consists of eight electrons in the
valence shell.
SIMPLE ION
• An atom that has acquired a net positive or negative
charge by losing or gaining one or more electrons.
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EXAMPLES OF SIMPLE IONS
• Magnesium, Mg, has two valence electrons which it loses to
form a simple ion with a +2 electrical charge. The ion is
written as Mg2+.
• Oxygen, O, has six valence electrons. It tends to gain two
electrons to form a simple ion with a -2 electrical charge.
The ion is written as O2-.
• Bromine, Br, has seven valence electrons. It tends to gain
one electron to form a simple ion with a -1 electrical charge.
The ion is written as Br -.
7
A SIMPLE WAY TO OBTAIN IONIC CHARGES FOR
REPRESENTATIVE ELEMENTS
• Representative metals will form ions having the same
positive charge as the number (Roman numeral) of the
group to which they belong.
• Representative nonmetals will form ions with a negative
charge equal to 8 minus the number (Roman numeral) of
the group to which they belong.
• For example, strontium, Sr, a group IIA metal forms Sr2+
ions and phosphorus, P, a group VA nonmetal forms P3ions.
8
IONIC BOND FORMATION
• Ions with positive charges are attracted to ions with
negative charges. The attractive force between such ions
holds them together and is called an ionic bond.
• Ionic bonds form when representative metal atoms lose
valence electrons. The electrons are gained by
representative nonmetal atoms. Both atoms are changed
into ions with noble gas configurations. The resulting ions
are then attracted to each other.
9
ISOELECTRONIC
• A term that literally means “same electronic,” used to
describe atoms or ions that have identical electronic
configurations
10
IONIC COMPOUNDS
• The substances formed when ionic bonds form between
positive and negative ions are called ionic compounds.
• When ionic compounds are formed by the reaction of only
two elements the resulting ionic compound is called a
binary ionic compound.
Cu2O
CuO
11
FORMULAS FOR BINARY IONIC COMPOUNDS
• Binary ionic compounds typically form when a metal and a
nonmetal react.
• The metal tends to lose one or more electrons and forms a
positive ion.
• The nonmetal tends to gain one or more electrons and
forms a negative ion.
• The symbol for the metal is given first in the formula.
• The formula for a binary ionic compound represents the
minimum number of each ion that when combined together
will provide equal numbers of positive and negative
electrical charges.
NaCl
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EXAMPLES OF FORMULAS FOR BINARY IONIC
COMPOUNDS
• Sodium and fluorine:
• Sodium, a group IA metal, will form sodium ions with the
symbol Na+.
• Fluorine, a group VIIA nonmetal, will form ions with the
symbol F-.
• The minimum number of ions needed to give the same
number of positive and negative charges is one of each.
• The one Na+ provides one positive charge and the one Fprovides one negative charge.
• The correct formula that results is NaF.
13
• Sodium and sulfur:
•Sodium is a group IA metal and will form sodium ions with
the symbol Na+.
•Sulfur is a group VIA nonmetal and will form ions with the
symbol S2-.
•The minimum number of ions required to give the same
number of positive and negative charges is two Na+ ions
and one S2- ion.
•The two Na+ ions provide two positive charges and the one
S2- ion provides two negative charges.
•The resulting formula is Na2S.
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• Aluminum and oxygen:
•Aluminum is a group IIIA metal and will form ions with the
symbol Al3+ .
•Oxygen is a group VIA nonmetal and will form ions with the
symbol O2-.
•The minimum number of ions required to give the same
number of positive and negative charges is two Al3+ ions
and three O2- ions.
•The resulting formulas is Al2O3.
15
NAMING BINARY IONIC COMPOUNDS
• Binary ionic compounds are named using the following
pattern:
• name = metal name + stem of nonmetal name + -ide
• The stem names and ionic symbols for some common
nonmetals are given in the following table:
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EXAMPLES OF BINARY IONIC COMPOUND NAMES
• Name K2O:
name = metal name + nonmetal stem + -ide
name = potassium + ox- + -ide = potassium oxide
• Name Mg3N2:
name = metal name + nonmetal stem + -ide
name = magnesium + nitr- + -ide = magnesium nitride
• Name BeS:
name = metal name + nonmetal stem + -ide
name = beryllium + sulf- + -ide + beryllium sulfide
• Name AlBr3:
name = metal name + nonmetal stem + -ide
name = aluminum + brom- + -ide
name = aluminum bromide
17
NAMING BINARY IONIC COMPOUNDS IN WHICH
METALS FORM IONS WITH MORE THAN ONE
CHARGE
• Some metal atoms, especially those of transition and
inner-transition elements form more than one type of
charged ion. (e.g. Cobalt forms both Co2+ and Co3+ ions.)
• The binary compounds containing such ions are named
following the pattern given earlier with one addition, the
number of positive charges on the metal ion is indicated by
a Roman numeral in parentheses following the metal
name.
• For example, the compounds CoO and Co2O3 contain
cobalt ions with 2+ and 3+ charges respectively. Their
names are cobalt (II) oxide and cobalt (III) oxide.
18
IONIC COMPOUND STRUCTURE
• The stable form of an ionic compound is not a molecule,
but a crystal in which many ions of opposite charge occupy
lattice sites in a rigid three-dimensional arrangement
called a crystal lattice.
19
IONIC COMPOUND FORMULAS AND WEIGHTS
• Formulas for ionic compounds represent only the simplest
combining ratio of the ions in the compounds, not the
precise numbers of atoms of each element found in a
molecule.
• Formula weight is the sum of the atomic weights of the
atoms shown in the formula of an ionic compound. This is
similar to molecular weight.
• One mole of an ionic compound contains Avogadro’s
number (6.022 x 1023) of the simplest combining ratio of
ions in the compounds.
20
COVALENT BONDING
• Covalent bonding is a type of bonding in which the octet
rule is satisfied when atoms share valence electrons. The
shared electrons are counted in the octet of each atom that
shares them as illustrated below for fluorine gas, F2.
• The atoms sharing one or more pairs of electrons are each
attracted to the shared electrons, and thus, are attracted to
each other. The attraction to each other is called a
covalent bond. The covalent bond may be represented by
the shared pair or by a single line between the bonded
atoms.
21
• The sharing of electrons takes place when electroncontaining orbitals of atoms overlap. This is shown below
for the formation of the H2 molecule.
22
• Electron sharing resulting in covalent bonding can occur
between identical atoms or between different atoms.
• Molecules such as Cl2, O2 and N2 are formed when electron
sharing occurs between identical atoms.
• Molecules such as H2O, and CH4 are formed when electron
sharing occurs between different atoms.
H2O
CH4
23
EXAMPLES OF COVALENT BONDING
24
DRAWING LEWIS STRUCTURES
FOR COVALENT MOLECULES
• Step 1:
•Use the molecular formula to determine how many atoms
of each type are in the molecule.
• Step 2:
•Use the provided connecting pattern of atoms to draw an
initial molecular structure with the atoms properly arranged.
• Step 3:
•Determine the total number of valence-shell electrons
contained in the atoms of the molecule.
25
• Step 4:
•Put one pair of electrons between each bonded pair of
atoms in the initial structure drawn in Step 2.
• Subtract the number of electrons used in this step from the
total number determined in Step 3.
•Use the remaining electrons to complete the octets of all
other atoms in the structure, beginning with the atoms that
are present in greatest number in the molecule.
•Remember, hydrogen atoms only require one pair of
electrons to achieve the electronic configuration of helium.
• Step 5:
•If all octets cannot be satisfied with the available electrons,
move pairs that are not located between atoms to positions
between atoms to complete octets. This will create double
or triple bonds between some atoms.
26
EXAMPLE OF DRAWING A LEWIS STRUCTURE
FOR A COVALENT MOLECULE
• Draw a Lewis structure for SO3.
•Step 1:
– The formula indicates one S and three O atoms are in the
molecule.
•Step 2:
– The connecting pattern is that each O is bonded only to
the S. Thus, the following arrangement is drawn:
O S O
O
•Step 3:
– Sulfur and oxygen are both in group VIA, and so each
atom has 6 valence electrons. The total number of
electrons is 24 (6 from the one S atom and 18 from the
three O atoms).
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•Step 4:
– One pair of electrons is put between each O atom and the
S atom of the arrangement drawn in step 2.
– This required 6 of the 24 available electrons. The
remaining 18 are used to complete the octets of the
atoms, beginning with the O atoms.
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•Step 5:
– After step 4, it is seen that the octet of S is not completed,
even though all available electrons have been used.
– One nonbonding pair from any of the three O atoms will
be moved to a location between the O and the S atoms.
This pair will continue to count toward the octet of the O,
but will also now count toward the octet of the S.
– The resulting correct Lewis structure contains one double
bond (two shared pairs) between the S and one of the O
atoms.
29
POLYATOMIC IONS
• Polyatomic ions are covalently-bonded groups of atoms
that carry a net electrical charge. Most common
polyatomic ions are negatively charged.
• Lewis structures can be drawn for polyatomic ions using
the same steps that were shown earlier for covalent
molecules with one change. In Step 3, one electron is
added to the total for each negative charge found on the
polyatomic ion and one electron is subtracted from the
total for each positive charge found on the polyatomic ion.
All other steps are used unchanged.
30
SHAPES OF MOLECULES AND POLYATOMIC IONS
• Most molecules and polyatomic ions are not flat twodimensional objects. Most have distinct three-dimensional
shapes.
• The shapes of molecules or polyatomic ions can be
predicted using a theory called the valence-shell electron
repulsion theory, or VSEPR theory (sometimes
pronounced "vesper" theory).
• According to the VSEPR theory, electron pairs in the
valence shell of an atom will repel each other and get as
far away from each other as possible.
• When the theory is applied to the valence-shell electron
pairs of the central atom in a molecule or ion, the shape of
the molecule or ion can be predicted. A central atom is an
atom that is bonded to other surrounding atoms such as
the S atom in the SO3 molecule whose Lewis structure
was drawn in the earlier example.
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• When the VSEPR theory is used, two rules are followed:
Rule 1. All valence-shell electron pairs around the central
atom are considered to behave the same regardless of
whether they are bonding or nonbonding pairs.
Rule 2. Double or triple bonds between surrounding
atoms and the central atom are treated like a single pair of
electrons when shapes are predicted.
32
ARRANGEMENTS OF ELECTRON PAIRS
• The VSEPR theory can be applied to molecules or ions
that have up to six pairs of electrons around the central
atom. This discussion will be limited to molecules with no
more than four pairs.
33
• According to the VSEPR theory, the arrangement of
electron pairs around the central atom (represented by E)
depends on the number of electron pairs.
•Two pairs locate opposite each other.
•Three pairs arrange themselves in a flat triangle around the
central atom.
•Four pairs become located at the four corners of a pyramidlike shape called a tetrahedron.
34
EXAMPLES OF USING THE VSEPR THEORY
• Draw a Lewis structure and determine the shape of the
molecule for CO2.
• Solution:
•The Lewis structure drawn according to the rules given
earlier is O=C=O.
•The central C atom is surrounded by two double bonds.
•Each double bond counts as a single electron pair, thus the
central atom has two pairs of electrons around it.
•They will take up positions on opposite sides of the C.
•The O, C, and O atoms are arranged in a line and the
molecule is linear.
35
• Draw a Lewis structure and determine the shape of the
molecule for NH3.
• Solution:
•The Lewis structure drawn according to the rules given
earlier is:
•The central atom is N.
•It has four electron pairs surrounding it.
•The four pairs will be located at the corners
of a tetrahedron, with the N in the middle:
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• The shape of the molecule is determined only by the
positions of the atoms, not by the positions of the unshared
electron pair. Thus, the NH3 molecule has the shape of a
pyramid with a triangular base. The N atom is at the peak of
the pyramid and an H atom is at each corner of the base as
shown below:
37
POLARITY OF COVALENT MOLECULES
• The shared electrons of covalent bonds are not always
shared equally by the bonded atoms.
• Electrons of a covalent bond are attracted toward atoms of
highest electronegativity.
38
• Unequal sharing of the bonding electrons of a covalent bond
cause the bond to become a polar covalent bond.
• For atoms bonded by a polar covalent bond, the more
electronegative atom acquires a partial negative charge (-)
and the less electronegative atom acquires a partial positive
charge (+).
• When the resulting partial charges are distributed
symmetrically in a molecule, the molecule is nonpolar.
When the partial charges are distributed nonsymmetrically,
the molecule is polar.
39
THE POLARITY OF MOLECULES
40
NAMING BINARY COVALENT COMPOUNDS
• The pattern used to name binary covalent compounds is
similar to that used to name binary ionic compounds:
name = name of least electronegative element
+ stem of more electronegative element + -ide
• In addition to the pattern, the number of each type of atom
in the molecule is indicated by means of the following
Greek prefixes:
• Note: The prefix mono is not
used when it appears at
the beginning of the name.
41
EXAMPLES OF NAMING BINARY COVALENT
COMPOUNDS
SO2: name = sulfur + di- + ox + -ide = sulfur dioxide
XeF6: name = xenon + hexa- + fluor + -ide
= xenon hexafluoride
H2O: name = di- + hydrogen + mono- + ox + -ide
= dihydrogen monoxide (also known as water)
(Note, the final o of mono- was dropped for
ease of pronunciation.)
42
WRITING FORMULAS OF IONIC COMPOUNDS THAT
CONTAIN POLYATOMIC IONS
• The rules for writing formulas for ionic compounds
containing polyatomic ions are essentially the same as
those used for writing formulas for binary ionic compounds.
• The symbol for the metal is written first, followed by the
formula for the negative polyatomic ion. Equal numbers of
positive and negative charges must be represented by the
formula.
• When more than one polyatomic ion is required in the
formula, parentheses are placed around the polyatomic ion
before the subscript is inserted.
Na3PO4
Mg3 PO4 2
NH4 3 PO4
43
EXAMPLES OF FORMULAS FOR IONIC COMPOUNDS
CONTAINING POLYATOMIC IONS
• Compound containing K+ and ClO3-: KClO3
• Compound containing Ca2+ and ClO3-: Ca(ClO3)2
• Compound containing Ca2+ and PO43-: Ca3(PO4)2
44
NAMING IONIC COMPOUNDS THAT CONTAIN A
POLYATOMIC ANION
• The names of ionic compounds that contain a polyatomic
anion are obtained using the following pattern:
name = name of metal + name of polyatomic anion
• Examples:
•KClO3 is named potassium chlorate
•Ca(ClO3)2 is named calcium chlorate
•Ca3(PO4)2 is named calcium phosphate
•CaHPO4 is named calcium hydrogen phosphate
45
A SUMMARY OF INTERPARTICLE FORCES
• Ionic and covalent bonds represent two of the forces that
occur between atomic-sized particles and hold the
particles together to form the matter familiar to us.
• Other forces also exist that hold the particles of some
types of matter together. These include:
•metallic bonding,
•dipolar forces,
•hydrogen bonding,
•dispersion forces.
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TYPES OF MATERIALS
•Ionic compounds (e.g. NaCl) are held together by ionic
bonds, which are attractive forces that hold together ions of
opposite charge.
•Polar covalent compounds (e.g. H2O and CO) are held
together by dipolar forces, which are attractive forces that
exist between the positive end of one polar molecule and the
negative end of another.
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•Some polar covalent molecules (e.g. H2O) experience
hydrogen bonding, which is the result of attractive dipolar
forces between molecules in which hydrogen atoms are
covalently bonded to very electronegative elements (O, N, or
F).
•Network solids are solids in which the lattice sites are
occupied by atoms that are covalently bonded to each other
(e.g. SiO2 and diamond – pictured below).
48
•Metals (e.g. Cu) are held together by metallic bonds, which
originate from the attraction between positively charged
atomic kernels that occupy lattice sites and mobile electrons
that move freely through the lattice.
•Nonpolar covalent molecules (e.g. O2 and CO2 – shown
below) are only held together by dispersion forces, which are
very weak attractive forces acting between the particles of all
matter that result from momentary nonsymmetric electron
distributions in molecules or atoms.
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50
RELATIVE STRENGTHS OF INTERPARTICLE
FORCES
51