Transcript Bonding

Bonding
Forces of attraction that
hold atoms together
making compounds
Chemical symbols



Symbols are used to represent elements
Either one capital letter, or a capital letter
with a lower case letter
Know names and symbols of elements:
– 1 – 30, plus
–Rb, Cs, Sr, Ba, Ag, Au, Cd,
Hg, Pt, Ga, Ge, As, Sn, Pb,
Se, Br, I, and U
Basic idea...
 All
chemical bonds form
because they impart
stability to the atoms
involved
 lower energy = greater
stability
Quick review
 All
types of chemical bonds
involve electrons
 Valence electrons, the electrons
in the outermost occupied
energy level of an atom, are
usually the electrons involved in
bonding
 The
representative elements
have the same number of
valence electrons as their family
number in the American system
–Example: Mg, column IIA, 2
valence electrons
 The
transition metals all have
two valence electrons
ns2(n-1)dx
 Lewis
dot structures are used to
represent the valence electrons
–each dot represents a valence
electron
.
–no more than 8 dots total
–no more than 2 dots on a side
.
–example = Mg: Na
Lewis dot structures of representative
elements
The Octet Rule
 Atoms
will gain, lose, or
share electrons in order to
2
6
achieve an ns np valence
configuration
Sizes of atoms

Periodic trend: atomic radii increase
moving down a group
– Increasing energy level

Periodic trend: atomic radii decrease
moving left to right in a period
– The charge felt by the valence electrons
becomes larger
Sizes of atoms
• There is a general decrease in atomic
radius from left to right, caused by
increasing positive charge in the
nucleus.
• Valence electrons are not shielded from
the increasing nuclear charge because
no additional electrons come between
the nucleus and the valence electrons.
• For metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the
element.
• For elements that occur as molecules, the
atomic radius is half the distance between
nuclei of identical atoms.
Atomic Radius
Atomic Radius
• Atomic radius generally increases as you
move down a group.
• The outermost orbital size increases down a
group, making the atom larger.
Sizes of ions

Periodic trend: anions are always larger
than the atom they were formed from
– Electrons repel each other

Periodic trend: cations are always
smaller than the atom they were formed
from
– Fewer electrons to share same positive
nuclear charge
Ionic Radius
• When atoms lose electrons and form
positively charged ions, they always
become smaller for two reasons:
1. The loss of a valence electron can leave an
empty outer orbital resulting in a small
radius.
2. Electrostatic repulsion decreases allowing
the electrons to be pulled closer to the
radius.
Ionic Radius
• When atoms gain electrons, they can
become larger, because the addition
of an electron increases electrostatic
repulsion.
Ionic Radius
• Both positive and negative ions increase in
size moving down a group.
Ionic Radius
• The ionic radii of positive ions generally
decrease from left to right.
• The ionic radii of negative ions generally
decrease from left to right, beginning with
group 15 or 16.
Bonding
Forces of attraction that
hold atoms together
making compounds
Ionization energy
 The
energy needed to remove a
valence electron from an atom
 A measure of how tightly the
electrons are being held
 periodic trend
–increases from the bottom up
–increases left to right
 In
general, metals have
lower IE than nonmetals
–alkali metals are the lowest IE
family
–noble gases are highest IE
family
Ionization energy
• The energy required to remove the first
electron is called the first ionization
energy.
• First ionization energy increases from
left to right across a period.
• First ionization energy decreases down
a group because atomic size increases
and less energy is required to remove
an electron farther from the nucleus.
Ionization energy
Ionization energy
Ionization energy
• Removing the second electron
requires more energy, and is called
the second ionization energy.
• Each successive ionization requires
more energy, but it is not a steady
increase.
• The ionization at which the large
increase in energy occurs is related to
the number of valence electrons.
Ionization energy
Electron affinity
A
measure of how strongly an
element would like to gain an
electron
 periodic trend
–increases from the bottom up
–increases left to right
–ignore the noble gases

Atoms that lose electrons easily
have little attraction for additional
electrons (and vice versa)
– metals have low IE, low EA
– Nonmetals have high IE, high EA


Octet rule: when atoms react, they tend
to strive to achieve a configuration
having 8 valence electrons
This results in some form of bond
formation
Periodic trends…
As you move from left to right along
a period…
 Atoms get
….
Smaller
 Ionization energy goes
….
Up
 Electron affinity goes
….
Up

Periodic trends…
As you move down a group/family
 Atoms get
….
Larger
 Ionization energy goes
….
Down
 Electron affinity goes
….
Down

Check your understanding
The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth
Check your understanding
The ionic radius of a negative ion
becomes larger when:
A. moving up a group
B. moving right to left across period
C. moving down a group
D. the ion loses electrons
Electron Configuration of Ions

Na 1s22s22p63s1
– will lose one e- to gain ns2np6
configuration
– Na+  1s22s22p6

S 1s22s22p63s23p4
– will gain 2 e- to gain ns2np6
configuration
– S2-  1s22s22p63s23p6
Ionic Bonding
 Metals
lose electrons easily,
nonmetals have a strong
attraction for more electrons
 metal atoms will lose electrons
to nonmetal atoms, causing
both to become ions
1.
2.
3.
Metals, having lost one or more
electrons, become cations (+)
Nonmetals, having gained one or
more electrons, become anions (-)
Opposites attract: the cations and
anions are held together
electrostaticly
– called “ionic bonds”
In summary...
 Ionic
bonds are electrostatic
attractions between cations
and anions formed when
electron(s) are transferred
from the low IE, EA metal to
the high IE, EA nonmetal
Ionic compound = crystalline solid
Cation
(+)
Ionic Compounds
 High
melting points
 brittle solids
 nonconducting as solids
 conduct electricity as liquids
or aqueous
Ionic Compounds

As solids, exist in a 3-D repeating pattern
called a crystal “lattice”

the lattice energy is the energy lowering
(stability) accomplished by the formation from
“free” ions

Also a measure of the energy required to
break apart the ionic compound once formed

The greater the lattice energy, the stronger
the force of attraction
Bonding
Forces of attraction that
hold atoms together
making compounds
Ion dissociation
 Many
ionic compounds will
dissolve in water if it results in
more stability (lower E) than in
the solid ionic compound
 the ions “dissociate” from each
other

Ex: CaCl2(s) + H2O  Ca2+(aq) + 2Cl-(aq)
Ionic Bond Strength
A
measure of the attractive
force between the ions
smaller atoms = stronger ionic
bonds
 fewer atom ratio = stronger bond

 evidence:
melting points
Compare the melting points:
 KCl
o
776 C
:
 KI : 723oC
 smaller atoms result in
stronger ionic bonds
Compare the melting points:
 CaCl2
o
772 C
:
 NaCl : 800oC
 fewer atoms result in
stronger ionic bonds
Bonding
Forces of attraction that
hold atoms together
making compounds
Covalent Bonding
 Covalent
bonding involves
the sharing of electron pairs
 usually between two high
EA, high IE nonmetals
–both want more e-’s, neither is
willing to lose the e-’s they
have
A
nonmetal will form as many
covalent bonds as necessary to
fulfill the octet rule
 example: C, with 4 valence e-’s,
will form 4 covalent bonds
–results in 8 valence e-’s around
the carbon atom at least part of
the time
 double
and triple covalent
bonding is a possibility
When does the
octet rule fail?
H, He and Li

Helium strives for 2 valence
electrons
– 1s2 configuration
Hydrogen will sometimes will share
its one electron with another atom,
forming a single covalent bond
 Lithium will lose its lone valence
electron, gaining the 1s2
configuration of He

Be
 Be
will sometimes lose its 2
valence electrons, gaining the
Is2 configuration of He
 Be will sometimes form 2
covalent bonds, giving it 4
valence electrons
–nuclear charge of +4 cannot
handle 8 valence electrons
B
 Boron
will often make three
covalent bonds using its three
valence electrons
–nuclear charge of +5 cannot
handle 8 valence electrons in a
stable manner
“organometallic”
compounds
 Some
metals will form covalent
compounds with nonmetals
–Hg, Ga, Sn, and others
 The
octet rule is not followed for
the metals,but is for nonmetals
 Form 2 or more covalent bonds
P, S, Cl, Se, Br, I
 Elements
in the third period and
lower have empty d orbitals
 there is room for more than 8
valence electrons
 These elements will at times
make more than 4 covalent
bonds
Rules for Drawing structural
formulas



1) Determine the central atom, place
the other atoms evenly spaced around
the outside
2) Count the total number of valence
electrons
3) Draw single bonds between the
central atoms and each of the outside
atoms



4) Complete the octet on the outside
atoms by placing electrons in pairs
around the outside atoms (lone pairs)
5) Place any remaining electrons on the
central atom in pairs
6) If the central atom does not have its
minimum number of electrons
(usually 8), form double bonds by
moving lone pairs off of the outside
atoms and drawing them as bonding
pairs
Binary Molecular
Nomenclature
 Two
nonmetals
 no charges to balance
 multiple subscripts possible
–ex: N2O, NO, NO2, N2O4,
N2O5
Use prefixes to represent
subscripts
 mono
 di
=1
=2
 tri = 3
 tetra = 4
 penta = 5
 Hexa
=6
 hepta = 7
 octa = 8
 nona = 9
 deca = 10
Rules, continued..
 Change
second name to end
in “ide”
 do not use prefixes on the
first word if the prefix is
“mono”
 always use prefixes on the
second name
Examples...
 CO2
 carbon
= first word
 subscript = 1, so no prefix
 oxide = second word
 subscript = 2, so prefix = di
 carbon dioxide
Examples...
 CO
 carbon
= first word
 subscript = 1, so no prefix
 oxide = second word
 subscript
 carbon
= 1, so prefix = mono
monoxide
Examples...
 SF6
1
sulfur, 6 fluorines
 sulfur hexafluoride
 P2O5
2
phosphorus, 5 oxygens
 diphosphorus
pentoxide
Examples...
 Dinitrogen
tetroxide
 di = 2, so two nitrogen’s
 tetra = 4, so 4 oxygens
 N2O4
 Dihydrogen
 H2O!
 DHMO.org
monoxide