Bonding and
Periodic Table
Trends
Honors Chemistry
Electron Configurations
• Stable Octet: 8 electrons in the outer level
is very stable (includes He)
• Ions – gain/lose electrons to achieve a
stable octet
• Isoelectronic – same electron configuration
• Examples: N, O, F, Na, Mg, Al are isoelectronic
with Ne – this is called an isoelectronic series
• Pseudoisoelectronic – same electron
configuration but includes the d orbitals
• Fe+2 is pseudoisoelectronic with Ar
Periodic Table Trends
Introduction
•
Properties are periodic
•
An element’s position & its properties are a result
of its electrons
•
The outermost electrons, aka valence electrons,
have the greatest influence on the properties of
the elements.
•
Adding an electron to an inner core orbital
results in less striking changes in properties than
adding an electron to an outer valence orbital
(higher energy).
•
Shielding Effect: electrons in the lower energy
levels (inner core electrons), shield electrons in
the outer levels from the full effect of the nuclear
charge.
Trends in the Periodic Table
A. Atomic Radius
1. The distance from the
center of the nucleus to the
outermost electron.
2. Bond Radius
3. Atoms get larger going
down a group and
smaller going across a
period.
Ex) Na is larger than Mg
Na is smaller than K
Ga vs. Al
Atomic Radii of the
Representative Elements
Atomic Radii vs Atomic Number
Positive Ion Size
• When atoms lose
electrons, they
become (positive) and
get smaller.
• The sizes of cations
increases down a
group.
• The sizes of cations
decreases across a
period.
Negative Ionic Size
•
•
•
When atoms gain
electrons, they
become (negative)
and get larger.
The sizes of anions
increases down a
group.
The sizes of anions
decreases across a
period.
Relative Sizes of Positive &
Negative Ions
The sodium ion lost an electron, and
therefore the positive protons in the
nucleus exert a stronger pull on the
remaining negative electrons,
shrinking the orbitals. Thus positive
ions are smaller than their atoms.
The chloride ion gained an electron, and
therefore the fewer positive protons in
the nucleus exert a weaker pull on the
extra negative electrons, increasing the
size of the orbitals. Thus negative ions
are larger than their atoms.
•
Within an isoelectronic series, radii
decrease with increasing atomic
number because of increasing
nuclear charge.
N-3 > O-2 > F-1 > Na+1 > Mg+2 > Al+3
How many electrons? Nuclear charge?
Ionic Radius
• Cations &
anions decrease
in size going
across a period
• Cations &
anions increase
in size going
down a group
Electron Attraction in a Bond &
Ion Size
Ionization Energy (IE):
1. The energy needed to remove one
electron from an atom. (kJ/mole)
2. IE measures how tightly electrons are
bound to an atom.
–
–
Elements that do not want to lose their
electrons have high ionization energies.
Elements that easily lose electrons have
low ionization energies.
X + energy  X+1 + 1 e-
1st Ionization Energy
3. I.E. decreases down a group.
4. I.E. increases across a period
–
Account for deviations across a period.
5. Metals tend to have low IE1.
6. Nonmetals tend to have high IE1.
Ionization Energy of the
1st 20 Elements
Successive Ionization Energies:
1. Energy required to remove electrons beyond
the 1st electron.
2. Ionization energies will increase for every
electron removed.
X + IE1  X+1 + 1 eX+1 + IE2  X+2 + 1 eX+2 + IE3  X+3 + 1 e-
Successive Ionization Energies:
kJ/mol
IE1
IE2
IE3
IE4
IE5
Na
496
4562
6912
9544
13 353
Mg
738
1451
7733
10 540
13 628
Al
578
1817
2745
11 578
14 831


Electron Configuration
Na: [Ne] 3s1
Mg: [Ne] 3s2
Al: [Ne] 3s23p1
3s
3s
 
3s
3p
Ionization Energy vs. Atomic Number
Notice the dips across the period… why?
Period 3
Na - [Ne] 3s1
Mg - [Ne] 3s2
Al - [Ne] 3s23p1
Si - [Ne] 3s23p2
P - [Ne] 3s23p3
S - [Ne] 3s23p4
Cl - [Ne] 3s23p5
Ar - [Ne] 3s23p6
___
___
___
___
___
___
___
___
3s
___ ___ ___
___ ___ ___
___ ___ ___
___ ___ ___
___ ___ ___
___ ___ ___
___ ___ ___
___ ___ ___
3p
Electronegativity (EN)
1. Reflects an atoms ability to attract electrons
in a chemical bond.
• Up to 4.0 for F
• Zero for He, Ne, Ar and Kr
2.
3.
4.
5.
Metals have low EN.
Nonmetals have high EN.
EN decreases down a group.
EN increases across a period.
Electron Affinity (EA)
1.
•
Energy change that occurs when a neutral gaseous atom
gains an electron. Units kJ/mol.
1 e- + X  X-1 + EA
Most elements have no affinity for an additional electron
and have an EA equal to zero.
He(g) + e-  HeEA = 0 kJ/mol
He will not add an electron
Cl(g) + e-  Cl- + 349 kJ/mol
EA = -349 kJ/mol
Exothermic!!!
Electron Affinity (EA)
2.
3.
4.
5.
Metals have low EA.
Nonmetals have high EA.
EA decreases down a group.
EA increases (becomes more negative)
across a period.
•
•
EXCLUDES noble gases
Exceptions: Groups IIA (~0) and VA (~0 for
N and smallerfor P to Bi)
•
Why? Filled s and half filled p
Metallic Character
1. Reflected by those elements that can
lose electrons easily.
2. Increases down a group.
3. Decreases across a period.
4. The most metallic metal is Cesium.
5. The most nonmetallic (least metallic)
metal is Aluminum.
Reactivity
•
Related to the ability of an element to lose
or gain an electron
Brainiac Alkali Metals
1. Reactivity
group.
2. Reactivity
period.
3. Reactivity
a group.
4. Reactivity
a period.
Alkali Metals Reactivity
of metals INCREASE down a
of metals DECREASE across a
of nonmetals DECREASE down
of nonmetals INCREASE across
8
Chemical Bonding
Chapter 6
Honors Chemistry
Introduction
• A chemical bond is an attractive force that holds atoms
together in elements or compounds (to function as a unit)
– intramolecular (within) vs. intermolecular (between)
• Bond energy is the energy needed to break or form 1 mole
of bonds in a gaseous substance (kJ/mol)
• Bonding usually involves only the valence electrons.
– In most compounds of the representative elements, the atoms have
an electron configuration that is isoelectronic or psuedoisoelectronic
with a noble gas
• The manner in which atoms are bound together in a given
substance has a profound effect on its chemical and
physical properties.
Octet Rule
• Many chemical compounds form such that
each atom, by gaining, losing, or sharing
electrons, has an octet of electrons in
its highest occupied energy level.
– Very stable
– There are exceptions: H, He, B (BF3)
• System achieves the lowest possible
energy.
Types of Chemical Bonds:
Metallic Bonds
Ionic Bonds
Covalent Bonds
Metallic Bonds
• Simplest crystalline solid – arranged in a very compact
and orderly pattern
• Sea of electrons – the valence electrons are mobile
around metal cations
– Electrons are delocalized
• Attraction of the metal atoms and the surrounding sea of
electrons
Metallic Bonds
• Explains metallic properties:
– High electrical and thermal conductivity (flow
of electrons)
– Luster (metals absorb wide range of  - excites
e- and fall back emitting E in form of light 
results in shiny appearance)
– Ductility & malleability (mobility of e-, metallic
bonding is same in all directions throughout
solid)
Metallic bonding visual on Holt
Metals vs. Ionic Crystals
• Metallic properties
due to sea of electrons
• Ionic compounds are
hard but brittle –
repulsions result from
shift and causes crystal
to break
Ionic bonding
•
•
•
Chemical bonding that results from an
electrostatic attraction between cations and
anions to form a neutral compound.
“salts”
Octet Rule
a) Atoms will transfer electrons (e-) to each other in
order to have a full set of valence electrons.
b) When electrons are transferred, ionic bonds are
formed.
Ionic bonding
Covalent Bonding
• Sharing one or more electron pairs
between 2 atoms
Characteristics of Ionic Compounds
(hundreds of compounds)
1. All are high melting solids
(>400°C).
a) Orderly 3D arrangements
(pattern) called crystalline solid
or crystal lattice.
b) Simplest arrangement =
formula unit
c) High mp reflects strong bonds –
large attractive forces are very
stable
d) Many are white.
e) Colored compounds usually
contain the transition elements
(Cu, Cr, Co, Ni, Mn)
Characteristics of Ionic Compounds:
2. Solubility
a) Many are soluble in polar solvents, such as water (aka
aqueous solutions)
b) Most are insoluble in nonpolar solvents, such as
hexane (C6H14)
Dissolving Salt Animation
Characteristics of Ionic Compounds:
3. Conductivity
a) Solids are non
conductive – ions
cannot move freely
b) Molten compounds
are conductive – ions
move freely (NaCl
mp ~800°C)
c) Aqueous solutions
are conductive – ions
free to move in
solution
Formation of Ionic Compounds
• Formation results from a transfer of
electrons and the electrostatic attractions of
the closely packed, oppositely charged
ions.
• Ionic substances are formed when an
atom that loses electrons relatively easily
reacts with an atoms that has a high
affinity for electrons.
• Forms between a metal and a nonmetal.
(large EN difference)
– Metal: low IE, EN, EA
– Nonmetal: high IE, EN, EA
– Metal is oxidized (loss of e-) and nonmetal is
reduced (gain of e-)
• The ion pair has lower energy than
separated ions.
Electron Configuration
Distribution of electron density
• Na: 1s22s22p63s1
– 186 pm
• Cl: 1s22s22p63s23p5
– 99 pm
• Na+1: 1s22s22p6
– 95 pm
• Cl-1: 1s22s22p63s23p6
– 181 pm
Lewis structures or
electron-dot structures
Lewis structures examples:
• Sodium and chlorine
• Potassium and phosphorus
Ionic, Nonpolar Covalent,
Polar Covalent
Character of Bonds
EN
Difference
0.00
0.65
0.94
1.19
1.43
1.67
1.91
2.19
2.54
3.03
% Ionic
Character
0%
10%
20%
30%
40%
50%
60%
70%
80%
90%
100%
90%
80%
70%
60%
50%
40%
30%
20%
10%
% Covalent
Character
HOLT
Nonpolar Covalent Bonds
• Electron pair is shared equally between the atoms
(ΔEN = 0 to ~0.4)
– Diatomic molecules (H O F Br I N Cl); allotropes (S8)
Nonpolar Covalent Bonds
Hydrogen
Polar Covalent Bonds
• Electron pair is shared unequally between
atoms (ΔEN = ~0.4 to ~1.9)
• Results in an electric dipole (2 poles)
• Equal but opposite charges that are separated by a
short distance
• Separation of charge between 2 covalently bonded
atoms
• Examples: HF, HBr, H2O
Polar and Nonpolar Bonds
Characteristics of Covalent Compounds
(~11 million compounds)
1. They are gases, liquids, or solids with
low melting points (<300°C)
2. Usually the simplest arrangement is
called a molecule (many sizes and
shapes)
Characteristics of Covalent Compounds
3. Solubility
– Many are insoluble in polar solvents
– Many are soluble in nonpolar solvents
4. Conductivity
– Liquid and molten compounds are
nonconductors
– Aqueous solutions are usually
nonconductors or poor conductors
Formation
1. Formation due to
sharing electrons
2. Often forms between
nonmetals
– metalloid and
nonmetals
– small electronegativity
differences
Potential Energy Changes
Potential Energy Changes
a) PE=0 separation of atoms do no affect
each other
b) PE decreases as atoms are drawn together
by attractive forces
c) PE is at MINIMUM when attractive forces
are balanced by repulsive forces = stable
d) PE is increasing when repulsion between
like charges outweighs attraction between
opposite charges
HOLT
Bond length & energy
• Bond Length = distance between 2
nuclei in most stable position (lowest
PE)
–1s orbitals overlap (head on) to form a
single covalent bond (sigma bond)
• Bond Energy = energy required to
break a bond
Formation
3. Number of covalent bonds likely to form for
nonmetals or metalloids depends upon the
number unpaired electrons
4. Electron configurations
Cl2:
1s22s22p63s23p23p23p1
1s22s22p63s23p23p23p1
1 bond: share 1 pair of e- = single bond
Electron config for
covalent compounds
O2
1s22s22p22p12p1
1s22s22p22p12p1
2 bonds: share 2 pairs of e- = double bond
7 Elements that can multiple bond:
C, N, O, Si, P, S, Se
Bond Energy for
Carbon Bonds
• Bond lengths (Å):
Single > double > triple
1.54
1.34
1.20
• Bond Energy (kJ/mol)
Single < double < triple
346
612
835
Multiple Bonds
Lewis Structures
Guidelines:
1. Select a reasonable skeleton for the molecule or
polyatomic ion. The central atom is the least
electronegative atom (excluding H)
2. Calculate the number of shared electrons (S):
S=N–A
– N = total number of valence electrons required (all 8,
except H)
– A = number of valence electrons available
3. Place shared pairs of electrons in skeleton
4. Place lone pairs (for octets)
5. NOTE: # of “dots” in the Lewis structure = A
Examples of Lewis Structures
for Covalent Bonds
• Cl2
• CO2
• ClO4-1
Structural formula (line structure) only
shows how the molecule or polyatomic ion
is bonded – NO “dots” shown
Coordinate Covalent Bonds
(Dative bonds)
• A bond formed when 1 atom provides
both electrons (to covalent bond)
• Perchlorate (ClO4-1) has 4 coordinate
covalent bonds
Equivalent Lewis Structures
• Also known as resonance structures
• CO3-2
Limitations of the Octet Rule
1. Less than an octet (electron deficient)
•
•
Be (4) & B (6)
BF3
2. More than an octet (expanded
valence shell or hypervalence)
•
•
8 elements: P, S, As, Br, Sb, Te, I, Xe
SF4
3. Odd number of electrons
•
NO
Overall, if you use N-A-S…
• If S leads to too many bonds
– 1st - look at central atom. Is it electron
deficient? (B, Be)
– 2nd - multiple bond
• If S lead to too few bonds
– Think hypervalence
Molecular Shapes
1. Valence shell electron pair repulsion theory
(VSEPR Theory): helps to predict the spatial
arrangement of atoms in a molecule or
polyatomic ion.
a) Introduction
i.
The central atom is any atom bonded to more than one
other atom.
ii. Unshared pairs (lone pair) of electrons and bonding pairs
on the central atom orient themselves to minimize
repulsions.
iii. Lone pairs of electrons occupy MORE space than
bonding pairs.
VSPER
b. Counting regions of high electron density
around the central atom.
i.
Each bonded atom is counted as ONE region of
high electron density, whether it is a single, a
double, or a triple bond.
ii. Each unshared pair of valence electrons on the
central atom is counted as ONE region of high
electron density.
Valence Bond (VB) Theory
• Describes HOW bonding occurs
• Usually atomic orbitals do not have the
correct energies or orientation to describe
where the electrons are when bonded to other
atoms
• Hybridization is the mixing of the atomic
orbitals to form new hybrid orbitals (s – p – d)
HOLT
VSPER
Regions of
High Electron Density
•
•
•
•
Two regions – LINEAR arrangement
2 regions e- density = sp
Bond angle = 180°
Example: BeH2
Animation of sp hybridization
Regions of
High Electron Density
• Three regions – TRIGONAL PLANAR
arrangement
• 3 regions e- density (all bonding) = sp2
• Bond angle = 120°
• Example: BF3
Animation of sp2 hybridization
Regions of
High Electron Density
• 3 regions e- density = sp2
– 2 bonding and 1 lone pair
• Electronic geometry – trigonal planar
• Molecular geometry – BENT or
ANGULAR
• Bond angle = 115°
• Example: NOCl
Regions of
High Electron Density
• Four regions – TETRAHEDRAL
arrangement
• 4 regions e- density (all bonding) = sp3
• Bond angle = 109.5°
• Example: CH4
sp3 hybrid orbital
Animation of sp hybridization
Regions of
High Electron Density
• 4 regions e- density = sp3
– 3 bonding and 1 lone pairs
• Electronic geometry – tetrahedral
• Molecular geometry – TRIGONAL
PYRAMIDAL
• Bond angle = 107.3°
• Example: NH3
Regions of
High Electron Density
• 4 regions e- density = sp3
– 2 bonding and 2 lone pairs
• Electronic geometry – tetrahedral
• Molecular geometry – BENT or
ANGULAR
• Bond angle = 104.5°
• Example: H2O
Molecular Polarity
Consider…
a) The presence of at least 1 polar bond
or 1 lone pair of electrons and
b) The molecular shape to determine
the overall molecular polarity
Examples:
HCl, BeCl2, BF3, CH4, NH3, H2O
Molecular Polarity
http://preparatorychemistry.com/Bishop_molecular_polarity.htm
Intermolecular Forces
(aka Van der Waals forces)
• Forces of attractions BETWEEN
molecules
– Intermolecular forces are weaker than
intramolecular forces
• Important for states of matter
• Boiling point is a measure of the
strength of these
– High bp – strong intermolecular forces
Intermolecular Forces
(aka Van der Waals forces)
1. London Dispersion Forces
•
•
•
Attractions caused by temporary
(instantaneous) dipoles and are
present between ALL atoms and
molecules
Increases with size
Holt
Boiling
Point
He
Ne
Ar
Kr
-269°
-246°
-186°
-153°
London Dispersion Forces:
induced dipole–induced dipole interactions
Intermolecular Forces
(aka Van der Waals forces)
2. Dipole-Dipole Forces
•
•
•
Boiling
Point
Attractions between polar molecules
Permanent dipoles
Holt
F2
HCl
BrF
CF4
CH3F
(nonpolar)
(polar)
(polar)
(nonpolar)
(polar)
-188°
-89°
-20°
-128°
-84 °
Intermolecular Forces
(aka Van der Waals forces)
3. Hydrogen Bonding
–
Attractions resulting when
hydrogen that is bonded to a
highly EN atom (F, N, O) is
attracted to an unshared pair of
electrons of an EN atom in a
nearby molecule
– H2S vs H2O
bp -59.6°
100°
– NH3 vs CH4
bp -33°
-161°
For each of the molecules below, list the
types of intermolecular forces which act
between pairs of these molecules.
(a)
(b)
(c)
(d)
(e)
CH4
PF3
CO2
HCN
HCOOH (methanoic acid)
website
Next: PF3
(a) CH4 is a tetrahedral molecule - it does not have a
permanent dipole moment.
–
•
•
•
The figure above shown CH4 in two views: one shows it as
it is commonly drawn, with one H at the top and three
H's at the bottom. The second figure shows CH4 rotated
to fit inside a cube. This might help to make clear why it
does not have a permanent dipole moment. The dipole
moments of the two C-H bonds pointing up exactly cancel
the dipole moments of the two C-H bonds pointing
downward.
CH4 does not contain N, O, or F and therefore
there are no hydrogen bonds between CH4
molecules.
Therefore only dispersion forces act between pairs
of CH4 molecules.
Other tetrahedral molecules like CF4 and CCl4 do
not have a permanent dipole moment.
Next: CO2
(b) PF3 is a trigonal pyramidal molecule (like
ammonia, the P has a single lone pair of
electrons); it does have a permanent dipole
moment.
• It does contain F, but it does not contain
any hydrogen atoms so there is no
possibility of forming hydrogen bonds.
• Therefore dispersion forces and dipoledipole forces act between pairs of PF3
molecules.
Next: HCN
(c) CO2 is a linear molecule; it does
NOT have a permanent dipole
moment
• It does contain O, however the oxygen
is not bonded to a hydrogen.
• Therefore only dispersion forces act
between pairs of CO2 molecules.
Next: HCOOH
(d) HCN is a linear molecule and it does
have a permanent dipole moment
• It does contain N, however the
nitrogen is not directly bonded to a
hydrogen.
• Therefore dispersion forces and
dipole-dipole forces act between pairs
of HCN molecules.
e) HCOOH is a non-linear molecule; it
does have a permanent dipole
moment;
• It does contain O, and the oxygen is
directly bonded to a hydrogen.
• Therefore dispersion forces, dipoledipole forces and hydrogen bonds act
between pairs of HCOOH molecules.
The intermolecular forces acting between
pairs of these molecules.
(a)
(b)
(c)
(d)
(e)
CH4
PF3
CO2
HCN
HCOOH
(methanoic acid)
(a)
(b)
(c)
(d)
(e)
dispersion
dispersion, dipole-dipole
dispersion
dispersion, dipole-dipole
dispersion, dipoledipole, hydrogen bonds