Title: Chemical Bonding
1Chemical Bonding
2- A chemical bond is a strong electrostatic force
of attraction between atoms in a molecule or
compound. - Bonding between atoms occurs because it creates a
more stable arrangement for the atoms.
3Lewis Symbols Dot Diagrams
- Convenient way to show the valence electrons
4Three types of bonding
- Metallic bonding results from the attraction
between metal atoms and the surrounding sea of
electrons - Ionic bonding results from the electrical
attraction between positive and negative ions. - Covalent bonding results from the sharing of
electron pairs between two atoms
5Ionic Bonding
- Many atoms transfer electrons and other atoms
accept electrons, creating cations (positive
metal ions) and anions (negative nonmetal ions). - The resulting ions are attracted to each other by
electrostatic force.
6Ionic Bonding
- The ions closely pack together in a crystal
lattice. - This arrangement maximizes the attractive forces
among cations and anions while minimizing
repulsive forces.
7- Because force is proportional to the charge on
each ion, larger charges lead to stronger
interactions. - Because force is inversely proportional to the
square of the distance between the centers of the
ions, smaller ions lead to stronger interactions.
8Ionic bonding between Na and Cl
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10Covalent bonding
- In many cases electrons do not completely
transfer from one atom to another. - The electrons between atoms are shared.
11Covalent bonding between H2
- Hydrogens electron configuration is 1s1
- Because both H atoms need 1 more electron to
become isoelectronic with He, it is unlikely that
either will give up an electron. -
12Covalent bonding between H2
1s
?
1s
They share the two electrons. H H ?
H H
13Types of Covalent Bonds
- When electrons are shared equally the bond is
called a NONPOLAR covalent bond. (i.e. H2) - Sometimes the electrons between two atoms are NOT
shared equally. The bond created is called a
POLAR covalent bond. - . . . .
- H Cl ? HCl
- . . . .
14Polar Covalent Bonding
- An example of this would be HCl.
HCl molecule Hydrogen atom
? 1s
Ne ?? ?? ?? ? 3s
3p
Chlorine atom
15How to classify bond types
- Electronegativity ability of an atom in a
molecule to attract shared electrons to it - Each element on the periodic table is assigned an
electronegativity value (see page 353) that
ranges from 0.7 to 4.0. - The difference in the electronegativity
determines the bonding type (ionic, polar
covalent, or nonpolar covalent).
16Electronegativity Values
17If the electronegativity difference is
- 1.7 and higher ionic
- 0.3 to 1.7 polar covalent
- 0.0 to 0.3 nonpolar covalent
18What if I get an electronegativity difference
that is 0.3 or 1.7?
- These cut-off numbers are guidelines.
- It is a gradual change not stair-step.
19Ionic Character
- As the electronegativity difference increases,
the ionic character increases as well!
20Practice Problems
- What type of bond will occur between iodine and
the following elements cesium, iron, and sulfur?
Bonding between I and Electronegativity difference Bond Type
Cesium
Iron
Sulfur
21Determine the type of bond between the following
pairs.
Bonding between Electronegativity difference Bond type
Li Cl
S O
Ca Br
P H
Si Cl
S Br
22Other ways to determine bonding
- Electronegativity is not the only factor in
determining bonding. - Generally, bonds between a metal and nonmetal are
ionic, and between two nonmetals the bonds are
covalent. - Examination of the properties of a compound is
the best way to determine the type of bonding.
23Ionic Bonding
- Ionic compounds are formed to maximize stability.
- Nonmetal - will gain electrons to become
isoelectronic with nearest noble gas called an
anion - Metal will lose electrons to become
isoelectronic with noble gas called a cation
24Transition Metals
- Zinc
- Electron configuration is 1s22s22p63s23p64s23d10
- When it forms the 2 ion, it loses the 2 valence
electrons in the 4s sublevel. - Zn2 configuration is 1s22s22p63s23p63d10
25Practice Problems
- Write the electron configurations for the
following ions. - Fe2
- S2-
- Mg2
- Use electron configurations to explain why the
most probable charge on the strontium ion is 2.
26Size of Ions
Does the size go up or down when losing an
electron to form a cation?
27Size of Ions
Forming a cation.
Li,152 pm
3e and 3p
- CATIONS are SMALLER than the atoms from which
they come. - The electron/proton attraction has gone UP and so
size DECREASES.
28Size of Ions
- Does the size go up or down when gaining an
electron to form an anion?
29Size of Ions
Forming an anion.
- ANIONS are LARGER than the atoms from which they
come. - The electron/proton attraction has gone DOWN and
so size INCREASES. - Trends in ion sizes are the same as atom sizes.
30Trends in Ion Sizes
Figure 8.13
31Which is Bigger?
- Cl or Cl- ?
- K or K ?
- Ca or Ca2 ?
- I- or Br- ?
32Which is Bigger?
- Cl or Cl- ? Cl-
- K or K ? K
- Ca or Ca2 ? Ca
- I- or Br- ? I-
33Lattice Energy Effects
- The change in energy when separated gaseous ions
are packed together to form an ionic solid. - M(g) X-(g) ? MX(s)
- Lattice energy is negative (exothermic) from the
point of view of the system.
34Lattice Energy
- To determine which compound will have the highest
lattice energy, take into consideration the
following - The size of the ions in the compound
- The smaller the size, the greater the lattice
energy - The charge of the ions in the compound
- The greater the charge, the greater the lattice
energy
35Calculating ?Hf
- We can take advantage of the fact the energy is a
state function and break the reaction into steps,
the sum of which is the overall reaction. - Lets do 41 Na(s) ½ Cl2 (g) ? NaCl(s)
- Given the following
- Lattice energy -786 kJ/mol
- Ionization energy for Na 495 kJ/mol
- Electron affinity for Cl -349 kJ/mol
- Bond energy of Cl2 239 kJ/mol
- Enthalpy sublimation for Na 109 kJ/mol
36Process
- Step 1 Sublimation of Na
- Na(s) ? Na(g) 109 kJ/mol
- Step 2 Ionization of Na
- Na (g) ? Na (g) e- 495 kJ/mol
- Step 3 Dissociation of Cl2
- ½ Cl2 (g) ? Cl(g) 119.5 kJ/mol
- Step 4 Formation of Cl- (Electron Affinity)
- Cl (g) e- ? Cl-(g) -349 kJ/mol
- Step 5 Formation of NaCl
- Na(g) Cl-(g) ? NaCl(s) -786 kJ/mol
Na(s) ½ Cl2 (g) ? NaCl(s) -411.5 kJ/mol
37Drawing Lewis Structures Valence Electron Review
- Valence electrons are in outermost level
- You can use periodic table or electron
configuration to determine valence electrons - Example Phosphorus
- Located in Group 15 or 5A
- Electron configuration is 1s22s22p63s23p3
- Contains 5 valence electrons
- Complete Exercise 1 on worksheet
38Drawing Lewis StructuresOctet Rule
- Most useful rule for creating Lewis structures
- Every atom usually has 8 valence electrons
- Exception hydrogen is good with 2 (like He)
- Lines are used to link atoms together (same as
using 2 dots)
Same as
39Steps to Drawing Lewis Structures
- Count valence electrons.
- Connect atoms together with bonds. In molecules
with a single atom of one element and several
atoms of another element, the single atom is
generally in the center with the other atoms
attached to it. - Add electrons around outside of atoms to give
each atom 8 electrons (or 2 in the case of
hydrogen). - Count electrons used. This number must be the
same as valence electrons.
40PCl3
5(37)26 e-
Complete Exercise 2.
Bonding Pairs
Lone Pairs (a.k.a. nonbonding electrons)
41Helpful Hints
- Carbon atoms form 4 bonds.
- Nitrogen atoms form 3 bonds.
- Oxygen atoms form 2 bonds.
- Hydrogen atoms form 1 bond.
- Fluorine atoms form 1 bond.
- Other halogens (Cl, Br, and I) frequently form 1
bond (but not always).
42Determining the Central Atom
- In a molecule, the atom that typically forms the
greatest number of bonds is in the center, with
other atoms attached to it. - Example CH3Cl
- Carbon forms 4 bonds
- Hydrogen forms 1 bond
- Chlorine forms 1 bond
- SO CARBON IS IN THE MIDDLE WITH HYDROGEN AND
CHLORINE AROUND IT! Dont forget electrons on
chlorine to make 8! - Complete exercise 3.
43Covalent Bonding
- Multiple Bonds
- It is possible for more than one pair of
electrons to be shared between two atoms
(multiple bonds) - One shared pair of electrons single bond (e.g.
H2) - Two shared pairs of electrons double bond (e.g.
O2) - Three shared pairs of electrons triple bond
(e.g. N2). - Generally, bond distances shorten with multiple
bonding.
Octet in each case
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45Resonance
- Occurs when more than one valid Lewis structure
can be written for a particular molecule.
46Odd Number of Electrons
NO
Number of valence electrons 11
Resonance occurs when more than one valid Lewis
structure can be written for a particular
molecule (i.e. rearrange electrons)
NO2
Number of valence electrons 17
Molecules and atoms which are neutral (contain no
formal charge) and with an unpaired electron are
called Radicals
47Beyond the Octet
- Elements in the 3rd period or higher can have
more than an octet if needed. - Atoms of these elements have valence d orbitals,
which allow them to accommodate more than eight
electrons.
48More than an Octet
Elements from the 3rd period and beyond, have ns,
np and unfilled nd orbitals which can be used in
bonding
P (Ne) 3s2 3p3 3d0 Number of valence electrons
5 (5 x 7) 40
PCl5
S (Ne) 3s2 3p4 3d0 Number of valence electrons
6 (4 x 7) 34
SF4
The Larger the central atom, the more atoms you
can bond to it usually small atoms such as F,
Cl and O allow central atoms such as P and S to
expand their valency.
49Formal Charge
- Difference between the of valence electrons in
the free atom and the of electrons assigned to
that atom in the Lewis structure. - FC formal charge G.N. Group Number
- BE bonding electrons LPE lone pair
electrons - If Step 4 leads to a positive formal charge on an
inner atom beyond the second row, shift electrons
to make double or triple bonds to minimize formal
charge, even if this gives an inner atom with
more than an octet of electrons.
50Formal Charge
51Molecular Shapes
- Lewis structures give atomic connectivity they
tell us which atoms are physically connected
together. They do not tell us the shape. - The shape of a molecule is determined by its bond
angles. - Consider CCl4 experimentally we find all Cl-C-Cl
bond angles are 109.5?. - Therefore, the molecule cannot be planar.
- All Cl atoms are located at the vertices of a
tetrahedron with the C at its center.
52Molecular Shape of CCl4
53VSEPR Theory
- In order to predict molecular shape, we assume
the valence electrons repel each other.
Therefore, the molecule adopts whichever 3D
geometry minimized this repulsion. - We call this process Valence Shell Electron Pair
Repulsion (VSEPR) theory.
54Why is VSEPR Theory Important?
- Gives a specific shape due to the number of
bonded and non-bonded electron pairs in a
molecule - Tells us the actual 3-D structure of a molecule
- In bonding, electron pairs want to be as far away
from each other as possible.
55VSEPR and Resulting Geometries
56How does VSEPR THEORY work?
- We can use VSEPR theory using 4 steps.
- Draw the Lewis Structure for the molecule.
- Example SiF4
57How does VSEPR THEORY work?
- We can use VSEPR theory using 4 steps
- Draw the Lewis Structure for the molecule.
- Tally the number of bonding pairs and lone
(non-bonding) pairs on the center atom.
Bonding pairs 4 Lone pairs on central atom 0
58How does VSEPR THEORY work?
- We can use VSEPR theory using 4 steps
- Draw the Lewis Structure for the molecule
- Tally the number of bonding pairs and lone pairs
on the center atom. - Arrange the rest of the atoms so that they are as
far away from each other as possible.
59How does VSEPR THEORY work?
- We can use VSEPR theory using 4 steps
- Draw the Lewis Structure for the molecule
- Tally the number of bonding pairs and lone pairs
on the center atom. - Arrange the rest of the atoms so that they are as
far away from each other as possible - Give the type of geometry the molecule has
Tetrahedral
60Another Example
To determine the electron pair geometry 1) draw
the Lewis structure 2) count the total number
of electron pairs around the central
atom. 3) arrange the electron pairs in one of
the geometries to minimize e--e-
repulsion. 4) multiple bonds count as one
bonding pair for VSEPR
61The VSEPR Model
Predicting Molecular Geometries
62The VSEPR Model
Predicting Molecular Geometries
63The VSEPR Model
Difference between geometry and
shape Geometry We determine the geometry only
looking at electrons. All the atoms that obey the
octet rule have the same tetrahedral-like
geometry. Shape We name the shape by the
positions of atoms. We ignore lone pairs in the
shape.
64The VSEPR Model
Predicting Shape
Shape
65The VSEPR Model
Predicting Shape
Shape
66The VSEPR Model
The Effect of Nonbonding Electrons and Multiple
Bonds on Bond Angles By experiment, the H-X-H
bond angle decreases on moving from C to N to
O Since electrons in a bond are attracted by
two nuclei, they do not repel as much as lone
pairs. Therefore, the bond angle decreases as the
number of lone pairs increase.
67The VSEPR Model
The Effect of Nonbonding Electrons and Multiple
Bonds on Bond Angles Similarly, electrons in
multiple bonds repel more than electrons in
single bonds.
68The VSEPR Model
Molecules with Expanded Valence Shells Atoms that
have expanded octets have AB5 (trigonal
bipyramidal) or AB6 (octahedral) electron pair
geometries. Examples PF5 trigonal
bipyramidal SCl6 octahedral
69The VSEPR Model
Molecules with Expanded Valence Shells
70The VSEPR Model
Molecules with Expanded Valence Shells
71The VSEPR Model
Molecules with More than One Central Atom In
acetic acid, CH3COOH, there are three central
atoms. We assign the geometry about each central
atom separately.
72Hybrid Orbitals
- In bonding, s and p orbitals are used in bonding.
It is easy to tell which ones are used by
looking at our molecule. - For example, CH4. Looking again at the Lewis
structure, we see that there are 4 bonds. We
call this sp3 hybridized.
73Hybrid Orbitals
- Regions of electron density-EACH BOND AND LONE
PAIR OF ELECTRONS ON THE CENTRAL ATOM IS KNOWN AS
A REGION OF ELECTRON DENSITY. - 2 regions of electron density-sp hybridized
- 3 regions of electron density-sp2 hybridized
- 4 regions of electron density-sp3 hybridized
74Hybridization
sp Hybrid Orbitals The two lobes of an sp hybrid
orbital are 180? apart.
75Hybrid Orbitals
sp2 Hybrid Orbitals Important when we mix n
atomic orbitals we must get n hybrid
orbitals. sp2 hybrid orbitals are formed with one
s and two p orbitals. (Therefore, there is one
unhybridized p orbital remaining.) The large
lobes of sp2 hybrids lie in a trigonal plane. All
molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central atom.
76Hybridization
77Hybridization
sp3 Hybrid Orbitals sp3 Hybrid orbitals are
formed from one s and three p orbitals.
Therefore, there are four large lobes. Each lobe
points towards the vertex of a tetrahedron. The
angle between the large lobes is 109.5? All
molecules with tetrahedral electron pair
geometries are sp3 hybridized.
78Hybridization
79Hybrid Orbitals
80Hybrid Orbitals
- Summary
- To assign hybridization
- Draw a Lewis structure.
- Assign the geometry using VSEPR theory.
- Use the geometry to determine the hybridization.
- Name the shape by the positions of the atoms.
81Hybridization and Multiple Bonds
- Multiple bonds overlap differently and are
called s-bonds and p-bonds - All single bonds are s
- Double bonds contain 1 s and 1 p bond
- Triple bonds contain 1 s and 2 p bonds
82Bond Energy
83Covalent Bonding Orbital Overlap
- As two nuclei approach each other their atomic
orbitals overlap. - As the amount of overlap increases, the energy of
the interaction decreases. - At some distance the minimum energy is reached.
- The minimum energy corresponds to the bonding
distance (or bond length).
84Covalent Bonding Orbital Overlap
- As the two atoms get closer, their nuclei begin
to repel and the energy increases. - At the bonding distance, the attractive forces
between nuclei and electrons just balance the
repulsive forces (nucleus-nucleus,
electron-electron).
85Bond Energies
- Bond breaking requires energy (endothermic).
- Bond formation releases energy (exothermic).
- ?H ?D(bonds broken) ? ?D(bonds formed)
86Bond Energies