Chemical Bonding Chapter 6

Introduction to Bonding Chemical bond – mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together Q: Why do atoms bond? A: By bonding with each other, the system decreases in potential energy, thereby creating more stable arrangements of matter

Types of Chemical Bonding Ionic – chemical bonding that results from the electrical attraction between large numbers of cations and anions

Covalent – results from the sharing of electron pairs between two atoms

Types of Chemical Bonding Ionic or covalent? Determined by the difference in electronegativities of the two atoms

Types of Chemical Bonding

Types of Chemical Bonding Ionic or covalent?

Types of Chemical Bonding Atoms

S and H Ca and Cl C and F K and O

N and N

Difference in Electronegativity

Bond Type

Formation of a Covalent Bond Nature favors chemical bonding because most atoms are at a lower potential energy when bonded to other atoms than they are as independent particles

Formation of a Covalent Bond Bond length – distance between two bonded atoms at their minimum potential energy

Formation of a Covalent Bond Bond length – distance between two bonded atoms at their minimum potential energy Bond

Bond Length (pm)

H-H

74

Cl-Cl

199

Br-Br

228

I-I

267

C-C

154

C=C

134

Formation of a Covalent Bond Bond energy – the energy required to break a chemical bond and form neutral isolated atoms

Formation of a Covalent Bond Bond energy – the energy required to break a chemical bond and form neutral isolated atoms Bond

Bond Length (pm)

Bond Energy (kJ/mol)

H-H

74

436

Cl-Cl

199

243

Br-Br

228

193

I-I

267

151

C-C

154

346

C=C

134

612

Covalent Bond Formation Molecule – neutral group of atoms that are held together by covalent bonds

CO2

H2O C6H12O6

Molecular formula – shows the number and types of atoms combined in a single molecule

Molecular Geometry The properties of molecules depend not only the bonding of the atoms but also on the molecular geometry The polarity of the bonds as well as the arrangement of the atoms in space determines the molecular polarity

Molecular Geometry VSEPR Theory Valence Shell Electron Pair Repulsion

VSEPR Theory Repulsion between the sets of valence electrons around an atom causes these sets to be oriented as far apart from each other as possible

Molecular Geometry

180°

Molecular Geometry Two areas of electron density Linear 180°

Molecular Geometry Three areas of electron density Trigonal planar 120°

Molecular Geometry

Molecular Geometry Four areas of electron density Tetrahedron 109.5°

Molecular Geometry

Molecular Polarity Bonds Non-polar covalent bond – two atoms involved share electrons equally Polar covalent bond – electrons are shared unequally Molecules Non-polar molecules – even distribution of electrons (electron density) Polar molecules – uneven distribution of electrons (electron density)

Molecular Polarity Molecular polarity depends on both the polarity of the individual bonds and their orientation in threedimensional space.

I2 CO2 HF

H2O

Molecular Polarity Draw the Lewis structures for the following molecules. For each molecule, predict its geometry and polarity. 1. 2. 3. 4.

IBr H2S SiCl4 PF3

Monoatomic Ions Ion –atom that has lost or gained electrons We must adjust the number of electrons to account for the charge Na is a metal and will lose its valence electron. Na → Na (atom) (ion) Cl is a non-metal and will gain an electron to fill its valence shell. Cl → Cl (atom) (ion)

Polyatomic Ions Polyatomic ion –group of covalently bonded atoms that has collectively lost or gained electrons We must adjust the number of electrons to account for the charge

NH4+ SO42-

Polyatomic Ions Draw the Lewis structure of the following ions OH-

CO32-

Ionic Compounds Composed of positive and negative ions Usually involves metals and non-metals Combine in a ratio where charges cancel out Li → Li+ and F → F- so they combine in a 1:1 ratio

Ionic Compounds Draw Lewis structures of the following compounds. KF MgO CaI2

Molecular vs Ionic The force that holds ionic compounds together is a very strong attraction between + and – ions

In a molecule, the forces that hold the atoms together within a molecule are very strong (covalent bonds) but the forces that hold one molecule to another are much weaker

Molecular vs Ionic Melting point, boiling point, and hardness of a compound depend on how strongly its particles are attracted to one another

Decreasing forces of attraction

Molecular vs Ionic Melting point, boiling point, and hardness of a compound depend on how strongly its particles are attracted to one another Because the forces of attraction between individual molecules are not very strong, many molecular compounds melt at low temperatures In contrast, many ionic compounds are held together by strong forces (ionic bonds) and have higher melting and boiling points than molecular compounds

Forces of Attraction As a liquid is heated, the kinetic energy of its particles increases At some temperature, the energy is sufficient to overcome the forces of attraction between the molecules and the liquid boils

Forces of Attraction Boiling point is therefore a good measure of the strength of the force of attraction between the particles of a liquid The higher the boiling point, the stronger the forces

Forces of Attraction The forces of attraction between molecules are called intermolecular forces The strongest intermolecular forces exist between polar molecules Polar molecules act as tiny dipoles because of their uneven charge distribution Dipole-dipole attraction

Forces of Attraction Some hydrogen containing molecules have unusually high boiling points This is explained by the presence of a particularly strong IM force Hydrogen bonding – IM force that results when a molecule contains a H atom bonded to a F, O, or N atom

Forces of Attraction  Even noble gases and molecules that are non-polar experience forces of a attraction  Electrons in molecules are in constant motion and create instantaneous dipoles  These temporary differences in charge distribution cause forces of attraction known as London dispersion forces  The more electrons a molecule has, the stronger its London dispersion interactions

Forces of Attraction The boiling points of alkanes (a type or organic compound) increases with increasing molar mass As the molecules get bigger, so does the size of their electron cloud This increases the London dispersion forces and thus more energy, or heat is required to pull them apart

Boiling Point (C) Methane

-164

Ethane

-88.6

Propane

-42.1

Butane

-0.50

pentane

36.1

Forces of Attraction If we compare the alkane to its corresponding alcohol we can see another trend

Boiling Point (C) Methane

-164

Methanol

64.7

Ethane

-88.6

Ethanol

78.3

Propane

-42.1

Propanol

97.2

Butane

-0.50

Butanol

117

Metallic Bonding Chemical bonding in metals is different from that in ionic or molecular compounds This difference is reflected in their unique properties

Results from the attraction between metal ions and the surrounding “sea” of mobile valence electrons