Chapter 8 – Covalent Bonding Ms. Wang Lawndale High School

Chapter 8.1 – Molecular Compounds In Chapter 7, we learned about electrons being “given up” or “stolen away” This type of “tug of war” bond between a metal and nonmetal is called an ionic bond In this chapter, you will learn about another type of bond Covalent Bond – atoms held together by sharing electrons between two nonmetals

Molecules We know that a metal cation and nonmetal anion are joined together by an ionic bond and called a SALT A neutral group of atoms joined together by a covalent bond is called a MOLECULE

Monatomic vs. Diatomic Molecules Molecules can be monatomic or diatomic Diatomic Molecule – a molecule consisting of two atoms There are 7 diatomic molecules on the periodic table (SUPER 7) – N2, O2, F2, Cl2, Br2, I2, H2

Properties of Molecular Compounds Lower Melting and Boiling Points than Ionic Compounds Gases or liquids at room temperature

Molecular Formulas Molecular Formula – the chemical formula of a molecular compound It shows how many atoms of each element a molecule contains Example H2O contains 3 atoms (2 atoms of H, 1 atom of O) C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)

Practice How many atoms total and of each do the following molecular compounds contain? H2 CO CO2 NH3 C2H6O

Chapter 8.2 – Covalent Bonding Remember that ionic compounds either gain or lose electrons in order to attain a noble gas electron configuration Covalent compounds form by sharing electrons to attain a noble gas electron configuration So the Octet Rule still applies to covalent bonds

Single Covalent Bond Single Covalent Bond - two atoms held together by sharing one pair of electrons Unshared Pair / Lone Pair / Nonbonding Pair – a pair of valence electrons that is not shared between atoms Let’s Practice (you can use dots or lines) F2, H2O, CH4

Double and Triple Covalent Bonds Double Covalent Bond – a bond that involves two shared pairs of electrons Triple Covalent Bond – a bond that involves three shared pairs of electrons Let’s Practice O2 N2

Polyatomic Ion Polyatomic Ion - tightly bound group of atoms that has a positive or negative charge and behaves as one unit Some examples are NH4+ and SO32-

Bond Dissociation Energy Bond Dissociation Energy - the energy required to break the bond between two covalently bonded atoms A large bond dissociation energy corresponds to a strong covalent bond For example, carbon-carbon has a strong bond dissociation energy so it is not very reactive

Chapter 8.3 - Bonding Theories So far, the orbitals we have been discussing are atomic orbitals (s, p, d, f) for each atom When two atoms combine, their atomic orbitals overlap and make molecular orbitals Molecular Orbitals – orbitals that apply to the entire molecule instead of just one atom

Molecular Orbitals Just as atomic orbitals belong to a particular atom, a molecular orbital belongs to a molecule as a whole Each orbital is filled with 2 electrons Bonding Orbital – a molecular orbital that can be occupied by two electrons of a covalent bond

S orbitals overlapping P orbitals overlapping end-to-end Sigma Bond () Sigma Bond - when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis S orbitals overlapping P orbitals overlapping end-to-end

P orbitals overlapping side-by-side Pi Bond () Bonding electrons likely to be found in a sausage-shape above and below the axis Pi bonds tend to be weaker than sigma bonds because pi bonds overlap less than sigma bonds P orbitals overlapping side-by-side

VSEPR Theory Explains the 3D shape of molecules According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence electron pairs stay as far apart as possible

A Few VSEPR Shapes

Nine possible molecular shapes

VSEPR Theory Unshared pairs of electrons (lone pairs) are very important in predicting the shapes of molecules Examples Methane (CH4) - tetrahedral Ammonia (NH3) - pyramidal Water (H2O) – bent Carbon Dioxide (CO2) - linear

Hybrid Orbitals VSEPR is good at describing the molecular shapes, but not the types of bonds formed Orbital hybridization provides information about both molecular bonding and molecular shape In hybridization, several atomic orbitals mix to form the same total number of hybrid orbitals

Bond Hybridization Hybridization Involving Single Bonds – sp3 orbital Ethane (C2H6) Hybridization Involving Double Bonds – sp2 orbital Ethene (C2H4) Hybridization Involving Triple Bonds – sp orbital Ethyne (C2H2)

Chapter 7.4 – Polar Bonds and Molecules There are two types of covalent bonds Polar Covalent Bonds Nonpolar Covalent Bonds Polar Covalent Bond – unequal sharing of electrons where one atom has a slightly negative charge and the other atom has a slightly positive charge (HCl, H2O) Nonpolar Covalent Bond – equal sharing of electrons between two atoms (Cl2, N2, O2)

Classification of Bonds You can determine the type of bond artificially by calculating the difference in electronegativity between elements Type of Bond Electronegativity Difference Nonpolar Covalent 0  0.4 Polar Covalent 0.5  1.9 Ionic 2.0  4.0

Let’s Practice Together What type of bond is HCl? (H = 2.1, Cl = 3.1) Difference = 3.1 – 2.1 = 1.0 Therefore it is polar covalent bond. Your Turn To Practice N(3.0) and H(2.1) H(2.1) and H(2.1) Ca(1.0) and Cl(3.0) Al(1.5) and Cl(3.0) Mg(1.2) and O(3.5) H(2.1) and F(4.0)

Dipole No bond is purely ionic or covalent … they have a little bit of both characters When there is unequal sharing of electrons a dipole exists Dipole is a molecule that has two poles or regions with opposite charges Represented by a dipole arrow pointing towards the more negative end.

Practice Drawing Dipoles P- Br P = 2.1 Br = 2.8 P –Br + - Practice H(2.1) – S(2.5) C(2.5) – F(4.0) Si(1.8) – C(2.5) N(3.0) – O(3.5)

Attractions Between Molecules Intermolecular attractions are weaker than either ionic or covalent bonds Van der Waals forces – consists of the two weakest attractions between molecules dipole interactions – polar molecules attracted to one another dispersion forces – caused by motion of electrons (weakest of all forces)

Hydrogen Bond Hydrogen Bonds - attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom

Hydrogen Bond This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen Hydrogen Bonds have about 5% of the strength of an average covalent bond Hydrogen Bond is the strongest of all intermolecular forces

Intermolecular Attractions A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide) Network Solid – solids in which all of the atoms are covalently bonded to each other Melting a network solid would require breaking covalent bonds throughout the solid

Chapter 8 Assessment Page 247 Homework Chapter 8 Assessment Page 247 #’s 39-41, 43-46, 51, 53, 54, 57-59, 61, 65, 68, 83, 85, 86, 89, 96, 99, 100