Chapter 8 – Covalent Bonding The unspoken hero: “Covalent Bond” 

Review of Chapter 7 In Chapter 7, we learned about electrons being transferred (“given up” or “stolen away”) This type of “tug of war” between a METAL and NONMETAL is called an IONIC BOND, which results in a SALT being formed

Chapter 8.1 – Molecular Compounds Covalent Bonds - atoms held together by SHARING electrons between NONMETALS

Salt versus Molecules Molecule - A group of atoms joined together by a covalent bond Compound - a group of two or more elements bonded together (Ionic or Covalent).

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

Properties of Molecular Compounds Liquids or gases at room temperature Lower Melting Points than Ionic Compounds (means weaker bonds than ionic)

Molecular Formulas Molecular Formula – formula of a molecular compound 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 2 2 3 4 9

Practice: True or False All molecular compounds are composed of atoms of two or more elements. All compounds are molecules. Molecular compounds are composed of two or more nonmetals. Atoms in molecular compounds exchange electrons. Molecular compounds have higher melting and boiling points than ionic compounds. Share Lower

Ionic versus Covalent IONIC COVALENT Bonded Name Salt Molecule Bonding Type Transfer e- Share e- Types of Elements Metal & Nonmetal Nonmetals Physical State Solid Solid, Liquid, or Gas Melting Point High (above 300ºC) Low (below 300 ºC) Solubility Dissolves in Water Varies Conductivity Good Poor

Chapter 8.2 – Covalent Bonding Remember that ionic compounds transfer electrons in order to attain a noble gas electron configuration Covalent compounds form by sharing electrons to attain a noble gas electron configuration Regardless of the type of bond, the Octet Rule still must be obeyed (8 valence electrons) All elements need 8 except Hydrogen, which needs 2.

Single Covalent Bond A Single Covalent Bond consists of two atoms held together by sharing 1 pair of electrons (2 e-)

Electron Dot Structure

Shared versus Unshared Electrons A Shared Pair is a pair of valence electrons that is shared between atoms An Unshared Pair is a pair of valence electrons that is not shared between atoms

Practice Lewis Dot Structures Chemical Formula # of Valence Electrons Single Line Bond Structure # of Remaining Electrons Lewis Dot Structure Octet Check All Atoms=8 Hydrogen=2 F2 H2O NH3 CH4 14 F-F 12 F, F = 8 O = 8 H, H = 2 8 H-O-H 4 N = 8 3 H = 2 8 2 C = 8 4 H = 2 8

Double Covalent Bonds Sometimes atoms need to share 2 or 3 pairs of electrons Hydrogen will NEVER form double or triple bonds. Double Bond - bond that involves 2 shared pairs of electrons (4 e-)

Triple Covalent Bond Triple Bond - bond that involves 3 shared pairs of electrons (6 e-)

Covalent Bonds

Practice Lewis Dot Structure Chemical Formula # of Valence Electrons Single Line Bond Structure # of Remaining Electrons Lewis Dot Structure Octet Check All Atoms=8 Hydrogen=2 O2 CO2 N2 HCN 12 O-O 10 2 O = 8 16 O-C-O 12 C = 8 2 O = 8 2 N = 8 10 N-N 8 C = 8 N = 8 H = 2 10 H-C-N 6

Drawing Lewis Structures of Molecules Chapter 8: Basic Concepts of Chemical Bonding Drawing Lewis Structures of Molecules If the compound contains more than 2 atoms: how are the atoms bonded and, if there are nonbonding electron, where are they?

Molecules with a central atom : NH3, PCl3, CHCl3 Chapter 8: Basic Concepts of Chemical Bonding Molecules with a central atom : NH3, PCl3, CHCl3 Central atom is generally the first in the molecular formula and the most electronegative one.

…unless the first element is Hydrogen : Chapter 8: Basic Concepts of Chemical Bonding …unless the first element is Hydrogen : H2O HCN (same order as in formula)

How to find the # of bonds in a lewis structure How to find the # of bonds in a lewis structure **Doesn’t work for molecule over an octet Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e- needed (all have 8 or 2 e-) - # of val e- available = ____/2 to find the # of bonds (# val e- – # val e- available) = # of bonds 2

a.) CO b.) C2F4 c.) C2H6 3 6 7 Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e- needed (all have 8 or 2 e-) - # of val e- available = ____/2 to find the # of bonds Ex: Find the number of bonds for each molecule or compound and write the lewis dot structure # of Bonds LDS a.) CO b.) C2F4 c.) C2H6 **Doesn’t work for molecule over an octet 3 6 7

(1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding Rules for Drawing Lewis Structures (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 8 NH3 (2) Connect atoms by covalent bonds: count electrons left 2 (3) Complete "octets" of atoms around central atom n/a (4) Place any leftover electrons on the central atom. Check that central atom has octet 2 leftovers (5) If there are not enough electrons to give the central atom an octet, try multiple bonds n/a

CO 10 8 C-O C-O C-O 2 left (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding CO (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 10 (2) Connect atoms by covalent bonds: count electrons left 8 C-O (3) Complete "octets" of atoms around central atom C-O (treat C as central) (4) Place any leftover electrons on the central atom. Check that central atom has octet C-O 2 left (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

SF2 20 F-S-F 16 4 left (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding SF2 (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 20 (2) Connect atoms by covalent bonds: count electrons left F-S-F 16 (3) Complete "octets" of atoms around central atom (4) Place any leftover electrons on the central atom. Check that central atom has octet 4 left (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

Lewis Structure for Ions If a molecule has a positive charge, subtract that many electrons from the total valence electrons available. 𝑁𝐻 4 + has 8 electrons available. If a molecule has a negative charge, add that many electrons to the total valence electrons available. 𝑁𝑂 3 − has 24 electrons available. 𝑆𝑂 4 −2 has 32 electrons available.

square brackets and the sign For ions, the charge is generally indicated by square brackets and the sign

NH4+ 8 0 left n/a n/a (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding Ions NH4+ (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 8 (2) Connect atoms by covalent bonds: count electrons left 0 left (3) Complete "octets" of atoms around central atom n/a (4) Place any leftover electrons on the central atom. Check that central atom has octet n/a (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

Chapter 8: Basic Concepts of Chemical Bonding Ions ClO2- (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 20 (2) Connect atoms by covalent bonds: count electrons left O-Cl-O 16 left (3) Complete "octets" of atoms around central atom (4) Place any leftover electrons on the central atom. Check that central atom has octet 4 left (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

If the molecule has an odd number of valence electrons Chapter 8: Basic Concepts of Chemical Bonding Exceptions to the Octet rule On occasion, an atom in a molecule does not have an octet of valence electrons: If the molecule has an odd number of valence electrons an atom may have less than an octet [mainly Be, B] an atom may have more than an octet [periods 3-7]

NO2 17 O-N-O 17 1 left 13 (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding Exceptions to the Octet rule: odd number of electrons NO2 (1) Sum valence electrons from all atoms: these are the ones that need to be distributed 17 (2) Connect atoms by covalent bonds: count electrons left O-N-O 17 (3) Complete "octets" of atoms around central atom (4) Place any leftover electrons on the central atom. Check that central atom has octet 1 left (5) If there are not enough electrons to give the central atom an octet, try multiple bonds 13

BF3 24 18 left 0 left (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding Exceptions to the Octet rule: less than an octet (B or Be) (1) Sum valence electrons from all atoms: these are the ones that need to be distributed BF3 24 (2) Connect atoms by covalent bonds: count electrons left 18 left (3) Complete "octets" of atoms around central atom 0 left (4) Place any leftover electrons on the central atom. Check that central atom has octet (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

More Than Eight Electrons The only way PCl5 can exist is if phosphorus has 10 electrons around it. Periods 3-7 can expand its orbitals.

More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

More Than Eight Electrons This eliminates the charge on the phosphorus and the charge on one of the oxygens. The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

BrF5 42 32 (1) Sum valence electrons from all atoms: Chapter 8: Basic Concepts of Chemical Bonding Exceptions to the Octet rule: more than an octet (1) Sum valence electrons from all atoms: these are the ones that need to be distributed BrF5 42 (2) Connect atoms by covalent bonds: count electrons left 32 (3) Complete "octets" of atoms around central atom 2 left (4) Place any leftover electrons on the central atom. Check that central atom has octet (5) If there are not enough electrons to give the central atom an octet, try multiple bonds

Bond Dissociation Energy Bond Dissociation Energy - energy required to break a bond between two atoms A large bond dissociation energy corresponds to a strong bond which makes it unreactive Carbon has strong bonds, which makes carbon compounds stable and unreactive

Chapter 8.3 - Bonding Theories Determining shape through bonds

VSEPR Theory VSEPR Theory predicts the 3D shape of molecules According to VSEPR, the shape of the molecule adjusts so that the electrons are far apart

A Few VSEPR Shapes

Nine possible molecular shapes

VSEPR Theory Unshared pairs of electrons are important in predicting the shapes of molecules Each bond (single, double, or triple) and unshared pair is considered a steric number Use the steric number to predict the molecular geometry VSEPR can only be used with the central atom

Bent 4 2 Linear 6 Molecule Lewis Dot Structure Steric # Shape H2O CO2   CO2 XeF4 Bent 4 2 Linear Square Planar 6

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 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)

Gets another orbital added for each atom and lone pair around atom of interest. Start with sp… …then sp2… (what is the orbital hybridization of C?) …then sp3

Chapter 8.4 – Polar Bonds and Molecules There are two types of covalent bonds Nonpolar Bonds (share equally) Polar Bonds (share unequally)

Polar Covalent Polar Bond - unequal sharing of electrons between two atoms (ex: HCl) In a polar bond, one atom typically has a negative charge, and the other atom has a positive charge

Nonpolar Covalent Bond Nonpolar Bond - equal sharing of electrons between two atoms (Cl2, N2, O2)

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

Practice Your Turn To Practice 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) Mg(1.2) and O(3.5) H(2.1) and F(4.0)

Polar molecules – All molecules with lone pairs (unless it’s linear) All molecules surrounded by different atoms. Nonpolar molecule – Linear molecules with the same atoms All molecules surrounded by the same atoms.

Dipole When there is unequal sharing of electrons, a dipole exists Dipole - a molecule with two poles with opposite charges Represented by an arrow pointing towards the more negative end

Practice Drawing Dipoles P- Br P = 2.1 Br = 2.8 P –Br + - Practice H(2.1) – Cl(3.0) C(2.5) - F(4.0) H(2.1)– F(4.0)

Attractions Between Molecules There are also attractions between molecules Intermolecular attractions are weaker than ionic, covalent, and metallic bonds There are 2 main types of attractions between molecules: Van der Waals and Hydrogen

Van der Waals Forces Van der Waals forces consists of the two weak attractions between molecules 2. dispersion forces – caused by the motion of electrons (weakest of all forces) 1. dipole interactions – polar molecules attracted to one another

Hydrogen Bond Hydrogen Bonds - forces where a hydrogen atom is weakly attracted to an unshared electron pair of another atom Strongest of all intermolecular forces