Chapter 8 Covalent Bonding The unspoken hero Covalent

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) – • H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2

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 n. H 2 O contains 3 atoms (2 atoms of H, 1 atom of O) n. C 2 H 6 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? 2 1. H 2 2 2. CO 3 3. CO 2 4 4. NH 3 9 5. C 2 H 6 O

Practice: True or False 1. All molecular compounds are composed of 2. 3. 4. 5. atoms of two or more elements. All compounds are molecules. Molecular compounds are composed of two or more nonmetals. Atoms in molecular compounds exchange Share electrons. Molecular compounds have higher melting and boiling points than ionic compounds. Lower

Ionic versus Covalent IONIC COVALENT Bonded Name Salt Molecule Bonding Type Transfer e- Share e- Types of Elements Metal & Nonmetals Physical State 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 F 2 H 2 O # of Valence Electrons 14 8 NH 3 8 CH 4 8 Single Line Bond Structure # of Remaining Electrons F-F 12 Lewis Dot Structure Octet Check All Atoms=8 Hydrogen=2 F, F = 8 4 O = 8 H, H = 2 2 N = 8 3 H = 2 H-O-H 0 C = 8 4 H = 2

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

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? Chemical Bonding

Chapter 8: Basic Concepts of Chemical Bonding Molecules with a central atom : NH 3, PCl 3, CHCl 3 Central atom is generally the first in the molecular formula and the most electronegative one. Chemical Bonding

Chapter 8: Basic Concepts of Chemical Bonding …unless the first element is Hydrogen : H 2 O HCN (same order as in formula) Chemical Bonding

How to find the # of bonds in a lewis structure **Doesn’t work for molecule over an octet 1. 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) 2 = # of bonds

1. 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. ) C 2 F 4 c. ) C 2 H 6 3 6 7 **Doesn’t work for molecule over an octet

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

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

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

Lewis Structure for Ions • If a molecule has a positive charge, subtract that many electrons from the total valence electrons available. • If a molecule has a negative charge, add that many electrons to the total valence electrons available. Chemical Bonding

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

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

Chapter 8: Basic Concepts of Chemical Bonding Ions (1) Sum valence electrons from all atoms: these are the ones that need to be distributed (2) Connect atoms by covalent bonds: count electrons left Cl. O 220 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 Chemical Bonding

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] Chemical Bonding

Chapter 8: Basic Concepts of Chemical Bonding Exceptions to the Octet rule: odd number of electrons (1) Sum valence electrons from all atoms: these are the ones that need to be distributed (2) Connect atoms by covalent bonds: count electrons left 17 O-N-O NO 2 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 Chemical Bonding

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 (2) Connect atoms by covalent bonds: count electrons left 24 BF 3 18 left (3) Complete "octets" of atoms around central atom (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 0 left Chemical Bonding

More Than Eight Electrons • The only way PCl 5 can exist is if phosphorus has 10 electrons around it. • Periods 3 -7 can expand its orbitals. Chemical Bonding

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. Chemical Bonding

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 3 rd row or below and expanding its octet eliminates some formal charges, do so. Chemical Bonding

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 (2) Connect atoms by covalent bonds: count electrons left (3) Complete "octets" of atoms around central atom 42 Br. F 5 32 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 Chemical Bonding

Bond Dissociation Energy • Bond Dissociation Energy - energy required to break a bond between two atoms l l 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 Chemical Bonding

Chapter 8. 3 - Bonding Theories • Determining shape through bonds

VSEPR Theory • VSEPR Theory predicts the 3 D shape of molecules l 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 l Unshared pairs of electrons are important in predicting the shapes of molecules l l l 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

Molecule Lewis Dot Steric # Structure H 2 O Shape 4 CO 2 Xe. F 4 Bent Linear 2 6 Square Planar

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 l In hybridization, several atomic orbitals mix to form hybrid orbitals

Bond Hybridization • Hybridization Involving Single Bonds – sp 3 orbital n Ethane (C 2 H 6) • Hybridization Involving Double Bonds – sp 2 orbital n Ethene (C 2 H 4) • Hybridization Involving Triple Bonds – sp orbital n Ethyne (C 2 H 2)

Gets another orbital added for each atom and lone pair around atom of interest. Start with sp… …then sp 2… (what is the orbital hybridization of C? ) …then sp 3

Chapter 8. 4 – Polar Bonds and Molecules • There are two types of covalent bonds Nonpolar Bonds (share equally) n Polar Bonds (share unequally) n

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 (Cl 2, N 2, O 2)

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 Nonpolar Covalent Electronegativity Difference 0 � 0. 4 Polar Covalent 0. 5 � 1. 9 Ionic 2. 0 � 4. 0

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 l N(3. 0) and H(2. 1) l. Mg(1. 2) and O(3. 5) l H(2. 1) and H(2. 1) l H(2. 1) and F(4. 0) l Ca(1. 0) and Cl(3. 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 l. When there is unequal sharing of electrons, a dipole exists l. Dipole - a molecule with two poles with opposite charges l. Represented by an arrow pointing towards the more negative end

Practice Drawing Dipoles P- Br P = 2. 1 Br = 2. 8 P –Br �+ �- Practice l H(2. 1) – Cl(3. 0) l C(2. 5) - F(4. 0) l H(2. 1)– F(4. 0)

Attractions Between Molecules • There 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 1. dipole interactions – polar molecules attracted to one another 2. dispersion forces – caused by the motion of electrons (weakest of all forces)

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