Unit 6: Covalent Bonding

Covalent Bonding A chemical bond is the force that holds two atoms together and makes them function as a unit Atoms form bonds to become most stable and obtain an octet Covalent Bonding: Electrons are shared between two or more elements. Always between 2 non-metals Metalloids can act as nonmetals Never involves polyatomic ions The bonding results from mutual attraction of the two nuclei for the shared electrons Not all electrons are shared in a covalent bond. The unshared electrons are called lone pairs. How is covalent bonding different than ionic bonding?

Properties of Covalent Molecules Poor conductors Low melting and boiling points than ionic Because covalent bonds are usually weaker than ionic bonds Soft Most are liquid or gas phase Typically more flammable than ionic compounds Many are insoluble or partially soluble (don’t dissolve in water well)

Prefixes for Covalent Bonding Number Indicated Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8

Rules for Naming Covalent Molecules The first element in the formula is named first, and the full element name is used. The second element is named as though it were an anion (ending gets ide). Prefixes are used to denote the numbers of atoms present. (AKA used to represent the subscripts) The prefix mono- is never used for naming the first element. For example, CO is called carbon monoxide, NOT monocarbon monoxide. Prefix(not mono)1st element space prefix2nd element- ide

Naming Covalent Compounds Examples BF3 Rule 1: Name the first element, using the full element name: boron. Rule 2: Name the second element as though it were an anion: Fluoride Rule 3 and 4: Use prefixes to denote numbers of atoms. One boron atom: do not use mono- in first position. Three fluorine atoms: use prefix tri-. The name is boron trifluoride.

Covalent Compounds Name the following covalent compounds: NO N2O5 CO2 SiO3 CCl4 IF5 PCl5 P4H6 Common Molecules to know: NH3 is ammonia CH4 is methane

Writing the Formulas for Covalent Use the prefixes as subscripts (do not crisscross) Do NOT simplify covalent molecules

Write the formulas for the following Covalent Compounds Carbon Monoxide Carbon tetrafluoride Dinitrogen Trioxide Selenium dioxide Nitrogen Monoxide

Diatomic Molecules Diatomic molecules: A molecule composed of two of the same atoms You must memorize the following diatomic molecules These are the only elements able to form diatomic molecules These elements exist naturally as diatomic gases Name Formula Hydrogen H2 Nitrogen N2 Oxygen O2 Fluorine F2 Chlorine Cl2 Bromine Br2 Iodine I2

Lewis Structures A Lewis Structure is a representation of a molecule showing how valence electrons are arranged among the atoms in the molecule or ion. In writing Lewis Structures, we ONLY include valance electrons Electrons involved in bonding are called bonding pair. Electrons not involved in bonding are called lone pairs or unshared pairs Keep in mind the octet rule when drawing Lewis structures Exceptions to the octet rule: Hydrogen and helium only need a duet (lithium and beryllium only need a duet as well, but they are present in ionic bonds) Boron follows the octet, but most of the time it can be satisfied with only 6 valance electrons

Steps for Writing Lewis Structures Calculate the sum of the valence electrons from all of the atoms. Do not worry about keeping track of which electrons come from which atoms. It is the total number of valence electrons that is important. The least electronegative element is the central atom (fluorine is the most electronegative) (central atom can never be hydrogen or a halogen) Use one pair of electrons to form a bond between each pair of bound atoms. For convenience, a line (instead of a pair of dots) is often used to indicate each pair of bonding electrons Arrange the remaining electrons in pairs to satisfy the duet rule for hydrogen and the octet rule for all other elements. Always start with the central atom. If you run out of electrons, take lone pairs from the central atom and turn them into double or triple bonds If the molecule has a charge, put the structure in brackets and indicate the charge as a superscript

Lewis Structures Write the Lewis Structure for water: Step 1: find the sum of the valence electrons for H2O: 1 + 1 + 6 = 8 valence electrons Using a pair of electrons per bond, we draw in the two O-H bonds, using a line to indicate each pair of boding electrons: H-O-H We arrange the remaining electrons around the atom to achieve a noble gas electron configuration for each atoms. Remaining electrons = number of total valance electrons – number of electrons used in bonds (2 per bond) Water has used 4 electrons for bonds so we have 4 remaining electrons (8-4=4) How many lone pairs does water have?

Lewis Structure Draw the Lewis Structure for the following molecules: CCl4 PH3

Lewis Structures with Multiple Bonds Let’s do the Lewis Structure for Carbon Dioxide Total valence electrons: 4 + 6 + 6 = 16 Form Bonds: O-C-O Remaining electrons: 16-4 = 12 Distribute remaining electrons: Is this correct? (Did we use 16 electrons? Is every octet filled?) No not all octet is filled We need a double or triple bond Correct Lewis Structures: Have to draw in the incorrect lewis structure; have to draw in the correct lewis structures (2 double or triple bond)

Bonds Single Bond- a covalent bond in which one pair of electrons is shared by two atoms Double Bond- A covalent bond in which two pairs of electrons are shared by two atoms (stronger than a single bond) Triple Bond- A covalent bond in which three pairs of electrons are shared by two atoms (stronger than single and double bonds) No more than 3 pairs can be shared Having multiple possible valid structures is referred to as having resonance

Examples Draw Lewis structure for the following molecules; Do any of the following molecules have resonance? HF N2 NH4+ SO2 CF4 NO CH2O NO3-

Valence Shell Electron Pair Repulsion Theory Planar triangular Valence Shell Electron Pair Repulsion Theory Tetrahedral Trigonal pyramidal Bent

VSEPR Theory VSEPR (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion VSEPR theory assumes that the molecular geometry or shape of a molecule is determined by the strong repulsion of electron pairs Electrons around the central atom repel each other in order to have the maximum bond angle as possible In other words in order to be as far apart as possible Based on Electron Dot (Lewis structures)

Electron Domains Electron domains are regions of high electron density (ex: bonds and lone pairs) A bond is considered ONE electron domain regardless if the bond is single, double, or triple A lone pair is considered an electron domain Lone pairs repel electrons more than bonded pairs In order to determine the VSEPR geometry or the shape of the molecule you must: Draw the correct Lewis structure Determine the number of electron domains around the central atom

VSEPR Overview You will need to know the following for each shape: Shape name What the shape looks like Number of electron domains Bond angle You will need to know the following shapes: Linear Trigonal planar Bent with 3 domains Bent with 4 domains Tetrahedral Trigonal pyramidal

Linear Electron Domains = 2 Bonds on central atom =2 Unshared electron pairs on central atom = 0 Bond angles = 180 ° Examples: BeI2, CO2, HCN

Trigonal Planar Electron Domains = 3 Bonds on central atom =3 Unshared electron pairs on central atom = 0 Bond angles = 120 ° Examples: BF3, SO3 , NO3-1

Bent (or V shape) with 3 domains Electron Domains = 3 Bonds on central atom = 2 Unshared electron pairs on central atom = 1 Bond angles = <120° Lone pairs repel more than bonded pairs Examples: SO2 , O3 , NO2-1

Tetrahedral Electron Domains = 4 Bonds on central atom = 4 Unshared electron pairs on central atom = 0 Bond angles = 109.5 ° Examples: CH4 , SO42- , NH4+1

Trigonal Pyramidal Electron Domains = 4 Bonds on central atom = 3 Unshared electron pairs on central atom = 1 Bond angles = 107 ° Examples: NH3 , SO32-

Bent or V shaped with 4 domains Electron Domains = 4 Bonds on central atom = 2 Unshared electron pairs on central atom = 2 Bond angles = 104.5 ° Examples: H20, OF2

VSEPR Summary Shape Picture # of Electron Domains (#Bonds and #LP) Bond Angle Linear Trigonal Planar V shape with 3 domains Tetrahedral Trigonal Pyramidal Bent with 4 domains

Bond Polarity Electronegativity- The ability of an atom to attract an electron in a bond Increases up and to the right fluorine has the highest electronegativity (ignore noble gases) Polar- having opposite ends Nonpolar bonds- Share electrons equally Polar bond- the electron is shared unequally Creating a partially positive and a partially negative end (lower case delta to indicate partial charge) Why would the electron not be shared equally? Different atoms have different electronegativity It is a tug-of-war with the electron and the more electronegative element is winning

Bond Polarity Polar bonds are indicated by a dipole moment (arrow with a tail pointing towards the more electronegative atom) For example the bond between hydrogen and chlorine is polar; the arrow shown is the dipole moment Lower case delta with a positive and negative indicates the partial charge Polar Covalent bond: electronegativity difference ≥ 0.5 and < 1.7 Nonpolar Covalent bond: electronegativity difference < 0.5 Ionic bond: electronegativity difference ≥ 1.7 Write these as one single inequality:

Bond Polarity Determine if the bond is polar, nonpolar, or ionic AND draw the dipole moment if there is one Carbon and Oxygen Hydrogen and Carbon Hydrogen and Oxygen Hydrogen and Fluorine Oxygen and Fluorine Electronegativity H 2.1 C 2.5 O 3.5 F 4.0

Molecular Polarity Molecules are considered to be polar if they have an overall dipole (partially positive end and partially negative end) Polar molecules must have one or more polar bonds Polar molecules have a net dipole; meaning all of the dipole moments add together to one dipole A molecule can have a polar bonds and be non- polar This happens when the polar bonds cancel each other out by symmetry Nonpolar molecules have no net dipole Water is VERY POLAR

Polar vs. Non-Polar No Net dipole moment This is a nonpolar molecule

Molecular Polarity Molecular polarity affects many properties of a compound Solubility (ability to dissolve) “Like dissolves like”: molecules of similar polarity and size will dissolve in each other Polar dissolves polar Nonpolar dissolves nonpolar Polar will not dissolve nonpolar Melting point Boiling point Surface tension