Ch. 8 Covalent Bonds

The Covalent Bond Covalent Bond: The bond that results from the sharing of valence electrons Both elements have a strong hold on the valence electron 2 nonmetals bonded together Molecule: a neutral group of atoms joined together by covalent bonds

Diatomic Molecules 7 elements that naturally form 2-atom molecules due to increased stability Found as pairs unless they are in another compound Five gases  H2, N2, O2, F2, Cl2 One Liquid  Br2 One Solid  I2 **Mr. BrINClHOF

Formation of a covalent bond Sharing of electrons to achieve an octet Example: Chlorine …7 valence electrons Notice: neither atom gives up an electron By sharing, each atom achieves an octet Sharing

Molecular Compounds Molecular Compounds: a compound composed of molecules Example: Water and Carbon Monoxide Molecular Formula: chemical formula for a molecular compound Shows how many atoms of each element a molecule contains

Formulas of Some Molecular Compounds

Types of Bonds Single Bond: Two atoms share one pair of electrons Bonding Pair: The shared pair of electrons can be represented by a pair of dots or a line H:H or H–H Lone pairs: non-bonding pairs of valence electrons

Types of Bonds Double Bond: Two atoms share two pairs of electrons Example: O::O O=O Triple Bond: Two atoms share 3 pairs of electrons Largest bond possible Example: N:::N N≡N

Bond Length & Strength Bond Length Bond Strength Distance between two bonding nuclei Single bond > double bond > triple bond Bond Strength Dissociation Energy: The energy needed to break a bond Indicates the strength of bonds Triple bond > double bond > single bond

Drawing Dot Structures of Molecules Draw the dot structure for each individual atom. Determine the number of valence electrons Determine the number of electrons each atom needs to share to become stable Choose your central atom (if more than one) - Typically the least electronegative element

Properties of Covalent Bonds Lower melting and boiling points (compared to ionic compounds) Most are gases or liquids at room temperature Composed of two or more nonmetals Poor conductor of electricity

Nomenclature for Molecules Naming Binary Molecular Compounds: - Composed of 2 nonmetals - Use Prefixes Number or atoms – prefix mono- hexa- di- hepta- tri- octa- tetra- nona- penta- deca-

Naming Guidelines First Element: Second Element: Use a prefix to indicate the number of atoms Name the element Note: Don’t use mono- for the first element Second Element: Name the root of the element name Add the –ide suffix Example: CO2 - Carbon Dioxide

Naming Practice P2O5 SO2 Dinitrogen tetrahydride

Naming Acids Acid: a substance whose molecules yield H+ when dissolved in water (H is first in the formula) Names are based on the suffix of the anion bonded to the hydrogen Anion Ending Formula Example Anion Name Naming System Acid -ide HCl chloride Hydro-anion root - ic Hydrochloric Acid -ate H2SO4 sulfate Anion root - ic Sulfuric Acid - ite HClO2 chlorite Anion root - ous Chlorous Acid

Acid Naming Practice Phosphoric Acid Hypochlorous Acid HBr H2CO3 HNO3

Acid Naming Practice Phosphoric Acid – H3PO4 Hypochlorous Acid - HClO HBr – Hydrobromic Acid H2CO3 – Carbonic Acid HNO3 – Nitric Acid

Molecular Structures Structural Formula: Uses symbols and bond lines to show the relative position and interactions of the elements in the formula Steps: Calculate the total number of available valence electrons For an anion, add e- equal to the negative charge For a cation, subtract e- equal to the charge Identify the central atom Usually the least electronegative Usually the first atom in the formula Never hydrogen

3) Connect the terminal atoms to the central atom(s) with single bonds 4) Complete the octet on the terminal atoms Hydrogen can only have a maximum of 2 electrons 5) Complete the octet on the central atom with remaining electrons **For emergency: Convert one or two lone pairs on the terminal atom to multiple bonds with the central atom if no electrons are left

Examples: CH3Br - Total # of Valence Electrons: C=4, H=1, H=1, H=1, Br=7 = 14 total valence electrons

Practice: CO2 SO42- H2CO

Practice: O C O O C O O C O O O S O O CO2 SO42- 4+6+6 = 16 total valence electrons SO42- 6+6+6+6+6+2(charge) = 32 total valence electrons O C O X O C O O C O 20 valence electrons – too many! Convert lone pairs to form a double bond 16 valence electrons! O 2- If the compound has a charge, we put brackets around the molecular structure and indicate the charge on the top right. O S O O

Practice: H C O H H C O H H C O H H2CO 1+1+4+6 = 12 total valence electrons H C O H H C O X H H C O H 14 valence electrons – too many! Convert lone pairs to form a double bond 12 valence electron – perfect!

Sigma & Pi Bonds Sigma ( ) bond: the first bond made with any other atom Pi () Bond: any 2nd or 3rd bond made with any other atom - 4 bonds 2 sigma bonds 2 pi bonds - 4 lone pairs O C O

Resonance Structure Resonance Structure: More than one Lewis structure can be drawn Atoms remain in the same location Electrons are moved Example: Nitrate Ion (NO3-) The structures vary based on the location of the double bond All resonance structures are considered equal

Octet Rule Exceptions In some cases, the octet rule can be broken If there is an odd number of valence electrons Central atom will not have an octet Examples: ClO2 and NO Central atom with less than 8 electrons These compounds tend to be very reactive B, Al, Be as the central atom Example: BH3 Central atom with more than 8 electrons Most common exception “Expanded octet” Extra electrons fill in the empty d-sublevel P, S, Cl, Br, I (3rd row or beyond) Examples: SF6 and XeF4

Molecular Orbits When two atoms combine, the molecular orbital model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule. Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole.

Molecular Shapes Molecular Shapes…VSEPR Valence Shell Electron Pair Repulsion Model Method of predicting molecular shapes Minimizes repulsion of the electron pairs 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. Predicts bond angles – angles formed by two terminal atoms and the central atom Lone pairs occupy a slightly large space Classification is based on a dot structure

Molecular Shapes Example: O3 Draw the Lewis (dot) structure Count the number of electron domains around the central atom Predict the shape and bond angles of the molecule based on the types and quantities of the domains 3 domains (regions where electrons are found) 2 bonding, 1 nonbonding Geometric Shape: BENT

Molecular Shapes ClF3 NH3

Molecular Shapes ClF3 28 Valence Electrons NH3 8 valence electrons 5 domains (regions where electrons are found) 3 bonding, 2 nonbonding Geometric Shape: T-Shaped 4 domains (regions where electrons are found) 3 bonding, 1 nonbonding Geometric Shape: Trigonal Pyramidal

Hybridization Hybridization Process of combining atomic orbitals to make equally shaped bonding orbitals In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. This enables the central atom to make identical bonds with terminal atoms Sigma bonds and lone pairs hybridize Carbon electron configuration: 1s2 2s2 2p2

Electronegativity and Bond Polarity Electronegativity: the ability to attract electrons in a chemical bond Trend: increases LR, decreases TB F is the most electronegative, Fr is the least Bond Character The bonding electron pairs are usually unequally shared The bonding pairs of electrons in a covalent bond are pulled (like tug-a-war) between the nuclei of the atoms sharing the electrons. Bond Polarity Describes the sharing of electrons

Bond Polarity Nonpolar Covalent Bond: Electrons are shared equally Example: two identical atoms (Cl2) Polar Covalent Bond: unequal sharing of electrons Also known as a dipole: one of the atoms exerts a greater attraction for the bonding electrons than the other The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. Example: HCl

Bond Character Bond Character: based on the electronegativity difference

H – Cl H – Cl Polar Bonds Polar Bonds can be represented 2 ways: #1 : Using () delta (meaning there is a partial charge) Example: HCl Electronegativity of Cl: 3.0 Electronegativity of H: 2.1 3.0 - 2.1 = 0.9 Polar Covalent Bond **The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. #2 : Using an arrow point to the more electronegative atom +  - H – Cl H – Cl

Practice CH4 CO2 H2CO HCN

Practice: Key CO2 C – H 2.5 – 2.1 = 0. 4 Nonpolar O – C 3.5 – 2.5 = 1  - + C – H 2.5 – 2.1 = 0. 4 Nonpolar O – C 3.5 – 2.5 = 1 Polar

Practice: Key HCN C – H N – C 2.5 – 2.1 = 0. 4 3.0 – 2.5 = 0.5 H2CO  - + C – H 2.5 – 2.1 = 0. 4 Nonpolar N – C 3.0 – 2.5 = 0.5 Polar  - + C – H 2.5 – 2.1 = 0. 4 Nonpolar O – C 3.5 – 2.5 = 1 Polar

Molecular Polarity All Nonpolar Bonds Polar Bonds If bonds are: Shape Keeps Polarity Shape Cancels Polarity Nonpolar Molecule Nonpolar Molecule Polar Molecule

Practice: Molecular Polarity H2O CS2 CO2

Practice: Molecular Polarity H2O CS2 O – H 3.5 – 2.1 = 1.4 Polar Bond Shapes Keeps Polarity – Polar Molecule C – S 2.5 – 2.5 = 0 Nonpolar Bond Nonpolar Molecule

Practice: Molecular Polarity O – C 3.5 – 2.5 = 1 Polar Bond Nonpolar Molecule CO2

Attractions between Molecules Intermolecular forces The attractive forces that causes the interactions between two molecules Weaker than ionic and covalent bonds Van der Waals Forces The weakest attractions between molecules Dipole Interactions Occurs when polar molecules are attract on another The electrical attraction involved occurs between the oppositely charged regions of polar molecules Dispersion Forces The weakest of all molecular interactions When the moving electrons happen to momentarily more on the side of a molecule closest to a neighboring molecule – attracts to the neighboring molecule

Forces of Attraction Hydrogen Bonds: The 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.

Characteristic Ionic Covalent Bond Formation Transfer of Electrons Sharing Electron Types of Elements Metal & Nonmetals Two or more nonmetals Internal Structure Crystalline Solid Geometric Shapes Melting & Boiling Points High Low Physical State Solids Liquids & Gases Solubility in Water Usually High High to Low Conductivity Naming techniques Cation (+) – Element name Anion (-) – Element root plus –ide suffix *Unless the anion is a polyatomic ion, then it ends in –ate or –ite Use PREFIXES based on the number of atoms * If mono is attached to the first element, drop it off. Add the suffix –ide to the end of the second element