Chapter 7 and 8: CHEMICAL BONDS 1. Ionic bonds Today, we will learn about: 2. Metallic bonds 3. Covalent bonds Introduction Scale of electron sharing in bonds 3 types of covalent bonds (single, double and triple) Polar covalent bonds and electronegativity Non-polar covalent bonds Coordinate covalent bonds Examples of diatomic molecules and other common molecules with covalent bonds Bond dissociation energy VSPER theory and molecular shapes Chemistry _ Notes Dr. Chirie Sumanasekera 11/28/ 2017

2. Metallic Bonds Metallic bonds =Inside metals, positively charged (cations) swim in a sea of valence electrons. These electrons are shared by all neighboring atoms. Metals are good electrical conductors as they have readily mobile electrons to produce a current Metals form compact cubic crystals The shared-electron mobility is also why metals are malleable and ductile Alloys = mixtures of two or more metals (Ex: sterling silver, cast iron, stainless steel, surgical steel)

3. Covalent bonds INTRO Covalent bonds are found in molecular compounds (compounds that form molecules) In covalent bonds, electrons are usually shared between non metal atom so that each can attain the noble gas electron configuration There can be single, double or triple covalent bonds between two atoms Based on the electronegativity values of each atom in a covalent bond, the resulting molecule may have a dipole (polarity based on negative and positive charged regions). Covalent bonds and ionic bonds exist in a spectrum going from zero to extreme electronegativity

Types of Covalent Bonds Sigma (Single) bond Cl H Cl HCl H + H Cl Double bond O O C O + C CO2 2 O C O Triple bond N2 N + N N N N N

Scale of Valence electron sharing Similar electronegativity Large difference in electronegativity Similar electronegativity Unequal Equal Non-polar Covalent bonds Polar Covalent bonds Ionic bonds

Covalent Bonds O H2O 2 H + O H H H O H

Water: Polar Covalent Bonds  +  - Electron Density shifts towards the more electronegative oxygen atom O H O H O H Oxygen is much more electronegative than Hydrogen, so Oxygen pulls the shared electrons towards itself This greater electronegativity of oxygen compared to hydrogen in the covalent bonds of water molecules result in a dipole or polarity. The result is a shifting of the electron density towards Oxygen. H gets a small positive charge Delta (+) and Oxygen gets a negative charge ( -) as electrons are less equally shared between O and H.

Nonpolar covalent bonds + N N N N2 N N If a molecule if formed by the same kind of atom or atoms with similar electronegativity then the covalent bond is non-polar (Ex: O2, N2, Cl2, O3, Br2) Carbon and H makes non polar covalent molecules such as hydrocarbons (CH4 , C2H6 , C3H8 and so on.)

Coordinate covalent bonds Both these electrons come from Oxygen In typical covalent bond, each atom in the bond provides an electron for a single bond. But in coordinate covalent bonds this is not the case. Here, one atom contributes both of the shared electrons for one covalent bond.

Molecules: formed by covalent bonding (electron sharing) between atoms Ionic bonded compounds: Do not form molecules. They are formed by atoms that stick close due to opposite charge (electrostatic) - attractions

**Bond length decreases when number of bonds increase **Bond Strength increases when number of bonds increase

Electronegativity difference & bond types: Table 8.3 (p238) Equal electron sharing Unequal electron sharing Increasing electronegativity Non-polar covalent bonds 0.0-0.4 Polar covalent bonds 0.4 - 2.0 Ionic bonds 2.0 and greater Electronegativity difference: Examples: H H Cl Na+ Cl- H F

VSPER: Valence-Shell Electron Pair Repulsion theory: VSPER Theory VSPER: Valence-Shell Electron Pair Repulsion theory: Repulsion between valence shell electron pairs cause molecular shapes to change so that valence electron pairs stay as far apart as possible Three bonds Four bonds Five bonds Six bonds Two bonds Pyramidal Tetrahedral Trigonal bipyramidal Octahedral Bent triatomic 180o Trigonal planar Linear triatomic Square planar

O O H H Effect of unpaired electrons on molecular shape Repulsion between unpaired electrons also effect 3D molecular shapes: Ex. H2O In water, the two lone pair electron repulsion and the electron clouds of these electron pairs cause the molecule to have a tetrahedral rather than a triatomic shape. Repulsive force between unpaired electrons O H O H

Diatomic Elements These are elements that exist in nature as molecules made up of two identical atoms. These atoms have the same electronegativity so these molecules are non-polar! Examples: Hydrogen gas (H2) Oxygen gas (O2) Nitrogen gas (N2) Gasses of Halogens (Cl2, Br2, I2) ***The subscript 2 is used to show the two identical atoms in the molecule Side Note: this is very different from mono-atomic gases. All of the noble gases are mono atomic gases. They are not molecules. They are single atoms that exist as a gas

Common molecular compounds (Table 8.2, Page 224) Nitrogen dioxide NO2 Hydrogen peroxide Name Chemical formula H2O2 Hydrochloric acid Hydrogen cyanide HCl SO3 HCN Sulfur trioxide Nitrous Oxide N2O CH4 Methane What do all these chemical formulas have in common based on the types of elements You see? (Clue: what types of elements are they? Find their location in Periodic Table…)

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Increasing Bond strength (kJ/mol) Covalent Bond Strength and Bond Dissociation Energy Increasing Bond strength (kJ/mol) Single bond Double bond Triple bond When two H-atoms combine to form H2 gas by covalent bonding, a large amount of heat is released (435 kJ). To break this covalent bond, the same amount of energy is required. The energy required to break a covalent bond = Bond Dissociation Energy. Bond dissociation energy depends on which molecules are involved the bonding and how many bonds are to be broken. If more electronegative atoms are present, the energy will be higher. It takes more energy to break a triple bond than a single bond. F H + 565 kJ of energy + 435 kJ of energy