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Covalent Bonding

Covalent Bonding. Chapter 8. 8.1 The Covalent Bond. In order for an atom to gain stability, it can gain, lose, or share electrons. Atoms that share electrons and form covalent bonds in order to complete their valence shells are called molecules. diatomic m olecules. Single Covalent Bonds.

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Covalent Bonding

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  1. Covalent Bonding Chapter 8

  2. 8.1 The Covalent Bond • In order for an atom to gain stability, it can gain, lose, or share electrons. • Atoms that share electrons and form covalent bonds in order to complete their valence shells are called molecules. • diatomic molecules

  3. Single Covalent Bonds • Lewis structures can be used to represent the arrangement of electrons in a molecule. • A line or a pair of vertical dots is used to represent a single covalent bond. • Single bonds (2 electrons shared) is also called a sigma bond (σ).

  4. Multiple Covalent Bonds • Double Bonds • Triple Bonds • A multiple covalent bond consists of one sigma and at least one pi (π)bond. • How many sigma and pi bonds are in methane? How many in a molecule of oxygen?

  5. The Strength of Covalent Bonds • Bond length • the distance between two bonded nuclei • as the number of shared pairs increases, the bond length decreases • the shorter the bond length, the stronger the bond • Bonds and Energy • Endothermic reaction: more energy required to break existing bonds than is released when the new bonds form • Exothermic reaction: more energy is released during product bond formation than is required to break bonds in the reactants

  6. 8.2 Naming Molecules • Binary Molecular Compounds • The first element in the formula is always named first, using the entire element name. • The second element in the formulat is named using its root and adding the suffix –ide. • Prefixes are used to indicate the number of atoms of each element that are present in the compound. • The first element never uses the mono- prefix. • If using a prefix results in two consecutive vowels, one is dropped.

  7. Practice How would you write sulfur trioxide? phosphorus pentafluoride?

  8. Naming Acids • Binary Acids • The first word has the prefix hydro- to name the hydrogen part of the compound. The rest of the first word consists of a form of the root of the second element plus the suffix –ic. • Oxyacids • Identify the polyatomic ion present. The first word consists of the root of the polyatomic ion. The first word has a suffix that depends on the polyatomic ion’s suffix. If the polyatomic ion ends with the suffix –ate, replace it with the suffix –ic. If the polyatomic ion ends in –ite, replace it with –ous.

  9. Practice • hydrochloric acid • sulfuric acid • H3PO3 • acetic acid • HBr • carbonic acid

  10. 8.3 Molecular Structure • Structural formulas use the symbols and bonds to show relative positions of atoms. • Lewis structures • Predict the location of certain atoms. Least electronegative is usually the center, and hydrogen is always terminal. • Determine the number of electrons available for bonding. • Determine the number of bonding pairs. • Place the bonding pairs. • Determine the number of electron pairs remaining. • Determine whether the central atom satisfies the octet rule. • You may have to convert bonds around the central atom to multiples to satisfy the octet. (CNOPS)

  11. Resonance and Exceptions • When more than one valid Lewis structure is possible, the compound is considered to have resonance. • nitrite, sulfur dioxide • Exceptions to the Octet Rule • Odd number of valence electrons (NO2) • Suboctets and coordinate covalent bonds (BH3 and NH3) • Expanded octets (PCl5)

  12. 8.4 Molecular Shapes • VSEPR Model – Valence Shell Electron Pair Repulsion • Bond angle • Electron pairs repulse each other in a molecule • Lone pairs take up more space than bonded pairs • Hybridization • Atomic orbitals can mix and form new, hybrid orbitals

  13. Practice • linear (BeCl2) • bent (H2O) • trigonal planar (AlCl3) • tetrahedral (CH4)

  14. 8.5 Electronegativity & Polarity • Electronegativity is the tendancy of an atom to attract electrons. • Bond character • Nonpolar (EN < 0.4) • Polar covalent (EN 0.4 – 1.7) • Ionic (EN > 1.7)

  15. Polar Covalent Bonds • Molecular polarity • partial charges occur in molecules (electrons are unequally shared) and are symbolized by the letter delta • Polarity and molecular shape • The shape of a molecule can determine polarity. • Water and carbon tetrachloride

  16. Dipoles and Polarity • dipole moment: the direction of the polar bond within a molecule (arrow) • arrow points at the more electronegative end • a molecule is polar if the dipole moments do NOT cancel out • a molecule is nonpolar if the dipole moments cancel • polarity of shapes: • always polar: trigonal pyramidal and bent • for others, polarity depends upon atoms attached to central atoms • to predict polarity • write electron dot • determine shape • determine dipole moment

  17. Solubility • Bond type and shape of a molecule determines its polarity. • LIKE DISSOLVES LIKE!!

  18. Properties of Covalent Compounds • Intermolecular Forces • Dispersion force, or induced dipole • Dipole-dipole force • Weak IMF result in highly volatile compounds (oxygen and carbon dioxide) and relatively soft solids (paraffin). • Covalent Network Solids • diamond and quartz • extremely hard solids and nonconductors

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