Chapter 6 Section1 Covalent Bonds Objectives • Explain the role and location of electrons in a covalent bond. • Describe the change in energy and stability that takes place as a covalent bond forms. • Distinguish between nonpolar and polar covalent bonds based on electronegativity differences.

Chapter 6 Section1 Covalent Bonds Objectives, continued • Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences.

Section1 Covalent Bonds Chapter 6 Sharing Electrons • When an ionic bond forms, electrons are rearrangedand are transferred from one atom to another to form charged ions. • In another kind of change involving electrons, the neutral atoms share electrons.

Section1 Covalent Bonds Chapter 6 Sharing Electrons, continued Forming Molecular Orbitals • A covalent bond isa bond formed when atoms _______________one or more pairs of electrons. • The shared electrons move within a space called a molecular orbital. • Amolecular orbital is the ______________of high probability that is occupied by an individual electron as it travels with a wavelike motion in the three-dimensional space around one of two or more associated nuclei.

Formation of a Covalent Bond Chapter 6

Visual Concepts Chapter 6 Chemical Bond

Section1 Covalent Bonds Chapter 6 Energy and Stability Energy Is Released When Atoms Form a Covalent Bond

Section1 Covalent Bonds Chapter 6 Energy and Stability, continued Potential Energy Determines Bond Length • When two bonded hydrogen atoms are at their lowest potential energy, the distance between them is 75 pm. • The ___________________isthe distance between two bonded atoms at their minimum potential energy. • However, the two nuclei in a covalent bond vibrate back and forth. The bond length is thus the average ______________________between the two nuclei.

Visual Concepts Chapter 6 Bond Length

Section1 Covalent Bonds Chapter 6 Energy and Stability, continued Bonded Atoms Vibrate, and Bonds Vary in Strength • The bond length is the average distance between two nuclei in a covalent bond. • At a bond length of 75 pm, the potential energy of H2 is –436 kJ/mol. • Thus 436 kJ of energy must be supplied to break the bonds in 1 mol of H2 molecules. • The energy required to break a bond between two atoms is the _________________________. • Bonds that have the higher bond energies (stronger bonds) have the shorter bond lengths.

Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding • In covalent bonds between two different atoms, the atoms often have different attractions for shared electrons. • Electronegativity values are a useful tool to predict what kind of bond will form.

Visual Concepts Chapter 6 Electronegativity

Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally • When the electronegativity values of two bonding atoms are similar, bonding electrons are shared equally. • A covalent bond in which the bonding electrons in the molecular orbital are shared equally is a _________________________covalent bond.

Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally, continued • When the electronegativity values of two bonding atomsare different, bonding electrons are shared unequally. • A covalent bond in which the bonding electrons in the molecular orbital are shared unequally is a _______________________covalent bond.

Predicting Bond Character from Electronegativity Differences Chapter 6

Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends • In a polar covalent bond, the ends of the bond have opposite partial charges. • A molecule in which one end has a partial positive charge and the other end has a partial negative charge is called a ________________________. • In a polar covalent bond, the shared pair of electrons is not transferred completely. Instead, it is more likely to be found near the more electronegative atom.

Section1 Covalent Bonds Chapter 6 Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends, continued • The symbol is used to mean partial. • + is used to show a partial _______________charge • – is used to show a partial _______________charge • example: H+F– • Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

Visual Concepts Chapter 6 Comparing Polar and Nonpolar Covalent Bonds

Chapter 6 Section1 Covalent Bonds Polarity Is Related to Bond Strength • In general, the greater the electronegativitydifference, the greater the polarity and the stronger the bond.

Section1 Covalent Bonds Chapter 6 Electronegativity and Bond Types • Differences in electronegativity values provide one model that can tell you which type of bond two atoms will form. • Another general rule states: • A covalent bond forms between two _______________. • An ionic bond forms between a _______________and a __________________.

Section1 Covalent Bonds Chapter 6 Properties of Substances Depend on Bond Type • The type of __________that forms (metallic, ionic, or covalent) determines the properties of the substance. • The difference in the strength of attraction between the basic units of ionic and covalent substances causes these types of substances to have different properties.

Properties of Substances with Metallic, Ionic, and Covalent Bonds Chapter 6

Section2 Drawing and Naming Molecules Chapter 6 Bellringer • Classify the following compounds according to the type of bonds they contain: • NO • CO • HF • NaCl • HBr • NaI

Section2 Drawing and Naming Molecules Chapter 6 Objectives • Draw Lewis structures to show the arrangement of valence electrons among atoms in molecules and polyatomic ions. • Explain the differences between single, double, and triple covalent bonds. • Draw resonance structures for simple molecules and polyatomic ions, and recognize when they are required.

Section2 Drawing and Naming Molecules Chapter 6 Objectives, continued • Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures • _________________electrons are the electrons in the outermost energy level of an atom. • A _________________is a structural formula in which valence electrons are represented by dots. • In Lewis structures, dot pairs or dashes between two atomic symbols represent pairs in covalent bonds.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons • As you go from element to element across a period, you add a dot to each side of the element’s symbol.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • You do not begin to pair dots until all four sides of the element’s symbol have a dot.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • An element with an octet of valence electrons has a stable configuration. • The tendency of bonded atoms to have octets of valence electrons is called the _______________rule.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • When two chlorine atoms form a covalent bond, each atom contributes one electron to a shared pair. • An ____________________pair, or a lone pair, is a nonbonding pair of electrons in the valence shell of an atom.

Section2 Drawing and Naming Molecules Chapter 6 Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued • A _______________bond is a covalent bond in which two atoms share one pair of electrons • The electrons can pair in any order. However, any unpaired electrons are usually filled in to show how they will form a covalent bond.

Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Single Bonds Sample Problem A Draw a Lewis structure for CH3I.

Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures for Polyatomic Ions Sample Problem B Draw a Lewis structure for the sulfate ion,

Section2 Drawing and Naming Molecules Chapter 6 Multiple Bonds • For O2 to make an octet, each atom needs two more electrons. The two atoms share four electrons. • A _______________bond is a covalent bond in which two atoms share two pairs of electrons.

Section2 Drawing and Naming Molecules Chapter 6 Multiple Bonds, continued • For N2 to make an octet, each atom needs three more electrons. The two atoms share six electrons. • A _______________bond is a covalent bond in which two atoms share three pairs of electrons.

Section2 Drawing and Naming Molecules Chapter 6 Drawing Lewis Structures with Multiple Bonds Sample Problem C Draw a Lewis structure for formaldehyde, CH2O.

Section2 Drawing and Naming Molecules Chapter 6 Resonance Structures • Some molecules, such as ozone, O3, cannot be represented by a single Lewis structure. • When a molecule has two or more possible Lewis structures, the two structures are called _______________structures.

Section2 Drawing and Naming Molecules Chapter 6 Multiple Bonds, continued Naming Covalent Compounds • The first element named is usually the first one written in the formula. It is usually the less-electronegative element. • The second element named has the ending -ide. • Unlike the names for ionic compounds, the names for covalent compounds must often distinguish between two different molecules made of the same elements.

Section2 Drawing and Naming Molecules Chapter 6 Naming Covalent Compounds, continued • This system of prefixes is used to show the number of atoms of each element in the molecule.

Section2 Drawing and Naming Molecules Chapter 6 Naming Covalent Compounds, continued • Prefixes can be used to show the numbers of each type of atom in diphosphorus pentasulfide.

Visual Concepts Chapter 6 Naming Compounds Using Numerical Prefixes

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