Chapter 6 Covalent Compounds Table of Contents -...

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Chapter 6 Covalent Compounds

Section 1 Covalent Bonds

Section 2 Drawing and Naming Molecules

Section 3 Molecular Shapes




Bellringer

•  Make a list of the elements that form ionic bonds. Note that most ionic bonds contain a metal and a nonmetal.

Chapter 6



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.

Section 1 Covalent Bonds Chapter 6



Objectives, continued

•  Compare the physical properties of substances that have different bond types, and relate bond types to electronegativity differences.




Sharing Electrons

•  When an ionic bond forms, electrons are rearranged and are transferred from one atom to another to form charged ions.

•  In another kind of change involving electrons, the neutral atoms share electrons.




Sharing Electrons, continued Forming Molecular Orbitals •  A covalent bond is a bond formed when atoms share

one or more pairs of electrons.

•  The shared electrons move within a space called a molecular orbital.

•  A molecular orbital is the region of high probability that is occupied by an individual electron as it travels around one of two or more associated nuclei.




Formation of a Covalent Bond

Chapter 6



Energy and Stability Energy Is Released When Atoms Form a Covalent Bond




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 bond length is the 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 distance between the two nuclei.




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 bond energy.

•  Bonds that have the higher bond energies (stronger bonds) have the shorter bond lengths.




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.




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 nonpolar covalent bond.

Chapter 6



Electronegativity and Covalent Bonding, continued Atoms Share Electrons Equally or Unequally, continued


•  When the electronegativity values of two bonding atoms are different, bonding electrons are shared unequally.

•  A covalent bond in which the bonding electrons in the molecular orbital are shared unequally is a polar covalent bond.

Chapter 6



Predicting Bond Character from Electronegativity Differences

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 dipole.

•  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.

Chapter 6



Electronegativity and Covalent Bonding, continued Polar Molecules Have Positive and Negative Ends, continued


•  The symbol δ is used to mean partial.

•  δ+ is used to show a partial positive charge

•  δ– is used to show a partial negative charge

•  example: Hδ+Fδ–

•  Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

Chapter 6



Polarity Is Related to Bond Strength

•  In general, the greater the electronegativity difference, the greater the polarity and the stronger the bond.




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 nonmetals.

•  An ionic bond forms between a nonmetal and a metal.

Chapter 6



Properties of Substances Depend on Bond Type


•  The type of bond 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.

Chapter 6



Properties of Substances with Metallic, Ionic, and Covalent Bonds

Chapter 6




Bellringer

•  Classify the following compounds according to the type of bonds they contain:

•  NO •  CO •  HF •  NaCl •  HBr •  NaI

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.

Section 2 Drawing and Naming Molecules Chapter 6



Objectives, continued

•  Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.




Lewis Electron-Dot Structures

•  Valence electrons are the electrons in the outermost energy level of an atom.

•  A Lewis structure 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.




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.




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.

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 octet rule.




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 unshared pair, or a lone pair, is a nonbonding pair of electrons in the valence shell of an atom.



Lewis Electron-Dot Structures, continued Lewis Structures Show Valence Electrons, continued •  A single 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.




Drawing Lewis Structures with Single Bonds

Sample Problem A Solution

Draw each atom’s Lewis structure, and count the total number of valence electrons.


number of dots: 14

Arrange the Lewis structure so that carbon is the central atom.



Multiple Bonds

•  For O2 to make an octet, each atom needs two more electrons. The two atoms share four electrons.


•  A double bond is a covalent bond in which two atoms share two pairs of electrons.



Multiple Bonds, continued

•  For N2 to make an octet, each atom needs three more electrons. The two atoms share six electrons.


•  A triple bond is a covalent bond in which two atoms share three pairs of electrons.



Naming Covalent Compounds, continued

•  This system of prefixes is used to show the number of atoms of each element in the molecule.




Naming Covalent Compounds, continued

•  Prefixes can be used to show the numbers of each type of atom in diphosphorus pentasulfide.




Predicting Molecular Shapes

Sample Problem D Solution

Draw the Lewis structure for H2O.

Section 3 Molecular Shapes Chapter 6

Count the number of shared and unshared pairs of electrons around the central atom.

H2O has two shared pairs and two unshared pairs.



Molecular Shape Affects a Substance’s Properties Shape Affects Polarity


•  One property that shape determines is the polarity of a molecule.

•  The polarity of a molecule that has more than two atoms depends on the polarity of each bond and the way the bonds are arranged in space.

Chapter 6



Molecular Shape Affects a Substance’s Properties, continued Shape Affects Polarity, continued


•  If two dipoles are arranged in opposite directions, they will cancel each other.

•  If two dipoles are arranged at an angle, they will not cancel each other.

•  Compare the molecules of nonpolar carbon dioxide, CO2, which has a linear shape, and polar water, H2O, which has a bent shape.

Chapter 6



Molecular Shape Affects Polarity

Chapter 6



Standardized Test Preparation

Understanding Concepts

Chapter 6

1. Which of these combinations is likely to have a polar covalent bond?

A. two atoms of similar size

B. two atoms of very different size

C. two atoms with different electronegativities

D. two atoms with the same number of electrons



1. Which of these combinations is likely to have a polar covalent bond?

A. two atoms of similar size

B. two atoms of very different size

C. two atoms with different electronegativities

D. two atoms with the same number of electrons


Standardized Test Preparation Chapter 6



3. How many electrons are shared in a double covalent bond?

A. 2

B. 4

C. 6

D. 8


Chapter 6 Standardized Test Preparation



9. If water is not a good conductor of electric current, why is it dangerous to handle an electrical appliance when your hands are wet or when you are standing on wet ground?

Reading Skills




9. If water is not a good conductor of electric current, why is it dangerous to handle an electrical appliance when your hands are wet or when you are standing on wet ground?

Answer: Because even a small amount of ionic compounds dissolved in water makes it a good conductor. The salts in your body or on the ground are enough to cause the water to carry a current.

Reading Skills


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