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How to use this presentation

How to Use This Presentation

  • To View the presentation as a slideshow with effects

    select “View” on the menu bar and click on “Slide Show.”

  • To advance through the presentation, click the right-arrow key or the space bar.

  • From the resources slide, click on any resource to see a presentation for that resource.

  • From the Chapter menu screen click on any lesson to go directly to that lesson’s presentation.

  • You may exit the slide show at any time by pressing the Esc key.


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Resources

Chapter Presentation

Bellringer

Transparencies

Sample Problems

Visual Concepts

Standardized Test Prep


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Covalent Compounds

Chapter 6

Table of Contents

Section 1Covalent Bonds

Section 2Drawing and Naming Molecules

Section 3Molecular Shapes


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Chapter 6

Section1 Covalent Bonds

Bellringer

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


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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.


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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.

  • Ignore pgs.192,193


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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.


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Section1 Covalent Bonds

Chapter 6

Sharing Electrons, continued

Forming Molecular Orbitals

  • A covalent bond isa bond formed when atoms share one or more pairs of electrons.

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

  • Amolecular orbital is the region 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.


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Formation of a Covalent Bond

Chapter 6


Chemical bond

Visual Concepts

Chapter 6

Chemical Bond


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Section1 Covalent Bonds

Chapter 6

Energy and Stability

Energy Is Released When Atoms Form a Covalent Bond


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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 bond length 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 distance between the two nuclei.


Bond length

Visual Concepts

Chapter 6

Bond Length


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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 bond energy.

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


Bond energy

Visual Concepts

Chapter 6

Bond Energy


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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.


Electronegativity

Visual Concepts

Chapter 6

Electronegativity


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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 nonpolar covalent bond.


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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 polar covalent bond.


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Predicting Bond Character from Electronegativity Differences

Chapter 6


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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 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.


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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 positive charge

    • – is used to show a partial negative charge 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


Comparing polar and nonpolar covalent bonds

Visual Concepts

Chapter 6

Comparing Polar and Nonpolar Covalent Bonds


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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.


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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 nonmetals.

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


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Section1 Covalent Bonds

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.


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Properties of Substances with Metallic, Ionic, and Covalent Bonds

Chapter 6


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Chapter 6

Section 6.1 review, pg 198

1. Describe the attractive forces and repulsive

forces that exist between two atoms as the

atoms move closer together.

The positive nucleus of each atom attracts the electrons of the other atom. At the same time, the nuclei repel each other, as do the electron clouds.

3. In what two ways can two atoms share electrons when forming a covalent bond?

The two atoms may share electrons equally, forming a nonpolar covalent bond, or unequally, forming a polar covalent bond.


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Chapter 6

5. How are the partial charges shown in a polar covalent molecule?

The symbol is written as a superscript on the element with the partial positive charge. The symbol is written as a superscript on the element with the partial negative charge.

6. What information can be obtained by knowing the electronegativity differences between two elements?

Knowing the electronegativity difference suggests what type of bond will form between the atoms of the two elements.


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Chapter 6

7. Why do molecular compounds have low melting points and low boiling points relative to ionic substances?

The attractive forces between individual molecules are weak, accounting for the low melting point of molecular compounds.


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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


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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.


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Section2 Drawing and Naming Molecules

Chapter 6

Objectives, continued

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


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Section2 Drawing and Naming Molecules

Chapter 6

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.


Valence electrons

Visual Concepts

Chapter 6

Valence Electrons


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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.


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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.


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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 octet rule.


The octet rule

Visual Concepts

Chapter 6

The Octet Rule


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


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Section2 Drawing and Naming Molecules

Chapter 6

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.


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures with Single Bonds

Sample Problem A

Draw a Lewis structure for CH3I.


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Section2 Drawing and Naming Molecules

Chapter 6

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.


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures with Single Bonds

Sample Problem A Solution, continued

Distribute one bonding pair of electrons between each of the bonded atoms. Then, distribute the remaining electrons, in pairs, around the remaining atoms to form an octet for each atom.

Change each pair of dots that represents a shared pair of electrons to a long dash.


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Chapter 6.2

Sample Problem A,practice pg.202

1)Draw the Lewis structures for H2S, CH2Cl2, NH3,

and C2H6.


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Chapter 6.2

Sample Problem A,practice pg.202

2)Draw the Lewis structures for methanol, CH3OH.


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures for Polyatomic Ions

Sample Problem B

Draw a Lewis structure for the sulfate ion,


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures for Polyatomic Ions

Sample Problem B Solution

Count electrons for all atoms. Add two additional electrons to account for the 2− charge on the ion.

number of dots: 30 + 2 = 32

Distribute the 32 dots so that there are 8 dots around each atom.


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures for Polyatomic Ions

Sample Problem B Solution, continued

Change each bonding pair to a long dash. Place brackets around the ion and a 2 charge outside the bracket to show that the charge is spread out over the entire ion.


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Chapter 6.2

Sample Problem B, practice pg.203

1)Draw a Lewis structure for ClO3-

2) Draw the Lewis structure for the hydronium ion,

H3O +


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


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


Comparing single double and triple bonds

Visual Concepts

Chapter 6

Comparing Single, Double, and Triple Bonds


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures with Multiple Bonds

Sample Problem C

Draw a Lewis structure for formaldehyde, CH2O.


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures with Multiple Bonds

Sample Problem C Solution

Draw each atom’s Lewis structure, and count the total dots.

number of dots: 12

Arrange the atoms so that carbon is the central atom.

O HC H


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Section2 Drawing and Naming Molecules

Chapter 6

Drawing Lewis Structures with Multiple Bonds

Sample Problem C Solution, continued

Distribute one pair of dots between each of the atoms and the rest, in pairs, around the atoms. C does not have an octet. To get an octet, move an unshared pair from the O to between the O and the C.

Change each bonding pair to a long dash. Two pairs of dots represent a double bond.


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Chapter 6.2

  • Sample Problem C, practice pg. 205

  • Draw the Lewis structures for carbon dioxide, CO2,

  • and carbon monoxide, CO.

  • 2) Draw the Lewis structures for ethyne, C2H2, and

  • Hydrogen cyanide, HCN.


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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 resonance structures.


Atomic resonance

Visual Concepts

Chapter 6

Atomic Resonance


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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.


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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.


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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.


Naming compounds using numerical prefixes

Visual Concepts

Chapter 6

Naming Compounds Using Numerical Prefixes


Homework section 6 2 review pg 207 questions 1 to 10

HomeworkSection 6.2 Review, pg. 207questions 1 to 10


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Section 6.2 Review, pg. 207

  • Which electrons do a Lewis structure show?

  • 2. In a polyatomic ion, where is the charge located?

  • 3. How many electrons are shared by two atoms that form a triple bond?

  • 4. What do resonance structures represent?

  • 5. How do the names for SO2 and SO3 differ?


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Section 6.2 Review, pg. 207

  • 6. Draw a Lewis structure for an atom that has

  • the electron configuration 1s22s22p63s23p3.

  • 7. Draw Lewis structures for each compound:

  • BrFc. Cl2O

  • b. N(CH3)3 d. ClO2


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Section 6.2 Review, pg. 207

  • 8. Draw three resonance structures for SO3.

  • 9. Name the following compounds.

  • SnI4 c. PCl3

  • b. N2O3 d. CSe2


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Section 6.2 Review, pg. 207

10. Write the formula for each compound:

a. phosphorus pentabromide

b. diphosphorus trioxide

c. arsenic tribromide

d. carbon tetrachloride


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Chapter 6

Section3 Molecular Shapes

Bellringer

  • Write a short paragraph telling what you think the “valence shell electron pair repulsion theory” might have to do with molecular shape.


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Chapter 6

Section3 Molecular Shapes

Objectives

  • Predict the shape of a molecule using VSEPR theory.

  • Associate the polarity of molecules with the shapes of molecules, and relate the polarity and shape of molecules to the properties of a substance.


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Chapter 6

Section3 Molecular Shapes

Determining Molecular Shapes

  • The three-dimensional shape of a molecule is important in determining the molecule’s physical and chemical properties.

A Lewis Structure Can Help Predict Molecular Shape

  • You can predict the shape of a molecule by examining the Lewis structure of the molecule.


How to use this presentation

Chapter 6

Section3 Molecular Shapes

Determining Molecular Shapes, continued

A Lewis Structure Can Help Predict Molecular Shape, continued

  • The valence shell electron pair repulsion (VSEPR) theory is a theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other.


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Chapter 6

Section3 Molecular Shapes

Determining Molecular Shapes, continued

Electron Pairs Can Determine Molecular Shape

  • According to the VSEPR theory, the shape of a molecule is determined by the valence electrons surrounding the central atom.

  • Electron pairs are negative, so they repel each other.

  • Therefore, the shared pairs that form different bonds repel each other and remain as far apart as possible.


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Chapter 6

Section3 Molecular Shapes

Determining Molecular Shapes, continued

Electron Pairs Can Determine Molecular Shape, continued

  • For CO2, the two double bonds around the central carbon atom repel each other and remain far apart.

  • For BF3, the three single bonds around the central fluorine atom will be at a maximum distance apart.


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Chapter 6

Section3 Molecular Shapes

Determining Molecular Shapes, continued

Electron Pairs Can Determine Molecular Shape, continued

  • The four shared pairs of electrons in CH4are farthest apart when each pair is positioned at the corners of a tetrahedron.


Vsepr and lone electron pairs

Visual Concepts

Chapter 6

Chapter 6

VSEPR and Lone Electron Pairs


Vsepr and basic molecular shapes

Visual Concepts

Chapter 6

VSEPR and Basic Molecular Shapes


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Chapter 6

Section3 Molecular Shapes

Predicting Molecular Shapes

Sample Problem D

Determine the shape of H2O.


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Chapter 6

Section3 Molecular Shapes

Predicting Molecular Shapes

Sample Problem D Solution

Draw the Lewis structure for H2O.

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

H2O has two shared pairs and two unshared pairs.


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Chapter 6

Section3 Molecular Shapes

Predicting Molecular Shapes

Sample Problem D Solution, continued

Find the shape that allows the shared and unshared pairs of electrons to be as far apart as possible.

The water molecule will have a bent shape.


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Chapter 6

Section3 Molecular Shapes

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.


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Chapter 6

Section3 Molecular Shapes

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.


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Molecular Shape Affects Polarity

Chapter 6


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 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


How to use this presentation

Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 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


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 2. According to VSEPR theory, which of these is caused by repulsion between electron pairs surrounding an atom?

  • F.breaking of a chemical bond

  • G.formation of a sea of electrons

  • H.formation of a covalent chemical bond

  • I.separation of electron pairs as much as possible


How to use this presentation

Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 2. According to VSEPR theory, which of these is caused by repulsion between electron pairs surrounding an atom?

  • F.breaking of a chemical bond

  • G.formation of a sea of electrons

  • H.formation of a covalent chemical bond

  • I.separation of electron pairs as much as possible


How to use this presentation

Chapter 6

Standardized Test Preparation

Understanding Concepts

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

  • A.2

  • B.4

  • C.6

  • D.8


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Chapter 6

Standardized Test Preparation

Understanding Concepts

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

  • A.2

  • B.4

  • C.6

  • D.8


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 4.How can the difference in number of valence electrons between nitrogen and carbon account for the fact that the boiling point of ammonia, NH3, is 130°C higher than that of methane, CH4.


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 4.How can the difference in number of valence electrons between nitrogen and carbon account for the fact that the boiling point of ammonia, NH3, is 130°C higher than that of methane, CH4.

  • Answer: Ammonia is a polar molecule because nitrogen has a pair of electrons that are not involved in a covalent bond, while methane is a nonpolar molecule. The attraction between polar ammonia molecules causes the higher boiling point.


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 5.Why don’t scientists need VESPR theory to predict the shape of HCl?


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 5.Why don’t scientists need VESPR theory to predict the shape of HCl?

  • Answer: Because HCl has two atoms, the shape can be only linear.


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 6.What are the attractive and repulsive forces involved in a covalent bond and how do their total strengths compare?


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Chapter 6

Standardized Test Preparation

Understanding Concepts

  • 6.What are the attractive and repulsive forces involved in a covalent bond and how do their total strengths compare?

  • Answer: Attractive forces exist between each electron and each nucleus. Repulsive forces exist between electrons and between nuclei. In a covalent bond, total attractive and repulsive forces are balanced.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • Read the passage below. Then answer the questions.

  • Although water is a polar molecule, pure water does not carry an electric current. It is a good solvent for many ionic compounds, and solutions of ionic compounds in water do carry electric currents. The charged particles in solution move freely, carrying electric charges. Even a dilute solution of ions in water becomes a good conductor. Without ions in solution, there is very little electrical conductivity.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 7.Why is a solution of sugar in water not a good electrical conductor?

  • F.Sugar does not form ions in solution.

  • G.The ionic bonds of sugar molecules are too strong to carry a current.

  • H. Not enough sugar dissolves for the solution to become a conductor.

  • I. A solution of sugar in water is not very conductive because it is mostly water, which is not very conductive.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 7.Why is a solution of sugar in water not a good electrical conductor?

  • F.Sugar does not form ions in solution.

  • G.The ionic bonds of sugar molecules are too strong to carry a current.

  • H. Not enough sugar dissolves for the solution to become a conductor.

  • I. A solution of sugar in water is not very conductive because it is mostly water, which is not very conductive.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 8.Why do molten ionic compounds generally conduct electric current well, while molten covalent compounds generally do not?

  • A.Ionic compounds are more soluble in water.

  • B.Ionic compounds have more electrons than compounds.

  • C.When they melt, ionic compounds separate into charged particles.

  • D.Most ionic compounds contain a metal atom which carries the electric current.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 8.Why do molten ionic compounds generally conduct electric current well, while molten covalent compounds generally do not?

  • A.Ionic compounds are more soluble in water.

  • B.Ionic compounds have more electrons than compounds.

  • C.When they melt, ionic compounds separate into charged particles.

  • D.Most ionic compounds contain a metal atom which carries the electric current.


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 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?


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Chapter 6

Standardized Test Preparation

Reading Skills

  • 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.


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Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • Use the diagram below to answer question 10.


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Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 10. The diagram above best represents which type of chemical bond?

  • F.ionic

  • G.metallic

  • H.nonpolar covalent

  • I.polar covalent


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 10. The diagram above best represents which type of chemical bond?

  • F.ionic

  • G.metallic

  • H.nonpolar covalent

  • I.polar covalent


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • The table below shows the connection between electronegativity and bond strength (kilojoules per mole). Use it to answer questions 11 through 13.


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Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 11.Which of these molecules has the smallest partial positive charge on the hydrogen end of the molecule?

  • A.HF

  • B.HCl

  • C.HBr

  • D.HI


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 11.Which of these molecules has the smallest partial positive charge on the hydrogen end of the molecule?

  • A.HF

  • B.HCl

  • C.HBr

  • D.HI


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 12.How does the polarity of the bond between a halogen and hydrogen relate to the number of electrons of the halogen atom?

  • F.Polarity is not related to the number of electrons of the halogen atom.

  • G.Polarity decreases as the number of unpaired halogen electrons increases.

  • H.Polarity decreases as the total number of halogen atom electrons increases.

  • I.Polarity decreases as the number of valence electrons of the halogen atom increases.


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 12.How does the polarity of the bond between a halogen and hydrogen relate to the number of electrons of the halogen atom?

  • F.Polarity is not related to the number of electrons of the halogen atom.

  • G.Polarity decreases as the number of unpaired halogen electrons increases.

  • H.Polarity decreases as the total number of halogen atom electrons increases.

  • I.Polarity decreases as the number of valence electrons of the halogen atom increases.


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 13.Based on the information in this table, how does the electronegativity difference in a covalent bond relate to the strength of the bond?


How to use this presentation

Chapter 6

Standardized Test Preparation

Interpreting Graphics

  • 13.Based on the information in this table, how does the electronegativity difference in a covalent bond relate to the strength of the bond?

  • Answer: A stronger bond is indicated by greater bond energy, so the strength of the bond increases as electronegativity increases.


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