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Section 9.1 Bonding of Atoms

Section 9.2 Molecular Shape and Parity

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Section 9.1

Bonding of Atoms

• Predict the type of bond that forms between atoms by calculating electronegativity differences.

• Compare and contrast characteristics of ionic, covalent, and polar covalent bonds.

• Interpret the electron sea model of metallic bonding.

Section 9.1

Bonding of Atoms

alkali metal: any element from Group 1: lithium, sodium, potassium, rubidium, cesium, or francium

Section 9.1

Bonding of Atoms

electronegativity

shielding effect

polar covalent bond

malleable

ductile

conductivity

metallic bond

The difference between the electronegativities of two atoms determines the type of bond that forms.

Section 9.1

A Model of Bonding

• Atoms form two different types of bonds—ionic and covalent.

Section 9.1

A Model of Bonding (cont.)

• The properties of any compound, particularly its physical properties, are related to how equally the electrons are shared.

Section 9.1

Electronegativity—An Attraction for Electrons• Electronegativity is a measure of the

ability of an atom in a bond to attract electrons.

Section 9.1

Electronegativity—An Attraction for Electrons (cont.)

• Electronegativity is a periodic property—it varies in a predictable pattern across a period and down a group on the periodic table.

Section 9.1

Electronegativity—An Attraction for Electrons (cont.)

Section 9.1

Electronegativity—An Attraction for Electrons (cont.)

• Noble gases are considered to have electronegativity values of zero and do not follow periodic trends.

• Shielding effect is the tendency for the electrons in the inner energy levels to block the attraction of the nucleus for the valence electrons.

Section 9.1

Electronegativity—An Attraction for Electrons (cont.)

Section 9.1

Electronegativity—An Attraction for Electrons (cont.)

• The farther the bonding atoms are from each other on the periodic table, the greater their electronegativity difference.

Section 9.1

Ionic Character

• Electronegativity difference (∆EN) is a measure of the degree of difference of ionic character in a bond.

• ∆EN is calculated by subtracting the smaller electronegativity from the larger, so the ∆EN is always positive.

Section 9.1

Ionic Character (cont.)

• When the electronegativity difference in a bond is 2.0 or greater, the sharing of electrons is so unequal that you can assume that the electron on the less electronegative atom is transferred.

Section 9.1

Ionic Character (cont.)

Section 9.1

Ionic Character (cont.)

• The greater the difference in the electronegatives of two atoms, the more ionic the bond between the atoms.

• The greater the distance between the bonding atoms on the periodic table, the more ionic the bond between the atoms.

Section 9.1

Covalent Character

• Two atoms of the same element form pure covalent bonds because the difference in electronegativity is zero and the valence electrons are shared equally.

• All diatomic elements have pure covalent bonds.

Section 9.1

Covalent Character (cont.)

• A bond in which the electronegativity is less than or equal to 0.5 is considered to be a covalent bond with electrons that are shared almost equally.

Section 9.1

Covalent Character (cont.)

Section 9.1

Polar Covalent Bonds

• Bonds in which the pair of electrons is shared unequally have electronegativity differences between 0.5 and 2.0.

• A bond that forms when electrons are shared unequally is called a polar covalent bond.

Section 9.1

Polar Covalent Bonds (cont.)

• Polar covalent bonds have a significant degree of ionic character.

– Represented by delta plus (+) and delta minus (–) to indicate a partial positive charge and partial negative charge

Section 9.1

Polar Covalent Bonds (cont.)

• Compounds with pure covalent bonds have different properties from compounds with polar covalent bonds, such as lower melting and boiling points.

Section 9.1

Polar Covalent Bonds (cont.)

• Unequal sharing causes an imbalance in the distribution of charge about the two bonding atoms.

Section 9.1

Polar Covalent Bonds (cont.)

• Water has polar bonds.

• A hydride is a compound formed between any element and hydrogen.

Section 9.1

Bonding in Metals

• When a metal can be pounded or rolled into thin sheets, it is malleable.

• Ductile metals can be drawn into wires.

Section 9.1

Bonding in Metals (cont.)

• Electrical conductivity is a measure of how easily electrons can flow through a material to produce an electric current.

• These properties—malleability, ductility, and electrical conductivity—are the result of the way that metal atoms bond with each other.

Section 9.1

Bonding in Metals (cont.)

Section 9.1

Bonding in Metals (cont.)

• A metallic bond is the bond that results when metal atoms release their valence electrons to a pool of electrons shared by all the metal atoms.

• This is referred to as the electron sea model.

Section 9.1

Section Assessment

With few exceptions, electronegativity values ___ as you move from left to right in any period of the periodic table.

A. increase

B. decrease

C. stay the same

Section 9.1

Section Assessment

Which element has the highest value of electronegativity?

A. carbon

B. oxygen

C. fluorine

D. iron

End of Section 9.1

Section 9.2

Molecular Shape and Polarity

• Diagram Lewis dot diagrams for molecules.

• Formulate three-dimensional geometry of molecules from Lewis dot diagrams.

• Predict molecular polarity from three-dimensional geometry and bond polarity.

Section 9.2

Molecular Shape and Polarity

electronegativity: a measure of the ability of an atom in a bond to attract electrons

Section 9.2

Molecular Shape and Polarity

double bond

triple bond

polar molecule

The shape of a molecule and the polarity of its bonds determine whether the molecule as a whole is polar.

Section 9.2

The Shapes of Molecules

• Molecular shape and polarity are related to each atom’s attraction for the electrons in a bond.

• Hydrogen molecules will always be linear.

Section 9.2

The Shapes of Molecules (cont.)

• Electron pairs repel each other and cause molecules to be in fixed positions relative to each other.

• Unshared electron pairs also determine the shape of a molecule.

• Electron pairs are located in a molecule as far apart as they can be.

Section 9.2

The Shapes of Molecules (cont.)

• In a tetrahedral arrangement, such as water, the repulsions between the electron pairs are minimized.

• The nonbonding electrons require more room so they distort the tetrahedral arrangement by squeezing the bonding pairs closer together and decreasing the bond angle.

Section 9.2

The Shapes of Molecules (cont.)

• A bond formed by sharing two pairs of electrons between two atoms is called a double bond.

• Carbon dioxide—the oxygen atom shares two pairs of electrons with the carbon atom

Section 9.2

The Shapes of Molecules (cont.)

• Ammonia—the arrangement is tetrahedral but the geometry of its four atoms is a triangular pyramid.

Section 9.2

The Shapes of Molecules (cont.)

• Methane is the simplest hydrocarbon and forms a perfect tetrahedron with bond angles of 109.5°.

Section 9.2

The Shapes of Molecules (cont.)

• Hydrocarbons are organic compounds composed only of carbon and hydrogen.

• Ethane is second member of hydrocarbon series known as the alkanes.

Section 9.2

The Shapes of Molecules (cont.)

• Alkanes are hydrocarbons that contain only carbon and hydrogen atoms with single bonds between all the atoms.

– A tetrahedral arrangement of bonding electron pairs around each carbon atom provides the most space for the electrons.

Section 9.2

The Shapes of Molecules (cont.)

• Ethene, also known as ethylene, is the simplest of the alkenes, or hydrocarbons in which one or more double bonds link carbon atoms.

– In order for the carbon atoms to acquire an octet of electrons, a double bond must exist between the carbons.

– flat, triangular arrangement of atoms

Section 9.2

The Shapes of Molecules (cont.)

• Ethyne, more commonly known as acetylene, is the first member of a hydrocarbon series called the alkynes, which are unsaturated hydrocarbons that contain a triple bond between two carbon atoms.

Section 9.2

• A bond formed by sharing three pairs of electrons between two atoms is called a triple bond.

• A triple bond makes the linear molecule rigid.

The Shapes of Molecules (cont.)

Section 9.2

• A polar molecule, or dipole, is a molecule that has a positive and a negative pole.

• Water molecules attract one another because they have positive and negative ends.

• Ammonia is another molecule with polar bonds.

Polar and Nonpolar Molecules

Section 9.2

Polar and Nonpolar Molecules (cont.)

Section 9.2

• Carbon dioxide is a nonpolar molecule since the linear arrangement of the molecules results in no separation of the positive and negative charge.

Polar and Nonpolar Molecules (cont.)

Section 9.2

• The force between dipole molecules is an attraction of the positive end of one dipole for the negative end of another dipole.

Polar and Nonpolar Molecules (cont.)

Section 9.2

• Properties of polar molecules differ from nonpolar molecules.

• Melting point and boiling points of polar substances tend to be higher than those of nonpolar molecules of the same size.

Polar and Nonpolar Molecules (cont.)

Section 9.2

Polar and Nonpolar Molecules (cont.)

Section 9.2

• Ionic compounds exhibit a narrower range of physical properties than covalent compounds.

Ions, Polar Molecules, and Physical Properties

Section 9.2

• Ionic compounds tend to be brittle, solid substances with high melting points.

• Covalent substances may be solids, liquids, or gases at room temperature.

Ions, Polar Molecules, and Physical Properties (cont.)

Section 9.2

Section Assessment

Which exhibit a narrower range of physical properties?

A. ionic

B. covalent

Section 9.2

Section Assessment

Carbon dioxide is a ___ molecule.

A. polar

B. nonpolar

C. dipole

D. triple bond

End of Section 9.2

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Study Guide

Chapter Assessment

Standardized Test Practice

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Concepts in Motion

Study Guide 1

Key Concepts

• Bond character varies from ionic to covalent. There is no clear-cut division between the types of bonds.

• Electronegativity — a measure of the attraction that an atom has for shared electrons— can be estimated from the periodic table.

• Electronegativity difference, ΔEN, is a measure of the degree of ionic character in a bond.

• A ΔEN = 2.0 or greater occurs when elements form ionic bonds; ΔEN = 0.5 – 2.0 reflects polar covalent bonds; and ΔEN < 0.5 reflects covalent bonds.

• Metal atoms bond by sharing in a sea of valence electrons.

Study Guide 2

Key Concepts

• Electron pairs about the central atom are either lone (nonbonding) pairs or bonding pairs.

• The polarity of the bonds and the shape of the molecule determine whether a molecule is polar or nonpolar.

• Interparticle forces determine many of the physical properties of substances.

Chapter Assessment 1

The noble gases are considered to have electronegativity values of zero.

A. true

B. false

Chapter Assessment 2

___ is the tendency for the electrons in the inner energy levels to block the attraction of the nucleus for the valence electrons.

A. Conductivity

B. Electronegativity

C. Shielding effect

D. Periodic property

Chapter Assessment 3

When can you assume that the electron on a less electronegative atom is transferred to a more electronegative atom?

A. when the element is a noble gas

B. when the electronegativity difference in a bond is 2.0 or greater

C. when the ∆EN is negative

D. when the electronegativity difference in a bond is less then 2.0

Chapter Assessment 4

Which of the following elements do not form pure covalent bonds?

A. Br2

B. I2

C. NaCl

D. N2

Chapter Assessment 5

A hydride is a compound formed between any element and ___.

A. hydrogen

B. helium

C. oxygen

D. a noble gas

STP 1

Which property is not the result of the way that metal atoms bond with each other?

A. malleability

B. ductility

C. electrical conductivity

D. solubility

STP 2

What kind of bond occurs within a molecule with unequal sharing of electron pairs?

A. ionic bond

B. polar covalent bond

C. non-polar covalent bond

D. hydride bond

STP 3

The two lone pairs of electrons of a water molecule do what to the bond angle between the hydrogen atoms and the oxygen atoms?

A. They attract the hydrogen atoms and increase the angle greater than 109.5°.

B. They occupy more space and squeeze the hydrogen atoms closer together.

C. They do no affect the bond angle.

D. They create structures with more than one correct angle.

STP 4

Which compound forms a perfect tetrahedron?

A. water

B. ammonia

C. methane

D. carbon dioxide

STP 5

___ are hydrocarbons that contain only carbon and hydrogen atoms with single bonds between all the atoms.

A. Alkanes

B. Alkenes

C. Tetrahedrals

D. Alkynes