Karen Timberlake Chapter 5 - PBworks

1

General, Organic, and

Biological ChemistryFourth Edition

Karen Timberlake

Chapter 5Compounds and

Their Bonds

© 2013 Pearson Education, Inc.

© 2013 Pearson Education, Inc. Chapter 3, Section 3

Electrons and Reactivity

Atoms contain

� a very small nucleus packed with neutrons and

positively charged protons.

� a large volume of space around the nucleus that

contains the negatively charged electrons.

It is the electrons that determine the physical and

chemical properties of atoms.

Elements react in order to achieve the same

electronic structure of the nearest noble gas

2

© 2013 Pearson Education, Inc. Chapter 3, Section 3

Electrons and Reactivity

Atoms contain

� a very small nucleus packed with neutrons and

positively charged protons.

� a large volume of space around the nucleus that

contains the negatively charged electrons.

It is the electrons that determine the physical and

chemical properties of atoms.

Elements react in order to achieve the same

electronic structure of the nearest noble gas

3 © 2013 Pearson Education, Inc. Chapter 5, Section 1 4

Ionic and Covalent Bonds

Atoms form octets

� to become more stable.

� by losing, gaining, or

sharing valence electrons.

� by forming ionic or

covalent bonds.

© 2013 Pearson Education, Inc. Chapter 5, Section 1 5

Covalent Compounds

Covalent compounds are formed when nonmetals

share valence electrons forming a covalent bond.

Covalent compounds consist of molecules, which are

discrete groups of atoms, such as:

� water, H2O 2 H atoms + 1 O atom

� carbon dioxide, CO2 1 C atom + 2 O atoms

� glucose, C6H12O6 6 C atoms + 12 H atoms

+ 6 O atoms


Ionic Compounds

Ionic compounds are formed when metal atoms

transfer electrons to nonmetal atoms resulting in the

formation of ionic bonds.

Ionic compounds we use every day include:

� table salt NaCl

� baking soda NaHCO3

� milk of magnesia Mg(OH)2

2


Ionic Compounds, Minerals

Precious and semiprecious gemstones are also

examples of ionic compounds called minerals.

Sapphires and rubies are made of aluminum oxide

(Al2O3), which is an ionic compound. Their brilliant

colors come from the presence of other metals such as

� chromium, which make rubies red and

� iron and titanium, which make sapphires blue


The octet rule is the tendency for atoms to share or transfer electrons to obtain a stable configuration of 8 valence electrons.

Octet Rule


Octet Rule

The octet rule is associated with the stability of thenoble gases except for He, which is stable with 2valence electrons (duet).

Valence Electrons

He 1s2 2

Ne 1s22s22p6 8

Ar 1s22s22p63s23p6 8

Kr 1s22s22p63s23p64s23d104p6 8


Metals Form Positive Ions

Metals form positive ions called

cations

� by a loss of their valence electrons.

� with the electron configuration of

the nearest noble gas.

� that have fewer electrons than

protons.

Group 1A (1) metals ion 1+




Formation of a Sodium Ion, Na+

Sodium achieves an octet by losing its one valence

electron.


Charge of Sodium Ion, Na+

With the loss of its valence

electron, a sodium ion has a 1+

charge.

Na atom Na+

ion

charge

3


Formation of Magnesium Ion, Mg2+

Magnesium achieves an octet by losing its two

valence electrons.


Charge of Magnesium Ion, Mg2+

With the loss of two valence

electrons magnesium forms a

positive ion with a 2+ charge.

Mg atom Mg2+ ion

charge


Formation of Negative Ions

In ionic compounds, nonmetals

� achieve an octet arrangement.

� gain electrons.

� form negatively charged ions called anions with

3–, 2–, or 1– charges.


Formation of a Chloride Ion, Cl–

Chlorine achieves an octet by adding one more

electron to its 7 valence electrons.


Charge of a Chloride Ion, Cl–

By gaining one electron, the

chloride ion has a 1− charge.

charge


Ionic Charge from Group

Numbers

Ions achieve the electron configuration of their

nearest noble gas by forming

� positive ions with a charge equal to its Group

number.

Group 1A (1) = 1+

Group 2A (2) = 2+

Group 3A (3) = 3+

� negative ions with a charge obtained by subtracting

8 or 18 from its Group number.

4


Some Typical Ionic Charges


Group Numbers for Some Positive

and Negative Ions


5.2

Writing names and

formulas for ionic

compounds


Their Bonds



Ionic compounds

� consist of positive and negative ions.

� have attractive forces between the positive and

negative ions called ionic bonds.

� have high melting and boiling points.

� are solid at room temperature.

Ionic Compounds


Salt Is an Ionic Compound

Sodium chloride, or “table salt,” is an example of an

ionic compound.


Formulas of Ionic Compounds

The chemical formula of an ionic compound

represents the element symbols and subscripts, which

represent the lowest whole-number ratio of ions.

In the formula of an ionic compound,

� the sum of positively and negatively charged ions is

always zero Rule of Zero Charge

� the charges are balanced.

5


Charge Balance for NaCl, “Salt”

In NaCl,

� a Na atom loses its valence electron.

� a Cl atom gains an electron.

� the symbol of the metal (sodium) is written first,

followed by the symbol of the nonmetal (chlorine).


Charge Balance In MgCl2

In MgCl2,

� a Mg atom loses two valence electrons.

� two Cl atoms gain one electron each.

� subscripts indicate the number of ions needed to

give charge balance.

� the symbol of the metal (magnesium) is written

first, followed by the symbol of the nonmetal

(chlorine).


Writing Formulas For Binary

Ionic Compounds

� Binary ionic compounds are made from 2

elements, a metal and a nonmetal.

� Charge balance is used to write the formula for

sodium nitride, a compound containing Na+ and

N3−. The metal ion is written first, followed by

nonmetal, subscripts indicate how many of each


Charge Balance in Na2S

In sodium sulfide, Na2S,

� two Na atoms lose one valence electron each.

� one S atom gains two electrons.

� subscripts show the number of ions needed to give

charge balance.

The group of ions with the lowest ratio of ions in an

ionic compound is called a formula unit .


Formula from Ionic Charges

Write the ionic formula of the compound containing

Ba2+ and Cl−.

� Write the symbols of each ion.

Ba2+ Cl−

� Balance the charges.

Ba2+ Cl− (two Cl−−−− needed)

Cl−

� Write the metal first, followed by the nonmetal,

, using a subscript to represent the number of

. Cl− ions.

BaCl2


Writing Names For Binary Ionic

Compounds

The name of an Binary ionic compound

� is made up of 2 elements.

� has the name of the metal ion written first.

� has the name of the nonmetal ion written second

using the first syllable of its element name.

followed by ide.

� has a space separating the name of the metal and

nonmetal.

6


Charges of Representative

Elements


Naming Binary Ionic Compounds


Step 1 Identify the cation and anion.

The cation from Group 2A (2) is Mg2+, and

the anion from Group 5A (15) is N3−.

Step 2 Name the cation by its element name.

Cation: Mg2+ is magnesium

Step 3 Name the anion by using the first syllable of

its element name followed by ide.

Anion: nitrogen becomes nitride, N3−

Step 4 Write the name for the cation first and the

name for the anion second.

Mg3N2 is magnesium nitride.

Naming Mg3N2


What do you notice about these

names and formulas?Chemical Formula Written Form

Fe2O3 iron(III) oxide

FeCl2 iron(II) chloride

PbS lead(II) sulfide

AlP aluminum phosphide

CuF2 copper(II) fluoride

SnI4 tin(IV) iodide

KBr potassium bromide

34


Metals with Variable Ion Charge

Most transition metals (3-12) and Group 4A (14) metals

form 2 or more positive ions, except Zn2+, Ag+, and Cd2+,

which form only one ion.


Metals with Variable Charge

The names of transition metals with two or more

positive ions (cations) use a Roman numeral after the

name of the metal to identify the ion charge.

7


Naming Ionic Compounds with

Variable Charge Metals


Naming FeCl2

Step 1 Determine the charge of the cation from the

anion.

Analyze the Problem.


Naming FeCl2

Step 2 Name the cation by its element name and

use a Roman numeral in parentheses for the

charge.

Fe2+ = iron(II)

Step 3 Name the anion by using the first syllable of

its element name followed by ide .

Cl− = chloride

Step 4 Write the name for the cation first and the

name for the anion second.

iron(II) chloride


Examples of Names of Compounds

with Variable Charge Metals


What do you notice about these

names and formulas?Chemical Formula Written Form

(NH4)2S ammonium sulfide

CoSO4 cobalt (II) sulfate

Fe(OH)3 iron (III) hydroxide

Ca3(PO4)2 calcium phosphate

NH4NO3 ammonium nitrate

41© 2013 Pearson Education, Inc. Chapter 5, Section 1 42

A polyatomic ion

� is a group of atoms.

� has an overall ionic charge.

Examples:

NH4+ ammonium OH− hydroxide

NO3− nitrate NO2

− nitrite

CO32− carbonate PO4

3− phosphate

HCO3− hydrogen carbonate

(bicarbonate)

Polyatomic Ions

8


The names of common polyatomic anions

� end in ate.

NO3− nitrate PO4

3− phosphate

� with one oxygen less, end in ite

NO2− nitrite PO3

3− phosphite

� with hydrogen attached, use prefix hydrogen (or bi).

HCO3− hydrogen carbonate (bicarbonate)

HSO3− hydrogen sulfite (bisulfite)

Some Names of Polyatomic Ions


Compounds Containing

Polyatomic Ions

Polyatomic ions

� must be associated with an ion of opposite charge.

� form ionic bonds with ions of opposite charge to

achieve charge balance.

Example:

Ca2+ NO3−

calcium nitrate ion

charge balance:

Ca(NO3)2

calcium nitrate


Some Compounds with Polyatomic

Ions


Guide to Naming Compounds with

Polyatomic Ions


Step 1 Identify the cation and polyatomic ion(anion).

Cation: K+ Anion: SO42−

Step 2 Name the cation, using a Roman numeral if needed.

K+ = potassium ion

Step 3 Name the polyatomic ion.

SO42− = sulfate ion

Step 4 Write the name or the compound, cation first and the polyatomic ion second.

K2SO4 = potassium sulfate

Name K2SO4


Guide to Writing Formulas for

Compounds with Polyatomic Ions

9


Write the Formula for

Aluminium Hydroxide

Step 1 Identify the cation and polyatomic ion

(anion).

Al3+ and OH−

Step 2 Balance the charges.


Write the Formula for

Aluminium Hydroxide

Step 3 Write the formula, cation first, using the

subscripts from charge balance.

Al(OH)3




Karen Timberlake

5.5

Covalent Compounds:

Sharing Electrons


Their Bonds

© 2013 Pearson Education, Inc.Lectures


Covalent bonds form when atoms

of nonmetals share electrons to

complete octets.

Valence electrons are not

transferred, but shared to achieve

stability.

Covalent Bonds


Formation of H2

In the simplest covalent molecule, H2 , the H atoms

� increase attraction as they move closer.

� share electrons to achieve a stable configuration.

� form a covalent bond.


Formation of H2

10


Electron-Dot Formulas of

Covalent Molecules

In a fluorine (F2) molecule, the F atoms

� share one of their valence electrons.

� acquire an octet.

� form a covalent bond.


Elements That Exist as Diatomic

Molecules

These seven elements share

electrons to form diatomic,

covalent molecules.

BrINClHOF!

Remember it!


The number of covalent bonds a nonmetal forms is

usually equal to the number of electrons it needs to

acquire a stable electron configuration.

Typical bonding patterns for some nonmetals are

shown in the table below.

Bonding Patterns of Some

Nonmetals


� Biological molecules tend to be made out of H,

O,N, and C mostly.

� Bonding tendency rule for H, O, N, C

HONC 1234 Rule

Hydrogen tends to form 1 bond

Oxygen tends to form 2 bonds

Nitrogen tends to form 3 bonds

Carbon tends to form 4 bonds

Bonding Patterns of Some

Nonmetals


Electron-Dot Formulas for Some

Covalent Compounds


Guide to Drawing Electron-Dot

Formulas

11


Step 1 Determine the arrangement of atoms.

In NH3, N is the central atom and is bonded to

three H atoms.

Step 2 Determine the total number

of valence electrons.

Total valence electrons for NH3 = 8 e−

Draw the Electron-Dot Formula

for NH3

H N H

H


Step 3 Attach each bonded atom to the central

atom with a pair of electrons.


for NH3

H N H

H


Step 4 Place the remaining electrons using single

or multiple bonds to complete the octets.

8 valence e− − 6 bonding e− = 2 e− remaining

Use the remaining 2 e− to complete the octet

around the N atom.


for NH3

H N H

H

H N H

H

or


Exceptions to the Octet Rule

Not all atoms have octets.

� Some can have less than an octet, such as

H, which requires only 2 electrons,

B, which requires only 3 electrons, and

Be, which requires only 4 electrons.

� Some can have expanded octets, such as

P, which can have 10 electrons,

S, which can have 12 electrons, and

Cl, Br and I, which can have 14 electrons


Single and Multiple Bonds

In many covalent compounds, atoms share two or

three pairs of electrons to complete their octets.

� In a single bond, one pair of electrons is shared.

� In a double bond, two pairs of electrons are

shared.

� In a triple bond, three pairs of electrons are

shared.


Step 1 Determine the arrangement of atoms.

In CS2, C is the central atom and is bonded to

two S atoms.


for CS2

S C S

12


Step 2 Determine the total number

of valence electrons.

Total valence electrons for


for CS2




A pair of bonding electrons (single bond) is

placed between each S atom and the central C

atom.


for CS2

S C S




16 valence e− - 4 bonding e− = 12 e− remaining

The remaining 12 electrons are placed as six

lone pairs of electrons on both S atoms.

However, this does not complete the octet for

the C atom.


for CS2

S C S


Step 4 Continued: Double and Triple Covalent

Bonds: To complete the octet for the C atom,

it needs to share an additional lone pair from

each of the S atoms, forming a double bond

with each S atom.


for CS2

S C S

or

S C S

S C S


A Nitrogen Molecule has

a Triple Bond

In a nitrogen molecule, N2,

� each N atom shares 3 electrons,

� each N atom attains an octet, and

� the sharing of 3 sets of electrons is called a triple

bond.


Resonance structures are

� two or more electron-dot formulas for the same

arrangement of atoms.

� related by a double-headed arrow ( ).

� written by changing the location of a double bond

between the central atom and a different attached

atom.

Resonance Structures

13


Sulfur dioxide has two resonance structures.

Step 1 Determine the arrangement of atoms. In SO2,

the S atom is the central atom.

O S O

Step 2 Determine the total number of valence

electrons.

Total valence electrons for

Writing Resonance Structures for

SO2





SO2

O S OO S O or




The remaining 14 electrons are drawn as lone

pairs of electrons to complete the octets of the

O atoms, but not the S atom.


SO2

O S O

O S O or


Step 4 Continued: To complete the octet for S, an

additional lone pair from one of the O atoms

is shared to form a double bond.

Because the shared lone pair of electrons can

come from either O atom, two resonance

structures can be drawn.


SO2

O S O

O S O


FNO2, a rocket propellant, has two resonance structures.

One is shown below. What is the other resonance

structure?

Learning Check


FNO2, a rocket propellant, has two resonance structures.

One is shown below. What is the other resonance

structure?

Solution

14




Karen Timberlake

5.6

Naming and Writing

Covalent Formulas


Their Bonds



Names of Covalent Compounds

When naming a covalent compound,

� the first nonmetal in the formula is named by its

element name, and

� the second nonmetal is named using the first syllable

of its element name, followed by ide.

Prefixes are used

� to indicate the number of atoms present for each

element in the compound, and

� because two nonmetals can form two or more

different compounds


Prefixes Used in Naming

Covalent Compounds


Some Covalent Compounds


Guide to Naming Covalent

Compounds with Two Nonmetals


Naming Covalent Compounds,

SO2


Step 1 Name the first nonmetal by its element

name. In SO2, the first nonmetal (S) is sulfur.

Symbols of

Elements

S O

Name sulfur oxygen

Subscripts 1 2

Prefix (none) understood di

15


Naming Covalent Compounds,

SO2

Step 2 Name the second nonmetal by using the first

syllable of its name followed by ide. The

second nonmetal (O) is named oxide.

Step 3 Add prefixes to indicate the number of

atoms (subscripts). Because there is

one sulfur atom, no prefix is needed. The

subscript two for the oxygen atoms is written as

the prefix di .

The name of SO2 is sulfur dioxide.


Learning Check

Name the covalent compound P2O5.


Solution

Name the covalent compound P2O5.


Step 1 Name the first nonmetal by its element

name. In P2O5, the first nonmetal (P) is

phosphorus.

Symbols of Elements P O

Name phosphorus oxygen

Subscripts 2 5

Prefix di penta


Solution

Step 2 Name the second nonmetal by using the first

syllable of its name followed by ide. The

second nonmetal (O) is named oxide.

Step 3 Add prefixes to indicate the number of

atoms (subscripts). The subscript for the two

P atoms is di; the subscript for the five oxygen

atoms is penta.

The name of P2O5 is phosphorus pentoxide.

When the vowels o and o, or a and o appear

together (pentaoxide), the first vowel is omitted.


Guide to Writing Formulas for

Covalent Compounds


Write the formula for sulfur hexafluoride.


Step 1 Write the symbols in order of the

elements in the name.

S F

Step 2 Write any prefixes as subscripts.

Prefix hexa = 6 Formula: SF6

Writing Formulas of Covalent

Compounds

Name sulfur hexafluoride

Symbols of Elements S F

Subscripts 1 6 (from hexa)

16




Karen Timberlake

The Shape of Molecules


Their Bonds



Valence-Shell Electron-Pair Repulsion

Theory (VSEPR)

In the valence-shell electron-pair repulsion (VSEPR)

theory, the electron groups around a central atom

� are arranged as far apart from each other as possible.

� have the least amount of electron-electron repulsion.

� are used to predict the molecular shape.


Molecular Shape – why does it

matter?

• Watson and Crick withtheir famous 3-D model of DNA

• How did they know? •Shape of molecules important as it influences their physical and chemical behavior


What Determines the 3-D Shape of

a Molecule?

� Need to understand electrostatic repulsion

(electron clouds around the central atom in a

molecule are going to repel each other)

� Need to draw Lewis structures

� Need to differentiate between lone pair

electrons and bonding pair electrons around

the central atom in a molecule

� Need to apply a theory based on all of the

above


Lewis Structure Lone pair e- (nonbonding e- domain)

Bonding pair e- (bonding domain)

• Electron pairs are also called domains

• Electron domains try to stay out of each

others way (to minimize repulsion)


Valence Shell Electron Pair

Repulsion Theory

� VSEPR, for short

� Proposed by English chemist Ron Gillespie in

the 1950s

� Based on Lewis structures of molecules and

electron repulsion

17


VSEPR Theory Rules

1. In a molecule, electron domains (pairs) will

orient themselves around the central atom

in an arrangement that minimizes the

repulsions among them.

2. The three dimensional orientation of the

outer atoms around the central atom

depends on the numbers and types of e-

domains

3. Lone pair electrons count as 1 electron

domain; single, double or triple bonding pair

electrons count as 1 electron domain.


Shapes of Molecules

The three-dimensional shape of a molecule can be

predicted by

� the number of bonding groups (bonding domains) and

lone-pair electrons (nonbonding domains) around the

central atom in the molecule

� VSEPR theory (valence-shell-electron-pair repulsion).


Molecules with Two Electron Domains

In a molecule of BeCl2,

� there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule).

� to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other.

� the shape of the molecule is linear.


Molecules With Three Electron

Domains

In a molecule of BF3,

� three electron groups are bonded to the central atom

B (B is an exception to the octet rule).

� repulsion is minimized with 3 electron groups at

angles of 120°.

� the shape is trigonal planar.


Three Electron Domains with a

Lone Pair

In a molecule of SO2,

� S has 3 electron groups; 2 electron groups bonded to

O atoms and one lone pair.

� repulsion is minimized with the electron groups at

angles of 120°, a trigonal planar arrangement.

� the shape is determined by the two O atoms bonded

to S, giving SO2 a bent (~120°°°°) shape.


Four Electron Domains

In a molecule of CH4,

� there are four electron groups around C.

� repulsion is minimized by placing four electron groups

at angles of 109°, which is a tetrahedral arrangement.

� the four bonded atoms form a tetrahedral shape.

18


Four Electron Groups with a Lone

Pair

In a molecule of NH3,

� three electron groups bond to H atoms, and the fourth

one is a lone (nonbonding) pair.

� repulsion is minimized with 4 electron groups in a

tetrahedral arrangement.

� the three bonded atoms form a pyramidal (~109°°°°)

shape.


Four Electron Groups with Two

Lone Pairs

In a molecule of H2O,

� two electron groups are bonded to H atoms and two

are lone pairs (4 electron groups).

� four electron groups minimize repulsion in a


� the shape with two bonded atoms is bent (~109°°°°).


Molecular Shapes for 2 and 3

Bonded Atoms


Molecular Shapes for 4 Bonded

Atoms


Guide to Predicting Molecular

Shape (VSEPR Theory)


Predicting the Molecular Shape of

H2Se

Step 1 Draw the electron-dot formula. In the

electron-dot formula for H2Se, there are four

electron groups, including two lone pairs of

electrons around Se.

19


Predicting the Molecular Shape of

H2Se

Step 2 Arrange the electron groups around the

central atom to minimize repulsion. The

four electron groups around Se would have a


Step 3 Use the atoms bonded to the central atom

to determine the molecular shape. Two

bonded atoms give H2Se a bent shape with a

bond angle of ~109°.




Karen Timberlake

A closer look at

bonding type – the

concept of

electronegativity


Their Bonds



General Rules for Bond Type

Nonmetal

Nonmetal

Covalent

Bond

Metal

Ionic

Bond

Negative complex

ions


Ionic or Covalent?

� So far, we have been employing the following

generalization for the type of bonding holding

atoms together in a substance

Generalization

� metal + nonmetal = ionic bond

� nonmetal + nonmetal = covalent bond

Reality

� Bonding type is on a continuum, from 100% ionic

to 100% covalent!

100% covalent 100% ionic


Bonding Conundrum

� Arose in the 1930’s from American Chemist

Linus Pauling’s work on the nature of the

chemical bond.

� Pauling realized that there was some “in-

betweeness” with the bonding in many

substances

� Calculated the ability of elements to attract

electrons to themselves when bonding

� Called this ability the Electronegativity of the

element


Determining Predominant Bond type

� Electronegativity is a relative measure of the

tendency for atoms of an element to attract

electrons to themselves in a chemical bond.

� Originated with American chemist Linus Pauling

(1901-1994), a 2x Nobel Prize winner who did

most of his work at Cal Poly Tech, but finished

his career at Stanford

20

© 2013 Pearson Education, Inc. Chapter 5, Section 1 115 © 2013 Pearson Education, Inc. Chapter 5, Section 1 116

Electronegativity (EN) Values of the

Elements

Scale ranges from 0.7 to 4.0


Electronegativity

� is a measure of an atom’s ability to attract

electrons to itself in a chemical bond

� increases from left to right, going across a period

on the periodic table.

� decreases going down a group on the periodic

table.

� is high for the nonmetals, with fluorine as the

highest.

� is low for the metals and transition metals.

Electronegativity


Trends in Periodic table


Periodic trends in electronegativity

� Atomic radius increases, valence e- further

away from nucleus, decreases pulling power

of nucleus

Electronegativity decreases down a group

because:

Electronegativity increases from left to right across

a period because:

� Atomic radius decreases, valence e- closer to

the nucleus, increases pulling power of nucleus


Bonding and Electronegativity

� According to Pauling, the difference in the electronegativity

values of the two atoms involved in a chemical bond can

be used to predict the type of bond that forms.

� Pauling Classified chemical bonds into the following three

types, based on differences in electronegativity

� Ionic

� Polar Covalent

� Nonpolar Covalent

21


Electronegativity and Bond Type

� It is the difference in the electronegativies (∆EN) of

two bonded atoms that determines the

predominant bond type

∆EN ≤ 0.4 � Nonpolar covalent bond

(e- shared equally between atoms)

∆EN 0.5 to 1.7 polar covalent bond,

(e- shared unequally)

∆EN > 1.7 ionic bond (e- transferred)


Bond Character Summary

a) Nonpolar covalent bond �

electrons shared equally

b) Polar covalent bond � e- not

shared equally, more

electronegative atom has greater

share of e- cloud

c) Ionic bond – e- transferred from

one atom to other, creates a + and

– ion


A nonpolar covalent bond

� occurs between nonmetals.

� has an equal or almost equal sharing of electrons.

� has almost no electronegativity difference (0.0 to

0.4).

Examples:

Atoms Electronegativity Type of BondDifference

Nonpolar Covalent Bonds


Example of Nonpolar Covalent

Bonds� Diatomic elements e.g. H2

� Bonds between P and H = 2.1-2.1=0


A polar covalent bond

� occurs between nonmetal atoms that do not share

electrons equally.

� has a moderate electronegativity difference

(0.5 to 1.7).

Examples:

Atoms Electronegativity Type of Bond

Difference

Polar Covalent Bonds


Comparing Nonpolar and Polar

Covalent Bonds

22


Result of differences in electronegativity:

• More electronegative element has greater

share of e- cloud, it is electron rich

•Less electronegative element is e- poor

•Creation of partial charges on the atoms in the

bond (bond dipoles) δδδδ+ and δδδδ -


H Cl

Example: Bond between H and Cl

• Electronegativity values H= 2.1, Cl = 3.0

• e- cloud pulled towards Cl atom = Polar Covalent

Bond

electron rich

region

electron poor

region

e- rich

e- poor

ClH

δ+ δ-

9.5

Direction e- cloud being pulled

2.1 3.0


Bond Polarity and Dipoles

Bonds become more polar as the difference in

electronegativity values of bonding atoms increases.

Polar covalent bonds have a separation of charges

called a dipole.

The positive and negative ends of the dipole are

indicated by the lowercase Greek letter delta with a

positive or negative sign, δ+ and δ-, or an arrow that

points from the positive to the negative charge.


Ionic Bonds

An ionic bond

� occurs between metal and nonmetal ions.

� is a result of electron transfer.

� has a large electronegativity difference (1.8 or more).

Examples:

Atoms Electronegativity Type of Bond Difference


Determining Predominant Bond Type

� Look at the electronegativity values of atoms

involved in bond

� Calculate electronegativity difference (∆EN)

� If the bond is polar covalent, draw arrow in

direction that e- cloud is pulled (more

electronegative atom)

� Also use lower case Greek symbol for delta, δ,

along with + or - sign


Electronegativity and Bond Types

23


Predicting Bond Types


Use the electronegativity difference to identify the type

of bond between the following atoms as nonpolar

covalent (NP), polar covalent (P), or ionic (I).

A. K–N

B. N–O

C. Cl–Cl

D. B–Cl

Learning Check


Use the electronegativity difference to identify the type

of bond between the following atoms as nonpolar

covalent (NP), polar covalent (P), or ionic (I).

A. K–N 2.2 ionic (I)

B. N–O 0.5 polar covalent (P)

C. Cl–Cl 0.0 nonpolar covalent (NP)

D. B–Cl 1.0 polar covalent (P)

Solution


Polar Molecules

A polar molecule

� contains polar bonds.

� has a separation of positive and negative charge. called a dipole, indicated with δ+ and δ–.

� has dipoles that do not cancel.δ+ δ–

H–Cl NH3

dipole

Dipoles do not cancel.


Nonpolar Molecules

A nonpolar molecule

� contains nonpolar bonds

Cl–Cl H–H

� or has a symmetrical arrangement of polar bonds.


Guide to Determination of

Polarity

24


Molecular Polarity, H2O

Determine the polarity of the H2O molecule.

Step 1 Determine if the bonds are polar covalent or

nonpolar covalent. From the electronegativity

table, O 3.5 and H 2.1 gives a difference of 1.4,

which makes the O — H bonds, polar covalent.


Molecular Polarity, H2O

Determine the polarity of the H2O molecule.

Step 2 If the bonds are polar covalent, draw the

electron-dot formula and determine if the

dipoles cancel or not. The four electron

groups of oxygen are bonded to two H atoms.

Thus, the H2O molecule has a net dipole, which

makes it a polar molecule.


Polar Molecules

A polar molecule

� contains polar bonds.

� has a separation of positive and negative charge. called a dipole, indicated with δ+ and δ–.

� has dipoles that do not cancel.δ+ δ–

H–Cl NH3

dipole

Dipoles do not cancel.


Learning Check

Determine the shape of each of the following molecules

and whether they are polar or nonpolar. Explain.

1. PBr3

2. HBr

3. Br2

4. SiBr4


Solution

Determine the shape of each of the following molecules

and whether they are polar or nonpolar. Explain.

1. PBr3 pyramidal; polar; dipoles don’t cancel

2. HBr linear; polar; one polar bond (dipole)

3. Br2 linear; nonpolar; nonpolar bond

4. SiBr4 tetrahedral; nonpolar; dipoles cancel




Karen Timberlake

5.9

Attractive Forces in

Compounds


Their Bonds


25


States of Matter

The fundamental difference between states of

matter is the distance between particles.


States of Matter

Because in the solid and liquid states

particles are closer together, we refer to them

as condensed phases.


The States of Matter

� The state a substance is

in at a particular

temperature and

pressure depends on

two antagonistic entities:

� The kinetic energy of the

particles

� The strength of the

attractions between the

particles


Ionic Compounds

In ionic compounds, ionic bonds

� require large amounts of energy to break.

� hold positive and negative ions together.

� explain their high melting points.


What about covalent compounds?

The attractions between molecules are not

nearly as strong as the intramolecular

attractions that hold compounds together.


Intermolecular Forces

They are, however, strong enough to control

physical properties such as the state a

molecular compound will be in at room

temperature, its melting and boiling point.

26



In addition, they are responsible for the

shapes of protein molecules, DNA, water’s

unique properties, the structure of the cell

membrane, etc


Covalent Compounds

In covalent compounds, the attractive forces between

solid and liquid molecules

� are weaker than ionic bonds.

� require less energy to break.

� explain why their melting points are lower than ionic

compounds.

These attractive forces include

� dipole-dipole attractions,

� dispersion forces, and

� hydrogen bonding.



They are, however, strong enough to control

physical properties such as boiling and

melting points, vapor pressures, and

viscosities.


Dipole-Dipole Attractions

In covalent compounds, polar molecules exert attractive

forces between molecules called dipole-dipole

attractions. dipole dipole forces


Dipole-Dipole Attractions,

Hydrogen Bonds

In covalent compounds, some polar molecules form

strong dipole attractions called hydrogen bonds, which

occur between the partially positive hydrogen atom of

one molecule and a lone pair of electrons on a nitrogen,

oxygen, or fluorine atom in another molecule.


Dispersion Forces

Dispersion forces are

� weak attractions between nonpolar molecules.

� caused by temporary dipoles that develop when

electrons are not distributed equally.

Nonpolar molecules form attractions when they form temporary

dipoles.

27


London Dispersion Forces

While the electrons in the 1s orbital of helium

would repel each other (and, therefore, tend

to stay far away from each other), it does

happen that they occasionally wind up on the

same side of the atom.



At that instant, then, the helium atom is polar,

with an excess of electrons on the left side

and a shortage on the right side.



Another helium nearby, then, would have a

dipole induced in it, as the electrons on the

left side of helium atom 2 repel the electrons

in the cloud on helium atom 1.



London dispersion forces, or dispersion

forces, are attractions between an

instantaneous dipole and an induced dipole.



� These forces are present in all molecules, whether they are polar or nonpolar.

� The tendency of an electron cloud to distort in this way is called polarizability.


Factors Affecting London Forces

� The shape of the molecule

affects the strength of dispersion

forces: long, skinny molecules

(like n-pentane tend to have

stronger dispersion forces than

short, fat ones (like neopentane).

� This is due to the increased

surface area in n-pentane.

28


Factors Affecting London Forces

� The strength of dispersion forces tends to

increase with increased molecular weight.

� Larger atoms have larger electron clouds,

which are easier to polarize.


Which Have a Greater Effect:Dipole-Dipole Interactions or Dispersion Forces?

� If two molecules are of comparable size and

shape, dipole-dipole interactions will likely be

the dominating force.

� If one molecule is much larger than another,

dispersion forces will likely determine its

physical properties.


Comparison of Bonding and

Attractive Forces


Intermolecular Forces Affect Many

Physical Properties

The strength of the

attractions between

particles can greatly

affect the properties

of a substance or

solution.


Melting Points and Attractive

Forces

The stronger the attractive force between ions or

molecules, the higher the melting points.

Ionic compounds, have the strongest attractive force

and, therefore the highest melting points.

Covalent molecules have less attractive forces than ionic

compounds and, therefore lower melting points.


Melting Points and Attractive

Forces

The attractive forces between covalent molecules vary

in magnitude; the stronger the attractive force, the

higher its melting point.

� Hydrogen bonds are the strongest type of dipole–

dipole attractions, requiring the most energy to

break, followed by dipole–dipole forces.

� Dispersion forces are the weakest, requiring even

less energy to break them, and therefore have lower

melting points than hydrogen bonds and dipole–

dipole forces.

29


Melting Points of Selected

Substances


Learning Check

Identify the main type of attractive forces for each of the

following compounds: ionic bonds, dipole–dipole,

hydrogen bonds or dispersion.

1. NCl3

2. H2O

3. Br2

4. KCl

5. NH3


Solution

Identify the main type of attractive forces for each of the

following compounds: ionic bonds, dipole–dipole,

hydrogen bonds or dispersion.

1. NCl3 dipole–dipole

2. H2O hydrogen bonds

3. Br2 dispersion

4. KCl ionic bonds

5. NH3 hydrogen bonds

Karen Timberlake Chapter 5 - PBworks

Documents

Transcript of Karen Timberlake Chapter 5 - PBworks