Chapter 3

Structure and Properties of Ionic and Covalent Compounds

3.1 Chemical Bonding

•

Chemical bond

- the force of attraction between any two atoms in a compound • This attractive force overcomes the repulsion of the positively charged nuclei of the two atoms participating in the bond • Interactions involving valence electrons are responsible for the chemical bond

Lewis Symbols

• Lewis symbol (Lewis structure) - a way to represent atoms using the element symbol and valence electrons as dots • As only valence electrons participate in bonding, this makes it much easier to work with the octet rule • The number of dots used corresponds directly to the number of valence electrons located in the outermost shell of the atoms of the element

Lewis Symbols

• Each “side” of the symbol represents an atomic orbital, which may hold up to two electrons • Using Lewis symbols o Place one dot on each side until there are four dots around the symbol o Now add a second dot to each side in turn o The number of valence electrons limits the number of dots placed o Each unpaired dot (unpaired electron of the valence shell) is available to form a chemical bond

Lewis Dot Symbols for Representative Elements

Principal Types of Chemical Bonds: Ionic and Covalent •

Ionic bond

- a transfer of one or more electrons from one atom to another • Forms attractions due to the opposite charges of the atoms •

Covalent bond

- attractive force due to the sharing of electrons between atoms • Some bonds have characteristics of both types and not easily identified as one or the other

Ionic Bonding

• Representative elements form ions that obey the octet rule • Ions of opposite charge attract each other creating the ionic bond • When electrons are lost by a metal and electrons are gained by a nonmetal o Each atom achieves a “Noble Gas” configuration o 2 ions are formed; a cation and anion, which are attracted to each other

Ionic Bonding

Consider the formation of NaCl Na + Cl  NaCl Sodium has a low ionization energy; it readily loses an electron Na   Na + + e Chlorine has a high electron affinity Propensity is to gain an electron to achieve a full octet.

: ..

Cl ..

  e     : ..

Cl ..

:   

Essential Features of Ionic Bonding

• Atoms with low I.E. and low E.A. tend to form positive ions • Atoms with high I.E. and high E.A. tend to form negative ions • Ion formation takes place by electron transfer • The ions are held together by the electrostatic force of the opposite charges • Reactions between metals and nonmetals (representative elements) tend to be ionic

Ion Arrangement in a Crystal

• As a sodium atom loses one electron, it becomes a smaller sodium ion • When a chlorine atom gains that electron, it becomes a larger chloride ion • Attraction of the Na cation with the Cl anion forms NaCl ion pairs that aggregate into a crystal

Covalent Bonding

Let’s look at the formation of H 2 : H + H  H 2 • Each hydrogen has one electron in its valance shell • If it were an ionic bond it would look like this: H   H   H      • However, both hydrogen atoms have an equal tendency to gain or lose electrons • Electron transfer from one H to another usually will not occur under normal conditions

• Instead, each atom attains a noble gas configuration by

sharing

electrons H    H  H : H Each hydrogen atom now has two electrons around it and attained a He configuration The shared electron pair is called a

Covalent Bond

Covalent Bonding in Hydrogen

Features of Covalent Bonds

• Covalent bonds form between atoms with similar tendencies to gain or lose electrons • Compounds containing covalent bonds are called

covalent compounds

or molecules • The diatomic elements have

completely covalent

bonds (totally equal sharing) – H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , I 2 : ..

F ..

   ..

F ..

:  : ..

F ..

: ..

F ..

: Each fluorine is surrounded by 8 electrons – Ne configuration

Examples of Covalent Bonding

2H    ..

O ..

  H : ..

O ..

: H 2e – 6e – from 2H from O 2e – 8e – for H for O 4H     .

C   H : H ..

C  H : 4e – 4e – from 4H from C 2e – 8e – H for H for C

Polar Covalent Bonding and Electronegativity

• The Polar Covalent Bond o Ionic bonding involves the

transfer

electrons of o Covalent bonding involves the

sharing

electrons of o

Polar covalent bonding

of

unequally

shared - bonds made up electron pairs

Electronegativity

•

Electronegativity

- a measure of the ability of an atom to attract electrons in a chemical bond • Elements with high electronegativity have a greater ability to attract electrons than do elements with low electronegativity • Consider the covalent bond as competition for electrons between 2 positive centers o The difference in electronegativity determines the extent of bond polarity

 somewhat positively charged H    ..

F :  ..

H : F : somewhat negatively charged • The electrons spend more time with fluorine • This sets up a polar covalent bond • A truly covalent bond can only occur when both atoms are identical These two electrons are not shared equally

Polar Covalent Bonding in Water

• Oxygen is electron rich =  • Hydrogen is electron deficient =  + • This results in unequal sharing of electrons in the pairs = polar covalent bonds • Water has 2 covalent bonds

Electronegativities of Selected Elements

• The most electronegative elements are found in the upper right corner of the periodic table • The least electronegative elements are found in the lower left corner of the periodic table electronegativity increases

3.2 Naming Compounds and Writing Formulas of Compounds

•

Nomenclature

- the assignment of a correct and unambiguous name to each and every chemical compound • Two naming systems: o ionic compounds o covalent compounds

Formulas of Compounds

• A formula is the representation of the fundamental compound using chemical symbols and numerical subscripts o The formula identifies the number and type of the various atoms that make up the compound unit o The number of like atoms in the unit is shown by the use of a subscript o Presence of only one atom is understood when no subscript is present

Ionic Compounds

• Metals and nonmetals usually react to form ionic compounds • The metals are cations and the nonmetals are anions • The cations and anions arrange themselves in a regular three-dimensional repeating array called a crystal lattice • Formula of an ionic compound is the smallest whole-number ratio of ions in the substance

Writing Formulas of Ionic Compounds from the Identities of the Component Ions • Determine the charge of each ion o Metals have a charge equal to group number o Nonmetals have a charge equal to the group number minus eight • Cations and anions must combine to give a formula with a net charge of zero o It must have the same number of positive charges as negative charges

Predict Formulas

Predict the formula of the ionic compounds formed from combining ions of the following pairs of elements: 1. sodium and oxygen 2. lithium and bromine 3. aluminum and oxygen 4. barium and fluorine

Writing Names of Ionic Compounds from the Formula of the Compound • Name the cation followed by the name of the anion • A positive ion retains the name of the element; change the anion suffix to -

ide

Writing Names of Ionic Compounds from the Formula of the Compound • • If the cation of an element has several ions of different charges (as with transition metals) use a Roman numeral following the metal name “Stock notation” • Roman numerals give the

charge

of the metal • • • Examples: FeCl 3 is iron (III) chloride FeCl 2 is iron (II) chloride CuO is copper (II) oxide

Common Nomenclature System

• Use

-ic

to indicate the higher of the charges that ion might have • Use -

ous

to indicate the lower of the charges that ion might have • Examples: • FeCl 2 is ferrous chloride • FeCl 3 is ferric chloride

Stock and Common Names for Iron and Copper Ions

Common Monatomic Cations and Anions

•

Monatomic ions

- ions consisting of a single charged atom

Polyatomic Ions

•

Polyatomic ions

- ions composed of 2 or more atoms bonded together with an overall positive or negative charge o Within the ion itself, the atoms are bonded using covalent bonds o The positive and negative ions will be bonded to each other with ionic bonds • Examples: • • NH 4 + SO 4 2 ammonium ion sulfate ion

Common Polyatomic Cations and Anions

Name These Compounds

1. NH 4 Cl 2. BaSO 4 3. Fe(NO 3 ) 3 4. CuHCO 3 5. Ca(OH) 2

Writing Formulas of Ionic Compounds from the Name of the Compound • Determine the charge of each ion • Write the formula so that the resulting compound is neutral • Example: Barium chloride: Barium is +2, Chloride is -1 Formula is BaCl 2

Determine the Formulas from Names

Write the formula for the following ionic compounds: 1. sodium sulfate 2. ammonium sulfide 3. magnesium phosphate 4. chromium(II) sulfate

Covalent Compounds

• • Covalent compounds are typically formed from nonmetals

Molecules

- compounds characterized by covalent bonding • Not a part of a massive three-dimensional crystal structure • Exist as discrete molecules in the solid, liquid, and gas states

Naming Covalent Compounds

1. The names of the elements are written in the order in which they appear in the formula 2. A prefix indicates the number of each kind of atom

3.

Naming Covalent Compounds

If only one atom of a particular element is present in the molecule, the prefix mono- is usually omitted from the first element 4.

5.

Example: CO is carbon

mon

oxide The stem of the name of the last element is used with the suffix –ide

The final vowel in a prefix is often dropped before a vowel in the stem name

Name These Covalent Compounds

1. SiO 2 2. N 2 O 5 3. CCl 4 4. IF 7

Writing Formulas of Covalent Compounds

• Use the prefixes in the names to determine the subscripts for the elements • Examples: • • nitrogen trichloride diphosphorus pentoxide NCl P 2 O 3 5 • Some common names that are used: H 2 O NH 3 C 2 H 5 OH C 6 H 12 O 6 water ammonia ethanol glucose

Provide Formulas for These Covalent Compounds

1. nitrogen monoxide 2. dinitrogen tetroxide 3. diphosphorus pentoxide 4. nitrogen trifluoride

3.3 Properties of Ionic and Covalent Compounds

• Physical State o Ionic compounds are usually solids at room temperature o Covalent compounds can be solids, liquids, and gases • Melting and Boiling Points o

Melting point

- the temperature at which a solid is converted to a liquid o

Boiling point

- the temperature at which a liquid is converted to a gas

Physical Properties

• Melting and Boiling Points o Ionic compounds have much higher melting points and boiling points than covalent compounds o A large amount of energy is required to break the electrostatic attractions between ions o Ionic compounds typically melt at several hundred degrees Celsius • Structure of Compounds in the Solid State o Ionic compounds are crystalline o Covalent compounds are crystalline or amorphous – having no regular structure

Physical Properties

• Solutions of Ionic and Covalent Compounds o Ionic compounds often dissolve in water, where they

dissociate

- form positive and negative ions in solution o

Electrolytes

- ions present in solution allowing the solution to conduct electricity o Covalent solids usually do not dissociate and do not conduct electricity -

nonelectrolytes

Comparison of Ionic vs. Covalent Compounds Composed of Electrons Physical state Dissociation Ionic Covalent Metal + nonmetal 2 nonmetals Transferred Solid / crystal Yes, electrolytes Boiling/Melting High Shared Any / crystal OR amorphous No, nonelectrolytes Low

3.4 Drawing Lewis Structures of Molecules and Polyatomic Ions

Lewis Structure Guidelines

1. Use chemical symbols for the various elements to write the skeletal structure of the compound o The least electronegative atom will be placed in the central position o Hydrogen and halogens occupy terminal positions o Carbon often forms chains of carbon-carbon covalent bonds

Lewis Structure Guidelines

2.

o Determine the number of valence electrons associated with each atom in the compound Combine these valence electrons to determine the total number of valence electrons in the compound o Polyatomic cations, subtract one electron for every positive charge o Polyatomic anions, add one electron for every negative charge o This total is the number of electrons that are available for bonding.

Lewis Structure Guidelines

• Add up the total number of electrons each atom needs to have a full octet.

o For hydrogen, this number is two o For any other representative element, the number is eight.

• The total of all is the number of electrons needed.

Lewis Structure Guidelines

• Subtract the number needed by the number available.

o This gives the number of electrons that must be shared.

• Divide the number to be shared by two and that is how many bonds the compound must have.

Lewis Structure Guidelines

3. Connect the central atom to each of the • • • • surrounding atoms using electron pairs Next, complete octets of all the atoms bonded to the central atom Hydrogen needs only two electrons Electrons not involved in bonding are represented as lone pairs Total number of electrons in the structure must equal the number of valence electrons in step 2

Lewis Structure Guidelines

4. Count the number of electrons you have and compare to the number you used • • If they are the same , you are finished If you used more electrons than you have, add a bond for every two too many you used • • Then, give every atom an octet If you used less electrons than you have ….

see later exceptions to the octet rule Recheck that all atoms have the octet rule satisfied and that the total number of valance electrons are used

Drawing Lewis Structures of Covalent Compounds

Draw the Lewis structure of carbon dioxide, CO 2 Draw a skeletal structure of the molecule 1.

Arrange the atoms in their most probable order 2.

3.

4.

C-O-O and/or O-C-O Find the electronegativity of O=3.5 & C=2.5

Place the least electronegative atom as the central atom, here carbon is the central atom Result is the O-C-O structure from above

Drawing Lewis Structures

5.

Find the number of valence electrons for each atom and the total for the compound 1 C atom x 4 valence electrons = 4 e 2 O atoms x 6 valence electrons = 12 e 16 e total 6.

Find the number of electrons needed.

1C atom x 8 e = 8 e 2O atoms x 8e = 16e 24e needed Number of shared electrons is: 24e - 16e = 8e A covalent bond always takes 2e so there must be four bonds in CO 2 .

Lewis Structures Practice

Using the guidelines presented, write Lewis structures for the following: 1. H 2 O 2. NH 3 3.NH

4 + 4. CO 3 2 5. N 2

Lewis Structure, Stability, Multiple Bonds, and Bond Energies

•

Single bond

- one pair of electrons are shared between two atoms •

Double bond

- two pairs of electrons are shared between two atoms •

Triple bond

- three pairs of electrons are shared between two atoms • Very stable

H : H or H H ..

 :: ..

 or O  O ..

N   ..

N or N  N

Bond energy

- the amount of energy required to break a bond holding two atoms together triple bond > double bond > single bond

Bond length

- the distance separating the nuclei of two adjacent atoms single bond > double bond > triple bond

Lewis Structures and Resonance

• Write the Lewis structure of CO 3 2 • Where should the double bond go?

• In some cases it is possible to write more than one Lewis structure that satisfies the octet rule for a particular compound

• : ..

O : : O : : ..

O : ..

 : : C :: O  : ..

 : :: C : ..

 :  : ..

O : :: C : : ..

O : Experimental evidence shows all bonds are the same length, meaning there is not really any double bond in this ion • None of theses three Lewis structures exist, but the

actual

structure is an average or

hybrid

of these three Lewis structures •

Resonance

- two or more Lewis structures that contribute to the real structure

Lewis Structures and Exceptions to the Octet Rule

1.

Incomplete octet - less then eight electrons around an atom other than H • Let’s look at BeH 2 1 Be atom x 2 valence electrons = 2 e 2 H atoms x 1 valence electrons = 2 e total 4 e • Resulting Lewis structure: H

:

Be

:

H or H

–

Be

–

H

Odd Electron

2.

• • Odd electron - if there is an odd number of valence electrons, it is not possible to give every atom eight electrons Let’s look at NO, nitric oxide It is impossible to pair all electrons as the compound contains an ODD number of valence electrons

N - O

Expanded Octet

• 3.

Expanded octet an element in the 3rd period or below may have 10 and 12 electrons around it • • Expanded octet is the most common exception Consider the Lewis structure of PF 5 Phosphorus is a third period element • 1 P atom x 5 valence electrons = 5 e 5 F atoms x 7 valence electrons = 35 e 40 e total Distributing the electrons results in this Lewis structure

Lewis Structures and Molecular Geometry: VSEPR Theory

• Molecular shape plays a large part in determining properties and shape • VSEPR theory -

R

epulsion theory

V

alance

S

hell

E

lectron

P

air • Used to predict the shape of the molecules • All electrons around the central atom arrange themselves so they can be as far away from each other as possible – to minimize electronic repulsion

VSEPR Theory

• In the covalent bond, bonding electrons are localized around the nucleus • The covalent bond is

directional ,

having a specific orientation in space between the bonded atoms • Ionic bonds have electrostatic forces which have no specific orientation in space

• Consider BeH 2 o Only 4 electrons surround the beryllium atom o These 2 electron pairs have minimal repulsion when located on opposite sides of the structure o

Linear structure

having bond angles of 180 °

• Consider BF 3 o There are 3 shared electron pairs around the central atom o These electron pairs have minimal repulsion when placed in a plane, forming a triangle o

Trigonal planar structure

with bond angles of 120 °

Basic Electron Pair Repulsion of

•

a Full Octet

• Consider CH 4 o There are 4 shared electron pairs around the central Carbon o Minimal electron repulsion when electrons are placed at the four corners of a tetrahedron o Each H-C-H bond angle is 109.5

°

Tetrahedron

is the primary structure of a full octet

Basic Electron Pair Repulsion of a Full Octet with One Lone Pair

Consider NH 3 • There are 4 electron pairs around the central Nitrogen • 3 pairs are shared electron pairs • 1 pair is a lone pair   A lone pair is more electronegative with a greater electron repulsion The lone pair takes one of the corners of the tetrahedron without being visible, distorting the arrangement of electron pairs • Ammonia has a

trigonal pyramidal

angles structure with 107 °

Basic Electron Pair Repulsion of a Full Octet with Two Lone Pairs

Consider H 2 O • There are 4 electron pairs around the central Oxygen • 2 pairs are shared electron pairs • 2 pairs are lone pairs o o All 4 electron pairs are approximately tetrahedral to each other The lone pairs take two of the corners of the tetrahedron without being visible, distorting the arrangement of electron pairs • Water has a

bent

angles or

angular

structure with 104.5

° bond

Predicting Geometric Shape Using Electron Pairs

Basic Procedure to Determine Molecular Shape

1. Write the Lewis structure 2. Count the number of shared electron pairs and lone pairs around the central atom 3. If no lone pairs are present, shape is: • 2 shared pairs - linear • • 3 shared pairs - trigonal planar 4 shared pairs - tetrahedral 4. Look at the arrangement and name the shape • Linear • • • • Trigonal planar Bent Trigonal pyramid Tetrahedral

More Complex Molecules

Consider dimethyl ether • Has 2 different central atoms: • oxygen • carbon o CH 3 (methyl group) has tetrahedral geometry (like methane) o

Portion of the molecule linking the two methyl groups would bond angles similar to water

Determine the Molecular Geometry

• PCl 3 • SO 2 • PH 3 • SiH 4

Molecular shape and Polarity

• A molecule is

polar

if its centers of positive and negative charges do not coincide • Polar molecules when placed in an electric field will align themselves in the field • Molecules that are polar behave as a dipole (having two “poles” or ends) • One end is positively charged the other is negatively charged • Nonpolar molecules will not align themselves in an electric field

Determining Polarity

To determine if a molecule is polar: • Write the Lewis structure • Draw the geometry • Use the following symbol to denote the polarity of each bond Positive end of the bond, the less electronegative atom Negative end of the bond, more electronegative atom attracts the electrons more strongly towards it

Practice Determining Polarity

Determine whether the following bonds and molecules are polar: 1.

2.

Si – Cl H – C 5.

6.

O 2 HF 3.

4.

C – C S – Cl 7.

8.

CH 4 H 2 O

3.5 Properties Based on Electronic Structure and Molecular Geometry

• Intramolecular forces – attractive forces

within

molecules – Chemical bonds • Intermolecular forces – attractive forces

between

molecules • Intermolecular forces determine many physical properties o Intermolecular forces are a direct consequence of the intramolecular forces in the molecules

Solubility and Intermolecular Forces

Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature • “Like dissolves like” o Polar molecules are most soluble in polar solvents o Nonpolar molecules are most soluble in nonpolar solvents • Does ammonia, NH 3 , dissolve in water?

• Yes, both molecules are polar

Interaction of Water and Ammonia

• The  end of ammonia, N , is attracted to the  + the water molecule, H • The  + end of ammonia, H , is attracted to the  the water molecule, O end of end of • The attractive forces, called hydrogen bonds, pull ammonia into water, distributing the ammonia molecules throughout the water, forming a homogeneous solution

Interaction of Water and Oil

• What do you know about oil and water?

o “They don’t mix” • Why?

o Because water is polar and oil is nonpolar • Water molecules exert their attractive forces on other water molecules • Oil remains insoluble and floats on the surface of the water as it is less dense

Boiling Points of Liquids and Melting Points of Solids

• Energy is used to overcome the intermolecular attractive forces in a substance, driving the molecules into a less associated phase • The greater the intermolecular force, the more energy is required leading to o Higher melting point of a solid o Higher boiling point of a liquid

Factors Influencing Boiling and Melting Points

• Strength of the attractive force holding the substance in its current physical state • Molecular mass • Larger molecules have higher m.p. and b.p. than smaller molecules as it is more difficult to convert a larger mass to another phase • Polarity • Polar molecules have higher m.p. and b.p. than nonpolar molecules of similar molecular mass due to their stronger attractive force

Melting and Boiling Points –

Selected Compounds by Bonding Type