Chapter 3

Structure and Properties of Ionic and Covalent Compounds

Denniston Topping Caret

7th Edition

Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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

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• 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

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• Each “side” of the symbol represents an atomic orbital, which may hold up to two electrons

• Using Lewis symbols– Place one dot on each side until there are four dots

around the symbol

– Now add a second dot to each side in turn

– The number of valence electrons limits the number of dots placed

– Each unpaired dot (unpaired electron of the valence shell) is available to form a chemical bond

Lewis Dot Symbols for Representative Elements

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3.1

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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– Each atom achieves a “Noble Gas”

configuration– 2 ions are formed; a cation and anion, which

are attracted to each other3.1

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3.1

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Na + Cl NaCl

Sodium has a low ionization energy itreadily loses this electron

Na Na+ + e-

When sodium loses the electron, it gains the Ne configuration

Chlorine has a high electron affinity

When chlorine gains an electron, it gains the Ar configuration

:

..

..Cl: e

..

..Cl:

Ionic Bonding

3.1

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• 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

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Let’s look at the formation of H2:

H + H H2

• Each hydrogen has one electron in its valance shell• If it were an ionic bond it would look like this:

• 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

:H H H H

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The shared electron

pair is called a Covalent Bond

Each hydrogen atom now has two electrons around it and attained a He configuration

• Instead, each atom attains a noble gas configuration by sharing electrons

H:H H H

Covalent Bonding in Hydrogen3.

1 C

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:..

..F:

..

..F: :

..

..F

..

..F:

Each fluorine is surrounded by 8 electrons – Ne configuration

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)– H2, N2, O2, F2, Cl2, Br2, I2

H:..

..O:H

..

..O 2H

H H:

..C:H

.C 4H

H

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2e– from 2H 2e– for H6e– from O 8e– for O

4e– from 4H 2e– for H4e– from C 8e– for C

3.1

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Electronegativity• The Polar Covalent Bond

– Ionic bonding involves the transfer of electrons

– Covalent bonding involves the sharing of electrons

– Polar covalent bonding - bonds made up of unequally shared electron pairs

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These two electrons are not shared equally

somewhat negatively chargedsomewhat positively charged

:..F:H :

..F H

• 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

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

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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– The difference in electronegativity determines

the extent of bond polarity3.1

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electronegativity increases elec

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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

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• The greater the difference in electronegativity between two atoms, the greater the polarity of their bond

• Which would be more polar, a H-F bond or H-Cl bond?

• H-F … 4.0 - 2.1 = 1.9

• H-Cl … 3.0 - 2.1 = 0.9

• The HF bond is more polar than the HCl bond

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:– ionic compounds– covalent compounds

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• A formula is the representation of the fundamental compound using chemical symbols and numerical subscripts– The formula identifies the number and type

of the various atoms that make up the compound unit

– The number of like atoms in the unit is shown by the use of a subscript

– Presence of only one atom is understood when no subscript is present

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• 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

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from the Identities of the Component Ions

• Determine the charge of each ion

– Metals have a charge equal to group number

– 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

• It must have the same number of positive charges as negative charges

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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

2

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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

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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

• Roman numerals give the charge of the metal

• Examples:

• FeCl3 is iron(III) chloride

• FeCl2 is iron(II) chloride

• CuO is copper(II) oxide

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• 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:

• FeCl2 is ferrous chloride

• FeCl3 is ferric chloride

Stock and Common Names for Iron and Copper Ions

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and Anions

• Monatomic ions - ions consisting of a single charged atom

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• Polyatomic ions - ions composed of 2 or more atoms bonded together with an overall positive or negative charge

– Within the ion itself, the atoms are bonded using covalent bonds

– The positive and negative ions will be bonded to each other with ionic bonds

• Examples:

• NH4+ ammonium ion

• SO42- sulfate ion

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Anions

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1. NH4Cl

2. BaSO4

3. Fe(NO3)3

4. CuHCO3

5. Ca(OH)2

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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 BaCl2

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NamesWrite the formula for the following ionic

compounds:

1. sodium sulfate

2. ammonium sulfide

3. magnesium phosphate

4. chromium(II) sulfate

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• 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

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

Example: CO is carbon monoxide

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

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

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1. SiO2

2. N2O5

3. CCl4

4. IF7

2

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Compounds• Use the prefixes in the names to determine the

subscripts for the elements• Examples:

• nitrogen trichloride NCl3

• diphosphorus pentoxide P2O5

• Some common names that are used:

– H2O water

– NH3 ammonia

– C2H5OH ethanol

– C6H12O6 glucose

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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– Ionic compounds are usually solids at room

temperature– Covalent compounds can be solids, liquids, and

gases

• Melting and Boiling Points– Melting point - the temperature at which a

solid is converted to a liquid– Boiling point - the temperature at which a

liquid is converted to a gas

Physical Properties

• Melting and Boiling Points– Ionic compounds have much higher melting points

and boiling points than covalent compounds

– A large amount of energy is required to break the electrostatic attractions between ions

– Ionic compounds typically melt at several hundred degrees Celsius

• Structure of Compounds in the Solid State– Ionic compounds are crystalline

– Covalent compounds are crystalline or amorphous – having no regular structure

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• Solutions of Ionic and Covalent Compounds– Ionic compounds often dissolve in water,

where they dissociate - form positive and negative ions in solution

– Electrolytes - ions present in solution allowing the solution to conduct electricity

– Covalent solids usually do not dissociate and do not conduct electricity - nonelectrolytes

Physical Properties3.

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Comparison of Ionic vs. Covalent Compounds

Ionic Covalent

Composed of Metal + nonmetal 2 nonmetals

Electrons Transferred Shared

Physical state Solid / crystal Any / crystal OR amorphous

Dissociation Yes, electrolytes No, nonelectrolytes

Boiling/Melting High Low

3.4 Drawing Lewis Structures on Molecules and Polyatomic Ions

Lewis Structure Guidelines

1. Use chemical symbols for the various elements to write the skeletal structure of the compound

– The least electronegative atom will be placed in the central position

– Hydrogen and halogens occupy terminal positions

– Carbon often forms chains of carbon-carbon covalent bonds

Lewis Structure Guidelines

2. 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

– Polyatomic cations, subtract one electron for every positive charge

– Polyatomic anions, add one electron for every negative charge

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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

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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

5. Recheck that all atoms have the octet rule satisfied and that the total number of valance electrons are used

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Drawing Lewis Structures of Covalent Compounds

Draw the Lewis structure of carbon dioxide, CO2

Draw a skeletal structure of the molecule1. Arrange the atoms in their most probable order

C-O-O and/or O-C-O2. Find the electronegativity of O=3.5 & C=2.5

3. Place the least electronegative atom as the central atom, here carbon is the central atom

4. Result is the O-C-O structure from above3.4

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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. Use electron pairs to connect the C to each O

with a single bond

O : C : O 7. Place electron pairs around the atoms : O : C : O :

This satisfies the rule for the O atoms, but not for C

Drawing Lewis Structures3.

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: :::

8. Redistribute the electrons moving 2 e- from each O, placing them between C:O

C::O::C

9. In this structure, the octet rule is satisfied• This is the most probable structure• Four electrons are between C and O

• These electrons are share in covalent bonds• Four electrons in this arrangement signify a double

bond

10. Recheck the electron distribution• 8 electron pairs = 16 valence electrons, number

counted at start• 8 electrons around each atom, octet rule satisfied

Drawing Lewis Structures of Covalent Compounds

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:: :

:

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Using the guidelines presented, write Lewis structures for the following:

1. H2O

2. NH3

3. CO2

4. NH4+

5. CO32-

6. N2

6

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Polyatomic Ions

• Prepare Lewis structures of polyatomic ions as for neutral compounds, except:

• The charge on the ion must be accounted for when computing the total number of valence electrons

Lewis Structure of Polyatomic Cations

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+

Draw a skeletal structure of the molecule1. Ammonium has this structure and charge:

2. The total number of valence electrons is determined by subtracting one electron for each unit of positive charge

1 N atom x 5 valence electrons = 5 e-

4 H atoms x 1 valence electron = 4 e-

- 1 electron for +1 charge = -1 e-

8 e- total 3. Distribute these 8 e- around the skeletal structure

Draw the Lewis structure of carbonate ion, CO32-

Draw a skeletal structure of the molecule1. Carbon is less electronegative than oxygen

• This makes carbon the central atom

• Skeletal structure and charge:

2. The total number of valence electrons is determined by adding one electron for each unit of negative charge

1 C atom x 4 valence electrons = 4 e-

3 O atoms x 6 valence electron = 18 e-

+ 2 negative charges = 2 e-

24 e- total 3. Distribute these e- around the skeletal structure

Lewis Structure of Polyatomic Anions

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Draw the Lewis structure of carbonate ion, CO32-

4. Distributing the electrons around the central carbon atom (4 bonds) and around the surrounding O atoms attempting to satisfy the octet rule results in:

5. This satisfies the octet rule for the 3 oxygen, but not for the carbon

6. Move a lone pair from one of the O atoms to form another bond with C

Lewis Structure of Polyatomic Anions

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• 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

Lewis Structure, Stability, Multiple Bonds, and Bond Energies

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NNor ..N

..N

OOor ..O::

..O H - Hor H:H

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

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• Write the Lewis structure of CO32-

• If you look around you, you will probably see the double bond put in different places

• Who is right? All of you!

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

• 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

:..O:C::

..O: :

..O:C:

..O: O::C:

..O:

: :: : :

..O: :O: :

..O:

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to the Octet Rule1. Incomplete octet - less then eight

electrons around an atom other than H

– Let’s look at BeH2

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

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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 - ON - O

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 PF5

• 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

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Geometry: VSEPR Theory• Molecular shape plays a large part in

determining properties and shape

• VSEPR theory - Valance Shell Electron Pair Repulsion theory

• 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

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• 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

Molecular Bonding

• Bonding pair = two electrons shared by 2 atoms– H:O

• Nonbonding pair = two electrons belonging to 1 atom, pair not shared– N:

• Maximal separation of bonding pairs = 4 corners of a TETRAHEDRON3.

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A Stable Exception to the Octet Rule

• Consider BeH2

– Only 4 electrons surround the beryllium atom

– These 2 electron pairs have minimal repulsion when located on opposite sides of the structure

– Linear structure having bond angles of 180°

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Another Stable Exception to the Octet Rule

• Consider BF3

– There are 3 shared electron pairs around the central atom

– These electron pairs have minimal repulsion when placed in a plane, forming a triangle

– Trigonal planar structure with bond angles of 120°

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Basic Electron Pair Repulsion of a Full Octet

• Consider CH4

– There are 4 shared electron pairs around the central Carbon

– Minimal electron repulsion when electrons are placed at the four corners of a tetrahedron

– Each H-C-H bond angle is 109.5°

• Tetrahedron is the primary structure of a full octet

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

Consider NH3

• 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 structure with 107° angles

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

Consider H2O

• There are 4 electron pairs around the central Oxygen• 2 pairs are shared electron pairs• 2 pairs are lone pairs

– 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 or angular structure with 104.5° bond angles

Predicting Geometric Shape Using Electron Pairs

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Molecular Shape1. 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

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Geometry• PCl3

• SO2

• PH3

• SiH4

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• 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

Positive end of the bond, the less electronegative atom

Negative end of the bond, more electronegative atom attracts the electrons more strongly towards it3.

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Determining PolarityTo determine if a molecule is polar:

• Write the Lewis structure

• Draw the geometry

• Use the following symbol to denote the polarity of each bond

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esPractice Determining PolarityDetermine whether the following bonds and molecules are polar:

1. Si – Cl 1. O2

2. H – C 2. HF

3. C – C 3. CH4

4. S – Cl 4. H2O

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– Intermolecular forces are a direct consequence

of the intramolecular forces in the molecules

Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature

• “Like dissolves like” – Polar molecules are most soluble in polar

solvents– Nonpolar molecules are most soluble in

nonpolar solvents

• Does ammonia, NH3, dissolve in water?• Yes, both molecules are polar

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Forces

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Ammonia

• The - end of ammonia, N, is attracted to the + end of the water molecule, H

• The + end of ammonia, H, is attracted to the - end of the water molecule, O

• The attractive forces, called hydrogen bonds, pull ammonia into water, distributing the ammonia molecules throughout the water, forming a homogeneous solution

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• What do you know about oil and water?

– “They don’t mix”

• Why?

– 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

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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

– Higher melting point of a solid

– Higher boiling point of a liquid

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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

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Selected Compounds by Bonding Type