Unit 5. INTRODUCTION TO BONDING Atoms are generally found in nature in combination held together by chemical bonds. A chemical bond is an electrical

Unit 5

CHEMICAL BONDING

INTRODUCTION TO BONDING

• Atoms are generally found in nature in combination held together by chemical bonds.

• A chemical bond is an electrical attraction between the nuclei and valence electrons of different atom that binds the atoms together so that they behave as one unit.

• There are two types of chemical bonds: ionic, and covalent.

HOW DO ATOMS INTERACT?

• There are 2 ways to look at atom-atom interactions

• 1) Balancing the opposing forces of repulsion and attraction

As the atoms come closer together there is a repulsion between the negative e- clouds of each atomSimultaneously there is an attraction between the positive nuclei and the negative electron clouds

ATOM-ATOM INTERACTIONS

As the optimum distance is achieved that balances these forces, there is a release of potential energy

The atoms vibrate within the window of maximum attraction/minimum repulsion

The more energy released the stronger the connecting bond between the atoms



• 2) Atoms want to achieve the most stable arrangement of valence electrons.

By rearranging the electrons so that each atom achieves a noble gas-like arrangement of its electrons creates a pair of stable atoms (only occurs when bonded)


• To achieve a stable arrangement one or more valence electrons are transferred between two atoms

Basis for ionic bonding• Sometimes valence electrons are

shared between two atoms

Basis for covalent bonding


• Gilbert Lewis formulated that atoms react in order to achieve a more stable electron configuration.

• Maximum stability occurs when an atom is isoelectronic with a noble gas. (has a full octet)

ns2np6, 8ve-

• When elements react it is only their valence shells that interact and therefore it is only the valence electrons that are of interest to us

• To keep track of how many valence electrons each element has and that the total number of valence electrons remains constant chemists use Lewis dot symbols

WHY DO ELEMENTS REACT?

• Valence electrons are electrons that exist the highest energy level.

These are electrons with the highest principle quantum number, n.

• Valence electrons are the electrons that participate in bonding

• Valence electrons can be determined by the electron configuration

Remember that all elements within the same column have the same number of valence electrons

i.e.. All alkali metals have 1 ve-

VALENCE ELECTRONS RECAP

• Lewis dot symbol consists of a symbol of an element an one dot for each valence electron in the element.

LEWIS DOT SYMBOLS

• A good predictor for which type of bonding will develop between a set of atoms is the difference in their electronegativity's.

• The greater the difference between the electronegativity's, the less equal the exchange of electrons between two atoms

• This leaves us with three different levels of interaction: pure covalent (nonpolar covalent), polar covalent, and ionic

PREDICTING THE BOND TYPE

• When the electronegativity difference between two atoms is greater than 2.0 the bond is ionic.

• When the difference is less than 0.3 the bond is considered pure (nonpolar) covalent.

• When the difference is between 0.3 and 1.6 the bond is considered polar covalent.

DETERMINING THE TYPE

So what about the range between 1.6 and 2.0?

If one of the atoms is a metal we will define it as ionic.

If both atoms are nonmetals we will define it as polar covalent.These ranges aren’t conclusive, they are our attempts to explain observed patterns in nature


• The take home lesson on electronegativity and bonding is this:

The closer together the atoms are on the P.T., the more evenly their e- interact, and so are more likely to form a pure covalent bond

The farther apart they are on the P.T., the less evenly their e- interact, and are therefore more likely to form an ionic bond.

In between the extremes exists varying degrees of polar covalent interactions


metal w/nonmetal = usually ionic

nonmetal

w/nonmetal = usually covalent

RULE OF THUMB

Let’s consider the compound Cesium Fluoride, CsF.

The electronegativity value (EV) for Cs is .70; the EV for F is 4.00.

This is a difference of 3.30

What type of bond is it?

WHAT BOND TYPE IS FORMED?

IONIC BOND

• Elements with low ionization energies (metals) tend to form cations by losing valence electrons

• Groups 1 and 2 are the elements that are most likely to form cations

• Elements with high electron affinities (non-metals) tend to form anions by gaining valence electrons

• Groups 16 and 17 are the elements that are most likely to form anions

GAINING/LOSING ELECTRONS

Reaction between Li and F

This occurs in 3 steps1. The ionization of Li2. The acceptance of the electron by F3. The electrostatic attraction between the ions

EXAMPLE

Li F LiF

Lewis dot symbol

Electron Configuration

An ionic bond holds together the ions in an ionic compound

An ionic compound is the product of a cation and an anion

An ionic bond is most often between a metal and a nonmetal

IONIC BONDS

• Ionic compounds tend to be crystalline solids.

The crystals are all different colors and textures, but they all tend to be hard and brittle

They can be broken to leave a clean smooth surface• Ionic compounds tend to have high m.p.

But when melted, they conduct electricity

PROPERTIES OF IONIC BONDING

• Many ionic solids are soluble in water where they are dissolved into separate ions.

• Dissolved ions in water make water conductive

These are called electrolytes.

Most covalent compounds are not electrolytes (acids are an exception)

PROPERTIES OF IONIC BONDING

Naming ionic compounds requires practice and a solid understanding of the rules

In order to help make this easier we will split ionic compounds into 3 types. Please note that these three types

are not real and are used only to help you learn how to name compounds.

Type 1 Binary compounds using group 1 and 2 metals

Type 2 Binary compounds using transition metals

Type 3 Ternary compounds using polyatomic ions

NAMING IONIC COMPOUNDS

TYPE 1

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Type 1 naming consists of binary compounds. Binary compounds contain two elements, a metal and a nonmetalThe metals in type one ONLY comes from the representative group metals, primarily those in groups 1 and 2Metals lose their valence electrons to become cationsCations are named by the element nameNonmetals gain valence electrons to become anionsAnions are the element name with the ending changed to -ide

TYPE 1

COMMON TYPE 1 IONS

Ion Formula Ba2+ Al3+ K+

Name

Ion Formula

Name Calcium Rubidium Gallium

Ion Formula N3 O2 F

Name

Ion Formula

Name phosphide sulfide chloride

LET’S TRY IT

1. Write the metal as an cation

2. Write the nonmetal as an anion

3. Balance the charge

(use the crisscross method)

Name to Formula Example: Potassium

Chloride

STEPS FOR WRITING FORMULAS TYPE 1

Name Cation formula

Anion formula

Formula

sodium chloride

potassium sulfide

magnesium oxide

calcium iodide

aluminum oxide

LETS TRY IT

1. Write the name of the metal

2. Write the name of the nonmetal

3. Change the ending of the nonmetal to –ide

4. Put the two names together

Formula to Name Example: BaS

STEPS FOR NAMING TYPE 1

Formula Cation name Anion name NameCaO

KBr

Al2O3

MgCl2

Sr3N2

LETS TRY IT

TYPE 2

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Type 2 naming consists of binary compounds.

The metals in type 2 ONLY come from the transition metals (groups 3-12)

These metals are trickier that those in groups previously discussed because they have multiple oxidation states.

This means that the number of electrons they can lose changes depending on :

what they are to and the conditions under which the bonding occurs

TYPE 2

An example of a transition element that has multiple oxidation states is Iron

Iron can lose between 1 and 6 electrons to become:

Fe+, Fe2+, Fe3+, Fe4+, Fe5+, Fe6+

How do we know which oxidation state Iron has when writing formulas or the names of compounds?????

We use the Stock system which uses Roman numerals to denote the oxidation state of the metal.

AS FAR AS WE ARE CONCERNED ONLY TRANSITION METALS, TIN, AND LEAD WILL HAVE MULTIPLE OXIDATION STATES EXEPT Zn2+,

Ag+, and Cd2+ WHICH ONLY FORM ONE ION.

MULTIPLE OXIDATION STATES

Roman Numeral

Arabic Numeral

Roman Numeral

Arabic Numeral

I VIII

II IX

III X

IV L

V C

VI D

VII M

ROMAN NUMERALS

COMMON TYPE 2 IONS

Ion Formula Fe2+ Au3+ Cu+

Name

Ion Formula

Name Tin (IV) Tin (II) Zinc

Ion Formula Ag+ V5+ W6+

Name

Ion Formula

Name Platinum (IV) Osmium (VII) Scandium (III)

LETS TRY IT

1. Write the metal as an cation using the Roman numeral as the charge

2. Write the nonmetal as an anion


(use the crisscross method, reduce if necessary)

Name to Formula Example: Manganese

(IV) oxide


Name Cation formula

Anion formula

Formula

Lead (II) oxide

Copper (I) Nitride

Zinc chloride

Gold (III) sulfide

Silver fluoride

LET’S TRY IT


2. Write the name of the nonmetal

3. Change the ending of the nonmetal to –ide

4. Determine the roman numeral using the anion

as a guide


Formula to Name Example: TiCl 4

STEPS FOR WRITING NAMES TYPE 2

Formula Cation name Anion name Name

SnCl4

TiO2

TiO4

Fe3P2

Cu3N

LETS TRY IT

TYPE 3

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Type 3 naming consists of ternary compounds.

Ternary compounds are made of 3 or more different elements.2 or more of these elements combine to form a polyatomic ion.

The cations are almost always metals with the exception of ammonium (NH4

+)

The anions are usually the polyatomic ions for example (CO3

2-; NO3- ; NO2

- ; OH- ; PO43-)

TYPE 3

Polyatomic Ion Charge = -1



bicarbonate

HCO3- thiosulfate S2O3

2- phosphate

PO43-

permanganate

MnO4- sulfite SO3

2- borate BO33-

perchlorate ClO4- sulfate SO4

2-

nitrite NO2- peroxide O2

2-

nitrate NO3- hydrogen

phosphateHPO4

2-

hydroxide OH- dichromate Cr2O72-

chlorate - ClO3- chromate CrO4

2-

cyanide - CN- carbonate CO32-

acetate - C2H3O2-

COMMON POLYATOMIC IONS

1. Write the metal as an cation (using the Roman numeral as the charge if

a type 2 metal)

2. Write the nonmetal as an anion. (Look up the

polyatomic ion on your polyatomic ion sheet)


(use the crisscross method, reduce if necessary)

Name to Formula Example: Aluminum

Acetate


Name Cation formula

Anion formula

Formula

Lead (II) phosphate

Calcium Nitrate

Zinc carbonate

Gold (III) sulfite

Ammonium hydroxide

LETS TRY IT


2. Write the name of the nonmetal. (find it on your polyatomic ion

sheet)

3. Determine the roman numeral using the anion as a guide (if the metals is a transition metal, tin

or lead)


Formula to Name Example: FePO 4


Formula Cation name Anion name Name

SnPO4

TiCO3

Fe(NO3)2

Mn(CO3)2

(NH4)2Cr2O7

LETS TRY IT

STABILITY

We can predict which elements will form ionic bonds based on electron affinities and ionization energies but how do we

evaluate how stable a bond is?

In the solid state each cation is surrounded by a specific number of anions

and vice versa.

LATTICE ENERGY

Overall stability is based on all these interactions and not just the interaction between one cation and one anion

A quantitative measure of stability of any solid ionic compound is its lattice energy.

Lattice energy is defined as the amount of energy needed to separate a mole of a solid ionic compound into gaseous ions

The greater the lattice energy the more stable the ionic compound

LATTICE ENERGY

LATTICE ENERGY TABLE

LETS TRY IT

Which of the following pairs of ionic compounds is most stable and why?

LiF and LiCl

Sodium bromide and potassium bromide

Magnesium chloride and LiF

CsCl and ScN

COVALENT BONDS

A covalent bond is one in which two (or more) electrons are shared by 2 atoms

Covalent bonding involves the sharing of valence electrons. The electrons that are involved in bonding are known as

bonding electrons

Not all valence electrons will participate in bonding. Those who do not participate in binding are known as nonbonding

or lone pair electrons

COVALENT BONDS

Covalent compounds are names using prefixes

A prefix is an affix which is placed before the root of a word

NAMING COVALENT MOLECULES

Prefix Number Prefix Number

1 6

2 7

3 8

4 9

5 10

1. Write the symbol of the first element

2. Using the prefix determine the number of the atoms

3. Write the symbol of the nonmetal.

4. Using the prefix determine the number of the atoms

5. Put the 2 together

Name to Formula Example: diphosphorus

pentoxide

STEPS FOR WRITING FORMULAS TYPE

LETS TRY IT

Name

Carbon monoxideCarbon tetrachlorideTetrabromine trifluorideSilicon dioxideDiphosphorus decoxide

1. Write the name of the first element

2. Add the prefix for the number of atoms

3. Write the name of the second element

4. Add the prefix for the number of atoms

5. Put it together

Formula to Name Example: P 7O5


LETS TRY IT

Formula

SO4

C3O8

CH4

NCl4

OF2

Lewis structures are representations of covalent

bonding using Lewis dot symbols in which shared electrons are shown as a

line between two atoms and lone pairs are shown as

pairs of dots on an individual atom

LEWIS STRUCTURES

1. Determine the total number of valence electrons

2. Determine the central atom.

(carbon will always occupy the central atom; if carbon is not present the least electronegative atom is the central atom)

3. Use 2 electrons to bond the central atom to each of the terminal atoms

If any electrons remain donate them to the terminal atoms until all have achieved a full octet

If electrons still remain donate them to the central atom

4. In the event that you do not have enough electrons double or triple bond

HOW TO DRAW LEWIS STRUCTURES

Draw the Lewis structure for NF3

LET TRY IT: SIMPLE STRUCTURE

Draw the Lewis structure for H2O

SIMPLE STRUCTURES

Draw the Lewis structure for O2

MULTIPLE BONDS

Draw the Lewis structure for N2

MULTIPLE BONDS

Draw the Lewis structure for HCO3

MULTIPLE CENTRAL ATOMS

Draw the Lewis structure for CO32-

POLYATOMIC ION

There are exceptions to the octet rule:

1. The incomplete octet

In some compounds the number of electrons needed to make it stable are less than 8. For example: Be, needs 4 ve-; Al and B need 6ve-

2. The expanded octet

1. In a number of compounds there are more than 8ve-. These are only available to elements in and beyond the third period, and is due to the unoccupied and available d orbitals. For example: SF6

EXCEPTIONS

MOLECULAR GEOMETRIES

Lewis structures only give us a 2 dimensional picture for a 3 dimensional molecule.

The study of geometry in molecules comes from the assumption that the electrons in the valence shell of an atom repel one another.

The geometry comes from the Valence Shell Electron Pair Repulsion (VSEPR) model

Make the following assumptions:

1. Double and triple bonds act as a single bond as far as electron repulsion

2. The VSEPR model may be applied to any resonance structure

3. Formal charges are not shown

3 DIMENSIONS

We can group geometries according to the number of electron pairs (both bonding and lone pairs) on the central

atom.

Electron Pair Geometries

1 or 2 electron pairs attached-linear

3 electron pairs attached-trigonal planar

4 electron pairs attached-tetrahedral

5 electron pairs attached-trigonal bipyramidal

6 electron pairs-octahedral

ELECTRON PAIR GEOMETRIES

LINEAR GEOMETRY

Number of Bonds

Number of lone pairs

Electron pair

Geometry

Molecular geometry

Bond angle

2 0 Linear Linear 180o

TRIGONAL PLANAR GEOMETRY

Number of Bonds


Electron pair

Geometry

Molecular geometry

Bond angle

3 0 Trigonal planar

Trigonal planar

120o

2 1 Trigonal planar

Bent <120o

TETRAHEDRAL GEOMETRY

Number of Bonds


Electron pair

Geometry

Molecular geometry

Bond angle

4 0 Tetrahedral Tetrahedral 109.5o

3 1 Tetrahedral Trigonal pyramidal

<109.5o

2 2 Tetrahedral Bent <<109.5o

TRIGONAL BIPYRAMIDAL GEOMETRY

Number of Bonds


Electron pair

Geometry

Molecular geometry

Bond angle

5 0 Trigonal bipyramidal

Trigonal bipyramidal

90o,180o


See-saw <90o,<180o


T-shaped <90o


Linear 180o

TRIGONAL BIPYRAMIDAL

OCTAHEDRAL GEOMETRY

Number of Bonds


Electron pair

Geometry

Molecular geometry

Bond angle

6 0 Octahedral Octahedral 90o

5 1 Octahedral Square pyramidal

<90o,<90o

4 2 Octahedral Square planar

90o

3 3 Octahedral T-shaped <90o

2 4 Octahedral Linear 180o

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Unit 5. INTRODUCTION TO BONDING Atoms are generally found in nature in combination held together by chemical bonds. A chemical bond is an electrical