L3 Chemical Bond September2014

CHEMICAL BOND &

MOLECULAR GEOMETRY

Naming of compounds

Types of bonding

• Ions: cations, anions, oxoanions• Ionic compounds• Covalent compounds: 4 rules

• Ionic bond: transfer of electron(s)• Covalent bond: sharing of electron(s)

Types of compounds• Ionic compounds: metals + nonmetals• Covalent compounds: between metals

SUMMARY OF LAST WEEK

Quantum numbers: n, l, ml and ms

n: determines the energy of the “orbital”

l: determines the shape of the orbital ( l= 0, 1,…..n-1); l= 0 is “s”, l=1 is “p”, l=2 is “d” etc

m (or ml): determines the orientation of the “orbital”; m= -l,…,0,…,+l;therefore: 1s orbital, 3p orbital, 5d orbital

ms: describes the spin of the electron, m= +1/2, -1/2



Multi-electron configuration:

Pauli Exclusion Principle: no 2 electrons in an atom may be in the same quantum state (n,l,ml,ms)

Aufbau Principle: electrons adopt the lowest possible energy

Shielding: orbitals of equal n nearest the nucleus have lowest energy

Hund’s Rule: in orbitals, electrons prefer to be unpaired first


PART 1: CHEMICAL BOND

Highlight:

• Covalent bond

• Ionic bond

• Metallic bond

THIS WEEK

CONCEPTS Valence electron Lewis dot formula Electronegativity Polar, non polar bond Polarity Dipole, dipole moment

Highlight:

• Lewis structure

• Octet rule

• Molecular

geometry

THIS WEEK

PART 2: MOLECULAR SHAPE

CONCEPTS How to draw a LEWIS structure Octet rule VSEPR (Valence-Shell Electron Pair Repulsion) theory Determine the molecular shape by VSEPR


THIS WEEK

CONCEPTSHighlight:

• Covalent bond

• Ionic bond

• Metallic bond

Valence electron Lewis dot formula Electronegativity Polar, non polar bond Polarity Dipole, dipole moment

CuSO4.5H2

O

NaCl

NiCl2.6H2

OK2Cr2O

7

CoCl2.6H2

O

IONIC COMPOUNDS

H2 O(water)

CO2

(carbon dioxide)

CH4

(methane)

C2H5OH(ethyl

alcohol)

COVALENT COMPOUNDS

Light bulb experiment

http://www.youtube.com/watch?v=aneshRgoQMk

Why are the properties of a substance different from another ?

• Properties of an atom depend on the electron configuration and the strength of the nucleus-electron attractions

Similarly,

• Properties of a substance depend on the type and strength of chemical bonds

CHEMICAL BONDS the forces that hold the atoms of elements

together in compounds

CHEMICAL BONDS the forces that hold the atoms of elements

together in compounds

Highlight:

• Ionic bond: metals and nonmetals

• Covalent bond: nonmetals and

nonmetals

• Metallic bond: metal with metal

3 types of bondings

To understand bonding, we should know:

• Valence electrons• How to draw Lewis formula• Octet rule

Does electron configuration has any effect on chemical bond?

YES. But only the electrons at outer most shell will determine bonding

Numbers and arrangements of valence electrons determine:• chemical and physical properties of elements• kinds of chemical bonds

They are called valence electrons


Electron configuration:

Valence electrons: the outer most electrons of atoms11Na: 1s2 2s2 2p6 3s1

8O: 1s2 2s2

2p4

valence electrons: 1valence electrons: 6

Electron configuration:

Valence electrons: the outer most electrons of atoms11Na: 1s2 2s2 2p6 3s1

8O: 1s2 2s2

2p4

valence electrons: 1valence electrons: 6


Lewis formulas (Lewis dot formulas): to describe the valence electrons

11Na: 1s2 2s2 2p6 3s1

8O: 1s2 2s2

2p4

valence electron: 1

valence electrons: 6

Na

O•

•

••••

•

Only valence electrons of s and p orbitals: shown in dots

Number of valence electrons of GROUP A = group number

-Use Lewis symbol to express valence electrons

-It’s easy to write the Lewis symbol for any main-group element:

1. Note its A-group number (1A to 8A), which equals the number of valence electrons.

2. Place one dot at a time on the four sides (top, right, bottom, left) of the element symbol.

3. Keep adding dots, pairing the dots until all are used up.

Now, can you quickly write the Lewis symbol of all elements in group 6A?

The Lewis symbol provides information about an element’s bonding behavior:

• For a metal, the total number of dots = the maximum number of electrons an atom loses to form a cation.

Na• Al •••

Ca••

The Lewis symbol provides information about an element’s bonding behavior:

• For a nonmetal, the number of unpaired dots equals

either the number of electrons an atom gains in becoming an anion

or the number it shares in forming covalent bonds.

O••

••••

3 types of bonding: a closer look

IONIC BONDING

•Ionic bonding: forms between atoms with LARGE differences in their tendencies to lose or gain electrons.

•These differences results in transfer of electron from one atom to another in large numbers to form a compound.

IONIC BONDING

Normally, ionic bonding forms between reactive metals and nonmetals

Reactive metals:

Group 1A&

Group 2A

NometalsGroup 7A

&the top of group 6A

+low ionisation energy easily loses e-

very negative affinity easily attract e-

•The electrostatic attraction of the ions draw them into a 3D array of an ionic solid

The strength of ionic bonds depends on: Charges & Sizes

Proportional to

Coulomb’s Law:

q: Charge(s) of the ion(s)d: Distance between ions

STRENGTH OF IONIC BOND

Ions with higher charges and smaller sizes will attract each other stronger

Higher charges q +, q - increase F increases

Smaller sizes D decreases F increases

COVALENT BONDING

• Covanlent bonding: forms between atoms with small differences in their tendencies to lose or gain electrons.

COVALENT BONDING

• Normally, covalent bonding forms between nonmetals.

• Each nonmetal has HIGH ionisation energy: hold onto its own electrons tightly

HIGHLY negative electron affinity: attract electrons from others

• The attraction of each nucleus & valence electron: draw atoms close to each other

METALLIC BONDING

•Metals: can easily loose electron at outer shell• Electron pool• Electrostatic attraction

3 types of bonding

Ionic bonding

Covalent bonding

Metallic bonding

MODELS OF BONDING

Model of ionic bonding

Model of covalent bonding

3 types of bonding

Ionic bonding

Covalent bonding

Metallic bonding

MODELS OF BONDING

Model of ionic bonding

Model of covalent bonding

COVALENT BONDING

Why does covalent bonding form?

Key concepts

• Covalent bonding

• Shared and unshared pair

• Bond energy and bond length

• Bond polarity

• Dipole moment

COVALENT BONDING of H2

COVALENT BONDING

the atoms are too far apart to attract each other.

each nucleus attractsthe other atom’s electron.

The combination of nucleus-electron attractions and electron-electron and nucleus repulsions gives the minimum energy of the system.

repulsions increase the system’s energy and force the atoms apart to point 3 again.

COVALENT BONDING

Distribution of electron density

Covalent bonding : results from sharing one or more pair electrons between atoms

Covalent bonding occurs when the electronegativity difference, (∆EN), between elements (atoms) is zero or relatively small.

When can a covalent bond be formed?

Valence electron(s): participate in covalent bonding

COVALENT BONDING

• Shared or bond pair (electron pairs)• Unshared or lone pair

SHARED SHARED

Unshared Unshared

BOND PAIR & LONE PAIR

• Shared or bond pair (electron pairs)• Unshared or lone pair

SHARED SHAREDUnshared Unshared


• Bonding pair (electron pair): Bonding pair A pair of electrons involved in a covalent bond.Also called shared pair.

• Lone pair: A pair of electrons residing on one atom and not shared by other atoms; unshared pair.


H• H•+ H • •H

H-Hor

H• +

H-For

F•••

•••• H •F•

••

••••

FORMATION OF COVALENT BOND

Covalent bonding : results from sharing one or more pair electrons between atoms

Single covalent bond: two atoms share one electron pair. SINGLE BOND

Double covalent bond: two atoms share two electron pairs. DOUBLE BOND

Triple covalent bond: two atoms share three electron pairs. TRIPLE BOND

BOND TYPES

BOND ENERGY AND BOND LENGTH

•The strength of a covalent bond depends on the magnitude of the mutual attraction between bonded nuclei and shared electrons.

•The bond energy (BE) (also called bond enthalpy or bond strength) is the energy required to overcome this attraction.



• Bond energies depend on characteristics of the bonded atoms

their electron configurations, nuclear charges atomic radii

• A covalent bond has a bond length, the distance between the nuclei of two bonded atoms.

Bond Energy and Bond LengthEnergy: (kJ) average bond energyLength: (pm) average bond length

BOND LENGTH AND ATOMIC RADIUS

Energy: (kJ) average bond energyLength: (pm) average bond lengthWithin a series of similar molecules, such as the diatomic halogen molecules, bond length increases as covalent radius increases

BOND STRENGTH, LENGTH AND BOND TYPE

Between single bond, double bond and triple bond

• Bond length: single bond > double bond > triple bond

• Bond strength: single bond < double bond < triple bond

BOND STRENGTH, LENGTH AND BOND TYPE

Remember: the longer the length the lower the bond energies the lower the bond STRENGTH

Another example:

Trend atomic size: I > Br > ClLets compare: C-I > C-Br> C-Cl

How the Model Explains the Properties of Covalent Substances

• Most covalent substances have low electrical conductivity because electrons are localized and ions are absent.

• The covalent bonding model proposes that electron sharing between pairs of atoms leads to strong, localized bonds, usually within individual molecules.


• Most covalent substances have low electrical conductivity because electrons are localized and ions are absent.

• The covalent bonding model proposes that electron sharing between pairs of atoms leads to strong, localized bonds, usually within individual molecules.

Pentane


• Substances that consist of separate molecules are generally soft and low melting because of the weak forces between molecules.

• Solids held together by covalent bonds extending throughout the sample are extremely hard and high melting.

COVALENT BONDING

Key concepts

Covalent bonding

Shared and unshared pair

Bond energy and bond length

Electronegativity

Bond polarity

Dipole moment

BONDING

The ionic and covalent bonding models portray compounds as being formed by either complete electron transfer or complete electron sharing.

Questions: are all ionic bonds are 0% covalent?YESQuestions: are all covalent bonds are 0% ionic?Let’s learn about electronegativity?

One of the most important concepts in chemical bonding is electronegativity (EN).

(EN), the relative ability of a bonded atom to attract the shared electrons.

ELECTRONEGATIVITY

Electronegativity different from electron affinity (EA), • Electronegativity refers to a bonded

atom attracting the shared electron pair; • Electron affinity refers to a separate

atom in the gas phase gaining an electron to form a gaseous anion.

ELECTRONEGATIVITY

•There are also trends in electronegativity in the periodic table

• Electronegativity is inversely related to atomic size.

Atomic size decreases nucleus of a bonded atom attract more SHARED electrons electronegativity increases

•For the main-group elements, electronegativity generally increases up a group and across a period.

ELECTRONEGATIVITY

Questions: are all covalent bonds are 0% ionic?YES in NONPOLAR BONDNO in POLAR BOND

Because:Difference in ELECTRONEGATIVITY affect an electrostatic (charge) contribution

POLAR AND NONPOLAR COVALENT BONDS

Let’s have a look at:

H-H and H-F

POLAR COVALENT BOND:The electron pair is shared unequally

NONPOLAR COVALENT BOND:

H-H or H : H• Both H atoms have the same electronegativity electrons spend equal amount of time near each H nucleausthe electron density is SYMETRICAL The electron pair is shared equally • this covalent bond is NONPOLAR

The covalent bonds in ALL homonuclear diatomic molecules must be nonpolar.

H2, O2, N2, F2 and Cl2

Now consider heteronuclear diatomic molecules: HF

H-F or H : F

• H and F are two different atoms different electronegativity uneven electron density polar

• The electron pair is shared unequally• Asymetric electron density

POLAR COVALENT BOND

POLAR AND NONPOLAR COVALENT BONDS

Different distribution of electron density

Partial positive charge

Partial negative charge

one way to indicate polarity

Another way to indicate polarity

POLARITY: indicate how polar a compound is

Compare POLARITY?Which on is more polar?

HF HCl HBr HI

The higher the difference in EN the more polar a compound the higher polarity

• The separation of charge in a polar covalent bond creates an electric dipole.

We can express bond polarities on a numerical scale as dipole moment

DIPOLE MOMENT µ

d: distance of separationq: charge

DIPOLE MOMENT

Different distribution of electron density

DIPOLE MOMENT

What happened if we put a polar compound in an electric field?

FIELD OFF FIELD ON

So WE understood:

• Covalent bonding : polar and nonpolar

bond

• Electronegativity: is the attraction of

electron more toward a bonded atom

• Difference in EN causes a bond polar

• Polarity can be expressed in number by a

value called DIPOLE MOMENT

Question: is the polarity of a bond the same as a molecule containing that bond?

• Dipole moments only associated with individual bonds

The polarity of the entire molecules depends on MOLECULAR SHAPE (GEOMETRY) next part

COVALENT BONDING

SUMMARY

Key concepts

• Covalent bonding

• Shared and unshared pair

• Bond energy and bond length

• Bond polarity

• Dipole moment

Highlight:

• Lewis structure

• Octet rule

• Molecular

geometry

THIS WEEK


CONCEPTS Octet rule How to draw a LEWIS structure VSEPR (Valence-Shell Electron Pair Repulsion) theory Determine the molecular shape by VSEPR

MOLECULAR SHAPE (GEOMETRY)

We will know to explain the geometries of the molecules in term of their electronic structures.

Molecular geometry: • the general shape of a molecules• determined by relative positions of the atomic nuclei

To see what a molecule look like:• Know the molecular formula• 2D structure with Lewis dot formula:

the position of bonding pair, the position of lone pair central atom

should know the OCTET RULE

LEWIS FORMULAS FOR MOLECULES

LEWIS FORMULAS FOR POLYATOMIC IONS

An example: NH4+

In most of their compounds, the representative elements achieve noble gas configuration

OCTET RULE: elements tend to reach a maximum 8 electrons in the outermost shell - lowest energy/stable configuration (Except for H: 2 electrons)

H •F•••

••••

This rule is not always correct, there are some exceptions

OCTET RULE

In most of their compounds, the representative elements achieve noble gas configuration

OCTET RULE: elements tend to reach a maximum 8 electrons in the outermost shell - lowest energy/stable configuration (Except for H: 2 electrons)

H •F•••

••••

OCTET RULE

How can we calculate the number of shared electrons in a compound?

How can we calculate the number of shared electrons in a compound?

S= N - AS: number of shared electrons

N: the total number of valence electrons needed by all the atoms in the molecule or ion to achieve noble gas configurations

A: the number of valence electrons of all of the atoms. (so it should be the group number?)

N = 8 numbers of atoms that are not H + 2 number of H atoms

Some examples?

HF:

N = 8×1 (1 atom F) + 2×1 (1 atom

H) = 10

A = 7×1 (1 atom F) + 1×1 (1 atom

H) = 8

S = N-A = 10-8 = 2 e- sharedH •F•

••

••••

CO2:

N = 8×2 (2 atom O) + 8×1 (1 atom

C) = 24

A = 6×2 (2 atom O) + 4×1 (1 atom

C) = 16

S = N-A = 24-16 = 8 e- shared

Some examples?

H2O: N =

A =

S =

NH4+: N =

A =

S =

How to draw a Lewis structure with single bonds

Select a skeleton for molecules or ionsStep 1

Calculate the shared electrons (S)

Put the shared electrons into the skeleton

Step 2

Step 4Put the unshared electrons into the skeleton to fulfill Octet rule

Step 3


A) The least electronegative element is usually the central element

B) Oxygen atoms do not bond to each other: don’t put them close to one another

How to draw a Lewis structure

C) Hydrogen usually bonds to an O atom, not to the central atom

D) For ions or molecules that have more than one central atom, the most symmetrical skeletons possible are used.

How to draw a Lewis structure with single bonds


A) The least electronegative element is usually the central element

B) Oxygen atoms do not bond to each other: don’t put them close to one another

E.g. : CS2 S C S

Some exceptions:a) O2 and O3

b) Peroxides : O22-;

c) Superoxides: O2

-

E.g. : SO4-

O S

O

O

O



C) Hydrogen usually bonds to an O atom, not to the central atom

Some exceptions:H3PO3 and H3PO2

D) For ions or molecules that have more than one central atom, the most symmetrical skeletons possible are used.

E.g. : Nitrous acid HNO2 H O N O

How to draw a Lewis structureHow to draw a Lewis structure with single bonds

Some examples: draw a Lewis structure of

a) H2SO4 b) ClO4 c)NO3

a) H2SO4

Step 1: Draw a skeleton

Step 2: Calculate number of shared electronsN= 8×4 (4O) + 8×1 (1S) + 2×2 (2H) = 44

A= 6×4 (4O) + 6×1 (1S) + 1×2 (1H) = 32

Number of shared electrons:S= N-A = 12 e-

How to draw a Lewis structureHow to draw a Lewis structure with single bonds

OH O S O H

O

a) H2SO4

Step 3: put the shared e- in the skeleton

Step 4: put the unshared e- in the skeleton..

..

..

..

..

..

...... ..

Some examples: draw a Lewis structure of

a) H2SO4 b) ClO4- c)NO3

-

How to draw a Lewis structure with MULTIPLE bonds


Calculate the shared electrons (S)

Put the shared electrons into the skeleton

Step 2

Sometimes after step 4, a central atom does not have an octet: MAKE MULTIPLE BOND by changing a lone pair into a bonding pair

change 1 pair: from single bond to doublechange 2 pairs: double to triple

Put the unshared electrons into the skeleton to fulfill Octet rule

Step 3

How to draw a Lewis structure with MULTIPLE bonds

Sometimes after step 4, a central atom does not have an octet: MAKE MULTIPLE BOND by changing a lone pair into a bonding pair

change 1 pair: from single bond to doublechange 2 pairs: double to triple

After STEP 4:Move a

lone pair to

bonding pair

Practice: draw a Lewis structure of some hydrocarbon


Lewis structures for Exceptions to the Octet Rule: self-study, further reading in Principle of General Chemistry.

CH4 C2H6 C2H4 C2H2

Sometimes, there COULD be 2 Lewis structure for 1 molecule

RESONANCE

They are not correct structure; the bonds should be between O-O and O=O “one-and-a-half” bond

Resonance structures

RESONANCE: DELOCALISATION

Resonance structures

• Resonance structures are not real bonding, so to depict something in between, RESONANCE HYBRID is used.

In Resonance hybrid, electron-pair delocalisation occur

RESONANCE: DELOCALISATION

Another example of electron-pair delocalisation

Highlight:

• Lewis structure

• Octet rule

• Molecular

geometry


CONCEPTS Octet rule How to draw a LEWIS structure VSEPR (Valence-Shell Electron Pair Repulsion) theory Determine the molecular shape by VSEPR


VSEPR MODEL : the valence-shell electron-pair repulsion.

VALENCE BOND THEORY

SHAPE

BONDING

• The valence-shell electron-pair repulsion model (VSEPR) can be used to construct the molecule shape from Lewis structure.

• VSEPR THEORY: each group of valence electrons around a CENTRAL ATOM is located as far away as possible from the others in order to minimize repulsions.

• Electron group: regions around the central atom where electrons are likely to be found


• So valence electron groups can be: Single bond Double bond Triple bond A lone pair Or even a lone electron

Electron group: regions around the central atom where electrons are likely to be found


• Each group REPEL each other to minimise energy increase BOND ANGLE

Is the structure of a covalent compound is ALWAYS flat?

NO, because there are repulsion between bond pair/ lone pair

• According the valence-shell electron-pair repulsion theory (VSEPR), these repulsions give rise to 5 geometric arrangements.


VSEPR MODEL

5 geometric arrangements


Valence electron groups defy the arrangement

Bond angle: is the angle formed by the nuclei of two surrounding atoms with the nucleus of the central atom at the vertex (top of angle).


• The valence-shell electron-pair repulsion model (VSEPR) can be used to predict the shape of an ABn molecule when A is a main group element.

AXnEwhere A = central atom, main group

elementX = outer atom(s), E: lone pair(s)n = # of “B” atoms

MOLECULAR SHAPE (GEOMETRY)• So remember we have 5 molecular shape

Linear Arrangement

• The molecular shape with 2 electron groups

• 2 electron groups are as far apart as each other

• Linear arrangement of electron groups

Linear shape bond angle of 180o

Linear Arrangement• The molecular shape with 2 electron

groups• Gaseous Beryllium Chloride (BeCl2)

• Carbon dioxide (CO2)

Remember: only electron groups around CENTAL atom influence shape

Trigonal planar Arrangement• The molecular shape with 3 electron

groups• 3 electron groups repel

each other• Trigonal arrangement of

electron groups Trigonal planar shape bond angle of 120o

• Electron groups can be: bonds (single and double) and lone pair


groups• Boron trifluoride (BF3)

• The nitrate ion (NO3-)


groups• Boron trifluoride (BF3)

• The nitrate ion (NO3-)

Trigonal planar Arrangement

• The bond angel change when 3 electron groups are not identical:

If one of them is a DOUBLE BOND

If one of them is a LONE PAIR

• Let’s examine

The effect of double bond on bond angle

The effect of a lone pair on bond angle


Example: formaldehyde (CH2O)

The effect of double bond on bond angle

The actual bond angles deviate from the ideal because • the double bond with its greater electron density repels the two single bonds more strongly than they repel each other.


The effect of lone pair on bond angle When one of the groups is a lone pair, the shape is BENT, or V-shaped

A lone pair can have a major effect on bond angle

Gaseous tin(II) chloride

a lone pair repels bonding pairs more strongly than bonding pairs repel each other. This stronger repulsion decreases the angle between bonding pairs.

Tetrahedral Arrangement• The molecular shape with 4 electron

groups

Tetrahedral Arrangement

• The molecular shape with 4 electron groups

• With 4 electron groups the molecule shape is in 3D.

• So, Lewis structures do not depict all shape

• Consider: Methane (CH4)

All molecules or ions with four electron groups around a central atom adopt the tetrahedral arrangement



Example: NH3


Example: H2O

Tetrahedral Arrangement: BOND ANGLE

BOND ANGLE

< <Electron-pair repulsions cause deviations from ideal bond angles in the following order:

Trigonal bibyramidal Arrangement• The molecular shape with 5 electron groups

• Examples can be found in text books

Octahedral Arrangement• The molecular shape with 6 electron

groups

VSEPR MODEL: MORE EXAMPLES

STEPS TO DETERMINE A MOLECULAR SHAPE by VESPR MODEL

• We understood how VESPR works for the molecule shape

• Now we learn to apply the model and determine a molecule shape from a molecular formula


Write the LEWIS structureStep 1

Choose suitable: Count the number of electron group then choose shape

Predict the ideal bond angle+ the direction of any deviation

Step 2

Step 4 Draw and name molecular shape

Step 3


Step 1

Example 1: PF3

Step 2

Step 3 Step 4

MOLECULAR SHAPE & DIPOLE MOMENT

•So we now know that molecules have different shapes•The shape influence the overall dipole moment

COVALENT BOND THEORY

Self study /further reading!

Hybrid orbital interacts

Hybrid orbitals

Sigma and Pi bonds


Highlight:

• Covalent bond

• Ionic bond

• Metallic bond

THIS WEEK

CONCEPTS Valence electron Lewis dot formula Electronegativity Polar, non polar bond Polarity Dipole, dipole moment

Highlight:

• Lewis structure

• Octet rule

• Molecular

geometry

THIS WEEK


CONCEPTS How to draw a LEWIS structure Octet rule VSEPR (Valence-Shell Electron Pair Repulsion) theory Determine the molecular shape by VSEPR

Documents

L3 Chemical Bond September2014