MME 2001 MATERIALS SCIENCE 1 8.10.2015

outline ● overview of atomic structure

● periodic table: organization/groups/periods

ionization energy

electronegativity

● Classification of elements

● Metals

● Nonmetals

● Metalloids

● interatomic bonding

● primary vs secondary bonding

● covalent/ionic/metallic

● bonding energy-bonding force

Atomic Structure HELIUM

ATOM

+ N

N

+ -

-

proton

electron neutron

Shell

nucleus

ATOMIC MASS NUMBER = number of protons + number of neutrons

ATOMIC NUMBER = number of protons

# electrons = # protons

Varying electrons: ions

Varying neutrons: isotopes

rules ● Heisenberg’s uncertainty principle

we cannot precisely measure the momentum and

the position of an electron at the same time.

● Pauli exclusion principle

No 2 Electrons in an atom Can Have the same 4

Quantum Numbers; the same n, l, ml, and ms.

● The Aufbau principle

states that electrons are added to the lowest energy

orbitals first before moving to higher energy

orbitals.

● Hund’s rule,

we must fill each shell with one electron each

before starting to pair them up.

The charge and mass number of an electron are:

a) charge = 0, Mass number = 1

b) charge = -1, Mass number = 0

c) charge = +1, Mass number = 1

d) charge = +1, Mass number = 0

The charge and mass number of a neutron are?

a) charge = +1, Mass number = 1

b) charge = 0, Mass number = 1

c) charge = +1, Mass number = 0

d) charge = -1, Mass number = 0

Atomic structure

The nucleus of the element having atomic number 25 and atomic weight 55 will contain?

a) 25 protons and 30 neutrons

b) 30 protons and 25 neutrons

c) 55 protons

d) 55 neutrons

Atomic structure

A beryllium atom has 4 protons, 5 neutrons, and 4 electrons. What is the mass number of this atom?

a) 4

b) 5

c) 8

d) 9

e) 13

Atomic structure

Which has the least mass in an atom?

a) nucleus

b) proton

c) neutron

d) electron

If uranium-235, has 92 protons, how many

neutrons does the isotope uranium-238 have?

a) 92

b) 95

c) 143

d) 146

Atomic structure

every electron in an atom is characterized by four

quantum numbers.

There are three quantum numbers necessary to

describe an atomic orbital.

The principal quantum number (n)

designates size

The angular moment quantum number (l)

describes shape

The magnetic quantum number (ml)

specifies orientation

quantum numbers

Angular moment Quantum Number (l)

l signifies the subshell

l describes the shape of the orbital.

l values range from 0 to n – 1

Example: If n = 2, l can be 0 or 1.

n 1 2 3 4 5 6

l

subshell

0 0,1 0,1,2 0,1,2,3 0 4 0 5

s s,p s,p,d s,p,d,f s,p g s,p h

energy state 1 3 5 7 9 11

Quantum Numbers To summarize quantum numbers:

principal (n) – size

angular (l) – shape

magnetic (ml) – orientation

electron spin (ms) direction of spin

Required to

describe an

atomic orbital

Required to describe an electron in an atomic orbital

2px

principal (n = 2)

angular momentum (l = 1)

related to the magnetic

quantum number (ml )

Electron Spin Quantum Number-ms

used to specify an electron’s spin.

There are two possible

directions of spin.

Allowed values of ms

are +½ and −½.

The lowest principal quantum number for

an electron is?

a) 0

b) 1

c) 2

d) 3

e) 4

Atomic structure

Quantum numbers An electron with

n = 2, ℓ = 1, ml = −1, and ms = +1/2

is found in the same atom as a second electron with

n = 2, ℓ = 1, ml = −1.

What is the spin quantum number for the second

electron?

Since the first three quantum numbers are identical for

these two electrons, we know that they are in the same

orbital. As a result, the spin quantum number for the second

electron cannot be the same as the spin quantum number

for the first electron. This means that the spin quantum

number for the second electron must be ms = −1/2.

Atomic structure Maximum number of electrons in a subshell with

l = 3 and n = 4 is

a) 10

b) 12

c) 14

d) 16

e) 18

Principal

Quantum No: n Subshell

l No. of energy States: ml

Number of

Electrons

Per Subshell

4

s / 0 1 / 0 2

p / 1 3 / -1,0,+1 6

d / 2 5 / -2,-1,0,+1,+2 10

f / 3 7 / -3,-2,-1,0,+1,+2,+3 14

Atomic structure

If n=3, and l=2, then what are the possible

values of ml ?

Since ml must range from –l to +l,

then ml can be: -2, -1, 0, 1, or 2.

Quantum Numbers: A Macroscale Analogy

n - indicates which train (shell)

l - indicates which car (subshell)

ml - indicates which row (orbital)

ms - indicates which seat (spin)

No two people can have exactly the same ticket (sit in the same seat).

Electron Configurations Q. the full electronic configuration of an element is

1s22s22p5.

How many electrons does it have in its outer

shell?

Q: the full electronic configuration of an element.

1s22s22p5.

What is its atomic number?

A. # of outer shell-valence electrons: 7

A. Atomic number: 9

Electronic Configurations

Fe-atomic # = 26

1s

2s 2p

K-shell n = 1; 2 electrons

L-shell n = 2; 10 electrons

3s 3p M-shell n = 3; 18 electrons

3d

4s

4p 4d

Energ

y

N-shell n = 4

1s2 2s2 2p6 3s2 3p6 3d 6 4s2

20 electrons

6 electrons left to be located

total # e-s

Electron Configurations rules for electron configurations:

● Electrons will reside in the

lowest possible energy orbitals

● Each orbital can accommodate

a maximum of two electrons.

● Electrons will not pair in

degenerate orbitals if an

empty orbital is available.

● Orbitals will fill in the order

..3p6/4s2/3d10/4p6/5s2/4d10/

5p6/6s2/4f14/5d10/6p6/7s2

Electron Configurations

1s22s22p5

1s

2s

2p 2p 2p

En

erg

y

The ground state electron

configuration of F

F has a total of 9 electrons

When there are one or more unpaired

electrons, as in the case of oxygen and

fluorine, the atom is called paramagnetic.

Electron Configurations

1s22s22p6

1s

2s

2p 2p 2p

En

erg

y

The ground state electron

configuration of Ne

Ne has a total of 10 electrons

When all of the electrons in an atom are

paired, as in neon, it is called diamagnetic.

learning check Write the electron configuration and give the orbital

diagram of a calcium (Ca) atom (Z = 20).

Z = 20, Ca has 20 electrons.

Each s subshell can contain a maximum of

two electrons, whereas each p subshell can

contain a maximum of six electrons.

Solution

Ca 1s22s22p63s23p64s2

1s2 2s2 2p6 3s2 3p6 4s2

Remember that the 4s orbital fills

before the 3d orbitals.

electron configuration for an arsenic atom (Z = 33) in the

ground state.

Z = 18 for Ar.

The order of filling beyond the noble gas

core is 4s, 3d, and 4p. Fifteen electrons go

into these subshells because there are 33 –

18 = 15 electrons in As beyond its noble gas

core.

2

2

2

2

6

6

3

10

Solution

As [Ar]4s23d104p3

Arsenic is a p-block element; therefore,

we should expect its outermost electrons

to reside in a p subshell.

learning check

They occupy the outermost shell.

They participate in the bonding between atoms

They dictate the physical and chemical properties

if the outermost or valence electron shell are

completely filled: stable electron configurations

occupation of the s and p states for the outermost

shell by a total of eight electrons, in neon (Ne),

argon (Ar), and krypton (Kr); inert, or noble, gases,

which are virtually unreactive chemically.

Valence electrons

Which one of the following is a proper orbital

configuration?

electron configuration?

Learning check The electrons with principle energy level n = 2 of a stable atom of

boron (atomic number = 5) would have an electron arrangement of

(a) ( ↑ ↓) ( ↑ ) ( ) ( )

(b) ( ↑ ) ( ↑ ) ( ↑ ) ( )

(c) ( ) ( ↑ ) ( ↑ ) ( ↑ )

(d) ( ) ( ↑ ↓ ) ( ↑ ) ( )

(e) ( ↑ ↓) ( ↑ ↓ ) ( ↑ ) ( ↑ )

Which of the following electron arrangements does not represent

an atom in its ground state?

(1s) (2s) (2p) (3s)

(a) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ )

(b) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ )

(c) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ) ( ↑ )

(d) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( )

beyond the d-orbitals

lanthanides

actinides

‘s’-groups ‘p’-groups

d-transition elements

f-transition elements

group

peri

od

1s2

2s2/2p6

3s2/3p6/

4s2/3d10/4p6

5s2/4d10/5p6/

6s2/4f14/5d10/6p6

you will have a quiz next week!

The Periodic Law

Mendeleev realized that:

When arranged by increasing atomic

number, the chemical elements display a

regular and repeating pattern of chemical

and physical properties.

what are these properties?

Metallic vs nonmetallic character

Atomic radius

Ionization energies (energy necessary to remove

the outermost electron from the atom)

Electron affinities (energy change when an

electron is added to a neutral atom)

Reactivity

Electronegativity

Organisation of the periodic table

The vertical columns: groups from 1 to 18.

Elements in the same group have similar valence

electron structures

and hence similar

chemical and

physical

properties.

groups

elements are situated, with increasing atomic

number, in seven horizontal rows

called periods.

Each contains

elements with

electrons in the

Same outer shell.

Organisation of the periodic table

periods

Periodic table

Periodic table

The first period contains elements with electrons in the first electron shell

only. Hydrogen and helium thus have behaviours very different to the lower

periods and are not easily classified into the groups used to describe the rest of

the table.

Periodic table

electrons fill the second electron shell which has both s and p orbitals. Elements

in this and the next period typically follow the so called octet rule, whereby

stable compounds are formed with 8 electrons in the outer shell.

Periodic table

elements with valence electrons in the 3s and 3p orbitals.

Elements in this period typically follow the so called octet rule, whereby stable

compounds are formed with 8 electrons in the outer shell.

Periodic table

elements with valence electrons in the 4s, 4p and 3d orbitals.

The octet rule is no longer applicable due to the introduction of the d subshell.

Periodic table

The fifth period contains elements with valence electrons in

the 5s, 5p and 4d orbitals.

Periodic table

32 elements with valence electrons in the 6s, 6p, 5d, and - including the

lanthanides - the 4f orbital. This period contains the last stable element, lead,

with all later elements being radioactive.

Periodic table

The seventh period contains 32 elements with valence electrons in the

7s,7p, 6d, and - including the actinides - the 5f orbital.

Periodic table

The alkali metals are soft, highly reactive metals with one electron in

their outermost s subshell. The reactivity of these elements increases

down the group

Periodic table

The alkaline earth metals are all reactive metals with two electrons in their

outermost s subshell. In general they are harder, denser, and have higher

melting points than their alkali metal analogues

Periodic table

A group of transition metal elements, the lightest two of which are exceptions

from the Aufbau principle (to determine the structure of the atom), showing

valence configurations of d5s1.

Periodic table

The chalcogens, or oxygen family is formed of non metals (oxygen and sulfur)

and metalloids and its elements are characterised by having 6 electrons in

their outer shell.

Periodic table

The halogens are a group of highly reactive elements with 7 electrons in their

outer shell. This is the only group which contains elements in all three states of

matter at room temperature and pressure.

Periodic table

The noble gases are typically relatively unreactive and are

characterised by a full outer electron shell.

Periodic table

The s block consists of elements with their valence electrons in s orbitals. Elements

within the s-block all behave fairly similarly, being soft, reactive metals. The s sub-shell

can contain a maximum of two electrons, and hence the block is two columns wide

Periodic table

The p block consists of elements with their valence electrons in p orbitals. The

characteristics of elements within the p-block are fairly varied, including metals and

non-metals and so called 'metalloids'. The p sub-shell can hold six electrons, in three

distinct orbitals known as px, py and pz

Periodic table

The d-block, also known as the ‘transition metals’ contains only metals,

typically capable of existing in at least two stable oxidation states. The d sub-

shell can hold up to 10 electrons in 5 distinct orbitals.

Periodic table

The f block consists of the lanthanides and actinides which are all soft metals

many of which are not found in nature. The f sub-shell can contain up to 14

electrons in seven distinct orbitals

● The elements of the rightmost group, are the inert

gases, with filled electron shells and stable

electron configurations.

Periodic table

● Group VIIA and VIA elements are one and two

electrons deficient from having stable structures.

● The Group VIIA elements (F, Cl, Br, I, and At) are

sometimes termed the halogens.

Periodic table

● The alkali and the alkaline earth metals (Li, Na, K,

Be, Mg, Ca, etc.) are labeled as Groups IA and IIA,

with one and two electrons in excess of stable

structures.

● The elements in the three long periods, Groups

IIIB through IIB, are termed the transition

metals, which have partially filled d electron

states and in some cases one or two electrons in

the next higher energy shell.

● Groups IIIA, IVA, and VA (B, Si, Ge, As, etc.)

display characteristics that are intermediate

between the metals and nonmetals by virtue of

their valence electron structures.

Periodic table

Learning check

The elements in each vertical column on the periodic

table usually have similar properties and are called

a(n)

a) period

b) group

c) element

d) Property

Elements on the periodic table are arranged in order of

a) increasing density.

b) decreasing density.

c) increasing atomic number.

d) decreasing atomic number.

An element has the electronic structure 2,8,4. Which

group is it in?

a) Group 3

b) Group 4

c) Group 5

d) Group 6

Learning check

Which of these electronic structures belongs to a noble

gas?

a) 2

b) 2,2

c) 2,8,2

d) 2,8,4

Two elements have these electronic structures: 2,1

and 2,8,1. What can you say about the elements?

a) They are both in group 1

b) They are both in group 2

c) They are both in period 1

d) They are both in period 2

Learning check

Periodic Table metallic character

nonmetallic character meta

llic

chara

cte

r

nonm

eta

llic

chara

cte

r

İoniz

ati

on e

nerg

y

Negati

ve e

lectr

on a

ffin

ity

İonization energy

Negative electron affinity

Atomic radii

Ato

mic

radii

Ionization Energy

IE = energy required to remove an

electron from an atom in the gas phase.

Mg (g) + 738 kJ ---> Mg+ (g) + e-

Atomic Size

Size goes UP on going down a group.

Because electrons are added farther from

the nucleus, there is less attraction.

Size goes DOWN on going across a period.

Atom size

increases!

Atom size

decreases!

atomic

radii (picometer//

10-12m)

● the tendency of an atom to attract electrons

towards itself.

● An atom's electronegativity is affected by both

its atomic number and the distance at which

its valence electrons reside from the charged

nucleus.

● The higher the associated electronegativity

number, the more an element or compound

attracts electrons towards it.

Electronegativity

electronegativity increases in moving from

left to right and from bottom up.

Atoms are more likely to accept electrons

if their outer shells are almost full, and if

they are less “shielded” from (i.e., closer

to) the nucleus.

Electronegativity

electronegativity

increases!

Electronegativity

Electronegativity

Electronegativity

Learning check

Which elements are the most electronegative

element of those shown in the diagrams below?

Learning check Fluorine has a lower electronegativity than

a) Oxygen

b) Chlorine

c) Lithium

d) None of the above

Metals, Nonmetals & Metalloids 1

2

3

4

5

6

7

Metals

Metalloids

Nonmetals

Metals 88 elements are metals or metal like element

Physical properties:

good conductors of heat and electricity

shiny

ductile (can be stretched into thin wires)

malleable (can be pounded into thin sheets)

High density (heavy for their size)

High melting point

Metals

Metals

chemical properties:

Easily lose electrons

Form positive (+) ions

Corrode easily

Non-metals Non-metals are on the right of the stairstep line.

Their characteristics are opposite to those of metals.

Physical Properties of Nonmetals:

No luster (dull appearance)

Poor conductor of heat and electricity

Brittle (breaks easily)

Not ductile

Not malleable

Low density

Low melting point

Many non-metals are gases.

Non-metals Chemical Properties of Non-metals:

Tend to gain electrons

metals that tend to lose electrons but nonmetals

that tend to gain electrons, to form compounds

with each other.

These compounds are called

ionic compounds.

When two or more

nonmetals bond with

each other, they form

a covalent compound.

Metalloids Metalloids (metal-like) have properties of both

metals and non-metals.

They are solids

can be shiny or dull

They conduct heat

and electricity better

than non-metals but

not as well as metals

They are ductile and

malleable

Non-metals

Learning check

Which of the following is a property of alkali metals?

a) They are so hard they cannot be cut.

b) They are very reactive.

c) They are stored in water.

d) They have few uses.

Most of the elements in the periodic table are

a) metals.

b) metalloids.

c) gases.

d) nonmetals.

Learning check The elements to the right of the zigzag line on the

periodic table are called

a) metalloids.

b) conductors.

c) metals.

d) nonmetals.

Transition metals are

a) good conductors of thermal energy.

b) more reactive than alkali metals.

c) not good conductors of electric current.

d) used to make aluminum.

Learning check

What do the elements on the far right of the

table (He, Ne, Ar, and Kr) have in common?

a)They are liquid in normal conditions

b)They are metals that rust easily

c)They are very reactive gases

d)They do not generally react with other elements

Learning check Which of the following statements describes most

metals?

a) They are easily shattered.

b) They are gases at room temperature.

c) They are dull.

d) They are good conductors of electric current.

Elements lying along the zigzag line on a periodic table

are

a) metals

b) nonmetals

c) metalloids

d) noble gases

Learning check Elements in a period have …….

a) a wide range of chemical properties

b) the same atomic radius

c) similar chemical properties

d) the same number of protons

The elements in Group 1 of the periodic table are

commonly called the…..

a) alkali metals

b) transition metals

c) alkaline earth metals

d) rare earth metals

Learning check Elements in a group have

a) a wide range of chemical properties

b) the same atomic radius

c) similar chemical properties

d) the same number of protons

What do the elements on the far right of the table

(He, Ne, Ar, and Kr) have in common?

a) They are liquid in normal conditions

b) They are metals that rust easily

c) They are very reactive gases

d) They do not generally react with other elements

● Some general behaviors of the various material

types (i.e., metals, ceramics, polymers) may be

explained by bonding type.

● For example, metals are good conductors of both

electricity and heat, as a consequence of their

free electrons.

● By way of contrast, ionically and covalently

bonded materials are typically electrical and

thermal insulators because of the absence of

large numbers of free electrons.

bonding and properties

interatomic bonding

● the bonding involves the valence electrons

● the nature of the bond depends on the electron

structures of the constituent atoms.

● There are three types of bonding: each bonding

type arises from the tendency of the atoms to

assume stable electron structures.

● Secondary or physical forces and energies are

weaker than the primary ones, but nonetheless

influence the physical properties of some

materials.

interatomic bonding Ionic

Metal (cation) with non-metal (anion)

Transfer of electron(s)

Strong bond high melting point

Covalent

Non-metal with non-metal

Sharing of electron(s)

Non-polar (equal distribution of electrons)

Polar (uneven electron distribution)

Weak bonds…low melting points

Metallic (nuclei in a “sea” of shared electrons)

Bonding forces

● physical properties of materials = f (interatomic

forces that bind the atoms together)

● two isolated atoms interact as they are brought

close together from an infinite separation.

● At large distances, interactions are negligible,

because the atoms are too far apart to have an

influence on each other; however, at small

separation distances, each atom exerts forces

on the other.

● The origin of an attractive force FA depends on

the particular type of bonding that exists

between the two atoms.

● Repulsive forces (FR) arise from interactions

between the negatively charged electron clouds

for the two atoms

● they are important only at small values of r as

the outer electron shells of the two atoms begin

to overlap.

Bonding forces

The minimum energy

corresponds to the

equilibrium spacing, r0.

the bonding energy for

these two atoms, E0,

corresponds to the

energy at this

minimum point;

it represents the energy

required to move these

two atoms to an infinite

separation.

Bonding force and bonding energy

Equilibrium

interatomic spacing

Bonding forces The net force FN between the two atoms is just the

sum of both attractive and repulsive components

FN = FA+ FR

When FA and FR balance, or become equal, there is

no net force; implying a state of equilibrium

FA + FR = 0

The centers of the two atoms will remain separated

by the equilibrium spacing r0

● Forms between metallic and nonmetallic

elements; elements at the horizontal extremities

of the periodic table.

● a metallic atom easily gives up its valence

electrons to the nonmetallic atoms.

● In the process all the atoms acquire stable

configurations and become ions.

● Ionic bonding is non-directional (magnitude of the

bond is equal in all directions around the ion)

● Ceramic materials exhibit ionic bonding

ionic bonding

• Occurs between + and - ions.

• Requires electron transfer.

• Large difference in electronegativity required.

Ionic Bonding

Na (metal)

Unstable

11 electrons

Cl (nonmetal) Unstable

17 electrons electron

+ -

Coulombic Attraction

Na (cation) stable

Cl (anion) stable

positive and negative ions, by virtue of their net electrical charge, attract one another

ionic bonding

Schematic

representation of

ionic bonding in

sodium chloride

(NaCl).

ionic bond : metal + nonmetal

donates accepts

electrons electrons

Dissimilar electronegativities

ex: MgO

Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4

Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6

ionic bond - Electronegativity

• Predominant bonding in Ceramics

Ionic Bonding - examples

Give up electrons Acquire electrons

NaCl

MgO

CaF 2

CsCl

ionic bonding - Periodic Table Columns: Similar Valence Structure

Electropositive elements:

Readily give up electrons

to become + ions.

Electronegative elements:

Readily acquire electrons

to become - ions.

giv

e u

p 1

e

giv

e u

p 2

e

giv

e u

p 3

e

iner

t g

ases

acce

pt

1e

acce

pt

2e

O

Se

Te

Po At

I

Br

He

Ne

Ar

Kr

Xe

Rn

F

Cl S

Li Be

H

Na Mg

Ba Cs

Ra Fr

Ca K Sc

Sr Rb Y

• Ranges from 0.7 to 4.0,

Smaller electronegativity Larger electronegativity

• Large values: tendency to acquire electrons.

Electronegativity

Ion Sizes

CATIONS are SMALLER than the atoms from

which they come.

The electron/proton attraction has gone UP

and so size DECREASES.

Li,152 pm 3e and 3p e

Li + , 78 pm 2e and 3 p

+ Forming

a cation.

Ion Sizes

ANIONS are LARGER than the atoms from

which they come.

The electron/proton attraction has gone DOWN

and so size INCREASES.

Trends in ion sizes are the same as atom sizes.

Forming

an anion. F, 71 pm 9e and 9p + e

F - , 133 pm 10 e and 9 p

-

learning check

Which atom or ion has the smallest radius?

a) O2+

b) O+

c) O

d) O2–

Learning check

How do the size of a negative ion compare to the

size of the atom that formed it?

a) it's smaller

b) it's larger

c) it's the same size

d) it varies

● for ionic materials to be stable, all positive ions

must have as nearest neighbors negatively

charged ions in a three dimensional scheme.

● The predominant bonding in ceramic materials

is ionic.

● Ionic materials are characteristically hard and

brittle and, electrically and thermally

insulative.

● These properties are directly related to

electron configurations and/or the nature of the

ionic bond.

ionic bonding

For two isolated ions, the attractive energy EA is a

function of the interatomic distance

EA = - A/r

An analogous equation for the repulsive energy is

ER = B/rn

A, B, and n are constants whose values depend on

the particular ionic system.

The value of n is approximately 8.

ionic bonding

Ionic Bonding Energy – minimum energy most stable

Energy balance of attractive and repulsive terms

Attractive energy EA

Net energy EN

Repulsive energy ER

Interatomic separation r

r A

n r B

EN = EA + ER =

Bonding Forces and Energies calculate the force of attraction between a K+ and

an O2- ion separated by r0 =1.5 nm.

The attractive force between two ions FA is the

derivative with respect to the interatomic separation

of the attractive energy expression,

FA

dEA

dr

= dr

d r

A (

)

= = 1

40

= Z1e ( ) Z2e (

) A r

A2

( )

0 is the permittivity of vacuum (8.85x10-12 F/m). Z1 & Z2 are the valences of the two ion types

e is the electronic charge (1.602x10-19 C).

FA = (Z1e) (Z2e)

40r2

= (1)(2)(1.602 1019 C)2

(4)() (8.85 1012 F/m) (1.5 109 m)2

=2.05 10-10 N

Since the valences of the K+ and O2- ions

(Z1 and Z2) are +1 and -2, respectively,

Z1 = 1 and Z2 = 2, then

Bonding Forces and Energies

covalent bonding ● stable electron configurations are assumed by the

sharing of electrons between adjacent atoms.

● Two atoms that are covalently bonded will each

contribute at least one electron to the bond, and

the shared electrons may be considered to belong

to both atoms.

● The covalent bond is directional; it is between

specific atoms and may exist only in the direction

between one atom and another that participates

in the electron sharing.

C: has 4 valence e-,

needs 4 more

H: has 1 valence e-,

needs 1 more

Electronegativities

are comparable.

Covalent Bonding ● similar electronegativity share electrons

● bonds are determined by valence –

● s & p orbitals dominate bonding

● Example: CH4 shared electrons

of carbon atom

shared electrons of hydrogen atom

H

H

H

H

C

CH 4 H

∙∙

H : C : H

∙∙

H

Covalent Bonding ● The bonds between oxygen and hydrogen in a

water molecule are covalent bonds.

● There are two covalent bonds in a water

molecule, between the oxygen and each of the

hydrogen atoms.

● Each bond represents one

electron.

● In a covalent bond,

electrons are shared

between atoms,

not transferred.

Fluorine has seven valence electrons

A second atom also has seven

By sharing electrons

Both end with full orbitals

F F 8 Valence

electrons

Covalent Bonding

Writing Lewis Structures The Lewis structure contains the element symbol with dots

representing electrons.

The only electrons

shown are those on

the outer energy

level or valence

electrons.

The electrons are

placed around the

element symbol,

one at a time,

clockwise or

counterclockwise,

and then grouped in pairs as more electrons are added.

Covalent bonding in H2 molecule

Covalent bonding in H2O molecule

covalent bonding ● Covalent bonds may be very strong, as in

diamond, which is very hard and has a very high

melting temperature, 3550 C, or they may be

very weak, as with bismuth, which melts at about

270 C.

● Polymeric materials typify this bond, the basic

molecular structure often being a long chain of

carbon atoms that are covalently bonded together

with two of their available four bonds per atom.

● The remaining two bonds normally are shared

with other atoms, which also covalently bond.

covalent bonding ● interatomic bonds may be partially ionic and

partially covalent.

● very few compounds exhibit pure ionic or covalent

bonding.

● the degree of either bond type depends on the

relative positions of the components in the periodic

table or the difference in their electronegativities.

● The wider the separation (the greater the difference

in electronegativity), the more ionic the bond.

● the closer they are (the smaller the difference in

electronegativity), the greater the degree of

covalency.

interatomic bonding

No electronegativity difference between two

atoms leads to a purely non-polar covalent bond.

A B

A small electronegativity difference leads to a

polar covalent bond.

A B

A large electronegativity difference leads to an

ionic bond.

ionic-covalent mixed bonding Ionic-Covalent Mixed Bonding

ionic character =

where XA & XB are Pauling electronegativities

%) 100 ( x

1e

(XAXB)2

4

ionic 70.2% (100%) x e1 characterionic % 4

)3.15.3(

2

Ex: MgO XMg = 1.3 XO = 3.5

metallic bonding ● found in metals and their alloys.

● Metallic materials have one, two, or at most,

three valence electrons.

● these valence electrons are more or less free to

drift throughout the entire metal. They may be

thought of as forming a “sea of electrons” or an

“electron cloud”.

● The remaining nonvalence electrons and atomic

nuclei form what are called ion cores, which

possess a net positive charge equal in magnitude

to the total valence electron charge per atom.

metallic bonding ● the metallic bond is nondirectional in character. In

addition, these free electrons act as a “glue” to

hold the ion cores together.

● Bonding may be weak or strong; energies range

from 68 kJ/mol (0.7 eV/atom) for mercury to 849

kJ/mol (8.8 eV/atom) for tungsten.

Their respective melting temperatures are 39 and

3410 C.

● Metallic bonding is found in the periodic table for

Group IA and IIA elements and, in fact, for all

elemental metals.

metallic bonding

Metallic bonding

• Covalent bonds can be strong

e.g, Diamond melting point >3550°C

or, covalent bonds can be weak

e.g, Bismuth melting point: 270°C

• Polymers: Covalent bonds

• Partially ionic + partially covalent: possible

• Wider separation in the periodic table: Ionic

• Closer together in the periodic table: Covalent

Primary Bonding-summary

Bonding energies correlate with melting points.

interatomic bonding

Secondary-van der waals-bonding Secondary, van der Waals, or physical bonds are

weak in comparison to the primary or chemical

ones; bonding energies are typically on the order

of only 10 kJ/mol (0.1 eV/atom).

Secondary bonding exists between virtually all

atoms or molecules, but its presence may be

obscured if any of the three primary bonding types

is present.

Secondary bonding is evidenced for the inert gases,

which have stable electron structures, and, in

addition, between molecules in molecular

structures that are covalently bonded.

Secondary-van der waals- bonding

Schematic illustration of van der Waals

bonding between two dipoles

Secondary bonding forces arise from atomic or

molecular dipoles. In essence, an electric dipole

exists whenever there is some separation of positive

and negative portions of an atom or molecule.

The bonding results from the coulombic attraction

between the positive end of one dipole and the

negative region of an adjacent one

Hydrogen bonding, a special type of secondary bonding, is found to exist

between some molecules that have hydrogen as one of the constituents.

Arises from interaction between dipoles

asymmetric electron

clouds

+ - + - secondary

bonding

H H H H

H 2 H 2

secondary bonding

ex: liquid H 2

Secondary bonding Molecular dipoles occur due to the unequal sharing of

electrons between atoms in a molecule.

More electronegative atoms pull the bonded electrons closer

to themselves.

This results in a molecular dipole in which one side of the

molecule possesses a partially negative charge and the other

side a partially positive charge.

● Melting Temperature, Tm

Tm is larger if Eo is larger.

Properties linked with Bonding:Tm

r o r

Energy

higher Tm

smaller Tm

• Coefficient of thermal expansion,

• a ~ symmetry at ro

is larger if Eo is smaller.

Properties From Bonding :

= ( T 2 - T 1 ) D L

L o

coeff. thermal expansion

D L

length, L o

unheated, T 1

heated, T 2

r o r

smaller

larger

Energ

y

unstretched length

Eo

Eo

L F

Ao = E

Lo

Elastic modulus

r

larger Elastic Modulus

smaller Elastic Modulus

Energy

ro unstretched length

E is larger if curvature is larger.

E similar to spring constant

Properties From Bonding: E

Summary ● Polymers, with weak, secondary, intermolecular

bonds (low melting points) have very high

expansion coefficients.

● Ceramics which are strongly bonded (i.e., ionic

or network covalent) have low thermal

expansion coefficients.

● Metals with high melting points (strong bonding)

have low thermal expansion coefficients. Low

melting point metals have high thermal

expansion coefficients.

Summary

The Periodic Table

● Elements in each of the columns (or

groups) of the periodic table have

distinctive electron configurations.

● For example, Group 0 elements (the inert

gases) have filled electron shells, and

● Group IA elements (the alkali metals) have

one electron other than a filled electron

shell.

Summary

● There is a sharing of valence electrons between

adjacent atoms when bonding is covalent.

● Polymers and some ceramic materials bond

covalently.

● The percent ionic character (%IC) of a bond

between two elements (A and B) depends on

their electronegativities (X’s).

● Relatively weak van der Waals bonds result from

attractive forces between electric dipoles, which

may be induced or permanent.

Type

Ionic

Covalent

Metallic

Secondary

Bond Energy

Large!

Variable large-Diamond small-Bismuth

Variable large-Tungsten small-Mercury

smallest

Comments

Nondirectional (ceramics)

Directional

(semiconductors, ceramics

polymer chains)

Nondirectional (metals)

Directional

inter-chain (polymer)

inter-molecular

Summary: Bonding

Ceramics

(Ionic & covalent bonding):

Metals

(Metallic bonding):

Polymers

(Covalent & Secondary):

Large bond energy

large Tm

large E

small a

Variable bond energy

moderate Tm

moderate E

moderate a

Directional Properties

Secondary bonding dominates

small Tm

small E

large a

Summary: Primary Bonds

Explain why covalently bonded materials are

generally less dense than ionically or

metallically bonded ones.

Learning check

because covalent bonds are directional in

nature whereas metallic and ionic bonds are

not; when bonds are directional, the atoms

cannot pack together in as dense a manner,

yielding a lower mass density.

Learning check If the difference in electronegativities between two atoms

is zero, the bonds are

a) Non polar covalent

b) Polar covalent

c) Mostly ionic

d) Slightly ionic

The bond between O and H in OH

a) Nonpolar covalent

b) Very polar

c) Slightly polar covalent

d) Mostly ionic

Learning check Electronegativity refers to

a) the degree of negative charge on an electron

b) the energy required to remove an electron from a gaseous

atom in the ground state

c) the ability of an atom to attract the electrons in a

covalent bond toward itself

d) the energy change that occurs when an electron is

accepted by a gaseous atom to form an anion

In general, electronegativity _______ going left to

right across a row in the periodic table

a) decreases

b) İncreases

c) Does not change

d) None of the above

Learning check

The most electronegative elements are

a) found in the upper right corner of the periodic table

b) the alkali metals

c) The alkaline earth metals

d) The transition elements

The level of attraction of one atom for electrons when

bonding with another atom is called

a) ionization energy

b) An ionic bond

c) A nonpolar covalent bond

d) electronegativity

Learning check When sodium and chlorine react, chlorine removes

sodium's valence electron and __________ forms

between them

a) a covalent bond

b) a nonpolar covalent bond

c) an ionic bond

d) A polar covalent bond

When an electron pair is shared between two atoms of

equal electronegativity,

a) a nonpolar covalent bond is formed

b) An ionic bond is formed

c) a polar covalent bond is formed

d) electron transfer occurs

Learning check A polar covalent bond results from

a) a transfer of electrons to the atom of least

electronegativity

b) an equal sharing of an electron pair between two atoms

c) the formation of oppositely charged ions

d) None of the above

The type of bond that involves a cation and an anion is

_____.

a) nonpolar covalent

b) Metallic

c) Polar covalent

d) ionic

Learning check Of the following, the most likely pair to form an ionic

bond is ….

a) an alkali metal and an alkaline earth metal

b) a halogen and an alkaline earth metal

c) a halogen and a metalloid

d) an alkaline earth metal and a transition element

Which of the following is in an ionic bond?

a) F2

b) MgCl2

c) NO

d) H2O

Learning check In a crystalline compound, each anion is surrounded

by…...

a) Negative ions

b) Molecules

c) Positive ions

d) Dipoles

When a metal forms an ionic bond with a non-metal, the

nonmetal atom will …………...

a) gain an electron and become a positive ion

b) lose an electron and become a positive ion

c) lose and electron and become a negative ion

d) gain an electron and become a negative ion

Learning check The less the electronegativity differences between two

bonded atoms, the greater the ……………...

a) polar character

b) ionic character

c) Metallic character

d) Covalent character

What type of bonding does NaCl have?

a) polar covalent

b) Metallic

c) Nonpolar covalent

d) ionic

Learning check How are bond length and bond energies related?

a) the higher the bond energy, the shorter the bond

length

b) the lower the bond energy, the shorter the bond length

c) they are not related

d) the higher the bond energy, the longer the bond length

What determines bond length?

a) the distance at which potential energy is at a minimum

b) the distance at which the two atoms are as close as

possible

c) the distance at whch potential energy is at a maximum

d) the point at which the attraction forces outweighs the

repulsion forces

Learning check How are thermal expansion coefficient () and bond

energies related?

a) the higher the bond energy, the smaller

b) the lower the bond energy, the smaller

c) they are not related

d) the higher the bond energy, the larger

Bonding between nonmetals and nonmetals primarily

involves?

a. interactions between protons, electrons, and neutrons

b. interactions between protons

c. interactions between protons and electrons

d. transfer of electrons

e. sharing of electrons

Learning check A bond in which electrons are shared equally is known

as ……………..

a) polar covalent

b) Metallic

c) İonic

d) Non-polar covalent

Mobile electrons within bonding networks best

describes which type of bond?

a) Metallic

b) İonic

c) Polar covalent

d) Non-polar covalent

see you next week!