CHEMICAL PERIODICITY

Development of Periodic TableDmitri Mendeleev, Russian Chemist in the mid-1800’s

arranged the 70 known elements into a systematic way.

– Put the names of elements on a card(Just like you arranged the pieces)

– He also put atomic masses, physical and chemical properties.

– When arranged in the order of atomic mass, he discovered a repeated, or periodic, pattern.

– He left blank spaces where he didn’t know the element for the spot.

– By doing this, he constructed the first periodic table for elements.

Mendeleev’s Problems

1. Why could most of the elements be arranged in the order of increasing atomic mass but a few could not?

2. What was the reason for chemical periodicity?

Henry Moseley’s Solution

1. Instead of arranging in atomic mass, in1913 Henry Moseley arranged in increasing order of nuclear charge.

His work led to modern definition and recognition of atomic number.

2. He discovers that it is the number of electrons that affect an atoms’ reactivity, or chemical property.

Parts of the Periodic Table

• Periods:– The horizontal rows

• Periodic Law:– When the elements are arranged in order of

increasing atomic number, there is a periodic repetition of their physical and chemical properties.

• Groups: (a.k.a. Families)– The vertical columns

Parts of the Periodic Table, cont.

• Representative Elements:–Group A elements. (The tall parts of the

periodic table.)

–These can be divided into 3 groups:• Metals• Nonmetals• Metalloids

Metals• Malleable-Very easy to be hammered or beaten into thin

sheets

• Ductile-Very easy to produce a wire

• High luster-De-excitation of electrons causes shiny appearance of metal surface

• High electrical conductivity - Due to the free movement of electrons

• High thermal conductivity• Most are solid at room temperature• Metallic Bond Strength

– Directly proportional to the heat of vaporization

• 80% of all elements are metals

Nonmetals• Upper right portion of the periodic table

• Generally nonlustrous

• Generally poor conductors of electricity

• Some are gases at room temperature and some brittle solids

• One is even a liquid at room temperature

Metalloids

• These border the stair step, except for aluminum

• These have properties of both metals and nonmetals

• Si, Ge, As, Sb, Te

Group 1 – Alkali Metal

• Very reactive!Not found in nature as free elements• Combine vigorously with most nonmetals• React strongly with water to produce hydrogen gas and

aqueous solutions of substances known as “alkalis.”• Usually stored in kerosene (paraffin oil)

Group 2 – Alkali Earth Metal• Harder, denser, and stronger than alkali metal

• Higher melting points

• Not so reactive compared to alkali metal;Still very reactive! Not found as free element

Group 3 – 12 Transition Metal• Good conductors of electricity and have a

high luster

• Less reactive than alkali and alkali-earth metal

• Some (e.g. platinum and gold) are so unreactive that they do not form compounds easily.

• Some are found as free element.

Group 17 - Halogens• Most reactive nonmetal• React vigorously with most metals to

form “salts”• Need one electron to obtain stable

noble-gas configuration• F and Cl are gaseous.• Br is liquid.• Iodine is solid.

Group 18 – Noble Gases

• Lack chemical reactivity

• Very stable

• Octet electronic configuration in the outermost energy level (or “shell”)

Categories of electrons• Core (inner) electrons – all those shared by the

previous noble gas

• Outer electrons – those in the highest occupied energy level– Similar chemical properties of elements in groups is a result

of similar outer electron configurations– In the main group, the group number equals the number of

outer electrons

• Valence electrons – those involved in bonding– In the main group, the outer electrons are valence– In the transition metals can include some d electrons

Trends in metallic behavior

Metallic Behavior• Metallic behavior increases left and

down• Metals tend to lose electrons in

reactions• Highly metallic elements, are likely to

make positive ions• Least metallic elements are likely to

make negative ions• In middle are more likely to make

covalent bonds

Ion (Cation and Anion)

Neutral atomAdd electron

Negative ion (Anion)

Neutral atomRemove electron

Positive ion (Cation)

# of electrons added/removed= # of charge

F+ 1e-

F-1 (Anion)

Ca-2e-

Ca+2 (Cation)

Main-group ions and the noble gas configurations

Octet (Duplet) Rule

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

8.2

Periodic LawAll the elements in a group have the same electron configuration in their outermost shells

Example: Group 2Be 1S2 2S2

2, 2

Mg 1S2 2S2 2p6 3S2

2, 8, 2

Ca 1S2 2S2 2p6 3S2 3p6 4S2

2, 8, 8, 2

Periodic Table and Electron Configuration

s1 s2 p1 p2 p3 p4 p5

p6

123 d1 - d10

456

f1 - f14

Groups 1-2 = s levelGroups 3-8 = p levelTransition = d levelLantanides = f level

Figure 8.13 The relation between orbital filling and the periodic table

Figure 8.12

© 1998 by Harcourt Brace & Company

s p

d (n-1)

f (n-2)

1234567

67

Periodic Patterns

Trends in Some Periodic Properties

• The physical and chemical behavior of the elements is based on the electron configurations of their atoms.

• e- configurations can be used to explain many of the repeating or “periodic” properties of the elements

Ionization energy: the energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)

Electronegativity is a measure of an atom’s attraction for another atom’s electrons.

Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).

Atomic Radius/Ionic Radius- size of atom/Ion

Definition of Some Periodic Properties

Atomic radii of the main-group and transition elements.

Vertically• The trend for atomic radius in a vertical

column is to go from smaller at the top to larger at the bottom of the family.

• With each step down the family, we add an entirely new energy level, making the atoms larger with each step.

Horizontally• Each step adds a proton and an

electron (and 1 or 2 neutrons).• Electrons are added to existing energy

level.The effect is that the more positive nucleus has a greater pull on the electron cloud.The nucleus is more positive and the electron cloud is more negative.The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

Electron Shells and Sizes of Atoms

• Size in main group elements follow 2 general rules– 1. AR increases going down in a group– 2. AR decreases going left to right in a period

• Why?• Moving down in a group:

– Zeff is constant because Z and S increase equally– Electrons in a higher n and therefore larger

• Moving across:– Z increases and S does not (electrons in same

shell don't shield each other well) so e- attracted more strongly and are closer/smaller

• In transition metals there is an initial decrease in size but then remains relatively constant

Figure 8.16

Periodicity of atomic radius

SAMPLE PROBLEM 8.3 Ranking Elements by Atomic Size

PLAN:

SOLUTION:

PROBLEM: Using only the periodic table (not Figure 8.15), rank each set of main group elements in order of decreasing atomic size:

(a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb

Elements in the same group increase in size and you go down; elements decrease in size as you go across a period.

(a) Sr > Ca > Mg These elements are in Group 2A(2).

(b) K > Ca > Ga These elements are in Period 4.

(c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr.

(d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.

Na

Mg

Al

Si

11

12

13

14

10

10

10

10

1

2

3

4

186

160

143

132

ZeffCoreZ Radius (pm)

Zeff = Z - s 0 < s < Z (s = shielding constant)

Zeff Z – number of inner or core electrons

Effective Nuclear Charge (Zeff)

increasing Zeff

incr

easi

ng Z

eff

Ions

• Here is a simple way to remember which is the cation and which the anion:

This is a cat-ion.This is Ann Ion.

He’s a “plussy” cat!She’s unhappy and negative.

+ +

Cation Formation

11p+

Na atom

1 valence electron

Valence e- lost in ion formation

Effective nuclear charge on remaining electrons increases.

Remaining e- are pulled in closer to the nucleus. Ionic size decreases.

Result: a smaller sodium cation, Na+

Anion Formation

17p+

Chlorine atom with 7 valence e-

One e- is added to the outer shell.

Effective nuclear charge is reduced and the e- cloud expands.

A chloride ion is produced. It is larger than the original atom.

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

The Radii (in pm) of Ions of Familiar Elements

Ionic Size• Cations are smaller than the parent atom

• Anions are larger than the parent atom.

• Ionic size increase down in a group

• In an isoelectronic series, the most negative ion is largest, the most positive is smallest.

• Atoms making more than one ion, most positive is smallest.

Ionization Energy• This is the second important periodic trend.• The atom has been “ionized” or charged.• The number of protons and electrons is no longer

equal.• The energy required to remove an electron from an

atom is ionization energy. (measured in kilojoules, kJ)• The larger the atom is, the easier its electrons are to

remove.• Ionization energy and atomic radius are inversely

proportional.• Ionization energy is always endothermic, that is

energy is added to the atom to remove the electron.

Ionization Energy (Potential)

• Draw arrows on your help sheet like this:

Ionization energy• Trends:

– Generally as size decreases, IE1 increases.– 1. IE1 generally increases from left to right, some

exceptions– 2. IE1 generally decreases going down in a group– 3. transition and f-block elements have much

smaller variances in IE1  • But Why?

– Across - increasing Zeff and smaller size(e- closer to nucleus)

– exceptions - Be to B because s shields p and lowers Zeff

– N to O because repulsions in first paired electron– Down - Zeff constant and larger so easier to ionize

Figure 8.17

Periodicity of first ionization energy (IE1)

Figure 8.18

First ionization energies of the

main-group elements

51

Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.

I1 + X (g) X+

(g) + e-

I2 + X+(g) X2+

(g) + e-

I3 + X2+(g) X3+

(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ionization energy

I1 < I2 < I3

SAMPLE PROBLEM 8.4 Ranking Elements by First Ionization Energy

PLAN:

SOLUTION:

PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE decreases as you proceed down in a group; IE increases as you go across a period.

(a) He > Ar > Kr

(b) Te > Sb > Sn

(c) Ca > K > Rb

(d) Xe > I > Cs

Group 8A(18) - IE decreases down a group.

Period 5 elements - IE increases across a period.

Ca is to the right of K; Rb is below K.

I is to the left of Xe; Cs is furtther to the left and down one period.

Ionization energy• Second ionization energy IE2 – to remove 2nd

e-– Always larger than IE1

– as e- are removed, Z remains constant and remaining e- are harder to remove

• A huge increases occurs in ionization energies when core electrons are reached because of the lower shielding/higher Zeff.

• Result - Rx only involve outer e-

Figure 8.19

The first three ionization energies of beryllium (in MJ/mol)

SAMPLE PROBLEM 8.5 Identifying an Element from Successive Ionization Energies

PLAN:

SOLUTION:

PROBLEM: Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5 IE6

1012 1903 2910 4956 6278 22,230

Look for a large increase in energy which indicates that all of the valence electrons have been removed.

The largest increase occurs after IE5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s23p3 and the element must be phosphorous, P (Z = 15).

The complete electron configuration is 1s22s22p63s23p3.

Electronegativity• Electronegativity is a measure of an atom’s

attraction for another atom’s electrons.• It is an arbitrary scale that ranges from 0 to

4.• The units of electronegativity are Paulings.• Generally, metals are electron givers and

have low electronegativities.• Nonmetals are are electron takers and have

high electronegativities. • What about the noble gases?

Electronegativity

• Your help sheet should look like this:

0

Overall Reactivity• This ties all the previous trends together

in one package.• However, we must treat metals and

nonmetals separately.• The most reactive metals are the largest

since they are the best electron givers.• The most reactive nonmetals are the

smallest ones, the best electron takers.

Overall Reactivity

• Your help sheet will look like this:

0

Electron Affinity

• What does the word ‘affinity’ mean?• Electron affinity is the energy change that

occurs when an atom gains an electron (also measured in kJ).

• Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

Electron Affinity

• Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy.

• If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic.

• This is true for the alkaline earth metals and the noble gases.

Electron Affinity

• Your help sheet should look like this:

++

64

Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

X (g) + e- X-(g)

F (g) + e- X-(g)

O (g) + e- O-(g)

DH = -328 kJ/mol EA = +328 kJ/mol

DH = -141 kJ/mol EA = +141 kJ/mol

electron affinities• energy required to put an e- on an atom

(making it negative or more negative)• EA1 usually negative • EA2 always positive

• Irregular trend: increases going right and up– Remember: Full shells best, full subshells good,

and half full is kinda cool– Cl and other halogens need only 1 e- so most

exothermic reaction– noble gases have endo because new electron in

higher n– Mg, Be endo because new e in higher l– N endo because new e paired

Figure 8.20 Electron affinities of the main-group elements

Figure 8.21

Trends in three atomic properties

Ato

mic rad

ii increases

Electronegativity Increases

Ele

ctro

neg

ativ

ity

Incr

ease

s

Atomic radii decreases