Periodicity

Guiding Questions

Why is the periodic table so important?

Why is the periodic table shaped the way it's shaped?

Why do elements combine? Why do elements react? What other patterns are there in the world and how do they help us?

Periodicity

• Elements in the PT are arranged in order of increasing atomic number.

• Elements in the same group - same chemical and physical properties.

• Across the period - repeating pattern of physical and chemical properties known as periodicity.

Periodic Trends

Properties such as• Atomic radii and ionic radii• First ionisation energy• Melting points• Electronegativityshow periodicity

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Atomic Radii (pm) of the Elements

Radius :

half the distance between neighbouring nuclei(from nucleus to outermost electons)

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Ionic Radius

• When an atom gains or loses electrons, the radius changes

• Cations are always smaller than their parent atoms (often losing an energy level)

• Anions are always larger than their parent atoms (increased e repulsions)

Atomic & ionic radii down a group

• Atomic radii is determined by 2 opposing factors- Shielding effect by the electrons of the inner shell(s)- Nuclear charge (due to protons) • Moving down the group, both the nuclear charge

and shielding effect increase. However, the outer electrons enter new shells. So, although the nucleus gains protons, the electrons are not only further away, but also more effectively screened by an addtional shell of electrons.

Ionic radii for ions of the same charge also increases down a group for the same reason.

Atomic radius increases down the group

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Isoelectronic Series– Two or more species having the same

electron configuration (same number of electrons) but different nuclear charges– Size varies significantly

Atomic Radius vs. Atomic Number

Li

Na

K

Rb

Cs

La

XeKr

Zn

Cl

FHe

H

3dtransition

series

4dtransition

series

0.3

0.25

0.2

0.15

0.1

0.05

00 10 20 30 40 50 60

atomic number

ato

mic

radiu

s

Ionic radii across a period• The radii of positive ions decrease from Na+ to Al 3+

• The radii of positive ions decrease from P3- (phosphide ion) to Cl -

• The ionic radii increase from the Al 3+ to P3- .

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Explain

• What do you notice about the atomic radius across a period? Why?

• What do you notice about the atomic radius down a column? Why?

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• What do you notice about the atomic radius across a period? Why?

Atomic radius decreases from left to right across a period due to increasing nuclear charge but no significant increase in the shielding effect. The force of attraction between the negatively charged valence electrons and the positive nucleus increases across the period.

• What do you notice about the atomic radius down a column? Why?

Atomic radius increases down a column of the periodic table because the distance of the electron from the nucleus increases as n increases.

Ionisation Energy (IE)

• The first ionisation energy is the energy required to remove one electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

• X(g) X+(g) + e-

152

Li

60

Li+

e

152

60

Li+

e

e

e

e

Li

Lithium atom

Lithium ion

+

152

Li

Lithium atom

Li

Energy

Things to remember about IE

• When talking about first ionization energies (IE), everything must be in gas form

• IE are measured in kilojoules per mole. • All elements have a first ionization energy, even

those that do not form cations. What can you conclude if their ionization energy is very

high ?It is difficult to lose an electron.

Successive ionisation energies

Na(g) Na+(g) + e

Na+(g) Na2+

(g) + e

Na2+(g) Na3+

(g) + e

IE of Aluminium

What pattern do you notice?

What does this suggest about the energy levels?

Multiple Ionization Energies

Al Al+ Al2+ Al3+

578

kJ/m

ole

-

1817

kJ/m

ole

-

2745

kJ/m

ole

-

The second, third, and fourth ionization energies of aluminum are higher than the first because the inner electrons are more tightly held by the nucleus.

1st Ionizationenergy

2nd Ionizationenergy

3rd Ionizationenergy

Smoot, Price, Smith, Chemistry A Modern Course 1987, page 190

1st IE

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IE1 (kJ/mol) Values for Main Group Elements

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Explain1. What do you notice about the 1st IE across a

period? Why? • When moving across the period from left to

right, the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell).

• Hence, the electron shells are pulled progressively closer to the nucleus and it is harder to remove the valence electron.

Explain2. What do you notice about the 1st IE down a

column? Why? • The atomic radius increases down the group

as additional electrons are added, causing the shielding effect to increase.

• The further the outer shell is from the nucleus, the smaller the attractive force exerted by the protons in the nucleus.

• More easily an outer electron can be removed, the lower the ionisation energy.

First ionisation across a period

• When moving across the period from left to right, the nuclear charge increases but the shielding effect only increases slightly (since electrons enter the same shell).

• Hence, the electron shells are pulled progressively closer to the nucleus.

BUT.........there are exceptions

18

18

8

8

2

1s

2s 2p

3s 3p

4s 4p 3d

5s 5p 4d

NUCLEUS

Be 1s22s21st I.E. = 900 kJ mol-

1

B 1s22s22px1

1st I.E. = 799 kJ mol-

1

Removal of electrons

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Adding of electrons

Discussion

Factors affecting ionisation energy

Limitation to Bohr’s Model

• The first IE of the elt 3 (Li) to 10 (Ne) do not increase evenly.

• There is a need for a more complex model of electron configurations than the Bohr model.

• Each main energy level is an atom is made up of sub energy levels (subshells).

Plenary - K U IAs a result of the lesson today I:

Know…

Understand…

Can use the information in the following other situations….

ElectronegativityThe electronegativity is the ability of an atom in a

covalent bond to attract shared pairs electrons to itself.

• The greater the electronegativity of an atom, the greater its ability to attract shared pairs of electrons to itself.

• Electronegativity value is based on the Pauling scale. A value of 4.0 is give to F (most electronegative atom). The least electronegative elements, Ce and Fr both have a value of 0.7

Trends in electronegativity

• There is an increasing distance between the nucleus and electrons down the group. Hence, the attractive force is decreased.

• Although the nuclear charge increases down the group, this is counteracted by the increased shielding effect due to additional electron shells.

Trends in melting pointGroup I• Metals are held together by metallic bonding.• The strength of metallic bonding decreases because the

attractive forces between the delocalised electrons and the nuclues decreases owing to the increase in the distance. The increase in the nuclear charge is counteracted by the increase in shielding.

Group 7• As the molecules become large, the attractive forces between

them increases with the number of electrons in atoms or molecules.

The melting point decreases down group 1 The melting point increases down group 7

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Periodic Trends in Chemical Properties of Main Group Elements

• IE and EA enable us to understand types of reactions that elements undergo and the types of compounds formed

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• General Trends in Chemical Properties– Elements in same group have same

valence electron configuration; similar properties – Same group comparison most valid if

elements have same metallic or nonmetallic character –Group 1A and 2A; Group 7A and 8A –Careful with Group 3A - 6A

Chemical Properties

Group I alkali metals• Li, Na and K contain 1 valence electron.• Reactive metals, stored under liquid paraffin

to prevent them from reacting with air.• Readily lose their valence electron -good

reducing agent• Reactivity increases down the group

• React with water to form an alkali solution of the metal hydroxide and hydrogen gas.

(i) 2Li(s) + 2H2O(l) LiOH(aq) + H2(g)

Lithium floats and reacts quietly(ii) 2Na(s) + 2H2O(l) NaOH(aq) + H2(g)

Sodium melts into a ball which darts around on the surface

(iii) 2Ks) + 2H2O(l) KOH(aq) + H2(g)

Heat generated from the reaction with potassium ignites the hydrogen.

• React readily with chlorine, bromine and iodine to form ionic salts, e.g.

(i) 2Na(s) + Cl2(g) 2NaCl(s)

(ii) 2K(s) + Br2(l) 2KBr(s)

(iii) 2Ks) + I2(g) 2LiI(s)

Chemical Properties • Chlorine is a stronger oxidizing agent than

bromine, so can remove the electron from bromide ions in solution to form chloride ions and bromine.

• Similarly, both chlorine and bromine can oxidize iodide ions to form iodine.

(i) Cl2(aq) + 2Br-(aq) 2Cl-(aq) + Br2(aq)

(ii) Cl2(aq) + 2I-(aq) 2Cl-(aq) + I2(aq)

(iii) Br2(aq) + 2I-(aq) 2Br-(aq) + I2(aq)

Test for halide ions• The presence of halide ions in solution can be detected by

adding silver nitrate solution.

Ag+(aq ) + X- (aq) AgX(s)

X = Cl,Br or I light

AgCl white Ag(s) + ½ X2

AgBr creamAgI yellow