The How and Why

2

HistoryElements known to ancients: C, Cu, Au, Fe,

Pb, Hg, Ag, S, Sn

Added before 1700: As, Sb, Bi, P, Zn

Dobereiner, Johann (1780-1849): arranged elements in triads (Ca, Sr, Ba; Cl, Br, I)

John Newlands (1837-1898): arranged elements in group of eight

Properties repeat every 8th elements:

Li, Be, B, C, N, O, F, Na

Na, Mg, Al, Si, P, S, Cl, K

3

History

Dmitri Mendeleev (1834-1907): used the masses of elements as most of the masses were determined in XIX century (1869)

Arranged elements in order of increasing atomic masses

Found a pattern of repeating properties

4

Mendeleev’s Table

Grouped elements in columns by similar properties in order of increasing atomic mass.

Found some inconsistencies - felt that the properties were more important than the mass, so switched order ( Te, I).

Found gaps in the trends- maybe undiscovered elements.

Predicted their properties before they were found

( eka boron – Sc; eka aluminum- Ga; eka silicon: Ge).

5

The Modern Table

Elements are still grouped by properties.Similar properties are in the same column.Order is in increasing atomic number

(Moseley, 1914).Added a column of elements Mendeleev

didn’t know about (Noble Gases).The noble gases weren’t found because they

didn’t react with anything.

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Periodic Law

Mendeleev (1869): Properties of elements are a function of the atomic masses of the elements.

Modern periodic Law (Mosley, 1914) properties of elements are a periodic function of their atomic numbers ( # of protons in the nucleus). Explains Mendeelev’s irregularities.

7

Horizontal rows are called periods There are 7 periods1

2

3

4

5

6

7

8

Vertical columns are called groups or families

Elements are placed in columns by similar properties.

9

1A

2A 3A 4A 5A 6A7A

8A0

The elements in the A groups are called the representative elements (also numbered 1-18)

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The group B are called the transition elements

Inner Transition elements

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Group 1A(1) are the alkali metals

Group 2A(2) are the alkaline earth metals

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Group 7A (17) is called the Halogens

Group 8A (18) are the noble or inert gases

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Why the similarities in Properties?

The part of the atom another atom sees is the electron cloud.These are the outside or valence

orbitals. The orbitals fill up in a regular pattern.

The outside orbital electron configuration repeats.

The properties of atoms repeat.

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1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10 5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s2

4f145d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

15

He2

Ne10

Ar18

Kr36

Xe54

Rn86

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

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Alkali metals all end in s1

Alkaline earth metals all end in s2

Have to include He but it fits better later.

He has the properties of the noble gases.

s2s1 S- block

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Transition Metals -d block

d1 d2 d3s1

d5 d5 d6 d7 d8s1

d10 d10

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The P-block p1 p2 p3 p4 p5 p6

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F - block inner transition elements

f1 f5f2 f3 f4

f6 f7 f8 f9 f10 f11 f12 f14

f13

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Each row (or period) is the energy level for s and p orbitals.

1

2

3

4

5

6

7

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d orbitals fill up after previous energy level … first d is 3d even though it’s in row 4.

1

2

3

4

5

6

7

3d

22

f orbitals start filling at 4f

1

2

3

4

5

6

7 4f

5f

Writing Electron configurations the easy way

Review NotesReview Notes

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Electron Configurations Repeat The shape of the periodic table is a

representation of this repetition of electron configurations.

When we get to the end of the column the outermost energy level is full.

This is the basis for our shorthand.

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Trends and Properties of Elements in the Periodic Table

The following properties will be examined:

• Radius of the atom

• Ionization energy

• Electron affinity

• Electron negativity

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Effective Nuclear Charge Effective nuclear charge is experienced

by an outer electron at the outer edge of an atom

Zeff = Z – S Z is the atomic numberS is the number of core electrons.See blackboard for examples.

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Atomic Size (Radius)First problem where do you start measuring.

The electron cloud doesn’t have a definite edge.

Determined by measuring more than 1 atom at a time.

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Atomic Size

Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

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Trends in Atomic Size

Influenced by two factors.

Energy Level

Higher energy level is further away from nucleus

Charge on nucleus ( Zeffective)

More charge pulls electrons in closer.

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Group Trends

As we go down a group

Each atom has another energy

level

So the atoms get bigger.

HLi

Na

K

Rb

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Periodic TrendsAs you go across a period the radius

gets smaller.Filling the same energy level.

***More nuclear charge (higher Zeff).Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

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Atomic Radii in the PT

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Overall

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

34

Ionization Energy (IE)The amount of energy required to completely remove an electron

from a gaseous atom.

Removing one electron makes a

+1 ion.

The energy required to remove the first e- is called the first ionization energy.

A(g) + IE → A(g) +1

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Ionization EnergyThe 2nd ionization energy is the energy required to remove the second electron.

2nd IE is always greater than 1st IE.

The 3rd IE is the energy required toremove a third electron.

3rd IE > 2nd IE > 1st IE

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Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

810 14840 3569 4619 4577 5301 6045 6276

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Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

810 14840 3569 4619 4577 5301 6045 6276

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What Determines IE

1. The greater the nuclear charge the greater IE.

2. Distance form nucleus influences IE.

3. Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE.

4. Shielding effect

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ShieldingThe electron on the outside

energy level has to look through all the other energy levels to see the nucleus. It is shielded from the nucleus by all the inner electrons

A second electron in the same energy level has the same shielding.

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Group Trends

As you go down a group first IE decreases because

The electron is further away

More shielding by inner electrons as there are

more energy levels.

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Periodic trendsAll the atoms in the same period have the

same energy level.

Same shielding.

Increasing nuclear charge increases the force of attraction between the nucleus and the electrons.

IE generally increases from left to right.

Exceptions at full and 1/2 fill orbitals.

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Firs

t Ion

izat

ion

ener

gy

Atomic number

He

He >IE than H

same shielding

greater nuclear charge H

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li < IE than H

more shielding

further away (>n)

outweighs greater nuclear charge

Li

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be > IE than Li

same shielding

greater nuclear charge

Li

Be

45

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

HeB < IE than Be

same shielding

greater nuclear charge

By removing an electron we make s-orbital half filled

Li

Be

B

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

47

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

48

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

Breaks the pattern because removing an electron gets to 1/2 filled p orbital

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

NeNe < IE than He

Both are full,

Ne has more shielding

Greater distance (>n)

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Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

NeNa < IE than Li

Both are s1

Na has more shielding

Greater distance

(>n)

Na

52

Firs

t Ion

izat

ion

ener

gy

Atomic number

53

Driving Force for Ionization

Full Energy Levels are very low energy.

Noble Gases have full orbitals.

Atoms behave in ways to achieve noble gas configuration.

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2nd Ionization Energy For elements that reach a filled or

half filled orbital by removing 2 electrons 2nd IE is lower than expected.

True for s2

Alkali earth metals form +2 ions.

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3rd IE Using the same logic s2p1

atoms have a low 3rd IE.

Atoms in the aluminum family form + 3 ions.

2nd IE and 3rd IE are always higher than 1st IE!!!

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Electron Affinity (EA)The energy change associated with adding an

electron to a gaseous atom.

Easiest to add to group 7A.Filled energy level.

EA increases from left to right atoms become smaller, with greater

nuclear charge.

EA decreases as go down a group.

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Ionic Size

Cations form by losing electrons.

Cations are smaller that the atom they come from.

Metals form cations.

Cations of representative elements have noble gas configuration.

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Ionic size

Anions form by gaining electrons.

Anions are bigger than the atom they come from.

Nonmetals form anions.

Anions of representative elements have noble gas configuration.

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Ionic Radii

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Configuration of IonsIons always have noble gas configuration.

Na is 1s12s22p63s1

Forms a +1 ion - 1s12s22p6

Same configuration as neon.

Metals form ions with the configuration of the noble gas before

them - they lose electrons.

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Configuration of IonsNon-metals form ions by gaining

electrons to achieve noble gas configuration.

They form the configuration of the noble gas after

them.

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Group trends

Adding energy level

Ions get bigger as you go down a column

Li+1

Na+1

K+1

Rb+1

Cs+1

63

Periodic TrendsAcross the period nuclear charge

increases so ions get smaller.

Energy level changes between anions and cations.

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

64

Size of Isoelectronic ions

Iso - same

Iso-electronic ions have the same # of electrons

Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3

10 electrons

configuration: 1s12s22p6

65

Size of Atoms and Ions

66

Size of Isoelectronic ionsPositive ions have more protons so

they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

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Atomic Radii and Ionic Radii Compared

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Comparison af Atomic and Ionic Radii

Electronegativity

70

Electronegativity The tendency for an atom to attract

electrons to itself when it is chemically combined with another element.

How fair it shares. Big electronegativity means it pulls the

electron toward it. Atoms with large negative electron

affinity have larger electronegativity. Scale designed by Linus Pauling

71

Group Trend The further down a group the

farther the electron is away and the more electrons an atom has.

More willing to share.

Low electronegativity.

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Periodic Trend Metals are at the left end. They lose electrons easily

Low electronegativity

At the right end are the nonmetals. They gain electrons.

High electronegativity.

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Ionization energy, electronegativity

Electron affinity INCREASE

74

Atomic size increases, shielding constant

Ionic size increases

75

Summary ot trends in the Periodic Table

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Oxidation States

Check Blackboard

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Metal, nonmetals, metalloids, & noble gases

78

Properties of Metals Elements with 3 or less electrons in outer

level Electrons loosely bound Become positive ions Good electric and heat conductors Malleable, ductile Most are solids at room temperature

(exception : mercury) Grayish, silvery in color (exceptions: Cu, Au)

79

Properties of Nonmetals Elements with 5 or more electrons in outer

level Gain electrons easily to become negative

ions Brittle Can be solid, liquid, or gas at room

temperature Good insulators Many have colors in naturals state (sulfur –

yellow, iodine: purple)

80

Transition Elements Located in B group elements (or groups

3-12) Many have multiple oxidation states

(because of the d-orbitals) Many form colorful compounds: Ni+2

green; Cu+2 – blue; Mn+7, purple All are metals, some with the highest

known melting points (Tungsten)

81

Properties of Metalloids Located on the zig-zag line Behave both as metals and

nonmetals with some exhibiting stronger metallic character (aluminum)

82

Noble Gases Located in group 8A(or 18) Most are inert (exceptions: Xe –

some compounds with oxygen and fluorine are known)

High ionization energies. Have octet of electrons (exception:

He with only 2 electrons)