The Periodic Table

Chapter 5

Video clip

What Does Periodic Mean?

To answer: Turn to the picture on pg. 133, Fig. 1

What things occur periodically? Moon phases, magazine publications

Keep these in mind as we learn about the periodic table

History

Russian scientist Dmitri Mendeleev taught chemistry in terms of properties.

Mid 1800 - molar masses of elements were known.

Wrote down the elements in order of increasing mass.

Found a pattern of repeating properties.

The Table

Compare this table with our modern table

The Table

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 Example: Iodine after Tellurium

Found some gaps. Must be undiscovered elements. Predicted their properties before they were

found.

Interestingly…

Mendeleev never won a Nobel Prize. He was nominated shortly before his death,

but lost to Henri Moissan, who discovered Fluorine in 1906

Periodic Law

Why could most of the elements be arranged in the order of increasing mass, but a few could not? Henry Moseley- discovered that atomic number

should be the basis for the table, not mass Example: I = 53, Te = 52

Physical and chemical properties of the elements are periodic functions of their atomic numbers

The Modern Table Elements are still grouped by properties. Similar properties are in the same column. Order is in increasing atomic number. Added a column of elements Mendeleev

didn’t know about. The noble gases weren’t found because they

didn’t react with anything. So a new group was formed

Horizontal rows are called periods There are 7 periods

Vertical columns are called groups. Elements are placed in columns by

similar properties. Also called families

1A

2A 3A 4A 5A 6A7A

8A0

The elements in the A groups are called the representative elements

The group B are called the transition elements

These are called the inner transition elements and they belong here

Group 1A are the alkali metals Group 2A are the alkaline earth metals

Group 7A is called the Halogens Group 8A are the noble gases

Review

Rewrite the periodic law in your own words What group are the Alkali Metals in? Halogens?

Why do families have the same properties?

The part of the atom another atom “sees” is the electron cloud.

More specifically the outside (valence) orbitals. A “family’s” orbitals fill up in a regular pattern. The outside orbital electron configuration

repeats. The properties of atoms repeat.

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10 5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

He2

Ne10

Ar18

Kr36

Xe54

Rn86

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

Alkali metals (column 1) all end in s1

Alkaline earth metals (column 2) all end in s2

really have to include He but it fits better later.

He has the properties of the noble gases.

Demo- Ca and Mg in test tube (pg. 142)

s2s1

S- block: Reactive Metals

Transition Metals -d block

d1 d2 d3 d4d5 d6 d7 d8 d9

d10

The P-block:Main Group Elements

p1 p2 p3 p4 p5 p6

F - block

inner transition elements

f1 f5f2 f3 f4

f6

f7 f8 f9 f10 f11 f12 f14

f13

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

1

2

3

4

5

6

7

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

1

2

3

4

5

6

7

3d

f orbitals start filling at 4f

1

2

3

4

5

6

7 4f

5f

Review

Create an acronym for remembering the position of the s, p and d blocks of the table

The s-Block Elements

1 s electron 2 s electrons

Alkali metals Alkaline-earth metals

Extremely reactive A little less reactive

silvery vary

soft Stronger, harder

Not found in nature Not found in nature

Melt at lower temps Melt at higher temps

Group 1 Group 2

Quantum Formula: ns1,2

Example problems Sample Problem A on pg. 143 Without looking at the periodic table, identify the

group, period, and block that [Xe] 6s2 is located Answer: group 2, sixth period, s block Write the electron config. For the Group 1 element

in the third period. Will it be more reactive or less? Answer: Grp 1, third period= 1s22s22p63s1

Must be more reactive, because it’s in

group 1

Quantum Formula: ns1,2

The d-Block5 orbitals, 10 e’ total

Transition metals

Extremely reactive

Conduct electricity

High luster

Less reactive

Not always the same outer e’configuration

Groups 3-12

d block

Quantum Formula: ns0-2 (n-1)d1-10

Example Problem

Sample Problem B on pg. 146 For d-block problems, identify group using

this formula: d + s Without using the periodic table, identify the

period, block and group of [Kr] 4d55s1

Answer: 5 period, d-block, group 6 Molybedenum

Quantum Formula: ns0-2(n-1)d1-10

p- Block

Group number – 10 for electrons

Main group elements

Properties vary

Nonmetals, metalloids, halogens

Not always the same outer e’configuration

Groups 13-18

Quantum Formula: ns2np1-6

Example

Sample Problem C on pg. 148 Write the outer electron configuration for the Group

14 element in the second period, name it and identify it as a metal, nonmetal or metalloid

Group number is higher than 12- so it’s the P block, second period makes n=2, group number -10 = 14-10= 4

So, you have 2 left over for the p’s 2s22p2

Quantum Formula: ns2np1-6

Halogens (p-block) Group 17 Gases, mostly Most reactive of nonmetals Seven valence electrons

Metalloids (p-block) Groups btw metals and nonmetals solids Electrical conductivity

Metals (p-block) Harder and denser than alkaline-earth

metals, but softer than d-block Found only in compounds

Problem Examples Without the table, write the outer electron

config. For Group 14 in second period. Name it and classify it as a metal, nonmetal or metalloid

Answer: 14= p block, 14-10= 4 so 2s22p2, carbon= nonmetal

Quantum Formula: ns2np1-6

f- Block

Sixth and seventh periods La-Hf = Cerium-Lutetium Ac-Rf = Thorium-Lawrencium Mostly lab made (sythetic)

f orbitals start filling at 4f

1

2

3

4

5

6

7 4f

5f

More Example Problems Name the block and group for each and identify as metal,

nonmetal or metalloid [Xe]4f145d96s1

Answer: d-block, group 10, Pt, metal (period 6) [Ne]3s23p6

Answer: p-block, group 18, Ar, nonmetal (period 3)

Group #

Group Config

Block Comments

1, 2 ns1,2 s One or two electrons in ns

3-12 ns0-1(n-1)d1-10 d Sum of electrons in ns and (n-1)d equals group number

13-18 ns2np1-6 p Number of electrons in np sublevel equals group number +/- 10

Nobel gasesHalogens

NonmetalsMetalloids

Metals

Alkali Alkali-Earth Transition

Driving Force of Atoms

Full Energy Levels are very stable Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas

configuration.

Atomic Size

First problem where do you start measuring. The electron cloud doesn’t have a definite

edge. They get around this by measuring more than

1 atom at a time.

Atomic Size

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

}Radius

What’s A Trend?

Name some fashion trends Color trends? Behavior trends?

Trends in Atomic Size

Influenced by two factors.Energy Level

Higher energy level is further away.Charge on nucleus

More charge pulls electrons in closer.

Periodic Trends As you go across a period the radius gets

smaller. They have the same energy level, though. More nuclear charge. Outermost electrons are closer.

Na Mg Al Si P S Cl Ar

Group trends

As we go down a group

Each atom has another energy level,

So the atoms get bigger.

HLi

Na

K

Rb

Ionic Size

Cations form by losing electrons. Form positive ions

Groups 1-3 Cations are smaller than the atom they come

from. Metals form cations

Ionic size

Anions form by gaining electrons. Form negative ions

Anions are bigger than the atom they come from.

Nonmetals form anions.

Ionization Energy- Pg. 153

The amount of energy required to remove an electron from an atom (only deals with losing an e’)

Removing one electron makes a +1 ion. The energy required is called the first ionization

energy. Measured in kilojoules per mole

If an atom has a low IE, it will release an electron easier than one with a high IE (making them more reactive) So, which group would have the highest IE?

Nobel gases

Ionization Energy

The second ionization energy is the energy required to remove the second electron.

Always greater than first IE. The third IE is the energy required to remove

a third electron. Greater than 1st of 2nd IE.

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

11810 14840 3569 4619 4577 5301 6045 6276

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

11810 14840 3569 4619 4577 5301 6045 6276

Group trends

As you go down a group first IE decreases because The electron is further away. More shielding occurs

What’s shielding?

Shielding

The electron on the outside energy level has to look through all the other energy levels to see the nucleus

Inner shell electrons “shield” nuclear charge from outer shell electrons

Period trends

All the atoms in the same period have the same energy level.

Same shielding. Increasing nuclear charge So IE generally increases from left to right

A higher charge will more strongly attract electrons, holding them “hostage”

Firs

t Ion

izat

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ener

gy

Atomic number

He

He has a greater IE than H.

same shielding greater nuclear charge H

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li has lower IE than H

more shielding further away outweighs greater

nuclear charge

Li

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE than Li

same shielding greater nuclear

charge

Li

Be

Firs

t Ion

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ion

ener

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Atomic number

H

He B has lower IE than

Be same shielding greater nuclear

charge By removing an

electron we make s orbital half filled Li

Be

B

Firs

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ener

gy

Atomic number

H

He

Li

Be

B

C

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

Firs

t Ion

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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

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

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ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Ne has a lower IE

than He Both are full, Ne has more

shielding Greater distance

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Na has a lower

IE than Li Both are s1

Na has more shielding

Greater distance

Na

Firs

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Atomic number

Electron Affinity- Pg 157

The energy associated with adding an electron to an atom (only deals with gaining e’) Easiest to add to group 7A. A highly negative number = a high EA, that means

the atom will gain electrons easily Trends:

EA increases from left to right because atoms become smaller, with greater nuclear charge.

EA decrease as we go down a group. Also measured in kJ/mol

Electronegativity- Pg. 161

The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares the electron with the atom its bonding

with

Big electronegativity means it pulls the electron toward it- it’s a bully!

Atoms with large negative electron affinity have larger electronegativity.

Flouine is the boss!

Na Cl

Valence Electron

Electronegativity= 0.9Electronegativity= 3.0

So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity

Na Cl

Valence Electron

Electronegativity= 0.9Electronegativity= 3.0

So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity

Group Trend

The further down a group, the farther the electron is away from the nucleus because there are more energy levels Therefore, atoms are more willing to share these

electrons. Low electronegativity.

Periodic Trend

Cations let their electrons go easily Low electronegativity

Anions want more electrons- try to steal them High electronegativity

Ionization energy, electronegativity

Electron affinity INCREASE

Atomic size increases, shielding constant

Ionic size increases