Periodic Law

The nucleus of the atom exerts forces upon its electrons that make

them have distinctive traits or trends.

What are these trends and how are they inter-related?

The atom - review• What are the three “players” in the inner

workings of an atom?

• Proton positively charged (+1)• Neutron neutrally charged (+0)• Electron negatively charged (-1)

• elementary charge = 1.60217646 × 10-19 coulombs

The atom - review• What is the geometry that atom takes?

Nucleus : protons and neutrons

Electron

Orbital or shell

Mendeleev’s Table• Meyer’s table grouped elements by the

atom’s valence electrons (outer electrons)

• Mendeleev said that not only is valence important but he also grouped elements by atomic mass and known chemical properties

• His claim to fame was that he predicted missing elements and “foretold” their properties.

4

20

40

84

131

1

19

35

80

127

1

23

39

85

133 137

7

88

40

24

9

128

11

79

32

16

115

27

1412

28

118

31

75

122

Period Table

Metals

Non-metals

Metalloids

Properties of MatterMetals

• Most are solid at room temperature

• Good conductor of heat and electric current

• Lustrous (shiny)• Malleable (easily flattened)• Ductile (easily stretched –

as into a wire)

Non-metals

• Most are gases at room temperature

• Poor conductors of heat and electricity

• Dull (as in not reflecting light well)

• Brittle

Metalloids: have properties common to both Metals and Non-metals

Period Table – Family Names

Period Table - Blockss-block

f-block

d-blockp-block

Modern Chart

• Instead of arranging elements on the chart by atomic mass (with special exceptions) the modern chart arranges elements primarily by atomic number (# of protons an element has)

• Which leads us to our first Periodic Trend

How we describe trends

• We express trends as we move along the periodic table

1.Top to Bottom

2.Left to Right

1st Trend – Nuclear Charge“The Master Trend”

• The nucleus holds the atom’s protons and neutrons

• The number of proton an atom has defines that element and it also influences all the other trends we will examine

1st Trend – Nuclear Charge“The Master Trend”

• Does the Nuclear Charge increase or decrease as we move from the top of the chart down?

Increases (follow the atomic number)

• Does the Nuclear Charge increase or decrease as we move from left to right?

Increases (follow the atomic number)

Electron Configuration

• Before we can examine more periodic trends we have to consider the electron

• The electron is actually the “employee” to the nucleus’s “management role”

• Except for nuclear chemistry, the interaction of the electrons between elements is where all of chemistry takes place

Electron Configuration

• Electrons reside in shells (or orbitals)

• Each “block” contains a different number of shells

• 4 types of blocks – s, p, d, f(some people dress funny)

Electron Configuration

• Each block has so many orbitals(s=1, p=3, d=5, f=7)

• Each orbital can hold up to 2 electrons

Block # of Orbitals Max Number of e-

s (ome) 1 (odd #s) 2 (Orb. X 2)

p (eople) 3 6

d (ress) 5 10

f (unny) 7 14

Electron Configuration

• How we label the Electron’s locations follows this pattern:

Row #Shell Letter

Electron # 2s2

Atomic Radius• Defined – the size of the atom (exactly the

same as the radius of a circle)

• Even though we will be using 2-D pictures to demonstrate trends – they are really in 3-D

and they have really “odd” shapes

Atomic Radius – Top to Bottom

• 3-D pictures of orbitals

• http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html

• As we go from Top to Bottom the Atomic Radius increases

Atomic Radius – Left to Right

• Remember: as we go left to right on the periodic chart the number of protons (or nuclear charge) is increasing

• Remember that nuclear charge is a positive charge and electrons have a negative charge

• The effect on the radius of an atom is demonstrated in the following animation

Atomic Radius

As we go left to right on the periodic chart, the effect of increasing nuclear charge decreases the radius of an atom

Ionization Energy• Defined: the amount of energy needed eject an

electron – thus making the atom an ion (or charged atom)

• As we move down the chart it gets easier to eject an electron because electrons are farther away from the nucleus and therefore held in the atom less securely

• As we move right it gets harder to eject an electron because electrons are closer to the nucleus and therefore held on more securely

New Way of Writing Electron Configuration

• Lets do Neon which has 18 electrons

2s1s 2p

↓↑ ↑↑ ↑ ↑↓ ↓↓

3s

↓ ↑

3p

↑↓ ↑ ↑↓ ↓ ↓

Electronegativity• Defined – ability of an atom to attract (or

steal) an electron from another atom or compound

• Nature “wants” to have a:

• Full shell first• Empty shell next• Half Empty third

Electronegativity• With regards to Electronegativity “s” shells

and “p” shells are considered one shell – we call these electrons valence electrons (meaning on the outside) as apposed to core electrons which are previously filled shells or energy levels

• Lets look at Fluorine:2s1s 2p

↓↑ ↑↑ ↑ ↑↓ ↓↓

↓

Core Valance

Full Shell – very happy

Electronegativity• Lets look at Sodium:

2s1s 2p

↓↑ ↑↑ ↑ ↑↓ ↓↓

Core Valance

Empty Shell – very happy

3s

↓ ↑

3p

Nuclear Charge’s Effect on Electronegativity

↑

Going from Oxygen to Fluorine

Orange – Nucleus

Blue – core Electrons

Red – Valance Electrons

Dotted Line – Nuclear influence

Electronegativity get greater as we go left to right

Nuclear Charge’s Effect on Electronegativity

Orange – Nucleus

Blue – core Electrons

Red – Valance Electrons

Dotted Line – Nuclear influence

↑

Going from Neon to Sodium

↑

Electronegativity get less as we go top to bottom

Likely Charges for Families

• Families of Elements have similar electron configurations

• Lets Look at the Alkali Metals:

2sLi [He]

↑

3sNa [Ne]

↑

4sK [Ar]

↑

5sRb [Kr]

↑

6sCs [Xe]

↑

7sFr [Rn]

↑

Likely Charges for Families

• Remember that the electronegativity for these elements is very low (electrons are easily taken by other elements or left behind)

• If these elements “lose” their valence (or outer) electrons easily, what charge do you think these elements might become?

Likely Charges for Families

• Sodium: 11 protons (+11) &

2s1s 2p

↓↑ ↑↑ ↑ ↑↓ ↓↓

Core Valance

3s

↓ ↑

3p

electrons

Total Charge:

11 (-11)10 (-10)

0+1

When Alkalai Metals become ions (a charged element) it will become +1 charged

Likely Charges for Families

• Lets Look at the Halogens:

F [He]2s 2p

↓ ↑↑ ↑ ↑↓↓

Cl [Ne]3s 3p

↓ ↑↑ ↑ ↑↓↓

Br [Ar]4s 4p

↓ ↑↑ ↑ ↑↓↓

I [Kr]5s 5p

↓ ↑↑ ↑ ↑↓↓

Likely Charges for Families

• Chlorine: 17 protons (+17) &

2s1s 2p

↓↑ ↑↑ ↑ ↑↓ ↓↓

Core Valance

3s

↓

3p

electrons

Total Charge:

17 (-17)18 (-18)

0-1

When Halogens become ions itwill become -1 charged

↑ ↓ ↑ ↓ ↑ ↓ ↑

↓

Likely Charges for Families

+1 +2 +3±4

-3-2

-10