unit 5 the periodic table the how and why. newlands -1865 u arranged known elements according to...

Unit 5The Periodic Table

The how and whyThe how and why

Newlands -1865 Arranged known elements

according to properties & order of increasing atomic mass

Law of Octaves – pattern of chemical & physical properties repeated every 8 elements

Mendeleev - 1869 Created 1st periodic table (63 elements) Ordered by increasing atomic mass Predicted pattern of missing elements Started new rows and lined up

columns to organize elements with similar properties

Rearranged elements so similar properties would line up correctly

The Modern Table Moseley- determined the atomic

number for each known element. Elements are still grouped by properties Similar properties are in the same

column Ordered by increasing atomic number Added a column of elements Mendeleev

didn’t know about – noble gases

Periodic Law When elements are arranged in

order of increasing atomic number, elements with similar properties appear at regular intervals

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 6A 7A

8A

3B 4B 5B 6B 7B 8B 8B 8B 1B 2B

1 2

13 14 15 16 17

18

3 4 5 6 7 8 9 10 11 12

IA IIA

IIIB

IVB

VB

VIB

VII

B

VII

IB

IIIA

IVA

VA

VIA

VII

A

VII

IA

IB IIB

Other Systems

1A

2A 3A 4A 5A 6A7A

8A0

The elements in the A groups are called the representative elements

Transition metals The Group B

elements

These are called the inner transition elements and they belong here

Three Classes of ElementsMetalsNonmetalsMetalloids

Metals

Metals Ductile – drawn into wires Malleable – hammered into sheets All solid at room temperature (except

Hg- Mercury) Conductors of heat and electricity Families

– 1 - Alkali– 2 - Alkaline Earth– Transition (B groups)

Group 1A are the alkali metals VERY reactive because one valence e-

• Found as compounds in nature• Not including H!

Group 2A are the alkaline earth metals Still highly reactive but not as much so

as alkali metals (2 valence e-)

Transition Metals The weird ones… May lose different #s of valence

electrons depending on the element with which it reacts

Less reactive than alkali or alkaline earth metals

Good conductors of electricity & heat, ductile, malleable

Inner Transition Metals 1st row = lanthanides

• Shiny metals similar in reactivity to alkaline earth metals

2nd row = actinides

• Unstable nuclei – all radioactive

Non-metals

Non-metals Most are gases, some solid, and 1

liquid (Br) More variation than metals Families

–Halogens (Group 17 or 7A)–Noble Gases (Group 18 or 8A)

Group 7A is called the Halogens Most reactive non-metals – 7 valence

React frequently with alkali metals

Group 8A are the noble gases

Low reactivity, very stable, inert

Metalloids or Semimetals

Metalloids Border the staircase between

metals and nonmetals Properties – similar to metals and

nonmetals

Part 2Periodic trends

Identifying the patterns

What we will investigate Atomic size

• how big the atoms are Ionization energy

• How much energy to remove an electron

Electronegativity

• The attraction for the electron in a compound

What we will look for Periodic trends

• How those things vary as you go across a period

Group trends

• How those things vary as you go down a group

Why?

• Explain why these variations exist

Atomic Size Where do you start measuring? The electron cloud doesn’t have a

definite edge. Scientists focused first on diatomic

elements -- measured more than 1 atom at a time

Atomic Size

Atomic Radius = half the distance between two nuclei of molecule

}Radius

Atomic Size - Periodic Trends The positive nucleus pulls on electrons Periodic trend

• As you move across a period, elements have more protons

• The charge on the nucleus gets bigger

• The outermost electrons of each element are in the same energy level

• So there is more pull on the outermost electrons as you move across

Periodic Trends As you go across a period, the radius

gets smaller. Same outermost energy level More nuclear charge Pulls outermost electrons closer

Na Mg Al Si P S Cl Ar

Atomic Size – Group Trends The positive nucleus pulls on electrons Group Trend

• As you go down a group, you add energy levels

• Outermost electrons not as attracted by the nucleus

+

Shielding Increasing numbers of

electrons between the nucleus and the valence electrons tends to decrease the force between the nucleus & the valence electrons

+

Shielding The electron on the

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

Shielding The electron on the

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

A second electron has the same shielding

In the same energy level (period) shielding is the same

+

Shielding As the energy levels

changes the shielding changes

Moving down the group

• More energy levels

• More shielding

• Outer electron less attracted

+

No shieldingOne shieldTwo shieldsThree shields

Group trends As we go down a

group

• Each atom has another energy level

• More shielding

• The atoms get bigger

HLi

Na

K

Rb

Overall

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

Atomic size increases,

IONIZATION ENERGY

It’s all about stability Alkali metals are more stable if

they lose an electron Example

• Sodium ([Ne] 3s1)

• Getting rid of the 3s1 electron makes sodium more stable and creates a sodium ion (Na1+)

Ionization Energy The amount of energy

required to completely remove an electron from a neutral atom.

The energy required for the 1st electron is called the first ionization energy

Ionization Energy The 2nd ionization energy is the

energy required to remove the second electron

Always greater than 1st IE The 3rd IE is the energy required to

remove a third electron Greater than 1st or 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

1181014840 3569 4619 4577 5301 6045 6276

Group trends As you go down a group first IE

decreases

• Valence e- farther from nucleus

• More shielding

Periodic trends First IE increases from left to right

across a period

• Increased nuclear charge from added proton

• Electron shielding not an issue b/c valence are all in same energy level

Exceptions at full and 1/2 full orbitals

• Lower IE b/c offer stability to atom

Ionization energy

INCREASE

How to remember?HILO

LO

Firs

t Ion

izat

ion

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 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

izat

ion

ener

gy

Atomic number

H

He B has lower IE than Be same shielding greater nuclear charge By removing an

electron we make s orbital full

Li

Be

B

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

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

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

izat

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

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

Na

Firs

t Ion

izat

ion

ener

gy

Atomic number

Electronegativity

Electronegativity There’s an electron tug of war

between atoms in a compound The tendency for an atom to attract

electrons to itself when it is chemically combined with another element

How “greedy” Large electronegativity means the

atom pulls the electron towards itself

Group Trend As you move down a group

• More shielding

• Less attraction for electrons

• Lower electronegativity

Periodic Trend As you move across a period

from left to right,

• Nuclear charge increases

• Greater electronegativity

Electronegativity

INCREASE

How to remember?HILO

LO

All 3 trendsHILO

LO