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Chapter 6 The Periodic Table. The how and why. History. 1829 German J. W. Dobereiner Grouped elements into triads Three elements with similar properties Properties followed a pattern The same element was in the middle of all trends Not all elements had triads. History.

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history
History
  • 1829 German J. W. Dobereiner Grouped elements into triads
    • Three elements with similar properties
    • Properties followed a pattern
    • The same element was in the middle of all trends
  • Not all elements had triads
history3
History
  • Russian scientist Dmitri Mendeleev taught chemistry in terms of properties
  • Mid 1800 – atomic masses of elements were known
  • Wrote down the elements in order of increasing mass
  • Found a pattern of repeating properties
mendeleev s 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.
  • Found some gaps
  • Must be undiscovered elements
  • Predicted their properties before they were found
the modern table
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.
slide7
Vertical columns are called groups.
  • Elements are placed in columns by similar properties.
  • Also called families
slide8

8A0

1A

  • The elements in the A groups are called the representative elements

2A

3A

4A

5A

6A

7A

slide9

VIIIB

IA

IIA

VIB

VIIB

IIIB

IVB

VB

1

1A

2

2A

8A

18

13

3A

14

4A

15

5A

16

6A

7A

17

VIIIA

VIA

VIIA

IIIA

IVA

VA

IB

IIB

3

3B

4B

4

5

5B

6B

6

7

7B

8

8B

9

8B

10

8B

1B

11

2B

12

Other Systems

metals11
Metals
  • Luster – shiny.
  • Ductile – drawn into wires.
  • Malleable – hammered into sheets.
  • Conductors of heat and electricity.
transition metals
Transition metals
  • The Group B elements
non metals
Non-metals
  • Dull
  • Brittle
  • Nonconductors- insulators
metalloids or semimetals
Metalloids or Semimetals
  • Properties of both
  • Semiconductors
slide17
Group 1A are the alkali metals
  • Group 2A are the alkaline earth metals
slide18
Group 7A is called the Halogens
  • Group 8A are the noble gases
slide19
Why?
  • The part of the atom another atom sees is the electron cloud.
  • More importantly the outside orbitals
  • The orbitals fill up in a regular pattern
  • The outside orbital electron configuration repeats
  • So.. the properties of atoms repeat.
slide20

H

1

Li

3

Na

11

K

19

Rb

37

Cs

55

Fr

87

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p63d104s24p65s1

1s22s22p63s23p63d104s24p64d105s2 5p66s1

1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p67s1

slide21

He

2

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p63d104s24p6

1s22s22p63s23p63d104s24p64d105s25p6

1s22s22p63s23p63d104s24p64d105s24f14 5p65d106s26p6

Ne

10

Ar

18

Kr

36

Xe

54

Rn

86

s block
S- block

s1

  • Alkali metals all end in s1
  • Alkaline earth metals all end in s2
  • really have to include He but it fits better later
  • He has the properties of the noble gases

s2

transition metals d block
Transition Metals -d block

s1 d5

s1 d10

d1

d2

d3

d5

d6

d7

d8

d10

the p block
The P-block

p1

p2

p6

p3

p4

p5

f block

f6

f13

f1

f2

f3

f4

f5

f7

f8

f10

f12

f14

f11

f9

F - block
  • inner transition elements
slide26

1

2

3

4

5

6

7

  • Each row (or period) is the energy level for s and p orbitals
slide27
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

slide28

1

2

3

4

5

6

7

  • f orbitals start filling at 4f

4f

5f

electron configurations repeat
Electron Configurations repeat
  • The shape of the periodic table is a representation of this repetition.
  • When we get to the end of the row the outermost energy level is full.
  • This is the basis for our shorthand
the shorthand
The Shorthand
  • Write the symbol of the noble gas before the element in brackets [ ]
  • Then the rest of the electrons.
  • Aluminum - full configuration
  • 1s22s22p63s23p1
  • Ne is 1s22s22p6
  • so Al is [Ne] 3s23p1
more examples
More examples
  • Ge = 1s22s22p63s23p63d104s24p2
  • Ge = [Ar] 4s23d104p2
  • Ge = [Ar] 3d104s24p2
  • Hf=1s22s22p63s23p64s23d104p64f14 4d105s25p65d26s2
  • Hf=[Xe]6s24f145d2
  • Hf=[Xe]4f145d26s2
the shorthand33
The Shorthand

Sn- 50 electrons

The noble gas before it is Kr

Takes care of 36

Next 5s2

Then 4d10

Finally 5p2

[ Kr ]

5s2

4d10

5p2

practice
Practice
  • Write the shorthand configuration for
  • S
  • Mn
  • Mo
  • W
electron configurations and groups
Electron configurations and groups
  • Representative elements have s and p orbitals as last filled
    • Group number = number of electrons in highest energy level
  • Transition metals- d orbitals
  • Inner transition- f orbitals
  • Noble gases s and p orbitals full
part 3 periodic trends

Part 3Periodic trends

Identifying the patterns

what we will investigate
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
  • Ionic size
    • How big ions are
what we will look for
What we will look for
  • Periodic trends-
    • How those 4 things vary as you go across a period
  • Group trends
    • How those 4 things vary as you go down a group
  • Why?
    • Explain why they vary
the why first
The why first
  • The positive nucleus pulls on electrons
  • Periodic trends – as you go across a period
    • The charge on the nucleus gets bigger
    • The outermost electrons are in the same energy level
    • So the outermost electrons are pulled stronger
the why first40
The why first
  • The positive nucleus pulls on electrons
  • Group Trends
    • As you go down a group
      • You add energy levels
      • Outermost electrons not as attracted by the nucleus
slide41

Shielding

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

+

shielding
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

+

shielding43
Shielding
  • As the energy levels changes the shielding changes
  • Lower down the group
    • More energy levels
    • More shielding
    • Outer electron less attracted

+

Three shields

No shielding

One shield

Two shields

atomic size
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 size45
Atomic Size

}

Radius

  • Atomic Radius = half the distance between two nuclei of molecule
trends in atomic size
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
group trends
Group trends

H

  • As we go down a group
  • Each atom has another energy level
  • More shielding
  • So the atoms get bigger

Li

Na

K

Rb

periodic trends
Periodic Trends
  • As you go across a period the radius gets smaller.
  • Same shielding and energy level
  • More nuclear charge
  • Pulls outermost electrons closer

Na

Mg

Al

Si

P

S

Cl

Ar

overall

Rb

Overall

K

Na

Li

Atomic Radius (nm)

Kr

Ar

Ne

H

10

Atomic Number

ionization energy
Ionization Energy
  • The amount of energy required to completely remove an electron from a gaseous atom.
  • Removing one electron makes a +1 ion
  • The energy required is called the first ionization energy
ionization energy52
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 or 2nd IE
slide53

Symbol First Second Third

1181014840 3569 4619 4577 5301 6045 6276

5247 7297 1757 2430 2352 2857 3391 3375 3963

1312 2731 520 900 800 1086 1402 1314 1681 2080

HHeLiBeBCNO F Ne

what determines ie
What determines IE
  • The greater the nuclear charge the greater IE.
  • Increased shielding decreases IE
  • Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE
group trends55
Group trends
  • As you go down a group first IE decreases because of
  • More shielding
  • So outer electron less attracted
periodic trends56
Periodic trends
  • All the atoms in the same period
    • Same shielding.
    • Increasing nuclear charge
  • So IE generally increases from left to right.
  • Exceptions at full and 1/2 full orbitals
slide57

He

  • He has a greater IE than H
  • same shielding
  • greater nuclear charge

H

First Ionization energy

Atomic number

slide58

He

  • Li has lower IE than H
  • more shielding
  • outweighs greater nuclear charge

H

First Ionization energy

Li

Atomic number

slide59

He

  • Be has higher IE than Li
  • same shielding
  • greater nuclear charge

H

First Ionization energy

Be

Li

Atomic number

slide60

He

  • B has lower IE than Be
  • same shielding
  • greater nuclear charge
  • By removing an electron we make s orbital full

H

First Ionization energy

Be

B

Li

Atomic number

slide61

He

C

H

First Ionization energy

Be

B

Li

Atomic number

slide62

He

N

C

H

First Ionization energy

Be

B

Li

Atomic number

slide63

He

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

N

O

C

H

First Ionization energy

Be

B

Li

Atomic number

slide64

He

F

N

O

C

H

First Ionization energy

Be

B

Li

Atomic number

slide65

Ne

He

F

  • Ne has a lower IE than He
  • Both are full,
  • Ne has more shielding

N

O

C

H

First Ionization energy

Be

B

Li

Atomic number

slide66

Ne

He

  • Na has a lower IE than Li
  • Both are s1
  • Na has more shielding

F

N

O

C

H

First Ionization energy

Be

B

Li

Na

Atomic number

slide67

Web elements

First Ionization energy

Atomic number

driving force
Driving Force
  • Full Energy Levels are very low energy
  • Noble Gases have full orbitals
  • Atoms behave in ways to achieve noble gas configuration
2nd ionization energy
2nd Ionization Energy
  • For elements that reach a filled or half-full orbital by removing 2 electrons 2nd IE is lower than expected
  • True for s2
  • Alkali earth metals form 2+ ions
3rd ie
3rd IE
  • Using the same logic s2p1atoms have an low 3rd IE
  • Atoms in the boron family form 3+ ions
  • 2nd IE and 3rd IE are always higher than 1st IE!!!
slide71

Web elements

Symbol First Second Third

1181014840 3569 4619 4577 5301 6045 6276

5247 7297 1757 2430 2352 2857 3391 3375 3963

1312 2731 520 900 800 1086 1402 1314 1681 2080

HHeLiBeBCNO F Ne

ionic size
Ionic Size
  • Cations are positive ions
  • Cations form by losing electrons
  • Cations are smaller than the atom they come from
  • Metals form cations
  • Cations of representative elements have noble gas configuration.
ionic size73
Ionic size
  • Anions are negative ions
  • 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.
configuration of ions
Configuration of Ions
  • Ions of representative elements have noble gas configuration
  • Na is 1s22s22p63s1
  • Forms a 1+ ion - 1s22s22p6
  • Same configuration as neon
  • Metals form ions with the configuration of the noble gas before them - they lose electrons
configuration of ions75
Configuration of Ions
  • Non-metals form ions by gaining electrons to achieve noble gas configuration.
  • They end up with the configuration of the noble gas after them.
group trends76
Group trends
  • Adding energy level
  • Ions get bigger as you go down

H1+

Li1+

Na1+

K1+

Rb1+

Cs1+

periodic trends77
Periodic Trends
  • Across the period nuclear charge increases so they get smaller.
  • Energy level changes between anions and cations

N3-

O2-

F1-

B3+

Li1+

C4+

Be2+

size of isoelectronic ions
Size of Isoelectronic ions
  • Iso - same
  • Iso electronic ions have the same # of electrons
  • Al3+ Mg2+ Na1+ Ne F1- O2- and N3-
  • all have 10 electrons
  • all have the configuration 1s22s22p6
size of isoelectronic ions79
Size of Isoelectronic ions
  • Positive ions have more protons so they are smaller

N-3

O-2

F-1

Ne

Na+1

Al+3

Mg+2

electronegativity81
Electronegativity
  • The tendency for an atom to attract electrons to itself when it is chemically combined with another element.
  • How “greedy”
  • Big electronegativity means it pulls the electron toward it.
group trend
Group Trend
  • The further down a group
    • More shielding
    • more electrons an atom has.
  • Less attraction for electrons
  • Low electronegativity.
periodic trend
Periodic Trend
  • Metals - left end
  • Low nuclear charge
  • Low attraction
  • Low electronegativity
  • Right end - nonmetals
  • High nuclear charge
  • Large attraction
  • High electronegativity
  • Not noble gases- no compounds
slide85

Atomic size increases,

Ionic size increases

slide86

Nuclear Charge

Energy Levels

& Shielding

how to answer why questions
How to answer why questions
  • Trend
    • Periodic
    • Group
  • Reason
    • Nuclear charge
    • Energy level and shielding
  • Result
    • What happens to which electron
trend
Trend

across a period

As you go

down a group

Reason

nuclear charge

the

increases

energy level and shielding

result
Result

electron ______ to remove

pull _____ on the other atom’s electron

Making the

outer electrons are _______.

so the ______ is _______.