Chapter 6

“The Periodic Table”

The Elements by Tom Lehrer

 Organizing the Elements used

properties

of elements to sort into groups.

 1829 J. W. Dobereiner arranged elements into triads – groups of 3 w/ similar properties  One element in each triad had

properties

intermediate of the other two elements  Cl, Br, and I look different, but similar chemically

Mendeleev’s Periodic Table  mid-1800s, about 70 elements known  Dmitri Mendeleev & teacher – Russian chemist  Arranged elements by increasing

atomic mass

Mendeleev 

blanks

for undiscovered elements  When discovered, his predictions accurate  Problems w/ order  Co to Ni  Ar to K  Te to I

 A better arrangement 1913, Henry Moseley – British physicist, arranged elements according to increasing

atomic number

The Elements by Tom Lehrer

Periodic Law  When elements arranged in order of increasing atomic #,

periodic repetition

of phys & chem props  Horizontal rows = periods  7 periods  Vertical column = group (or family)  Similar phys & chem prop.

 ID’ed by # & letter (IA, IIA)

Areas of periodic table   3 classes of elements:

1) Metals

: electrical conductors, have luster, ductile, malleable    

2) Nonmetals

: generally brittle and non-lustrous, poor conductors of heat and electricity Some gases (O, N, Cl) some brittle solids (B, S) fuming red liquid (Br)



3) Metalloids

: border the line-2 sides  Properties are

intermediate

between metals and nonmetals

Section 6.2

Classifying the Elements  OBJECTIVES: 

Describe the information in a periodic table.

Section 6.2

Classifying the Elements  OBJECTIVES: 

Classify elements based on electron configuration.

Section 6.2

Classifying the Elements  OBJECTIVES: 

Distinguish representative elements and transition metals.

Groups of elements

- family names

 Group IA – alkali metals  Forms “base” (or alkali) when

reacting

w/ H 2 O (not just dissolved!)  Group 2A – alkaline earth metals  Also form bases with H 2 O; don’t dissolve well, hence “earth metals”  Group 7A – halogens  “salt-forming”

Electron Configurations in Groups  Elements sorted

based on e configurations

: 1) Noble gases 2) Representative elements 3) 4) Transition metals Inner transition metals Let’s now take a closer look at these.

Electron Configurations in Groups 1) Noble gases in Group 8A called Group 18)  very stable = don’t react  (also e- configuration w/ outer s & p sublevels

full

Electron Configurations in Groups 2) Representative Elements Groups 1A - 7A   wide range of properties “Representative” of all elements  s & p sublevels of highest energy level NOT filled  Group # equals # of e- in highest energy level

Electron Configurations in Groups 3) Transition metals in “B” columns    outer s sublevel full Start filling “d” sublevel “Transition” btwn metals & nonmetals

Electron Configurations in Groups 4) Inner Transition Metals below main body of PT, in 2 horizontal rows     outer s sublevel full Start filling “f” sublevel Once called “rare-earth” elements not true b/c some abundant

1A

2A  Elements 1A-7A groups called representative elements outer s or p filling 3A 4A 5A 6A 7A 8A

The group B called transition elements 

These are called the inner transition elements, and they belong here

Group 1A called alkali metals (but NOT H) Group 2A called alkaline earth metals

H

 Group 8A are noble gases  Group 7A called halogens

 Let’s take a quick break…… Periodic table rap

H 1 1 s 1 Li 3 1s 2 2 s 1 Na 11 Do you notice any similarity in these configurations of the alkali metals?

1s 2 2s 2 2p 6 3 s 1 K 19 1s 2 2s 2 2p 6 3s 2 3p 6 4 s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5 s 1 Rb 37 Cs 55 1s 2 2s 5p 6 2 6 2p s 1 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Fr 87 1s 2 2s 2 2p 6 3s 2 3p 6 4s s 2 4f 14 5d 10 6p 6 7 s 1 2 3d 10 4p 6 5s 2 4d 10 5p 6 6

Do you notice any similarity in the configurations of the noble gases?

1 s 2 He 2 1s 2 2s 2 2p 6 Ne 10 1s 2 2s 2 2p 6 3s 2 3p 6 Ar 18 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Kr 36 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 Xe 54 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Rn 5p 6 6s 2 4f 14 5d 10 6p 6 86

s 1 Elements in the s - blocks s 2  Alkali metals end in s 1  Alkaline earth metals end in s 2  should include He, but…  He has properties of noble

gases

 has a full outer level of e ’s 

group 8A.

He

Transition Metals - d block Note the change in configuration.

d 1 d 2 d 3 s 1 d 5 d 5 d 6 d 7 d 8 s 1 d 10 d 10

The P-block p 1 p 2 p 3 p 4 p 5 p 6

F - block  Called “inner transition elements” f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14

Period Number

5 6 7 1 2 3 4  Each row (or period) is energy level for s & p orbitals.

 “d” orbitals fill up in levels 1 less than period #  first d is 3d found in period 4.

4d 5d

1 2 3 4 5 6 7 3d

5 6 7 1 2 3 4

4f



5f

f orbitals start filling at 4f….2 less than period #

Demo p. 165

Section 6.3

Periodic Trends  OBJECTIVES:  Describe

trends

among the elements for

atomic size

.

Section 6.3

Periodic Trends  OBJECTIVES:  Explain how

ions

form.

Section 6.3

Periodic Trends  OBJECTIVES:  Describe periodic

trends

for first ionization energy, ionic size, and electronegativity.

Trends in Atomic Size Radius  Measure Atomic Radius - half distance btwn 2 nuclei of diatomic molecule (i.e. O 2 )  Units of picometers (10 -12 m… 1 trillionth)

ALL

Periodic Table Trends  Influenced by 3 factors: 1. Energy Level  Higher energy levels further away from nucleus.

2. Charge on nucleus (# protons)  More charge pulls electrons in closer. (+ and – attract each other)  3. Shielding effect

What do they influence?

 Energy levels & Shielding have effect on

GROUP

(  )  Nuclear charge has effect on

PERIOD

(  )

#1. Atomic Size - Group trends  Going down a group, each atom has another energy level (floor) H Li Na  atoms get

b i g g e r

K Rb

#1. Atomic Size - Period Trends  left to right across period:  size gets s m a l l e r  e ’s occupy same energy level  more nuclear charge  Outer e ’s pulled closer Here is an animation to explain the trend.

Na Mg Al Si P S Cl Ar

Rb Period 2 K Na Li Ar Ne H 3 10 Atomic Number Kr

Trends of Atomic Radius

Ions  Some compounds composed of “ions” 

ion

is atom (or group of atoms) w/ + or -

charge

 Atoms are neutral because the number of protons = electrons  + & - ions formed when e gained) btwn atoms

transferred

(lost or

Ions  Metals LOSE electrons , from outer energy level  Sodium loses 1 e  more p+ (11) than e- (10)  + charge particle formed…“

cation ”

 Na + called “

sodium ion

”

Ions  Nonmetals GAIN one or more electrons  Cl gains 1 e  p+ (17) & e- (18), so charge of -1  Cl

1-

called “chloride ion” 

anions

#2. Trends in Ionization Energy  Ionization energy - energy required to

completely remove e-

(from gaseous atom)  energy required to remove only 1st e called first ionization energy.

Ionization Energy  second ionization energy is E required to remove 2nd e  Always greater than first IE.

 third greater than 1st or 2nd IE.

Symbol First

Table 6.1, p. 173

Second H 1312 He 2731 5247 Li 520 7297 Be 900 1757 B 800 2430 C 1086 2352 N 1402 2857 O 1314 3391 F 1681 3375 Ne 2080 3963 Third 11810 14840 3569 4619 4577 5301 6045 6276

Symbol First H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 Second Third Why did these values 5247 increase

so much

?

7297 11810 1757 14840 2430 3569 2352 4619 2857 4577 3391 5301 3375 6045 3963 6276

What factors determine IE  greater nuclear charge =

greater

IE  Greater distance from nucleus

decreases

IE  Filled & half-filled orbitals have lower energy  Easier to achieve (lower IE)  Shielding effect

Shielding  e ’s in outer energy level “looks through” all other energy levels to see nucleus

Ionization Energy - Group trends  going down group  first IE decreases b/c...

 e- further away from nucleus attraction  more shielding

Ionization Energy - Period trends  Atoms in same period:  same energy level  Same shielding  Increasing nuclear charge  So IE generally increases left - right  Exceptions…full & 1/2 full orbitals

H He  He has greater IE than H.

 Both have same shielding (e- in 1st level)  He = greater nuclear charge Atomic number

H He Li  Li lower IE than H  more shielding  further away  These outweigh greater nuclear charge Atomic number

He H Be  Be higher IE than Li  same shielding  greater nuclear charge Li Atomic number

He H Be Li B  B has lower IE than Be  same shielding  greater nuclear charge  By removing an electron we make s orbital half-filled Atomic number

He H Be C Li B Atomic number

He N H Be C Li B Atomic number

H He Be Li B N C O  Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital Atomic number

He N F H Be C O Li B Atomic number

H He Li Be B N F Ne  Ne has a lower IE than He C O  Both are full,  Ne has more shielding  Greater distance Atomic number

H He Li Be B N F C O Ne  Na has a lower IE than Li  Both are s 1  Na has more shielding Na  Greater distance Atomic number

Atomic number

Trends in Ionization Energy (IE)

Driving Forces 

Full Energy Levels

to remove e require high E  Noble Gases = full orbitals  Atoms want noble gas configuration

2nd Ionization Energy  For elements w/ filled or ½ filled orbital by removing 2 e-, 2nd IE lower than expected.

 True for s 2  Alkaline earth metals form 2+ ions.

3rd IE  Using the same logic s 2 p 1 atoms have an low 3rd IE.

 Atoms in the aluminum family form 3+ ions.

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

Trends in Ionic Size: Cations  Cations form by losing electrons.

 metals  Cations are smaller came from –  they lose electrons than the atom they  they lose an

entire energy level

.

 Cations of representative elements have noble gas configuration

before

them.

Trends in Ionic size: Anions  Anions gain electrons  Anions

bigger

came from – than the atom they  same energy level  greater area the nuclear charge needs to cover  Nonmetals

Configuration of Ions  Ions always have noble gas configurations (full outer level)  Na atom is: 1s 2 2s 2 2p 6 3s 1  Forms a 1+ sodium ion: 1s 2 2s 2 2p 6  Same as Ne

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.

Ion Group trends  Each step down a group is adding an energy level  Ions get bigger going down, b/c of extra energy level Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

Ion Period Trends  Across period  nuclear charge increases  Ions get smaller.



energy level changes

anions and cations.

between Li 1+ B 3+ N 3 O 2 F 1 Be 2+ C 4+

Size of Isoelectronic ions  Iso means “the same”  Isoelectronic ions have the same # of electrons  Al 3+ Mg 2+ Na 1+ Ne F 1 O 2 and N 3  all have 10 electrons  all have the same configuration: 1s 2 2s 2 2p 6 (which is the noble gas: neon)

Size of Isoelectronic ions?

 Positive ions that have more protons would be

smaller

(more protons would pull the same # of electrons in closer) F 1 O 2 N 3 Al 3+

13 12

Na 1+ Ne

11 10

Mg 2+

9 8 7

#3. Trends in Electronegativity  Electronegativity is tendency for atom to attract e ’s when atom in a compound  Sharing e-, but how equally do they share it?

 Element with big electronegativity means it pulls e- towards itself strongly!

Electronegativity Group Trend  Further down a group, farther e- is away from nucleus, plus the more e ’s an atom has  more willing to share  Low electronegativity

Electronegativity Period Trend  Metals let e ’s go easily  low electronegativity  Nonmetals want more electrons  take them away from others  High electronegativity.

Trends in Electronegativity 0

Chemistry Song "Elemental Funkiness" Mark Rosengarten