Chapter 6 “The Periodic Table” The Elements by Tom Lehrer The Elements by Tom Lehrer.

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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 propertiesOne element in each triad had properties intermediate of the other two elementsCl, Br, and I look different, but similar chemically

Mendeleev’s Periodic Table mid-1800s, about 70 elements

known Dmitri Mendeleev – Russian chemist

& teacher Arranged elements by

increasing atomic mass

Mendeleevblanks for undiscovered

elementsWhen discovered, his predictions

accurate

Problems w/ orderCo to NiAr to KTe to I

A better arrangement1913, 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 = periods7 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 6.2Classifying the ElementsOBJECTIVES:

Describe the information in a periodic table.

Classify elements based on electron configuration.

Distinguish representative elements and transition metals.

Groups of elements - family names

Group IA – alkali metalsForms “base” (or alkali) when

reacting w/ H2O (not just dissolved!)

Group 2A – alkaline earth metalsAlso form bases with H2O; don’t

dissolve well, hence “earth metals”Group 7A – halogens

“salt-forming”

Electron Configurations in Groups

Elements sorted based on e- configurations:

1) Noble gases2) Representative elements

3) Transition metals

4) Inner transition metals

Let’s now take a closer look at these.

1) Noble gases in Group 8A (also called Group 18)

very stable = don’t react e- configuration w/ outer s & p

sublevels fullfull

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

3) Transition metals in “B” columns outer s sublevel full Start filling “d” sublevel “Transition” btwn metals &

nonmetals

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

2A 3A 4A 5A 6A7A

8A Elements 1A-7A groups called

representative elements

outer s or p filling

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

Group 8A are noble gases Group 7A called halogens

Periodic table rap

Let’s take a quick break……

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10

5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

Do you notice any similarity in these configurations of the alkali metals?

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

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

Elements in the s - blocks

Alkali metals end in s1

Alkaline earth metals end in s2

should include He, but… He has properties of noble gases has a full outer level of e-’s

group 8A.

Transition Metals - d block

d1 d2 d3s1

d5 d5 d6 d7 d8s1

d10 d10

Note the change in configuration.

The P-blockp1 p2 p3 p4 p5 p6

F - block

Called “inner transition elements”

f1 f5f2 f3 f4

f6 f7 f8 f9 f10 f11 f12 f14

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

Period Number

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

4d4d5d5d

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

7 4f4f

Demo p. 165

Section 6.3Periodic Trends

OBJECTIVES:

Describe trends among the elements for atomic size.

OBJECTIVES:

Explain how ions form.

OBJECTIVES:

Describe periodic trends for first ionization energy, ionic size, and electronegativity.

Trends in Atomic Size

Measure Atomic Radius - half distance btwn 2 nuclei of diatomic molecule (i.e. O2)

Units of picometers (10-12 m… 1 trillionth)

}Radius

ALLALL Periodic Table Trends Influenced by 3 factors:

1. Energy LevelHigher 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)

atoms get

bbiiggggeerr

#1. Atomic Size - Period Trends left to right across period:

size gets smaller

e-’s occupy same energy levelmore nuclear chargeOuter e-’s pulled closer

Na Mg Al Si P S Cl Ar

Here is an animation to explain the trend.

Atomic Number

Period 2

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- transferred (lost or

gained) btwn atoms

IonsMetals LOSE electrons, from outer

energy level Sodium loses 1 e-

more p+ (11) than e- (10)+ charge particle formed…“cation”

Na+ called “sodium ion”

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

Cl1- 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 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

Table 6.1, p. 173

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

Why did these values increase so much?

What factors determine IE greater nuclear charge = greater IE Greater distance from nucleus

decreases IE Filled & half-filled orbitals have lower

energyEasier to achieve (lower IE)

Shielding effect

Shieldinge-’s in outer

energy level “looks through” all other energy levels to see nucleus

Ionization Energy - Group trendsgoing 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 levelSame shielding Increasing nuclear charge So IE generally increases left - right

Exceptions…full & 1/2 full orbitals

Atomic number

He He has greater IE than H.

Both have same shielding (e- in 1st level) He = greater nuclear

charge

Atomic number

Li lower IE than H more shielding further away These outweigh

greater nuclear charge

Atomic number

Be higher IE than Li

same shielding greater nuclear

charge

Atomic number

He B has lower IE

than Be same shielding greater nuclear

charge By removing an

electron we make s orbital half-filled

Atomic number

Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital

Atomic number

Ne Ne has a lower

IE than He Both are full, Ne has more

shielding Greater

distance

Atomic number

Ne Na has a lower

IE than Li Both are s1

Na has more shielding

Greater distance

Atomic number

Trends in Ionization Energy (IE)

Driving Forces

Full Energy Levels require high E to remove 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 s2

Alkaline earth metals form 2+ ions.

3rd IE

Using the same logic s2p1 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 than the atom they

came from – they lose electrons 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 than the atom they came from – same energy levelgreater area the nuclear charge needs to

cover Nonmetals

Configuration of Ions Ions always have noble gas

configurations (full outer level) Na atom is: 1s22s22p63s1 Forms a 1+ sodium ion: 1s22s22p6 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

Ion Period Trends Across period

nuclear charge increases Ions get smaller.

energy level changes between anions and cations.

N3-O2- F1-

Size of Isoelectronic ions Iso- means “the same” Isoelectronic ions have the same # of

electrons Al3+ Mg2+ Na1+ Ne F1- O2- and N3-

all have 10 electrons all have the same configuration:

1s22s22p6 (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)

Na1+ Ne F1- O2- N3-

13 12 11 10 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 TrendFurther down a group,

farther e- is away from nucleus, plus the more e-’s an atom has

more willing to shareLow electronegativity

Electronegativity Period Trend Metals let e-’s go easily

low electronegativity

Nonmetals want more electronstake them away from othersHigh electronegativity.

Trends in Electronegativity

Chemistry Song "Elemental Funkiness" - Mark Rosengarten

Chapter 6 “The Periodic Table” The Elements by Tom Lehrer The Elements by Tom Lehrer.

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