3.1 org, clssify and trend periodic

8/10/2019 3.1 Org, Clssify and Trend Periodic

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Chapter

The Periodic Table


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Section 1

Organizing the Elements

OBJECTIVES:

Explain how elements are

organized in a periodic table.


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Section 1


OBJECTIVES:

Compare early and modern

periodic tables.


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Section 1


OBJECTIVES:

Identify three broad classes

of elements.


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Section 1

Organizing the ElementsA few elements, such as gold and

copper, have been known for thousands

of years- since ancient timesYet, only about 13 had been identified

by the year 1700.

As more were discovered, chemistsrealized they needed a way to organize

the elements.


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Section 1

Organizing the ElementsChemists used the propert iesof

elements to sort them into groups.

In 1829 J. W. Dobereiner arrangedelements into triadsgroups of threeelements with similar properties

One element in each triad hadpropertiesintermediate of the other twoelements


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Mendeleevs Periodic Table

By the mid-1800s, about 70 elementswere known to exist

Dmitri Mendeleeva Russian chemist

and teacherArranged elements in order of increasing

atomic mass

Thus, the first Periodic Table


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Mendeleev

He left b lanksfor yetundiscovered elements

When they were discovered,he had made good predictions

But, there were problems:

Such as Co and Ni; Ar and

K; Te and I


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A better arrangement

In 1913, Henry MoseleyBritish physicist,arranged elements according to increasingatomic number

The arrangement used today

The symbol, atomic number & mass arebasic items included-textbook page 162 and163


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Another possibility:

Spiral Periodic Table


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The Periodic Law says:

When elements are arranged inorder of increasing atomic number,there is aperiodic repetitionof their

physical and chemical properties.Horizontal rows = periods

There are 7 periods

Vertical column = group (or family) Similar physical & chemical prop.

Identified by number & letter (IA, IIA)


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Areas of the periodic table

Three classes of elements are:1) metals, 2) nonmetals, and

3) metalloids

1) Metals: electrical conductors, have

luster, ductile, malleable

2) Nonmetals: generally brittle andnon-lustrous, poor conductors of

heat and electricity


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Areas of the periodic table

Some nonmetals are gases (O, N,

Cl); some are brittle solids (S); one

is a fuming dark red liquid (Br) Notice the heavy, stair-step line?

3) Metalloids: border the line-2 sides

Properties are intermediate

between metals and nonmetals


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Section 2

Classifying the ElementsOBJECTIVES:

Describe the informationin a periodic table.


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Section 2


Classify elements basedon electron

configuration.


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Section 2


Distinguishrepresentative elements

and transition metals.


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Squares in the Periodic Table

The periodic table displays thesymbols and names of the

elements, along with

information about the structure

of their atoms:

Atomic number and atomic mass Black symbol = solid; red = gas;

blue = liquid

(from the Periodic Table on our classroom wall)


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Groups of elements - family names

Group IAalkali metals Forms a base (or alkali) when reacting

with water (not just dissolved!)

Group 2Aalkaline earth metalsAlso form bases with water; do not

dissolve well, hence earth metals

Group 7Ahalogens Means salt-forming


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Electron Configurations in Groups

Elements can be sorted into 4different groupings based on

their electron configurations:1) Noble gases

2) Representative elements

3) Transition metals

4) Inner transition metals

Lets

now

take a

closer

look at

these.


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1) Noble gasesare the elementsin Group 8A (also called Group18 or 0)

Previously called inert gasesbecause they rarely take part in a

reaction; very stable = dont react

Noble gases have an electronconfiguration that has the outer s

and p sublevels completely full


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2) Representative Elementsarein Groups 1A through 7A

Display wide range of properties,

thus a good representative

Some are metals, or nonmetals,or metalloids; some are solid,

others are gases or liquids Their outer s and p electron

configurations are NOT filled


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4) Inner Transition Metalsarelocated below the main body ofthe table, in two horizontal rows

Electron configuration has theouter s sublevel full, and is nowfilling the f sublevel

Formerly called rare-earthelements, but this is not truebecause some are very abundant


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1A

2A 3A 4A 5A 6A7A

8AElements in the 1A-7A groups

are called the representative

elements

outer s or p filling


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The group B are called thetransition elements

These are called the inner

transition elements, and they

belong here


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Group 1A are the alkali metals (but NOT H)

Group 2A are the alkaline earth metalsH


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Group 8A are the noble gases

Group 7A is called the halogens


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1s1

1s22s1

1s22s22p63s1

1s

2

2s

2

2p

6

3s

2

3p

6

4s

1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10

5p66s1

1s22s22p63s23p64s23d104p65s24d105p66

s24f145d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

Do you notice any similarity in these

configurations of the alkali metals?


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He2

Ne10

Ar18

Kr36

Xe54

Rn86

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p6

5s2

4d10

5p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

Do you notice any similarity in the

configurations of the noble gases?


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Alkali metals all end in s1

Alkaline earth metals all end in s2

really should include He, but it fits

better in a different spot, since He

has the properties of the noble

gases, and has a full outer level

of electrons.

s2s1

Elements in the s - blocksHe


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Transition Metals - d block

d1 d2 d3s1

d5 d5 d6 d7 d8s1

d10 d10

Note the change in configuration.


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The P-blockp1 p2 p3 p4

p

5 p6


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F - block

Called the inner transition elements

f1 f5f2 f3 f4 f6 f7 f8 f9 f10 f11f12 f14f13


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Each row (or period) is the energylevel for s and p orbitals.

1

2

3

4

5

6

7

Period

Number


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The d orbitals fill up in levels 1 lessthan the period number, so the first dis 3d even though its in row 4.

1

2

3

4

5

6

7

3d

4d5d


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f orbitals start filling at 4f, and are 2less than the period number

1

23

4

5

6

7 4f

5f


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Section 3

Periodic Trends

OBJECTIVES:

Describe trendsamong the

elements for atomic size.


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Section 3

Periodic Trends

OBJECTIVES:

Explain how ionsform.


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Section 3

Periodic Trends

OBJECTIVES:

Describe periodic trendsfor

first ionization energy, ionic

size, and electronegativity.


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Atomic Size

Measure the Atomic Radius - this is half thedistance between the two nuclei of a diatomicmolecule.

}

Radius


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ALLPeriodic Table Trends

Influenced by three factors:1. Energy Level

Higher energy levels are further

away from the nucleus.

2. Charge on nucleus (# protons)

More charge pulls electrons incloser. (+ andattract each other)

3. Shielding effect (blocking effect?)


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What do they influence?

Energy levels and Shielding have

an effect on the GROUP( )

Nuclear charge has an effect on a

PERIOD ( )


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#1. Atomic Size - Period Trends

Going from left to right across a period,the size getssmaller.

Electrons are in the same energy level.

But, there is more nuclear charge.Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

Rb


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Atomic Number

Ato

micRadius(pm)

H

Li

Ne

Ar

10

Na

K

Kr

3

Period 2


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Ions

Some compounds are composed ofparticles called ions

An ionis an atom (or group of atoms)

that has apositive or negative chargeAtomsare neutral because the number

of protons equals electrons

Positive and negative ions are formedwhen electrons are transferred(lost or

gained) between atoms


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IonsMetals tend to LOSE electrons,

from their outer energy level

Sodium loses one: there are now

more protons (11) than electrons(10), and thus a positively charged

particle is formed = cation

The charge is written as a numberfollowed by a plus sign: Na1+

Now named a sodium ion


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#2. Trends in Ionization Energy

Ionization energy is the amount ofenergy required to complete lyremove an electron(from a gaseousatom).

Removing one electron makes a 1+ion.

The energy required to remove only

the first electron is called the firstionization energy.


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Ionization Energy

The second ionization energy isthe energy required to remove

the second electron.

Always greater than first IE.

The third IE is the energy

required to remove a thirdelectron.

Greater than 1st or 2nd IE.


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Symbol First Second Third

HHe

Li

Be

B

C

N

OF

Ne

13122731

520

900

800

1086

1402

13141681

2080

5247

7297

1757

24302352

2857

33913375

3963

11810

14840

35694619

4577

53016045

6276


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Symbol First Second Third

HHe

Li

Be

B

C

N

OF

Ne

13122731

520

900

800

1086

1402

13141681

2080

5247

7297

1757

24302352

2857

33913375

3963

11810

14840

35694619

4577

53016045

6276

Why did these values

increase so much?


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What factors determine IE

The greater the nuclear charge,the greaterIE.

Greater distance from nucleus

decreasesIE

Filled and half-filled orbitals have

lower energy, so achieving themis easier, lower IE.

Shielding effect


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Shielding

The electron on theoutermost energy

level has to look

through all the otherenergy levels to see

the nucleus.

Second electron hassame shielding, if it

is in the same period


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Ionization Energy - Group trends

As you go down a group, thefirst IE decreases because...

The electron is further away fromthe attraction of the nucleus, and

There is more shielding.


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Ionization Energy - Period trends

All the atoms in the same periodhave the same energy level.

Same shielding.

But, increasing nuclear charge

So IE generally increases from left to

right.

Exceptions at full and 1/2 fullorbitals.


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FirstIoniza

tionener

gy

Atomic number

HeHe has a greater IE

than H.Both elements have

the same shielding

since electrons areonly in the first level

But He has a greater

nuclear charge

H


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FirstIoniza

tionener

gy

Atomic number

H

He

Li has lower IE

than H

more shielding

further away These outweigh

the greater

nuclear chargeLi


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FirstIoniza

tionener

gy

Atomic number

H

He

Be has higher IE

than Li

same shielding

greater nuclearcharge

Li

Be


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FirstIoniza

tionener

gy

Atomic number

H

HeB has lower IE

than Be same shielding

greater nuclear

charge

By removing an

electron we makes orbital half-filled

Li

Be

B


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FirstIoniza

tionener

gy

Atomic number

H

He

Li

Be

B

C


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FirstIoniza

tionener

gy

Atomic number

H

He

Li

Be

B

C

N


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FirstIoniza

tionenergy

Atomic number

H

He

Li

Be

B

C

N

O

F


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FirstIoniza

tionenergy

Atomic number

H

He

Li

Be

B

C

N

O

F

NeNe has a lower

IE than He

Both are full,

Ne has moreshielding

Greater

distance


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FirstIoniza

tionenergy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne

Na has a lower

IE than Li

Both are s1

Na has moreshielding

Greater

distanceNa


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FirstIoniza

tionenergy

Atomic number


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Driving Forces

Ful l Energy Levelsrequirelots of energy to remove theirelectrons.

Noble Gases have fullorbitals.

Atoms behave in ways to tryand achieve a noble gasconfiguration.


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2nd Ionization Energy

For elements that reach afilled or half-filled orbital by

removing 2 electrons, 2ndIE is lower than expected.

True for s2

Alkaline earth metals form

2+ ions.


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3rd IE

Using the same logic s2p1atoms have an low 3rd IE.

Atoms in the aluminumfamily form 3+ ions.

2nd IE and 3rd IE arealways higher than 1st IE!!!


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Trends in Ionic Size: Cations

Cations form by losing electrons.Cations are smallerthan the atom

they came fromnot only do

they lose electrons, they lose anentire energy level.

Metals form cations.

Cations of representativeelements have the noble gasconfiguration beforethem.


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Ionic size: AnionsAnions form by gaining electrons.

Anions are biggerthan the atomthey came fromhave the same

energy level, but a greater area thenuclear charge needs to cover

Nonmetals form anions.

Anions of representative elementshave the noble gas configuration

afterthem.


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Configuration of Ions

Ions always have noble gasconfigurations ( = a full outer level)

Na atom is: 1s22s22p63s1

Forms a 1+ sodium ion: 1s22s22p6

Same configuration as neon.

Metals form ions with theconfiguration of the noble gas

beforethem - they lose electrons.


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Ion Group trends

Each step down agroup is adding an

energy level

Ions therefore getbigger as you go

down, because of

the additionalenergy level.

Li1+

Na1+

K1+

Rb1+

Cs1+


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Ion Period Trends

Across the period from left toright, the nuclear charge

increases - so they get smaller.

Notice the energy level changesbetween anions and cations.

Li1+

Be2+

B3+

C4+

N3- O2-

F1-


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Size of Isoelectronic ions

Iso- means the sameIsoelectronic 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)


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Size of Isoelectronic ions?

Positive ions that have more protonswould be smaller (more protons would

pull the same # of electrons in closer)

Al3+

Mg2+

Na1+Ne F

1- O2- N

3-

13 12 11 10 9 8 7


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#3. Trends in Electronegativity

Electronegativity is the tendency foran atom to attract electrons to itself

when it is chemically combined with

another element.They share the electron, but how

equally do they share it?

An element with a bigelectronegativity means it pulls the

electron towards itself strongly!

El t ti it G T d


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Electronegativity Group TrendThe further down a group, the farther

the electron is away from thenucleus, plus the more electrons anatom has.

Thus, more willing to share.Low electronegativity.


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Electronegativity Period Trend

Metals are at the left of the table.They let their electrons go easily

Thus, low electronegativity

At the right end are the nonmetals.They want more electrons.

Try to take them away from others

High electronegativity.


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The arrows indicate the trend:

Ionization energy and Electronegativity

INCREASEin these directions

At i i d I i i i


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Atomic size and Ionic size increase

in these directions:


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3.1 org, clssify and trend periodic

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