Ch. 5 “The Periodic Table”

Why is the Periodic Table important to me?

• The periodic table is the most useful tool to a chemist.

• It organizes lots of information about all the known elements.

• You get to use it on every test.

Pre-Periodic Table Chemistry …

• …was a mess!!!• No organization of

elements.• Imagine going to a

grocery store with no organization!!

• Difficult to find what you need.

• Chemistry didn’t make sense.

(CHAOS)

During the nineteenth century, chemists began to categorize

the elements according to similarities in their physical and chemical properties. The end result of these studies was our

modern periodic table.

How did chemists begin to organize the known elements?

Chemists used the properties of elements to sort them into groups Ex. Chlorine, bromine, and iodine have very similar chemical properties.

As the number of elements increased, chemists inevitably began to find patterns in their properties.

1780 - 1849

Model of TriadsIn 1817, Johann Dobereiner classified some elements into groups of three, which he called triads.The elements in a triad had similar chemical and physical properties.

Law of Octaves

1838 - 1898

In 1865, John Newlands suggested that elements be arranged in “octaves” because he noticed when he arranged the elements in order of increasing atomic mass certain properties repeated every 8th element.

LiBeBCNOF

NaMgAlSiPS

Cl KCa??

AsSeBr

He called this the “Law of Octaves”

because of its similarity to

musical octaves

Lightest to heaviest.LiBeBCNOF

NaMgAlSiPS

Cl

KCa??

AsSeBr

ABCDEFGABCDEFGABCDEFG

John Newlands Law of Octaves

Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society.

The Modern Periodic Table

1834 - 1907

In 1869 Dmitri Mendeleev published a table of the elements organized by increasing atomic mass.He was trying to organize elements so his students could learn them more easily!

A. Mendeleev and Chemical Periodicity

• Mendeleev placed known information of elements on cards (atomic mass, density, etc…). He arranged them in order of increasing atomic masses, certain similarities in their chemical properties appeared at regular intervals. Such a repeating pattern is referred to as periodic.

1830 - 1895

At the same time, Lothar Meyer published his own table of the elements organized by increasing atomic mass.

• Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass.

• Both left vacant spaces where unknown elements should fit.

So why is Mendeleev called the “Father of the Periodic Table” and not Meyer, or both?

Mendeleev published first!Could it be his dashing good looks?!

• Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match. The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces. When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery.

and the elusive element 32…

Predicted Properties

Observed Properties

Atomic weight 72 72.61

Density 5.5 g/cm3 5.32 g/cm3

Melting point 825 C 938 C

Oxide formula RO2 GeO2

Density of oxide 4.7 g/cm3 4.70 g/cm3

Dates (predicted and found)

1871 1886

Color Dark gray Gray-white

Gallium

Germanium

After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions were amazingly close to the actual values, his table was generally accepted.

The Father of the Periodic Table

However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights were determined.

By looking at our modern periodic table, can you identify what problems might have caused chemists a headache?

Ar and KCo and NiTe and I

Th and Pa

1887 - 1915

In 1913, through his work with X-rays, Henry Moseley determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.*“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.”

Remember This…?!

Henry Moseley

His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

Periodic LawWhen elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

The Current Periodic Table

• Mendeleev wasn’t too far off.• Now the elements are put in rows by

increasing ATOMIC NUMBER!!

• The vertical columns are called groups or families and are labeled from 1 to 18 (modern)

• or in A & B Groups (with Roman numerals)

Groups…Here’s Where the Periodic Table Gets Useful!!

• Elements in the same group have similar chemical and physical properties!!

• (Mendeleev did that on purpose.)

Why??• They have the

same number of valence electrons.

• They will form the same kinds of ions.

Groups in the Periodic Table

Elements in groups react in similar ways!

Periods in the Periodic Table

11

1

2

3

4

5

6

7

The horizontal rows are called periods and are labeled from 1 to 7.All elements in a period have the same number of energy levels (= to period #)

Energy LevelsEnergy Levels

n = 1

n = 2

n = 3

n = 4

In addition to Group Labels, many of the groups have Family Names

Group 1A: Alkali Metals

Group 1A: Alkali Metals

Cutting sodium metal

lithium

potassium

Alkali Metals • They are the most reactive metals.

• They react violently with water.

• Alkali metals are never found as free elements in nature - they are always in compounds with other elements.

• Only 1 valence electron

• Soft metals• Must be stored under

mineral oil, etc.

Magnesium

Group 2A: Alkaline Earth Metals

Group 2A: Alkaline Earth Metals

Only 2 valence electronsToo reactive to be uncombined in nature.

calcium

strontium

barium

Group 7A: The Halogens

7 valence electronsAll non-metalsVery reactiveAll physical states representedColored gases (always poisonous!)Occur as diatomic molecules when pure

fluorine

F2 Cl2Br2

chlorine bromine

Iodine I2

The Noble Gases

Noble Gases

• Noble Gases are colorless gases that are extremely un-reactive.(inert)

• They are inactive because their outermost energy level is full. (8 valence electrons – except He which has 2)

• Having 8 valence electrons is low in energy and, therefore, very stable.

Hydrogen• The hydrogen square sits atop

Family IA, but it is not a member of that family. Hydrogen is in a class of its own. (An orphan?)

• Like the Alkali metals, it only needs to lose one electron to be stable. (but it is not a metal!)

• Sometimes it’s shown above 7A.• Like the Halogens, it only needs to

gainone electron to have the stable Noble Gas electron configuration. (but it is not a Halogen!)

Hey Cameron, why are those elements by themselves on the bottom of the Periodic Table?!

I’ll handle this one, Cam! If they weren’t put on the bottom, the Periodic Table wouldn’t fit very nicely on a page! In fact, the table would look like this.

In fact, we have Glen Seaborg to thank for the fact that my Periodic Table doesn’t stick out of my notebook in a truly tasteless manner!

Glenn T. SeaborgAfter co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became knownas the Actinide series.

1912 - 1999

“I was warned at the time that it was professional suicide to promote this idea, which has since been called one of the most significant changes in the periodic table since Mendeleev’s 19th century design. Luckily, I stuck to my guns and have seen the actinide concept become the foundation for many significant discoveries in heavy element research.”

Seaborgium

Glenn T. SeaborgHe is the only person to have an element named after him while still alive.

1912 - 1999

"This is the greatest honor ever bestowed upon me - even better, I think, thanwinning the Nobel Prize." 106

SgSeaborgium

271

There are many ways that we can break the PeriodicTable up into sections!

MetalsMetals

Metals

Metals are good conductors of heat. That's why a branding iron is made from metal. The heat transfers quickly to the animal's hide.

Metals also conduct electricity.

Notice that the Tesla coil sparks seek out metallic objects because they conduct electricity better than the

nonmetallic materials such as wood or soil.Metals are also malleable and can be bent or hammered

into various shapes.

Metals• What comes to mind?• Most elements are metals• Loosely held valence e-’s• Properties of metals:

1. Good conductors of heat and electricity (p)

2. High density (p)3. High melting points (p)4. Luster (p)5. Malleable (p)6. Ductile (p)7. 1, 2, or 3 valence electrons

Nonmetals

Nonmetals do not conduct heat well. The insulating tiles from the Space Shuttle are made from fibers of silicon and oxygen (silica=sand).

Nonmetals

• Opposite of metals• Properties of nonmetals:

1. Dull (no luster)

2. Do not conduct heat/elec.

3. Not ductile

4. Not malleable

5. All phases

6. Have 5, 6, or 7 valence electrons

• Form many compounds with metals

Metalloids(Semi-Metals)

• Means “metal-like”• Dividing line

between metals and nonmetals

• Al is the exception• Properties of both

metals and nonmetals

Four Main Categories of Elements

• Noble Gases- group 18 or 0 or 8A– s & p sublevels filled– 8 valence __s2…__p6

– Inert- not reactive- because of full outer shell of electrons

• Representative Elements also called main group elements- Groups 1A-7A– s & p partly filled– Includes alkali metals, alkaline earth

metals, and halogens

1A

2A 3A 4A 5A 6A7A

8A0

The elements in the A groups are called the representative elements

Four Main Categories of Elements

• Transition Metals – – Unfilled inner shells – outermost s & inner d sublevels contain

electrons – Hard & brittle

• Inner transition metals- – outermost s & nearby f sublevel contain

electrons – Lanthanides (4f) and actinides (5f)

RepresentativeRepresentative

Inner Transition Elements

Noble Gases

Using the Diagonal Rule is just so bothersome! I wish there was an easier way to figure out electron configurations!

Oh, but thereis! Watch this!

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p63d104s24p65s1

1s22s22p63s23p63d104s24p64d105s2 5p66s1

1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

1

2

3

4

5

6

7

Group 2A

Be

Mg

Ca

Sr

1s22s2

1s22s22p63s2

1s22s22p63s23p64s2

1s22s22p63s23p64s23d104p65s2

1H

2He

2 3Li

4Be

3 11Na

12Mg

4 19K

20Ca

5 37Rb

38Sr

6 55Cs

56Ba

7 87Fr

88Ra

s-block

Always for row you are on!

B

Al

Ga

Group 3A

1s22s22p1

1s22s22p63s23p1

1s22s22p63s23p64s23d104p1

The P-block p1 p2 p3 p4 p5 p6

Always for row you are on!

Electron Configurations in Groups• In atoms of the Group 1A elements below, there

is only one electron in the highest occupied energy level.

• In atoms of the Group 4A elements below, there are four electrons in the highest occupied energy level.

6.2

It’s always s1 forthe row it’s on!

Always s2p2 for the row they’re on!

Electron Configurations in Groups

– The Noble Gases• In atoms of the Group 8A elements below, there

are eight electrons in the highest occupied energy level.

6.2

Except for He, always s2p6 for the row they are on!

Chemical elements in d-block

Group →

3 4 5 6 7 8 9 10 11 12

↓ Period

4 21Sc

22Ti

23V

24Cr

25Mn

26Fe

27Co

28Ni

29Cu

30Zn

5 39Y

40Zr

41Nb

42Mo

43Tc

44Ru

45Rh

46Pd

47Ag

48Cd

6 71Lu

72Hf

73Ta

74W

75Re

76Os

77Ir

78Pt

79Au

80Hg

7 103Lr

104Rf

105Db

106Sg

107Bh

108Hs

109Mt

110Ds

111Rg

112Uub

Always 1 level in (back) from the row you’re on!

3d

4d

5d

6d

F - block inner transition elements

f1 f5f2 f3 f4

f6 f7 f8 f9 f10 f11 f12 f14

f13

Always go back 2 energy levels from the row you’re on!

4f

5f

s-block

Blocks of Elements

d-block p-block

f-block

Periodic Table e- configuration from the periodic

periodic table(To be covered in future chapters)

•B•2p1

1IA

18VIIIA

12

IIA13

IIIA14

IVA15VA

16VIA

17VIIA

2

33

IIIB4

IVB5

VB6

VIB7

VIIB8 9

VIIIB10 11

IB12IIB

4

5

6

7

H1s1

Li2s1

Na3s1

K4s1

Rb5s1

Cs6s1

Fr7s1

Be2s2

Mg3s2

Ca4s2

Sr5s2

Ba6s2

Ra7s2

Sc3d1

Ti3d2

V3d3

Cr4s13d5

Mn3d5

Fe3d6

Co3d7

Ni3d8

Zn3d10

Cu4s13d10

B2p1

C2p2

N2p3

O2p4

F2p5

Ne2p6

He1s2

Al3p1

Ga4p1

In5p1

Tl6p1

Si3p2

Ge4p2

Sn5p2

Pb6p2

P3p3

As4p3

Sb5p3

Bi6p3

S3p4

Se4p4

Te5p4

Po6p4

Cl3p5

Be4p5

I5p5

At6p5

Ar3p6

Kr4p6

Xe5p6

Rn6p6

Y4d1

La5d1

Ac6d1

Cd4d10

Hg5d10

Ag5s14d10

Au6s15d10

Zr4d2

Hf5d2

Rf6d2

Nb4d3

Ta5d3

Db6d3

Mo5s14d5

W6s15d5

Sg7s16d5

Tc4d5

Re5d5

Bh6d5

Ru4d6

Os5d6

Hs6d6

Rh4d7

Ir5d7

Mt6d7

Ni4d8

Ni5d8

Periodicity

When one looks at the chemical properties of elements, one notices a repeating pattern of properties when the elements are in order of

increasing atomic number.

1. Nuclear charge- the number of protons in the nucleus.

More protons = increased nuclear charge so increased attraction between the nucleus and electrons.

Think of the nucleus as a magnet – each extra proton makes the magnet more powerful at attracting electrons & holding them tight!

Explaining Periodic Trends

Why a property is higher/lower, bigger/smaller, etc.!

2. Shielding- lessens the attractive force of the nucleus for the valence electrons– caused by electrons in energy levels

between the nucleus and the valence electrons

Shielding increases as you go down a group

because there are more energy levels (more core electrons). Shielding stays the same as you move across a period because the number of energy levels is

staying the same.Which atom has more shielding? (A) K or Ca (B) Na or KWhich atom is smaller? (A) N or P (B) Li or K

Metallic Character

Atomic Radius

• Atomic Radius- half the distance between the nuclei of two atoms of the same element in a diatomic molecule

Atomic Radius

• Trend for atomic size- – Down a group, size increases

• Occurs because # of energy levels increases

*Makes a BIG difference in size!!• shielding also increases.

– Across a period, size decreases• # of protons increases (nuclear

charge increases), pulling electrons closershielding doesn’t change because electrons are added to the same energy level

Atomic Radius

Ionization Energy• Ionization Energy- energy needed to

remove an electron from an atom. • Outer shell electrons are easier to remove than

‘core’ electrons so it takes less energy to remove them!

Highest toward upper right corner of PT since

these atoms are smaller & their valence

electrons are closer to the nucleus

-so held more tightly

Trends in Ionization Energy

Periodic Trends

• Ionization energy– Down a group- decreases – because electrons are held more loosely

due to increased # of energy levels & increased shielding

– Across a period- increases because electrons are held more tightly due to increased nuclear charge (increased # of protons in the nucleus)

Ionization Energy

Ionization Energy

• There are big jumps in ionization energy whenever you try to remove an electron from an inner energy level!

Ionization Energy

Electron Affinity

• Electron affinity of an element is the energy given off when an atom (in the gas phase) gains an electron to form an ion

– Example: F(g) + e- F-(g) – Ho (ENERGY) = -328.0 kJ/mol

Trends in Electron Affinity

• It decreases down a group, because electron shielding blocks some of the attraction from the nucleus

• It increases across a period, because nuclear charge increases, attracting electrons more strongly.

Periodic Trends

• Electronegativity- tendency for the atoms of the element to attract electrons when the atoms are part of a compound

• Noble gases- no electronegativity values- don’t form compounds

• In general, metals have low EN and nonmetals have high EN.

• The actual amount of EN an atom has is indicated by a number on the Pauling Electronegativity Scale that goes from 0 to 4.

• Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element.

• Flourine is the most electronegative element at 4. (3.98) and Francium is the least electronegative at 0.7.

Periodic Trends

• Electronegativity Trends- – Down a group – decreases- electron shielding

results in less attraction for electrons by the nucleus

– Across a period- increases- higher atomic number and consistent electron shielding result in more attraction for electrons

• Electronegativity allows you to predict bond type: covalent- polar vs. nonpolar and ionic

Ionization Energy (IE)

Electron Affinity (EA) Electronegativity ( )c

Atomic & Ionic Radii

General Trends in the Periodic Table:

Decreasing

Increasing

IncreasingIncreasing

IE, EA, and c are useful concepts used to characterize different types of bonding and estimate bond energies.

Summary of Trends

6.3

Atomic Size Increases

Incr

ease

s

Decreases

Dec

reas

es

Size of cationsShieldingNuclear ChargeElectronegativityIonization energySize of anionsIonic size Constant

IONS • Remember – Atoms are neutral

• But…atoms can gain or lose electrons

(*# of protons NEVER changes during

reactions!)

• IONS are atoms or groups of atoms with

a charge.

• To tell the difference between an atom and an ion, look to see if there is a charge in the superscript!

• Examples: Na Ca I O• Na+ Ca+2 I- O-2

• When an atom loses an electron it gets a

positive charge (because it now has more

protons than electrons)

Mg --> Mg+2 + 2 e-

When an atom gains an electron it forms a negative ion (because it now has more electrons than protons)

F + e- --> F-

Atom versus Ion

Forming Cations & AnionsForming Cations & AnionsA CATION forms when an atom loses one or more electrons.

An ANION forms when an atom gains one or more electrons

Mg --> Mg2+ + 2 e-F + e- --> F-

Now has 12 protons & 10 electrons

Now has 9 protons & 10 electrons

–Metals have 1, 2, or 3 valence electrons so tend to lose electrons (to get an octet)- forming cations. (+ charge)

–Non-Metals have 5, 6, or 7 valence electrons so tend to gain electrons (to get an octet) - forming anions. (- charge)

Periodic Trends

• Ionic Radii Trends– Cations- smaller than neutral atom because

fewer electrons result in greater attraction by nuclei

– Anions- larger than neutral atom because more electrons result in less attraction by nuclei

– Within period- size decreases – Down a group – size increases

Forming cations

Forming anions

What would the charge be on a sodium ion?

EXAMPLE

Since sodium in in Group IA it has 1 valence e-

and so it would LOSE an electron

So it gets a charge of +1Remember an electron is negatively charged. When an atom loses electrons it forms positively charged ions.

When electrons are gained negatively charged ions form

It goes from 11 protons & 11 electrons to11 protons & 10 electrons

How would you write the symbol for the sodium CATION?

EXAMPLE

Na+1

How many outer electrons does sodium have before it loses one?

It has 1…remember the group number!

5

– 1. Which of the following sequences is correct for atomic size?• Mg > Al > S• Li > Na > K• F > N > B• F > Cl > Br

6.3 Section Quiz

6.3 Section Quiz

– 2. Metals tend to• gain electrons to form cations.• gain electrons to form anions.• lose electrons to form anions.• lose electrons to form cations.

6.3 Section Quiz

– 3. Which of the following is the most electronegative?• Cl• Se• Na• I

The Periodic Table

Summary of Trend• Periodic Table and Periodic Trends• 1. Electron Configuration

2. Atomic Radius: Largest toward SW corner of PT

3. Ionization Energy: Largest toward NE of PT4. Electron Affinity: Most favorable NE of PT

Periodic Table: electron behavior

• The periodic table can be classified by the behavior of their electrons1IA

18VIIIA

12

IIA13

IIIA14

IVA15VA

16VIA

17VIIA

2

33

IIIB4

IVB5

VB6

VIB7

VIIB8 9

VIIIB10 11

IB12IIB

4

5

6

7

West (South) Mid-plains East (North)METALS

AlkaliAlkaline

Transition

METALLOID NON-METALSNoble gasHalogensCalcogens

These elementstend to give up

e- and formCATIONS

These elementswill give up e- or

accept e-

These elementstend to accept

e- and formANIONS

ELEMENTS THAT EXIST AS DIATOMIC MOLECULES

ELEMENTS THAT EXIST AS DIATOMIC MOLECULES

Remember:

BrINClHOFThese elements only exist as

PAIRS. Note that when they

combine to make compounds, they

are no longer elements so they are no longer in

pairs!

1 2 3 4 5 6 7 8

Valence electrons

Select an element

= Internet link( )

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 Ions

This can explain why metals are shiny. This is the surface of copper at a ridiculously high magnification. The surface shows a lake of electrons along with ripples. The two islands are imperfections on the surface. Most likely a couple of atoms that aren't copper.