THE PERIODIC TABLETHE PERIODIC TABLETHE PERIODIC TABLETHE PERIODIC TABLE

Objectives:

Examine the progression of periodicity

Alkali metals

alkaline earth metals

“s” groupNonmetals

“p” block

Metalloids (semimetals)

Transition metals

“d” blockInner transition metals

“f” group

Noble gases

Halogens

Periodic PatternsPeriodic Patterns Periodic PatternsPeriodic Patterns

ALKALI METALS (part of the “s” group of elements)~ all are in shiny solid form but are quite soft~ form the 1st group of metals on the periodic table~ highly reactive elements based upon their electron configurations (ns1)

Other Characteristics: malleable and ductile; low density and melting points good conductors of electricity; very soluble as comps.

Alkaline Earth MetalsAlkaline Earth MetalsAlkaline Earth MetalsAlkaline Earth Metals

Belong to the second group of metals on the periodic table.

Harder, more dense, and stronger than there group 1 counterparts

Not as reactive as the Alkali metals due to the metals in this group having 2 electrons in their valence shell.

This gives them the configuration of “ns2” for these metals.

PART OF THE “S” GROUP JUST LIKE ALKALI METALS

Be

Sr

Ba

Mg

Ra

Ca

Transition Transition MetalsMetalsTransition Transition MetalsMetals

Transition metals begin in the 4th period after the alkaline earth metals.

Metallic elements with varying properties.

Not nearly as reactive as group 1 and 2 elements.

Fill their sublevels differently than do the Main group elements.

The “d” - block elements

Valuable as structurally useful materials!

Important in living organisms!

Lanthanoids & ActinoidsLanthanoids & Actinoids

LANTHANOIDS

Composed of the elements with atomic numbers 58 through 71

Electrons are being added to the “ 4f “ sublevel

Shiny reactive metals with practical uses

ie. dots in TV tubes

ACTINOIDS

Composed of the elements with the atomic numbers 90 through 103

They fill the “ 5f “ sublevel

All are radioactive with an unstable nucleus

“Y”The ‘f’-group is broken into two classifications

Nonmetals & Metalloids (semi-metals)Nonmetals & Metalloids (semi-metals)

NONMETALS! Generally are gases at room temperature (or brittle solids) Poor conductors of heat and electricity Have more electrons in their outer level than metals

METALLOIDS! Properties of both metals and nonmetals Will give up (electron donor) electron(s) when reacted with a

nonmetal, and will accept (electron acceptor) electron(s) when reacted with a metal

In general, more like nonmetals than metals Considered semiconductors

Periodic TrendsPeriodic Trends

Trends (we will study) – atomic radius (ionic radius), ionization energy, electronegativity, electron affinity

Trends are looked at from top to bottom of a column and from left to right in a period (row)

Trends show patterns of atoms properties (relationships among elements)

ATOMIC RADIUS

Li

Na

K

Rb

Cs

Fr

Atomic radius is the half the distance between the nuclei of two like atoms.

Trend Number 1Trend Number 1Trend Number 1Trend Number 1

the trend for atomic radius shows us the size of the atom will increase as we move down a column

WHY: more levels and orbitals, greater distances from the nucleus

TREND:

TREND: the atomic radii will decrease from the left to

the right in a period

WHY: Effective Nuclear Charge (also applies to what

takes place from top to bottom of a column) positive charge felt by the outermost

electrons of an atom atomic # - # of inner complete level electrons The larger the ENC, the greater the

attraction of electrons to the nucleus

Shielding - the ability of other electrons,especially inner

electrons, to lessen the nuclear charge of the outer electron(s)

ATOMIC RADIUSATOMIC RADIUSATOMIC RADIUSATOMIC RADIUS

THE TREND!THE TREND!

The trend shows the increase of radii down a group and decrease of radii across a period.

IONIZATION ENERGYIONIZATION ENERGYIONIZATION ENERGYIONIZATION ENERGY

IONIC BOND bond formed between two ions by the

transfer of electrons

Ions: How do they form? In certain types of bonding, the atom

will “lose” or “gain” an electron(s)

When an atom loses or gains electrons, it is called an ion

Magnesium

NUMBER TWO!NUMBER TWO!NUMBER TWO!NUMBER TWO!

Atoms that lose electrons have a positive charge

Atoms that gain electrons have a negative charge

Magnesium

BOINK!

BOINK!

For the most part, the metals will lose electrons and the nonmetals will accept the electrons

The atoms gain or lose electrons to reach outer shell (valence) stability

CHLORINE

Electron from magnesium

Ionic Bonds: One Big Greedy Thief Dog!

Ion SizesIon SizesIon SizesIon Sizes

Li,152 pm3e and 3p

Li+, 60 pm2e and 3 p

+

Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

Ion SizesIon SizesIon SizesIon Sizes

CATIONSCATIONS are are SMALLERSMALLER than the atoms from than the atoms from which they come.which they come.

The electron/proton attraction has gone UP and The electron/proton attraction has gone UP and so size so size DECREASESDECREASES..

Li,152 pm3e and 3p

Li +, 78 pm2e and 3 p

+Forming Forming a cation.a cation.Forming Forming a cation.a cation.

Ion SizesIon SizesIon SizesIon Sizes

F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-

Does the size go up or down when Does the size go up or down when gaining an electron to form an gaining an electron to form an anion?anion?

Ion SizesIon SizesIon SizesIon Sizes

ANIONSANIONS are are LARGERLARGER than the atoms from which they than the atoms from which they come.come.

The electron/proton attraction has gone DOWN and so The electron/proton attraction has gone DOWN and so size size INCREASESINCREASES..

Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm

9e and 9pF-, 133 pm10 e and 9 p

-

Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes

Figure 8.13Figure 8.13

IONIZATION ENERGY

the energy required to remove the most loosely held electron from an atom

ionization energy decreases as the size of the atom increases (top to bottom of a column)

“Y”? Because the outer most electron is farther from the nucleus and the electrical attraction to the protons.

More Details!More Details!

Energy is absorbed by the atom to free the electron(s)

Ionization is endothermicendothermic, meaning that the atom or molecule increases its internal energy ( takes energy from an outside source)

A + energy A+ + e-

Ionization Energy is affected by three factors:

1. Effective Nuclear Charge

2. Number of Energy Levels

3. Shielding

Ionization EnergiesIonization Energies The first ionization energy, I1, is the energy needed to

remove the first electron from the atom:

Mg Mg+ + 1e-

The second ionization energy, I2, is the energy needed to remove the next (i.e. the second) electron from the atom

Mg+ Mg2+ + 1e-

•The The higherhigher the value of the the value of the ionization energy, the more ionization energy, the more difficultdifficult it is to remove the it is to remove the electronelectron

1st IE 2nd IE 3rd IE 4th IE 5th IE 6th IE 7th IE

Na 496 4,560

Mg 738 1,450 7,730

Al 577 1,816 2,881 11,600

Si 786 1,577 3,228 4,354 16,100

P 1,060 1,890 2,905 4,950 6,270 21,200

S 999.6 2,260 3,375 4,565 6,950 8,490 27,107

Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000

Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000

Ionization Energies in kJ/mol

• Within each period ( row) the ionization Within each period ( row) the ionization energy energy increases with atomic number.increases with atomic number.

•Y?Y?-Electrons are being added to the same Electrons are being added to the same energy level (ENC)energy level (ENC)

- increasing valence electrons as increasing valence electrons as approaching the nonmetalsapproaching the nonmetals

Na Mg Al Si P S Cl Ar

The TrendThe Trend

ElectronegativityElectronegativity The tendency for an atom to attract electrons to itself when

in combination with another atom

Defined differences in electronegativity determine the bonding character of a compound

• Ionic or Covalent bonds

Linus Pauling scale is used to determine electronegativity differences

•

COVALENT BONDCOVALENT BOND

bond formed by the bond formed by the sharing sharing of electron cloudsof electron clouds• Between nonmetallic elements of similar electronegativity.

•Formed by sharing electron pairs

when electron clouds are shared when electron clouds are shared equallyequally

NONPOLAR NONPOLAR COVALENT BONDSCOVALENT BONDS

HH22 or Cl or Cl22

2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.

Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (OOxygen Molecule (O22))

• when electron clouds are shared but when electron clouds are shared but shared shared unequallyunequally

POLAR COVALENT BONDSPOLAR COVALENT BONDS

HH22OO

Polar Covalent Bonds: Unevenly matched, but willing to share.

- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

Electronegativity Differences and Bond Type

nonpolar

covalent

polar covalent

ionic

• If the electronegativity difference is less than 0.2 then the bond is a nonpolar covalent

• If the difference is between 0.2 and 1.6, the bond is polar covalent

•If the difference is greater than 2, the bond is ionic

?

? between 1.6 and 2, if a metal is involved, the bond is ionic. If only nonmetals are involved the bond is polar covalent

0

4

2

1.6

0.2

Trend of ENTrend of EN

decrease

increase

Electron AffinityElectron Affinityelements elements GAINGAIN electrons to form electrons to form anionsanions..

Electron affinity is the energy change when an Electron affinity is the energy change when an electron is added:electron is added:

A(g) + e- ---> AA(g) + e- ---> A--(g) E.A. = ∆E(g) E.A. = ∆E

Electron Affinity of OxygenElectron Affinity of Oxygen

∆∆E is E is EXOEXOthermic thermic because O has an because O has an affinity for an e-.affinity for an e-.

[He] O atom

EA = - 141 kJ

+ electron

O [He] - ion

Affinity for electron increases Affinity for electron increases across a period (EA becomes across a period (EA becomes more negative).more negative).

Affinity decreases down a group Affinity decreases down a group (EA becomes less negative).(EA becomes less negative).

Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Atom EAAtom EAFF -328 kJ-328 kJClCl -349 kJ-349 kJBrBr -325 kJ-325 kJII -295 kJ-295 kJ

Trends in Electron AffinityTrends in Electron Affinity

Trends in Electron AffinityTrends in Electron Affinity

Practice with Comparing Practice with Comparing Ionization EnergiesIonization Energies

For each of the following sets of atoms, decide which has the highest and lowest ionization energies and why.

a. Mg, Si, S

b. Mg, Ca, Ba

c. F, Cl, Br

d. Ba, Cu, Ne

e. Si, P, N

Answers to Comparing Ionization Energies

Here are answers to the exercises above.

a. Mg, Si, Sa. Mg, Si, S

All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear charge.

b. Mg, Ca, Bab. Mg, Ca, Ba

All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of energy levels.

c. F, Cl, Brc. F, Cl, Br

All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of energy levels.

d. Ba, Cu, Ned. Ba, Cu, Ne

All are in different groups and periods, so both factors must be considered. Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear charge and uses the lowest number of energy levels.

e. Si, P, Ne. Si, P, N

Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for having the highest effective nuclear charge.

BECAUSE...The relative stability of an atom can be predicted by its electron configuration

Rule of ThumbRule of Thumb

• As a general rule, elements with three or fewer electrons in their outer level are considered to be metals.

Lets Review!Lets Review!1. What is the periodic Law?

2. How is an element’s outer electron configuration related to its position in the periodic table?

3 Indicate which element in each of the following pairs has the greater atomic radius.

a. sodium & lithium b. strontium & magnesium

c. carbon & germanium d. selenium & oxygen

4. In general, would you expect metals or nonmetals to have higher ionization energies?

More review!More review!

5. Arrange the following elements in order of increasing ionization energies.

a. Be, Mg, Sr b. Bi, Cs, Ba c. Na, Al, S

6. How does the ionic radius of a typical metallic atom compare to its atomic radius?

7. Explain why it takes more energy to remove a 4s electron from an atom of zinc than from than from an atom of calcium.

8. Give the symbol of the element found at each of the following locations in the periodic table.

a. group 1, period 4 b. group 13, period 3

c. group 2, period 6 d. group 10, period 2

Test

Friday

Even more review!Even more review!9. What was Newland’s Law of Octaves all about?

10. How was Mendeleev’s periodic table of elements better than the previous attempts by others?

11. What property do the noble gases share? How does this property relate to the electron configuration of the noble gases?

12. How do the electron configurations of the transition metals differ form the electron configurations of the metals in groups 1 and 2?

13. What group numbers make up the main-block elements?

This test will be a bear if you forget to study!

Are you kidding me, more good stuff!Are you kidding me, more good stuff!Are you kidding me, more good stuff!Are you kidding me, more good stuff!

14. Define what ionization energy and electron affinity are.

15. What periodic trends exist for ionization energy? How about for electron affinity? What about atomic radius and its trend?

16. Why does the first period of the periodic table contain only two elements while all the other periods have eight or more element in them?

17 What feature of electron configuration is unique to actinoids and lanthanoids?

That should just about cover it!

TO REINFORCE YOUR ALREADY EXTENSIVE KNOWLEDGE OF THE PERIODIC TABLE, YOU CAN READ THROUGH THE PAGES IN YOUR BOOK OF CHAPTER 14.

Lastly, if you need to out any last minute problems you can show up at 7:00 in room 224 for some last minute brushing up.

Transition cont. Transition cont.

IT’S ALL ABOUT ELECTRONS!IT’S ALL ABOUT ELECTRONS!

Transition elements fill their sublevels differently than do the Main group elements.

For the most part, there are a few exceptions, these “d” block metals will place the 2 electrons into a higher s-sublevel before the electrons go into a “d” energy sublevel.

Inner Transition MetalsInner Transition Metalsthe “ f ”-groupthe “ f ”-group

The ‘f’-group is broken into two classifications

~lanthanides

~actinides