THE PERIODIC TABLE
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Transcript of THE PERIODIC TABLE
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