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Periodic TrendsPeriodic Trends
Chapter 6
Octet RuleOctet Rule
Atoms tend to achieve electron configuration of Noble Gases
Octet = EightNoble Gases have eight electrons in their
highest energy levelGeneral Equation for Noble Gases is S2P6
IONSIONS
Ion- is an atom or group of atoms that have a positive or negative charge
A typical atom is electrically neutral because it has an equal amount of protons and electrons
Positive and Negative Ions are formed when at atom donates or receives an electron
An Ion with a positive charge is called a Cation An Ion with a negative charge is called an Anion
Effective Nuclear ChargeEffective Nuclear Charge
Force of attraction between an electron and the nucleus depends on the magnitude of the net nuclear charge acting on the electron and the average distance between the nucleus and the electron.
Force of attraction increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.
Effective Nuclear Charge Cont.Effective Nuclear Charge Cont.
Valence electron in an atom is attracted to the nucleus of the atom and is repelled by the other electrons.
Inner electrons (core) partially shield the outer electrons from the attraction of the nucleus
The effective nuclear charge increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.
Periodic Properties: Effective Periodic Properties: Effective Nuclear ChargeNuclear Charge
Mg has a greater effective nuclear
charge than Na, and is smaller than Na.
Trends in Atomic SizeTrends in Atomic Size
Atomic Radius- ½ the distance between the nuclei of two like atoms in a diatomic molecule
Radius is measured in Picometers1pm = 1 picometer = 1 x 10-12 m
Atomic SizeAtomic Size
Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
Group Trends of Atomic SizeGroup Trends of Atomic Size
Atomic Size generally increases as you move down a group on the periodic table
As you descend, electrons are added to higher principle energy levels and the nuclear charge increases
The outermost orbital is also larger as you move down a group
The shielding of the nucleus by electrons also increases as you move down a group
Group trendsGroup trends
As we go down a group
Each atom has another energy level,
So the atoms get bigger.
HLi
Na
K
Rb
ShieldingShielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus
ShieldingShielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus.
A second electron has the same shielding.
Increasing Atomic SizeIncreasing Atomic Size
Periodic TrendsPeriodic Trends
Atomic Size generally decreases as you move from left to right across a period
As you go across a period, the energy level remains
Each element has one more proton and electron then the preceding
The electrons are added to the same principle energy level
The effect of the increasing nuclear charge on the outermost electrons is to pull them closer to the nucleus
Atomic Size therefore decreases
Periodic TrendsPeriodic TrendsAs you go across a period the radius gets
smaller.Same energy level.More nuclear charge.Outermost electrons are closer.
Na Mg Al Si P S Cl Ar
OverallOverall
Atomic Number
Ato
mic
Rad
ius
(nm
)
H
Li
Ne
Ar
10
Na
K
Kr
Rb
3) Would you expect the atomic radius to be larger or smaller for element 37 than for element 36? Give the reason for your answer.
• We would expect that the atomic radius for element 37 to be larger than that of element 36. The trend is for the elements at the beginning of the period (Alkali metals, such as Rb) to have bigger atomic radii than the elements at the end (e.g., Kr). Therefore, the atomic radii diminish as one goes to the right in the periodic table.
CONCLUDING QUESTIONS:
Trends in Ionization EnergyTrends in Ionization Energy
When an atom gains or loses an electron, it becomes an ion
Ionization Energy- The energy required to overcome the attraction of the nuclear charge and remove an electron from an atom
The energy required to remove the first outermost electron is called the first ionization energy
The energy required to remove the second outermost electron is called the 2nd ionization energy
Ect……
Ionization EnergyIonization EnergyThe second ionization energy is the energy
required to remove the second electron.Always greater than first IE.The third IE is the energy required to
remove a third electron.Greater than 1st of 2nd IE.
The Noble Gases are at the top showing they don’t want to form an Ion
The Alkali are at the bottom of the peaks, showing their ease to form an Ion
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
PROCEDURE "A" QUESTIONS
• For any period, the first ionization energy is highest for the noble gas (Group VIIIA) and lowest for the alkali metals (Group IA). The ionization energy then increases from the alkali metals to the noble gases in a period.
8) If there is a periodic variation between first ionization energy and atomic numbers of the elements, how would you describe it?
What determines IEWhat determines IEThe greater the nuclear charge the greater
IE.Distance form nucleus increases IEFilled and half filled orbitals have lower
energy, so achieving them is easier, lower IE.
Shielding
Group TrendsGroup Trends
1st Ionization energy generally decreases as you move down a group
The size of the atoms increases as you descend, so the outermost electron is farther from the nucleus
The outermost electron should be more easily removed and the element should have a lower ionization energy
Periodic TrendsPeriodic Trends
The 1st ionization energy generally increases as you move from left to right across a period
The nuclear charge increases and the shielding effect is constant as you move across
A greater attraction of the nucleus for the electron leads to the increase in ionization energy
Exceptions at Full and ½ fill orbitals
Firs
t Ion
izat
ion
ener
gy
Atomic number
He
He has a greater IE than H.
same shielding greater nuclear charge H
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li has lower IE than H
more shielding further away outweighs greater
nuclear charge
Li
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Be has higher IE than Li
same shielding greater nuclear
charge
Li
Be
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He B has lower IE than
Be same shielding greater nuclear
charge By removing an
electron we make s orbital half filled Li
Be
B
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
Breaks the pattern because removing an electron gets to 1/2 filled p orbital
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
NeNe has a lower IE
than HeBoth are full,Ne has more
shieldingGreater distance
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Na has a lower
IE than Li Both are s1
Na has more shielding
Greater distance
Na
Firs
t Ion
izat
ion
ener
gy
Atomic number
Driving ForceDriving ForceFull Energy Levels are very low energy.Noble Gases have full orbitals.Atoms behave in ways to achieve noble gas
configuration.
2nd Ionization Energy2nd Ionization Energy
For elements that reach a filled or half filled orbital by removing 2 electrons 2nd IE is lower than expected.
True for s2
Alkali earth metals form +2 ions.
3rd IE3rd 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!!!
Electron AffinityElectron AffinityThe energy change associated with adding
an electron to a gaseous atom.Easiest to add to group 7A.Gets them to full energy level.Increase from left to right atoms become
smaller, with greater nuclear charge.Decrease as we go down a group.
Electron AffinityElectron Affinity
The greater the attraction between a given atom and an added electron, the more negative the atom’s electron affinity
The more negative the E.A., the greater the attraction of the atom for the electron
The trends in E.A. are not very evident.
Difference between I.E. & E.A.Difference between I.E. & E.A.
Ionization Energy measures the ease with which an atom loses an electron
Electron Affinity measures the ease with which an atom gains an electron
Trends in Ionic SizeTrends in Ionic Size
Atoms of metallic elements have low ionization energies. They form positive ions easily
Atoms of nonmetallic elements readily form negative ions.
How does the lose or gain of electrons affect the size of the ion formed?
Group TrendsGroup TrendsPositive Ions are always smaller than the
neutral atoms from which they form.The loss of outer-shell electrons results in
increased attraction by the nucleus for the fewer remaining electrons
Negative Ions are always larger than the neutral atoms from which they form
The effective nuclear attraction is less for an increased number of electrons
Group trendsGroup trends
Adding energy levelIons get bigger as you
go down.
Li+1
Na+1
K+1
Rb+1
Cs+1
The Sodium Atom is larger than the Sodium Cation.
Why is this true?
The Chlorine Atom is smaller then the Chlorine Anion.
Why is this true?
Sodium Cation is smaller than the Sodium Atom
Chlorine Anion is larger than the Chlorine Atom
Periodic TrendsPeriodic Trends
Going from left to right across a row, there is a gradual decrease in the size of the positive ions.
Beginning with group 5A, the negative ions, which are much larger, gradually decrease in size an you continue to move right.
Periodic TrendsPeriodic TrendsAcross the period nuclear charge increases so
they get smaller.Energy level changes between anions and
cations.
Li+1
Be+2
B+3
C+4
N-3O-2 F-1
Trends in ElectronegativityTrends in ElectronegativityElectronegativity- is the tendency for the
atoms of the element to attract electrons when chemically combined with atoms of another element.
Electronegativities have been calculated for elements and are expressed in arbitrary units on the Pauling electronegativity scale
The scale is based on a number of factors
Group TrendsGroup Trends
Electronegativity generally decreases as you move down a group
The metallic elements have a low electronegativity meaning they don’t want to want attract electrons
Periodic TrendsPeriodic Trends
As you go across a period from left to right, the electronegativity of representative elements increases
The non-metallic elements (excluding Noble Gases) have high electronegativities
The trends in electronegativities among transitional metals are not so regular
Electronegativity values help predict the type of ionic or covalent bonding that can exist between atoms in compounds
