Periodic Trends Presentation student notes.notebookyooschem1314.pbworks.com/w/file/fetch/70754887/Periodic Trends … · B diamagnetism C electronegativity D electron changetomass

Periodic_Trends_Presentation student notes.notebook

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October 05, 2012

Periodic Trends:

Atomic RadiusIonization EnergyElectronegativityMetallic Character

Ionic Radius

Periodic Trends

Five main trends in the periodic table will be discussed:The sizes of atoms Ionization energyElectronegativityMetallic characterThe sizes of ions

These five periodic trends are affected by these main factors:Effective nuclear chargeCoulomb's LawShielding from inner electrons

Effective Nuclear Charge

• In a manyelectron atom, electrons are both attracted to the nucleus and repelled by other electrons.

• The nuclear charge that an electron experiences depends on both factors.

For example, here's sodium.

The effective nuclear charge, Zeff, is found this way:

Zeff = Z − S

Where Z is the atomic number and S is a screening constant, usually about equal to the number of inner electrons.

In this example, the outer electron of sodium is attracted towards the nucleus by an effective charge (Zeff) of 1 proton.


Moving across the periodic table increases Zeff, and therefore increasing the force of attraction between the electrons and the nucleus.


• For any given atom, the electrons closest to the nucleus experience the greatest effective nuclear charge (they have the least shielding).

• The electrons farthest from the nucleus experience the least effective nuclear charge (they are shielded the most).

As you move across the periodic table (from left to right), the effective nuclear charge felt by the outermost electron increases.

Atomic Radius

• In general, as you move from left to right across the periodic table the atoms have a smaller radius

• As you move down from row to row, the radius increases

Helium has the smallest radius

Francium has the largest radius


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Atomic radius decreases going left to right across a period.

• Across a period, effective nuclear charge increases.

• So the force of attraction between the nucleus and outer electron(s) gets stronger.

• Outer electrons are pulled in more strongly, so radius gets smaller.

• Size of orbitals stays approximately the same (across a given period)

Summary of Atomic Radius Trends

Atomic radius increases going down a group. • The size of orbitals increases significantly.

• Outer electrons are located farther from the nucleus in each successive period.

• The "r" in Coulomb's law increases so the force of attraction between nucleus and outer electrons weakenss

• Electronelectron repulsion increases with greater atomic number.

Summary of Atomic Radius Trends1 The effective nuclear

charge acting on an electron is larger than the actual nuclear charge.

True

False

2 A selenium atom has 34 electrons. Electrons in the __________ subshell experience the lowest effective nuclear charge.

A 1s

B 3p

C 3d

D 4p

E 5p

3 In which orbital does an electron in an arsenic (As) atom experience the greatest shielding?

A 2p

B 4p

C 3p

D 3s

E 1s


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4 In which orbital does an electron in a calcium atom experience the greatest effective nuclear charge?

A 1s

B 2s

C 2p

D 3s

E 3p

5 Atomic radius generally increases as we move __________.

A down a group and from right to left across a period

B up a group and from left to right across a period

C down a group and from left to right across a period

D up a group and from right to left across a period

E down a group; the period position has no effect

6 Which one of the following atoms has the smallest radius?

A O

B F

C S

D Cl

E B

When a neutral atom gains or loses electrons, it becomes an ion.

Any atom losing electrons becomes a positive ion:

Na Na+ + e (neutral (sodium (electron)

sodium) ion)

Any atom gaining electrons becomes a negative ion:

F + e F (neutral (electron) (fluorine

fluorine) ion)

Ions

• Cations are positive and are formed by elements on the left side of the periodic chart.

• Anions are negative and are formed by elements on the right side of the periodic chart.

Ions

Atoms tend towards having complete outer shells of electrons (remember stability).

A full outer shell has:

• 2 electrons in the s subshell and• 6 electrons in the p subshell

Thus, atoms tend towards having a total of:

8 Valance Electrons This is the Noble Gas electron configuration.

Ions


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The outer most electrons in an atom are called valence electrons.

Valence Electrons

Number of valence electrons in neutral atoms:1 2 3 4 5 6 7 8

1 4

7 Which ion below has a noble gas electron configuration?

A Li 2+

B Be 2+

C B 2+

D C 2+

E N 2

Ionization Energy

• The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.

• The first ionization energy is that energy required to remove first electron.

Ca Ca+ + e

• The second ionization energy is that energy required to remove second electron, etc.

Ca+ Ca2+ + e

Trends in First Ionization Energies

• As one goes down a column, less energy is required to remove the first electron.

• For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus, "r" increases, so it is easier to remove the outermost electron


• Generally, as one goes across a row, it gets harder to remove an electron.

• As you go from left to right, Zeff increases, making it harder to remove an electron


However, there are two apparent discontinuities in this trend.


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• The first is between Groups 2 and 3.

• In this case the electron is removed from a p1 orbital rather than an s orbital.

• The electron removed is farther from nucleus, there is a small amount of repulsion by the s electrons, and the atom gains stability by having a full s subshell.


• The second is between Groups 15 and 16.

• The electron removed comes from doubly occupied p orbital.

• Repulsion from the other electron in the orbital aids in its removal.

• The atom gains stability by having a half full p orbital

Summary of Trends in Ionization Energy

• By adding one proton and one electron you are increasing q1 and q2 in Coulomb's Law

• Therefore, the force of attraction between the nucleus and outermost electrons is strengthened

• Thus, it takes more energy to remove the outermost electron

Ionization Energy increases left to right across a period • The size of orbitals increases significantly

• The distance between the nucleus and outer electrons increases

• So the force of attraction between the nucleus and outer electron is less

• Also, there is more shielding by inner electrons as you go down a group

• And more electronelectron repulsion with increasing number of electrons

Summary of Trends in Ionization Energy

Ionization Energy decreases going down a group

Ionization Energy

• It requires more energy to remove each successive electron.

ie: second ionization energy is greater than first, third ionization energy is greater than second, etc.

• When all valence electrons have been removed, the ionization energy takes a huge jump.

8 Which noble gas has the lowest first ionization energy? Give the atomic number.


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9 Of the following atoms, which has the largest first ionization energy?

A Br

B O

C C

D P

E I

10 Of the following elements, which has the largest first ionization energy?

A Na

B AlC Se

D ClE Br

11 Of the elements below, __________ has the largest first ionization energy.

A Li

B KC Rb

D Na

E H

Electronegativity

Recall that atoms gain stability when they have a full orbitals.

Fluorine has 7 valence electrons.

Neon has 8 valence electrons.

An atom of fluorine would be much more stable if it gained an electron, and became the fluoride ion (with the same electron configuration as neon).

Fluorine atoms "like" to acquire electrons

Electronegativity

Electronegativity is a measure of the ability of atoms in a molecule to attract electrons to themselves.

On the periodic chart, electronegativity increases as you go…

• …from left to right across a row.

• …from the bottom to the top of a column.

Electronegativities


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Electronegativities

In general we will not be concerned with

the electronegativites of transition metals.

12 The ability of an atom in a molecule to attract electrons is best quantified by the __________.

A paramagnetism

B diamagnetism

C electronegativity

D electron changetomass ratio

E first ionization potential

13 Electronegativity __________ from left to right within a period and __________ from top to bottom within a group.

A decreases, increases

B increases, increases

C increases, decreases

D stays the same, increases

E increases, stays the same

14 Of the atoms below, __________ is the most electronegative.

A BrB OC ClD NE F

15 Of the atoms below, __________ is the most electronegative.

A SiB ClC Rb

D Ca

E S

16 Of the atoms below, __________ is the least electronegative.

A Rb

B F

C Si

D Cl

E Ca


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17 Which of the elements below has the largest electronegativity?

A SiB Mg

C P

D SE Na

Metallic Character

The metallic character of an element is a measure of how loosely it holds onto its outer electrons.

For a metal to conduct electricity or heat, it needs to have electrons that are free to move through it, not tightly bound to a particular atom.

So the metallic character of an element is inversely related to its electronegativity. On the periodic chart, metallic character increases as you go…

• …from right to left across a row.

• …from the top to the bottom of a column.

Metallic Character Metals, Nonmetals, and Metalloids

• All atoms can be classified as metals, nonmetals or metalloids (also called semimetals)

• There are 7 semimetals: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, and Astatine (85)

18 Of the elements below, __________ is the most metallic.

A sodium

B barium

C magnesium

D calcium

E cesium

19 Of the elements below, __________ is the most metallic.

A NaB MgC AlD K

E Ar


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Ionic size depends upon:

• The nuclear charge.

• The number of electrons.

• The orbitals in which electrons reside.

But...

Cations are always smaller than their parent atom.

Anions are always larger than their parent atoms

Sizes of IonsCations are always smaller than their parent atom.

Sizes of Ions

3+ 3+

Neutral Lithium

Li

Lithium CationLi+

Free electron

e+

+

Sizes of Ions

Neutral Fluorine

F

Free electron

e

Fluorine AnionF

+

Anions are always larger than their parent atoms

9+ 9++

Ionic size depends upon:

• The nuclear charge.

• The number of electrons.

• The orbitals in which electrons reside.

Sizes of Ions

Sizes of Ions

• Cations (pink) are smaller than their parent atoms (gray).

• The outermost electron is removed and repulsions between electrons are reduced.

• The more electrons removed, the smaller the cation becomes

Sizes of Ions

• Anions (blue) are larger than their parent atoms (gray).

• Electrons are added and repulsions between electrons are increased.

• The more electrons added, the larger the anion becomes


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• Ions increase in size as you go down a group.

• This is due to increasing value of n (adding energy levels).

Sizes of Ions

• In an isoelectronic series, ions have the same number of electrons.

• Ionic size decreases with an increasing nuclear charge.

• The following ions are isoelectronic with Neon.• Z is the nuclear charge (atomic number)

Sizes of Ions

Z = 8 Z = 9 Z = 10 Z = 11 Z = 12 Z = 13

neon

20 _____ is isoelectronic with argon and _____ is isoelectronic with neon.

A Cl , F

B Cl , Cl +

C F+, F

D Ne , Kr+

E Ne , Ar +

21 Which of the following is an isoelectronic series?

A B 5, Sr 4, As3 , Te 2

B F, Cl, Br , I

C S, Cl, Ar, K

D Si2, P2, S2, Cl2

E O2 , F, Ne, Na+

22 Which of the following has the largest radius?

A Ar

B Kr

C Br

D Sr 2+

E Rb +

Documents

Periodic Trends Presentation student notes.notebookyooschem1314.pbworks.com/w/file/fetch/70754887/Periodic Trends … · B diamagnetism C electronegativity D electron changetomass