c h e m i s t r y 1 a : c h a p t e r 8 p a g e | chapter...

C h e m i s t r y 1 A : C h a p t e r 8 P a g e | 1

Chapter 8: Periodic Properties of the Elements

Homework: Read Chapter 8. Work out sample/practice exercises

Check for the MasteringChemistry.com assignment and complete before due date

The Periodic Table:

1869 Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) classified known

elements by organizing similar physical and chemical properties. Mendeleev’s periodic

table was an attempt to organize the known data at the time in a way that made sense.

Elements were arranged by increasing atomic mass and grouped together by chemical

reactivity. Where problems with mass positions occurred ( Te and I), he re-ordered by

other properties. Several holes led to predictions of elements and their properties that

were not yet discovered “eka-aluminum” (Ga) and “eka-silicon”(Ge).

1913 *Henry Moseley improved the periodic table by ordering the elements by

increasing atomic number. More “holes” were found, which led to the discovery of more

elements and the family of noble gases.

The periodic table gives a great amount of information in an organized manner. Vertical

columns are called groups or families. If you are aware of the properties of a couple

elements in a group, you can make a good guess at the properties of the other elements in

the same group. Periods are the horizontal rows in the periodic table. Many patterns can

be seen or predicted following periods and groups.

Electron Configurations and Orbital Diagrams:

A. Energy Levels:


i. In the Bohr Atom (one electron systems) there is a one-to-one correspondence

between an orbit and its energy level (En = -B/n2). Example: the 3s, 3p, 3d

orbitals in Hydrogen all have the same energy.

ii. In the Quantum Mechanical version of the atom, the energy level of multi-

electron atoms depend on both the size (1, 2, 3, 4…) and shape (s, p, d, f).

B. Ground State: Filling orbitals of an atom where electrons go into the lowest

energy orbitals first.

i. Pauli Exclusion Principle, no two electrons in an atom may have all four

quantum numbers alike.

ii. Hund's Rule, when filling degenerate energy levels, each orbital fills one

electron, spin unpaired, before any orbital fills with two electrons.

C. Aufbau Principle: Electronic Configuration by energy

i. 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2 5f146d107p6

ii. Isoelectronic Same Electronic Configuration


Watch for exceptions:

Electron Configurations (complete/short) Core electrons

Valence electrons Pseudocore electrons

Ground state Excited state

By size By energy

Pauli’s rule Hund’s Rule

Paramagnetic Diamagnetic

Quantum numbers (m, l ml, ms)


Quantum Numbers:

A. Each electron in an atom is defined by four Quantum Numbers.

The principle quantum number, n, is related to the size (90% probability that the

electron is within a given radius)

B. The values of

the

azimuthal quantum number, , determines the shape of the orbital. In the

designation of an orbital, this quantum number is represented by a letter.

C. The magnetic quantum number, ml, identifies the three dimensional orientation in

space. For an s orbital it is 0, for p it can be -1, 0, +1, d has 5 orientations, f has 7

Three quantum numbers are required to specify an orbital: principle, azimuthal and

magnetic.

An orbital is a place in an atom to hold electrons. An orbital may contain a

maximum of two electrons

D. The fourth quantum number is the electron spin, ms. Values are ½. Stern-Gerlach

experiment split a beam of silver atoms (47 electrons-odd number-paramagnetic) in

two by a magnetic field. As electrons spin they generate a magnetic field. About

have half the electrons point “north” (spin up) and others point “south” (spin down)

Name Symbol Values Significance

principle n 1, 2, 3, ... size

azimuthal 0, 1, ...(n-1) shape

magnetic ml 0, 1, 2,.... orientation

spin ms ½ electron spin

Value Letter Shape

0 s spherical

1 p dumbbell

2 d four-lobed

3 f eight-lobed


Example 1: Predict the ground state short electron configuration for each. Identify electrons

core, [valence], pseudocore. Draw the orbital diagram for the outershell electrons,

identify if species is diamagnetic or paramagnetic, write out the quantum numbers

for each of the outershell electrons.

Ag Ag+1

Ga Ga+3

Fe Fe+2 Fe+3

Sn Sn+2 Sn+4

P P-3 P+3 P+5

S S-2

Ce Ce +2 Ce+3 Ce+4


Coulombs Law: F = kQ1Q2/d2 Since Energy is force times distance, E = kQ1Q2/d

Coulombs Law describes attractions and repulsions between charged particles.

Attraction is stronger as atomic sizes decrease and charge differences increase.

Effective Nuclear Charge:

Negatively charged electrons are attracted to the positively charged nucleus and

repelled by other electrons in the atom.

The force of attraction depends on the magnitude of the net nuclear charge

acting on an electron and the average distance between the nucleus and the

electron.

We can estimate the net attraction and average environment of a single electron

through Zeff (effective nuclear charge), which is always smaller than the total

charge of the nucleus. Zeff = Ztotal - Sscreening constant

S is generally close to the # of core electrons, (i.e. for Na, 10 core electrons, 1

valence electron. The Zeff = 11-10 = +1 for the valence electron 3s1)


Penetration:

• The closer an electron is to the

nucleus, the more attractions it

experiences.

• The degree of penetration is related

to the orbitals radial distribution

function

• The radial distribution function

shows that the 2s orbital penetrates

more than the 2p

• The weaker penetration of the 2p

sublevel means that electrons in

the 2p sublevel are more shielded

from the attractive force of the

nucleus

• The deeper penetration of the 2s

electrons means electrons in the 2s

sublevel experience a greater

attractive force to the nucleus and

are not shielded as effectively

• Penetration causes the energies of

sublevels in the same principal

level to not be degenerate (2s and 2p are different energies)

• In the 4th and 5th principle levels, the effects of penetration cause the s orbital to be

lower in energy than d orbitals of the previous principal level (4s is lower than 3d)

• The energy separations between one set of orbitals and the next become smaller

beyond the 4s so the ordering can vary among elements causing variations

(exceptions) in the electron configurations of the transition metals and their ions

Periodic Trends:

Properties of the elements follow a periodic pattern: Same column have similar

properties and in a period a pattern repeats. This is explained with the

quantum-mechanical model because the number of valence electrons and

types of orbitals they occupy are periodic.


Size (atomic and ionic radii):

Atomic radii increase from right to left; top to bottom of periodic table.

Zeff is the effective nuclear charge. Zeff = Zactual – electron shielding

This is the charge felt by the outer electrons that are shielded from the full power

of the positive nuclear charge

Along a period the Zeff increases left to right pulling in electrons closer to the nucleus

and causing the atoms to decrease in size.

Transition metals in the same d block are roughly the same size

Vertically, the size of the orbitals (quantum number n) increases from top to bottom


Ionic Radii:

Cations lose electrons and are therefore smaller than the original atom

Anions gain electrons and are larger than the original atom

Isoelectronic series (all have the same number of electrons, same electron

configuration) the size increases as the charge of the nuclei decreases.

(smallest Sr+2, Rb+1, Kr, Br-1, Se-2 largest)


First Ionization Energy, Ei:

Energy required to remove the outermost ground state electron, endothermic

Ionization Energy decreases from right to left; top to bottom of periodic table.

The small nonmetals require the highest ionization energy, they do not want to

lose electrons

Large metals have lowest ionization energy, they want to lose electrons and

become positively charged cations.

Minor irregularities occur

a. Ei of Be is larger than B and that of Mg is larger than Al. An

explanation is that Be and Mg lose an s2 electron while B and Al are

losing the p1 electron. An s electron spends more time closer to the

nucleus and is therefore harder to remove. Additionally, the p electrons

are shielded somewhat by the s electrons and feel a smaller Zeff.

b. Ei of N is larger than O. The explanation lies in the difference between

losing an electron from a half filled orbital verses a filled orbital.

Oxygen’s last filled electron is 2p4 . Electrons repel each other and

when electrons are forced to share space in a filled orbital they are

slightly higher in energy, so it is slightly easier to remove one giving O

a smaller ionization energy compared to N.


Higher Ionization Energies:

Energy required removing the second, third

of even more electrons from an atom

Larger amounts of energy are required to

remove each successive electron. It is

relatively easier to remove from partially

filled valence shell and harder to remove

from filled d shells or core electrons.

Electron Affinity, Eea:

Energy given away when adding an

electron, exothermic:

The greatest negative value (most

preferred) electron affinity is for F.

Small nonmetals.

Ignore noble gases.

-Eea generally decreases from right to

left; top to bottom of periodic table.


Electronegativity: (Chapter 9: page 394-395)

The ability to attract electrons toward the atom

Electronegativity is important in the covalent bonding or ionic transfer of

electrons in molecules and compounds. (Ch 9: Lewis Structure, Valence Shell

Electron Pair Repulsion, determining polarity of a molecule)

Increases from left to right; bottom to top of periodic table. Small nonmetals are

much better at attracting electrons.

Ignore the noble gases as most do not attract electrons (except xenon which may

make a few compounds such as XeOF4, XeCl2 or XeF4)

http://www.green-planet-solar-energy.com/electronegativity-values.html

http://www.green-planet-solar-energy.com/electronegativity-values.html


Metallic Character: Increases from right to left; top to bottom of periodic table

Octet Rule:

Main group elements tend to undergo reactions that leave them with 8 outer shell

electrons, obtaining the noble gas configuration.

Chemistry by Group: Alkali Metals:

Metallic, soft enough to cut with a knife, silver color, low melting points,

malleable, conductive, reactive and must be stored under oil to prevent reaction

with air and moisture.


Reaction of alkali metal (M) with a halogen (X) MX

Reactions:

2 M + X2

2 M + H2

6 M + N2

4 M + O2

2 M + 2 H2O

Alkaline Earth Metals: Metallic, silver color, malleable, conductive, can lose 2 electrons easily

causing them to be powerful reducing agents.

Reactions:

M + X2 (X is a halogen)

M + H2

2 M + O2

M + 2 H2O

Halogens:

Nonmetals, diatomic molecules, high electron affinities (tendency to gain

electrons), powerful oxidizing agents

Reactions: (X is a halogen)

2 M + n X2 2MXn

H2 + X2


Noble Gases: Unreactive nonmetals, low melting and boiling points, colorless, odorless,

filled core electron configuration.

He and Ne undergo no known reactions. Argon is known to form HArF. Kr

and Xe may react with fluorine (XeF2, XeF4, XeF6, XeOF4)

Problems:

1. Use the concepts of effective nuclear charge, shielding, and n value of the valence orbital to

explain the trends in atomic radius as you (a) move across the periodic table, (b) move down the

periodic table.

2. Is the order of electron removal upon ionization simply the reverse of electron addition upon

filling? Why or why not. Complete the short electron configurations for thalium…Tl, Tl+1, Tl+3

3. Arrange the elements as follows: C, Mg, He, Sr, O, Fr

a) Increasing metallic character

b) Increasing atomic size

c) Increasing Ionization energy

d) Increasing electronegativity

4. Identify 2 cations and 2 anions that are isoelectronic with Xe. Place them in order of increasing

atomic radii.

5. Estimate the normal melting point of Br from the given melting points of atoms in the same

group… F = -219°C, Cl = -101°C, I = 114°C

6. Estimate the density of Kr from the given density of atoms in the same group at STP…

Ne = 0.90 g/L, Ar= 1.78 g/L, Xe = 5.86 g/L

7. Use Coulomb’s Law to arrange the ionic compounds by expected increasing normal melting

points. BaO, CsI, NaCl, Fe2O3

c h e m i s t r y 1 a : c h a p t e r 8 p a g e | chapter...

Documents