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BICOL UNIVERSITY College of Science

Department of Chemistry

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CHEM 1 LECTURE HANDOUT 3 Ver. 1.2 α 20110814 (mbab ver) GENERAL CHEMISTRY Php14 michaelvmontealegre

2 http://buchem.weebly.com michaelvmontealegre

periodic table & quantum numbers 3

An atom is composed of three types of subatomic particles: the proton, neutron, and electron.

Particle Mass (g) Charge

Proton 1.6727 x 10-24 +1 Neutron 1.6750 x 10-24 0 Electron 9.110 x 10-28 -1

Here, charge is given in multiple of 1.602 x10 -19 coulombs.

Protons and neutrons have similar masses and electrons are much lighter (over 1,000 times lighter).

Protons and electrons have equal and opposite charges while neutrons have no charge.

GROUP OR

FAMILY

PER

IOD


We have the following simple picture of the atom.

The atom is comprised of a positively charged nucleus composed of

protons and neutrons. This small nucleus is surrounded by orbiting

electrons. Because the protons and neutrons are so much more massive

than the electrons, virtually all the mass of the atom is located in the

nucleus. The light negatively charged electrons move around in an orbit in

the space around the nucleus.

We use the following symbol to describe the atom:

A= Z + N, where N is the number of neutrons.

If you add or subtract a proton from the nucleus, you create a new

element.

If you add or subtract a neutron from the nucleus, you create a new

isotope of the same element you started with.


In a neutral atom, the number of positively charged protons in the nucleus is equal to the number of orbiting electrons.

When we add neutrons to the nucleus of we can make the isotopes of

hydrogen. Here are three common isotopes of hydrogen.

If we add a proton to the hydrogen nucleus we would get helium (a different element). Here are two common isotopes of helium.

Electron Orbitals (s, p, d, f) An orbital is a region within an energy level where there is a probability of finding an electron. Orbital shapes are defined as the surface that contains about 90% of the total electron probability. The labels s, p, d and f have their origins in the words ‘sharp’, ‘principal’, ‘diffuse’ and ‘fundamental’. These originally referred to the characteristics of the lines observed in the atomic spectrum of hydrogen

s orbital shape The s orbital has a spherical shape centered on the origin of the three axes in space.


p orbital shape There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.

Degenerate orbitals possess the same energy. Orbitals that differ only in their orientation in space, such as the 2px, 2py, and 2pz orbitals, are therefore degenerate.

d orbital shapes Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells” …and a “dumbell with a donut”!


f orbitals


Electrons enter and fill orbitals according to four rules:

1. Pauli Exclusion Principle Orbitals can contain a maximum of two electrons which must be of opposite spin.

He (Z = 2): 1s2

Incorrect:

Correct:

2. Aufbau or Build-up Principle

Electrons enter and fill lower energy orbitals before higher energy orbitals.

Maximum number of electrons in each subshell: s=2, p=6, d=10, f=14. A superscript number after the letter of the subshell shows how many electrons occupy that subshell.

Shown is a sample orbital notation for sulfur where the orbitals are arranged based on their energies.

Writing the Electronic configuration H (Z=1)

1s1 Orbital

(1s)

Number of electrons in orbital


3. Hund's Rule When there are degenerate (equal energy) orbitals available, electrons will enter the orbitals one-at-a-time to maximize degeneracy, and only when all the orbitals are half filled will pairing-up occur. This is the rule of maximum multiplicity.

For degenerate orbitals, the electrons are distributed following Hund’s Rule. The electrons in the 2p orbitals on carbon can be as follows.

C (Z = 6): 1s2 2s

2 2p

2

Writing the Noble Gas Notation and Orbital Notation

Element Electronic

Configuration Noble Gas Notation

Orbital Notation

He (Z=2)

1s2 -

N (Z=7)

1s2 2s2 2p3 [He] 2s2 2p3


When we get to N (atomic number, Z = 7), we have to put one electron into each of the three degenerate 2p orbitals.

N (Z = 7): 1s2 2s

2 2p

3

Because each orbital in this subshell now contains one electron, the next electron added to the subshell must have the opposite spin quantum number, thereby filling one of the 2p orbitals.

The ninth electron fills a second orbital in this subshell.

F (Z = 9): 1s2 2s

2 2p

5

1. The neon atom has 10 electrons.

2 electrons in the s subshell of the first energy level: 1s2

8 electrons in second energy level, made up of 2 electrons in an s

subshell and 6 electrons in a p subshell: 2s2 2p

6

Its subshell electron configuration is 1s2 2s

2 2p

6

2. The argon atom has 18 electrons.

2 electrons in the s sub-shell of the first energy level: 1s2

8 electrons in the second energy level made up of 2 electrons in an

s subshell and 6 electrons in a p subshell: 2s2 2p

6

8 electrons in the third energy level made up of 2 electrons in an s

subshell and 6 electrons in a p subshell: 3s2 3p

6

Its subshell electron configuration is: 1s2 2s

2 2p

6 3s

2 3p

6

Electron Configuration of Ions

Atom Z No.

Electrons = Z

Subshell Electron Configuration

Ion No. Electrons = Z - charge

Subshell Electron Configuration

O 8 8 1s2 2s

2 2p

4 O

2- 8 - (-2) = 10 1s

2 2s

2 2p

6

F 9 9 1s2 2s

2 2p

5 F

- 9 - (-1) = 10 1s

2 2s

2 2p

6

Na 11 11 1s2 2s

2 2p

6 3s

1 Na

+ 11 - 1 = 10 1s

2 2s

2 2p

6


Mg 12 12 1s2 2s

2 2p

6 3s

2 Mg

2+ 12 - 2 = 10 1s

2 2s

2 2p

6

P 15 15 1s2 2s

2 2p

6 3s

2 3p

3 P

3- 15-(-3)=18 1s

2 2s

2 2p

6 3s

2 3p

6

Cl 17 17 1s2 2s

2 2p

6 3s

2 3p

5 Cl

- 17-(-1)=18 1s

2 2s

2 2p

6 3s

2 3p

6

Special Electronic Configurations

When two electrons occupy the same orbital, they not only have different

spins (Pauli exclusion principle), the pairing raises the energy slightly. On

the other hand, a half filled subshell and a full filled subshell lower the

energy, gaining some stability. Bearing this in mind, you will be able to

understand why we have the following special electronic configurations.

Cr [Ar]4s1 3d

5 <=All s and d subshells are half full

Cu [Ar]4s1 3d

10 <=Prefers a filled d subshell, leaving s with 1

Nb [Kr]5s1 4d

4 <=5s and 4d energy levels are close

Mo [Kr]5s1 4d

5 similar to Cr above

Tc [Kr]5s2 4d

5 (not special, but think of Hund's rule)

Ru [Kr]5s1 4d

7 <= Only 1 5s electron

Rh [Kr]5s1 4d

8 <= in both

Pd [Kr]5s0 4d

10 <= Note filled 4d and empty 5s

Ag [Kr]5s1 4d

10 <= partial filled 5s, but filled d

Exceptions to Electron Configuration Trends

Period 4: Period 5: Chromium: Z:24 [Ar] 3d

54s

1 Niobium: Z:41 [Kr] 5s

1 4d

4

Copper: Z:27 [Ar] 3d104s1 Molybdenum: Z:42 [Kr] 5s1 4d5

Ruthenium: Z:44 [Kr] 5s1 4d

7

Rhodium: Z:45 [Kr] 5s1 4d

8

Palladium: Z:46 [Kr] 4d10

Silver: Z:47 [Kr] 5s1 4d

10

Period 6: Period 7:

Lanthanum: Z:57 [Xe] 6s2 5d

1 Actinium: Z:89 [Rn] 7s

2 6d

1

Cerium: Z:58 [Xe] 6s2 4f

1 5d

1 Thorium: Z:90 [Rn] 7s

2 6d

2

Gadolinium: Z:64 [Xe] 6s2 4f

7 5d

1 Protactium: Z:91 [Rn] 7s

2 5f

2 6d

1

Platinum: Z:78 [Xe] 6s1 4f

14 5d

9 Uranium: Z:92 [Rn] 7s

2 5f

3 6d

1

Gold: Z:79 [Xe] 6s2 4f

14 5d

10 Neptunium: Z:93 [Rn] 7s

2 5f

4 6d

1

Curium: Z:96 [Rn] 7s2 5f

2 6d

1

Lawrencium: Z:103 [Rn] 7s2 5f

14 7p

1


Solid lines surrounding elements designate filled (d10 or f 14) or half-filled (d6 or f 7) subshells.

Dashed lines surrounding elements designate irregularities in sequential orbital filling, which is also

found within some of the solid lines.

4. Madelung's Rule

Orbitals fill with electrons as n + l, where n is the principal quantum number and l is the subsidiary quantum number. This rule 'explains' why the 4s orbital has a lower energy than the 3d orbital, and it gives the periodic table its characteristic appearance.


1. Complete the table below

Element Electronic

Configuration Orbital notation

Noble gas

notation

Neon

____ ____ ____ ____ ____

1s 2s 2p

[He]2s2p6

Ca2+

Noble gas notation is a short electron configuration replacing certain notation

with the noble gas symbol in square brackets - [He] = 1s2, [Ne] = 1s

2 2s

2 2p

6,…

2. For an element with atomic number 20, which is the last or highest

occupied subshell of atomic orbitals?

1s 2s 2p 3s 3p 3d 4s 5s

3. Choose the electronic configuration for palladium, Pd (Z = 46).

a [Kr]5s1 4d

7

b [Kr]5s1 4d

8

c [Kr]5s0 4d

10

d [Kr]5s1 4d

10


1. If the valence electron is less than or equal to the main energy level, the element is METAL. Example: Li-3 1s22s1 (1<2)

Be-4 1s22s2 (2=2)

2. If the valence electron is greater by 1 or equal to the main energy element is a METALLOID. Example: B-5 1s22s22p1 (3>2 by 1)

Al-13 1s22s22p63s23p1 (3=3)

3. If the valence electron is greater by 2 or more than the main energy level, the element is a NONMETAL. Example: C-6 1s22s22p2 (4>2 by 2)

O-8 1s22s22p4 (6>2 by 4)

4. If the valence electron is equal to 8, the element is a NOBLE or INERT GAS. Example: Ne-10 1s22s22p6

Valence Electrons = Group Number

Valence Electrons = Group Number-10

(except He = 2)

d and f block Valence Electrons = commonly 1 or 2

Valence electrons are electrons in the outmost

shell (energy level).


Half of the distance between nuclei in covalently bonded diatomic

molecule

Radius decreases across a period

Increased effective nuclear charge due to decreased

shielding

Radius increases down a group

Each row on the periodic table adds a “shell” or energy

level to the atom


Cations

Positively charged ions formed when an atom of a metal loses

one or more electrons

Smaller than the corresponding atom

Anions

Negatively charged ions formed when nonmetallic atoms gain

one or more electrons


Larger than the corresponding atom

(or Ionization Potential)

Definition: the energy required to remove an electron from an atom


Tends to increase across a period

As radius decreases across a period, the electron you are

removing is closer to the nucleus and harder to remove

Tends to decrease down a group

Outer electrons are farther from the nucleus and easier to

remove

The energy required to remove one valence electron is the first ionization energy, the second ionization energy is the energy required to remove a second valence electron, and so on.

1st I.E.: Na(g) →Na+(g)+ e-

2nd I.E.: Na(g) →Na2+(g) + 2e-

Ionization Energies of certain elements (1st IE, 2nd IE, etc) Element 1st 2nd 3rd 4th 5th 6th 7th

Na 496 4562

Mg 738 1451 7733

Al 577 1817 2745 11580

Si 786 1577 3232 4356 16090

P 1060 1903 2912 4957 6274 21270

S 999.6 2251 3361 4564 7013 8496 27110

Cl 1256 2297 3822 5158 6542 9362 11020

Ar 1520 2666 3931 5771 7238 8781 12000


These are the ionization energies for the period three elements. Notice how Na after in the second I.E, Mg in the third I.E., Al in the fourth I.E., and so on, all have a huge increase in energy compared to the proceeding one. This occurs because the proceeding configuration was in a stable octet formation; therefore it requires a much larger amount of energy to ionize.

Electron affinity (E.A.) is the energy change that occurs when an electron is added to a gaseous atom. Electron affinity can further be defined as the enthalpy change that results from the addition of an electron to a gaseous atom. It can be either positive or negative value. The greater the negative value, the more stable the anion is.

X(g) + e- ---> X- + Energy (Exothermic) The electron affinity is positive

X(g) + e- + Energy ---> X- (Endothermic) The electron affinity is negative

It is more difficult to come up with trends that describe the electron affinity. Generally, the elements on the right side of the periodic table will have large negative electron affinity. The electron affinities will become less negative as you go from the top to the bottom of the periodic table. However, Nitrogen, Oxygen, and Fluorine do not follow this trend. The noble gas electron configuration will be close to zero because they will not easily gain electrons.


Definition: A measure of the ability of an atom in a chemical compound to

attract electrons

o Electronegativity tends to increase across a period

o As radius decreases, electrons get closer to the bonding

atom’s nucleus

o Electronegativity tends to decrease down a group or remain the

same

o As radius increases, electrons are farther from the bonding

atom’s nucleus


1. Which member of each pair has the greater negative electron

affinity?

a. Se or Br

b. Si or Al

c. Na or Cl

2. Which member of each pair has the larger radius?

a. Ca or Mg

b. Al or P

c. F- or F

3. Which member of each pair has the higher ionization energy?

a. He or Ne

b. Na or Mg

c. Ca or Ca2+

4. Which of the alkali elements has the highest ionization energy

(IE): Li, Na, K, Rb, or Cs?

5. Which of the following elements (Cs, C, O, Cl, or F) releases the

most energy when acquiring an electron? In other words, which

has the largest absolute value of electron affinity?


Each electron has a set of four numbers, called quantum numbers, that

specify it completely. The four quantum numbers are:

n : principal quantum number

l : azimuthal quantum number

ml : magnetic quantum number

ms : spin quantum number

The principal quantum number (n)

describes the size of the orbital.

Orbitals for which n=2 are larger than

those for which n=1, for example.

Because they have opposite electrical

charges, electrons are attracted to the

nucleus of the atom. Energy must

therefore be absorbed to excite an

electron from an orbital in which the

electron is close to the nucleus (n=1)

into an orbital in which it is further

from the nucleus (n=2). The principal

quantum number therefore indirectly

describes the energy of an orbital.

The principal quantum number is always a positive integer and tells us the

energy level or shell that the electron is found in.

The maximum number of subshells permitted for a particular shell is equal

to n2.

The maximum number of electrons permitted in a particular shell is equal

to 2n2.


*Shells are designated K for 1st energy level, L for 2nd, …

For each value of n, the angular momentum (azimuthal) quantum

number (l) for an electron can have integral values from zero to n-1 it

cannot be as large as n. The azimuthal quantum number tells us which

subshell the electron is found in, and therefore it tells us the shape of the

orbital. The number of orbitals permitted for a particular subshell is equal

to 2l + 1.


Energy Levels, Orbitals, Electrons

Energy

Level

(n)

Orbital type in

the energy level

(types = n)

Number of

orbitals

Number of

Electrons

Number of

electrons per

Energy level

(2n2

)

1 S 1 2 2

2 S

p

1

3

2

6

8

3 S

p

d

1

3

5

2

6

10

18

4 S

p

d

f

1

3

5

7

2

6

10

14

32

The magnetic quantum number (ml) is an integer with a value from -l to

+l. It defines the orientation of an orbital in the space around the nucleus

of an atom

It is not always possible to associate a value of ml with a particular orbital.

You do not assign -1 for the p subshell as the px orbital. It can either be

one of the orbitals.

value of l subshell values of ml possible orbitals

0 s 0 s

1 p -1, 0, 1 px, py, pz

2 d -2, -1, 0, 1, 2 dxy, dxz, dyz, dx2

-y2, dz

2


The spin quantum number, ms, tells us the spin of the electron.

Possibilities for electron spin: + ½, - ½

Permitted Values of the Quantum Numbers through n=4


The argon atom has 18 electrons.

The quantum numbers for each of the 18 electrons is shown below:

electron n (shell) l (subshell) ml (possible orbital)* ms

1 1 (K) 0 (s) 0 (1s) -½

2 1 (K) 0 (s) 0 (1s) +½

3 2 (L) 0 (s) 0 (2s) -½

4 2 (L) 0 (s) 0 (2s) +½

5 2 (L) 1 (p) -1 (2px) -½

6 2 (L) 1 (p) -1 (2px) +½

7 2 (L) 1 (p) 0 (2py) -½

8 2 (L) 1 (p) 0 (2py) +½

9 2 (L) 1 (p) +1 (2pz) -½

10 2 (L) 1 (p) +1 (2pz) +½

11 3 (M) 0 (s) 0 (3s) -½

12 3 (M) 0 (s) 0 (3s) +½

13 3 (M) 1 (p) -1 (3px) -½

14 3 (M) 1 (p) -1 (3px) +½

15 3 (M) 1 (p) 0 (3py) -½

16 3 (M) 1 (p) 0 (3py) +½

17 3 (M) 1 (p) +1 (3pz) -½

18 3 (M) 1 (p) +1 (3pz) +½

*for degenerate orbitals such as px, py and pz, the assignments of -1, 0 and

+1 are arbitrary. That is -1 can be any of the orbitals. The convention is

used only to demonstrate filling up of orbitals. Any electron may occupy

any of the degenerate orbitals as long as they are following the rules for

filling up orbitals.


1. Describe the allowed combinations of the n, l, and m quantum

numbers when n = 3.

2. How many possible orbitals are there if n = 3?

3. How many possible orbitals are there in the subshell [n=5, l=4]?

4. How many electrons can be accommodated in the subshell 4f?

5. How many atomic orbitals are there for the subshell with [n = 3, l

= 2]?

6. What is the symbol representing the set of orbitals in the subshell

with [n = 3, l = 2]?

7. Consider the following sets of quantum numbers. Which ones

could not occur? For the valid sets, identify the orbital involved.

a. 3, 1, 0, +½

b. 1, 1, 0, -½

c. 2, 0, 0, +½

d. 4, 3, 2, +½

e. 2, 1, 0, 0


Book References: Chang, Raymond. Chemistry. 9

th edn. Digital Content Manager. 2007.

Housecroft, Catherine E. and Edwin Constable. Chemistry 3rd

edn. Pearson

Education International. England. 2006.

Silberberg, Martin S. Chemistry: The molecular nature of matter and change. 5th

edn. The McGraw-Hill Companies, Inc. New York, NY USA. 2009

Tro, Nivaldo J. Chemistry: A Molecular Approach. 2nd edn. Pearson Education.

Inc. Upper Saddle River, New Jersey. 2011

Internet Resources: http://antoine.frostburg.edu/chem/senese/101/measurement/

http://www.sciencegeek.net/Chemistry/index.shtml

http://www.visionlearning.com/library/cat_view.php?cid=1

http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html

http://www.ausetute.com.au/index.html

http://www.science.uwaterloo.ca/~cchieh/cact/c120/quantnum.html

http://www.meta-synthesis.com/webbook/34_qn/qn_to_pt.html

http://www.grandinetti.org/Teaching/Chem121/Lectures/TheAtom

http://antoine.frostburg.edu/chem/senese/101/measurement/

http://www.sciencegeek.net/Chemistry/index.shtml

http://www.visionlearning.com/library/cat_view.php?cid=1

http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html

http://www.ausetute.com.au/index.html

http://www.science.uwaterloo.ca/~cchieh/cact/c120/quantnum.html

http://www.meta-synthesis.com/webbook/34_qn/qn_to_pt.html

http://www.grandinetti.org/Teaching/Chem121/Lectures/TheAtom

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