electronic structure of matter

1. QUARTER 2
MODULE 1

2. Lesson 1: Electronic
Structure of Matter

3. 1.History of the Atom
2.Quantum Mechanical Model of the
Atom
3.Electron Configuration

4. 1. On the basis of Rutherford’s model of an atom, which subatomic particle is present in the
nucleus of an atom?
A. neutron and electron C. proton and neutron
B. proton and electron D. proton only
2. If the first and second energy levels of an atom are full, then what would be the total
number of electrons in the atom?
A. 6 C. 10
B. 8 D. 18
3. Which atomic model is proposed by Schrodinger?
A. nuclear model C. raisin bread model
B. planetary model D. quantum mechanical model
4. Which electron transition results in the emission of energy?
A. 1s to 2s C. 3p to 4p
B. 2s to 2p D. 3p to 3s
5. The symbol “n” in the Bohr Theory of atomic structure refers to A. the energy of electron
B. the total energy of the atom
C. the orbit in which an electron is found
D. the number of electrons in an energy level

5. 12. A space around the nucleus where the electron is most likely to be found is called an
________________.
A. Atomic Orbital C. p-orbital
B. d-orbital D. s-orbital
13. How many types of orbitals are present in principal energy level four?
A. 6 C.12
B.8 D.16
14. Who discovered the electron?
A. Bohr C. J.J Thompson
B. Dalton D. Rutherford
15. Who proposed the planetary model of an atom?
A. Bohr C. J.J Thompson
B. Dalton D. Rutherford

6. HISTORY OF THE
ATOM

7. An atom is the
smallest unit of
matter

8. SCIENTIST WHO
PROPOSE THE
MODEL OF THE ATOM

9. JOHN DALTON

10. 1. JOHN DALTON
 English chemist.
 All elements are composed (made up) of atoms. It is impossible to
divide or destroy an atom.
 All atoms of the same elements are alike. (One atom of oxygen is
like another atom of oxygen.)
 Atoms of different elements are different. (An atom of oxygen is
different from an atom of hydrogen.)
 Atoms of different elements combine to form a compound. These
atoms have to be in definite whole number ratios. For example,
water is a compound made up of 2 atoms of hydrogen and 1 atom of
oxygen (a ratio of 2:1). Three atoms of hydrogen and 2 atoms of
oxygen cannot combine to make water.

11. DALTONS ATOMIC MODEL

12. JJ THOMPSON

13. 2. JJ Thompson
1. an English scientist.
2. He discovered the
electron when he was
experimenting with gas
discharge tubes. He
noticed a movement in a
tube. He called the
movement cathode rays.
The rays moved from the
negative end of the tube
to the positive end. He
realized that the rays
were made of negatively
charged particles –
electrons. His model of
the atom is known as
“Plum Pudding Model”

14. Summary of J.J Thompson’s Model of the
Atom:
-What did J.J. Thompson discover? -electron
-What is the charge of an electron? -
negatively charge
-What are cathode rays made of? -negatively
charge electrons

15. JJ Thompson Plum Pudding Model of the
Atom

16. ERNEST RUTHERFORD

17. GOLD FOIL EXPERIMENT

18. He also used special equipment to
shoot alpha particles (positively
charged particles) at the gold foil. Most
particles passed straight through the
foil like the foil was not there. Some
particles went straight back or were
deflected (went in another direction) as
if they had hit something.

19. Summary of Rutherford’s Model of the Atom:
-What is the charge of an alpha particle? -Positively
charged particles
-Why is Rutherford’s experiment called the gold
foil experiment? -He used a thin sheet of gold
 How did he know that an atom was mostly empty
space? - Most particles passed straight through the foil
like the foil was not there.
-What happened to the alpha particles as they
hit the gold foil? -Alpha particles were deflected
(went in another direction) as if they had hit
something.

20. ERNEST RUTHERFORD

21. NIELS BOHR
4. Niels Bohr (Early
1900s):
Niels Bohr was a Danish physicist. He
proposed a model of the atom that is similar
to the model of the solar system. The
electrons go around the nucleus like
planets orbit around the sun. Electrons in
each orbit have a definite energy, which
increases as the distance of the orbit from
the nucleus increases. Each energy level
can hold a certain number of electrons.
Level 1 can hold 2 electrons, Level 2 - 8
electrons, Level 3 - 18 electrons, and level
4 – 32 electrons. The energy of electrons
goes up from level 1 to other levels. When
electrons release (lose) energy they go
down a level. When electrons absorb (gain)
energy, they go to a higher level.

23. energy level of an electron

24. QUANTUM MECHANICAL MODEL OF THE ATOM
Lesson 1.2 Quantum Mechanical Model of the
Atom

 Bohr’s model of the atom was not sufficient to describe atoms with more than one electron. His idea
that electrons are found in definite orbits around the nucleus was rejected when three physicists developed
a better model of the atom—quantum mechanical model.
These were Louie de Broglie, Erwin Schrodinger, and Werner Karl Heisenberg.
 De Broglie proposed that the electron (which is thought of as a particle) could also be thought as a
wave.
 Schrodinger used this idea to develop a mathematical equation to describe the hydrogen atom.
 Heisenberg discovered that for a very small particle like the electron, its location cannot be exactly
known and how t is moving. This is called the uncertainty principle.

 Instead, these scientists believed that there is only a probability that the electron can be found in a
certain volume in space around the nucleus. This volume or region of space around the nucleus where the
electron is most likely to be found is called an atomic orbital

26. Orbitals have specific energy
values. They have particular
shapes and direction in space. The
s orbitals are spherical, and p
orbitals are dumbbell-shaped, as
shown in Figure 5

27. Table 1: Principal Energy Levels and
Sublevels of Electrons
Principal energy
level
Number of
Sublevels
Type of Sublevel and number of
orbitals
Maximum number of
electrons
1 1 1s (1 orbital) 2
2 2 2s (1 orbital)
2p (3 orbitals)
8
3 3 3s (1 orbital)
3p (3 orbitals)
3d ( 5 orbitals)
18
4 4 4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals)
32
5 5 5s (1 orbital)
5p (3 orbitals)
5d (5 orbitals)
5f ( 7 orbitals)
5g (9 orbitals)
50

28. ELECTRON CONFIGURATION
 The electron configuration is the standard notation used to describe the
electronic structure of an atom. Under the orbital approximation, we let each
electron occupy an orbital, which can be solved by a single wavefunction. In doing
so, we obtain three quantum numbers (n,l,ml), which are the same as the ones
obtained from solving the Schrodinger's equation for Bohr's hydrogen atom.
Hence, many of the rules that we use to describe the electron's address in the
hydrogen atom can also be used in systems involving multiple electrons.

29. . Principal Quantum Number (n)
 The principal quantum number n indicates the shell or energy level in which
the electron is found. The value of n can be set between 1 to n, where n is the
value of the outermost shell containing an electron. This quantum number can
only be positive, non-zero, and integer values. That is, n=1,2,3,4,..
For example, an Iodine atom has its outmost electrons in the 5p orbital.
Therefore, the principle quantum number for Iodine is 5.


30. 2. Orbital Angular Momentum Quantum Number (l)
 The orbital angular momentum quantum number, l, indicates the subshell of
the electron. You can also tell the shape of the atomic orbital with this quantum
number. An s subshell corresponds to l=0, a p subshell = 1, a d subshell = 2, a f
subshell = 3, and so forth. This quantum number can only be positive and
integer values, although it can take on a zero value. In general, for every value
of n, there are n values of l. Furthermore, the value of l ranges from 0 to n-1.
For example, if n=3, l=0,1,2. So in regards to the example used above, the l
values of Iodine for n = 5 are l = 0, 1, 2, 3, 4.

31. Magnetic Quantum Number (ml)
 The magnetic quantum number, ml, represents the orbitals of a given
subshell. For a given l, ml can range from -l to +l. A p subshell (l=1), for
instance, can have three orbitals corresponding to ml = -1, 0, +1. In other
words, it defines the px, py and pzorbitals of the p subshell. (However, the ml
numbers don't necessarily correspond to a given orbital. The fact that there are
three orbitals simply is indicative of the three orbitals of a p subshell.) In
general, for a given l, there are 2l+1 possible values for ml; and in a n principal
shell, there are n2 orbitals found in that energy level.
 Continuing on from out example from above, the ml values of Iodine are ml = -4,
-3, 2, -1, 0 1, 2, 3, 4. These arbitrarily correspond to the 5s, 5px, 5py, 5pz, 4dx
2-
y2, 4dz
2, 4dxy, 4dxz, and 4dyz orbitals

32. Spin Magnetic Quantum Number (ms)
 The spin magnetic quantum number can only have a value of either +1/2 or -1/2. The value
of 1/2 is the spin quantum number, s, which describes the electron's spin. Due to the spinning
of the electron, it generates a magnetic field. In general, an electron with a ms=+1/2 is called an
alpha electron, and one with a ms=-1/2 is called a beta electron. No two paired electrons can
have the same spin value.
 Out of these four quantum numbers, however, Bohr postulated that only the principal
quantum number, n, determines the energy of the electron. Therefore, the 3s orbital (l=0) has
the same energy as the 3p (l=1) and 3d (l=2) orbitals, regardless of a difference in l values. This
postulate, however, holds true only for Bohr's hydrogen atom or other hydrogen-like atoms.
 When dealing with multi-electron systems, we must consider the electron-electron
interactions. Hence, the previously described postulate breaks down in that the energy of the
electron is now determined by both the principal quantum number, n, and the orbital angular
momentum quantum number, l. Although the Schrodinger equation for manyelectron atoms is
extremely difficult to solve mathematically, we can still describe their electronic structures via
electron configurations.


33. What do the s, p, d and f orbitals look like?
The electron orbitals shown below represent a
volume of space within which an electron would have a
certain probability of being based on particular energy
states and atoms. For example, in a simple lowest-
energy state hydrogen atom, the electrons are most
likely to be found within a sphere around the nucleus of
an atom. In a higher energy state, the

34. The number of electrons each subshell can
hold

35. 1. Aufbau’s Principle is also known as the “building-up” principle, states that electron's
occupyorbitals in order of increasing energy. As soon as an energy level is filled with electrons,
any additional electron is thrown to the next outer or higher energy level.
It follows this mnemonic in filling up the orbital:
Example:

36. How to write an electron configuration:
Table 2: Distribution of electrons in the atomic orbitals of the first 10 elements.
Elements
atomic orbitals
Electron
Configuration
1s 2s 2px 2py 2pz 3s 3px 3py 3pz
H1 ↑ 1s1
He2 ↑↓ 1s2
Li3 ↑↓ ↑ 1s2
2s1
Be4 ↑↓ ↑↓ 1s2
2s2
B5 ↑↓ ↑↓ ↑ 1s2
2s2
2p1
C6 ↑↓ ↑↓ ↑ ↑ 1s2
2s2
2p2
N7 ↑↓ ↑↓ ↑ ↑ ↑ 1s2
2s2
2p3
O8 ↑↓ ↑↓ ↑↓ ↑ ↑ 1s2
2s2
2p4
F9 ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s2
2s2
2p5
Ne10 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s22s22p6

37. LEARNING ACTIVITIES

38. Activity 1: History of the Atom
Directions: Supply the needed information. Write your
answers directly on a separate sheet of paper.
1. What is the name of John Dalton’s theory?
_______________________
2. What are elements made of?
____________________________________
3. An atom of hydrogen and an atom of carbon are
_____________________
4. What are compounds made of?
__________________________________
5. The ratio of atoms in HCl is: a) 1:3 b) 2:1 c) 1:1

39. 1. What did J.J. Thompson discover?
______________________________
2. What is the charge of an electron?
______________________________
3. What are cathode rays made of?
________________________________
4. Why do electrons move from the negative end of the tube to the
positive end?
____________________________________________________
____
5. What was Thompson working with when he discovered the
cathode rays?
_____________________________________________
_________________

40. Lord Ernest Rutherford (1871 –
1937):
1. What is the charge of an alpha particle? _________________________
2. Why is Rutherford’s experiment called the gold foil experiment?
___________________________________________________________
3. How did he know that an atom was mostly empty space?
___________________________________________________________
4. What happened to the alpha particles as they hit the gold foil?
___________________________________________________________
4. How did he know that the nucleus was positively charged?

41.  Activity 2. Atomic Orbitals in Principal Energy Level
 Refer to Table 1: Principal Energy Levels and Sublevels of Electrons p.7 in answering the following questions.

 Example:
 Q: How many types of orbitals are in principal energy level three?
 A: There are three types of orbitals (s, p and d) in the principal energy level three.
 Q: How many atomic orbitals are in the highest sublevel of principal energy level three? A: There are five atomic
orbitals in the highest sublevel of the principal energy level three.

 1. How many types of orbitals are in the principal level two?
 ___________________________________________________________________
 2. How many atomic orbitals are in the highest sublevel of the principal energy level two?
 _______________________________________________________________
 3. How many types of orbitals are in the principal level four?


 4. How many atomic orbitals are in the highest sublevel of the principal energy level four?
 _______________________________________________________________

42. Elemen
ts
Orbital Diagram Electron Configuration
Be4
↑↓ ↑↓
1s 2s
1s2 2s2
Mg12
Al13
Si14
P15
Activity 3. Electron Configuration
Write the electron configuration of the elements and show the orbital diagram using the Hund’s
Rule and Pauli’s Principle