Module 1: Electronic Structure of Matter Objectives: Know atom and its sub - particles determine the characteristics colors that metal salts emit what is atom? atom is the basic unit of chemical element it composes three subatomic particle proton with a positively electric charge electron with a negatively electric charge neutron no electric charge What minerals produce the color in fireworks? Mineral elements provide color in fireworks. Barium produces bright greens; strontium yields deep reds;' copper produces blues/ and sodium yields yellow. other colors can be made by mixing elements; strontium and sodium produce brilliant orange; titanium, zirconium, and magnesium alloys make silvery white; copper and strontium make lavender. gold sparks are produced by iron fillings and small pieces of charcoal. bright flashes and loud bangs come from aluminum powder.

1. Module 1: Electronic
Structure of Matter

2. Objectives:
• Know atom and its sub – particles.
• Determine the characteristic colors that metal salts emit

3. What is atom?

4. Atom – is the basic unit of a
chemical element.

5. What’s Inside Atom?

6. 3 Sub – Atomic Particles Of Atom
a. Electron – negatively charge
b. Proton - positively charge
c. Neutron – has no electric charge

9. • Mineral elements provide the color in fireworks. Barium
produces bright greens; strontium yields deep reds;
copper produces blues; and sodium yields yellow. Other
colors can be made by mixing elements: strontium and
sodium produce brilliant orange; titanium, zirconium,
and magnesium alloys make silvery white; copper and
strontium make lavender. Gold sparks are produced by
iron filings and small pieces of charcoal. Bright flashes
and loud bangs come from aluminum powder.

10. Red: Sr - Strontium
Orange: Sr - Strontium, Na - Sodium
Yellow: Na - Sodium
Green: Ba - Barium
Blue: Cu - Copper
Purple: Sr - Strontium, Cu - Copper
Greys and White: Ti - Titanium, Zr - Zirconium, Mg
- Magnesium

11. Flame Test

12. Metal Salt Tested
Element
Producing Color
Color of the Flame
Boric Acid Boron Green
Calcium Chloride Calcium Yellow
Sodium Chloride Sodium Orange
Potassium
Chloride
Potassium Light Violet
Cooper (II) Sulfate Copper Blue green
Flame Test
1. Why do you think are there different colors
emitted?
Answer: Metal salts emitted different colors because
of the absorption of heat from the flame.
2. What particles in the heated compounds are
responsible for the production of the colored light?
Answer: The outermost particle in the metallic element
are responsible for the production of colored light.
3. How did the scientists explain the relationship
between the colors observed and the structure of the
atom?
Answer: The colors of observed is an indication that
definite energy transformations occurs inside the atom
emitting light. It follows that electrons must occupy
orbits of fixed energy.

13. Flame Test
• Is a form of qualitative analysis that is used to visually determine the
identity of an unknown metal or metalloid ion based on the color
emission.
• A distinctive color is emitted because the heat of the flame excites
the electrons of the metal ions, causing them to emit visible light.
• The outermost

14. Bohr’s Atomic Model Theory

15. S Y M B O
FREE

16. Metal Salt Tested Element Producing Color Color of the Flame
Boric Acid Boron Green
Calcium Chloride Calcium Yellow
Sodium Chloride Sodium Orange
Potassium Chloride Potassium Light Violet
Cooper (II) Sulfate Copper Blue green

17. • You have observed that each of the substances you tested showed a specific
color of the flame.
Why do certain elements give off light of specific color when heat is applied?
These colors given off by the vapors of elements can be analyzed with an instrument called
spectroscope.
Niels Bohr

18. Figure 1: An atomic Spectroscope

19. ROYGBIV
• A glass prism separates the light given off into its
component wavelength.
• The spectrum produced appears as a series of sharp
bright lines with characteristic colors and wavelength on a
dark background instead of being continuous like the
rainbow.
• We call this series of lines the atomic spectrum of the
element. The color, number and
• position of lines produced is called the “fingerprint” of an
element. These are all
• constant for a given element. See Fig. 2.
Figure 2. Atomic spectra of H, Na, and Ne

20. • How did Bohr explain what you observed in
Activity 1 and the findings about the elements in a
spectroscope?

21. • Individual lines in the atomic spectra of elements indicate definite energy transformations within the
atom.
• Bohr considered the electrons as particles moving around the nucleus in fixed circular orbits. These
orbits are found at definite distances from the nucleus.
• The orbits are known as the energy levels, n where n is a whole number 1, 2, 3…and so forth.

22. • Electrons in each orbit have a definite energy, which
increases as the distance of the orbit from the
nucleus increases.
• As long as the electron stays in its orbit, there is no
absorption or emission of energy. As shown in Figure
3, when an electron of an element absorbed extra
energy (from a flame or electric arc), this electron
moves to a higher energy level.
• At this point the electron is at its excited state.
• Once excited, the atom is unstable.
• The same electron can return to any of the lower
energy levels releasing energy in the form of light
with a particular color and a definite energy or
wavelength.
• Bohr’s model explained the appearance of the bright
line spectrum of the hydrogen atom but could not
explain for atoms that has more than one electron.
Figure 3. Excited state of an electron

23. • The energy levels of electrons are like the steps of a ladder.
• The lowest step of the ladder corresponds to the lowest
energy level.
• A person can climb up and down by going from step to step.
• Similarly, the electrons can move from one energy level to
another by absorbing or releasing energy.
• Energy levels in an atom are not equally spaced which
means that the amount of energy are not the same.
• The higher energy levels are closer together.
• If an electron occupies a higher energy level, it will take less
energy for it to move to the next higher energy level.
• As a result of the Bohr model, electrons are described as
occupying fixed energy levels at a certain distance from the
nucleus of an atom.

24. • However, Bohr’s model of the atom was not sufficient to
describe atoms with more than one electron.
• The way around the problem with the Bohr’s model is to know
the arrangement of electrons in atoms in terms of the probability
of finding an electron in certain locations within the atom.

25. • Bohr’s idea that electrons are found in definite
orbits around the nucleus was rejected.

26. Three physicists led the development of a
better model of the atom.
• Louie de Broglie - De Broglie proposed that the
electron (which is thought of as a particle) could also
be thought of as a wave
• Erwin Schrodinger - used this idea to develop a
mathematical equation to describe the hydrogen
atom
• Werner Karl Heisenberg - discovered that for a very
small particle like the electron, its location cannot be
exactly known and how it is moving. This is called
the uncertainty principle.

27. • 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. Thus, we could only guess the most probable location of the electron at a
certain time to be within a certain volume of space surrounding the nucleus.
• The quantum mechanical model of the atom comes from the mathematical solution to the
Schrodinger equation.
• The quantum mechanical model views an electron as a cloud of negative charge having a certain
geometrical shape.
• This model shows how likely an electron could be found in various locations around the nucleus.
• However, the model does not give any information about how the electron moves from one position
to another.

29. Signature
Maria Jane Dela Cruz
11/29/2022

31. • The principal quantum number always equals the number of
sublevels within that principal energy level. The maximum number of
electrons that can occupy a principal energy level is given by the
formula 2𝑛2.