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1. Overview
In Grade 8, you have learned the Rutherford’s atomic model which pictures the atom
as mostly empty space and its mass is concentrated in the nucleus, where you find the
protons and the neutrons. This model has worked well during his time, but it was only able
to explain a few simple properties of atoms. However, It could not explain why metals or
compounds of metals give off characteristic colors when heated in a flame, or why objects–
when heated to much higher temperatures first glow to dull red, then to yellow, and then to
white. A model different from Rutherford’s atomic model is necessary to describe the
behavior of atoms
Niels Bohr refined Rutherford’s model of an atom. Based on his experiments, Bohr
described the electron to be moving in definite orbits around the nucleus. Much later,
scientists discovered that it is impossible to determine the exact location of electrons in an
atom. In Activity 1, you will learn about the evidence that Bohr used to explain his model of
the atom. In Activity 2, you will do a task that will help you understand that there is a certain
portion of space around the nucleus where the electron is most likely to be found.
In addition, you will know more about the present model of the atom, which is called
the quantum mechanical model of the atom. It is important for you to understand that the
chemical properties of atoms, ions and molecules are related to how the electrons are
arranged in these particles of matter. You will find out the answers to the following
questions as you perform the activities in this module.
How does the Bohr atomic model differ from Rutherford’s model? What is the
basis for the quantum mechanical model of the atom? How are electrons arranged
in the atom?

2. Activity 1: The Flame Test
Objectives:
o determine the characteristic colors that metal salts emit; and
o relate the colorsemittedbymetal saltstothe structure of the atom.
Materials:
o 0.50 grams of each of the following metal salts:
o Calcium chloride 6 pcs watch glass
o Sodium chloride 1 pc 10-ml graduated cylinder
o Copper(II) sulfate 1 pc dropper
o Potassium chloride safety matches
o Boric acid
o 100 mL 95% Ethanol (or ethyl alcohol)
o 100 mL 3 M hydrochloric acid
Procedure:
1. Place eachmetal salton a watch glassand add 2 to 3 drops of 3 M hydrochloricacid.
2. Pourabout 3 - 5 mL or enoughethyl alcohol tocoverthe size of a 1 peso-coininthe firstwatchglass.Light
witha match and observe the colorof the flame.(Thiswill serveasreferenceforcomparisonof the flame
color).Waitfor the flame tobe extinguishedorputout onits own.
3. Repeatprocedure No.2 foreach salt.Observe the colorof the flame.

3. 4. Write your observation in a table similar to the one below.
Table 1. Color of flame of metal salts
Metal salt tested
Element
producing color
Color of the flame
Boric acid boron
Calcium chloride calcium
Sodium chloride sodium
Potassium chloride potassium
Copper(II) sulfate copper
Q1. Why do you think are there different colors emitted?
Q2. What particles in the heated compounds are responsible for the production of the colored light?
Q3. How did the scientists explain the relationship between the colors observed and the structure of
the atom?
________________________________________________________________
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. See Figure1.
Precautions:
1. Wear goggles, gloves and a safety apron while performing the
activity.
2. Do this activity in a well-ventilated area.
3. Handle hydrochloric acid with care because it is corrosive.
4. Ethyl alcohol is flammable.
5. Be careful to extinguish all matches after use.

4. A glass prism separates the light given off into its component wavelength. The spectrum produced appears
as a seriesof sharp bright lineswith characteristic colors and wavelength on a dark background instead of
being continuous like the rainbow. We call this series of linesthe atomic spectrum of the element.The
color, number and position of linesproduced is called the “fingerprint” of an element.These are all
constant for a given element.See Fig. 2.
How did Bohr explain what you observed in Activity 1 and the findings about the elements in a
spectroscope? 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.
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 elementabsorbed 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 levelsreleasingenergy in the
form of light with a particular color and a definite energy or wavelength. Bohr’s model explainedthe
appearance of the bright line spectrum of the hydrogen atom but could not explainfor atoms that has
more than one electron.

5. Q4. Explain how your observation in Activity 1 relates to Bohr’s model of the atom. You can explain using an
illustration.
Q5. Which illustration below represents the energy of the electron as described by Bohr? Explain your answer.
The energy levelsof 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 levelsinan atom are not equally spaced which means
that the amount of energy are not the same. The higher energy levelsare closer together. If an electron occupies a
higher energy level,itwill take lessenergy for it to move to the nexthigher energy level.As a result of the Bohr model,
electrons are described as occupying fixedenergy levelsat a certain distance from the nucleus of an atom.
However, Bohr’s model of the atom was not sufficientto 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. In the next activity, you will use an analogy to
understand the probability of findingan electron in an atom.
Activity 2:
Predicting the Probable Location of an Electron
Objective:
o Describe how it is likely to find the electron in an atom by probability.
Materials:
o One sheet of short bond paper or half of a short folder
o pencil or colored marker with small tip
o compass
o graphing paper
o one-foot ruler
o Procedure:
o 1. Working with your group mates, draw a dot on the center of the sheet of paper or folder.
o
o 2. Draw 5 concentric circles around the dot so that the radius of each circle is 1.0 cm, 3 cm, 5 cm, 7
cm and 9 cm from the dot

6. 3. Tape the paper on the floor so that it will not move.
4. Stand on the opposite side of the target from your partner.(Target is the center which
represent the nucleus of an atom). Hold a pencil or marker at chest level above the center of
the circles you have drawn.
5. Take turns dropping the pencil or marker so that it will leave 100 dots on the circles drawn
on paper or folder.
6. Count the number of dots in each circle and record that number on the data table.
7. Calculate the number of dots per square centimeter (cm2).
8. Using a graphing paper, plot the average distance from the center on the x-axis and
number of dots per sq.cm on the y-axis.

7. Data Table:
Circle
Number
(A)
Average
Distance
from
Center
cm
(B)
Area of
Circle,
cm2
(C)
Difference of
Areas of the
Two
Consecutive
Circles, cm2
(D)
Number of
Dots in
Circle
(E)
Number of
Dots per
cm2
(E)/(D)
(F)
Percent
Probability
of Finding
dots,
%
(G)
1 1.0 3.14 25.13 5 0.1920 19.20
2 3.0 28.27 50.27
3 5.0 78.54 75.40
4 7.0 153.94 100.53
5 9.0 254.47 125.66
Q1. What happens to the number of dots per unit area as the distance of the dots go farther from the
center?
Q2. Determine the percent probability of finding a dot in each of the circle drawn on the target by
multiplying No. of dots /cm2 (column D) by the total number of dots (100). For example: In circle 1(A)
Percent probability = No. of dots /cm2 X 100
= [0.1920 / 100 ] X 100 = 19.20%
Q3. Based on your graph, what is the distance with the highest probability of finding a dot? Show this
in your graph.
Q4. How many dots are found in the area where there is highest probability of finding dots?
Q5.How are your results similar to the distribution of electrons in an atom?
Activity 1 is an analogy to show you that it is notpossible to know the exact position of the electron. So, Bohr’s idea that
electrons are found in definite orbits around the nucleus was rejected. Three physicistsled the development of a better model
of the atom. These were Louie de Broglie,Erwin Schrodinger,and Werner Karl Heisenberg. De Broglieproposed that the
electron (which is thought of as a particle) could also bethought of as a wave. Schrodinger used this idea to develop a
mathematical equation to describethe hydrogen atom. Heisenberg discovered that for a very small particlelikethe electron,
its location cannotbe exactly known and how it is moving. This is called the uncertainty principle.
Instead,these scientists believed that there is only a probability thatthe electron can be found in a certain volume in space
around the nucleus.This volume or region of spacearound the nucleus where the electron is most likely to be found is called
an atomic orbital.Thus,we could only guess the most probablelocation of the electron at a certain time to be within a certain
volume of spacesurroundingthe 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 havinga certain geometrical shape. This model
shows how likely an electron could be found in various locationsaround thenucleus.However, the model does not give any
information abouthow the electron moves from one position to another.