Post on 01-Jun-2020
Chapter 2 Atom
1. Atom and History of Atom
2. Subatomic Particles
3. Isotopes
4. Ions
5. Atomic Terminology
Table of Contents
Chapter 2
Warm up
The History of The Atom
• Make a list of inferences about any properties of objects in the box.
• How could you learn more about the objects in the box without opening the box?
• Scientist face these same questions as they try to learn more about atoms.
Chapter 2 The History of The Atom
1. Atomic Models
2. Modern Atomic Model
3. Quantum Numbers
4. Electron Configuration
Table of Contents
Chapter 2
Warm up
The History of The Atom
• Make a list of inferences about any properties of objects in the box.
• How could you learn more about the objects in the box without opening the box?
• Scientist face these same questions as they try to learn more about atoms.
Chapter 2 The History of The Atom
• The idea of an atomic theory is more than 2000 years old.
• Until recently, scientists had never seen evidence of atoms. • Democritus (460-400BC), a Greek philosopher, first thought
that the universe was made up of very small particle and
named them “atomos”, meaning indivisible.• Aristotle, however, thought that matter was continuous, that
there was no limit on how finely you could cut it up.
• In the twentieth century, the scientists studied about atom are
Dalton, Rutherford, Planck, Einstein, Bohr and Schrödinger.
Chapter 2 1. Atomic Models
1. Dalton’s Atomic ModelIn 1803 John Dalton proposed some ideas about atom;
1. Atom or atom groups have all the characteristic properties
of substances.
2. Atoms of an element are completely identical.
3. An atom is a filled sphere like a billiard ball.
4. Different types of atoms have different masses.
5. Atoms are the smallest unit of substances and cannot be
further divided.
6. Atoms form molecules in a definite numerical ratio.
Chapter 22. Thomson’s Atomic Model
In 1897 J. J Thomson discovered the electron.
1. Protons and electrons are charged particles.
2. In neutral atoms since the number of electron and number
of protons are equal, net charge is zero.
3. An atom has a shape of sphere with 10-8 cm radius. Protons
and electron are distributed in arbitrary positions.
4. The mass of an electron is so small that it can be neglected.
1. Atomic Models
Chapter 22. Thomson’s Atomic Model
1. Atomic Models
Chapter 22. Thomson’s Atomic Model
1. Atomic Models
Chapter 23. Rutherford’s Atomic Model
In 1911 Ernest Rutherford discovered the nucleus and in 1919
proton . He proposed α-particle experiment. His ideas;
1. There is a small positively charged, dense region called
nucleus in an atom.
2. The mass of the atom approximately equal to the mass of
the protons and electrons.
3. Protons are in the nucleus and electrons are distributed
around nucleus.
1. Atomic Models
Chapter 23. Rutherford’s Atomic Model
1. Atomic Models
Chapter 23. Rutherford’s Atomic Model
1. Atomic Models
Chapter 23. Rutherford’s Atomic Model
1. Atomic Models
Chapter 24. Bohr’s Atomic ModelIn 1913 Niels Bohr proposed places of electrons around the
nucleus of atoms. He concluded his ideas as follows;
1. Electrons are found in certain places so called “energy
levels” or “shell” around nucleus with certain energies.
2. Electrons move in each stationary state in circular paths.
3. When an electron falls back to a lower energy level from
a higher one, it emits a quantum of light that is equal to the
energy difference between these two energy levels.
1. Atomic Models
Chapter 24. Bohr’s Atomic Model
4. The possible stationary energy levels of electrons are
named either by letters, K, L, M, N, O ... or by positive integer
numbers starting from the lowest energy level. These
numbers are generally denoted by “n” where n = 1, 2, 3, ...
Discovery of NeutronIn 1930s James Chadwick thought that since protons have
same charges in the nucleus and they found together, there
must have been some neutral particles which hold protons
together, then he called them “neutron”.
1. Atomic Models
Chapter 2
• Although Bohr’s model was valid for atoms of H or He+1 or
Li+2, it did not work for atoms having more than one electron.
But his ideas lead to a step forward in the development of
modern atomic theory.
2. Modern Atomic Model
• The pioneers to modern atomic theory are Lois de Broglie,
Heisenberg and Schrödinger.
• In 1924 Lois de Broglie proposed that small particles
sometimes show wave-like properties.
• In 1920 Werner Heisenberg stated his uncertainty principle
which explains position of electrons.
Chapter 2 2. Modern Atomic Model
Chapter 2 2. Modern Atomic Model
Chapter 2
• Modern atomic theory explains the probability of finding
electrons around the nucleus by virtue of quantum numbers
and orbitals.
2. Modern Atomic Model
• The quantum numbers are the integer numbers designating the energy levels of the electrons in an atom, and the orbitalsare the probable regions in which the electrons might be found around the nucleus.
Chapter 2 2. Modern Atomic Model
Chapter 2 2. Modern Atomic Model
Comparison of Bohr’s Atomic Model and Modern Atomic Model
Chapter 2 1. Atom
• What is all matter made up? And to what do they look like?
Warm up
• Make a list of inferences about any properties of objects in the box.
• How could you learn more about the objects in the box without opening the box?
• Scientist face these same questions as they try to learn more about atoms.
Chapter 2 2.1. Atom
• Scientist have accepted that the smallest parts of
substances are called atoms.
• Atom has basically two parts, nucleusand electrons. Nucleus is located in the center of atom and electrons are rotating around nucleus with high speed.
• Atom means indivisible derived from
atomos in Greek language because of
its very small size.
Chapter 2 2. Subatomic Particles
• Scientist believed that atoms were indivisible up to 20th
century. Today it is well known that atoms have
subatomic particles, called protons, neutrons and
electrons.
• Protons are positively charged particles found in the
nucleus of an atom, and denoted by “p”. Each element
has certain number of protons which differ the element
from others.
Chapter 2 2. Subatomic Particles
• Neutrons are neutral particles found in the nucleus of
an atom, and denoted by “n”.
• Electrons are negatively charged particles placed
around the nucleus of an atom, and shown by “e-”.
• Protons and neutrons almost have the same masses,
but electrons have negligible mass with respect to
protons and neutrons.
Chapter 2 2. Subatomic Particles
Chapter 2 2. Subatomic Particles
Chapter 2 2. Subatomic Particles
• Neutral atoms have the equal number of protons and
electrons.
• Electrons are rotating in certain places called orbit,energy level or shell. Energy levels are represented by letters, K, L, M, N, O…etc, or numbers, 1, 2, 3, 4, …etc.
• Each shell can hold a certain number of electrons calculated by the equation of “2n2” where n refers to number of shell.
• The electrons located in the outermost shell of atoms are
called valance electrons.
Chapter 2 2. Subatomic Particles
• In the 1st shell, No. of e- = 2x12 = 2e-
In the 2nd shell, No. of e- = 2x22 = 8e-
In the 3rd shell, No. of e- = 2x32 = 18e-
In the 4th shell, No. of e- = 2x42 = 32e-
Example 1Show the electron configuration of 6C and 13Al atoms.
6C: 2) 4)
13Al: 2) 8) 3)
Solution
Chapter 2 2. Subatomic Particles
Chapter 2 3. Isotopes• Isotope atoms have the same number of protons but
different number of neutrons.
• They have similar chemical properties but different
physical properties.
Example 2
C6
12C
6
13C
6
14atoms are isotopes.
All they have 6 protons but 6, 7 and 8 neutrons respectively.
Chapter 2 3. Isotopes
Chapter 2
• Isotones, atoms with the same number of neutrons, but
different numbers of protons.
Example 3
P31
15S
32
16atoms are isotones.
Each have 16 neutrons but 15 and 16 protons respectively.
• Isotone atoms are completely different atoms, they have
different chemical and physical characteristics.
3. Isotopes
Chapter 2 4. Ions
• Electrically charged atoms are called ions.
• When an atom loses electrons it becomes positively
charged ion, called cation.
• When an atom gains electrons it becomes negatively
charged ion, called anion.
• Charge of an atom, q , can be found with q = p - e.
Chapter 2 4. Ions
Chapter 2 4. Ions
Atom
Neurtal Atom Ion(Charged Atom)
Cation (+) Anion (-)(p>e) (p<e)
p=e
Chapter 2 4. Ions
Example 4Find the charge and ion type of atom.
Atom Proton Electron Charge IonFe 26 24Al 13 10O 8 10P 15 15Cl 17 18
Chapter 2 4. Ions
Solution
Atom Proton Electron Charge IonFe 26 24 +2 CationAl 13 10 +3 CationO 8 10 -2 AnionP 15 15 0 NeutralCl 17 18 -1 Anion
Chapter 2
Example 5
Li+1, Ca+2, Al+3, Pb+4 are cations.
F-1, O-2, P-3 are anions.
OH- NO3-
CO3-2
PO4-3NH4
+1are polyatomic ions
4. Ions
Chapter 2 5. The Atomic Terminology
• Atomic number = Number of Protons
• Each type of atom has different number of protons.
For a neutral atom,
• Atomic number = Number of protons = Number of electrons
Z = p = e
1. Atomic Number, Z
Chapter 2 5. The Atomic Terminology
• Atomic mass number = Number of Protons + Number of Neutrons
A = p + n
2. Atomic Mass Number, A
Example 6Fill in the blanks in the table below.
Atom p n Z ATi 22 48Al 14 27S 16 16Br 45 35
Chapter 2 5. The Atomic Terminology
Solution
Atom p n Z ATi 22 26 22 48Al 13 14 13 27S 16 16 16 31Br 35 45 35 80
Chapter 2 5. The Atomic Terminology
XAtomic Mass Number, A
Atomic Number, ZProtons, p
Neutrons, n
Electron, e
Charge, q
q = p - e
A = p + n
Chapter 2 5. The Atomic Terminology
Example 7
What is the number of protons and atomic mass
number of Zn.
Zn+2
28
?
?35
Chapter 2 5. The Atomic Terminology
Solution
Zn+2
28
65
3035
q = p - e
A = p + n
+2 = p – 28
p = 30
A = 30 + 35
A = 65
Chapter 2 5. The Atomic Terminology
Example 8
Cr+3 ion has 21 electrons and its atomic mass number is 52.
What is the number of neutrons for Cr ?
Solution
q = p - e
A = p + n3 = p – 21 p = 24
52 = 24 + n n = 28
Chapter 2 5. The Atomic Terminology
• 12C is accepted as a standard atom. The mass of the 12C
isotope atom is accepted as 12 amu ( atomic mass unit) and
atomic masses of other atoms were calculated accordingly.
For example relative atomic masses of H atom is 1.008
(≈1 amu) and oxygen atom becomes 15.9994 (≈16 amu)
with respect to 12C.
3.Relative Atomic Mass and Relative Formula Mass
Chapter 2 5. The Atomic Terminology
• Relative formula mass is sum of the relative atomic masses of
atoms found in a compound.
Example 9Relative formula mass of CH4 and NO2
CH4 = (1x12) + (4x1) = 16 amu
NO2 = (1x14) + (2x16) = 46 amu
Chapter 2 5. The Atomic Terminology
• Most of the elements in nature are found as a mixture of
isotope atoms.
• The average atomic mass is the average masses of natural
isotopes of an element.
4.Average Atomic Mass
Average Atomic Mass =
A1x% of 1st isotope + A2x% of 2nd isotope + ……
A1 and A2 are atomic mass numbers of natural isotopes.
Chapter 2 5. The Atomic Terminology
Example 10
Naturally occurring Ga consists of 60% 69Ga and 40% 71Ga.
What is the average mass of Ga?
SolutionAverage Atomic Mass =
A1x% of 1st isotope + A2x% of 2nd isotope + ……
Average Atomic Mass = 69x60 + 71x40100 100
= 69.80 amu
Chapter 2 5. The Atomic Terminology
Example 11
Find the average atomic mass of Pb?
End of the chapter 2