John Dalton suggested that all matter is made up of tiny spheres called atoms in the early 1800s. J.J. Thomson discovered electrons through cathode ray tube experiments in the late 1800s. Ernest Rutherford determined that the nucleus occupies a very small portion of the atom through nuclear scattering experiments. Niels Bohr developed the Bohr model of the atom to explain atomic spectra. Atoms have a small, dense nucleus containing protons and neutrons, surrounded by electrons in an electron cloud. The number of protons determines the element, and electrons normally equal protons for neutral atoms. Isotopes are atoms of the same element with different numbers of neutrons.

1. Submitted by
Nihal tyagi
Xi a5

3. 1808 john dalton
 Suggested that all the matter
was made up of tiny spheres
that were able to bounce back
around perfecty with elasticity
 https://www.britannica.com/
video/153020/John-Dalton-
development-atomic-theory
Regarded atom as ultimate particle

4. 1898 JJ thompson
 Carried out cathode ray tube
experiment
 In his studies he founded that:
 Cathode start from cathode
and move toward towards
anode
 rays are not visible
 He named them as electron
Discovered electron

5. Erenest Rutherfofd
 Conducted nuclear scatterind
experiment
 Gave conclution that:
 Most of the space in atom is
empty
 His calculation shows that
nucleus occupy negligble
mass
 He show that positive charge
is not scaterred

6. Neil bohr
 Bohr model of atom
 Dual character of electonic
radiation
 Experiment result regarding
atomic spectra

7. Electron around nucleus

8. What is an atom?
 Atom: the smallest unit of matter that retains
the identity of the substance

9. Atomic Structure
 Atoms have 2 regions
1) Nucleus: the center of the atom that contains
most of the mass of the atom.
2) Electron cloud: surrounds the nucleus & takes up
most of the space of the atom.
Nucleus
Electron
Cloud

10. What’s in the Nucleus?
 In the nucleus we find:
 Protons: positively charged subatomic particles
 Mass of 1 amu
 Neutrons: neutrally charged subatomic particles
 Mass of 1 amu

11. What’s in the Electron Cloud?
 In the electron cloud we find:
 Electrons: the subatomic particle with a negative
charge and relatively no mass
 Mass of ~ 1/1836 amu

12. Subatomic Particles
Particle Charge Mass (g) Location
Electron
(e-) -1 9.11 x 10-28 Electron
cloud
Proton
(p+) +1 1.67 x 10-24 Nucleus
Neutron
(no) 0 1.67 x 10-24 Nucleus

13. How do we know the number of
protons in an atom?
 Atomic number (#)= # of protons in an atom
 Ex: Hydrogen’s atomic # is 1
 hydrogen has 1 proton
 Ex: Carbon’s atomic # is 6
 carbon has 6 protons
**The number of protons identifies the atom-it’s
an atom’s fingerprint.

14. How do we know the number of
neutrons in an atom?
 Mass #: the # of protons plus neutrons in the
nucleus
 # of neutrons = mass # - atomic #
Example
 Li has a mass # of 7 and an atomic # of 3
 Protons = 3 (same as atomic #)
 Neutrons= 7-3 = 4 (mass # - atomic #)

15. Mass # vs. Atomic Mass
Mass # ? = The Atomic mass on the periodic table
rounded either up or down

16. How do we find the number of
electrons in an atom?
 Most atoms are neutral (have no overall
charge)
 Because the only charged subatomic
particles are the protons and electrons…
they must balance each other out in an
electrically neutral atom.
 Therefore..
 # Electrons = # Protons *
* (in a neutral atom..)

17. Examples
 He has a mass # of 4 and an atomic # of 2
 p+ = 2 no = 2 e- = 2
 Cl has a mass # of 35 and an atomic # of 17
 p+ = 17, no = 18, e- = 17

18. How exactly are the particles
arranged?
 Bohr Model of the atom: electron
configurations
All of the
protons and
the neutrons
The 1st ring can
hold up to 2 e-
The 2nd ring can
hold up to 8 e-
The 3rd ring
can hold up
to 8 e-

19. What does carbon look like?
Mass # = 12 atomic # = 6
p+ = 6 no = 6 e- = 6
6 p and 6 n live
in the nucleus

20. Isotopes
Dalton’s 1st postulate was wrong.
Atoms of the same element can
be different (they can have
different # of neutrons)
Thus, different mass numbers.
These are called isotopes.

21. Isotopes
Isotopes are atoms of the same element
having different masses, due to varying numbers
of neutrons.

22. Isotope Protons Electrons Neutrons Nucleus
Hydrogen–1
(protium) 1 1 0
Hydrogen-2
(deuterium) 1 1 1
Hydrogen-3
(tritium)
1 1 2

23. Naming Isotopes
We name the isotope based on
its mass number
carbon-12
carbon-14
uranium-235

24. Isotopes
• Elements
occur in
nature as
mixtures of
isotopes.

25. Atomic Mass
 How heavy is an atom of oxygen?
 It depends.. b/c there are different oxygen
isotopes.
 We are more concerned with the average atomic mass.
 This is determined based on the abundance of each isotope
 We don’t use grams for this mass because the
numbers would be too small.

26. Measuring Atomic Mass
Instead we use the Atomic Mass
Unit (amu)
defined as one-twelfth the mass of a
carbon-12 atom.
Each isotope has its own atomic
mass, thus we determine the
average from percent abundance.

27. Atomic Mass
Isotope Symbol Composition of
the nucleus
% in nature
Carbon-12 12C 6 protons
6 neutrons
98.89%
Carbon-13 13C 6 protons
7 neutrons
1.11%
Carbon-14 14C 6 protons
8 neutrons
<0.01%
Atomic mass is the average of all the
naturally occurring isotopes of that element.
Carbon = 12.011