it explains the model of an atom along with all necesary details it contains all the theorems orbitals etc complete giant slide file of and on atoms

2. Objectives
 Be able to define an atom
 Distinguish between a molecule and a compound

3. Atoms
 Definition
- the smallest particle that has the properties of
an element, basic unit of matter
- surprising the number of atoms is small
- the number of combinations these atoms make
are huge
ex. color print: 3 colors can make
numerous colors
- 119 distinct atoms as of 1999, form elements

4. Atoms Cont.
- hydrogen makes up more than 90% of the atoms in
the universe
- 1st direct evidence was inadvertently discovered in
1827, by Scottish botanist Robert Brown while he
was studying pollen
- grains were in a constant state of agitation
- thought they were moving life forms, later
discovered a perpetual jiggling of particles known as
brownian motion; collisions between visible
particles and invisible atoms

5. Atomic Symbols
- each element has its own name, accompanied by
a symbol
- usually two/three letters (first is always
capitalized)
ex. Iron: Fe
- Fe represents 1 atom of iron,
- 2Fe represents 2 atoms of iron etc…
- can also be written as Fe2

6. Molecules
 Definition
- the smallest unit of a substance that exhibits all the
properties characteristics of that substance
- two or more atoms
ex. H8, O2

7. Compound
 Definition
- a substance that is made from two or more simpler
substances and can be broken down into those
simpler substances
ex. H2SO4

8. Objectives
 Describe ancient Greek models of of matter
 List the main points of Dalton’s atomic theory and
describe his evidence for the existence of atoms
 Explain how Thomson and Rutherford used data
from experiments to produce their atomic models

9. Democritus Theory
 Greek Philosopher in the 4th century B.C
- believed that all matter consisted of extremely
small particles
- suggested these particles are made of invisible
units called atoms
- term atom is derived from a Greek word
meaning “unable to divide”
- believed there were different types of atoms,
liquids: round,smooth solids: rough, prickly
- unable to provide evidence that an atom existed,
therefore many people were very skeptical

10. Dalton’s Atomic Theory
 John Dalton
- interested in predicting the weather SO.. studied
the behavior of gases in the air, concluded that a gas
consists of individual particles
 Evidence
- masses of elements as they combined to form
compounds always produced the same ratio no
matter what the size of the sample
ex. carbon dioxide
- 1 carbon 2 oxygen: 1:2 ratio

11. Dalton’s Atomic Theory
 Theory
- used a Greek concept of the atom and the 3 laws to
give the atomic theory a scientific basis

12. Dalton’s Atomic Theory Cont.
 5 Principles
1. All matter is made of indivisible and
indestructible atoms
2. All atoms of a given element are identical in
their physical and chemical properties
3. Atoms of different elements differ in
their physical and chemical properties
4. Atoms of different elements combine in
simple whole-numbers ratios to form
compounds

13. Dalton’s Atomic Theory Cont.
5. Chemical reactions consist of the
combination, separation, or rearrangement of
atoms
- theory explained most of the chemical data of
the day and was readily accepted
- evidence since has shown the first two
principles are not valid; overlooked that most
atoms will combined with other of their own
kind
- NOT discarded only modified

14. Objectives
 Explain JJ Thompson’s experiment in detail
 Understand Thompson’s observations and
conclusions
 Draw a Thomson’s model of the atom

15. Thomson’s Experiment
 J.J Thomson
1st experiment
- pumped most the air out of a glass tube,
placed a metal plate at each end
- applied a voltage to two metal plates one
became positively charged:anode
other became negatively charged: cathode

16. Thomson’s Experiment Cont.
2nd experiment
- placed a charged metal plate on either side
of the glass tube

17. Thomson’s Experiment Cont.
 Observations
1st experiment
- glowing ray emerged between the cathode and
anode
2nd experiment
- charged plates caused the beam to deflect/bend
- repelled from the negative
- attracted to the positive
 Conclusion
- beam of light (stream of charged particles)
- negative

18. Thomson’s Experiment Cont.
Uses
- TV screens
- computer monitors
- radar displays
*Later became known as cathode rays due to their
origin now known as an electron beam
- 1st to provide evidence that atoms are made of
smaller particles
- revised Dalton’s model

19. Thomson’s Model
* If there is a negative charge there must also be a
positive charge
 Why?
- believes the atom is neutral
 Plum Pudding
- negative particles are evenly scattered throughout
an atom with a positively charged mass of matter
- similar to that of chocolate chip ice cream
- later proved to be incorrect

20. Objectives
 Explain Rutherford’s experiment in detail
 Understand Rutherford’s observations and
conclusions
 Draw Rutherford’s model of the atom
 Compare and contrast Thomson’s and Rutherford’s
models

21. Rutherford’s Theory
 Ernest Rutherford
- a former student of Thomson came up with a more
accurate picture of the atom in 1909
-oversaw the now famous gold foil experiment
 Gold Foil Experiment
 Hypothesis
- alpha particles are thousands of times more
massive, hence they would not be impeded as it
passed through the “atomic pudding”
- beam of positively charged particles, alpha particles
from a radioactive source was directed through a
sheet of very thin gold foil

22. Gold Foil Experiment

23. Gold Foil Experiment Cont.
 Observations
- nearly all passed through undeflected and produced
spots of light
- some were widely deflected, and a few bounced
straight back
What massive object did they hit?
atomic nucleus, an extremely dense positively
charged center of the atom

24. Gold Foil Experiment Cont.
Why did the others pass through then?

25. Gold Foil Experiment Cont.
Conclusion
- atom is mostly empty space
- most of its mass concentrated in the central
region, atomic nucleus

26. Gold Foil Experiment Cont.
- the nucleus and surrounding electrons occupy only a
tiny fraction of the atomic volume
- diameter of an atom is generally about 10,000
times greater than the diameter of its nucleus
* If the nucleus were the size of the period at the end
of this sentence, the outer edges of the atom would
be located some 3.3 meters away *

27. Gold Foil Experiment Cont.
 Rutherford’s Model
- all of the atoms positive charge is concentrated in
the nucleus, which only takes up a very small
amount of the atom
Can we then say we are mainly empty space?

28. Objectives
 Identify three subatomic particles
 Understand how subatomic particle was discovered
 Compare the properties of the subatomic particles
 Distinguish between atomic number and mass
number
 Calculate the number of protons, electrons and
neutrons in an atom

29. Subatomic Particles
 Subatomic particles
- 3 important to chemistry
- protons, neutrons, electrons

30. Protons
 Definition
- a positively charge subatomic particle that is found
in the nucleus of an atom
 About Protons
- proton is nearly 2000 times more massive than the
electron, but equal in charge and opposite in sign to
the electron
- number of protons in the nucleus is electrically
balanced by an equal number of electrons
ex. oxygen atom: contains 8 electrons
and protons: neutral atom, no net charge

31. Electron
 Definition
- a negatively charged subatomic particle that is
found in the space outside the nucleus
- name comes from the Greek word for amber
- Amber: material discovered by early Greeks that
was found to exhibit the effects of electrical
charging
ex. Ben Franklin: Key/Kite
- lead others to experiment with electric currents
through gases in sealed tubes

32. Neutrons
 Definition
- is a neutral subatomic particle that is found in the
nucleus of the atom
- mass almost exactly equal to that of the proton

33. Comparing Subatomic Particles

34. Atomic Number
 Definition
- number of protons in the atom
ex. Oxygen 8p + 8n = 16
- elements are classified by this number
- continues up to 119
- unique to a given element
- all atoms are electrically neutral, meaning the
number of electrons must equal the number of
protons
- this arrangement of elements by their atomic
numbers makes up the periodic table
- Usually located at the upper left hand corner

35. Mass Number
 Definition
- the total number of protons and neutrons in the
nucleus of an atom
- mass number – atomic number = neutrons
ex. N: mass number of 14
atomic number of 7
7 neutrons

36. Mass Number Cont.
- usually found at the bottom of the atomic symbol,
and sometimes found written at the bottom left of an
atomic symbol
ex. 16O

37. Mass Number
- although a given type of atom will usually contain a
certain number of neutrons in the nucleus, a small
percentage will not
ex. most hydrogen atoms contain no neutrons
- a small percentage contain one neutron and a smaller
percentage two neutrons
 What do we call atoms with a different number of
neutrons?
- isotopes

38. Isotopes
 Definition
- the number of neutrons in the nucleus of a given
element may vary, protons remain the same
ex. H contains 1 protons (1H)
H contains 1 protons and 1 neutrons
(2H) deuterium
H contains 1 protons and 2 neutrons
(3H) tritium
ex. 14C: Carbon-14

39. Atomic Mass
 Definition
- mass of an atom in atomic mass units (amu)
- atoms have very little mass
- equal to 1/12th of the mass of carbon
- often an average mass
- weighted mass
 AMU or the Dalton (Da)
- equal to 1.6605402 x 10-27 kg

40. Atomic Mass Number Cont.
ex. 99% of all carbon atoms are the isotope
containing 6 neutrons, the remaining 1% is the
heavier isotope containing 7 neutrons, which
raises the average mass of carbon from 12.000 to
12.011

41. Objectives
 Describe Bohr’s model of the atom and the evidence
for energy levels
 Explain how the electron cloud model represents the
behavior and locations of electrons in atoms

42. Bohr’s Model of the Atom
 Bohr’s Model
- an early conceptual model of the atom
- classic planetary model in which electrons whirl
around the small but dense nucleus: like planets
orbiting the Sun
- developed by the Danish physicist Niels Bohr in
1911
- each electron has a certain energy that is
determined by it’s path around the nucleus
- explains how atoms (lose) emit or (gain) absorb
energy
resulting energy: energy level

43. Bohr’s Model: Energy Levels
 Definition
- any of the possible energies an electron may have
in an atom
 Evidence
- measured amount of energy gained or lost
- when energy is lost we often see it as a light
ex. fireworks
- 1925 this model no longer explained all
observations being made by scientists and new
models were created
ex. Electron cloud model

44. Electron Cloud Model
- visual model of the most likely locations for
electrons in an atom
 Orbitals
- a region in an atom where there is a high
probability of finding an electron
ex. propeller on a helicopter (you know its there
you see a blur, can’t pinpoint exact location)
- 4 orbitals
s – 2 houses 2 e-
p – 3 houses 6 e-
d – 5 houses 10 e-

45. Orbitals
- electrons occupy the lowest energy levels first
- electrons in the outermost energy levels of an atoms
are called valence electron

46. Electron Configurations
 Definition
- arrangement of electrons in the orbitals of an atom
(similar to seating assignments on an airplane)
- when all electrons at their lowest energies this is
called ground state