1. Chem 101
Ppts
Week ending 2/3/12

2. Chapter 2
continued

3. John Dalton
• An English scientist and teacher
• He knew of these observations and offered an
explanation for them
• His explanation is known as:

4. DALTON’S ATOMIC THEORY
The main ideas of his theory include:
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical. Atoms of a given
element are different from those of any other element.
3. Atoms of one element can combine with atoms of other
elements to form compounds. A given compound always
has the same relative number and types of atoms
4. Atoms are indivisible in chemical processes.
** Atoms are not created or destroyed in chemical reactions.
A chemical reaction simply changes the way atoms are
grouped together.

5. * Dalton’s Theory offered simple explainations for some basic
laws of chemistry such as:
The Law of Conservation of Mass
 Mass is neither created or destroyed.
If atoms are conserved in a reaction then mass must also be
conserved
The Law of Constant Composition
 Tells us that a cmpd regardless of its origin or method of
preparation always contains the same elements in the same
proportions by weight
Law of Multiple Proportions
 When 2 elements combine to form more than 1 cmpd the
masses of one element which combines with a fixed mass of
the other elelment are in a ratio of small whole numbers
such as 2:1

6. Law of Multiple Proportions
Example:
Element A + B combine to form 2 cmpds.
Cmpd 1 Cmpd 2
A2B AB
The weight of A combined with a fixed amt of
B in cmpd 1 would be 2x that of the second.

7. Dalton’s Atomic Teory
 Like all new ideas Dalton’s theory was not
accepted immediately
- this did not dissuade Dalton
- Dalton predicted that N and O could combine
to form several compounds:
NO , N2O and NO2
- when the existence of these compounds was
verified…..
Dalton’s Theory became widely accepted

8. A Look at Compounds
A compound is a distinct substance composed
of atoms of two or more elements that always
contains exactly the same relative masses of
those elements
• H2O always has two H atoms and one O atom

9. Chemical Formulas
 Chemical formulas- used to show the types and
number of each type of atom
• Atoms are indicated by the elements symbol
• The number of each type of atom is represented by a
subscript (appearing to the right and below the
element symbol)
ie. SO3 contains 1 atom of sulfur and 3 atoms of oxygen
Write the formula for the compound containing 2
atoms of nitrogen and 5 atoms of oxygen.
Ans. N2O5

10. Chemical Formulas continued
• Glucose has 6 carbon atoms, 12 hydrogen
atoms and 6 atoms of oxygen.
Write its formula.
Ans. C6H12O6
• The unique formula or arrangement of atoms
in a molecule makes one compound
different from another

11. Compounds with the same chemical
formula:
CH3– CH2 –OH CH3-O-CH3
C2H6O C2H6O
ethyl alcohol dimethyl ether
• The difference lies in the order of the bonded atoms
• Formulas that show the order and arrangement of
the specific atoms are called structural formulas

12. • Dalton’s Theory provided a convincing explanation for
the composition of compounds
• Scientists believed that elements consisted
atoms but…….
What did the atom look like?

13. The Structure of the Atom
• Much pondering about the structure occurred
during the 1800’s
• It was not until 1900 that convincing evidence
regarding the structure became available

14. First Convincing Evidence for Subatomic
Particles
• Came from experiments involving the
conduction of electricity through gases at low
pressures using a device called a Cathode Ray
Tube

15. • In 1897 J. J. Thomson (English physicist) took this apparatus
partially evacuated it and connected it to a high voltage
source (spark coil)
• An electric current flows thru the tube
• Associated with this flow are colored rays of light which
originate at the (-) end of the cathode
* Thomson found that these rays were bent by both electric
and magnetic fields
* Careful study of the nature of this deflection demonstrated
that the the rays consisted of a stream of negatively charged
particles which he called electrons.

16. 2.2
• In the absence of any field, the cathode ray (which is -) strikes
at B.
• Cathode ray strikes at A in the presence of a magnetic field
• Electric plates create an electric field perpendicular to the
direction of the cathode rays (cathode ray is -) causing them
to strike the screen at C when the field is on.

17. Cathode Ray Tube
2.2
Passing an
electric current
makes a beam
appear to move
from the
negative to the
positive end.

18. •Went on to measure the mass to charge ratio of the
electron finding it to be:
m/e = 5.69 x 10-9
g/C
e/m = -1.76 x 108
C/g
•Since he found the ratio to be the same regardless
of what gas was in the tube, this implied that the
electron was a fundamental particle common
to all atoms
•Provided the first evidence that atoms are made of
even smaller particles
J.J. Thomson
and the Cathode Ray Tube
1897
(1906 Nobel Prize in Physics)

19. J. J. Thomson
• Thomson showed that atoms of any element can be
made to emit tiny negative particles
• He called these negative particles electrons
• Thomson new that although the atoms contained
negative particles the atoms overall charge had to be
zero
• He postulated that the atom must also contain
positive particles to balance exactly the negative
charge of the electrons

20. Thompson’s Model
• Found the electron
– 1 unit of negative charge
– Mass 1/2000 of hydrogen atom
– Later refined to 1/1840
• Concluded that there must be a
positive charge since atom was
neutral
• Atom was like plum pudding
– A bunch of positive stuff, with
electrons able to be removed.

21. Millikan
1909
• Millikan determined the charge on an electron
with his famous Oil Drop Experiment

22. • In this experiment small droplets of oil which had picked up
extra electrons were allowed to fall between 2 electrically
charged plates
• The drops were observed:
 When the voltage between the plates was increased the
negatively charged drop fell more slowly (it was attracted to
the + plate)
 At some point the drop will be balanced and stationary.

23. Millikan’s Experiment continued:
• Knowing this voltage and the mass of the drop, it
was possible to calculate the charge on the drop
• Millikan found the charge to always be an integral
multiple of a smallest charge
• Assuming the smallest charge to be that on an
electron he arrived at a value of 1.60 x 10-19
C
• Combining this value with the charge/mass ratio he
calculated the mass of an electron

24. e-
charge = -1.60 x 10-19
C
Thomson’s charge/mass of e-
= -1.76 x 108
C/g
e-
mass = 9.10 x 10-28
g
Measured mass of e-
(1923 Nobel Prize in Physics)
2.2

25. Henri Becquerel
1896
• Studying a uranium mineral called pitchblende he
discovered that it spontaneously emitted high
energy radiation.
• Further studies by the Curies and Ernest
Rutherford revealed 3 types of Radiation
1. Alpha (α) - positively charged(+2) and heavy
2. Beta (β)- negatively charged (-1)
3. Gamma (γ)- neutral

26. (Uranium compound)
2.2

27. • With the growing evidence that the atom was
composed of even smaller particles, attention
was being given to how they fit together
• Recall, early in the 1900’s JJ Thomson
proposed the “plum pudding” model of the
atom

28. Ernest Rutherford
• A student of J. J.
Thomson
• Believed in the “Plum
Pudding” Model
• In 1911 he performed
his now famous
“Gold Foil Experiment”

29. “Plum Pudding” Model of the Atom
• Developed by J.J.
Thomson
• The atom was a uniform
pudding of positive
charge with enough
negative electrons
scattered within to
counter balance the
positive charge
Plum Pudding Model or
Raisin Bun Model

30. Rutherford’s Experiment
• Rutherford and his students bombarded a
piece of gold foil using alpha particles
• Here’s what it looked like…….

32. Rutherford’s “Gold Foil” Experiment
• Rutherford directed α-
particles 7500 times the
mass of an electron at a
thin sheet of gold foil
only a few atoms thick
• The detector produced
tiny flashes if it was hit
by an α- particle

33. • What he expected to see

34. – Alpha particles should pass through without
change in direction
– Positive charges were spread out evenly.
Alone they were not enough to stop an alpha
particle

36. • What he got

37. • How he explained it
– Atom is mostly empty
– Small dense, positive piece at the center
– Alpha particles are deflected if they get close
enough to positive center

38. A picture of the Atom Evolves
• Based on his alpha-
particle scattering
experiment on gold,
Rutherford concluded
that the atom consisted
of a hard central core
where most of the mass
of the atom rested.

39. Other Subatomic Particles
• Protons were discovered by Rutherford in
1919.
• Neutrons were discovered by James Chadwick
in 1932.

40. Subatomic Particles
• Protons and electrons are the only particles that
have a charge.
• Protons and neutrons have essentially the same
mass.
• The mass of an electron is so small we ignore it.

41. Nuclear Structure
• Facts about the nucleus:
Protons and neutrons have
roughly the same mass
(1amu=1.67x10-4
g), and each is
about 2000 times as massive as
the electron.
The number of protons is the
same as the number of electrons
(not shown) which orbit the
nucleus at a distance of about 10-
8
cm.
• **If the nucleus were a grape the
electrons would be 1 mile
away!!!
No charge.
Help hold
protons in
the nucleus.
Without
neutrons
the + charges
would repel
one another.
+ charge

42. If all atoms contain the same components
(protons, neutrons and electrons) why do different
atoms have different chemical properties?
The answer lies in the number and arrangement
of the electrons.
***It is the electrons that are responsible for the
chemical properties of an atom of an element.

43. Representing Elements
• Each and every atom of an element is described on the basis
of the number and types of nuclear particles
A Q
X – atomic symbol
Z
A = Mass number
The number of protons and the number of neutrons
Z = Atomic number
The number of protons
In a neutral atom the number of protons = the number
of electrons
Q = Charge if not neutral

44. Atomic Symbols
23
ie. Na
11
Sodium-23 contains 11 electrons, 11 protons
and 12 neutrons.
24
Sodium 24: Na contains 11 electrons, 11 protons
11 and 13 neutrons. It is an
example of an isotope

45. Isotopes
• Atoms of the same element that contain an
identical number of protons but differ in the
number of neutrons
12 13 14
• C(carbon -12) C(carbon-13) C(carbon-14)
6 6 6
6e, 6p, 6n 6e, 6p, 7n 6e, 6p, 8n
- three different forms of carbon
- differ only in the number of neutrons in
the nucleus
- illustrates the concept of isotopes

46. More Isotopes
1 2 3
H H H
1 1 1
Hydrogen-1 Hydrogen-2 Hydrogen-3
deuterium tritium
• In nature elements are usually found as a mixture of isotopes
• An isotopes mass is determined by comparison to a standard,
carbon-12 (has a mass of 12 atomic mass units or amu)
• An elements atomic mass is obtained by taking a weighted
average of the atomic masses of all isotopes of that element
present in nature

47. Putting it all together
• How many protons , neutrons and electrons
are present in 96
Mo ? 42
Plan: The number of p, n and e are determined from
the atomic # and mass #
Atomic # = 42 and Mass # = 96
Atomic # = # protons → 42p
# protons = # electrons in a neutral atom → 42e
Mass # = # protons + # neutrons → 96 – 42 = 54n

48. In any room where chemistry is taught,
or practiced
you are certain to find
a chart called the……….

49. Periodic Table:
• A systematic catalog
of elements.
• Elements are
arranged in order of
atomic number.

50. The Periodic Table
• Shows all the known
elements
• Each box contains the
element’s atomic number
(# of protons = # of
electrons) written over the
one or two letter symbol
• Each box contains the
element’s atomic mass (#
of protons + # of neutrons)
written below the symbol

51. Symbols and much, much more
• Not only does the
Periodic table provide
symbols, atomic
number and atomic
mass of each element,
……
• The periodic table tell
us a good deal about
each element

52. As you peruse the table you cannot help
but note that:
• The elements are
arranged in order of
increasing atomic
number in nice
horizontal and vertical
columns

53. Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
reactivities.

54. Dimitri Mendeleev
• In 1869 he arranged the
elements based on
similarities in chemical
properties into
“families”
• Mendeleev listed the
families vertically and
called them groups
- groups are referred to by
the number over the
column
- many groups also have
names

55. • Mendeleev called the
horizontal rows periods
they are designated with
numbers
• Elements in a period each
have common
characteristics
• Each period ends with a
member of the family of
elements called the Noble
Gases
Noble Gases - chemically un-
reactive elements that exist
in nature as individual
atoms

56. Groups
These five groups are known by their names.

57. Periodic Table
Nonmetals are on the
right side of the
periodic table (with
the exception of H).

58. Periodic Table
Metalloids border the
stair-step line (with
the exception of Al
and Po).

59. Periodic Table
Metals are on the left
side of the chart.

60. Chemical Formulas
The subscript to the right of
the symbol of an element
tells the number of atoms
of that element in one
molecule of the compound.

61. Molecular Compounds
Molecular compounds are
composed of molecules and
almost always contain only
nonmetals.

62. Diatomic Molecules
These seven elements occur naturally as
molecules containing two atoms.

63. Types of Formulas
• Empirical formulas give the lowest whole-
number ratio of atoms of each element in a
compound.
• Molecular formulas give the exact number of
atoms of each element in a compound.

64. Types of Formulas
• Structural formulas show the
order in which atoms are
bonded.
• Perspective drawings also show
the three-dimensional array of
atoms in a compound.

65. Ions
• When atoms lose or gain electrons, they become
ions.
– Cations are positive and are formed by elements on
the left side of the periodic chart.
– Anions are negative and are formed by elements on
the right side of the periodic chart.

66. Ionic Bonds
Ionic compounds (such as NaCl) are generally
formed between metals and nonmetals.

67. Writing Formulas
• Because compounds are electrically neutral, one
can determine the formula of a compound this
way:
– The charge on the cation becomes the subscript on the
anion.
– The charge on the anion becomes the subscript on the
cation.
– If these subscripts are not in the lowest whole-number
ratio, divide them by the greatest common factor.

68. Common Cations

69. Important groups in the
Periodic Table
• Group IA – VIIA – known as The Main Group or
Representative Elements
Group IA- Alkali Metals
most highly reactive metals
Group IIA – Alkaline Earth Metals
highly reactive metals
Groups IIIA, IVA and VA - are not generally referred to by a
family name
Group VIA – Chalcogens
Group VIIA – Halogens
highly reactive nonmetals

70. Group VIIIA
• Group VIIIA – known as the Noble or Inert
Gases
- Noble Gases form few chemical compounds
- He, Ne, and Ar do not form any compounds…
all exist in nature as individual atoms

71. Group B Elements
• Called the Transition Metals
• They include many familiar metals ie. Fe, Cr, Ni, Sn…
• They also include the noble metals Cu, Ag, and Au (Grp IB)
-Noble metals are the rare metals of coins and jewelry.
-Noble metals are comparatively chemically inert to rust and corrosion

72. Inner Transition Metals
• Lanthanides – group between 57
La and 72
Ha
• Actinides – group between 89
La and 104
Rf

73. The Periodic Table at a Glance
• Allows us to classify an element very broadly into two
classes:
1. Metals
2. Non-metals
- the point of separation is the heavy stair step line
- B, Si, Ge, As, Sb and Te which border the line are called
metalloids or semimetals

74. Physical State and the
Periodic Table
• the periodic table tells us about an elements physical state at T= 25o
C or
standard reference temperature
• Except for hydrogen all gaseous elements are found at the extreme right
of the table
He
N O F Ne
Cl Ar
Kr
Xe
Rn
• There are only two liquids….metal Hg and non-metal Br2
• ****all other elements are solids

75. • Of the non-metals many exist as diatomic molecules
rather than individual atoms
• This includes all gaseous elements except the Noble
gases
H2, N2,O2, F2, Cl2, Br2, I2
• Natural form of P is P4
• Most common form of S is S8
• Carbon exists in three different forms called
allotropes
- all three allotropes of carbon have different
properties
- the 3 allotropes are diamond, graphite and
buckminsterfullerene

76. A molecule is an aggregate of two or more atoms in a
definite arrangement held together by chemical forces
H2 H2O NH3 CH4
A diatomic molecule contains only two atoms
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
2.5

77. Formation of Cations from a Neutral Atom
• Cations or positively charged ions are formed when
one or more e-
are lost from a neutral atom to create
a species with a + charge
-1e-
Na → Na+
+ 1e-
11p+
, 11e-
11p+
, 10e-
sodium ion
-2e-
Mg → Mg+2
+ 2e-
12p+
, 12e-
12p+
,10e-
magnesium ion
Cations are named by retaining the elements original name

78. Formation of Anions from Neutral Atoms
• Anions are negative ions formed when a neutral
atom gains extra electrons to create a species with a
- charge
Cl + 1e-
→ Cl –
17p+
, 17e-
17p+
, 18e-
chloride ion
• Anions are named by taking the root name of the
atom and changing the ending to –ide
Br -
bromide ion O2-
oxide ion S2-
sulfide ion

79. A monatomic ion contains only one atom
A polyatomic ion contains more than one atom
2.5
Na+
, Cl-
, Ca2+
, O2-
, Al3+
, N3-
OH-
, CN-
, NH4
+
, NO3
-

80. 13 protons, 10 (13 – 3) electrons
34 protons, 36 (34 + 2) electrons
Do You Understand Ions?
2.5
How many protons and electrons are in ?Al27
13
3+
How many protons and electrons are in ?Se78
34
2-

81. Ionic Compounds
• Compounds formed from the strong
electrostatic attraction between oppositely
charged ions (a cation and an anion)
• The strong electrostatic attraction between
oppositely charged ions is called an ionic bond
ie. Na+
+ Cl-
→ NaCl
Sodium Chloride
Formula represents the simplest ratio of cation to anion present in sodium
chloride. It is called a formula unit.

82. Formula Units
• For Ca+2
and Cl-
For every Ca +2
ion two Cl-
ions are needed
CaCl2 is the formula unit for calcium chloride

83. For ionic compounds, the formula is always the same as the
empirical formula
An empirical formula shows the simplest whole-
number ratio of the atoms in a substance
• the sum of the charges on the cation(s) and anion(s) in each
formula unit must equal zero
The ionic compound NaCl
2.6

84. Polyatomic Ions
• Groups of atoms with a charge
NH4
+
SO4
2-
NO3
-

85. Metals vs. Nonmetals
• Metals have a tendancy
to lose electrons and
form cations
• Most Non-metals
(except the Noble
Gases) gain electrons
and form anions
** The Periodic Table can tell us:
The charge of the cation formed by a metal
and
The charge of the anion formed by a non-metal

86. Group Numbers and Ion Charge
• Group I Metals- all form cations with a +1
charge
• Group II Metals- all form cations with a +2
charge
• Group III Metals- all form cations with a
+3 charge
 For Groups I-III charges of cation formed are
identical to the Group Number

87. Transition Metals
• All the transition metals form cations with
various positive charges
• There is no easy way to predict the charge of
the cation that will be formed
• Common charges of the transition metals are
found in Table 2.3 on pg 61 ….they must be
memorized

88. Non-Metals
• Non-metals form negative ions by gaining
electrons
• Group VII- all gain one electron to form
1-
ions
• Group VI- atoms of elements in this group
gain two electrons to form 2-
ions
• Group V- atoms gain 3 electrons forming
anions with a 3-
charge

89. It is important to Remember:
• Isolated atoms do not
form ions on their own
• Ions form when metallic
elements combine with
non-metallic elements to
form compounds called
ionic compounds
• Chemical compounds
must have a net charge of
zero
 Therefore if a compound
contains ions:
1. There must be both
cations and anions present
2. the numbers of cations and
anions must be such that
the net charge is zero
**These simple rules are integral to your ability to write
Formulas for Ionic Compounds

90. Formulas of Metal / Non-metal Binary
Ionic Compounds
• When writing the formula for a binary compound it is
important to remember:
Total charge + total charge → cmpd with zero
on cation on anion net charge
ie. Magnesium + chlorine → ?
Plan: use periodic table to find charges on ions of elements
• Metal and nonmetal combine to neutralize charge
• Always write the metal first and the non-metal second
• Use subscripts to indicate the relative number of ions or atoms
• Consider - Mg+2
, Cl1-
– cross multiply charges
– Mg2+
+ 2Cl1-
= MgCl2magnesium chloride

91. More Examples of
Metal /Nonmetal Compounds
• Barium + Oxygen → ?
– barium an alkaline earth metal - +2
– Ba - +2
– O - -2
• Ba+2
+ O-2
→ BaO (barium oxide)
• Lithium + Nitrogen → ?
- lithium an alkali metal - +1
- Li- +1
- N- +3
• Li +1
+ N-3
→ Li3N (lithium nitride)