AHS Chemistry Unit 2 Atomic Structure

1
AHS Chemistry Unit 2Atomic Structure

• The Basics of Atoms and the Periodic Chart

2
Democritus
Nothing exists except atoms and empty space
everything else is opinion
Born about 460 BC in Abdera, Thrace,
GreeceDied about 370 BC
atomic theory Atomic theory is the theory that
matter and space are not infinitely divisible.
3
Atomists hold that the absolute existents are the
ultimate parts or components of visible bodies,
though these absolute existents can exist
separately from these visible bodies. These
existents are completely full or solid they
contain no gaps, no holes, no empty places. They
cannot be subdivided, cut, cracked, split, or
penetrated for this would imply that nonbeing
has gained a foothold in them. These absolute
existents are called atoms, from the Greek word
atomon, which means "that which cannot be
divided." Atoms do not come into being or pass
away for if they come into being (or pass away),
then they would have to partially exist and
partially not exist but partial existence makes
no sense for ultimate things
4
John Dalton
1803 John Dalton proposes the atomic theory of
matter.
5
1780's, Lavoisier ushered in a new chemical era
by making careful quantitative measurements which
allowed the compositions of compounds to be
determined with accuracy
By 1799 enough data had been accumulated for
Proust to establish the Law of Constant
Composition ( also called the Law of Definite
Proportions).
In 1803 Dalton noted that oxygen and carbon
combined to make two compounds. Of course, each
had its own particular weight ratio of oxygen to
carbon (1.331 and 2.661), but also, for the
same amount of carbon, one had exactly twice as
much oxygen as the other. This led him to propose
the Law of Simple Multiple Proportions, which was
later verified by the Swedish chemist Berzelius.
6
In an attempt to explain how and why elements
would combine with one another in fixed ratios
and sometimes also in multiples of those ratios,
Dalton formulated his atomic theory.
Assumption One His atomic theory, stated that
elements consisted of tiny particles called atoms
Assumption Two He said that the reason an
element is pure is because all atoms of an
element were identical and that in particular
they had the same mass
Assumption Three He also said that the reason
elements differed from one another was that atoms
of each element were different from one another
in particular, they had different masses
7
Assumption Four Atoms of different elements may
have nearly identical masses Assumption
Five Atoms are not subdivided in a chemical
reaction
On Compounds He also said that compounds
consisted of atoms of different elements combined
together Compounds have constant composition
because they contain a fixed ratio of atoms and
each atom has its own characteristic weight, thus
fixing the weight ratio of one element to the
other.
In addition he said that chemical reactions
involved the rearrangement of combinations of
those atoms.
8
Dalton's model was that the atoms were tiny,
indivisible, indestructible particles and that
each one had a certain mass, size, and chemical
behavior that was determined by what kind of
element they were.
9
Atoms Molecules

• Atoms
• can exist alone or enter into chemical
  combination
• the smallest indivisible particle of an element
• Molecules
• a combination of atoms that has its own
  characteristic set of properties

10
Law of Conservation of Matter

• In an ordinary chemical reaction matter is
  neither created nor destroyed.
• The sum of the masses of the reactants equals the
  sum of the masses of the products.

11
Law of Constant Composition

• A chemical compound always contains the same
  elements in the same proportions by mass.

Law of Multiple Proportions compounds can be
formed by the same elements combining in
different ratios. For Example CO2 and CO or H2O
and H2O2
Movie
12
The particles in the beam are negative, what will
happen as they pass through the electrically
charged plates?
13
The cathode ray is Negative, describe what is
happening here?
14
Another look at the Cathode Ray set up.
15
Describe what is happening in each situation a,
b, and c.
16
Watch the Movie
Thomson discovered the ratio of charge over mass
of an electron to be 1.7588196 x 1011 C kg-1.
Later, the concept of the electron was furthered
by the American physicist Robert Millikan, who
defined e, the charge of the electron, as
1.6021773 x 10-19 Coulombs. Knowing e and the
ration e/me, the mass of the electron was found
to be 9.109390 x 10-31kg Watch the Movie
17
Animated Model
Movie
18
Animated Model
19
Rutherfords Model of the Atom

• atom is composed mainly of vacant space
• all the positive charge and most of the mass is
  in a small area called the nucleus
• electrons are in the electron cloud surrounding
  the nucleus

20
Structure of the Atom Composed of

• protons
• neutrons
• electrons

21
Structure of the Atom

• Composed of
• protons
• neutrons
• electrons

• protons
• found in nucleus
• relative charge of 1
• relative mass of 1.0073 amu

22
Structure of the Atom

• Composed of
• protons
• neutrons
• electrons

• neutrons
• found in nucleus
• neutral charge
• relative mass of 1.0087 amu

23
Structure of the Atom

• Composed of
• protons
• neutrons
• electrons

• electrons
• found in electron cloud
• relative charge of -1
• relative mass of 0.00055 amu

24
Animated Model
25
Radioactivity
A movie about radioactive particles!
26
Alpha Radiation

• composed of 2 protons and 2 neutrons
• thus, helium-4 nucleus
• 2 charge
• mass of 4 amu
• creates element with atomic number 2 lower

27
Beta Radiation

• composed of a high energy electron which was
  ejected from the nucleus
• neutron converted to proton
• very little mass
• -1 charge
• creates element with atomic number 1 higher

28
Gamma Radiation

• nucleus has energy levels
• energy released from nucleus as the nucleus
  changes from higher to lower energy levels
• no mass
• no charge

29
Ions

• charged single atom (monatomic ions)
• charged cluster of atoms (polyatomic ions)

Atoms become ions when they gain or lose an
electron. When Na gives up an electron it has one
more protons than electrons thus you get
Na1 When Cl gains an electron you have one more
electron than protons thus you get Cl-1
30
Ions

• cations
• positive ions
• anions
• negative ions
• ionic compounds
• combination of cations and anions
• zero net charge

31
Atomic number, Z

• the number of protons in the nucleus
• the number of electrons in a neutral atom
• the integer on the periodic table for each element

32
Isotopes

• atoms of the same element which differ in the
  number of neutrons in the nucleus
• designated by mass number

33
Mass Number, A

• integer representing the approximate mass of an
  atom
• equal to the sum of the number of protons and
  neutrons in the nucleus

34
Masses of Atoms

• Carbon-12 Scale

35
Isotopes of Hydrogen H-1, 1H, protium

• 1 proton and no neutrons in nucleus
• only isotope of any element containing no
  neutrons in the nucleus
• most common isotope of hydrogen

36
Isotopes of Hydrogen H-2 or D, 2H, deuterium

• 1 proton and 1 neutron in nucleus

37
Isotopes of Hydrogen H-3 or T, 3H, tritium

• 1 proton and 2 neutrons in nucleus

38
(No Transcript)
39
Isotopes of Oxygen

• O-16
• 8 protons, 8 neutrons, 8 electrons
• O-17
• 8 protons, 9 neutrons, 8 electrons
• O-18
• 8 protons, 10 neutrons, 8 electrons

40
The radioactive isotope 14C has how many
neutrons? 6, 8, other
41
The identity of an element is determined by the
number of which particle? protons, neutrons,
electrons
42
The mass spectrometer measures the relative
atomic mass of atoms
43
The percentage of atoms at certain masses are
collected by the mass spectrometer and the
abundance of each isotope is then known by the
frequency of the hits on the collector plate at
different mass numbers
44
Measurement of Atomic Masses

• Mass Spectrometer
• a simulation is available at
• http//www.colby.edu/chemistry/
• OChem/DEMOS/MassSpec.html

45
Atomic Masses andIsotopic Abundances

• natural atomic masses
• sum(atomic mass of isotope)
• (fractional isotopic abundance)

46
Amadeo Avagadro Amadeo Avagadro was born on June
9, 1776 in Turin - the capital city of the
independent nation of Piedmont. He began his
career as a lawyer, earning his doctorate of law
in 1796. With Gay-Lussac's discovery in 1802 that
all gases expand at the same rate with a change
in temperature, Avogadro gave up law and turned
to the world of science. In 1811, he showed that
Dalton's Atomic theory and Gay-Lussac's Gas Law
could exist together if there was, at a constant
temperature and pressure, an equal number of
particles in a given volume of a gas. This became
known as Avogadro's Hypothesis. In developing
this idea, he stated that the particles did not
need to be separate atoms but could exist as a
combination of atoms, which he called a molecule.
He also proposed that certain gases (hydrogen and
oxygen) existed as diatomic molecules.
47
Unfortunately, Avogadro's theories were not
readily accepted due to the new terminology he
used and the concepts developed by Berzelius,
which were the accepted norm at the time. Not
until Stanislao Cannizzaro presented Avogadro's
theories at a convention of chemist in 1860 did
the world begin to accept Avogardro's Hypothesis.
Unfortunately Avogadro did not live to receive
this acclaim, dying in obscurity on July 9, 1856
in his hometown of Turin.
48
The Mole

• a unit of measurement, quantity of matter present
• Avogadros Number
• 6.022 x 1023 particles
• Latin for pile

49
(No Transcript)
50
(No Transcript)
51
Molar Mass

• Sum
• atomic masses
• represented by formula
• atomic masses gt gaw (gram atomic weight)
• molar mass gt MM

Atomic mass of Oxygen 16g/mol Molar Mass of
water H2O H H O or 1g/mol 1g/mol
16g/mol 18g/mol
52
Dimensional Analysis
dimensional analysis a general problem-solving
method that uses the units associated with
numbers as a guide in setting up the calculation.
53
How many meters in 6.20 miles?
We can use the density of a material as a
conversion factor between mass and volume rather
than a formula in this manner The density of
mercury is 13.6 g/cc. What is the mass of two
liters of mercury?
54
W5P METHOD (WILD WONDERFUL WAY TO WORK WORD
PROBLEMS) GIVEN - List all pertinent information
with dimension symbol, number and unit. FIND -
List the dimension of the quantity requested in
problem. FORMULA - With the dimensions in GIVEN
and FIND, list the formula of formulas that fit.
SOLVE - Solve the formula for what you are
looking for (FIND), substitute the number values
in GIVEN, and perform the math on both the units
and the numbers. ANSWER - Check the answer for
likeliness, make sure the units are appropriate,
express the answer in scientific notation and to
the accuracy required, and draw a box around it
so it is obvious which number your answer is.
55
MoleculeMoleGramConversions
56
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.

57
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• state questions in mathematical form
• S atoms

58
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• set equal to quant. data particular to problem
• S atoms (1.000g)

59
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• multiply by conversion factor
• S atoms (1.000g)(1 mole/ 32.07 g)

60
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• cancel units
• S atoms (1.000g)(1 mole/ 32.07 g)

61
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• cancel units
• S atoms (1.000)(1 mole/ 32.07)

62
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• another conversion factor
• S atoms (1.000)(1 mole/ 32.07)
• (6.022x1023atoms/mole)

63
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• cancel units
• S atoms (1.000)(1 mole/ 32.07)
• (6.022x1023atoms/mole)

64
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• cancel units
• S atoms (1.000)(1 / 32.07)
• (6.022x1023atoms)

65
Example

• Calculate the number of sulfur atoms in 1.000g
  sample of the element.
• solve mathematics
• S atoms (1.000)(1 / 32.07)
• (6.022x1023atoms)
• 1.878 x 1022 S atoms

66
Example

• How many moles of silicon, Si, are in 30.5g of Si?

67
Example

• How many moles of silicon, Si, are in 30.5g of
  Si?
• mol Si

68
Example

• How many moles of silicon, Si, are in 30.5g of
  Si?
• mol Si (30.5g)

69
Example

• How many moles of silicon, Si, are in 30.5g of
  Si?
• mol Si (30.5g)(1 mol/28.0855g)

70
Example

• How many moles of silicon, Si, are in 30.5g of
  Si?
• mol Si (30.5g)(1 mol/28.0855g)

71
Example

• How many moles of silicon, Si, are in 30.5g of
  Si?
• mol Si (30.5)(1 mol/28.0855)
• 1.09 mol Si

72
Example

• How many grams of Cu are there in 2.55 mol Cu?

73
Example

• How many grams of Cu are there in 2.55 mol Cu?
• g Cu

74
Example

• How many grams of Cu are there in 2.55 mol Cu?
• g Cu (2.55 mol)

75
Example

• How many grams of Cu are there in 2.55 mol Cu?
• g Cu (2.55 mol)(63.546g/mol)

76
Example

• How many grams of Cu are there in 2.55 mol Cu?
• g Cu (2.55 mol)(63.546g/mol)

77
Example

• How many grams of Cu are there in 2.55 mol Cu?
• g Cu (2.55)(63.546g)
• 162 g

78
Development of Periodic Table

• Newlands - English
• 1864 - Law of Octaves - every 8th element
  has similar properties

79
Development of Periodic Table

• Mendeleev - Russian
• 1869 - Periodic Law - allowed him to
  predict properties of unknown elements

80
Mendeleevs Periodic Table

• the elements are arranged according to increasing
  atomic weights

81
Missing elements 44, 68, 72, 100 amu
82
(No Transcript)
83
Modern Periodic Table

• Moseley, Henry Gwyn Jeffreys
• 18871915, English physicist.
• Studied the relations among bright-line spectra
  of different elements.
• Derived the ATOMIC NUMBERS from the frequencies
  of vibration of X-rays emitted by each element.
• Moseley concluded that the atomic number is equal
  to the charge on the nucleus.
• This work explained discrepancies in Mendeleevs
  Periodic Law.

84
Modern Periodic Table

• the elements are arranged according to increasing
  atomic numbers

85
Periodic Table of the Elements
86
Organization of Periodic Table

• period - horizontal row
• group - vertical column

87
Family Names

• Group IA alkali metals
• Group IIA alkaline earth metals
• Group VIIA halogens
• Group VIIIA noble gases
• transition metals
• inner transition metals
• lanthanum series rare earths
• actinium series trans-uranium series

88
Types of Elements

• metals
• nonmetals
• metalloids - semimetals

89
Niels Bohr, the Bohr Model
Bohr built upon spectroscopic observations of
atoms. Spectroscopists noticed that an atom can
only absorb certain energies (colors) of light
(the absorption spectrum) and once excited can
only release certain energies (the emission
spectrum) and these energies happen to be the
same. Bohr used these observations to argue that
the energy of a bound electron is "quantized."
Animated Model
90
Electron Cloud
Colorado Physics Site
Erwin Schrödinger built upon the thoughts of Bohr
yet took them in a new direction. He developed
the probability function for the Hydrogen atom
(and a few others). The probability function
basically describes a cloud-like region where the
electron is likely to be found. It can not say
with any certainty, where the electron actually
is at any point in time, yet can describe where
it ought to be. Clarity through fuzziness, is one
way to describe the idea. The model based on this
probability equation can best be described as the
cloud model.
91
The Current View of the Atom
Its not so much like a solar system but a
beehive of electron activity
92
Mass is concentrated in a nucleus which is
1/10,000 the over-all size of the atom and
consists of protons and neutrons. Most of an
atom's volume is filled by wandering electrons.
93
Atomic Number
Atoms attract or repel electrons until they have
no unbalanced charge. The number of protons in
the nucleus equals the number of electrons in a
neutral atom and determines the chemical identity
of the atom. Number of protons Atomic Number
Z
94
Each value of Z, corresponds to a unique chemical
element. Z1 Hydrogen --- an electron orbiting
a single proton. Z-1 Deuterium --- chemically
identical to hydrogen. Z2 Helium-3 --- two
tightly bound electrons orbiting two protons and
a neutron. Z2 Helium-4 --- the normal form of
Helium. Z3 Lithium-7 --- two tightly bound
electrons and one loosely bound electron orbiting
three protons and four neutrons.
95
The Mass of Protons, Neutrons and Electrons
96
Atomic mass
Mass- The mass of an atom, expressed in atomic
mass units (AMU), is roughly equal to the number
of protons plus the number of neutrons. This is
because both the protons and the neutrons in an
atom have a relatively equal mass. The mass of an
electron is so insignificant that it is not
represented in the atomic mass. Since not all
atoms have only one
97
Average Atomic Mass
isotope1, the atomic mass is the average of all
isotopes, once abundance is computed. For
example, if you took a container of the element
hydrogen (H), 99.984 of it would be H-1, 0.0156
of it would be H-2, and 0 of the hydrogen would
be H-3. Since H-1 has one proton and no neutrons,
its mass is 1. Because H-2 has one proton and one
neutron, its mass is 2. Therefore, when you
compute the percentages of the isotopes of H in
any container, you find that the atomic mass of H
is actually 1.0079. If the atomic mass of a
particular element is shown in parentheses, such
as (145) for Promethium (Pm), the atomic mass
reflects that of the most stable isotope1, and is
not the average atomic mass for all isotopes of
the element.
98
Electron Arrangement
Ground state - atoms whose electrons are in their
normal / lowest energy state.
Excited state - atoms having higher potential
energy than normal (electrons have moved to a
higher orbital)
99
How do you excite an electron?
http//www.colorado.edu/physics/2000/applets/schro
edinger.html
Spectrum of Hydrogen We can look at either
absorption or emission spectra
100
Using our equation for the energy of the hydrogen
levels we can write an equation for the change in
energy of an electron that charges orbitals and
emits or absorbs a photon.
101
With this equation we can calculate the frequency
of light emitted or absorbed when an electron
moves between orbitals of different principal
quantum numbers
102
The Visible Spectrum Typically, superheated
solids and liquids will emit light of all colors.
We see it as white light but if we use a
spectroscope or prism to separate the light by
different frequencies, we see it as a spectrum
like the one to the left.
103
The Absorption Spectrum When atoms are present
as a gas, we find that they will absorb only
certain colors of light. So if we shine white
light through a gas then view it with a
spectroscope we find that there are some colors
missing. These missing colors are the colors that
were absorbed by the gas as the light passed
through it. Different atoms absorb different
colors. This technique is used to identify what
the outer layers of stars are made of.
104
The Emission Spectrum When we excite an atom,
we find that it gives off (emits) only certain
colors of light. Notice the position and colors
of the absorption spectrum line up with those of
the emission spectrum. This means that the colors
of light emitted by an atom are the same colors
that can be absorbed by that atom.
105
The Bohr Model with emission spectrum frequencies.
106
Heisenberg Uncertainty Principle
When examining electrons the more precisely the
position is determined, the less precisely the
momentum is known in this instant, and vice
versa.--Heisenberg, uncertainty paper, 1927
107
Quantum Theory

• Rules governing the combinations of quantum
  numbers
• Three quantum numbers n, l, and m are integers.
• The principal quantum number (n) cannot be zero.
• n must be 1, 2, 3, etc.
• The angular quantum number (l) can be any integer
  between 0 and n - 1.
• For n 3, l can be either 0, 1, or 2.
• The magnetic quantum number (m) can be any
  integer between -l and l.
• For l 2, m can be either -2, -1, 0, 1, or 2.
• The spin quantum number (s) is arbitrarily
  assigned the numbers 1/2 and -1/2

108
Quantum Numbers Numbers that describe the
energies of electrons in atoms derived from
quantum mechanical treatment.
Energy-Level Diagram for the 1-D Electron
109
n3, l2, m2
n1, l0, m0
n3, l2, m1
Quantum Orbitals
n4, l2, m2
110
Quantum Orbital Shapes s, p, and d
Orbital Axis, x, y, and z
111
The P orbitals, Px, Py, and Pz
112
The S Orbital or Spherical Orbital
113
Electron Configurations and the Periodic Table
When electron configuration data are arranged so
that we can compare elements in one of the
horizontal rows of the periodic table, we find
that these rows typically correspond to the
filling of a shell of orbitals. The second row,
for example, contains elements in which the
orbitals in the n 2 shell are filled. Li (Z
3) He 2s1 Be (Z 4) He 2s2 B (Z
5) He 2s2 2p1 C (Z 6) He 2s2 2p2 N (Z
7) He 2s2 2p3 O (Z 8) He 2s2 2p4 F
(Z 9) He 2s2 2p5 Ne (Z 10) He 2s2
2p6