Transcript Unit 5 Atomic Structure

Unit 5: Atomic
Structure
Trimble
CDO CP Chemistry
2014-2015
Dalton’s Atomic Theory
• 1803
– John Dalton linked the existence
of elements, which cannot be
decomposed chemically, to the idea
of atoms, which are indivisible.
Postulates
• All
matter is made of atoms. These
indivisible and indestructible objects are
the ultimate chemical particles.
• Atoms cannot be created or destroyed
• Atoms of the same element are alike in
every way
Postulates cont.
Atoms of different elements are different
• Atoms can combine together in small
numbers to form molecules (compounds)
•
Basic Components of the Atom
• There
are 3 fundamental subatomic
particles
• electron relative mass = 0.0005 amu
relative charge = -1
5
• proton
relative mass = 1 amu
relative charge = +1
• neutron
relative mass = 1 amu
charge = 0
Atomic Structure
• Nucleus
is found at the center of
the atom. Contains the protons
and neutrons.
• The number of
protons in the nucleus
determines
what element an
atom is.
6
Atomic Structure
Electrons are found in
the space around the
nucleus.
7
Electrically Neutral Atoms
• For
any neutral atom, the number
of negatively charged electrons
around the nucleus equals the
number of positively charged
protons in the nucleus
8
Important Note
• To
chemists the electrons are
the most important part of
the atom, because they are
the first part of the atom that
contacts another atom.
9
Isotopes
10
Atomic Number
• All
atoms of the same element
have the same number of
protons.
• This number is called the atomic
number and is given the symbol
Z.
• The atomic number is the whole
number in each element box on
the periodic table.
11
Mass Number
• The
sum of the number of
protons and neutrons in an
atom is called the mass number
and is represented with the
symbol A.
• This
number is not listed on the
periodic table
12
Isotopes
• Atoms
having the same atomic
number Z but a different mass
number A.
• Isotopes are atoms of the same
element, but they have
different masses because they
have different numbers of
neutrons.
13
Nuclear Symbol of an Atom
14
Hyphen Notation
• Hyphen
notation reports the
element name, then the
mass number of the isotope
after a dash.
Helium-4
15
Example: Counting Sub Atomic
Particle Isotopes
How many protons and
electrons are in found in each of
the following neutral elements:
• Li
•P
• Ag
• Xe
•U
16
Example: Counting Sub Atomic
Particle Isotopes
• Determine
the number of protons,
electrons, and neutrons in each of the
following:
17
Example: Counting Sub Atomic
Particle Isotopes
• Determine
the number of protons,
electrons, and neutrons in each of the
following:
• Oxygen-18
• Calcium-40
• Lead-204
• Mercury-196
18
Example: Counting Sub Atomic
Particle Isotopes
Write the nuclear symbol and hyphen
notation name for each of the following
elements:
• 7 protons, 7 electrons, 8 neutrons
• 28 protons, 28 electrons, 30 neutrons
• 47 protons, 47 electrons, 62 neutrons
• 50 protons, 50 electrons, 64 neutron
19
Ions
20
Ion Formation
• Atoms
can gain or lose electrons to
form ions
• Cation – an atom that has lost an
electron and forms a positive
charge
• Anion – an atom that has gained
an electron and is negatively
charged
21
Ionic Symbols
• Ions are indicated by writing the
element symbol with charge value
written in the upper right hand
corner
• The number one is never written
+
•Na
2•O
22
Ion Formation
To determine the charge on
an ion subtract the number
electrons from the number of
protons
Charge = protons - electrons
23
Example: Counting Sub Atomic
Particle - Ions
•
Give the number of protons and electrons in
the following:
+
• Li
3+
• Al
3•N
4+
• Pb
• Br
24
Example: Counting Sub Atomic
Particle - Ions
• Write
the ionic symbol for each of the
following ions.
• 17
protons, 18 electrons
• 38 protons, 36 electrons
• 16 protons, 18 electrons
• 31 protons, 28 electrons
25
Example: Counting Sub Atomic
Particle - Ions
Determine the number of protons,
electrons, and neutrons in the following:
O
64
30
Zn
120
50
26
2
16
8
Sn
4
Example: Counting Sub Atomic
Particle - Ions
•
Write the nuclear symbol, with charge is
needed, for the following:
• 15
protons, 18 electrons, 17 neutrons
• 12 protons, 10 electrons, 14 neutrons
• 16 protons, 16 electrons, 20 neutrons
• 35 protons, 36 electrons, 46 neutrons
27
Average
Atomic
Mass
28
•Average Atomic Mass
• The
weighted average of
the naturally occurring
isotopes of an element.
• Found by averaging the
natural abundances of
its isotopes
•Calculating Average Atomic Mass
(amu)
• To
calculate the average atomic mass
multiply the percent abundance of an
isotope by its mass or mass number.
• Do this for each isotope and then add the
results together
(Mass of Isotope 1)(% abundance)  (Mass of Isotope 2)(% abundance)  ....
100
•Average Atomic Mass
Rubidium has two common isotopes, Rb-85
and Rb-87. If the abundance of 85Rb is
72.2% and the abundance of 87Rb is 27.8%,
what is the average atomic mass of
rubidium?
Uranium has three common isotopes. If
the abundance of 234U is .01%, the
abundance of 235U is .71%, and the
abundance of 238U is 99.28%, what is the
average atomic mass of uranium?
Calculating Atomic Masses
• For boron, 19.9% occurs as 10B and
80.1% occurs as 11B. The isotopic
mass of 10B is 10.013 and 11B is
11.009
33
Practice
• Gallium
consists of two isotopes: Ga-69
with a mass of 68.9256 amu accounts
for 60.11% and Ga-71 with a mass of
70.9247 amu accounts for the rest.
What is the average mass of gallium?
34
Electron
Arrangement
Electromagnetic spectrum
• Electromagnetic
radiation – comes in
differing forms
• All EMR travels at the same speed (c)
but can be distinguished by their
different wavelengths (l)
36
Electromagnetic Radiation
•Visible Spectrum
38
•Wavelengths of Visible Spectrum
• Purple
– 400 nm
• Blue – 475 nm
• Green – 510 nm
• Yellow – 570 nm
• Orange – 590 nm
• Red – 650 nm
39
Electromagnetic spectrum
• Frequency
(f) – the number of waves
which pass through a particular point.
•
Measured in Hertz (Hz), which is equal to
1/s.
• Wavelength
(l)  the distance between
crests in a wave.
• The shorter the l the higher the
frequency (f)
40
Calculating Wavelength or Frequency
c=fl
C = 3.0 x 108 m/s
41
•Practice
• The
distinctive green color of Aurora
Borealis is caused by the interaction of
the radiation with oxygen and has a
frequency of 5.38 x 10 14 Hz. What is the
wavelength of this light?
42
• Light
with a frequency of 7.26 x 1014 Hz
lies in the violet region of the visible
spectrum. What is the wavelength of
this frequency of light?
43
•A
certain violet light has a wavelength
of 413 nm. What is the frequency of this
light?
44
Energy of Light
• The
energy of light can be calculated
using the following equation
E = hf
E = energy of light in Joules (J)
h = Planck’s constant = 6.626 x 10-34
f = frequency
45
Energy of Light
• Calculate
the energy of a photon of
radiation with a frequency of 8.5 x 1014
Hz.
46
• Calculate
the energy of a photon of
radiation with a wavelength of 6.4 x 10-7
m.
47
• Calculate
the energy of a photon of
radiation with a wavelength of 6.4 x 10-7
m.
48
• What
is the energy of light whose
wavelength is 4.06 x 10-11 m?
49
ELECTRON CONFIGURATION
50
•Orbital Diagram
• An
orbital is a potential
space for an electron.
• Orbitals are represented
by boxes or lines
grouped by sublevel
with small arrows
indicating the electrons.
•Pauli Exclusion Principle
• An
atomic orbital can hold a maximum
of 2 electrons and those 2 electrons
must have opposite spins.
• An electron is represented by an arrow.
• Spin is represented by the arrow facing
up or down.
•Aufbau Principle
• Electrons
are placed
the lowest energy
level first.
in
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p
•Hund’s Rule
• When
filling sublevels other than s,
electrons are placed in individual
orbitals first, before they are paired up.
• They must be placed singly before
doubly.
•Sublevels
• The
“s” sublevel can hold 2 electrons.
• The “p” sublevel can hold 6 electrons.
•
2 electrons in each of the 3 orbitals (x, y, z)
• The
•
2 electrons in each of the 5 orbitals.
• The
•
“d” sublevel can hold 10 electrons.
“f” sublevel can hold 14 electrons.
2 electrons in each of the 7 orbitals.
•Practice Problems
• Write
the orbital diagram for Fluorine.
•Practice Problems
• Write
the orbital diagram for
Magnesium.
•Electron Configuration
• Shows
the arrangement of electrons in
an atom.
Electron Configuration
• The
rules for electron configurations is
the same as orbital diagrams
• However…
• Instead of drawing in orbitals and
arrows, write the number of electrons in
the sublevel in superscript after the sub
level name.
59
•Practice Problems
• Write
the electron configuration and
the orbital diagram for Sulfur.
•Practice Problems
• Write
the electron configuration and
the orbital diagram for Potassium.
•Organization of Orbitals
• The
periodic table has organized the
orbitals.
• The “s”
orbitals
• The “p”
orbitals
• The “d”
orbitals
• The “f” orbitals
•Organization of Orbitals
The first row is Principal Energy Level 1.
• The second row is Principal Energy Level 2.
• Principal
Energy Level
3 begins in
the 3rd row.
• Principal
Energy Level 4 begins
in the 4th row.
•
•and so the pattern continues…
•Noble Gas Configuration
• Is
an abbreviated version of electron
configuration.
• Uses the noble gas that precedes the
element, then the electron
configuration that comes after the
noble gas.
• Used for elements with larger atomic
numbers.
• Example:
Nitrogen
•Noble Gas Configuration
• Is
important because it shows the
valence electrons present in an atom.
• Nitrogen
has an atomic number of 7. It
has 7 total electrons. If you look at the
electron configuration, you can count
7 electrons.
•Noble Gas Configuration
• But
if you look at the Noble Gas
Configuration, you can count 5 electrons.
• These
5 electrons are the valence
electrons, the electrons found in the
outermost energy level. These are the
electrons available for bonding.
•Valence Electrons
• The
periodic table organizes valence
electrons.
• The number
of valence
electrons
are written
above each
column in
the diagram.
•Practice Problems
• Write
the noble gas configuration and
the orbital diagram for Iron.
•Practice Problems
• Write
the noble gas configuration and
the orbital diagram for Tin.
IONIZATION ENERGY
•
•
•
71
Ionization energy, Ei: minimum energy
required to remove an electron from the
ground state of an atom in the gas phase.
M(g) + h  M+ + e.
Ei related to electron arrangement
Sign of the ionization energy is always
positive.
Ionization Energy: Periodic Table
Ionization Energy vs atomic #
72
Successive Ionization Energies
• Atoms
can lose more than one electron
• Each electron lost requires more energy
than the one before it
• Electrons removed from levels closer to
the nucleus requires more energy
73
Ionization Energy of Aluminum
74