Download Structure of the Atom

General Chemistry
Name:
Period:
Unit 3 Note Packet – Atomic Structure, Nuclear Processes, and the Periodic Table
History of Atomic Theory
Carefully read pages 101 to 108 and 127 to 129 (time-line) in your textbook. Answer the following questions
as you go. You are responsible for learning this material on your own; it will not be covered during lectures.
Greeks(~450 BC)
What was Democritus’ main contribution to atomic theory?
Other important Greek philosophers disagreed with Democritus. What did they think?
John Dalton (1766-1844)
What are the four postulates of Dalton’s atomic theory of matter?
a.
b.
c.
d.
Benjamin Franklin (1706-1790)
Benjamin Franklin investigated electricity. He concluded there are two kinds of electric charges an object can
have and he called them positive and negative. He concluded that like (same) charges repel each other while
opposite charges attract each other.
Michael Faraday (1791-1867)
Early scientists wondered where these charges come from and had questions about their physical properties.
Michael Faraday thought that the structure of atoms
was related to electricity. Faraday and other
scientists investigated this possibility using a cathode
ray tube.
Cathode Ray Tube (mid-1800s)
Describe a cathode ray tube and draw a diagram of
one:
Scientists found that a kind of radiation streams from the
to the
of a
cathode ray tube. One experiment showed that a cathode ray could spin a small paddle wheel, which suggested
that it was actually a
. They also discovered that a magnet deflects the
cathode ray in the direction expected for
charged particles.
JJ Thomson (1856-1940)
In 1896, JJ Thompson began to systematically study cathode rays. He measured the ratio of electricity charge
to mass of the electrons flowing through a cathode ray tube by deflecting the beam of electrons with a magnetic
field. Thomson discovered that magnetic and electric fields deflected the ray’s path in a mathematically
predictable way. Thomson concluded that a cathode ray is composed of
.
This meant that atoms were not indivisible balls but instead had a substructure. Thomson named these negative
582731445
-1-
particles
. He determined the ratio of an electron’s electrical charge to its mass, which is
.
JJ Thomson also developed the “plum pudding” model of an atom (see pp. 106 and 128). In this model, negative
charges (electrons) are distributed evenly throughout an atom’s positively charged interior, similar to the way
chocolate chips are distributed throughout cookie dough.
Robert Millikan (1868-1953)
Millikan was able to determine the quantity of charge carried by the electron. He was able to balance the oil
drop against the force of gravity to determine that charge.
Millikan further built on Thomson’s experiments by using the values Thomson obtained for the charge-to-mass
ratio. With these values he was able to successfully calculate the mass of a single electron by using his oil
droplet experiment.
The electron carries exactly
unit of
Millikan calculated the mass of an electron to be
charge.
.
Ernest Rutherford (1871-1937)
Studying the nature of radiation gave scientists further clues about the substructure of the atom. Ernest
Rutherford continued the study of radiation.
In 1909, Rutherford and his colleagues performed a famous experiment called the gold foil experiment. In this
experiment, alpha particles (with a +2 charge) from radioactive polonium were aimed at a thin piece of gold foil.
Most of the particles passed through the foil, but some were scattered in all directions. This led Rutherford to
reject Thomson’s plum pudding model. Instead, he proposed that
.
What name, originally proposed by Rutherford, is still used to describe the core of an atom?
Draw a diagram showing the apparatus used
the gold foil experiment.
in
Draw a diagram showing the details of what
went on in the gold foil experiment.
582731445
-2-
Atom – the smallest part of matter that retains the characteristics of that type of matter




considered to be the building blocks of matter
atoms consist of a positively charged center, or nucleus, surrounded by negatively charged particles
nucleus consists of protons and neutrons
negatively charged particles are called electrons existing outside the nucleus in mathematical “regions
of probability”
particle
proton
neutron
electron
mass (g)
mass (amu)
relative charge
location in atom
Types of Atoms  elements











each different type of atom is an element
elements (different types of atoms) are distinguished by the # of protons
every atom of the same element has the same # of protons (ex: every atom of hydrogen has one
proton, every atom of boron has 5 protons, etc.)
the # of protons =
elements are organized on the periodic table according to atomic number
each element has a chemical symbol
the symbol is the first letter of an element’s name capitalized and may have another letter from the
name in lower case (examples: C for carbon and Al for aluminum)
some symbols are derived from Latin names (example: Argentum is Latin for silver; the symbol for
silver is Ag)
# protons = # electrons in a neutral atom
# protons ≠ (do not equal) # electrons in a charged (positive or negative) atom (called an ion)
# neutrons can vary (not the same for every atom of an element)
Ions  the result of a neutral atom gaining or losing an electron
 cation 
 anion 
 charge of an ion = number of protons – number of electrons
Isotopes atoms that have the same number of protons but different numbers of neutrons
 for example…..hydrogen has three isotopes:
hydrogen 1
hydrogen 2 (deuterium)
hydrogen 3 (tritium)
Isotope notation:
Mass
number
Charge
15N37
Atomic
number
582731445
Example
7 protons
8 neutrons (15-7)
10 electrons (7+3)
-3-
Isotopes are named by their Mass number. In the example above the isotope is:
Nitrogen 15
or it could also be written as
Nitrogen-15
But if you look at your periodic table, the atomic mass of nitrogen is 14.01 atomic mass units. Rounding that to
the nearest whole number, we see that the most common isotope of nitrogen is nitrogen-14. The example just
happens to be talking about nitrogen-15.
What is the most common isotope of Phosphorous? (round what you see on the periodic table)
What is the most common isotope of Boron?
Use your periodic table to write the symbols for the following:
Mass
number
P
__ protons
0
Charge
B
__ neutrons
Atomic
number




Mass
number
Charge
__ protons
0
__ neutrons
Atomic
number
__ electrons
__ electrons
mass number = # protons plus (+) neutrons
atomic mass or atomic weight takes into account all the different isotopes
atomic mass/weight is an AVERAGE of all the different isotopes
an example calculation for atomic mass is show below:
107
Ag and 48.65%
108
Example:
A sample of silver is 51.35%
atomic mass/weight?
Ag. What is its average
Answer:
0.5135 x 107 = 54.9445
0.4865 x 108 = 52.542
Total
= 107.4865 amu = average atomic mass/weight
(107.49 amu)
CHANGES IN THE NUCLEUS
 chemical reactions involve interactions between outer (valence) electrons
 nuclear reactions change the composition of the atom’s nucleus
radiation =
Nuclear stability
Most atoms have a stable nucleus  ____________
Protons are positively charged  like charges repel
 question: why doesn’t the nucleus fly apart since all the protons are repelling each other?
 Answer: another force, called the ________________________, is holding the nucleus together
Strong nuclear force = the force that helps to hold the nucleus together (not a force we encounter in daily
life  only exists between subatomic particles that are extremely close)
Protons
Neutrons
582731445
- have both repulsive forces due to like charges, and the strong nuclear force
- have only the strong nuclear force acting on them
- help to balance the electromagnetic repulsion that the positive protons experience
- help to hold the nucleus together  ___________________________________
-4-
Elements 1-20  _____________________________
Beyond element 20  ____________________________________________
Beyond element 83 (bismuth)  _____________________________________
Interesting to note:
 There are “magic numbers” of protons and neutrons (2, 8, 20, 28, 50, and 82) that are stable
 Atoms with even numbers of protons and neutrons tend to be stable
Types of Radioactive Decay
Symbol(s)
Explanation
Alpha
Emission of a helium nucleus (2 protons and 2
(α)
neutrons)
Penetrating ability
Low, stopped by paper
Beta
(β)
Emission of a high-speed electron; equivalent
to the conversion of a neutron to a proton
Medium, stopped by heavy
clothing
Gamma
(γ)
Emission of a high energy; similar to an x-ray
High, stopped by lead
Example of alpha emission
example of beta emission
example of gamma emission
*notice that the superscripts and subscripts to the left of the arrow equal the superscripts and subscripts to
the right!
Write the nuclear equation for the alpha decay of gold-185:
Decay rate
 Different radioactive elements decay at different rates
 Rate is described as a _________________, the amount of time (usually in years) for ½ of all the
radioactive atoms (parent isotope) in a sample to decay to the more stable isotope (daughter product).
Nuclear Fission & Fusion
When nuclei of certain isotopes are bombarded with neutrons, they undergo ________________, the splitting
of the nucleus into smaller fragments.
There are only two fissionable isotopes, draw the symbols for each below its name (both are atoms, not ions):
Uranium-235
Plutonium-239
The equation for the fission of Uranium-235 is:
Notice that there are the same number of protons on each side (how many? ______) and the same number of
neutrons on each side (how many? _________)
582731445
-5-
In a chain reaction ____________________________________________________________________
__________________________________________________________________________________
An atomic bomb is a device that starts an _______________________ nuclear chain reaction.
To be used for energy, _______________ must be controlled so that _____________ is released more
_______________.
Nuclear Fusion
Nuclear fusion can be thought of as the opposite of nuclear fission, that is ___________ nuclei ___________
to produce a nucleus with a greater mass.
This releases much ___________ energy than a fission reaction, but only happens above _____________ ˚C.
Write the formula for 4 hydrogen nuclei combining to create one helium nucleus:
MODELS OF THE ATOM (1900 to present)
Rutherford proposed an atom with a nucleus and negative particles moving around the nucleus.
The next major step forward on the model of the atom came after some important breakthroughs in our
understanding of



light behaves like a wave, but also like a stream of extremely tiny, fast-moving particles (wave-particle
duality)
light is a form of electromagnetic radiation
electromagnetic radiation can be described in terms of:
amplitude
wavelength ()
frequency ()
speed of light (c)
Notice: if  is large, then v must be small since c is constant; the reverse is also true
(therefore, long  light is low energy while short  light is high energy)


visible spectrum of light is just part of the electromagnetic spectrum – the part detectable by the human
eye

the energy absorbed by and given off by objects was thought to be continuous

Max Planck, in 1900, proposed that energy is only available in discrete packets, or
(singular is
); the size of the quantum is related to the frequency of the radiation by a
simple equation
582731445
-6-
Bohr Model of the Atom (1913)
 Bohr applied Planck’s idea of quantized energy to the model of the atom.
 Electrons move around the nucleus in specific orbits, or energy levels. These energy levels are said to be
quantized.
 An atom has several orbits, each representing a specific energy level. The energy levels are like steps of a
ladder. Only specific energy values (steps) can exist – there are none in between.

When an atom is not excited (
nucleus. The electrons are at the lowest energy level.

If an atom gains energy (
), an electron is displaced farther away
from the nucleus to one of the higher energy levels.
An atom emits energy when an electron falls from a higher energy level to a lower energy level in one sudden
drop, or transition. The energy released is electromagnetic energy.
The frequency () of the radiation emitted depends on the difference between the higher and lower energy
levels involved in the transition. Remember that a specific frequency correlates with a specific wavelength
(and color!) because c = .


), its electrons are in orbits close to the
Bohr’s model of the hydrogen atom
Bohr used this model to explain
the discontinuous line spectrum
of excited hydrogen atoms.




When an electron is close to the nucleus
of the atom 
Further away from the nucleus 
Absorption of energy increases the PE
of the electron as it moves further away
from the nucleus.
Thus, an electron of an excited atom
absorbs energy in the form of heat or
electricity when it moves to a higher
energy level. When the electron returns to
ground state, it emits the energy it absorbed.
The packet of energy emitted corresponds to a
specific color on the line spectrum.
Problems with Bohr’s Model:
1. it only explained the line spectrum for hydrogen
2. he couldn’t explain why the electrons, which he pictured as circling the nucleus like tiny planets, did not fall
into the nucleus every time they emitted light
Wave/Particle Paradox:




Light sometimes acts as a particle and sometimes acts as a wave.
Thomas Young noticed that when electrons are forced through a narrow slit, a pattern of wave interference
emerged. He was familiar with Rutherford’s experiment that showed electrons are particles and needed to
reconcile the two points of view.
This led to the development of the Quantum Mechanical Model of the atom.
The Quantum Mechanical Model is based on complex mathematical equations developed by Erwin Schrodinger
that describe the wave nature of the atom and the locations of the electrons.
582731445
-7-
Bohr Model vs. Quantum Theory
Bohr








Quantum Theory
QUANTUM THEORY
based partly on Heisenberg’s Uncertainty Principle
 the position and the momentum of a moving object cannot simultaneously be measured and known exactly
 there is an inherent limitation to knowing both where a particle is at a particular moment and how it is moving
in order to predict where it will be in the future
an electron is in an
– probability space where an electron can be found a
certain percentage of the time as defined by Schrodinger’s equations
An orbital can hold _____ electrons (this means the bigger the energy level, the more orbitals it has!).
Schrodinger’s equations (once all the mathematics has been done) describe where a particular electron is
likely to be found
We can know 4 pieces of information about the location of any electron within an atom. The following table
summarizes what we can know.
What we know about the location of an electron within an atom
Type of information
Label
What it represents
How many options?
Energy Level
n
n=1 through n=7
Sublevel
l
Distance from the
nucleus
Shape
Magnetic
lm
Orientation of Orbital
(which direction it is
pointing)
Spin
ls
How electron is
spinning
582731445
-8-
n = 1  1 sublevel (s)
n = 2  2 sublevels (s, p)
n = 3  3 sublevels (s, p, d)
n = 4  4 sublevels (s, p, d, f)
Depends on the number of
sublevels.
p-orbitals follow the x, y, and z
axes (one orbital per axis)
Up = + ½
Down = - ½
The spinning of electrons
generates an electric field.
How many e- can it
hold?
2n2
Level 1 = ______
Level 2 = ______
Level 3 = ______
Level 4 = ______
Level 5 = ______
(each orbital can hold
2 electrons)
For 2 electrons to
occupy the same
orbital they must
have opposite spin.
Each type of orbital has a distinct shape. The same type of orbital always has the same shape, but each orbital is
bigger in higher energy levels.
s sublevels look like:
p sublevels look like:
d sublevels look like:
f sublevels are weirder still… (check out http://winter.group.shef.ac.uk/orbitron/)
DESCRIBING THE LOCATION OF ELECTRONS IN AN ATOM:
To determine the electron configuration of an atom, 3 guidelines are followed:
Pauli exclusion principle:


Taken together, the four quantum numbers describe the state of a particular electron. For example, an
electron may have these quantum numbers: 1,0,0,-1/2 (corresponding to n,l,lm, and ls). Think of these as a
social security number for an individual, or a street address for a house. No two are identical.
Aufbau principle:

 1s is filled first, then 2s, 2p, 3s, 3p …
Hund’s rule:


equal energy orbitals are “degenerate orbitals”; for example, p x , py, and pz are degenerate orbitals
Orbital Diagrams
582731445
-9-
The locations of electrons in an atom are described using an orbital notation. In order to draw an orbital
notation diagram, you must understand the quantum numbers (energy levels, sublevels, orientation, and spin).
Summary of energy levels, sublevels, and orbitals
Principal energy Sublevel
Orbitals
level
n=1
1s
1s (one)
n=2
2s, 2p
2s (one) + 2p (three – x,y,z)
n=3
3s, 3p, 3d
3s (one) + 3p (three – x,y,z) + 3d (five – xy,xz, etc.)
n=4
4s, 4p, 4d, 4f 4s (one) + 4p (three – x,y,z) + 4d (five – xy,xz, etc.) + 4f (seven – xyz, etc.)
Number of electrons in each sublevel
Sublevel
Number of orbitals
Maximum # of electrons
(two per orbital)
s
p
d
f
This diagram helps show the differences in energy for each sublevel within an electron cloud. Working from the
bottom to the top, you can see how each atom’s electrons fill.
Use this to write in the electrons for a neutral atom of Phosphorous.
582731445
- 10 -
Assignment of electrons to an atom
Please summarize the three things you know about how we assign electrons to an atom.



Electrons fill orbitals in the following sequence:
Electron Configuration: The electron configuration is another way of describing the positions of electrons in an
atom using only letters and numbers. This is like a summary of the orbital notation. The number of electrons
in each orbital is written as a superscript above the orbital designation. The electron configuration for sulfur is
as follows:
Write the electron configurations for the following atoms:
Element
H
Electron configuration
Element
Al
He
O
Li
C
Cl
Mg
N
Ar
Electron configuration
At atomic number greater than 18, the sublevels begin to fill out of order. A good approximation of the order of
filling can be determined using either the diagonal rule or the location of the element on the periodic table.
1s
2s 2p
3s 3p 3d
4s 4p 4d
5s 5p 5d
6s 6p 6d
7s 7p 7d
Write the electron configurations for the following atoms.
4f
5f
6f
7f
Element
K
Electron configuration
V
Co
Zr
Cu
Looking for patterns:
582731445
- 11 -
Write the electron configurations for the following elements:
He
Li
Ne
Mg
Ar
Ca
Ti
Ni
Br
Describe the pattern(s) in the electron configurations above:
The Noble Gas Shortcut –
Write the electron configurations for the following atoms using the shortcut you just learned:
Be
Ne
Al
Ca
Sc
Ni
I
Ga
Valence Electrons
The electrons in the outermost energy level are active in bond formation. They are only in the outermost s and p
sublevels (between 1 and 8 electrons  see the periodic table). Since only s and p electrons are valence
electrons, the maximum number of valence electrons is eight. For the elements below, do the following:
1. Write the electron configuration
2. Determine the number of valence electrons
3. Write the number of valence electrons and circle it
Example:
Carbon
Electron configuration: 1s22s22p2
Number of valence electrons: 4
Example:
Oxygen
Electron configuration:
Number of valence electrons:
582731445
- 12 -
For the following atoms, write the electron configuration and circle the valence electrons.
1. fluorine
7. zinc
2. phosphorus
8. carbon
3. iron
9. oxygen
4. argon
10. aluminum
5. potassium
11. hydrogen
6. magnesium
12. helium
Electron Configurations for Ions
Follow all the same rules that you learned in writing electron configuration for atoms. Before you begin,
determine which noble gas is closest to the atom/ion in questions, then count how many electrons will be lost or
gained to make the ion. Subtract or add that number of electrons to the number of electrons in the original
atom. Write the electron configuration for the ion using the number of electrons in the ion. Note that it will
usually be the same as the electron configuration for the nearest noble gas.
For example: Magnesium atom  magnesium ion: Magnesium is closest to the noble gas
so a
magnesium atom will gain/lose
electrons when it becomes an ion. As an ion, magnesium has the same
electron configuration as
.
Write the electron configurations for the following elements and their most common ions:
Sodium
sodium ion
Oxygen
oxide (ion)
Phosphorus
phosphide (ion)
Beryllium
beryllium ion
Iodine
iodide (ion)
Lewis electron dot diagrams
Lewis diagrams or electron dot diagrams emphasize the valence electrons. Such diagrams show:
Kernel - the nucleus plus the inner electrons (all non-valence electrons)
- represented by the element’s symbol
Valence electrons – electrons in the outermost s and p orbitals
- represented by a series of dots around the kernel (element’s symbol)
- s orbital is 2 dots at 12 o’clock, px, py, and pz orbitals are 2 dots each at 3, 6, and 9 o’clock
Example: sulfur
Draw Lewis dot diagrams for the following atoms to show the valence electrons for each element:
1. calcium
582731445
5. bromine
- 13 -
9. phosphorus
2. potassium
6. carbon
10. hydrogen
3. argon
7. helium
11. vanadium
4. aluminum
8. oxygen
12. silver
The Periodic Table
Mendeleev (Russian) –
Moseley (English) –
Periodic Law –
Basic Organization
Groups / Families – vertical column –
Row / Period – horizontal row Periodic Characteristics:
1. atomic size
2. ionic size
3. metallic properties
luster – shiny
conductivity – able to transfer heat or electrons
malleability – can be rolled or hammered into sheets
ductility – can be drawn (pulled) into a wire
explained by….
4. ionization energy
5. electronegativity
582731445
- 14 -
nonmetallic properties
Explained by…
Groups: Elements in the same group have more similarities than elements in the same period because
.
Alkali Metals (Group 1)
Alkaline Earth Metals (Group 2)
Transition Metals
Metalloids – elements that have both metallic and nonmetallic properties
Diatomic elements – HONClBrIF elements – found combined with self – Br2, I2 N2, Cl2, etc.
Halogens (Group 17)
Noble / Inert Gases (Group 18)
582731445
- 15 -