Download UNIT 1 Atomic Structure

Chemistry
Objectives
Goals : To gain an understanding of :
1. Atoms and their structure.
2. The development of the atomic theory.
3. The quantum mechanical model of the atom.
4. The electron configuration of the first twenty elements.
5. Periodicity.
Objectives : The student will be able to :
1. Define "atom" and state John Dalton's atomic theory.
2. Name and describe the particles which make up an atom and their arrangement in an
atom.
3. State who discovered the electron and how.
4. State who discovered the nucleus and how.
5. State who discovered the neutron.
6. State the relative size and mass of the nucleus as compared to the entire atom.
7. Define and be able to determine the following given the data from the periodic table :
atomic number, mass number, and atomic mass.
8. Define "isotope" and "atomic mass unit."
9. State which discovery shattered Dalton's theory of atoms as indivisible particles.
10. Contrast and evaluate Thomson's, Rutherford's, Bohr's and the quantum mechanical
models of the structure of an atom.
11. State which of the above models is accepted today.
12. Define "quanta" and "energy level."
13. Describe the quantum mechanical model of the atom and use it to explain how the
electrons of an atom are arranged into energy levels and orbitals (sublevels).
14. Describe the s,p,d, and f orbitals.
15. Determine the electron configuration for the any element in its ground state.
16. Name and describe the three rules which govern the filling of atomic orbitals by
electrons within their given energy levels.
17. Explain how and why flame tests can be used to identify certain elements.
18. Describe and explain the science of spectroscopy, stating how qualitative and
quantitative data can be obtained.
19. Distinguish between absorption spectra and emission spectra.
20. Describe the contributions of Dmitri Mendeleev and Henry Moseley to the
development of the modern periodic table.
21. Describe the arrangement of elements on the periodic table.
22. Define "group" and "period" and describe how the properties of elements change in
groups and periods.
23. Locate the metals, metalloids and nonmetals on the periodic table.
24. Describe the metals, metalloids and nonmetals.
24. State the periodic law.
25. State which subatomic particle plays the greatest role in determining the chemical and
physical properties of the different elements.
26. Describe the noble gases, representative elements, transition metals and inner
transition metals in terms of their electron configurations.
27. Identify the s, p, d, and f blocks on the periodic table.
28. Describe and explain the groups and periodic trends in atomic size, ionization energy,
electron affinity, ionic size and electronegativity.
29. Define "ionization energy," " electron affinity," and "electronegativity."
30. Name and describe three forces affecting the periodicity of elements.
31. Describe the following groups and state several uses of each : noble gases, alkali
metals, alkaline earth metals, aluminum group, carbon group, nitrogen group, oxygen
group and the halogens.
32. Name and explain the accepted theory about the origin of the elements which make
up the universe.
Notes
1. An atom is the smallest part of an element that retains the properties of that element.
The concept of an atom goes a long way back. It was first suggested by an ancient Greek
named Democritus (the Greek word "atomos" means indivisible). Democritus theorized
that if you took an object and cut it in half again and again you would eventually end up
with some particle which could not be further divided.
In the early 1800's an English scientist by the name of John Dalton started relating
what chemists could see to the concept of the atom. He came up with an atomic theory
which could be stated as follows :
o
o
o
o
All elements are composed of tiny indivisible particles called atoms.
Atoms of the same element are identical and differ from atoms of other
elements.
Atoms of different elements can combine together in simple whole
number ratios to form compounds.
Chemical reactions are the rearranging of the combinations of atoms of
elements in compounds. The atoms themselves remain unchanged.
2. Atoms are composed of three primary particles, protons, neutrons and electrons.
Table1-1 Subatomic Particles and Their Properties
Relative electrical Approximate relative Actual mass in
charge
mass in amu
grams
Particle
Symbol
Location
Electron
e-
orbits the
nucleus
1-
1/1840
9.11 x 10-28
Proton
p+
nucleus
1+
1
1.66 x 10-24
Neutron
nº
nucleus
0
1
1.66 x 10-24
amu (atomic mass unit = 1.66 x 10-24 grams)
3. The electron was discovered by the English physicist Sir Joseph J. Thomson around
1897 with the use of a cathode ray tube. A cathode ray tube is similar to your TV. It has
an anode (negative electrode) and a cathode (positive electrode). These are enclosed in an
evacuated (air removed) glass container and when a charge is applied the electrons flow
from anode to cathode through the open space of the glass container. Thomson observed
these particles and determined that the particles :




move at a very high speed
have a negative charge
have a mass of about 1/2000 of a hydrogen atom (smallest atom)
were the same regardless of which gas was used in the container or the metal
used as the electrode
The particles were eventually named "electrons." This discovery shattered
Dalton's notion of an atom. To Dalton atoms were tiny, solid particles, not
containing smaller particles.
4. The nucleus was discovered by Ernest Rutherford in 1911. Rutherford set up an apparatus in
which he aimed alpha particles (type of radiation made up of helium nuclei) at a very thin sheet
of gold foil. If the atoms were made up of evenly dispersed protons and electrons, as believed at
the time, the alpha particles should go straight through unhindered. What happened was that
some of the alpha particles went through, some were deflected as they passed through and some
bounced back. Rutherford concluded that the positive mass of the atom (protons) must be
concentrated in a very small area and most of the rest of the atom must be empty space. The area
where the protons were located was called the nucleus.
5. Sir James Chadwick discovered the neutron in 1932.
6. The nucleus of an atom is extremely small as compared to the entire size of the atom. It can be
compared to the size of a marble in a football stadium where the marble represents the nucleus
and the stadium the entire atom. Most of the space of the atom is occupied by the orbiting
electrons.
7. Atomic number - indicates the number of protons and defines the element (atomic
number 6 is always carbon, atomic number 7 is always nitrogen etc.).
Mass number - equals the total number of neutrons and protons in the nucleus of an
atom for the most common isotope this equals the atomic mass rounded off to the nearest
whole number
Atomic mass - the average mass of an atom of an element (in amu)
Calculation of the number of particles in an atom of an element :
- number of protons equals the atomic number
- number of neutrons equals the mass number minus the atomic number
(remember virtually all the mass is from the neutrons and protons -each with an
amu of 1)
- number of electrons equals the number of protons in a neutral atom (same
number of positive charges as negative charges)
Consider the following periodic table information for carbon, nitrogen and
sodium :
Carbon's atomic number is 6, has an average mass of 12.011 amu and carbon's most common
isotope has a mass number of 12 amu. Therefore, the most common type of carbon atom has 6
protons, 6 neutrons and 6 electrons.
Nitrogen's atomic number is 7, has an average mass of 14.007 amu and nitrogen's most common
isotope has a mass of 14 amu. Therefore the most common type of nitrogen atom has 7 protons,
7 neutrons and 7 electrons.
Sodium's atomic number is 11, has an average mass of 22.990 amu and nitrogen's most common
isotope has a mass of 23 amu. Therefore the most common type of sodium atom has 11 protons,
12 neutrons and 11 electrons.
http://science.widener.edu/svb/tutorial/protons.html - for practice calculating the above.
8. An isotope is a particular form of an atom of an element. Isotopes have a different number of
neutrons and therefore differ in mass (that is why there is a non-whole number atomic mass
which is an average of the various isotopes). The number of protons is always the same in
isotopes since they are different forms of the same element (must be same atomic number). For
example, the following chart shows three isotopes of hydrogen :
Isotope
Atomic
Number
Number of
protons
Number of
Neutrons
Number of
electrons
mass
(amu)
Hydrogen-1
1
1
0
1
1
Hydrogen-2
(deuterium)
1
1
1
1
2
Hydrogen-3
(tritium)
1
1
2
1
3
Note hydrogen has three isotopes, each with a whole number mass, yet hydrogen as an element
has a characteristic average mass called the atomic mass (1.0079 amu).
An atomic mass unit (amu) is 1/12 the mass of a carbon-12 atom, or 1.66 x 10-24 grams
9. Thomson's discovery of the electron shattered Dalton's theory of the atom. Dalton thought that
atoms were solid, indivisible particles. The discovery of smaller particles within the atom
changed this view of the atom.
10. The different historical models are described as follows :





Dalton's model of the atom - solid, tiny, indivisible particles.
Thomson's model - often describe as the "plum pudding" model - electrons
are scattered throughout the atom.
Rutherford's model - includes the solid nucleus in the center of the atom.
Niels Bohr's model - electrons are in fixed orbitals, or energy levels, at
certain distances from the nucleus much the way the planets orbit the sun.
Quantum mechanical model - the electrons have certain probabilities of
being located at certain places and their positions are described according
to four quantum numbers. The circle in the diagram below shows where
the electrons have a 90% probability of being within the electron cloud.
11. The quantum mechanical model is the most accepted model today. The diagrams you are
probably most familiar with are that of the Bohr model, which is useful in teaching basic electron
configuration (distribution in energy levels).
12. Quanta (s. quantum) are the discrete amounts of energy that an electron absorbs as it moves
up an energy level or releases when it moves down an energy level. Energy levels are certain
regions surrounding the nucleus of an atom where electrons are likely to be found.
13. The quantum mechanical model of the atom is a mathematical model which predicts the
probability of electron locations and paths in electron clouds. The positions and orbits of
electrons are referred to as energy states and are described by four quantum numbers :




the principle quantum number (n) - indicates the energy level the electrons are in
(there are seven energy levels, therefore n may equal 1,2,3,4,5,6 or 7)
the orbital quantum number - indicates the shape of the atomic orbital (region in
space where there is a high probability of finding an electron). the first four
orbital numbers are indicated by the letters s (has a spherical shape), p (has a
dumbbell shape), d, and f (d and f have complex shapes).
magnetic quantum number - indicates the position of the orbital about the x, y,
and z axes
spin quantum number - indicates the direction of spin of an electron - clockwise
or counterclockwise
These quantum numbers can be summarized as follows :
Principle
energy level
Number
Total number
Type of sublevel (number of
Total number of
of
of sublevels (=
orbitals per sublevel)
electrons(= 2n2)
sublevels
n2)
n=1
1
1s (one orbital only)
1
2
n=2
2
2s (1 orbital), 2p (3 orbitals)
4
8
n=3
3
3s (1 orbital), 3p (3 orbitals),
3d (5 orbitals)
9
18
n=4
4
4s (1 orbital), 4p (3 orbitals),
4d (5 orbitals), 4f (7 orbitals)
16
32
Note : Each quantum position as described by the first three quantum numbers can be occupied
by only two electrons with opposite spins. For example, if the position of a pair of electrons are
described as 4px, this means they are in the fourth energy level (n=4), and have a dumbbell
shaped orbital (orbital quantum number is p) around the x axis (third quantum number). These
two electrons in this orbital, as stated above, would have opposite spins.
Electrons must be placed in the lowest possible energy levels first. this is referred to as the
ground state of an atom - the state with the lowest possible energy level.
14. Above.
15 and 16. There are three rules which govern electron configuration (the arrangement of the
electrons of an atom into its orbitals).

Aufbau (German for "building up') Principle - electrons enter orbitals of the
lowest energy levels first. The following is the aufbau help diagram. The arrows
indicate the lowest energy levels which need to be filled first. The 1s is always
filled first then the 2s, 2p, 3s etc.. Note the 4s is filled before the 3d.

Pauli Exclusion Principle - an atomic orbital may describe at most two electrons ;
electrons in the same orbital must have opposite spins.
Hund's rule - when electrons occupy orbitals of the same energy level (e.g. the
three 3p orbitals or the five 4d orbitals), one electron will enter each orbital with a
parallel spin until all orbitals have one electron. Then and only then will a second
electron be added to those orbitals.

The following diagram shows four elements and their electron configurations. They are
explained below.
Hydrogen has only one electron so it is in the lowest orbital, the 1s orbital. The upwards arrow
indicates a clockwise spin. Hydrogen's electron configuration can then be described as 1s1. This
means there is one electron in the first energy level in the s orbital (spherical shape).
Helium has two electrons. The first orbital will hold two electrons so its summary electron
configuration is 1s2. Note the arrows indicating opposite spins (Pauli exclusion principle).
Lithium has three electrons so one has to move into the next lowest orbital, the 2s. Its summary
electron configuration is 1s22s1.
Nitrogen has seven electrons. The 1s and 2s fill up first (Aufbau principle). Now one electron
will go into each of the 2p orbitals with the same spin according to the Hund's rule. Nitrogen's
summary electron configuration is 1s22s22p3.
Oxygen's electron configuration is similar to nitrogen, only now there are four electrons to put
into the three 2p orbitals. Is summary electron configuration is 1s22s23p4.
17. Flame tests are tests used to identify metal or metallic ions. The substance to be identified is
usually put on a wire loop and then placed into a flame. The flame will then show a color which
can be used to help identify the metal. The colors in fireworks are produced by metals or metallic
ions added to the fireworks. The color comes from energy being released by the metallic atoms
or ions. When energy is added to the atoms or ions electrons jump up energy levels as they
absorb the energy. When they fall back down energy levels they release energy in the form of
light of a characteristic wavelength or color for that element. For example barium gives off a
pale green color and strontium gives off a brilliant red color.
18 and 19. Spectroscopy is the study of spectra - the types of light (wavelengths) that are
absorbed or emitted by elements. Spectra allow scientists to, in a sense, see into the atomic
structure of matter. The colors emitted or absorbed indicate changes in the positions of electrons
in their orbitals. Emission spectra are the spectra released by substances that have been excited
by the addition of energy. Absorption spectra are spectra that show the wavelengths (colors) of
light that are absorbed by substances.
Spectroscopy is a very important tool in the hands of an analytical chemist. Qualitative data
(what is present) can be determined by the wavelengths of light that are produced - much like the
flame tests. Quantitative data (how much is present) can be determined by the brightness or
thickness of the spectral bands that are produced
Goals : To gain an understanding of :
1. Valence electron and electron dot notation.
2. Stable electron configurations.
Objectives : The student will be able to :
1. Define "valence electron" and state the number of valence electrons in the groups of
representative elements.
2. Use electron dot formulas to represent the valence electron of given elements.
3. Describe a stable electron configuration and use this information to explain how and why
some atoms are inert and others react readily to form compounds.
Notes
1. Valence electrons are the electrons in the highest energy level of an atom. For example, in the
calcium atom (electron configuration 1s22s22p63s23p64s2) the 4s2 electrons are the valence
electrons. In the titanium atom (electron configuration 1s22s22p63s23p64s23d2) The 4s2 electrons
are still the valence electrons -they are in the highest energy level. In the phosphorus atom
(electron configuration 1s22s22p63s23p3) the 3s23p3 are the valence electrons.
The valence electron number of the representative elements are :








Group 1A - 1 valence electron
Group 2A - 2 valence electrons
Group 3A - 3 valence electrons
Group 4A - 4 valence electrons
Group 5A - 5 valence electrons
Group 6A - 6 valence electrons
Group 7A - 7 valence electrons
Group 0 - 8 valence electrons
3. Dot formulas use dots surrounding the symbol of the element to represent the valence
electrons. The dot formulas for period 2 and 3 would appear as follows.
Note electrons are usually shown as far apart as possible -they have the same charge and
therefore repel each other.
3. The Noble gases (Group 0) have a stable electron configuration (s2p6) with 8 electrons filling the
outer s and p orbitals. This stability comes from the low energy state of this configuration and
also accounts for the low reactivity of these elements (most elements react with other elements to
get to a lower, more stable energy state). For example the halogens (Group 7A) have 7 valence
electrons (s2p5) and want to gain one electron to get the low energy, stable electron configuration
of the noble gases. The elements in group 6 (s2p4) want to gain 2 electrons to get the low energy,
stable electron configuration of the noble gases. The Group 1A elements (s1) want to lose their
outer electron to empty their outer shell and get a stable electron configuration. For example if
sodium (1s22s22p63s1) loses its 3s1 electron it will have filled s and p orbitals in its outer energy
level.
4. Gilbert Lewis, in 1916, proposed the octet rule : Atoms react by changing their number of
electrons so as to acquire the stable electron configuration of a noble gas (s2p6).
5. An exception to the octet rule is the electron configuration of helium. Helium(1s2) is a noble gas,
only it has only one orbital, the s orbital. It is filled and therefore stable and elements close to it
(lithium, beryllium and sometimes hydrogen) try to acquire its electron configuration by losing
or gaining electrons).
6. The pseudo noble-gas electron configuration has the outer three orbitals filled, the s, p and ds2p6d10 (18 electrons total) and so is fairly stable. Elements that attain this electron configuration
are at the right side of the transition metals (d-block).
7. An ion is a charged atom that is formed by the gaining or losing of electrons. When an atom
gains or loses electrons it is no longer neutral because the number of electrons (negative charges)
and protons (positive charges) are not equal. For example when a sodium atom loses an electron
to get the noble gas electron configuration of neon it gets a charge of +1and becomes the sodium
ion, because it now has 11 protons (+) and only 10 electrons(-). When a chlorine atom gains one
electron to get the noble gas electron configuration of argon it gets a charge of -1 and becomes
the chloride ion, because it has 17 protons (+) and 18 electrons (-).
Cations are positive ions and are formed by the losing of electrons. Anions are negative ions and
are formed by the gaining of electrons.
8. The charge on an ion depends on the number of electrons gained or lost. An element that loses
two electrons becomes an ion with a +2 charge. An element which gains three electrons becomes
an ion with a -3 charge.