Download I. History of the Atomic Theory

UNIT 3 and 4 – THE ATOM
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I. History of the Atomic Theory:
Atom - the smallest particle of matter which will exhibit the properties of that element. When broken down smaller than an
atom, the parts (protons, electrons, and neutrons) of different elements look exactly the same. You cannot tell a proton in a
gold atom from a proton in oxygen gas.
Atoms are very small—typically about 1 x 10-8 cm in diameter. 1.0 gram of lead contains 2.9 x 10 21 atoms of lead. By
comparison, the earth’s entire population is only 5 x 10 9 people.
Models – People use models to describe understanding of truth. Models evolve over time as technology improved to study
the atom. There were MANY scientists involved in this process. We will focus on only a few.
A. JOHN DALTON was an Englishman who was the first to develop and publish a theory about how atoms looked
and behaved. He conceived of the atom as a solid sphere, much like a billiard ball. The following are statements
of John Dalton’s ATOMIC THEORY with “*” beside the parts which are known NOT true today.
(1) Proposed: all elements are composed of very small particles called atoms which are indivisible. WRONG: Today we
know that atoms can be divided into protons, electrons, neutrons and almost 200 other subatomic particles.
(2) Proposed: All atoms of the same elements are identical. WRONG: We know that this is not true due to the presence of
isotopes. Isotopes are atoms with the same number of protons (same element) but different numbers of neutrons (different
masses). For example, “regular” carbon has 6 protons and 6 neutrons, but carbon-14 (radioactive carbon used in carbon
“dating”) has 6 protons and 8 neutrons.
(3) Proposed: Atoms of different elements are different TRUE
(4) Proposed: Atoms of different elements can combine with each other only in simple whole number ratios to form
compounds. TRUE
(5) Proposed: Chemical reactions occur when atoms are separated, joined or arranged. However, atoms of one element ARE
NOT changed into atoms of another element by a chemical reaction. (Only by nuclear reactions)
B. J. J. THOMPSON discovered both electrons and protons by means of his CATHODE RAY EXPERIMENT
(1) Proposed: Electrons are negatively charged particles, abbreviated e-; very lightweight –their mass is considered negligible
when describing the “mass of an atom” because; they actually weigh about 1/1839 the mass of a proton or neutron. TRUE
CATHODE RAY TUBE EXPERIMENT
+ Magnetic field
Note: Since the e- stream could be produced
from any metal, then, Thomson suggested
that all atoms had electrons and they were
negatively charged.
- Magnetic field
Electron Stream w/o applied electric field
Electron Stream after applying field, + on top, - on bottom
Know WHO, HOW and WHAT was discovered
 Electrodes are small pieces of metal which have electrical power running through them
 (-) negative electrode is called the cathode
 (+) positive electrode is called the anode
 Before their true nature was known or understood; electrons were called “cathode rays.” Now we know that
they are small negatively- charged particles which move at a very high speed. A stream of electrons is called
electricity.
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UNIT 3 and 4 – THE ATOM
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(2) Proposed: Thomson conceived of the atom as a plum pudding model—the protons and electrons were mixed together like
a fruit salad (or plum pudding). The larger pieces of the “plum” were the protons swimming in a pudding of electrons.
WRONG
Some models show the pudding model as a
mixture of positive and negative charges. –
that is incorrect
Correct model is with negative electrons in a
positive cloud
C. ERNEST RUTHERFORD worked for Thomson in his lab for a while. He performed the GOLD FOIL
EXPERIMENT
(1) Proposed: the atom as mostly space and that all of the positive charge was located in a very small central nucleus. TRUE
GOLD FOIL EXPERIMENT
Results: 1.
2.
3.
Conclusions:
Most particles went through
Some were slightly deflected
Some were deflected back and missed the fluorescent screen
Dense positive nucleus (i.e. plum pudding incorrect)
Negative electrons were moving around the nucleus
D. NIELS BOHR – Danish scientist who originally worked for
Thomson and Rutherford
(1) Proposed: Electrons were in energy levels. The further an electron was
from the nucleus, the higher its energy. TRUE
(2) Proposed: t the “planetary model” of the atom in which the electrons
orbit the nucleus much as the planets orbit the sun. WRONG
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E. ERWIN SCHRODINGER – an Austrian physicist, along with Werner Heisenberg and Louis de Broglie,
postulated the quantum (wave) mechanical model of the atom which we believe is true today – at least the math
works! In this concept the electrons do not actually “orbit” the nucleus but are found only in definite areas based
on the amount of energy they have and move according to wave functions. Key – we don’t really know how an
electron moves – just the probability of finding it in a particular region of space around the nucleus.
II. THE ATOM
F. PROTON- positively-charged particle which is found in the nucleus; it is considerable more massive than the
electron (1837 times heavier)
G. ELECTRON- very light negatively-charged particle which is found somewhere outside the nucleus. Its mass is
considered negligible when determining the mass of an atom. It weighs only 1/1837 that of a proton, but its
negative charge is as powerful as the positive charge of a proton.
H. NEUTRON- a particle found in the nucleus which is approximately the same mass as a proton but does not have
an electrical charge associated with it – it is neutral.
I. ATOMS ARE ALWAYS NEUTRAL PARTICLES—that is, they contain the same number of protons as
electrons (and it doesn’t matter how many neutrons they have)
J. IONS – particles which do not have the same number of protons and electrons so therefore they do have an
electrical charge associated with them. If they have more protons than electrons, they will have a positive charge.
If they have more electrons than protons, they will have a negative charge. # of Protons – # of e- = charge
K. ATOMIC NUMBER – the number of protons found in the nucleus of an atom. On the Periodic Chart it is the
large whole number in the upper corner. Notice that the only thing which makes on element different from another
one is the number of protons it contains. An atom with 6 protons is carbon; an atom with 5 protons is boron and an
atom with 7 protons is nitrogen.
NOTE THE DIFFERENCE BETWEEN ATOMIC MASS AND MASS NUMBER!
L. MASS NUMBER – the number of protons and neutrons found in the nucleus of an atom, this DOES NOT show
on the Periodic Chart. The atomic mass which does show on the chart is an average of all the different isotopes of
that element.
# of neutrons + # of protons = mass #
M. ATOMIC MASS – The atomic mass that shows on the Periodic Table for each element is really an average of the
masses of all the isotopes of that element, weighted by their percentage of abundance. Every element on the
Periodic Table has at least 3 isotopes; some of them have 20 or 30 or more isotopes.
III.
Determining Atomic Mass:
Avg. Atomic Mass = (fraction isotope #1) (mass of #1) + (fraction isotope #2) (mass of #2) + (fraction of isotope 3) (mass
of #3) +…….
The fraction abundance is simply the percent ÷ 100
Example 3-1. Naturally occurring chlorine is 75.53% chlorine-35 and 24.47% chlorine-37. What is the average atomic mass
which should be placed on the Periodic Table for the element?
Example 3-2. The element neon consists of three isotopes with masses of 19.99, 20.99 and 21.99 amus. These three isotopes
are present in nature to the extent of 90.92%, 0.25% and 8.83% respectively. Calculate the atomic weight of neon.
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Example 3-3. The element silver consists in nature of two isotopes, Ag-107 and Ag-109. The accepted atomic mass of silver
is 107.87 amu. What are the percentage of abundance of Ag-107 and Ag-109?
IV. CHEMICAL CONFIGURATION is a way of writing atoms or ions which gives
you lots of information.
Mass number
Ionic charge
16 2O
Atomic number 8 2 Number in a compound
We won’t be dealing with the number of the atoms in a compound until a later unit, so
leave the lower RHS blank for now.
Example 3-4. Write the chemical configuration for a particle containing 5 protons, 6 neutrons and 5 electrons.
Write the chemical configuration for a particle which contains 8 neutrons, has a mass number of 6 and contains 6 electrons.
Write the chemical configuration for a particle which contains 18 electrons, has a mass number of 32 and has a -2 charge.
Write the chemical configuration for a particle which has a charge of +3 and contains 10 electrons and 14 neutrons.
Write the chemical configuration for a particle which has a charge of +2, an atomic number of 24 and contains 28 neutrons.
Write the chemical configuration for a particle which has a mass number of 31, contains 18 electrons and has a -3 charge.
Sometimes you are not given enough information to write the entire chemical configuration for a particle but you can at
least identify as to what element it is and whether it is an atom or a positively-charged or negatively-charged ion.
Example 3-5. What particle has +2 charge and has 12 protons?
What particle has 50 protons and 48 electrons?
What particle has 44 protons and 44 electrons?
What particle has a -3 charge and contains 36 electrons?
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LIGHT AND ELECTRON CONFIGURATION
Much of what we know about the atom has been learned through experiments with light; thus, you need to know some
fundamental concepts of light in order to understand the structure of the atom, especially the placement of the
electrons.
I. Light
A. Characteristics of Light
1. Has "Dual" nature (or split personality)
2. There are times in which it behaves like a WAVE and other times when it behaves like a
PARTICLE
B. Light as a WAVE: You must be able to describe its wavelength, frequency, and velocity
1. Speed of light in a vacuum is the most accurately known constant, c, in the universe.
c = 3.0 x 108 meters/sec
OR 3.0 x 1017 nm/sec (assuming a vacuum)
2. Wavelength is represented by the Greek letter, lambda (  ): It is the length between
corresponding parts of adjacent crests and can be expressed in ANY units of length you
choose (feet, inches, meters, kilometers, etc)
3. Frequency is represented by the Greek letter, nu (  ): It is the number of wave crests which
pass a given point in 1 second. Its units are, most commonly, cycles / second , simply sec-1.
The unit, Hertz (Hz) is also used as a unit label.
II. Electromagnetic Spectrum:
A. Visible Light: considered to have a wavelength of between 700 nm and 400 nm. The longest
wavelength in visible is red (about 700 nm) and the shortest visible wavelength is violet (about 400 nm) The only
thing that makes one color of light different from another is its  and , the velocity is always the same.
B. Continuous Spectrum: Has all the energies of visible light. The colors of visible light in order from
longest wavelength to shortest spell out: ROY G BIV.
C. Bright-line Spectrum: Has distinct energies of light. Each element has its own distinct bright-line
spectrum. When an atom gains energy, the outer electrons become "excited" and JUMP to a higher
energy level. When the electron "falls back" to lower energy levels, the energy released is seen as
color. The amount of energy released corresponds to the color seen.
Example 4-3: Why is a red sweater red?
D. Entire Range of light energy: Includes ALL energies of light. Become very familiar with the
electromagnetic spectrum!!!
I better
memorize this
by the next
quiz… arghh!
nm
nm
III. Light as a WAVE:
c=

where c = speed of light (3.0 x 108 m/sec OR 3.0 x 1017 nm/sec)
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Note: BE CAREFUL OF THE UNITS YOU USE!!! If speed of light is in meters, then wavelength must be in
meters.
Example 4-1: Do you think a longer or a shorter wavelength would have the greater frequency?
Example 4-2: What mathematical relationship can you draw about wavelength and frequency?
Example 4-3: If the wavelength is known to be 550 nm, what is its frequency?
Example 4-5: If the frequency of light is known to be 9.45 x 10 14 s-1, is the light visible? If not, is it UV or IR?
Example 4-6: If the wavelength of light is 7.23 x 10-5 cm, is it visible? If so, what color? If not, is it UV or IR?
IV. Light as Particles:
When light behaves as a particle, it is called a PHOTON. We are interested in how much energy they
have. The unit of energy we will use is the Joule, J, and the equation which allows us to calculate the
energy of one photon is:
E1 photon = h  OR Eph = (hc)/ where h= Planck's Constant = 6.63 x 10-34 J . sec
Example 4-7: What is the energy of a photon of light whose frequency is 7.85 x 10 15 Hz? Is this visible?
Example 4-8: If a light has a wavelength of 550 nm, what is the energy of one photon of this light?
Example 4-9: If one photon of light is known to have energy of 3.33 x 10 -19 J, is it visible?
VI. Bohr Atom
A. electrons possess certain fixed amounts of energy
B. electron cloud made up of energy levels which are like the rungs of a ladder and are NOT
EQUALLY SPACED
C. amount of energy possessed by an electron determines its energy level
D. levels farther from nucleus represent higher energy
E. in order to move from one energy level to another, electrons must gain or lose an exact amount
of energy known as a Quantum.
n= ___
n= ___
n= ___
n= ___
n= ___
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e- ____ energy
Process___________
e- ____ EMR (energy)
Process___________
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V. Schrodinger Model: Quantum Mechanical Model (1926)
A. AGREES with Bohr about quantized energy levels
B. DIFFERS: does not define an EXACT path called an orbit. The math describes a region in space with a high
PROBABILITY of finding an electron. Unfortunately, we still call this an ORBITAL.
ELECTRON LOCATION: Electrons are distributed around the nucleus in specific ways and each electron has four
coordinates to describe its most probable location or distribution: 4 quantum numbers, n, l, m, and s.
A. Principle Quantum Number:
 ENERGY LEVEL
 Indicate distance from the nucleus


#e- = 2n2
designated by the letter n. ↑n  ↑E
The number of electrons in any energy level = 2n2; (n=energy level)
Example 4-10:. How many electrons could the 7th energy level hold? The 11th?
B. Second Quantum Number
 SUBLEVELS:

Indicate the shape of the region.

Designated by the letter ℓ.

The number of sublevels contained in an energy level is related to the principal quantum number of that
energy level. Sublevels are labeled by the letters s,p,d,f,g,h,I,j,k,l,m,m,etc.
The first energy level will have 1 sublevel and its label is "s"
The 2nd energy level will have 2 sublevels; "s", "p"
The 3rd energy level will have 3 sublevels; "s" , "p", and "d"
Example 4-11: What are the labels for the sublevels within the 4th energy level?
C. Magnetic Quantum Number
 ORBITALS

Indicate the orientation of the region in space

designated by the, mℓ.

The surface of the orbital is drawn to represent area where an electron with a particular energy can be found
90% of the time.


The number of orbitals is related to the quantum number, ℓ.
Each orbital holds only 2 electrons
s sublevel - 1 orbital; spherical shaped
p sublevel - 3 orbitals; dumb-bell shaped
d sublevel - 5 orbitals;
Energy Level
1
2
Example 4-12: How many orbitals would the "f" sublevel have?
3
Example 4-13: Complete the table:
4
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Sublevel
Orbitals
# e- 's
UNIT 3 and 4 – THE ATOM
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D. Spin Quantum Number:
 When two electrons are in a magnetic field, they will align opposite of one another. This is call the “spin”.
 The electrons within an orbital must have opposite spins to overcome their repulsion.


designated by, ms – can be either +1/2 or -1/2.
Use arrows to indicate the spin. +1/2: ↑, -1/2: ↓
VI. Rules for Arranging Electrons
A. Aufbau Principle
 electrons are placed in orbitals of _____________________ energy first
 s<p < d < f
B. Pauli Exclusion Principle: NO TWO ELECTRONS CAN HAVE THE SAME 4 QUANTUM NUMBERS.
 Remember: An orbital may hold only ___ electrons.

If they have the same energy (n), are in the same sublevel (ℓ), and are in the same orbital (mℓ), then they
must have opposite spins.
C. Hund’s Rule: when filling an orbital of equal energy, one electron enters each orbital until all contain ________,
then they __________.
VII.
Electron Configurations and Orbital Diagrams: Knowing the rules for electron arrangement allows us to
write what is called electron configuration which gives a description of where in that atom the electrons are actually located.
We can also draw an orbital diagram which gives the same information with more detail.
Orbital Diagrams
This is SO abstract.
If I don’t get this, I
better attend
tutorials!
Elements with unpaired electrons in their
outer energy level are called
_____________________
Paramagnetism is defined as being weakly attracted to a magnet due to unpaired electrons.
Example 4-14: Would zinc be attracted to a magnet?
Would chromium be attracted to a magnet?
Is palladium (#46) paramagnetic?
Example 4-15: Draw the orbital diagram for cobalt. Is cobalt paramagnetic? You may draw the diagram showing energy
increasing from left to right instead of up and down.
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Electron Configurations
Orbital Diagrams provide a lot of information, but they become cumbersome to draw. Chemists use electron
configurations to help model that structure of an atom or an ion. This in turn will help us with bonding in the next unit..
NOTE: After the 3p fills up you would think the 3d would fill, but it is not so. Electrons will fill the areas which require the
least amount of energy, and it takes less energy to fill the 4s than it does to fill the 3d, soooooo after
the 3p fills, the 4s will fill and then the 3d and then the 4p. THE ENERGY LEVELS THEMSELVES ARE
NOT OVERLAPPING--THERE IS SOME OVERLAPPING OF THE ENERGY OF EACH OF THE LEVELS. There are two
methods for predicting the filling order. The first uses the periodic table and the second a “diagonal tree” model.
PERIODIC TABLE and ELECTRON CONFIGURATION:
Most students do not realize that the Periodic Chart provides information on the various energy levels and
sublevels. Look at the Periodic Chart: On the LEFT side there are 2 columns which "stick up" above the
rest, on the RIGHT side there are 6 columns which "stick up" above the rest, in the MIDDLE there are 10
columns that seem to form a "short squatty" section all to itself, and down at the bottom there are 14
columns of elements which do not "appear" to "fit" into the chart at all (but they do)
DO THESE NUMBERS SOUND FAMILIAR TO YOU? 2, 6, 10, 14?
They should! You can actually read the electron configuration from the chart without using the diagonal
tree.
1.
2.
3.
4.
5.
Horizontal row (called a period) corresponds to an energy level
Read the Periodic Chart like a paragraph.
The "s" and "p" electrons will have an energy level number same as the period they are on
"d" electrons will have an energy level ONE LESS than the period they are on
"f" electrons will have an energy level two less than the period they are on
Label the Periodic Table with s, p, d, and f blocks
Example 4-16: Use the Periodic Chart to read the electron configuration for V (#23)
For In (#49)
For W (#74)
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Example 4-17: A certain element's electron configuration is known to be 1s 22s22p63s23p5. What element are we describing?
The outermost electrons are the electrons that are in the highest energy level, “n”.
The highest energy electrons are in the last sublevel that was filling.
Not always
the same!
Abbreviation: You can only use this if specified in the question! Start with the previous noble gas symbol in brackets and
write the configuration from there. For example: [Ar]4s23d5 is Mn.
Example 4-18: Using only the Periodic Chart, read the abbreviated electrons configuration for the following:
a. selenium What is the highest energy electron? How many? Outermost? How many?
b. Au (#79) What is the highest energy electron? How many? Outermost? How many?
DIAGONAL TREE:
Example 4-19: Write the electron configuration for iron based on the diagonal rule.
Draw the orbital diagram.
Write the electron configuration for tin.
Write the electron configuration Bi.
Example 4-20: Using either the periodic table or the Diagonal Rule, determine which is higher-energy electron:
3d or 4s
How do you know?
5d or 6p
How do you know?
6p or 7s
How do you know?
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UNIT 3 and 4 – THE ATOM
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Example 4-21: For the element chlorine:
 which electron is the highest-energy electron?
 Which electron is physically farthest from the nucleus?
 How many unpaired electrons are in chlorine? (it helps to do a quick orbital diagram of the highest sublevel to do
this!)
Example 4-22: For the element iron:
 In iron which electron is the highest energy electron?
 Which electron is physically farthest from the nucleus?
 How many pairs of electrons and unpaired electrons are in iron?
VIII. PERIODIC TABLE - Label hand-out with group and series names.
IX. IRREGULAR CONFIGURATION: Sometimes for a variety of reasons, electrons do not occupy the sublevel
they "should".
Note: Sublevels that are FULL are stable. The "d" and "f" sublevels are "fairly" stable when they are
exactly one-half full. (i.e. each orbital has one electron)
Example 4-23: Write the abbreviated electron configurations for Cr and Cu.
These are the actual configurations:
XIII. IONIC CONFIGURATIONS
Most metals will ____________ electrons to become ISOELECTRONIC with the nearest noble gas
Non-metals will ____________ electrons when they are with metals to become ISOELECTRONIC with the nearest noble
gas.
N3-
O2-
F-
Ne
Na+
Na+
Mg2+
Al3+
Example 4-24: Write the electron configurations for Ar, Ca, and S.
Example 4-25:
Predict the charge for Group 1 metals:
Predict the charge for Group 2 metals:
Predict the charge for Group 13 metals:
Predict the charge for Group 17 non-metals:
Predict the charge for Group 16 non-metals:
Predict the charge for Group 15 non-metals:
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“Iso” means “same” and
“electronic” refers to electrons. SO
isoelectronic must mean species
(ions or atoms) that have the same
number of electrons!
UNIT 3 and 4 – THE ATOM
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X. PERIODIC TRENDS
A. ATOMIC SIZE:
1. Decreases across a period from left to right:
a) Electrons are being added to the same energy level so they are the same distance from
the nucleus. BUT
b) There are more protons in the nucleus and they pull the electrons in closer.
2. Increases down a group
a) Electrons are being added to energy levels farther from the nucleus
b) Inner electrons shield the outer electrons from feeling the full effect of the nuclear
charge.
SHIELDING – a barrier made of inner-shell electrons which serves to decrease the pull of the nucleus
Shielding INCREASES as you go DOWN a column because there are more electron shells, and so more innershell electrons
ex: Lithium has 2 shells of electrons, so only ONE INNER SHELL of shielding, but cesium
has 6 shells of electrons, so FIVE INNER SHELLS of shielding
Shielding is considered to be CONSTANT as you move ACROSS a period because the number of inner shells is
staying the same
ex: Sodium and chlorine both have 3 shells of electrons, so both have TWO INNER
SHELLS of shielding
**The size of an atom (atomic radius) cannot be measured exactly because an atom does not have a sharplydefined boundary
Example 4.26: Which atom is larger, Zr (#40) or Sn (#50)? Why?
Which atom is larger, Li (#3) or Cs (#55)? Why?
B. IONIZATION ENERGY
Ionization Energy – the amount of energy required to remove the most loosely-held electron from gaseous atom or
an ion (usually the outermost electron)
1. Increases across a period
a) Electrons are closer to the nucleus
b) Electrons feel a stronger nuclear charge.
2. Decreases down a group.
a) Outer electrons are in higher energy levels
b) Outer electrons are further away from the nucleus
c) Inner electrons shield outer electrons from full nuclear charge.
Example 4.27: Which has higher ionization energy, Mg (#12) or Ba (#56)? Why?
Which has higher ionization energy, Al (#13) or Si (#14)? Why?
Which has higher ionization energy, Sb (#52) or Te (#)? Why?
Which has higher ionization energy, Mg (#12) or Cl (#17)? Why?
Multiple ionization energies: Removal of a second, third, fourth, etc, electron from an atom involves different
ionization energies. As more and more electrons are removed, it is noticed that there are places where there seems
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to be a big JUMP in the energy required to remove an electron. This indicates
(1) disruption of a stable ½ full condition
(2) a change in sublevels
(3) a change in orbits, or
(4) a change into the 1st orbit
C. ELECTRONEGATIVITY – SEE PERIODIC TABLE HAND-OUT
Electronegativity – the tendency of an atom to pull electrons WHICH ARE BONDED closer to itself and perhaps
add them to its outer orbit.
**Dr. Linus Pauling created an electronegativity scale from 0-4 to provide a way to compare electronegativities of
different elements with numbers. The most electronegative element on the table (Fluorine) has a value of 4. All the
other elements have lower numbers than this, with large metals having the lowest numbers of all.
1. INCREASES as you move ACROSS a period
a) Electrons feel a stronger nuclear attraction
2. DECREASES as you move DOWN a column
a) Electrons further from positive nucleus
b) More shielding of positive nucleus
Example 4-28: Which of the following elements is the most electronegative, Si(#14) or
Cl(#17)? Why?
Which would be the most electronegative, oxygen, with a Pauling
Electronegativity value of 3.44, or chlorine, with a EN value of 3.14?
D. REACTIVITY – the ability to gain, lose, or share electrons.
1. Reactivity in metals INCREASES as you move DOWN a column because the ionization
energy in decreasing – the atoms are getting larger and there is more shielding so it is easier to
remove and electron.
2. Reactivity in metals DECREASES as you move ACROSS a period from left to right because
it becomes harder and harder to lose an electron.
3. Reactivity in non-metals DECREASES as you move DOWN a column because the atoms are
larger with more shielding, so electron affinity decreases and it gets harder for atoms to gain
electrons.
4. Reactivity in non-metals INCREASES as you move ACROSS a period from left to right
because it gets easier and easier to attract electrons to the atom.
Example 4-29: Which is more active, Ca (#29) or Ba (#56)? Why?
Which is more active, Sr (#38) or Sn (#50)? Why?
Which is more active, P (#15) or Cl (#17)? Why?
Which is more active, Cl (#17) or Br (#35)? Why?
E. ION SIZE
1. The METALLIC ION is SMALLER than its corresponding atom because all metal ions
LOSE the electrons in the outer orbit, so they are one orbit smaller. Also, the proton/electron ration
is LARGER in positive ions, and the higher the proton/electron ration, the SMALLER the
particle.
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2. The NON-METALLIC ION is LARGER than its corresponding atom because all non-metal
ions GAIN electrons in the outer orbit and these extra electrons create repulsion and expansion in
the outer orbit. Also, the proton/electron ratio is SMALLER in negative ions, and the lower the
proton/electron ratio, the LARGER the particle.
3. For BOTH METAL AND NON-METAL IONS, the size INCREASES as you move DOWN
a column due to increased shielding and additional orbits.
4. For BOTH METAL AND NON-METAL IONS, the size DECREASES as you move
ACROSS a period from left to right due to the increasing proton/electron ratios.
Example 4-33: Which would be larger, and Al atom or and Al ion? Why?
Which would be larger, a F atom or a F ion? Why?
Which would be larger, oxide ion or sulfide ion? Why?
Arrange the following in order of increasing size: Cl ion, Al ion, P ion, Mg ion, Na ion
REMEMBER THAT THE LARGER THE PROTON/ELECTRON RATIO, THE SMALLER THE
PARTICLE AND VICE VERSA
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