Chapter 3 - Atoms and Elements

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					           Chapter 3 – Atoms and Elements
Study Goals

 Classification of Matter

 Elements and Symbols

 The Periodic Table

 The Atom

 Atomic Number and Mass Number

 Isotopes and Atomic Mass

 Electron Energy Levels

 Periodic Trends
All things around us that we can see or touch or feel but don't see
(e.g. air current) are made up of what chemists call matter − the
material that makes up a substance. Before we look into the stuffs that make up matter at the
atomic level, let us define certain terms that chemists use to describe matter macroscopically.

A. Matter Can Be Of A Pure Form.
 i. Pure matters can be made of a single type of atoms (e.g. copper metal, made of 100% copper),
    in this case they are called pure elements;
ii. or they can be made of 2 or more different types of atoms that are chemically combined in a
    definite proportion (e.g. water consists of 2 H atoms and 1 O atom), in this case they are called
    pure compounds.
B. Matter Can Also Be Of A Mixture
A mixture consists of 2 or more substances that are physically mixed but not chemically

                                      Air is a mixture of several gases:
                                      N2, O2, CO2, and traces of other gases.

              Cup made of brass. Brass is an alloy metal of copper and
              zinc (an alloy metal is one that is made of 2 or more
              different metals that are homogeneously mixed together).

                                      Tea is a mixture of water and substances
                                      (caffeine, minerals, and others) that
                                      dissolve from the tea powder.

                           Ocean water (a mixture of water, salt
                           (NaCl) and other dissolved minerals).
 In a mixture, the different types of substances that are found in it are not chemically combined;
as a result they can be separated by physical processes.
Ex: Place ocean water in a cooking pan and heat it until all the water molecules
    have evaporated; left behind is a residue of dry, white powder consisting of
    salt (sodium chloride) and other minerals.

 Piles of salt surround an evaporation pond in the
 valley near Glendale, Arizona.

            Gold panning separates the heavier
            gold nuggets from sand and dirt.
            Homogeneous vs Heterogeneous Mixtures
A mixture can be classified as homogeneous or heterogeneous.
A) In a homogeneous mixture (also called a solution), two or more
substances are mixed together in the same phase. They
are uniformly dispersed throughout each other even down
to the molecular level.

         IV saline water is a homogeneous mixture of water                   A bronze cauldron.
         and dissolved electrolytes such as sodium chloride.                 Bronze is a homogeneous
                                                                             mixture of copper and tin.

B) By contrast to a homogeneous mixture, a heterogeneous mixture does not have a uniform
composition throughout the sample. The uneven texture of the material can often be detected.
Heterogeneous mixtures may appear uniform macroscopically but on closer examination (i.e.
under the microscope), their composition is not uniform.

                                                Blood may appear homogeneous until
                                                examined under the microscope, which
                                                reveals red and white blood cells.

                                 Granite is a heterogeneous mixture that contains discrete
                                 regions of different minerals (feldspar, mica, quartz). These
                                 minerals do not mix together at the molecular level.
         Pure Element, Compound, or Mixture ?
1) Carbon dioxide (CO2) gas we exhale

2) A salad dressing of oil and vinegar

3) Herbal tea

4) Water and sand

5) A chocolate chip cookie

6) Vinegar (consists of acetic acid and water)

7) Ice

8) Oxygen (O2)

9) Mecury (Hg) in a thermometer

10) Air in a scuba tank
                     Elements and Their Symbols
Elements are primary substances from which all things are built and can not be broken down
into simpler substances that still retain their identity.
- currently 116 elements are known. Of these, only about 90 are found in nature, the rest have
been created in the laboratory.

           Hg                 S                  Cu                  Fe             Al
                          Naming of Elements
Many elements are named after planets, mythological figures, minerals, colors,
famous people, etc.

Elements                     Source of Name

Uranium                     The planet Uranus
Titanium                    Titans (mythology)
Chlorine                     Chloros (Gr. for greenish-yellow)               At room temperature,
                                                                             chlorine is a pale yellow-
Iodine                       Ioeides (Gr. for violet)                        green gas. It is denser
                                                                             than air and will settle to
Californium                  California                                      the bottom of a flask.
Curium                       Marie Curie

                                       Iodine is a dark-gray/purple-black solid
                                       that sublimes at room temperature into a
                                       purple-pink gas that has an irritating odor.

              Marie Curie was a Polish-born French
              chemist and pioneered in the early field
              of radiology and a two-time Nobel laureate.
                       Chemical Symbols
Chemical symbols are either one-letter or two-letter abbreviations:
One-Letter Symbols                      Two-Letter Symbols
C carbon                                   Co cobalt
S sulfur                                   Si    silicon
N nitrogen                                 Ne neon
I   iodine                                 Ni    nickel                             Nickel

Note that in a two-letter symbol, only the 1st letter is capitalized, and the
2nd letter is in lower case.

● Many names and symbols are derived from their ancient Greek or Latin names:
Name                            Symbol                     Silver necklace
copper (cuprum)                    Cu
gold (aurum)                       Au                                                 Gold bar
iron (ferrum)                      Fe                                At RT, sodium is soft and can be cut
                                                                     with a knife. Exposed to humid air,
lead (plumbum)                     Pb                                the silver surface is quickly
                                                                                  oxidized to a chalky
potassium (kalium)                 K                                              white sodium oxide
silver (argentum)                  Ag                                             powder. To prevent
                                                                                  exposure to moisture,
sodium (natrium)                   Na                                             sodium is stored often
                                                                                  under oil.
                Physical Properties of Elements
Every element has physical properties that include shape, color, odor, taste, density,
hardness, melting point, and boiling point.

                                    COPPER                     Physical properties

                                    Color                         reddish-orange
                                    Odor                          odorless
                                    Melting point                 1083C
                                    Boiling point                 2567C
                                    State at 25C                 solid
         copper kettle
                                    Luster                        very shiny
                                    Conduction of electricity     excellent
                                    Conduction of heat            excellent

                         copper tubes
                              (Main Group Elements)

                Periodic Table of Elements Elements
                           Periodic Table of

                                                                                 Dmitri Mendeleev
                                                                                 (1834 -1907)

By the late 1800s, chemists began to recognize that the elements can be arranged into groups or
families based on similar chemical or physical properties. Dmitri Mendeleev, a Russian scientist,
was credited for arranging the 60 elements known at that time into what is known today as the
Periodic Table. Currently, about 116 elements are known in the Periodic Table.
                                The Alkali Metals
Several groups are known by their special names as a result of their characteristic chemical
The group 1A elements are also known as the ALKALI METALS. The word "alkali" is derived
from an Arabic word meaning "ashes". Many sodium and postassium compounds were isolated
from wood ashes (Na2CO3 and K2CO3 are still occasionally referred to as "soda ash" and
"potash“, respectively).
The alkali metals are: - soft, shiny metals
                       - good conductors of heat and electricity
                       - relatively low melting points
                       - react vigorously with water
                       - form white products when combined with Oxygen
(Exception: Hydrogen, although listed as a group 1A element, is not a metal.)
                           Examples of Alkali Metals

                                          4 Li(s) + O2 (g)  2 Li2O(s) (lithium oxide, left)
                                          2 Na(s) + O2 (g)  Na2O2(s) (sodium peroxide, center)
                                          K(s) + O2 (g)  KO2(s) (potassium superoxide, right)

a) Lithium reacts with water to give off hydrogen gas (H2): Li(s) + 2 H2O(l)  2 Li(OH)(aq) + H2(g)
b) Sodium reacts with water to give off hydrogen gas which can ignite spontaneously.
c) Potassium reacts with water and ignites the evolved hydrogen gas very readily.
The Group 2A elements (Be, Mg, Ca, Sr, Ba, Ra) are called
the ALKALINE EARTH METALS. They are silvery-colored,
                                                                            The Alkaline
soft and shiny metals like group 1A metals but not as reactive              Earth Metals
as the alkali metals. The reactivity of these elements
increases going down the group.

- they are found in compounds which form important rocks in
the Earth's crust. This gives rise to the group's alternative
name, the "alkaline earth" metals. The emerald
(Be3(Al,Cr)2Si6O18) is the green variety of beryl, the mineral
that is the principal source of beryllium.
                                                         Radium is rare, white, luminescent, highly radio-
                                                         active metallic element found in very small
                    Emerald                              amounts in uranium ores. It is used in cancer
                                                         radiotherapy, as a neutron source for some
                                                         research purposes, and as a constituent of
                                                         luminescent paints.
Beryl                                                    A radium watch dial with radio-
                                                         luminescent paints. Left image shows
                                                         the areas of the watch that are the
                                                         most radioactive (dark shadings).

                   Solid Lumps of Beryllium metal. Beryllium has low
                   aqueous solubility, which means it is rarely available
                   to biological systems  it has no known role in living
                   organisms, and when encountered by them, is
                   generally highly toxic.
The Group 7A elements are called the HALOGENS, which include
F, Cl, Br, I, and At. The halogens are strongly reactive, especially    The Halogens
fluorine and chlorine, and they form compounds with many other
elements. In particular, the halogens form salt compounds with
metals (e.g. NaCl, MgBr2, etc).
- in their elemental forms, the halogens exist as diatomic molecules
(e.g. F2, Cl2) but these only have a fleeting existence in nature and
are much more common in the laboratory and in industry. At room
temperature and pressure, fluorine and chlorine are gases, bromine
is a liquid and iodine and astatine are solids; Group 7A is therefore
the only periodic table group exhibiting all three states of matter.        Bromine
Halogens are highly reactive, and as such can be
harmful or lethal to biological organisms in
sufficient quantities. Both chlorine and bromine
are used as disinfectants for drinking water and
swimming pools.                                         Chlorine
           Chlorine is the active ingredient of most     (Cl2)
           fabric bleaches and is used in the produc-
           tion of most paper products.

                                                                        Iodine (I2)
                          The Noble Gases
One of the eight main groups of elements, Group 8A is known as the Noble Gas
group. Elements in this family have such a low reactivity that they were formerly
known as the inert gases. Thus they are chemically very stable due to having the
maximum number of valence electrons their outer shell can hold. The Noble Gases
include helium, neon, argon, krypton, xenon, and radon.
- under normal conditions, the noble gases occur as odorless,
colorless, monatomic gases. Each of them has its melting and
boiling point close together, so that only a small temperature range
exists for each noble gas in which it is a liquid (e.g. Argon has
mp: –189.6°C, bp: –185.8°C). The Noble gases have numerous
important applications in lighting, welding and space technology.
Argon is often used as a safe and inert
                                                                            Helium blimp
atmosphere for the inside of filament light bulbs.
Some of the noble gases glow distinctive colors           Krypton laser
                  when used inside lighting tubes
                  (neon lights).

                                                          Helium, due to its nonreactivity (compared
                                                          with flammable hydrogen) and lightness, is
                                                          often used in blimps and balloons. Krypton
                                                          is also used in lasers, which are used by
                                                          doctor for eye surgery.
                           The Transition Metals
The transition metals are the subgroups of elements intervening between groups IIA and IIIA
in the periodic table. They are classified separately because of the filling of their d subshell
orbitals.                                All the transition elements are metallic, but unlike the
                                        representative metals, they are likely to be hard, brittle,
                                        and have high melting points. Mercury (Hg) is an exception
                                        because it is a liquid at room temperature.
                                        - most have high electrical conductivity and can have a
                                        variety of oxidation states when forming compounds. Many
                                        transition metals are colored and as a result cause some
                                        of their ionic compounds to be colored.

                                       Naturally occurring copper:
                                       Copper occurs as the metal and
                                       as minerals such as blue azurite
                                       [2 CuO3·Cu(OH)2] and malachite

Small amounts of transition metal compounds are used to color
glass: Cobalt oxide (Co2O3, blue); Copper or Chromiun oxides
(green); Nickel oxides (purple); Uranium oxides (iridescent green).
The inner transitional metals contain 2 main               The Inner Transition Metals
groups, the lanthanide series, and the actinide
- like the Groups 1A and 2A elements, the inner
transition metals are highly reactive. They catch
fire in air easily, and react with water to liverate H2
gas. Their reactivity makes them of very limited                                             actinides
use structurally.

 The lanthanide series include the 14 elements                           Gadolinium (Gd), a lanthanide metal,
that proceed from atomic numbers 58 to 71. They                           is relatively stable in dry air. Its
                                                                          compounds are used in metallurgy,
are shiny and silvery white metals, and tarnish                           optical instruments, and as
easily when exposed to air. All lanthanides form                          contrasting agents in medical
                                                                          magnetic resonance imaging (MRI). MRI
+3 ions as their principal chemical species.                                                                colonography

 The actinide series include the 14 elements that proceed from atomic numbers 90 to 103.
These elements are unstable and their nuclei are prone to undergo radioactive decay. Unlike the
lanthanides, the actinides show a variety of oxidation states. Uranium, for example,
has compounds in each of the states, +3, +4, +5, and +6.
Uranium (U) occurs naturally in soil and rock in trace amounts. When                   Uranium ore
refined, it is silvery white and weakly radioactive. It is 70% denser
             than lead and is used in military high-density projectiles
             to destroy heavily armored targets.
                                 Glass containing uranium has a yellow green hue. Under UV light,
                                 uranium glass glows fluorescent green. Due to its radioactive
                                 nature, Uranium glass is used only in decorative items.
                   Metals, Nonmetals, and Metalloids
The heavy zig-zag line separates the elements into the metals and the nonmetals:
- all elements on the left of the zig-zag line are metals (except Hydrogen).
- the nonmetals are on the right of the zig-zag line.
- the elements located along the zig-zag line are metalloids (B, Si, Ge, As, Sb, Te).

Inner Transition
  Some Characteristics of a Metal, a Nonmetal, and a Metalloid

Silver (Ag)             Antimony (Sb)            Sulfur (S)
 Metal                  Metalloid               Non-metal
 Shiny                  Blue-grey, shiny        Dull, yellow
 Very ductile           Brittle                 Brittle
 Malleable              Shatters when           Shatters when
                          hammered                 hammered
 Good conductor of      Poor conductor of       Poor conductor,
 heat and electricity     heat and electricity     good insulator
 Density = 10.5 g/ml    Density = 6.7 g/ml      Density = 2.1 g/ml
 Melting point 962C    Melting point 630C     Melting point 113C
               Essential Elements
The essential elements are vital to life that a deficiency in any one
will result in either death, severe developmental abnormalities, or
chronic ailments.
- of the 116 elements known, 11 are predominant in biological
systems. In humans, these 11 elements constitute 99.9% of the
total number of atoms present, but 4 of these elements --- C, H,
N, and O --- account for 99% of the total. The other 7 elements
(Na, K, Ca, Mg, P, S, and Cl) comprise only 0.9% of the total            atoms
in the body; these generally occur as ions, such as Na+,                K +,
Mg2+, Ca2+, Cl–, and HPO42–.
- the 11 essential elements are all “light” elements, having atomic
numbers less than 21. Another 17 elements are also required by
most by not all biological systems. These are generally “heavier”
elements with atomic numbers greater than 18. They are about
evenly divided between metals, nonmetals, and metalloids.
         Some Important Trace Elements In The Body
Some metals and nonmetals known as trace elements are essential to the proper functioning of
the body. Although they are required in very small amounts, their absence can disrupt major
biological processes and cause illness. The trace elements listed below are present in the body
combined with other elements.

DV = Adult Daily Value
Anemia, meaning "without blood", is a condition in        Iron Deficiency Anemia
which there is a deficiency of red blood cells (RBCs)
and/or hemoglobin, the RBC protein that binds oxygen (O2) for transport to tissues.
Without adequate RBCs or hemoglobin, blood has reduced ability to deliver O2 to
local tissues, causing tissue hypoxia (low oxygen level in tissue). An anemic person
often looks pale and fatigues easily.

- a deficiency in body iron can lead to iron deficiency anemia (IDA), a condition
in which blood lacks adequate healthy red blood cells. The body needs the element
iron to make hemoglobin, an iron-rich substance that gives blood its red color and enables
RBCs to carry O2 from the lungs to all parts of the body. In IDA, the RBCs are smaller than
normal (microcytic) and have an increased zone of central pallor (hypochromic).
                                                                      Normal RBCs with biconcave
                                                                      shape and central pallor.

 Each RBC contains several hundred
 hemoglobin (Hb) molecules, each
 carrying four Fe atoms.                       RBCs from IDA patient are
                                               mirocytic and hypochromic.
Symptoms of Iron Deficiency Anemia may include:
 - extreme fatigue, muscle weakness
 - pale skin, often cold hands and feet
 - shortness of breath, lightheadedness
 - headache
 - hair loss
Other IDA symptoms may include:
 - inflammation of the tongue (glossitis)                    koilonychia
 - brittle nails that are spoon-shaped (koilonychia)
 - scaling and fissure formation on the lips (cheilosis)
 - unusual cravings for non-nutritive substances such as ice, dirt, or pure starch (pica)
 - poor appetite, especially in infants and children with iron deficiency anemia

 IDA is common in 3 groups of population:
1) Pregnant women and growing children (esp. in developing countries).
Increased body demands for iron in pregnant women and growing children are
not being met with daily diet.
2) Menstruating women, especially those with menorrhagia (heavy monthly periods).
Loss of iron due to heavy menstruation is not being compensated by dietary intake.
3) Older population. Chronic occult (not readily visible) blood loss can occur as a result of
gastric ulcers, stomach and colon polyps, stomach and colon cancers, hemorrhoids, and
others. It only takes about 1 to 2 teaspoons of blood loss daily to exceed iron absorption.
         Zinc Deficiency Leads to Acrodermatitis Enteropathica
Zinc is an essential component of the diet. It is found in milk, meat, shellfish and wheat germ.
Foods of plant origin are mostly low in zinc. Zinc is needed to assist metalloenzymes (proteins
containing a metal atom) that are involved in many cellular processes throughout the body and
in the normal functioning of the brain.
 A deficiency in dietary zinc can lead to the following characteristic manifestations:
- skin inflammation (dermatitis) that evolve into crusted, blistered, pus-filled and eroded
  lesions. This occurs particularly around body openings such as the mouth, anus, and eyes,
  and the skin on elbows, knees, hands, and feet.
- diffuse hair loss on the scalp, eyebrows and, eyelashes.
- impaired wound healing; secondary bacterial infection.
- in severe cases, irritability and emotional disturbances are
  evident due to wasting (atrophy) of the brain cortex.
Treatment: Supplemental zinc usually
           eliminates the symptoms.

Clinical Case: a 6-week-old male infant developed a widespread red crusted and eroded
dermatitis that was most prominent on the distal extremities, creases, and around body
orifices. He was listless and had frequent watery stools. His skin lesions and systemic
symptoms cleared within several days of initiating high dose oral zinc supplementation.
Zinc levels obtained at the time of admission to the hospital were extremely low.
                      Dalton’s Atomic Theory
All the elements in the Periodic Table are made up of atoms, which are the smallest
particles of matter that retain the characteristics of the elements.
- in the 19th century, John Dalton developed the Atomic Theory, propounding the idea that
atoms are responsible for the combinations of elements found in compounds. The Dalton's
Atomic Theory has the following postulates:
1. All matter is made up of tiny particles called atoms.
    Example: Salt is made up of Sodium atoms and
             Chlorine atoms.

             2. All atoms of a given element are similar to one another and different from
                atoms of a different element.

3. Atoms of 2 or more different elements combine to form compounds.
.  A particular compound is always made up of the same kinds of atoms
.  that are in the same proportions.
    Ex: Water is made up of 1 Oxygen atom and 2 Hydrogen atoms.

4. A chemical reaction involves rearrangement, separation, or combination of atoms.
.  Atoms are never created nor destroyed in a chemical reaction.
                                           The total number of atoms for each element
          4 Fe + 3 O2  2 Fe2O3
                                           is preserved in a chemical reaction.
                    Particles in the Atom
          In the early 20th century, growing experimental evidence indicated
          that atoms are not solid spheres as scientists have imagined but
          composed of smaller particles called subatomic particles.

● 3 subatomic particles are of interest to us:
 i. Proton has a (+) charge and a mass = 1.7 x 10–24 g.
ii. Neutron has no electrical charge and is therefore neutral; it has the
    same mass as the proton.
iii. Electron has a (–) charge; its mass is about 1/2000 the mass of either the proton or neutron
     and is usually ignored in atomic mass calculations.

                Subatomic particles of like charges
                repel one another, whereas particles
                of opposite charges attract one
                                                                             When dry hair is brushed
                another.                                                     with a balloon on a dry day,
                                                                             it is attracted to the
                - electrical charges have been                               balloon.
                noted early in the development of
                the theory of the atom. Examples of
                electric charge phenomena that we
                see in everyday life are shown
                                               Rub a glass rod with a silk cloth,
                at right.                      it will attract a pile of confetti.
                   Because the value 1.7 x 10–24 g is at such a tiny scale, scientists defined
                   another convention for measuring the mass of matter at the atomic level
                   – the atomic mass unit (amu).

                   Definition: One atomic mass unit (1 amu) is defined as one-twelfth of
                   the mass of an atom of carbon with 6 protons and 6 neutrons.

- such a carbon atom is assigned an exact mass of 12.000 amu and is called Carbon-12.
Masses of other atoms are assigned relative to the mass of carbon-12.
 Ex: an Oxygen atom with 8 protons and 8 neutrons has a relative mass of 16.000.

 The atomic mass (A) of an atom can be estimated by summing its number of protons and
                   A = number of protons + number of neutrons
 Ex: a sodium atom (Na) with 11 protons and 12 neutrons has a mass of A = 23.

 The amu is also called dalton in honor of the 19th century chemist John Dalton. On the
atomic scale:
                mass of a proton = mass of a neutron = 1 amu =
               Summary of Subatomic Particles

Subatomic        Symbol        Electrical       Mass           Location
particle                        charge          (amu)          in atom

Proton           p or p+           +1             1            Nucleus
Neutron          n or n0           0              1            Nucleus
Electron           e–              –1           1/2000         Outside nucleus

- the mass of a proton is ~1.7 x 10–24 g
- 1 amu = 1/12 the mass of carbon-12 (which has 6 protons
                                            and 6 neutrons).
- on the amu scale: mass of proton = mass of neutron.
- 1 amu is also known as 1 dalton.
- mass of an electron is usually ignored.
        Rutherford’s Gold Foil Bombardment Experiment
While various experiments have indicated the existence of subatomic particles
such as the electron and the positively charged particles, scientists were not
certain how these particles are arranged in an atom. Further, if electrons are
very much lighter than atoms, then the positively charged particles must carry
the mass of the atom. JJ Thomson suggested that atoms are spheres of
positive charge in which light, negatively charged electrons are embedded, much as raisins
might be embedded in the surface of a pudding. This came to be known as the Raisin Pudding
Model of the Atom.
● In the early 1900s, Ernest Rutherford
performed a gold foil bombardment experiment
using a beam of positively charged particles that
have the same mass as the Helium nucleus (e.g.
 particles). Where he expected the  particles to
pass undeflected through the gold foil, to his
              surprise he found a few were
                deflected at various angles. Some
                even came backwards.

                           Rutherford’s Gold Foil Experiment. A beam of (+)charged  particles was
                           directed at a thin gold foil. A fluorescent screen was used to detect
Ernest Rutherford          particles passing through. Most particles passed through the foil, but some
                           were deflected from their path. A few were even deflected backward.
 Rutherford reasoned that only a very concentrated positive charge in a tiny space within the gold
atom could possibly repel the fast-moving,  particles enough to reverse the  particles’ direction.
He hypothesized that the mass of this positively charged region, which he called the nucleus,
must be larger than the mass of the  particle. He went on to argue that the reason most of the 
particles were undeflected was that most parts of the atoms in the gold foil were empty space.

                         Top: Expected results of Rutherford's gold foil experiment:
                         Alpha particles passing through the raisin pudding model of
                         the atom undisturbed.

                         Bottom: Observed results: Some of the particles were deflected,
                         and some by very large angles. Rutherford concluded that the
                         positive charge of the atom must be concentrated into a very
                         small location: the atomic nucleus.

      A New Model Of The Atom
      From the results of the gold foil experiment, Rutherford proposed
      a new model of the atom in which all of the positive charges
      and most of the mass of the atom are concentrated in a very
      small core, the nucleus. The electrons occupy the rest of the
      space in the atom. Most of the atom is actually “empty space”.
                            Structure Of The Atom
As a result of Rutherford’s gold foil experiment (and other scientists’), the general structure of
the atom has come to accepted as having a small, dense core region known as the nucleus,
where the subatomic particles protons and neutrons are located. The nucleus is situated at
the center of the atom and has a positive charge due to the presence of protons.
- most of the rest of the atom is empty space, occupied by
fast moving electrons.
                 Atomic Number and Mass Number
An element is defined by the number of protons in its nucleus. It is the number of protons that
give an element its characteristics and distinguish it from another element. The number of
protons in the nucleus of an atom is also referred to as the atomic number, which is the whole
number written above the atomic symbol in the Periodic Table.
● The Periodic Table arranges the elements in order of increasing atomic number.
Example: Lithium has an atomic number = 3  lithium has 3 protons
         Carbon has an atomic number = 6  carbon has 6 protons
● An atom is electrically neutral, that is, all the (+) charges of the
protons in the nucleus are balanced by an equal number of (–) charges
in the electrons. Thus, there are as many electrons in a neutral atom                 as
there are protons.
- in other words, in a neutral atom, the atomic number, which is the
number of protons in an atom, is also the number of electrons.

Atomic Mass: The mass of an atom is the sum of its numbers of protons and neutrons:

                    Atomic Mass = # Protons + # Neutrons
 - Oxygen (O) atom has 8 protons and 8 neutrons  atomic mass = 16
 - Iron (Fe) atom has 26 protons and 30 neutrons  atomic mass = 56
 - Gold (Au) atom has 79 protons and 118 neutrons  atomic mass = 197
             Atoms and their Subatomic Particles
                       Atomic       Mass      Number         Number         Number
Element     Symbol     number      number     of protons   of neutrons     of electrons

Hydrogen       H          1           1           1             0              1

Nitrogen       N           7         14           7             7              7

Chlorine       Cl        17          37           17           20              17

Iron           Fe        26          56           26           30              26

Gold           Au        79          197          79          118              79

- note that the atomic number and the number of electrons are identical.
- the mass number is the sum of the numbers of protons and neutrons.
- atoms with increasingly higher atomic mass tend to have increasingly higher number
 of neutrons relative to the number of protons.
        Number of protons ? neutrons ? electrons ?

 An atom of phosphorus (P) has an atomic mass = 31
- what are the number of protons ?
- what are the number of electrons ?
- what are the number of neutrons ?

 Would you use atomic number, atomic mass, or both to obtain the
 following ?
- the number of protons in an atom ?
- the number of neutrons in an atom ?
- the number of electrons in a neutral atom ?
- the number of particles in the nucleus ?
                           Writing Atomic Symbol
An atom is written with a convention that gives information on its atomic number and mass

      24    top number indicates mass number (# protons + # neutrons)
      12    bottom number indicates atomic number (# protons)

- note that this writing convention is not the same way as the Periodic Table presents the
elements. In the Periodic Table, the top number represents the number of protons instead of
the mass, but this is the exception that is found in the Periodic Table.

                                       In the Periodic Table, the atomic number is commonly
                                       listed as a superscript or at the top of the symbol of
                                       the element (and the mass number at the bottom).
It was found that not all atoms of the same element have exactly the same mass number.
That is, if you can examine a sample of most any element at the atomic level, you will find
that while the majority of the atoms in the sample has the same mass number, a small
percentages of these atoms have slightly different mass numbers.
 Ex: Magnesium (atomic number 12) can have 3 slightly different masses that differ
     in their number of neutrons:
         24                  25                   26
        Mg (79%)            Mg (10%)             Mg (11%)
         12                  12                   12
- all atoms of the element Magnesium has the same number of protons (12 protons). The
percentages given in parentheses indicate the average natural abundances of these atoms
in a pure sample of magnesium.
- however, some has: 12 neutrons ( to give atomic mass 24)
                     13 neutrons ( to give atomic mass 25)
                     14 neutrons ( to give atomic mass 26)
- we called atoms of the same element that have different masses      Purified magnesium is a
                                                                      silvery, white metal. By
due to variable numbers of neutrons isotopes. While isotopes of       mass, it makes up 2% of the
                                                                      Earth’s crust and and the
the same element have different masses, their chemical behavior       4th most abudant element in
is very similar.                                                      our body.
 Examples of Isotopes
Atoms that have the same number of protons but
different numbers of neutrons are called isotopes.
The element Hydrogen, for example, has three
commonly known isotopes ─ protium (1H),
deuterium (2H), and tritium (3H). Helium and Lithium
each has two stable isotopes.

- some isotopes are classified as stable isotopes
and others are classified as unstable, or radioactive,
isotopes. Stable isotopes maintain constant
concentrations on Earth over time.

                                          - by contrast, unstable isotopes are atoms that
                                          disintegrate (or decay) at predictable and measurable
                                          rates to form other isotopes by emitting a nuclear
                                          electron or a helium nucleus and radiation. These
                                          isotopes continue to decay until they reach stability. As
                                          a rule, the heavier an isotope is than the
                                          most common isotope of a particular
                                          element, the more unstable it is and the
                                          faster it will decay.
                     Average Atomic Mass (Weight)
The percentages of the different isotopes of an element are
experimentally determined. Because a sample of an element
often has a mixture of 2 or more isotopes of different abundances,
the mass of an element as shown on the Periodic Table is actually
the average of all the masses of the isotopes that are found for
the given element. This is one reason why the atomic masses
listed on the Periodic Table are not a whole number.

 Ex: Calculate for the average mass of Magnesium (Mg) based on their isotope abundances
     (magnesium-24: 79%; magnesium-25: 10%; magnesium-26: 11%).

Solution: Weighted average mass: 24 (0.79) + 25 (0.10) + 26 (0.11) = 24.32
- the weighted average mass is calculated by multiplying the mass of each isotope by its
fractional abundance and summing all the isotopic masses. Because the average mass 24.32 is
closer to isotope 24 than isotope 25 or 26, Mg-24 isotope is the most prevalent in a magnesium
                                                                       Atomic         % Natural
                                                                        Mass         Abundance

  Isotopes of Neon. Based on natural abundances,         20Ne         19.99          90.9%
  the average atomic mass of neon is closest to          21Ne         20.99          0.3%
  which whole number ?                                   22Ne         21.99          8.8%
                             Practice Problems
1. Copper contains two isotopes, copper-63 and copper-65. Are there more atoms
   of copper-63 or copper-65 in a sample of copper ?

2. What are the number of protons, neutrons, and electrons in the following
   isotopes ?                                          Aluminum metal
        27                 106
         Al                  Cd
        13                  48

3. Write the atomic symbols for isotopes with the following:
   i. 26 electrons and 30 neutrons                             Pure cadmium metal is soft,
                                                               malleable, ductile, and has a
  ii. a mass number 24 and 13 neutrons                         white bluish color. It is
                                                               similar to zinc.
                                                     We learn that an atom consists of a central
                                                     nucleus made of protons and neutrons, and
                                                     about which revolve electrons like planets of
                                                     the solar system revolve about the sun.
          n=1                                        - however, electrons don't just move about
                                                     randomly within the space around the nucleus,
        Niels Bohr’s Model of the                    but rather there are distinct levels of energy
                                                     in which electrons with certain energies occupy
              Nucleus                                and move about.
                                                     - Niels Bohr was one of the first physicists
                                                     early in the 20th century who propounded the
                                                     idea of quanta of energy to explain the atomic
                                                     line spectrum of the Hydrogen atom.

 According to Niels Bohr, there are energy levels called principal quantum energy levels (also
called main energy levels, or energy shells), conventionally given the symbol n, that are closer
to the nucleus in which electrons with lower energy occupy, and there are quantum energy levels
that are farther away from the nucleus for electrons with higher energies to occupy. The above
diagram shows an atom with the first two principal energy levels n = 1 and n = 2.
- for an electron to move from a lower energy level to the next higher level requires that it gains or
absorbs energy. Conversely, an electron drops from a higher energy level to a lower one by losing
energy so that it now moves closer to the nucleus.
                           We can think of the quantum energy levels of an atom as similar to
                          the rungs on a ladder. The lowest energy level would be the 1st rung,
                          the second energy level the 2nd rung, and so on. As one climbs up or
                          down the ladder, one must step from one rung to the next and can not
                          stop at a level between the rungs.
                          - similarly, electrons in atoms exist in the available energy levels. To
                          move up to the next higher energy level (n+1), an electron must
                          absorb a definite quantum of energy (E = h). Conversely, to fall to
      E = h              the next lower energy level (n–1), an electron must lose the same
                          amount of quantum of energy.
                          - unlike the ladder, the lower energy levels in an atom are far apart
                          compared to the higher energy levels that are closer together.

The maximum number of electrons allowed
in each energy level is given by the formula:        2n2

  Using this formula, one can calculate the total number of
  electrons in any given energy level as shown in the Table
  at right. The lowest energy level (n = 1) can hold up to 2
  electrons; energy level n = 2 can accommodate up to 8
  electrons; etc. In the atoms of the elements known today,
  electrons occupy 7 different energy levels.
 When atoms of an element are energized, such as in the red glow of a metal strip
produced by a hot flame or the white radiance of a bulb filament produced by electrical
energy, the electrons in those atoms absorb a certain amount of energy and jump to
higher energy levels. When some of the energized electrons drop back to lower, more
stable energy levels, energy is emitted in the form of radiation.
                             The “radiation rays” emitted by
                             energized atoms can have
                             different wavelengths (often
                             measured in nm). If these
                             radiation wavelengths fall within
                             the visible range of the
                             electromagnetic spectrum, our
                             eyes can detect them as visible
                             light emission. Depending on
                             the wavelengths of radiation, the
                                    light emission takes on the
                                    characteristic color
                                    produced by the energized
                                    atoms of the element.

                             Radiation within the visible region of the electro-
                             magnetic spectrum produced by hydrogen gas.
                  Line Spectrum of Hydrogen
                            Line Spectrum of Hydrogen
 When H2 gas is excited by an electrical current, it emits radiation that we
 can see as a glow of light. This light is actually a mixture of light energy
 emitted at different wavelengths. When this light is passed through a
 glass prism, it is resolved into four discrete bands of light, each having a
 different wavelength and associated color. Together, the bands of light
 are called the visible line spectrum of hydrogen.

Sunlight is diffracted
in a rainbow of colors
in a glass prism.

                         A glass prism resolves light emitted by energized hydrogen
                         gas into four discrete bands, each with a different
                         wavelength and color.
    Orbitals and Electron Arrangements in Atoms
The principal quantum energy levels (or energy
shells) are equivalent to the Periods of the
Periodic Table.

 Ex: Shell n = 1 corresponds to Period 1, which
contains the elements Hydrogen and Helium
and holds a maximum of 2 electrons (in He).
- similarly, shell n = 2 corresponds to Period 2,
which contains the eight elements Li, Be, B, C, N,
O, F, and Ne and holds a maximum of 8 electrons (in Ne).

Electron Orbitals
The view of electrons orbiting in energy shells around the nucleus of an atom
as planets of the solar system orbiting the sun is a simplistic model of the atom.
An implied assumption of this model is that electrons belonging to the same
energy shell all have equal energy. However, early in the 20th century, with the development of
quantum mechanics – a series of complex mathematical equations that describe the position of
an electron around a nucleus – scientists began to realize that electrons
belonging to a given energy shell can have different energies and, depending
partly on their energies, tend to occupy certain spatial regions with peculiar
shapes called orbitals.
Electron Orbitals. An orbital is defined as a region in space around the nucleus in which
electrons is most likely to be found. Each orbital can hold a maximum of 2 electrons.

There are several types of orbitals, of which four are of interest to us – the s, p, d, and f orbitals.

- an s orbital is spherical with the nucleus at the center.

- a p orbital has 2 lobes, each lobe being the region of space having the highest probability of
finding an electron. The 2 lobes ended at the node where the nucleus is located. There is a set
of three p orbitals arranged in the three directions x, y, and z around the nucleus, forming px, py,
and pz orbitals. These three p orbitals can accommodate up to 6 electrons.

- there are a set of five d orbitals and seven f orbitals; however, their geometry is much too
complex for our purpose. The five d orbitals can house up to 10 electrons, and the seven f
orbitals up to 14 electrons.
Divisions of the s-, p-, d-, and f-Orbitals in the Periodic
The lowest energy shell (n = 1) consists only of a single s orbital.
Shell n = 2 can have both s and p orbitals.
Shell n = 3 can have s, p, and d orbitals.
Higher energy shells (n = 4 and above) can have s, p, d, and f orbitals.
The distribution of the orbitals based on electron arrangements of elements in the Periodic Table
is given in the chart below.
      Energy shell

                Orbitals in Energy Shells n = 1 to n = 4
Orbital distribution for the first 4 energy shells is given the Table below. Recall that the maximum
number of electrons that can be found in a given shell is given by the formula 2n2.
     Electron Level Arrangements for Elements in the First 4 Shells
The electron level arrangement of an atom gives the number of electrons in each energy level.
When a shell is completely filled to its maximum allowed number of electrons, additional
electrons are placed in the next higher shell.
                                            Number of electrons
                             Atomic        in Energy Shell
  Element        Symbol      number        1        2         3            4

  hydrogen         H           1           1
  helium           He          2           2 (1st shell completed)
  lithium          Li          3           2          1
  beryllium        Be          4           2          2
  boron            B           5            2         3
  neon            Ne          10           2          8 (2nd shell completed)
  sodium          Na          11           2         8          1
  magnesium       Mg          12           2          8         2
  aluminum        Al          13           2          8         3
  argon           Ar          18           2          8         8
  potassium       K           19           2          8         8          1
  calcium         Ca          20           2          8         8          2
  scandium        Sc          21           2         8          9          2
  zinc            Zn          30           2         8         18          2 (3rd shell completed)
  gallium         Ga          31           2         8         18          3
  krypton         Kr          36           2         8         18          8
Periodic Table
 The chemical properties of representative elements are
mostly due to their valence electrons, which are the electrons
                                                                   Valence Electrons
in the outermost energy level. The valence electrons are the electrons that participate in
chemical reactions.
- in Chem109, we define the valence electrons as the s and p electrons. The Group Numbers
indicate the valence (outer) electrons for the elements in each vertical columns.
Ex: Group 1A elements (e.g. Li, Na, K, etc) all have 1 valence electron in the outer energy shell.
    Group 2A elements (e.g. Be, Mg, Ca, etc) all have 2 valence electrons in the outer shell.
    Group 7A elements (e.g. F, Cl, Br, etc) all have 7 valence electrons in the outer shell.
Electron-Dot Symbols and the Octet Configuration
The valence electrons are represented as dots placed along the 4 edges of the symbol of the
element. When there are more than 4 electrons, the electrons are paired up. For any given shell,
the valence electrons fill up the s and p orbitals until reaching the Noble Gas configuration with 8
valence electrons. The atom is said to achieve the Octet configuration and chemically becomes
very stable. For this reason, the Noble Gases are unreactive and usually do not form compounds
with other elements.
Trends in Atomic Sizes for the Representative Elements
                                Atomic size can be taken approximately
                                as the distance from the nucleus to the
                                valence electrons on the outermost shell.


                                - atomic size tends to increase going
                                down a group. Each successive
                                increase in the energy level puts the outer
                                electrons further away from the nucleus.

                                - atomic size typically decreases from
                                left to right across a period. Going
                                across a period, as the number of
                                valence electrons increases, the number
                                of protons in the nucleus also increases,
                                thereby pulling the valence electrons
                                closer to the nucleus. This makes the
                                atom shrinking smaller in size.
    First Ionization Energy of the Representative Elements
                                                      In a neutral atom, the (–)charged e–
                                                      move about the (+)charged central
                                                      nucleus. The neutral atom can acquire
                                                      a net (+) charge when an e– is
                                                      removed from the atom. An atom that
                                                      has a positive charge is called a

                                                      To pull an e– away from the
                                                      (+)charged nucleus requires energy.
                                                      The ionization energy is defined as
                                                      the energy needed to remove an
                                                      electron from an atom in the
                                                      gaseous state.

                                                      Ex: Na(g) + energy  Na+(g) + e–

                                                      The ionization energy generally
The ionization energy generally increases going       decreases going down a group
across a period from left to right because as the     because less energy is needed to
number of protons increases, the nuclear attraction   remove an e– as nuclear attraction for
for e– becomes increasingly stronger.                 e– decreases farther from the nucleus.