Chapter 3 Atoms and the
Periodic Table
Reactivity
• Groups reactivity METALS
• As orbits are added the electrons move
further away from nucleus and become
easier to lose , thus as you move down the
group the elements become more reactive
• Nonmetals need electrons therefore the
elements closer to the nucleus are more
capable of accepting electrons and the
elements moving up the group become
more reactive
More information
• Elements moving from left to right in a
period becoming smaller due to the pull of
the nucleus.
• Elements moving down a group become
larger in size due to adding orbits
• The most reactive families have fewer
electrons to lose or gain, therefore groups
1 and 17, why not 18?
• Group 2 more reactive than 3 ;;group 16
more reactive than 15
Gases on periodic table
• All of 18 are gases known as the noble
gases
• Other gases that are found naturally
• H, N, O, F, Cl,
• When these gases appear by themselves
they are written as diatomic gases
• What are Atoms?
– defined - are tiny units that determine the
properties of all matter
• an atom is the smallest part of an element that still
has the element's properties
– introduction
• Democritus
– Greek philosopher
• lived in the 4th century B.C.
• suggest that the universe made of invisible units
called atoms
• defined - Greek word meaning "unable to divide"
• believed that the changes he observed was due to
the movement of the atoms
• unable to provide the evidence needed to convince
people that atoms existed
– Atoms are the building blocks of molecules
• John Dalton atomic theory in 1808
– English school teacher
– widely supported due to supporting evidence
– three parts
• every element is made of tiny unique particles called
atoms that cannot be subdivided
• atoms of the same element are exactly alike
• atoms of different elements can join to form
molecules
• What is in an atom?
– introduction
• less than a 100 years after Dalton published his
atomic theory scientist determined that atoms
could be split further
• today we know there are many different parts of an
atom but only three are used in everyday
chemistry of most substances
– In the nucleus - dense center of the atom
• protons - 1 positive charge with a mass of 1 amu
(atomic mass unit)
• neutrons - 0 charge (neutral) with a mass of 1 amu
– Electron cloud - made of very tiny moving particles
• electrons - 1 negative charge with very little mass 0
amu
• Atoms have no over all charge because they have
an equal number of protons and electrons
– example He (helium) atom
• 2 protons
• 2 neutrons
• 2 electrons
• charge of 2 protons +2
• charge of 2 neutrons 0
• charge of 2 elections -2
• total charge
0
• Models of the Atom
– introduction
• like most scientific models and theories the model of the
atom has been revised many time to explain each new
discovery
– Bohr's model
• Niels Bohr - Danish scientist in 1913
– electrons move in set paths around the nucleus like the planets
orbit the sun
– each electron has a certain energy that is determined by its
path around the nucleus
– energy level
• the path of the possible energies an electron may have in
an atom
– electrons must gain energy to move to a higher energy level
– Modern theory
• by 1925 Bohr's model no longer explained all
observations
– electrons no longer moved in definite paths
– electrons behave like waves vibrating on a string than
like particles
– impossible to determine the exact location, speed, and
direction
• like a fan blade
– try to determine location by shading
• the darker the shading the better the chance to find
an electron
• the whole shaded region is called an electron cloud
• electrons are found in an orbital within each energy
level (orbit or shell)
– orbital - the region in an atom where electrons are found
• exist only when an electron occupies it
– four different kinds of orbitals
– "s" orbital
• simplest
• shaped like a sphere
• only 1 orbital or orientation per orbit
• can contain 2 electrons maximum
– "p" orbitals
• dumbbell shaped
• 3 different orbitals or orientations per orbit
• x, y, and z axis
• 2 electron in each orbital
• 6 electrons maximum
– "d " orbitals
• 5 possible orbitals or orientations per orbit
• 2 electrons each orbital
• 10 electrons maximum
– "f " orbitals
• 7 possible orbitals or orientations per orbit
• 2 electrons each orbital
• 14 electrons maximum
• electrons start off occupying the lowest
level then are added to the next highest
energy level or orbit
– 1st energy level or orbit
– contains only the "s" orbital
– 2 electrons maximum
– 2nd energy level or orbit
• contains
– "s" orbital
• 2 electrons maximum
– "p" orbitals
• 6 electrons maximum
– 8 electrons maximum for the energy level
– 2 in the s and 6 in the p
– 3rd energy level or orbit
• contains
– "s" orbital
• 2 electrons maximum
– "p" orbitals
• 6 electrons maximum
– "d " orbitals
• 10 electron maximum
• 18 electrons maximum for the energy level
– 2 in the s, 6 in the p, and 10 in the d
– 4th energy level
• contains
– "s" orbital
• 2 electrons maximum
– "p" orbitals
• 6 electrons maximum
– "d " orbitals
• 10 electrons maximum
– "f " orbitals
• 14 electrons maximum
• 32 electrons maximum for the energy level
– 2 in the s, 6 in the p, 10 in the d, and 14 in the f
– every atom has one or more valence electrons
• valence electron - is an electron in the outer most
energy level of an atom
• hydrogen has 1 valence electron (the least number)
• neon has 8 valence electrons (the maximum
number)
• Page 76
• Questions 1-7
• Write questions and answers
3.2 A Guided Tour of the
Periodic Table
•
Objectives
– Relate the organization of the periodic table to the
arrangement of electron within an atom.
– Explain why some atoms gain or lose electrons to
form ions.
– Determine how many protons, neutrons, and
electrons an isotope has, given its symbol, atomic
number, and mass number.
– Describe how the abundance of isotopes affects
and element’s average atomic mass.
• Historical prospective
– Developed by
• Dimitri Mendeleev
–
–
–
–
Russian chemist
in 1869
based on repeating properties and atomic mass
he arranged the known elements and left blank spaces
for unknown elements
– Henry Mosley
• the first to group atoms by protons
• Organization of the Periodic Table
– similar elements grouped together
– makes it easier to predict the properties of an
element based on where it is in the periodic
table
– elements represented by their symbols
– order based on the number of protons the
atom has in its nucleus
– Hydrogen has one proton and is the first
elements listed in the Periodic Table
– Period law
• properties of elements tend to change in regular
pattern when elements are arranged in order of
increasing atomic number, or number of protons in
their nucleus
• atomic numbers equals the number of protons
– increases from left to right and top to bottom
• Using the Periodic Table to determine
electronic arrangement
– Periods
•
•
•
•
horizontal rows
1-7
indicates the outer most energy level
Period 1
– 2 elements - H and He
– has only 1 s orbital
• maximum of 2 electrons
• Period 2
– starts with Li and ends with Ne
– contains 1 s and 3 p orbitals
• Period 3
– starts with Na and ends with Ar
– contains 1 s and up to 3 p orbitals
• Periods 4 and 5
– contains 1 s and up to 3 p and 5 d orbitals
• Periods 6 and 7
– contains 1 s and up to 3 p, 5d, and 7 f orbitals
– Groups
• vertical columns
– 1-18
– have similar properties
– have the same number of valence electrons
• Groups 1 and 2
– electrons are going in the s orbital
– Group 1
• H down to Fr
• only one electron in the outer most energy level
• 1 electron in the s orbital
• these elements have 1 valence electron
– Group 2
• Be down to Ra
• 2 electrons in the outer most energy level
• 2 electrons in the s orbital and it is now full
• these elements have 2 valence electrons
• Groups 13 to 18
– are placing electrons in the p orbitals of the outer most
energy level
– these elements have 3 to 8 valence electrons
• Group 13 elements have 3 valence electrons
• Group 18 elements have 8 valence electrons
• this is the maximum number of valence electrons
• the p orbitals are full
• Valence electrons equals the last digit of the Group
number
• Groups 3 to 12
– are placing electrons in the d orbital of the next lower
energy level
– they have 2 valence electrons as far as this class is
concerned
• The Lanthanoid Series and Actinoid Series
– are placing electrons in the f orbital of the energy level 2
place back
– they have 2 valence electrons as far as this class is
concerned
• Atoms in Group 18 have full outer energy levels
– 8 is the maximum for an outer level
– except for level 1 He has 2
– non reactive (inert)
• united video Elements of Chemistry: The Periodic
Table (20:00 min.)
http://www.unitedstreaming.com/search/assetDetai
l.cfm?guidAssetID=F59A819C-DB1B-48E6-90D7DF3829C74230
• Elements are reactive because their outer
most energy levels are only partially filled.
• Some Atoms Form Ions
– Ionization
• defined - atoms that may gain or lose valence
electrons so that they have a full outermost
energy level
• no longer the same number of protons and
electrons
• it has a net electrical charge
– ion
• defined - an atom or group of atoms that has lost
or gained one or more electrons and therefore has
a net electric charge
• cation
– defined - an ion with a positive charge
– example: Li has 1 valence electron
• 2 electrons in the 1st energy level
• 1 electron in the 2nd energy level
• when the valence electron is removed Li becomes a
positive ion Li+
• Li+ ion
Li atom
3 protons +3
3 protons +3
2 electrons -2
3 electrons -3
charge
+1
0
• the other elements in Group 1 form +1 cations by
having only one valence electron
• anion
– defined - an ion with a negative charge
– example: F has 7 valence electrons
• 2 electrons in the 1st energy level
• 7 electrons in the 2nd energy level
• easier to gain 1 electron than lose 7 electrons to
become a negative ion.
• F- ion
F atom
9 protons +9
9 protons +9
10 electrons -10
9 electrons -9
charge
-1
0
• the other elements in Group 17 form -1 anions by
having 7 valence electrons
• How Do the Structures of Atoms Differ
– Atomic number
• defined - the number of protons in the nucleus of
an atom
– remember atoms are always neutral because they have
equal number of protons and electrons
• the simplest atom H has only 1 proton and 1
electron
– atomic number is 1
• the largest naturally occurring atom U has 92
protons and 92 electrons
– atomic number is 92
– Mass number
• defined - the total number of protons and neutrons
in the nucleus of an atom
• F has 9 protons and 10 neutrons for a mass
number (A) = 19
• the mass number can vary from atom to atom of
the same element
– Isotopes
• defined - atoms having the same number of
protons but different number neutrons
– example: H has 2 isotopes
• the first is the protium
– the atom of H (the most common)
– has only one proton and 0 neutrons
– a mass number of 1
– the second isotope is Deuterium
• sometimes called "heavy Hydrogen"
• 1 proton and 1 neutron
• a mass number of 2
• only 1 out of every 6000 H are Deuterium
– the third isotope is Tritium
• 1 proton and 2 neutrons
• mass number of 3
– All three are hydrogen, only one proton, but have
different masses due to the neutrons.
– Calculating the number of neutrons in an
atom
• average atomic mass
– defined - the weighted average of the masses of all
naturally occurring isotopes of an element
• This is found under the Symbol on the Periodic
Table
– round this number to the nearest whole number
– subtract the atomic number
– example: C
• average atomic mass 12.011 = 12 mass number
• atomic number
-6
number of neutrons
6
• this is for the most common Carbon atoms (carbon 12)
• the isotopes for C will be those with different number
of neutrons like carbon - 14
• Mass number 14
atomic number - 6
neutrons
8
• Rules for Electron configuration
– Find the total number of electrons (atomic
number).
– Find the number of energy levels (the period
number).
–
Draw the orbits
– Find the electrons in the last energy level.
• For Groups 1& 2 use the Group number.
• For Groups 13 - 18 use the last digit of the Group
number (3 - 8).
• For He always 2 electrons.
• For Group 3 - 12 assign 2 electrons
– Subtract the electrons from the total as you
place them in their energy level.
– Fill in the inner energy levels with the
remainder of the electrons starting with the
first energy level.
• Use the following pattern when they are the inner
energy levels.
–
–
–
–
–
–
1st energy level - 2 electrons
2nd energy level - up to 8 electrons
3rd energy level - 8 or 18 electrons
4th energy level - 8, 18, or 32 electrons
5th energy level - 8, 18, or 32 electrons
6th energy level - 8 or 18 electrons
• Remember to subtract as you add them to their
energy levels.
• Examples: Br K and Bi
• Page 85
• Questions 1-7
• Questions and answers
3.3 Families of Elements
• Objectives
– Locate alkali metals, alkaline-earth metals,
and transition metals in the periodic table.
– Locate semiconductors, halogens, and noble
gases in the periodic table.
– Relate an element’s chemical properties to
the electron arrangement of its atoms.
• Groups are sometimes called families
– each is unique yet share certain similarities
– elements have common chemical and
physical properties
• they have the same number of valence electrons
• How elements are classified
• Metals
– the majority of all elements
– most are
•
•
•
•
•
•
solids
luster - shiny
ductile - can be stretched
malleable - can be shaped
good conductors of heat and electricity
form cations only
• Alkali Metals
– the highly reactive metallic elements located
in Group 1 of the Periodic Table
•
•
•
•
•
Li, Na, K , Rb, Cs, and Francium
form cations with a +1 oxidation number
highly reactive
soft
luster - shiny
– Uses
• NaOH (sodium hydroxide) used to manufacture
–
–
–
–
•
•
•
•
paper
soap
synthetic fabrics
petroleum refining
NaCl -table salt
KCl - table salt substitute
K - used in fertilizers
Na+ and K+ are important for proper functioning of
nerves in our bodies.
• Alkaline Earth Metals
– reactive metallic elements located in Group 2
• Be, Mg, Ca, Sr, Ba, and Ra
• form cations with a +2 oxidation number
• less reactive than the Alkali Metals
– requires more energy to remove the 2nd electron than
the 1st electron from an energy level
• light
• good structural strength
– uses
• Ca - animals shells, limestone, marble, bones, and
teeth
• Mg - airplane frames, activates enzymes, flares,
Epson salt, and milk of magnesia
• Transition Metals or Elements
– located in Groups 3 - 12
– conducts heat and electricity like other metals
– form multiple cations
• some up to 4 different cations
– frequently form colorful compounds such as
rubies and emeralds
– uses
• Ag – a better conductor than Au
• Au does not corrode or tarnish under ordinary
conditions
– great for connectors for computers and other electronic
devices
• Fe, Co, Cu and Mn play important roles in our
body chemistry
• Hg the only metal that is a liquid at room
temperature
– flows easily
– does not stick to glass makes it good for thermometers
• Fe, Co, and Ni
– in the same period
– the only metals that can be magnetized
• Cu, Ag, and Au
– in the same family
– called the coinage metals
• radioactive isotopes
– defined - nuclei are continually decaying to produce
different elements
– used at times to detect cancer in the soft tissue of our
bodies
• Nonmetals
– Except for H they are found on the right hand
side of the Periodic Table
•
•
•
•
•
some of the elements in Groups 13 - 16
all the elements in group 17 - 18
do not conduct electricity or heat
brittle
no luster - dull
–C
• occurs in three forms naturally
– graphite
– diamonds
– fullerenes
• combine with other elements to form millions of
compounds
– sugars, chlorophyll, gasoline, and rubber to name a few
– O, N, and S
• common nonmetals
• form anions of oxide-2, sulfide-2, and nitride-3
• most plentiful gases in the atmosphere are N and
O
• S is an odorless yellow solid
– many S compounds are known for their terrible smell
• rotten eggs, H2S
• skunk spray
• Halogens
– Group 17
– slightly reactive
– forms salts with metals
– form ions with a -1 oxidation number
– 4 of the 7 diatomic elements
• 2 atoms per molecules as an element
• the other 3 are H2, N2, and O2
– Cl2 yellowish green poisonous gas used to kill
bacteria in water
– F2
• most reactive non metal
• a poisonous yellowish gas
– Br2 a dark red liquid
• Noble Gases
– Group 18
– non reactive gases exist as single atoms and
not as compounds
– He lighter than air and used in balloons
– Ne used in signs because of its reddish
orange glow
– Ar used in light bulbs
• Metalloids or Semiconductors
– have properties of metals and nonmetals
– weak conductors of electricity and heat
– solids
– white or gray in color
– B, Si, Ge, As, Sb, Te, and Po form a stair step
downward from left to right
– metals are to the left of the metalloids and the
nonmetals are to the right
• Page 94
• Questions 1-7
• Questions and answers
3.4 Using Moles to Count Atoms
• Objectives
– Explain the relationship between a mole of a
substance and Avogadro’s constant.
– Find the molar mass of a n element by using
the periodic table.
– Solve problems converting the amount of an
element in moles to its mass in grams, and
vice versa.
• Counting Units
– dozen (12 items)
– bushel (32 qt container)
– reams (500 sheets)
– pairs (2)
• Mole is used for counting very small
particles
– abbreviated mol.
– a collection of 602 213 670 000 000 000 000
000 particles
– usually written as 6.022 x 1023 particles per
mole
– known as Avagodro's number or constant
• named for Amedeo Avagodro
– an Italian that lived from 1776 - 1856
– a lawyer interested in mathematics and physics
– 1st to make a distinction between atoms and molecules
– This constant was determined by Joseph
Loschmidt
• German physicist
• in 1865
– 1 mole of popcorn kernels would cover the
entire US to a height of 500 km (310 mi)
• not a good way to count popcorn
• Molar mass
– defined - the mass in grams of 1 mol of a
substance
– The molar mass of an element in grams is the
same as its average atomic mass in amu
– conversion factor
• defined - a ratio equal to one that expresses the
same quantity in two different ways
• 10 gumballs = 21.4 g
• can be written as 10 gumballs / 21.4 g or 21.4 g /
10 gumballs
• What is the mass of 50 gumballs?
– 50 gumballs x 21.4 g / 10 gumballs = 107 g
– p. 98 practice factors 1 - 3
– Relating moles to grams
• 1 molar mass of element or
1 mol of element
•
1 mol of element
1 molar mass of element
• Fe has 55.85 amu therefore 55.85 g Fe
•
1 mol Fe
– determine the mass in grams of 5.5 mol of iron.
• 5.50 mol Fe x 55.85 g Fe = 307 g Fe
•
1 mol Fe
• p. 99 practice converting Amount to mass
1&4
– Converting mass to amount
• Determine the amount of iron present in 352 g of
iron.
• 352 g Fe x 1 mol Fe = 6.3 mol Fe
•
1
55.85 g Fe
– How many moles are in 536 g of copper?
• 536 g Cu x 1 mol Cu
= 8.44 mol Cu
• 1
63.55 g Cu
– How many moles are present in 12.1 g of
sulfur.
• 12.1 g S x 1 mol S = .377 mol S
•
1
32.07 g S
Page 100
questions 1 – 9
• Write questions and answers (show work
on the problems)