Download Periodic Table Ch4 Honors

Do Now
• Define an element.
• What relationship exists between atomic
number, protons and electrons?
• Do you notice any pattern or relationships
between elements of the same row or
column?
The Periodic Table
Mendeleev’s Periodic Table
Dmitri Mendeleev
History of the Periodic Table
• By 1860 more than 60 elements had been
discovered.
• Chemists learned about the new elements by
reacting them with other elements to form new
compounds.
• In 1865, John Newlands, an English chemist,
arranged the known elements according to their
properties in order of increasing atomic mass.
• The properties of elements seemed to repeat every
eight elements. This pattern is known as the law of
octaves.
History of the Periodic Table
• In 1869, Dmitri Mendeleev, a Russian chemist, used Newland’s
observations and other observations to produce the first PERIODIC Table.
• He placed each element on a card with the element’s atomic mass,
chemical and physical properties. He arranged the elements in many
different ways according to various properties and looked for trends.
• When he arranged the elements according to increasing atomic mass, he
noticed that certain similarities occurred at regular intervals. This was
considered a “periodic pattern”.
• He arranged all the elements in a table according to increasing atomic
mass, starting a new row every time the pattern repeated, and the
elements with similar properties fell in the same vertical columns (with a
few exceptions).
• He had some blank spaces in his table where he thought new elements
would be discovered. These blank spaces were used to predict the
existence and discovery of new elements.
History of the Periodic Table
• By 1886 three new elements were discovered
that followed Mendeleev’s predication.
• Most chemists were persuaded to accept his
table, labeling Mendeleev the founder of the
periodic law.
History of the Periodic Table
• In order to keep the pattern of properties, Mendeleev had to
switch some elements out of order based on atomic mass.
• In 1911, Henry Moseley, an English chemist, was examining
the spectra of 38 different metals. He noticed that the
wavelengths of spectra lines correlated to atomic numbers,
not atomic mass.
• Moseley discovered a new pattern and organized the
elements according to their increasing nuclear charge, or
atomic number, i.e. increasing number of protons.
• When the elements were arranged by increasing atomic
number, the discrepancies in Mendeleev’s table disappeared.
Modern Russian Table
Chinese Periodic Table
Stowe Periodic Table
A Spiral Periodic Table
Triangular Periodic Table
“Mayan” Periodic Table
Meet the Elements Music Video
• http://www.youtube.com/watch?v=d0zION8xj
bM
• https://www.youtube.com/watch?v=Uy0m7jn
yv6U
• http://www.privatehand.com/flash/elements.
html
Do Now
• Take out your Reference Tables and open to
the Periodic Table of Elements
• What do you think of when you hear the word
metal?
• What is a nonmetal?
• Do you know of any examples of each?
The Periodic Table consists of:
• METALS
• NONMETALS
• METALLOIDS
Properties of Metals
• Metals are good
conductors of heat and
electricity
• Metals are malleable
(can be hammered or
rolled into sheets)
• Metals are ductile (can
be made into wire)
• Metals have high
tensile strength
• Metals have luster
Examples of Metals
Potassium, K reacts with water
and must be stored in mineral oil
Copper, Cu, is a relatively soft
metal, and a very good electrical
conductor.
Examples of Metals
Zinc, Zn, is more stable than
potassium
Mercury, Hg, is the only metal
that exists as a liquid at room
temperature
Mixing metals
• Alloy = mixture of a metal with another
element, usually another metal
• Alloys have properties different from the
individual elements, usually eliminating some
disadvantages.
– Alloys are usually harder and more resistant to
corrosion
•
•
•
•
Ex: brass = copper + zinc
Ex: sterling sliver = copper + silver
Ex: steel = iron + carbon + manganese + nickel
Ex: Stainless steel = iron + chromium
Properties of Nonmetals
• Carbon, the graphite in “pencil lead” is a great
example of a nonmetallic element.
• Nonmetals are poor conductors of heat and
electricity
• Nonmetals tend to be brittle
• Many nonmetals are gases at room
temperature
Examples of Nonmetals
Sulfur, S, was once known as
“brimstone”
Microspheres of phosphorus, P, a
reactive nonmetal
Examples of Nonmetals
Graphite is not the only pure form
of carbon, C. Diamond is also
carbon; the color comes from
impurities caught within the crystal
structure
Bromine is a nonmetal that
exists as a liquid at room
temperature.
Properties of Metalloids
• Metalloids straddle the border between
metals and nonmetals on the periodic
table.
• They have properties of both metals and
nonmetals.
• Metalloids are more brittle than metals,
less brittle than most nonmetallic solids
• Metalloids are semiconductors of
electricity
• Some metalloids possess metallic luster
*note them on your
Periodic Table in your
Reference Table!
Silicon, Si – A Metalloid
Silicon has metallic luster
Silicon is brittle like a nonmetal
Silicon is a semiconductor of
electricity
Other metalloids include:
Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te
Elements classified as metals and
nonmetals
http://www.ptable.com/
Lab
• Alien Nation
Do Now
• Take out your Reference Table and open to the
Periodic Table of Elements
• Do you notice any patterns as you look at the
properties of elements?
– Hints: scan left to right, and scan from the top
down
The Periodic Table
• PERIODIC LAW
– The physical and chemical properties of the
elements are periodic functions of their atomic
numbers (elements with similar properties appear
at regular intervals).
• THE MODERN PERIODIC TABLE
– An arrangement of the elements in order of
increasing atomic number so that the elements
with similar properties fall in the same column or
group.
Periodic Table with Group Names
PERIODS= horizontal rows
GROUPS= vertical columns
VALENCE ELECTRONS
• Elements within the same group have the
same number of valence electrons.
VALENCE ELECTRONS
• Most loosely bound electron is called the VALENCE
ELECTRON!
• Valence electrons participate in chemical reactions,
so elements with similar valence electrons react in
similar ways.
• Noble gases are stable with _8_ valence electrons.
• All elements with the exception of Hydrogen and
Helium want _8_ valence electrons to be stable.
• Hydrogen and Helium want _2_ valence electrons to
be stable.
Orbital filling table
Orbital filling table
s-BLOCK and p-BLOCK are considered
“MAIN GROUP ELEMENTS”
s-Block: Includes GROUP 1 and GROUP 2
• GROUP 1 - Alkali Metals
• GROUP 2 – Alkaline Earth Metals
p-Block: Elements vary greatly in properties,
includes GROUP 13 through GROUP 18
• GROUP 17 – Halogens
• GROUP 18 – Noble Gases
GROUP 1
Alkali Metals
• Properties:
–
–
–
–
–
–
–
–
Metal
Easily lose valence electron
1 valence electron
Chemically reactive – do not
occur as free elements in
nature
Soft, silvery
Good conductor of electricity
React violently with water
React with halogens to form
salts
EXCEPTION:
Hydrogen is in group 1,
but has different properties than the alkali metals
*note that Hydrogen is a nonmetal
on your Periodic Table in your
Reference Table!
Hydrogen
• Most common element in the universe.
• It consists of one proton and one electron.
• It can react with many other elements:
– With oxygen to make water
– With carbon to make organic compounds
– With nitrogen to make ammonia
GROUP 2
Alkaline Earth Metals
• Properties:
– They are metals, but
have less metallic
characteristics than
Alkali Metals
– 2 valence electrons
– Chemically reactive – do
not occur as free
elements in nature
– Harder, denser, stronger,
higher boiling points,
and slightly less reactive
than alkali metals
Group 13
• Properties:
– 3 valence electrons
– Lose 3 electrons to form
3+ ions
– Boron is a metalloid and
does not form a 3+ ion
readily
Group 14
• Properties:
– 4 valence electrons
– Carbon is a nonmetal
with two distinct
crystalline forms known
as allotropes (graphite
and diamond)
– Silicon and Germanium
are metalloids used in
the computer industry
– Tin and Lead are metals
Group 15
• Properties:
– 5 valence electrons
– Nitrogen is a nonmetal
that exists naturally as
N2, a diatomic molecule
with a triple bond.
– Phosphorus is a
nonmetal that exists in
two allotropic forms (red
and white)
Group 16
• Properties:
– 6 valence electrons
– Oxygen is a nonmetal
that exists naturally in
two allotropic forms
(O2 and O3)
– Sulfur is a nonmetal that
exists naturally in several
allotropic forms
– Selenium is a nonmetal
– Tellurium is a metalloid
– Polonium is a metal
GROUP 17
Halogens (“Salt makers”)
• Properties:
– Most reactive nonmetals.
– 7 valence electrons – need
1 electron to become
stable
– They vigorously react with
metals to form salts.
•
•
•
•
F2 - most reactive element
Cl2
Gases at STP
Br2
Liquid at STP
I2
Solid at STP
GROUP 18
NOBLE GASES (or INERT GASES)
• Properties:
– Full set of valence electrons:
most elements have 8
valence electrons, except
Helium with 2 valence
electrons, but it is still
associated with this group
because its properties match
these elements.
– Extremely stable and occur as
monoatomic gases in nature
– Although they do not readily
combine with other
elements, but compounds of
Krypton and Xenon have been
prepared.
Monoatomic Atoms
• All of the noble/inert gases are monatomic.
• Monatomic means one atom.
Magic seven
*note them on your
Periodic Table in your
Reference Table!
• Diatomic molecules= always exist as 2 of the
same atoms bonded to each other
H2
N2
O2
F2
Cl2
Br2
I2
Transition Metals (d-Block)
• They have “typical metallic properties”
• Luster, ductile, malleable, good conductors of heat and
electricity
• Less reactive than Group 1 and 2
• Many are unreactive (for example, palladium, platinum and
gold are found as pure elements in nature)
• These elements begin in Period 4 and include Groups 3-12.
Transition Metals (d-Block)
• As ions, the transition elements form colorful
solutions.
Lanthanides and Actinides (f-Block)
• Lanthanide Series- shiny metals similar in
reactivity to Alkaline Earth Metals
• Actinides Series- all radioactive (mostly
synthetic)
Synthetic Elements
• Synthetic elements do not occur in nature
– Ex: Technetium (Tc): atomic number 43
– Ex: All elements with atomic numbers > 92
(transuranium elements), which are all radioactive
• These elements are produced artificially by
bombardment or as by-products of fission
reactors.
What did you learn today?
What did you learn today?
•
•
•
•
•
•
The placement or location of elements on the Periodic Table gives an
indication of physical and chemical properties of that element. The elements
on the Periodic Table are arranged in order of increasing atomic number.
The number of protons in an atom (atomic number) identifies the element.
The sum of the protons and neutrons in an atom (mass number) identifies an
isotope.
Elements can be classified by their properties and located on the Periodic
Table as metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases.
Elements can be differentiated by their physical properties. Physical properties
of substances, such as density, conductivity, malleability, solubility, and
hardness, differ among elements.
Elements can be differentiated by chemical properties. Chemical properties
describe how an element behaves during a chemical reaction.
For Groups 1, 2, and 13-18 on the Periodic Table, elements within the same
group have the same number of valence electrons (helium is an exception)
and therefore similar chemical properties.
Quiz
• Quiz on Intro to the Periodic Table
Lab
• Periodic Trends lab
Do Now
• What is a trend?
• Have you noticed any trends in history? Have
some trends become more abundant, while
others have become more scarce?
• List some examples.
Periodic Trends
•
•
•
•
•
•
ELECTRONEGATIVITY
ATOMIC RADIUS
IONIC RADIUS
IONIZATION ENERGY
ELECTRON AFFINITY
METALLIC CHARACTER
The trend depends on 3 factors:
1. NUCLEAR CHARGE= atomic number= # of
protons- the higher the nuclear charge the more
the electrons are pulled toward the nucleus.
2. PRINCIPAL ENERGY LEVEL- Principal energy level is
determined by the period an atom is located; the
higher the principal energy level, the higher the
potential energy, the larger the atom.
3. ELECTRON CLOUD SHIELDING EFFECT- inner
electrons shield the outer electrons so they are
less attracted to the nucleus so the atom is larger
Electronegativity
• A measure of the ability of an atom in a chemical
compound to attract electrons.
• FLUORINE is the most electronegative atom on the Periodic
Table
• PERIOD TREND: Electronegativity increases across a period
(the closer to 8 valence electrons the more they will pull on the
electrons)
• GROUP TREND: Electronegativity decreases down a group
(the larger the atomic radius the less ability of that atom to attract
electrons due to electron shielding)
Periodic Table of Electronegativities
http://www.youtube.com/watch?v=Kj3o0XvhVqQ
http://faculty.ucc.edu/chemistry-pankuch/ElectNeg/EN_Bond.dcr
Electronegativity Review
• Element with most Electronegativity:
• Element with least Electronegativity:
Atomic Radius
• Atomic radius = half of the distance between nuclei of
identical atoms bonded together (bond radius).
• This is determined by how strongly an atom’s nucleus
(positive charge) is attracted to its outermost electrons.
• PERIOD TREND: Radius decreases across a period
• Nuclear charge increases across a period so electrons are more
closely attracted to the nucleus
• GROUP TREND: Radius increases down a group
• Addition of principal energy levels, so inner electrons shield
outer electrons and the electrons are farther away from
nucleus, so atom is larger.
Table of Atomic Radii
Review Ions
• Ions are formed when atoms gain or lose
electrons
• Ions have a positive or negative charge
• Ionic Radii:
The loss or gain of electrons by an atom
causes a corresponding change in size.
Ionic Radii
Cations
• Positively charged ions
• Smaller than the
corresponding atom
• Metal tend to form CATIONS
by LOSING an electron.
• Metal cations are smaller
than their corresponding
neutral atoms (they lost
electrons).
Anions
• Negatively charged ions
• Larger than the
corresponding atom
• Nonmetals tend to for
ANIONS by GAINING an
electron.
• Nonmetal anions are larger
than neutral atoms because
they gained electrons.
Here is a simple way to remember a
cation and an anion:
• This is a cat-ion.
• He’s a “plussy” cat!
+
+
• This is Ann Ion.
• She’s unhappy and
negative.
Ionic Radius Review
• The loss or gain of electrons by an atom causing a
corresponding change in size.
Metals
• Tend to ____ their valence electrons.
• What happens to their valence electrons?__________
Before:
After:
• Therefore the atomic radius of metals is ___________ than
their ionic radius.
Ionic Radius Review
• The loss or gain of electrons by an atom causing a
corresponding change in size.
Nonmetals
• Tend to _________ electrons.
• What happens to their radius?__________
Before:
After:
• Therefore the atomic radius of nonmetals is
________________ than their ionic radius.
Table of
Atomic
Radii
and
Ionic
Radii
Ionic Radius
Metals
form
CATIONS
Loss of electron
Get a Positive Charge
Nonmetals
Form
ANIONS
Gain of electron
Get a Negative charge
Radius Decreases
when it becomes an
ion
Radius Increases
when it becomes an
ion
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf
http://faculty.ucc.edu/chemistry-pankuch/IonSizePer_TEST/IonSizePer.dcr
Do Now
• Period 2: Li
Be
C
N
O
F
• Which elements do not want to lose their
electrons? ____________
Ne
Ionization Energy
• Ionization energy = the energy required to remove an
electron from an atom
A + ionization energy  A+ + e• PERIOD TREND: Tends to increase across a period
– Within same principle energy level: due to increased
nuclear charge the electrons are more closely attracted to
nucleus
• GROUP TREND: Tends to decrease down a group
– Addition of principle energy levels, so outer electrons are
farther from the nucleus due to electron shielding, so less
attracted to the nucleus
Another Way to Look at Ionization
Energy
Ionization of Magnesium
• Ionization energy increases for successive electrons taken from
the same atom
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Ionization Energy Review
• Element with the most Ionization Energy:
• Element with the least Ionization Energy:
• Why do noble gases have such high ionization
energy but no electronegativity?
Electron Affinity
• The energy change that occurs when an electron is
acquired by a neutral atom is called the atom’s
electron affinity.
A + e-  A- + energy
A + e- + energy  A- (unstable)
• PERIOD TREND: Tends to increases across a period
– due to increasing nuclear charge (remember: the Halogens
gain electrons most readily)
• GROUP TREND: Tends to decrease down a group
– due to increasing electron shielding from addition of
principle energy levels
REACTIVITY
• METALS- most reactive in lower left corner
• NONMETALS- most reactive in top right
corner, but it does not include noble gases
http://video.google.com/videoplay?docid=-2134266654801392897&q=alkali+me
http://www.youtube.com/watch?v=m55kgyApYrY
Reactivity Review:
• Place the following elements in increasing
order of reactivity: (least reactive first)
– Sodium, potassium, calcium
– Group 1, Group 1, Group 2, so…
– Calcium, sodium, potassium
Metallic Character
• As you go down a group, the metallic
character increases.
• As you go across a period, the metallic
character decreases.
Melting Point
• Melting points are the amount of energy
required to break a bond(s) to change the
solid phase of a substance to a liquid.
• The stronger the bond between the atoms of
an element, the higher the energy
requirement in breaking that bond.
• Metals generally possess a high melting point.
• Most nonmetals possess low melting points.
Boiling Point
• Boiling points are the amount of energy
required to break a bond(s) to change the
liquid phase of a substance to a gas.
• The stronger the bond between the atoms of
an element, the higher the energy
requirement in breaking that bond.
• Metals generally possess a high boiling point.
• Most nonmetals possess low boiling points.
Summation of Periodic Trends
Electronegativity
Electronegativity
http://www.youtube.com/watch?v=h7XWqwgZII0&feature=related
Periodic Trends
Definition
Trend in Period
Electronegativity The measure of
Increases across
the ability of an
a period
atom in a
chemical
compound to
attract electrons
Trend in Group
Decrease down a
group
Atomic Radius
Half the distance Decreases across Increases down a
between nuclei
a period
group
of identical
atoms bonded
together
Ionization
Energy
The energy
required to
remove an
electron from an
atom
Increases across
a period
*Use Table S in your Reference Tables!
Decreases down
a group
Lab
• Reactivity lab
• Videos:
• http://www.youtube.com/watch?v=m55kgyApYrY
• http://www.youtube.com/watch?v=l0U7VDSxGHk
Quiz
• Quiz on Periodic Table Trends
Test
• Test on the Periodic Table