Transcript Periodic Table - MunterChemistry

Periodic Table
And the Periodic Law
Dmitri Mendeleev
• Russian chemist
• Created a table by arranging elements
according to atomic masses
• Noticed that chemical properties of the
elements followed a repeating pattern
Periodic Law
– Henry Moseley who worked with Ernest
Rutherford was the scientist who improved
on Mendeleev’s periodic table by ordering
the elements by the number of protons –
the atomic number.
• The physical and chemical properties
of the elements are periodic functions
of their atomic numbers.
Mendeleev left spaces in his
periodic table and predicted the
existence of 3 elements and their
1.
2.
3.
4.
Atomic numbers
Colors
Properties
Radioactivity
Arrangement of the table
• The periodic table is an arrangement of
the elements in order of their atomic
numbers so that elements with similar
properties fall in the same column or
group.
Groups of the Periodic Table
• 1st column (gold)
alkali metals
• 2nd column (purple)
alkaline earth metals
• 3rd-12th Transition
Metals
• 17th column –
halogens
• 18th – Noble Gases
Electron Configuration and the
Periodic Table
• The periodic table can be divided into
blocks which indicate which orbitals are
filling.
s Block – Groups 1and 2
• All are chemically reactive metals
• Metals in group 1 (alkali metals) are more
reactive than metals in group 2 (alkaline
earth metals)
• Neither are found in nature as elements
since they react readily with water and
nonmetals.
Which is more reactive,
magnesium or sodium?
1. Magnesium (Mg)
2. Sodium (Na)
Name the element [Ar] 4s2
1.
2.
3.
4.
Sodium
Potassium
Calcium
Strontium
Alkali Metals Video
Hydrogen and Helium
• Hydrogen is not an alkali metal.
• Hydrogen is a unique element.
• Helium is placed with the noble gases
• Helium has a full outer shell of electrons.
• Helium is a colorless gas, not a reactive
metal.
d Block – Groups 3-12
• Some of these elements don’t follow the
diagonal rule exactly in electron
distribution
• Metals with typical metallic properties.
– Good conductors of electricity
– High luster
• Called the transition elements.
• Less reactive than s block metals
Name the element [Ar]4s23d5
1.
2.
3.
4.
Bromine
Iron
Magnesium
Manganese
p block – groups 13 – 18
except helium
• p block and s block together are called the
main group elements
• Properties of this group vary
– On the right side – nonmetals
– On the left – metals
– In between – metalloids (boron, silicon,
germanium, arsenic, antimony, tellurium)
Metals, Nonmetals, Metalloids
• Metal- element that is shiny and malleable,
and conducts heat and electricity
• Nonmetal – Conducts heat and electricity
poorly and is brittle
• Metalloid – Element that has the
properties of both metals and nonmetals,
sometimes called a semiconductor
An element which is shiny and
brittle is likely a
1.
2.
3.
4.
Metal
Nonmetal
Metalloid
s-block element
p-block
• Group 17 – Halogens (fluorine, chlorine,
bromine, iodine, astatine)
– Most reactive nonmetals
• Group 18 – Noble Gases (helium, neon,
argon, krypton, xenon, radon)
– Nonreactive and stable
f-block: Lanthanides and Actinides
• Lanthanides are shiny metals, with
reactivity similar to the alkaline earth
metals
• Actinides are all radioactive, many are
man-made
Name the group
1. Alkali metals
2. Alkaline earth
metals
3. Halogens
4. Noble Gases
Name the group
1. Alkali metals
2. Alkaline earth
metals
3. Halogens
4. Noble gases
Name the section of the table
1.
2.
3.
4.
S block
P block
D block
F block
What types of elements make up
the p block?
1.
2.
3.
4.
Metals
Nonmetals
Metalloids
All of the above
Periodic Trends
• The arrangement of the periodic table
shows directional trends for various
properties of the atoms of each element.
Nuclear Charge and Atomic Radius
• Atoms decrease in size from left to right on
the periodic table because of the increase
in nuclear charge.
IN A ROW IN THE PERIODIC TABLE, AS
THE ATOMIC NUMBER INCREASES, THE
ATOMIC RADIUS
1. Decreases
2. Remains constant
3. Increases
4. Becomes
immeasurable
Within a group of elements,as the
atomic number increases, the
atomic radius
1.
2.
3.
4.
Increases
Remains constant
Decreases
Varies
unpredictably
Ionization Energy
• The energy required
to remove an electron
from a neutral atom of
an element.
– A large ionization
energy shows that the
electrons of an atom
are bound more tightly
to the nucleus and it is
more difficult to
remove the electron.
Which group of elements has the
highest ionization energies?
1.
2.
3.
4.
Alkali metals
Halogens
Noble gases
Alkaline earth
metals
The energy required to remove an
electron from an atom is called the
1.
2.
3.
4.
Electron affinity
Electron energy
Electronegativity
Ionization energy
Ionization Energy (IE) Trends
• IE generally increases across each period.
– As the nuclear charge increases, the
electrons are held more tightly
• IE decreases down a group
– As the electrons reside farther from the
nucleus, they can be removed more easily.
Electron Affinity
• The energy change that occurs when an
electron is acquired by a neutral atom is called
the atom’s electron affinity.
• Halogens have high electron affinities because
acquiring an electron will give them a full outer
shell which increases the stability of the atom.
• Half filled orbitals also give increased stability, so
that the electron affinity of carbon is greater than
the electron affinity of nitrogen.
Electron Affinity Trend on the
Periodic Table
• Generally increases across a period.
• Generally decreases down a group.
When an electron is added to a
neutral atom, a certain amount of
energy is
1. Always absorbed
2. Always released
3. Either released or
absorbed
4. Transferred to the
more
electronegative
element
Which represents a neutral atom
acquiring an electron in a process
where energy is released?
1.
2.
3.
4.
A + e- + energy AA + e- A- + energy
A + e- A- - energy
A- + energy A + e-
Ionic Radii
• Cation – positive ion
– Formation of a cation decreases the atomic
radius
• Anion – negative ion
– Formation of an anion increases the atomic
radius
As the atomic number of the metals
of Group 1 increases, the ionic
radius25% 25% 25% 25%
1. Increases
2. Decreases
3. Remains the
same
4. Cannot be
determined
0 of 30
1
2
3
4
Valence Electrons
• The electrons available to be lost, gained,
or shared in the formation of chemical
compounds are referred to as the valence
electrons.
• Valence electrons are those filling s and p
orbitals.
Valence electrons are those s and
p electrons
1. Closest to the
nucleus
2. In the lowest
energy level
3. In the highest
energy level
4. Combined with
protons
The number of valence electrons in
Group 17 elements is
1.
2.
3.
4.
7
8
17
Equal to the
period number
The electrons available to be lost,
gained or shared when atoms form
compounds are called
1.
2.
3.
4.
Ions
Valence electrons
d electrons
Electron clouds
Electronegativity
• A measure of the ability of an atom in a
chemical compound to attract electrons
from another atom in the compound.
Fluorine is
assigned a value
of 4 and all other
elements have
values relative to
it.
The element that has the greatest
electronegativity is
1.
2.
3.
4.
Oxygen
Sodium
Chlorine
Fluorine