Exploring the Periodic Table

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Exploring the Periodic Table

• Modern Chemistry Holt, Rinehart, Winston

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Chapter 5 Section 1history of the periodic
table

• In the late 1800s, scientists had identified over
  60 elements. Certain characteristic physical and
  chemical properties were associated with each
  element. The physical property called atomic mass
  provided chemists with a convenient way to
  organize the elements. At the same time, it was
  recognized that there were certain elements that
  had similar chemical properties. Mendeleev
  arranged the elements in rows according to atomic
  weight and kept elements with similar chemical
  properties in the same columns. Today elements
  are ordered according to atomic number rather
  than atomic mass.

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Learning Targets

• I can explain the roles of Mendeleev and Moseley
  in the development of the periodic table.
• I can describe the modern periodic table.
• I can explain how the periodic law can be used to
  predict the physical and chemical properties of
  elements.
• I can describe how the elements belonging to a
  group of the periodic. table are interrelated in
  terms of atomic number.

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Stanislao Cannizzaro (1826-1910)

• Italian chemist
• Determined a method for accurately measuring the
  relative masses of atoms
• His method allowed chemists to search for a
  relationship between atomic mass and other
  properties of elements

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Dmitri Mendeleev (1834-1907)

• Russian chemist
• Credited as being the creator of the first
  version of the periodic table of elements
• Arranged his periodic table according to atomic
  mass so that elements with similar properties
  were in the same group
• Some elements could not be arranged according to
  atomic mass in order to keep the elements
  arranged according to properties
• Predicted the properties of elements that had not
  yet been discovered using his periodic table

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Mendeleevs Periodic Table

• I began to look about and write down the
  elements with their atomic weights and typical
  properties, analogous elements and like atomic
  weights on separate cards, and this soon
  convinced me that the properties of elements are
  in periodic dependence upon their atomic
  weights. --Mendeleev, Principles of Chemistry,
  1905, Vol. II

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Henry Moseley (1887-1915)

• English chemist
• Worked with Rutherford
• Proved Mendeleevs arrangement of the periodic
  table to be correct only, the periodic table
  was arranged according to atomic number, not
  atomic mass

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The Periodic Law

• States that when elements are arranged in order
  of increasing atomic number, their physical and
  chemical properties show a periodic pattern

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Chapter 5 Section 2electron configuration and
the periodic table

• The modern periodic table has 112 squares, which
  represent a unique element. The distinctive shape
  of the periodic table comes in part from the
  periodic law. Elements in the same column have
  similar properties. These columns are referred to
  as groups or families of elements. The horizontal
  rows of the periodic table are called periods.
  The elements in the periodic table are also
  grouped as metals, nonmetals, and semimetals.
  Metals make up most of the periodic table and are
  located in the center and at the left of the
  table. With the exception of hydrogen, nonmetals
  are on the right side, and semimetals are located
  between the metals and nonmetals. The periodic
  table can also be viewed in terms of orbital
  blocks. These orbital blocks refer to the
  orbitals (s, p, d, and f ) which contain the
  elements incompleted sublevels of electrons.

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Learning Targets

• I can describe the relationship between electrons
  in sublevels and the length of each period of the
  periodic table
• I can locate and name the four blocks of the
  periodic table and explain the reasons for these
  names
• I can discuss the relationship between group
  configurations and group numbers
• I can describe the locations in the periodic
  table and the general properties of the alkali
  metals, the alkaline-earth metals, the halogens,
  the transition metals, the noble gases, the
  actinides, the lanthanides, the metals, the
  nonmetals, the metalloids, and the main group
  elements

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Periodic Law Demonstrated in Groups

• Why do elements in groups have similar physical
  and chemical properties?

• They have the same number of valence electrons in
  their outer energy levels.
• Generally, the configurations of the outermost
  electron shells of elements within the same group
  are the same.

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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
METALS METALLOIDS NONMETALS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
ALKALI METALS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
ALKALINE-EARTH METALS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
HALOGENS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
NOBLE GASES
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
TRANSITION METALS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
INNER TRANSITION (Rare Earth) METALS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
LANTHANIDES
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
ACTINIDES
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
PERIODS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
GROUPS
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In the periodic table below, indicate the
location of the groups, periods, alkali metals,
alkaline earth metals, halogens, noble gases,
lanthanides, actinides, transition metals, inner
transition metals, main group elements, metals,
nonmetals and metalloids.
MAIN GROUP ELEMENTS
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Lets Compare!

• Metals

• Nonmetals

• Metalloids

• Good conductors of heat and electricity
• Malleable
• Ductile
• Luster
• Typically solids at room temperature

• Solids, liquids and gases at room temperature
• Solids are brittle and dull
• Poor conductors of heat and electricity

• Have properties of both metals and nonmetals
• Mostly brittle solids
• Intermediate conductors of electricity- AKA
  semiconductors

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Properties of Alkali Metals

• Extremely reactive
• Readily react with water and air
• Silvery in appearance
• Soft enough to cut with a knife
• Lower densities than other metals
• Lower melting points than other metals

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Properties of Alkaline-Earth Metals

• Harder stronger than alkali metals
• Higher densities melting points than alkali
  metals
• Less reactive than alkali metals

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Properties of Halogens

• Most reactive nonmetals
• React readily with most metals to form salts
• Most electronegative elements

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Properties of Noble Gases

• Least reactive elements because their highest
  occupied energy levels are completely filled with
  an octet of electrons (except He, which only
  requires 2 electrons to be filled).

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Properties of Transition Metals

• High densities
• High melting points
• Good conductors of heat electricity
• High luster
• Less reactive than alkali and alkaline-earth
  metals

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Properties of p Block Metals

• Harder and more dense than the s block metals
• Softer and less dense than the d block metals.

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Properties of Lanthanides

• Soft, silvery metals
• Similar reactivity to alkaline-earth metals

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Properties of Actinides

• All radioactive
• The first 4 have been found naturally on Earth

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Did you know?

• Oxygen, carbon, hydrogen and nitrogen make up 96
  of the human body mass
• Calcium and phosphorous make up 3
• Sodium, potassium, chloride and magnesium make up
  0.7
• Iron, cobalt, copper, zinc, selenium, cyanide and
  fluorine are found in trace amounts

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Chapter 5 Section 3electron configurations and
periodic properties

• Many of the properties of the elements change in
  predictable ways as you move across a period or
  move down a group of the periodic table. The
  predictable changes in these properties are
  called periodic trends. There are periodic trends
  for properties such as atomic radius, ionic size,
  ionization energy, electron affinity, and
  electronegativity. Knowledge of these trends
  helps develop a better understanding of the
  periodic table and of the patterns of behavior of
  the elements.

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Learning Targets

• I can define the term periodic trend.
• I can define atomic radius, ionic radius,
  ionization energy, electron affinity and
  electronegativity.
• I can describe the general trends on the periodic
  table for atomic radius, ionic radius, electron
  affinity, ionization energy and
  electronegativity.
• I can apply the trends on the periodic table to
  answer questions regarding size, electron
  affinity, ionization energy and electronegativity.

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Atomic Radii

• Atomic radius one-half the distance between the
  nuclei of identical atoms that are bonded together

Atomic Radius
Distance between nuclei
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Period Trends

• Decreases across a period

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Why?

• Protons are added to the nucleus moving across a
  period from left to right
• This increases the charge of the nucleus
  (effective nuclear charge Zeff)
• As Zeff increases, the electrons are pulled
  closer to the nucleus

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Period Trends
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Group Trends

• Increase down a group

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Why?

• The addition of shells increases the electrons
  distance from the nucleus and the size of the atom

• Electron-electron repulsion plumps up the atom
• Zeff decreases the further the electrons are from
  the nucleus

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Variations in Atomic Radii
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Atomic Radii Trends
DECREASES
DECREASES
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Ionization Energy

• The energy required to remove one electron from a
  neutral atom of an element creating an ion
• A Energy ? A e-

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Period Trends

• Increase across a period
• Why?

• Zeff increases across the period

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Group Trends

• Decrease down the group
• Why?

• Electron shielding causes a decrease in effective
  nuclear charge
• Electron-electron repulsion forces increase

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Variations in Ionization Energies
Draw the orbital notation for Group 5A and Group
6A.
Can you explain the dips in the chart for these 2
groups?
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Variations in Ionization Energies
If removing an electron will create an empty or ½
filled subshell, ionization energy will decrease.
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Successive Ionization Energies

• Each successive electron removed from an ion
  feels an increasingly stronger effective nuclear
  charge (Zeff) therefore, successive ionization
  energies are larger than 1st ionization energies
• A large jump in ionization energy occurs when
  removing an electron from an ion that assumes a
  noble gas configuration

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Ionization Energy Trends
INCREASES
INCREASES
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Electron Affinity

• The change in energy that a neutral atom
  undergoes when an electron is acquired (the
  ability to attract an e -)
• A e- ? A- energy
• negative energy value (exothermic)
• A e- energy ? A-
• positive energy value (endothermic)

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Period Trends

• Increase across a period
• Why?

• Zeff increases across the period

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Group Trends

• Decrease down the group
• Why?

• Electron shielding causes a decrease in effective
  nuclear charge
• Electron-electron repulsion forces increase

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Variations in Electron Affinities
Why is there such a large decrease in energy for
groups 2A and 5A?
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Electron Affinity Trends
INCREASES
INCREASES
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Ionic Radii

• Cation positively charged ion
• Cations are smaller than their parent atom
  why?
• Anion negatively charged ion
• Anions are bigger than their parent atom why?

Removal of an electron creates an unbalanced
positive charge increasing Zeff and decreasing
the radius of the ion.
Addition of an electron creates an unbalanced
negative charge decreasing Zeff and increasing
the radius of the ion.
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Ionic Radii Trends
DECREASES
DECREASES
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Valence Electrons

• Electrons available to be gained, lost or shared
  in the formation of a chemical compound
• Located in the outer energy level

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Electronegativity

• A measure of the ability of an atom in a chemical
  compound to attract a bonding pair of electrons
• NOTE Electronegativity is a property of atoms
  in compounds and thus differs from ionization
  energy and electron affinity, which are
  properties of isolated atoms

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Trends

• Increase across a period
• Effective nuclear charge increases
• Decrease down a group
• Increase in atomic size and increase in electron
  shielding decreases the effective nuclear charge
• Electronegativity depends upon
• The number of protons in the nucleus
• The distance from the nucleus
• Electron shielding

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Electronegativity Trends
INCREASES
INCREASES