The Periodic Table And the Periodic Law. wSKFBpo wSKFBpo .

The Periodic Table

And the Periodic Law

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Chapter 6: The Periodic Table and Periodic Law

• Section 6.1 Development of the Modern Periodic Table

Antoine Lavoisier (1743-1794)• Composed a list in 1790 of the 23 elements known at the time.

• The advent of electricity (used to break down compounds into their component elements- electrolysis) and the development of the spectrometer, used to identify elements lead to the discovery of many new elements (70 total by 1870)

Nuclear Magnetic Resonance

(NMR) spectrometer

John Newlands

• Noticed that when elements were arranged by increasing atomic mass their properties repeated every eight element.

John Newlands

• This type of pattern is called periodic because it repeats in a specific manner.

• Newlands called this relationship the Law of Octaves.

Dmitri Ivanovich Mendeleev(1834-1907)

• Organized elements increasing atomic mass

• Noticed that there was a repetition or periodic pattern in their properties.

Dmitri Mendeleev

• Elements with similar properties were arranged in columns- this was the first periodic table

Dmitri Mendeleev

• Mendelev predicted the existence and properties of undiscovered elements and left spaces where they should go.

Henry Moseley

•Arranging the elements by mass resulted in several elements being placed in groups of elements with different properties.

Henry Moseley(1887-1915)

•Discovered that atoms of each element contained a unique number of protons in their nuclei, the number of protons being equal to the atom’s atomic number.

Henry Moseley

• By arranging elements by increasing atomic number, the problems with the order of the elements in the periodic table were solved.

The Periodic Law

When elements are arranged by increasing atomic number, there is a periodic repetition of their chemical and physical properties.

The Modern Periodic Table

• Consists of boxes that contain the elements name, symbol, atomic number and atomic mass.

The Modern Periodic TableGroups

• The boxes are arranged in order of increasing atomic number in a series of columns called groups or families.

The Modern Periodic TableGroups

• Each group is numbered 1 through 18.

The Modern Periodic TableGroups

• Each group is numbered 1 through 8, followed by the letter A or B.

Periods• The boxes are arranged in order of

increasing atomic number in a series of rows called periods.

• Beginning with hydrogen in period 1, there are 7 periods.

Main Group Elements• Groups 1,2,and 13-18 are referred to as

the main group (or representative) elements because they possess a wide range of chemical and physical properties.

Transition elements

• Transition elements

• Groups 3-12

Classifying the Elements

• Metals

• All elements to the left side (except H) of the stair-step line from B to At

Metals

• good conductor of heat and electricity

• shiny luster

• solid at room temperature (except Hg)

• Malleable and ductile

• high boiling and melting points

• High densities

Metals (continued)

• large atomic radius

• low ionization energy

• low electronegativity

Alkali Metals

• The group 1 elements (except for H)

• Most reactive group of metals !!!

Alkaline Earth Metals

• The group 2 elements

• Second most reactive group of metals 

 Transition Elements

• These are the group B elements (all are metals)

Transition Metals

• Transition Metals

• Also known as the d-block elements

Inner transition Metals• Located along the bottom of the

periodic table

• Known as the lanthanide- actinide series

• Lanthanides are also called rare earth elements

• Many are radioactive

• Also know as the f- block elements

Non- Metals• Occupy the upper right side of the

periodic table• (right of the stair-step line from B to At)

Non-Metals (Characteristics)

• Generally gasses at room temperature

• Poor conductors of heat and electricity

• Brittle dull-looking when solids

• lower boiling and melting points than metals (except carbon)

• usually have lower densities than metals

Non-Metals ( Other Characteristics)

• high electronegativity

• Small atomic radius

• higher ionization energy than metals

Halogens

• Group 17 elements

• Most highly reactive group of the non-metals

• Diatomic (Mr. BrINClHOF)

Nobel gases

• Group 18 elements

• Extreamly unreactive

• also called Inert gases

Metalloids• Elements border the stair-step line from

B to At

• Physical and chemical properties of both metals and non-metals

Section review 6.1

• Laviosier,- Composed a list in 1790 of the 23 known elements.Newlands: Noticed that when elements were arranged by increasing atomic

mass their properties repeated every eight element Law of Octaves.

Mendeleev : Organized elements increasing atomic mass Elements with similar properties were arranged in columns- this was the

first periodic tablepredicted the existence and properties of undiscovered elements and left

spaces where they should go

Moseley:By arranging elements by increasing atomic number, the problems with the order of the elements in the periodic table were

solved.

• .

Classification of the Elements

By noting an atoms position on the periodic table, you can determine an its electron configuration and number of valence electrons

Valence Electrons

electrons in an atoms highest principle energy level

•elements have similar chemical properties because they have the same number of valence electrons

Valence Electrons and Period

The energy level of an elements valence electrons indicates the period on the periodic table in which it is found

•Ex. Li - found in period 2; valence electrons found in second energy level

•Ex. Gallium (Ga) found in period 4; valence electrons found in fourth energy level

Valence electrons and group number

• A main group (representative) element and the number of valence electrons it contains are related.

• Group 1 has one valence electron, Group 2 has 2 valence electrons and so on. ( Exception He is in group 18, but only has 2 valence electrons

• Lewis electron dot structures illustrate this connection.

The s- p- d- and f- block elements

• The periodic table is divided into sections, or blocks representing the atoms energy sub levels being filled with valence electrons.

s- block elements

Consist of groups 1 and 2, H and He

s- block elements

• Valence electrons occupy only s orbitals

• group 1 elements have partially filled s orbitals s1

• group 2 elements have partially filled s orbitals s2

s- block elements

• Group 1 elements have partially filled s orbitals s1

• Group 2 elements have filled s orbitals s2

p- block elements

• After the s sublevel is full, the valence electrons next occupy the p sublevel with its three p orbitals.

p- block elements

• The p-block is comprised of groups 13 through 18, with filled or partially filled p orbitals

p- block elements

• There are no p block elements in the period 1, the first p-block element B is in period 2

p- block elements• Together the

s- and p- blocks compromise the main group elements.

p- block elements

• Group 18 are the Nobel gases that are unreactive with any other elements.

p- block elements

• Nobel gases have their highest PEL completely full, this is a very stable electron configuration.

d- block elements

• Largest of the blocks

• Contains the transition metals

d- block elements

• Characterized by a filled outermost s orbital and filled or partially filled d orbitals

• d- block spans

10 groups on

the periodic table

f- block elements

• Contains the inner transition metals (lanthanide-actinide series)

f- block elements

• Have filled or partially filled outermost s orbital, and filled or partially filled 4f and 5f orbitals

F- block

• Electrons do not fill their orbitals in a predictable manner.

F- block http://freezeray.com/flashFiles/discoveryDates.htm

• There are 7 seven f orbitals that hold a maximum of 14 electrons, so the f-block spans 14 columns of the periodic table.

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• Period 1: s-block elements only

• Periods 2 and 3: s- and p- block elements

• Periods 4 and 5: s-, p-, and d- block elements

• Periods 6 and 7: s-, p-, d- and f- block elements

Section 6.3 Periodic Trends

• As you move across a period or down a group elements tend to change in a predictable way, this is known as a periodic trend.

Atomic Radius

• Defined as half the distance between adjacent nuclei in a crystal of the element

• For diatomic atoms AR is defined as half the distance between nuclei of identical atoms that are chemically bonded together.

Trends within periods

• In general there is a decrease in atomic radii as you move left-to-right across a period

Trends within periods

• In general there is a decrease in atomic radii as you move left-to-right across a period

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• The result is that the increased nuclear charge pulls the outermost electrons closer to the nucleus.

Trends within groups

• Atomic radii generally increase as you move down a group.

• Electrons are added to higher PELs, so the outermost orbitals increase the size of the atom.

• The valence electrons are also farther from the nucleus

• Largest: Na

• Smallest: S

•  Largest: Xe

• Smallest: He

• No, with only this information, you will be unable to determine the specific groups and periods That the elements are in; so you can’t apply the periodic in atomic size to determine which element has the larger atomic radius.

• No, with only this information, you will be unable to determine the soecific groups and periods That the elements are in; so you can’t apply the periodic in atomic size to determine which element has the larger atomic radius.

Ionic Radius 

•When atoms lose electrons they form positively charged ions, they always become smaller.  

Ionic Radius 

•This is due to the loss of electrons that are always valence electrons, thus the outer orbital is empty resulting in a smaller radius.  

Ionic Radius• Atoms can gain or lose one or more

electrons to form ions and gain a net charge.  

Ionic Radius• Atoms can gain or lose one or more

electrons to form ions and gain a net charge. An ion is an atom or bonded group of atoms that has a positive or negative charge. 

Ionic Radius

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•When atoms gain electrons, they have a negative charge resulting in a larger radius.

Trends within periodsIonic Radius

As you move left to right across the period the ionic radius decreases, then when you get to group 5A or 6A, the size of the larger negative ions also decreases.

• Trends within groups

• As you move down a group, an ion’s outer electrons are in a higher PEL resulting in a gradual increase in size.

Ionic radius

• Atoms can gain or lose one or more electrons to form ions and gain a net charge. An ion is an atom or bonded group of atoms that has a positive or negative charge.

Ionic radius

• When atoms lose electrons they form positively charged ions, they always become smaller. This is due to the loss of electrons that are always valence electrons, Thus the outer orbital is empty resulting in a smaller radius.

Ionic radius

 When atoms gain electrons, they have a negative charge resulting in a smaller radius.

Ionic radius

• Trends within periods• As you move left to right

across the period the ionic radius decreases, then when you get to group 5A or 6A, the size of the larger negative ions also decreases.

Ionic radius

Trends within groups

•As you move down a group, an ion’s outer electrons are in a higher PEL resulting in a gradual increase in size.

Ionization energy

• To form a positive ion, an electron must be removed from a neutral atom.

• Ionization energy is defined as the energy required to remove an electron (from the gaseous state).

• The energy required to remove the first electron is called the first ionization energy (energy required to remove the second electron, is called the second ionization energy)

Ionization energy

• A high ionization energy value indicates that an atom has a strong hold on electrons.

• IE is a measure of how strongly an atom’s nucleus will hold onto valence electrons.

• Atoms with large IE are less likely to form positive ions than atoms with low IE.

Ionization energy

• Trends within periods

• First ionization energies generally increase as you move l to r across a period.

Ionization energy

Trends within groups

•First ionization energies generally decrease as you move down a group.

•With the valence electrons farther away from the nucleus, less energy is required to remove them.

Octet Rule

• Atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electron, (noble gas configuration).

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Octet Rule

• This is a very stable electron configuration. Elements to the right of the periodic table tend to gain electrons (form negative ions)

• Elements to the left of the periodic table tend to lose electrons (form positive ions to attain this configuration.

Electronegativity

• The relative ability of an atom to attract elements in a chemical bond.

• Expressed in terms of a numerical value of 3.98 or less.

• Units called Paulings (see Linus Pauling 1901-1994)

Electronegativity

• The relative ability of an atom to attract elements in a chemical bond.

Electronegativity

• Noble gases have been left out

• F is the most electronegative element (3.98); Cs and Fr are the least electronegative (.79 and .7)

• In a chemical bond the atoms with the greater electronegativity more strongly attract the bond’s electrons

Electronegativity

Trends within groups and periods

•Electronegativity generally decreases as you move down a group and increases as you move left to right across a period.

Electronegativity

Trends within groups and periods

•Elements with the higher electronegativity are at the upper right and elements with the lower electronegativity are at the bottom left of the periodic table.

AR increase down a group as the electrons are added to higher energy levels and inner core electrons shields the valence electrons from the increased nuclear charge. AR decrease across a period as increased nuclear charge coupled with unchanging shielding by inner core electrons pulls the valence electrons (being added to the same energy level) closer to the nucleus.

•Largest: antimony (Sb) Smallest: (nitrogen (N)

a. F c. Br

b. Br d. F

•Lithium’s second removed electron is an inner core electron, not a valence electron. Carbon’s fourth removed electron is still a valence electron.