The Periodic Table
Content Review:
Links: The Periodic Table Periodic Trends
Periodicity Practice Problems
History:
Antoine Lavoisier was the first to list the known elements during his time (mid 1700s). However, Dmitri Mendeleev is credited with organizing the periodic table as we know it in 1871, but he was not the only one trying to understand the behavior of the elements. Lothar Meyer was one of many scientists also working on arranging the elements at the same time, but Mendeleev published his results first.
John Newlands was the first to notice a pattern in element behavior when they were arranged in groups of eight in 1863. The chemical properties of certain elements were very similar when arranged in this manner. Newlands called this the Law of Octaves. It is similar to a musical scale, which repeats every eighth tone. However, the law fails for elements beyond calcium.
Mendeleev's periodic table organized the elements by atomic mass, but it caused problems with the placement of some of the known elements. Henry Moseley, using X-ray spectroscopy (around 1913), proposed ordering the elements by atomic number, and the discrete nature of atomic numbers made it easier for the elements to be placed in their correct positions.
Structure:
The periodic table has:
· 118 elements
· 18 groups/families (columns)
· 7 periods (rows)
· 4 blocks (sections) S P D F
· 2 classifications (Representative Elements [A] and Transition Metals [B])
3 types: metals, nonmetals, metalloids
Elements in each group have the same number of valence (outermost) electrons. For example: Group 1 elements have 1 valence electron, and Group 18 (or 8A) elements have 8 valence electrons (a full octet).
Elements in each period have the same number of energy levels or shells. For instance, all the elements in period 3 have three energy levels. All elements in period 6 have six energy levels.
Hydrogen is a very special element, which is why it isn’t connected to the other elements. It is classified as a Group 1 element because of its one valence electron, but that’s the only thing it has in common with those elements.
Group 1 elements are the Alkali Metals. They are extremely reactive because of their one valence electron. They are silvery and soft enough to be cut with a knife.
Group 2 elements are the Alkaline Earth Metals. They are slightly less reactive than their neighbors. They are also silvery, being metals, and are denser than Group 1 elements.
Groups 1-2 and 13-18 are the Representative Elements (aka “A” elements). They have a wide range of physical and chemical properties.
Groups 3 – 12 are the Transition Metals (aka “B” elements). They have a moderate range of reactivity and a wide range of properties. They are shiny and are good conductors, with higher densities and melting points than Groups 1 and 2.
Lanthanides and Actinides: These are the “f-block” transition metals. They share many properties. The lanthanides are shiny and reactive. The actinides are all radioactive and unstable.
Elements 95-118 do not exist in nature; they were created in a laboratory.
Group 13 elements are headed by boron (but I named them the Boronites because they need a name). This group contains one metalloid and four metals. They all have 3 valence electrons.
Group 14 elements are headed by carbon (and I call them the Carbonicles). There is one metal, two metalloids and two metals. This group has 4 valence electrons and varied reactivity.
Group 15 elements are headed by nitrogen. Their family name is Pricogens. This group has varied reactivity and 5 valence electrons.
Group 16 elements, the Chalcogens, are headed by oxygen. They have six valence electrons and is a very reactive group.
Group 17 elements are the Halogens (“salt former”). They are all nonmetals and extremely reactive. They are poor conductors. They have 7 valence electrons.
Group 18 elements are the Noble Gases. They are all inert, which means that they don’t react. They have a full octet of 8 valence electrons, and the shell is filled as a result, which is why these elements are so stable.
PowerPoint
Reading Guide
Student Missions:
Mission 1: What We Already Know.
Mission Objectives. You should be able to...
1. Explain how Mendeleev organized the first periodic table.
2. Describe the ways in which the periodic table is organized.
3. Name the families of the periodic table.
The PT has a long and sordid (I wish) history and is the result of the work of many scientists. Dmitri Mendeleev gets the most credit because he found an organizational system that worked better than others and could make predictions about future elements. The Royal Society of Chemistry provides an in-depth review which you should read outside of class. However, Mendeleev's version of the PT didn't quite work (which element was the one to cause him problems?) because he ordered the elements by increasing atomic mass. We now know, thanks to Henry Moseley, that elements are ordered by increasing atomic number.
Links: The Periodic Table Periodic Trends
Periodicity Practice Problems
History:
Antoine Lavoisier was the first to list the known elements during his time (mid 1700s). However, Dmitri Mendeleev is credited with organizing the periodic table as we know it in 1871, but he was not the only one trying to understand the behavior of the elements. Lothar Meyer was one of many scientists also working on arranging the elements at the same time, but Mendeleev published his results first.
John Newlands was the first to notice a pattern in element behavior when they were arranged in groups of eight in 1863. The chemical properties of certain elements were very similar when arranged in this manner. Newlands called this the Law of Octaves. It is similar to a musical scale, which repeats every eighth tone. However, the law fails for elements beyond calcium.
Mendeleev's periodic table organized the elements by atomic mass, but it caused problems with the placement of some of the known elements. Henry Moseley, using X-ray spectroscopy (around 1913), proposed ordering the elements by atomic number, and the discrete nature of atomic numbers made it easier for the elements to be placed in their correct positions.
Structure:
The periodic table has:
· 118 elements
· 18 groups/families (columns)
· 7 periods (rows)
· 4 blocks (sections) S P D F
· 2 classifications (Representative Elements [A] and Transition Metals [B])
3 types: metals, nonmetals, metalloids
Elements in each group have the same number of valence (outermost) electrons. For example: Group 1 elements have 1 valence electron, and Group 18 (or 8A) elements have 8 valence electrons (a full octet).
Elements in each period have the same number of energy levels or shells. For instance, all the elements in period 3 have three energy levels. All elements in period 6 have six energy levels.
Hydrogen is a very special element, which is why it isn’t connected to the other elements. It is classified as a Group 1 element because of its one valence electron, but that’s the only thing it has in common with those elements.
Group 1 elements are the Alkali Metals. They are extremely reactive because of their one valence electron. They are silvery and soft enough to be cut with a knife.
Group 2 elements are the Alkaline Earth Metals. They are slightly less reactive than their neighbors. They are also silvery, being metals, and are denser than Group 1 elements.
Groups 1-2 and 13-18 are the Representative Elements (aka “A” elements). They have a wide range of physical and chemical properties.
Groups 3 – 12 are the Transition Metals (aka “B” elements). They have a moderate range of reactivity and a wide range of properties. They are shiny and are good conductors, with higher densities and melting points than Groups 1 and 2.
Lanthanides and Actinides: These are the “f-block” transition metals. They share many properties. The lanthanides are shiny and reactive. The actinides are all radioactive and unstable.
Elements 95-118 do not exist in nature; they were created in a laboratory.
Group 13 elements are headed by boron (but I named them the Boronites because they need a name). This group contains one metalloid and four metals. They all have 3 valence electrons.
Group 14 elements are headed by carbon (and I call them the Carbonicles). There is one metal, two metalloids and two metals. This group has 4 valence electrons and varied reactivity.
Group 15 elements are headed by nitrogen. Their family name is Pricogens. This group has varied reactivity and 5 valence electrons.
Group 16 elements, the Chalcogens, are headed by oxygen. They have six valence electrons and is a very reactive group.
Group 17 elements are the Halogens (“salt former”). They are all nonmetals and extremely reactive. They are poor conductors. They have 7 valence electrons.
Group 18 elements are the Noble Gases. They are all inert, which means that they don’t react. They have a full octet of 8 valence electrons, and the shell is filled as a result, which is why these elements are so stable.
PowerPoint
Reading Guide
Student Missions:
Mission 1: What We Already Know.
Mission Objectives. You should be able to...
1. Explain how Mendeleev organized the first periodic table.
2. Describe the ways in which the periodic table is organized.
3. Name the families of the periodic table.
The PT has a long and sordid (I wish) history and is the result of the work of many scientists. Dmitri Mendeleev gets the most credit because he found an organizational system that worked better than others and could make predictions about future elements. The Royal Society of Chemistry provides an in-depth review which you should read outside of class. However, Mendeleev's version of the PT didn't quite work (which element was the one to cause him problems?) because he ordered the elements by increasing atomic mass. We now know, thanks to Henry Moseley, that elements are ordered by increasing atomic number.
Mission 2: It's PERIODIC Because...
Mission Objective. You should be able to...
1. Explain the relationships between atomic number and ionization energy, atomic radius, electronegativity, and ionic radii.
...the organization of elements via atomic number showed properties (with a magical number of 8, of course) that were...wait for it...periodic! Elements in the same columns tend to behave the same way. We now know this is because of their valence electrons.
Because of periodicity, trends in elements can be predicted. We discuss four types of trends: ionization energy, electronegativity, atomic radius and ionic radius.
Mission Objective. You should be able to...
1. Explain the relationships between atomic number and ionization energy, atomic radius, electronegativity, and ionic radii.
...the organization of elements via atomic number showed properties (with a magical number of 8, of course) that were...wait for it...periodic! Elements in the same columns tend to behave the same way. We now know this is because of their valence electrons.
Because of periodicity, trends in elements can be predicted. We discuss four types of trends: ionization energy, electronegativity, atomic radius and ionic radius.
Atomic Radius (AR)
As atoms gain protons, their sizes decrease because of the increased positive pull of the nucleus on the valence electrons. Think in terms of a game of tug of war; with the nucleus on one end and the valence electrons on the other. Adding more protons pulls the valence electrons closer. Atomic number increases from left to right, so atoms decrease in size from left to right. Atomic number increases from top to bottom, so atoms increase in size from top to bottom.
Therefore, atomic radii decrease across a period and increase down a group.
As atoms gain protons, their sizes decrease because of the increased positive pull of the nucleus on the valence electrons. Think in terms of a game of tug of war; with the nucleus on one end and the valence electrons on the other. Adding more protons pulls the valence electrons closer. Atomic number increases from left to right, so atoms decrease in size from left to right. Atomic number increases from top to bottom, so atoms increase in size from top to bottom.
Therefore, atomic radii decrease across a period and increase down a group.
Ionic Radius (IR)
Metals form cations easily, and nonmetals form anions easily. Cations are always smaller than the neutral atoms from which they form, because the stripping of the valence electrons increases the attraction from the nucleus. As a result, the nucleus pulls the remaining electrons closer. Anions are always larger than the neutral atoms from which they form, because adding electrons decreases the nuclear pull.
Therefore, just like atomic radii, ionic radii decrease across a period and increase down a group.
Metals form cations easily, and nonmetals form anions easily. Cations are always smaller than the neutral atoms from which they form, because the stripping of the valence electrons increases the attraction from the nucleus. As a result, the nucleus pulls the remaining electrons closer. Anions are always larger than the neutral atoms from which they form, because adding electrons decreases the nuclear pull.
Therefore, just like atomic radii, ionic radii decrease across a period and increase down a group.
Ionization Energy (IE)
When an atom gains or loses an electron, it becomes an ion. Positively charged ions are called cations and negatively charged ions are called anions. The energy required to overcome the attraction of the nuclear charge and remove an electron is called ionization energy.
Removing one electron results in the formation of a cation with a +1 charge (and the energy required to do so is called the first ionization energy). Removing two electrons (which is more difficult) results in a cation with a +2 charge (second ionization energy). As the ionization energy increases, it becomes more and more difficult to remove electrons. Larger atoms' valence electrons are further away from the nucleus, and therefore have lower IEs.
Similarly, adding electrons increases the negative charge. Adding one electron results in the formation of an anion with a -1 charge. Adding two electrons results in an anion with a -2 charge, and so on and so forth.
Therefore, ionization energy increases across a period and decreases down a group.
When an atom gains or loses an electron, it becomes an ion. Positively charged ions are called cations and negatively charged ions are called anions. The energy required to overcome the attraction of the nuclear charge and remove an electron is called ionization energy.
Removing one electron results in the formation of a cation with a +1 charge (and the energy required to do so is called the first ionization energy). Removing two electrons (which is more difficult) results in a cation with a +2 charge (second ionization energy). As the ionization energy increases, it becomes more and more difficult to remove electrons. Larger atoms' valence electrons are further away from the nucleus, and therefore have lower IEs.
Similarly, adding electrons increases the negative charge. Adding one electron results in the formation of an anion with a -1 charge. Adding two electrons results in an anion with a -2 charge, and so on and so forth.
Therefore, ionization energy increases across a period and decreases down a group.
Electronegativity (EN)
The electronegativity of an element is the tendency of its atoms to attract electrons when they are chemically combined with another element. Noble gases do not have electronegativity values because they do not readily form compounds; they are inert. Electronegativity follows the same pattern as ionization energy and the opposite pattern of atomic radius.
Therefore, electronegativity increases across a period and decreases down a group.
The electronegativity of an element is the tendency of its atoms to attract electrons when they are chemically combined with another element. Noble gases do not have electronegativity values because they do not readily form compounds; they are inert. Electronegativity follows the same pattern as ionization energy and the opposite pattern of atomic radius.
Therefore, electronegativity increases across a period and decreases down a group.
For chemical properties of the elements, go HERE.