Chapter 6 The Periodic Table

What makes a group of elements?

What makes a group of elements  Some elements are gases, some are liquids, and others are solids.  Some are colored and some are colorless.  Though some elements look dissimilar, they react similarly.  Samples of the compounds NaCl, NaBr, and NaI show similar physical characteristics.

Periodic Pattern  In 1865, the English chemist John Newlands arranged the first 16 elements known at the time in order of increasing atomic mass.  When Newlands placed the elements in two rows, he observed that the elements in each column had similar chemical and physical properties.

Li Be B C N O F Na Mg Al Si P S Cl Because the chemical and physical properties repeated with the eight element, Newlands called this pattern the law of octaves.

 In 1870, the Russian chemist Dmitri Mendeleev made use of this scheme and other information to produce the first orderly arrangement of all 63 elements known at the time.  Mendeleev wrote the symbol for each element on a card along with its relative atomic mass. Arranged elements in order of increasing atomic mass.  He is considered the “father of the periodic table”

 Periodic law- states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.  The most common periodic table is based on the periodic law.  The periodic table contains a wealth of information about individual elements.

The modern periodic table  The rows of the periodic table are called periods.  Elements in a period have similar electron configurations.  For example, the first period has two elements, hydrogen and helium, and the electrons of these elements occupy the 1s orbital.

 The columns of the periodic table are called groups.  The elements within a group have properties in common but with some gradation.  For example, all Group I elements are solid at room temp., and are good conductors of electricity.

METALS  It is convenient to divide the periodic table into two distinct regions, metals and nonmetals.  The metals include all members of Groups 1 through 12 as well as some elements of Group 13 through 16.

 All metals are good conductors of electricity, and their conductivity increases as temperature decreases.  Except for Mercury, all metals are solid at room temperature.  Elements in Group 3 through 12 including the two long rows below the main table, are called transition metals.

 As you move from left to right across the transition metals, you will see that the electrons are usually added to d orbitals.  For this reason the transition metals are sometimes referred to as the d- block elements.

Metals that are good conductors of electricity are also good conductors of heat.  In general, poor electrical conductors are poor heat conductors.  Think about it.  This means that a mechanism by which electricity is conducted must also be closely connected with the mechanism for conduction of heat.

 You don’t have to look that far for the agent that causes conductivity in metals.  Ever since the discovery of electrons it has been known that electrons are responsible for the conduction of electricity by metals.

 They are responsible for the conduction of heat as well.  However, metals are the only elemental substances that are good conductors of electricity.  This must mean that at least some of the electrons in metals must be in a different configuration than in nonmetals.  Those electrons are free to move through the metal in all directions.

 Metals can have extremely high melting points, and some can have very low melting points.  Metals can be extremely reactive while others do not react at all.  Metals can be strong and durable.

 Metals can also be ductile (wire) and malleable (metal sheets).

NONMETALS  The second region of the periodic table contains the nonmetals.  The nonmetals include all of Groups 17 and 18 as well as some members of Groups 14 through 16.  The characteristics shared by all nonmetals is that they are poor conductors of electricity.

 Nonmetals may be gases, liquids, or solids at room temperature.  Along the stair-step line separating metals from nonmetals are the elements known as semiconductors, or metalloids.  Metalloids are solids at room temp.

Main-group elements  Groups 1, 2, and 13 through 18 are referred to as the main-group elements.  The electron configuration of elements within each group is quite consistent.  For example, all of the elements in Group 14 have four electrons in their outermost shell.

 The main-group elements include gases, liquids, solids, metals, and nonmetals.  The main group elements silicon and oxygen account for four of every five atoms found on or near the Earth’s surface.

 Four groups within the main-group elements have special names.  These are alkali metals (Group 1), the alkaline-earth metals (Group 2), the halogens (Group 17), and the noble gases (Group 18).

Group 1 (Alkali Metals)  Alkali metals- react with water to produce alkaline solutions and because they have metallic properties.  The term alkali dates back to ancient times, when people discovered that wood ashes mixed with water produces a slippery solution that can remove grease.

 The alkali metals are so soft that they can be cut with a knife. The freshly cut surface of an alkali metal is shiny, but it dulls quickly as the metal reacts with oxygen and water in the air.  All of the alkali metals are excellent conductors of electricity.

Group 2 (Alkaline-Earth Metals)  The elements of Group 2 are called the alkaline-earth metals.  Compared with the alkali metals, the alkaline-earth metals are harder, denser, stronger, and have higher melting points.  The best known alkaline-earth metal is calcium. Calcium compounds such as those in limestone and marble, are common in the Earth’s crust.

Group 17 (Halogens)  The elements of Group 17 are the halogens.  Fluorine, chlorine, bromine, iodine, and astatine.  The halogens combine with most metals to produce the compounds known as salts.

 The word halogen is derived from Greek and means “salt former”  In common table salt, sodium chloride, the halogen chlorine has reacted with the alkali metal sodium to form sodium and chloride ions.

Group 18 (Noble Gases)  The Group 18 elements, are called noble gases.  Noble gas atoms are characterized by an octet of electrons in the outermost energy level.

Hydrogen  Hydrogen is in a class by itself  Hydrogen is the most common element in the universe. It behaves unlike other elements because it has just one proton and one electron. This distinguishes hydrogen from all of the other elements.

6-2, 6-3 What trends are found in the period table?

 You have read that elements are arranged in the periodic table in order of increasing atomic number.  The elements are further organized into groups and periods.  The arrangement of the periodic table also reveals trends in the chemical and physical properties of the elements.

 What is a trend?  A trend is a predictable change in a particular direction.  For example, as you move down group 1, reactivity increases for each element.

Periodic Trends in atomic radii  The exact size of the atom is difficult to determine.  Bond radius-half the distance from center to center in two like atoms bonded together.

 The Vand der Waals radius is seldom used to state the size of atoms, but data are only available for only a few main group elements.  Van der Waals radius- half the distance between the nuclei in adjacent non-bonded molecules.

 We will be using bond radius to determine the size of atoms because there is more information available.  Measuring bond radius is a useful way to compare sizes of atoms.

Electron Shielding  The electrons in the inner energy levels are between the nucleus and the outermost valence electrons  This shields the valence electrons from the full charge of the nucleus.  This phenomenon is called electron shielding.  Because the valence electrons are not subject to the full charge of the nucleus, they are not held as close to the nucleus.

Atomic radius increases as you move down a group.  There is a trend toward larger radii as you proceed down a group.  This is caused by the addition of another main energy level as you move from one period to the next.

Atomic radius decreases as you move across a period.  From left to right across a period, each atom has one more proton and one more electron that the element before it.

 The additional electrons are going into the same energy level.  Electrons in an outer energy level do not screen the other electrons in that energy level very effectively.  Meanwhile the nuclear charge is increasing as protons are added and the electrons are pulled closer to the nucleus, reducing the size of the atom.

In other words…  There is a gradual decrease in the atomic radii across the second period from Li to Ne  The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus  Increased pull results in a decrease in atomic radii

Ionization energy, Electron affinity, and Electronegativity.  Recall that atoms are electrically neutral.  But if you add enough energy, the atom may lose an electron to become a positive ion.  Imagine that you can reach into an atom and remove on of its valence electrons, creating an ion.

 The energy you used to remove that electron called ionization energy.  Ionization energy- the amount of energy needed to remove an outer electron from a specific atom or ion in its ground state.

Ionization energy increases across a period and decreases down a group  Group 1 elements have the lowest first ionization energy They lose electrons most easily  Major reason for the high reactivity of the Group 1  The Group 18 elements, the noble gases, have the highest ionization energies. They do not lose their electrons easily; the low reactivity is partly due to the difficulty to removed an electron

 In general, ionization energies of the main-group elements increase across each period  Increase is due to increase in nuclear charge of the protons  Higher charge attracts electrons in the same energy level  In general, nonmetals have higher ionization energies than metals do

 As you move down a group, the number of energy levels between the nucleus and the valence electrons increases and the outermost electrons are farther from the nucleus.  The nuclear charge stays the same so the electrons are held less tightly to the nucleus, and less energy is required to remove one of them.

Periodic Trends in Electron Affinity  The ability of an atom to attract and hold an electron is called electron affinity.  You may wonder why a neutral atom would attract electrons in the first place.  The answer is that electrons in the orbitals generally do not shield the nuclear charge to a full 100%.  An approaching electron may experience the a net pull because the nuclear charge is greater than 0.

 Across a period, shielding remains the same, but nuclear charge increases.  Therefore, the atoms attraction for extra electrons increases.

 Going down a group, both shielding and nuclear charge increase.  However, the shielding effect offsets the increase in nuclear charge.  Therefore, the atom’s attraction for extra electrons decreases.

Electronegativity  Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons  The most electronegative element is Fluorine  The values are assigned 0-4, Fluorine is 4

Period Trends  Electronegativities tend to increase across each period, although there are exceptions  Electronegativities tend to either decrease down a group or remain about the same  Noble gases – some do not form compounds so they cannot be assigned electronegativities

Positive Ions  A positive ion is known as a cation  The formation of a cation by loss of one or more electrons always leads to a decrease in atomic radius because of the removal of the highest-energy-level electrons results in a smaller electron cloud  The remaining electrons are drawn closer to the nucleus

Negative Ions  A negative ion is known as an anion.  The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius  Electron cloud spreads out

Period Trends cont..  The metals on the left form cations  Nonmetals on the right tend to form anions  Cationic radii decreases across a period The electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level  Anionic radii decrease across a period

Group Trends  The outer electrons in both cations and anions are higher in energy levels as you read going down a group  Atomic radii gradually increases going down a group  Ionic radii gradually increases going down a group