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Unit 5: Periodicity History of the PT Dimitri Mendeleev (1836-1907) – Russian chemistry professor In writing a textbook of general chemistry, he compiled vast quantities of data on the elements. He realized as he examined the data that the chemical properties of the elements repeated themselves when the elements were placed in order of atomic mass. Based on his realization, Mendeleev arranged the elements in order of increasing atomic mass, and in such a way that elements with similar chemical properties fell in the same column. He published a primitive version of today’s periodic table in 1869. It contained the 62 elements that had been discovered at that time. Genius of Mendeleev’s Work Left spaces for elements not yet discovered. He predicted that some still-unknown elements must exist to fit in the holes. Dmitri Mendeleev (1869) http://www.chem.msu.su/eng/misc/mendeleev/welcome.html Inductive vs. Deductive Reasoning Inductive reasoning is the use of detailed facts to form a general principle or model (going from specific to general). How did Mendeleev use inductive reasoning? Deductive reasoning is the use of a general principle or model to draw specific inferences (going from general to specific). How did Mendeleev use deductive reasoning? Mendeleev’s Periodic Table Contained 1 inconsistency. – He placed the elements in order of atomic mass Forced to break pattern a couple times to preserve the patterns he had discovered. Henry Moseley Shortly after Rutherford’s discovery of the proton in 1911, Henry Moseley (1887-1915) did experiments with xrays to determine the number of protons in various elements. When Moseley arranged the elements of Mendeleev’s periodic table according to increasing atomic number and not atomic mass, the inconsistencies associated with Mendeleev's table were eliminated. The modern periodic table is based on Moseley's arrangement by atomic number. At age 28, Moseley was killed in action during World War I. As a direct result, Britain adopted the policy of exempting scientists from fighting in wars. Periodic Law From Mendeleev’s and Moseley’s work comes the Periodic Law: The properties of the elements are periodic functions of their atomic numbers. What this means is that if we arrange the elements in order of increasing atomic number, we will periodically encounter elements that have similar chemical and physical properties. These elements appear in the same vertical column (group). Elements Science has come along way since Aristotle’s theory of Air, Water, Fire, and Earth. Scientists have identified 90 naturally occurring elements, and created about 28 others. Periodic Table The periodic table organizes the elements in a particular way. A great deal of information about an element can be gathered from its position in the period table. For example, you can predict with reasonably good accuracy the physical and chemical properties of the element. You can also predict what other elements a particular element will react with chemically. Understanding the organization and plan of the periodic table will help you obtain basic information about each of the 118 known elements. Key to the Periodic Table Elements are organized on the table according to their atomic number, usually found near the top of the square. – The atomic number refers to how many protons an atom of that element has. – For instance, hydrogen has 1 proton, so it’s atomic number is 1. – The atomic number is unique to that element. No two elements have the same atomic number. What’s in a square? Different periodic tables can include various bits of information, but usually: – atomic number – symbol – atomic mass – number of valence electrons – state of matter at room temperature. Periodic Table Expanded View The way the periodic table usually seen is a compress view, placing the Lanthanides and actinides at the bottom of the stable. The Periodic Table can be arrange by sublevels. The s-block is Group I and & 2, the p-block is Group 13 - 18. The d-block is the transition metals, and the f-block are the Lanthanides and Actinide metals Valence Electrons The number of valence electrons an atom has may also appear in a square. Valence electrons are the electrons in the outer energy level of an atom. These are the electrons that are transferred or shared when atoms bond together. Nonmetals Metals Properties of Metals Metals are good conductors of heat and electricity. Metals are shiny. Metals are ductile (can be stretched into thin wires). Metals are malleable (can be pounded into thin sheets). A chemical property of metal is its reaction with water which results in corrosion. Properties of Non-Metals Non-metals are poor conductors of heat and electricity. Non-metals are not ductile or malleable. Solid non-metals are brittle and break easily. They are dull. Many non-metals are gases. Sulfur Properties of Metalloids Metalloids (metal-like) have properties of both metals and non-metals. They are solids that can be shiny or dull. They conduct heat and electricity better than nonmetals but not as well as metals. They are ductile and malleable. Silicon Columns of elements are called groups or families. Elements in each family have similar but not identical properties. For example, lithium (Li), sodium (Na), potassium (K), and other members of family 1 are all soft, white, shiny metals. All elements in a family have the same number of valence electrons. Each horizontal row of elements is called a period. The elements in a period are not alike in properties. In fact, the properties change greatly across even given row. The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas. Hydrogen The hydrogen square sits atop Family 1, but it is not a member of that family. Hydrogen is in a class of its own. It’s a gas at room temperature. It has one proton and one electron in its one and only energy level. Hydrogen only needs 2 electrons to fill up its valence shell. Alkali Metals The alkali family is found in the first column of the periodic table. Atoms of the alkali metals have a single electron in their outermost level, in other words, 1 valence electron. They are shiny, have the consistency of clay, and are easily cut with a knife. Alkali Metals They are the most reactive metals. They react violently with water. Alkali metals are never found as free elements in nature. They are always bonded with another element. Alkaline Earth Metals They are never found uncombined in nature. They have two valence electrons. Alkaline earth metals include magnesium and calcium, among others. Transition Metals Transition Elements include those elements in the B families. These are the metals you are probably most familiar: copper, tin, zinc, iron, nickel, gold, and silver. They are good conductors of heat and electricity. Transition Metals The compounds of transition metals are usually brightly colored and are often used to color paints. Transition elements have 1 or 2 valence electrons, which they lose when they form bonds with other atoms. Some transition elements can lose electrons in their next-tooutermost level. Transition Elements Transition elements have properties similar to one another and to other metals, but their properties do not fit in with those of any other family. Many transition metals combine chemically with oxygen to form compounds called oxides. Boron Family The Boron Family is named after the first element in the family. Atoms in this family have 3 valence electrons. This family includes a metalloid (boron), and the rest are metals. This family includes the most abundant metal in the earth’s crust (aluminum). Carbon Family Atoms of this family have 4 valence electrons. This family includes a nonmetal (carbon), metalloids, and metals. The element carbon is called the “basis of life.” There is an entire branch of chemistry devoted to carbon compounds called organic chemistry. Nitrogen Family The nitrogen family is named after the element that makes up 78% of our atmosphere. This family includes nonmetals, metalloids, and metals. Atoms in the nitrogen family have 5 valence electrons. They tend to share electrons when they bond. Other elements in this family are phosphorus, arsenic, antimony, and bismuth. Oxygen Family Atoms of this family have 6 valence electrons. Most elements in this family share electrons when forming compounds. Oxygen is the most abundant element in the earth’s crust. It is extremely active and combines with almost all elements. Halogen Family The elements in this family are fluorine, chlorine, bromine, iodine, and astatine. Halogens have 7 valence electrons, which explains why they are the most active non-metals. They are never found free in nature. Halogen atoms only need to gain 1 electron to fill their outermost energy level. They react with alkali metals to form salts. Noble Gases Noble Gases are colorless gases that are extremely unreactive. One important property of the noble gases is their inactivity. They are inactive because their outermost energy level is full. Because they do not readily combine with other elements to form compounds, the noble gases are called inert. The family of noble gases includes helium, neon, argon, krypton, xenon, and radon. All the noble gases are found in small amounts in the earth's atmosphere. Rare Earth Elements The thirty rare earth elements are composed of the lanthanide and actinide series. One element of the lanthanide series and most of the elements in the actinide series are called trans-uranium, which means synthetic or man-made. Octet Rule Is 8 valence electrons. Is associated with the stability of the noble gases. Helium (He) is stable with 2 valence electrons Valence Electrons – – – – He 1s2 Ne 1s2 2s2 2p6 Ar 1s2 2s2 2p6 3s2 3p6 Kr 1s2 2s2 2p63s2 3p6 4s2 3d10 4p6 2 8 8 8 Electron Shielding Shielding electrons: electrons in the energy levels between the nucleus and the valence electrons. They are called "shielding" electrons because they "shield" the valence electrons from the force of attraction exerted by the positive charge in the nucleus. Video Periodic Trends Trends in properties of the elements that follow a pattern down a group and across a period in the periodic table. 1 2 3 4 5 6 7 Trends in Atomic Radius (size) atomic radius: distance from center of nucleus to edge of electron cloud group trend: increases going down a group. – As you move down a group, energy levels are added, thus increasing the size of the electron cloud, so the atoms get larger. periodic trend: decreases going left to right across a period. – [NOTE: from now on, “across a period” will refer to the left-to-right direction] Atomic Radius Atomic Radius – Increases to the LEFT and DOWN 1 2 3 4 5 6 7 Atomic Radius Trends in Ionization Energy ionization energy: the measure of the energy required to remove an electron from the outermost energy level of an atom. group trend: decreases going down a group – This is due to the shielding effect - an electron in the outer energy level of a large atom is easier to remove because it is well-shielded from the pull of the nucleus by the inner electrons. periodic trend: increases going across a period – This is due to nuclear charge - across a period, nuclear charge increases, so it becomes more difficult to remove an electron (held tighter). – Note that this periodic trend supports the idea that metals have a much greater tendency to lose electrons than nonmetals do. Ionization Energy First Ionization Energy – Increases UP and to the RIGHT 1 2 3 4 5 6 7 Trends in Electron Affinity electron affinity (electron-liking): the energy change that accompanies the addition of an electron to an atom. – *If energy is released in the process of adding an electron, EA is negative. A negative EA means that the atom wants to gain the electron. The larger the negative number, the more it wants to gain the electron. It is a favorable process. Nonmetal atoms have large negative EAs. – *If energy is absorbed in the process of adding an electron, EA is positive. A positive EA means that the atom does not want to gain the electron. The larger the positive number, the more it does not want to gain the electron. It is an unfavorable process. Metal atoms have small negative or positive EAs. group trend: EA decreases (less favorable/value is more +) going down a group. – It is harder to add the electron when it is farther from the nucleus. The nucleus can’t “grab onto it” as well periodic trend: EA increases (more favorable/value is more - ) going across a period. – This is due to nuclear charge - across a period, nuclear charge increases, so it becomes easier to add an electron. – Note that this periodic trend supports the idea that nonmetals have a much greater tendency to gain electrons than metals do. Electron Affinity Electron Affinity – Increases UP and to the RIGHT 1 2 3 4 5 6 7 Trends in Electronegativity electronegativity: the tendency of an atom to attract electrons to itself when it is chemically bonded with another element. – numerical scale which combines both ionization energy and electron affinity. – can be used to predict whether atoms will form ionic or covalent bonds in molecules. Diagram of water molecule: – In H2O, oxygen is more electronegative than hydrogen, so it pulls the electrons closer, and thus obtains a partially negative charge. group trend: decreases down a group. – Larger atoms have more energy levels, so it is harder for them to attract electrons to the nucleus (shielding effect). periodic trend: increases across a period. – Nonmetallic character increases across a period, and nonmetals attract electrons more than metals do. Electronegativity Electronegativity – Increases UP and to the RIGHT 1 2 3 4 5 6 7 Summary of Periodic Trends Ionization Energy, Electron Affinity, and Electronegativity – Increases UP and to the RIGHT Atomic Radius – Increases to the LEFT and DOWN IE 1 2 3 IE AR EA 4 EN 5 6 7 EA AR EN