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The Periodic Table History of the Table Periodic Law Periodic Trends.

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Presentation on theme: "The Periodic Table History of the Table Periodic Law Periodic Trends."— Presentation transcript:


2 The Periodic Table History of the Table Periodic Law Periodic Trends

3 Development of the Periodic Table Dmitri Mendeleev organized the first periodic table according to atomic masses. He noticed that chemical properties repeated at regular intervals, thus the term periodic. A few elements contradicted the pattern, so he rearranged them so their properties were aligned, even though the atomic mass was not in increasing order.

4 Predicting the Future? The most amazing thing about Mendeleev ’ s work is that he predicted the properties of elements that did not exist yet. He left empty spaces for these elements, which were found and named Scandium, Gallium and Germanium. The properties that he predicted were very similar to the actual properties observed.

5 Moseley and the Periodic Law Though most of the elements could be arranged according to atomic mass, not all fit this pattern. Moseley worked with the spectra of elements and determined that they fit better into the pattern when arranged according to atomic number. Moseley ’ s periodic table still adhered to the chemical periodicity that Mendeleev observed. Therefore it is Mendeleev who is credited with periodic law, which states that the physical and chemical properties of elements are periodic functions of their atomic numbers.

6 Additions to the Table More than 40 new elements have been added since Moseley ’ s time. Most significant additions: Noble gases - Group 18 Lanthanides (#58-71) Actinides (#90-103)

7 Organization of the periodic table Rows are referred to as periods. Columns are referred to as groups or families. We have discussed the organization of the periodic table as a result of electron configuration. There are also patterns of periodicity between elements in a group.

8 Organization of the periodic table Valence electrons play an important role in reactivity of elements. Valence electrons are defined as the electrons in the outermost energy level. You will notice a pattern between the groups on the periodic table and number of valence electrons.

9 The Octet Rule Elements want to have 8 electrons in their valence shell. To get 8 electrons, some atoms gain electrons and others lose electrons. When this happens, an ion is formed. Ions are atoms that have gained or lost one or more electrons. The only exceptions to the octet rule are Hydrogen and Helium…they desire 2 electrons to fill their valence shell.

10 Main group elements The term main group elements refers to elements in the s and p blocks only. Remember the trend that allows you to determine charge of an ion for elements in these groups: Group 1=+1 Group 14=+/- 4 Group 2=+2Group 15= -3 Group 17=-1 Group 13=+3 Group 16=-2

11 Group 1- Alkali Metals This group of metals is extremely reactive! This is due to the fact that they all have only 1 valence electron. In their pure state, all members of this group are silvery in appearance and soft enough to be cut with a butter knife. Because they are so reactive, they are not found in nature as “ free ” elements. They react readily with water, even moisture in the air. Therefore, they must be stored in kerosene or mineral oil.

12 Group 2- Alkaline Earth Metals These metals are harder, denser and stronger than alkali metals. They have 2 valence electrons. They have higher melting points and are less reactive than group 1 metals. However, they are still too reactive to be found “ free ” in nature.

13 Special Cases: Hydrogen and Helium Even though H has 1 valence electron, it does not share the same properties as the other group 1 elements. Can you think of some differences? Helium has 2 valence electrons, but you will notice that it is not placed above group 2. It is placed over group 18 because it has a full valence shell.

14 Transition Elements These are elements belonging to groups 3-12. The number of valence electrons varies for these elements. These elements have metallic properties that you might suspect: good conductivity of electricity, high luster and good conductivity of heat. They are also less reactive than Group 1 and 2 metals, so they are found in nature in free form.

15 P-block Elements Groups 13-18, excluding Helium The p-block and s-block elements are collectively called the main group elements. For the p-block elements, the number of valence electrons is equal to the group number minus 10. The groups vary drastically in their properties.

16 The P-block Divide The p-block elements are divided into three sections by the stair-step line located in the center of the block. Elements that border the line are called metalloids or semiconductors. These include boron, silicon, germanium, arsenic, antimony and tellurium. To the left of those elements are metals. To the right of the metalloids are nonmetals.

17 The Halogens Halogens are elements that belong to group 17 on the periodic table. They are the most reactive nonmetals, so they react vigorously with metals to form salts. Their reactivity is based on the fact that they have 7 valence electrons. Since all elements desire a full valence shell (halogens desire to get one more!)

18 The Halogens Fluorine and chlorine are gases at room temp. Bromine is a reddish liquid and can burn your skin! Iodine is a dark purple solid. Astatine is rare but is known to be a solid. Only small quantities exist.

19 Justifying Periodic Trends It is REALLY important that you understand how to justify a trend observed on the periodic table. This is a really major topic in AP Chem. AP will not accept that an element is more (insert term here… electronegative, for example) because it is further to the right on the periodic table. You MUST explain in terms of nuclear charge, or more energy levels. I have given you the proper justification for each trend on the slides to follow!

20 Periodic Trends Atomic Radius- used to determine the size of an atom Defined as 1/2 the distance between the nuclei of two identical atoms that are bonded together The trend for atomic radius on the periodic table is: Size of atoms decrease across a period (as a result of increasing positive charge of nucleus, which attracts electrons more closely) Size of atoms generally increase down a group (electron energy levels are further from the nucleus and increase in number)

21 Periodic Trends Ionization energy- the energy required to remove an electron from an atom Any time an electron is removed from an atom, it is referred to as ionization. The trend for ionization energy on the periodic table is: IE of main group elements increases across a period (this is due to increased nuclear charge). IE of main group elements decreases down a group (electrons are further from the nucleus requiring less energy to remove an electron.

22 Shielding Part of the reason that large atoms can give up valence electrons so easily to form ions is an effect called shielding. Shielding refers to the fact that the inner shell electrons shield the positive nuclear charge from the outer electrons, making them easier to remove. An atom can have more than 1 IE if more than one electron is removed!!!

23 Periodic Trends Electron Affinity- the energy change that occurs when a neutral atom acquires an electron Trends on the periodic table: Generally, EA increases across the p block period. An exception to this is between groups 14 and 15. It decreases slightly. Hard to establish group trends for EA.

24 Periodic Trends Ionic radii- the radius of an atom once it has become an ion Positive ions are called cations, which are formed from the loss of electrons. Negative ions are called anions, which are formed from the gain of electrons. Trends on the periodic table: Generally, metals, to the left of the stair- step line, form cations. Cationic radii decrease across a period, due to increased nuclear charge.

25 Ionic Radii Continued Nonmetals, to the right of the stair-step line, form anions. Generally, ionic radii of anions decreases across a period. For groups on the periodic table: ionic radii increase down a group, due to adding more energy levels.

26 Periodic Trends Electronegativity- measure of the ability of an atom in a chemical compound to attract electrons The reason this is necessary is that the negative charge of electrons in a compound is not evenly distributed between the atoms. Most electronegative atom is Fluorine, and has an electronegativity value of 4 assigned to it. Values for the other elements are assigned in relation to F.

27 Electronegativity Continued Trends on the periodic table: Electronegativities tend to increase across a period. They tend to decrease or remain the same down a group. Noble gases are unique because many are not assigned electronegativity values, since they are relatively unreactive. When a noble gas does form a compound, the electronegativity value is quite high, like F.

28 Periodic Trends in the d and f block Atomic radii- decrease across a period in the d block Ionization energies- generally decrease across the period for both f and d, BUT d block elements have IE that increase down each group Ionic radii and ion formation- order in which electrons are removed is exactly the reverse order from their electron configuration Electronegativity- d block ranges from 1.1-2.54 F block range from 1.1-1.5

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