The Periodic Table Chapter 6

Introduction The periodic table represents an organizing principle which allows the prediction of the properties of each element based on their position in the periodic table. The elements of the periodic table can be classified into different categories. There are trends that appear in the periodic table that allow us to make predictions about an atoms size, ionization energy, and electronegativity.

Organizing Elements (Section 6.1) Searching for an Organizing Principle Mendeleev’s Periodic Table The Periodic Law Metals, Nonmetals, and Metalloids

Defining the Periodic Table Periodic Table: An arrangement of elements in which the elements are separated into groups based on a set of repeating properties. · A periodic table allows you to easily compare the properties of one element (or group of elements) to another element (or group of elements).

I.) Searching for an Organizing Principle Copper, gold, and silver have been known for thousands of years. Only 13 elements were known of by the year 1700. Between 1765-1775 five more elements were discovered. Problem was how to ID new elements and how many new ones there actually were.

The Dobereiner System Organized known elements into triades. Recognized a relationship between atomic weights and chemical properties Triade: A set of three elements with similar properties (ex. Cl, Br, I) Not all the known elements could be grouped into triades J.W. Dobereiner German Chemist 1780-1849

The Dobereiner Triades 1 element in each trade tended to have properties with values that fell midway between the other two. Here we see atomic and mass numbers

II.) Mendeleev’s Periodic Table From 1829-1869 many other systems were proposed but none gained wide acceptance. Mendeleev created his table while working on a text book for his students Beat a competitor because he was better able to explain the table’s usefulness. Dmitri Mendeleev Russian Chemist/Teacher 1834-1907

Arranged the elements in his table in order of increasing atomic mass. There was a close match between the predicted properties of unknown elements and the actual properties of the elements. This organizational method had its problems and does not account for all elements (ex. atomic masses for Te and I).

III.) The Periodic Law There was a problem with organizing elements by atomic masses. Organizing elements by increasing atomic number was more useful. Moseley determined the atomic number for the known elements at the time Henry Moseley British Physicist 1887-1915

Expression of the Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

The Modern Periodic Table Horizontal Rows = Periods Vertical columns = Groups

Periods in the Periodic Table These are 7 rows extending horizontally across the periodic table. Period 1 = 2 elements Period 2/3 = 8 elements Period 4/5 = 18 elements Period 6/7 = 32 elements Each period corresponds to a principle energy level. More elements in the higher periods because there are more orbitals in the higher energy levels Properties of elements within elements within a period change as you move from left to right.

Groups in a Periodic Table These are the 18 columns that run up and down the periodic table. There are 3 different ways that the groups are numbered. These groups also possess names. Elements within a group in the periodic table have similar physical and chemical properties.

IV.) Metals, Nonmetals, & Metalloids We saw how the periodic table can be divided into 7 periods and 18 groups. We can also divide the table into three broad classes based on the general properties of the elements. Metals Nonmetals Metalloids Across the periods, the properties of elements become less metallic and more nonmetallic.

The Three Classes of Elements

Metals This is the most numerous class. Characteristics of metals: Good conductors of heat and electric current Possess luster and sheen Solids at room temperature (except Hg) Ductile (i.e. can be drawn into wires) Most are malleable (i.e. can be hammered into thin sheets)

Examples of Metals Copper is ductile. Aluminum is malleable and has luster and sheen.

Nonmetals Less numerous than the metals There is greater variation in the characteristic of nonmetals. One general characteristic: They are not metals. Poor conductors of heat and electric current (carbon is an exception) Solids tend to be brittle. Most nonmetals are gases at room temperature, a few are solids, and 1 is a liquid (Br).

Examples of Nonmetals Chlorine is a gas Bromine is a liquid Carbon is a solid and is a good conductor of electricity.

Metalloids Least numerous elements Have properties that are similar to those of metals and nonmetals. The behavior is often controlled by changing the conditions.

Examples of Metalloids Silicon is not a good conductor of electricity until mixed with boron. Arsenic has luster and sheen like metals.

Classifying Elements (Section 6.2) Squares in the Periodic Table Electron Configuration in Groups Transition Elements

I.) Squares in the Periodic Table All periodic tables display at least the symbol, the atomic number, and the mass number of the elements. Some periodic tables provide more information such as physical state, electron configuration, and classification of each element.

Noble Gases Alkali Metals Alkaline Earth Metals Halogens

II.) Electron Configuration Groups Elements can be sorted into separate groups based on their electron configuration Noble Gases Representative Elements Transition Metals Inner Transition Metals

The Noble Gases Let’s show this for helium, neon, argon, and These are the elements located in Group 18 (the farthest column to the right) These are the “inert” gases because they rarely participate in reactions. The “s” and “p” orbitals of the highest occupied energy levels are filled for all noble gases. Let’s show this for helium, neon, argon, and krypton by writing out the electron configuration for each.

Helium Neon Argon Krypton

The Representative Elements These elements display a wide range of physical and chemical properties. The atoms of the representative elements have “s” and “p” orbitals of the highest occupied energy levels that are not full. For any element of this group, the group number (the American and European numbering system) equals the number of electrons in the highest occupied energy level.

Let’s Look at the Electron Configuration of Some Representative Elements Lithium Sodium Carbon Silicon

Transition Elements Two kinds of transition metals: transition and inner transition metals – classification is based on the electron configuration of an element. Transition Metals: Atoms have highest occupied sublevels that have electrons in the “s” and “d” orbitals. Inner Transition Metals: Atoms of these metals have the highest occupied “s” orbital and nearby “f” orbitals that contain electrons.

Let’s Look at the Electron Configuration of Some Transition Elements Iron Silver Nickel Chromium

Let’s Look at the Electron Configuration of Some Inner Transition Elements Cerium Uranium

The Divisions Based on Electron Configuration The electron configuration and the position of an element in the periodic table gives a particular pattern. This pattern are the blocks we see here.

Using the periodic table to write electron configurations. Based upon the blocks that were described in the previous slide, we can write the electron configuration for any element based on its location in the periodic table. The steps: Find the element on the periodic table. Start counting from hydrogen. Move towards the right At the end of each row drop down one row Begin counting towards the right again Each row represents an energy level. Each square represents an electron.

Noble Gas Configuration This is a shorter way to write the shorthand notation for electron configurations. There are four easy steps: 1.) Locate the element in the periodic table. 2.) Find the noble gas that precedes it. 3.) Place the symbol for this gas in brackets ([ ]). 4.) Write the remaining electron configuration. Write the noble gas configuration for rubidium (Rb).

Let’s try this. Write the extended and Noble gas configuration for the following elements. 1.) chlorine 2.) lead

Section 6.3 Periodic Trends Trends in Atomic Size Ions Trends in Ionization Energy Trends in Ionic Size Trends in Electronegativity

Atomic Size Atomic radius: One half the distance between the nuclei of two atoms of the same element when the atoms are joined. Distance between Nuclei Atomic Radius

Ionization Energy Ions: An atom or a group of atoms that has a The energy required to remove an electron from an atom. Ions: An atom or a group of atoms that has a positive or negative charge resulting from a loss or gain of electrons, respectively. Cation: An ion with a positive charge Anion: An ion with a negative charge

Electronegativity This property is related to chemical bonding The ability of an atom of an element to attract electrons when the atom is in a compound. This property is related to chemical bonding and is best understood by examining chemical bonds.

I.) Periodic Trend in Atomic Radius In general, atomic size increases from top to bottom within a group and from right to left across a period.

These trends depend upon the number of protons and electrons being added as we move through the periodic table.

Question: Based on the data for alkali metals and halogens, how does the atomic size change within a group? Why?

Within a group: Increasing # of occupied orbitals. Shielding of outer electrons increases the atomic radius. Within a period: Electrons are being added to the same energy level. Shielding is constant Increasing nuclear charge pulls all electrons closer

II.) Ions An atom or a group of atoms that has a positive or negative charge resulting from a loss or gain of electrons, respectively.

Cations These are positively charged ions, resulting from a loss of an electron. Metals tends to lose electrons from their highest occupied energy levels to become cations. The charge for a cation is written as a number followed by a plus sign. Na1+ Ca2+

Representing the Formation of Cations Atoms tend to lose their electrons from the outer most energy levels to become cations. Na → Na+ + e- Ca → Ca2+ + 2e- Why would the atoms lose their electrons to become cations?

Anions These are negatively charged ions resulting from a gain of electrons. Nonmetals tend to add electrons into their highest occupied energy levels to become anions. The charge for an anion is written with a number followed by a negative sign. Cl1- F1-

Representing the Formation of Anions Atoms tend to accept electrons into their highest occupied energy levels to become anions. e- + Cl → Cl- e- + F → F- Why would these atoms accept electrons to become anions?

Ionic Compounds Fe2(SO4)3 FeS2 Fe2O3

III.) Trends in Ionization Energy The energy required to remove an electron from an atom. This energy is measured when an element is in the gaseous state. There are ionization energies for every electron in an atom The 1st ionization energy is the energy needed to remove the first electron from an atom. Each successive ionization energy increases dramatically.

Periodic Trend in 1st Ionization Energy

Group Trend: As size of the atom increases the nuclear charge has a smaller effect on the outer most electrons. Period Trend: Nuclear charge increases as we move across a period, thus nuclear attraction increases.

III.) Trends in Ionic Size

IV.) Electronegativity The ability of an atom of an element to attract electrons when the atom is in a compound. Electronegativity values are calculated from ionization energies.

In general, electronegativity values increase from bottom to top within a group. For representative elements, the values tend to increase from left to right across a period.

The Periodic Table Chapter 6 The End