Ch 6 The Periodic Table Development of the

Ch. 6: The Periodic Table

Development of the Modern Periodic Table • Why is it called the periodic table? - Properties of the elements repeat in a periodic way - Periodic: predictable

History of the Table’s Development • Antoine Lavoisier - 1700 s - Compiled a list of all the elements known at the time - 33 elements in 4 categories • John Newlands - 1864 - Proposed an arrangement where elements were ordered by increasing atomic mass - Law of Octaves: when elements were arranged by increasing atomic mass, the properties repeat every 8 th element

History of Table’s Development • Dmitri Mendeleev - 1864 Created first periodic table Arranged elements in atomic mass order Major contribution: left blank spaces and predicted the properties of some undiscovered elements - Stated the 1 st periodic law - Periodic Law: when the elements are arranged in increasing atomic number order, there is a periodic repetition of their chemical and physical properties

In this early version of Mendeleev’s periodic table, the rows contain elements with similar properties.

History of Table’s Development • Lothar Meyer - 1800 s - Also demonstrated a connection between atomic mass and elemental properties - Arranged elements in order of increasing atomic mass - Not as credited as Mendeleev because Mendeleev’s work was published first • Henry Moseley - 1913 - Realized arranging the elements by atomic mass was not the best way—some elements ended up in columns with elements of different properties - Established the atomic number - Rearranged the table by increasing atomic number which resulted in a clear periodic pattern

The Modern Periodic Table - Organization • Periods (series): horizontal rows - The elements in each period have the same number of shells • Groups (families): vertical columns - The elements in each group have the same number of valence electrons - Group numbers 1 -18 (new system) or A/B system (older) • Representative Elements - elements in the s and p blocks - Groups 1, 2 and 13 -18 - Possesses a wide variety of chemical and physical properties • Transitional Elements - elements in groups 3 -12

Periodic Table Organization Special Group Names • Alkali Metals: group 1 except for Hydrogen • Very reactive • Alkaline Earth Metals: group 2 • Highly reactive • Transitional Metals: group 3 -12 • Inner Transitional Metals: 2 series below the periodic table • Lanthanide Series: 1 st row in f-block • Actinide Series: 2 nd row in f-block • Halogens: group 17 • Highly reactive, “salt former” • Noble Gases: group 18 • Extremely unreactive, stable

Special Group Names Inner Transitional Metals

Periodic Table Organization Standard State • Solids- most elements • Liquids- Br and Hg • Gases- upper right corner, H, N, O, F, Cl, and nobles • Synthetic Elements- not found in nature, made in lab - Tc (#43), Pm (#61) & all elements after uranium (#93 and higher)

Standard State

Periodic Table Organization Classes of Elements • Metals • Shiny, solid at room temperature, good conductors, ductile, malleable, high melting points • Nonmetals • Generally gases or brittle, dull-looking solids, poor conductors, low melting points • Metalloids (B, Si, Ge, As, Sb, Te, Po, At) --staircase • Intermediate properties

Metalloids – along step line (except Aluminum)

• A company plan to make an electronic device. They need to use an element that has a chemical behavior similar to that of Silicon (Si) and lead (Pb). The element must have a atomic mass greater than sulfur (S), but less than that of cadmium (Cd). Use the periodic table to determine which element the company can use.

• Identify each of the following as metal, nonmetal, and metalloid - Oxygen - Barium - Germanium - Iron

• If the periodic table were arranged by atomic mass, which of the first 55 elements would be ordered differently than they are in the existing table?

Periodic Trends Many properties of the elements tend to change in a predictable way, known as a trend, as you move across a period or down a group.

Atomic Radius • One-half the distance between the nuclei of identical atoms that are bonded together • How closely an atom lies to a neighboring atom Trends: • generally DECREASES across a period (L R) • generally INCREASES down a group

Atomic Radius WHY the trends occur: • Trend across a period: As electrons are added to the s and p sublevels, they are pulled closer to the nucleus, decreasing the radii • Trend down a group: As electrons occupy sublevels in successively higher energy levels located farther from the nucleus, the sizes of the atoms increase

Ionic Radius • Ion: an atom or bonded group of atoms that has a positive or negative charge • When atoms lose electrons and become positive ions, they always become smaller (compared to the neutral atom) • Loss of valence electron can leave an empty outer orbital resulting in a small radius • When atoms gain electrons and become negative ions, they become larger

Octet Rule • Octet Rule: atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons • Used to predict what types of ions an element is likely to form • Exception: Hydrogen does not want 8 valence electrons, it will usually give up its one electron to make a positive ion

Ionic Radius Trends: • generally positive ions DECREASE across a period (L R) • generally negative ions (beginning in group 15 or 16) DECREASE across a period (L R) • generally ionic size INCREASES down a group WHY the trends occur: • Trend across a period: The loss of a valence electron can leave a completely empty outer orbital, which results in a smaller radius. The addition of an electron to an atom increases the repulsion between outer electrons, forcing them apart • Trend down a group: extra inner shell of electrons is added period causing the radius to increase

Ionic Radius

Ionization Energy • The energy required to remove one electron from a neutral atom of an element • Energy required to remove the outermost electron from a gaseous atom • Measured in k. J/mol • First Ionization Energy: the energy required to move the first electron • Removing the second electron requires more energy (second ionization energy)

Ionization Energy Trends: • generally INCREASES across a period (L R) • generally DECREASES down a group WHY the trend occurs: • Trend across a period: radii has been decreased, meaning a very strong attraction between the electrons and the nucleus. This means that it will be harder to remove an electron. • Trend down a group: the ionization energy decreases due to a larger radius. Attraction between the nucleus and the outer electrons is also decreased because there are more orbits in between blocking the way.

Electronegativity • Ability to attract electrons in a chemical bond • Values range from 0. 7 to 4. 0 Pauling units (arbitrary units) • Most noble gases have no values Trends: • generally INCREASES across a period (L R) • generally DECREASES down a group

Density • The amount of mass per unit volume Trends: • generally it INCREASES from left to the middle, and then DECREASES (more rapidly) from the middle to the right across a period • generally INCREASES down a group Why the trend occurs: • Increases as you go down because the atomic radius increases, volume increases, as well as mass at a larger rate, so it will be more dense.

Trends in Density Generally Increases