Unit 2 The Periodic Table

Section 1 The Periodic Table

Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were put in the same row and/or column.

Mendeleev In mass order, some properties didn’t match up. But he predicted properties of undiscovered elements.

Moseley Henry Mosely (1913, British) Organized elements by increasing atomic number. Properties vary in regular ways up and down and across the periodic table. Resolved discrepancies in Mendeleev’s arrangement. Tellurium & Iodine

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

Modern Periodic Table Family/Group Period Numbered 1-18 from left to right. Elements in each column have similar chemical and physical properties. Period Numbered 1-7 from top to bottom. Closer elements have more similar properties. Chemical and physical properties change somewhat regularly across a row.

The Modern Periodic Table Ninety naturally occurring elements plus an additional twenty-eight synthetic elements for a total of one hundred & eighteen elements. Most elements are solids at room temperature.

The Modern Periodic Table Lanthanide Series & Actinide Series Metallic elements that fit into the table just after elements 56 & 88 Placed at the bottom to keep the table from being too wide.

METALS Good conductor of heat and electricity. Ductile (drawn into wire) High tensile strength (ability to resist breaking when pulled) Malleable (hammered into a thin sheet) Most have a silver or grayish-white luster. Solid at room temperature, except Hg

NONMETALS Poor conductor of heat and electricity. Often gases, but not all of them. Solids tend to be brittle

METALLOIDS Have some characteristics of both metals and nonmetals. All are solids at room temperature Tend to be semiconductors of electricity Do not unconditionally conduct electricity like metals but do not impede electric flow like non-metals

Section 2 Periodic Groups

Groups of the Periodic Table

Alkali Metals – Group 1 Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium As pure elements: Soft, silver, metals. So soft that the metal can easily be cut with a dull knife. But… they are SO reactive that they are never found in nature as pure elements. React vigorously with most non-metals React strongly with water (even as vapor) to produce hydrogen Usually stored in kerosene to keep them from reacting Melt at low temperatures.

Alkaline Earth Metals – Group 2 Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium Harder, denser, stronger than the Alkali Metals and have higher melting points Less reactive than the Alkali Metals – not found in nature as free elements but when purified in the lab they can be kept pure in a jar.

Transition Metals – Group 3-12 Mercury, Gold, Silver, Copper, Zinc, Nickel, etc. 1 or 2 valance electrons Less reactive then Alkali or Alkaline Earth metals Good conductors of electricity High luster Posses metallic properties

Halogens – Group 17 Fluorine, Chlorine, Bromine, Iodine, and Astatine Are the most reactive non-metals. React vigorously with metals. Do not exist in nature as single atoms of the pure element. F2, Cl2, Br2, and I2 exist because they’ve formed molecules. F2 and Cl2 are gases at room temperature, Br2 is a liquid, and I2 is a dark purple solid. Astatine is a synthetic element.

Noble Gases – Group 18 Helium, Neon, Argon, Krypton, Xenon, and Radon All are gases at room temperature He is less dense than air (It “floats” in air) Ne, Ar, Kr, and Xe are used in lighting Rn is radioactive All have 8 electrons in the valence shell without forming compounds, so … Considered to be un-reactive (or “inert”) so natural compounds of these elements do not exist In 1962 Xenon tetrafluoride was produced at extremely high temperatures and pressures.

Section 3 Periodic Trends

Valence Electrons The number of valence electrons in an atom is equal to the number of electrons in the highest numbered energy level. In any period, the number increases moving from group 1 to 2, and then from group 13 to 18. In any group, all atoms have the same number of valence electrons.

Valence Electrons 1ve- 8ve- 2ve- 3ve- 4ve- 5ve- 6ve- 7ve- 1ve- or 2ve-

Trends in Atomic Radius The atomic radius is ½ the distance between the centers of the nuclei of two identical atoms bonded together. SIZE OF THE ATOM Done this way because the “edge” of an atom is somewhat “fuzzy”

Trends in Atomic Radius PERIOD TREND Decreases going across left to right Why? More positively charged nucleus (more protons) pulls electrons in tighter. GROUP TREND Increases going down Why? Increasing number of energy levels (more electrons)

Fig. 14.10, p401

Ionic Radii Positive ions (cations) are SMALLER than the corresponding neutral atom The removal of the outer layer of electrons makes it smaller There are the same number of protons, but fewer electrons. Each electron is pulled closer by the unbalanced charge

Ionic Radii Negative ions (anions) are LARGER than the corresponding neutral atom There are more electrons, but no more protons. Each electron is not drawn to the nucleus as strongly, electrons in the cloud repel one another

Trends in Ionization Energy The amount of energy needed to remove an electron from a neutral atom is the ionization energy. It may be easy to remove the first electron, but it takes more energy to remove each one after that. Ionization energy generally decreases as you move down a group because the electrons that you are removing are further away from the positive nucleus – the “pull” isn’t as great. Generally, Ionization energy increases as you move across a period because the valence level is getting more full. Increasingly, it would rather gain electrons.

Trends in Electron Affinity (NOT in your book – but it is on the PA Keystone Standards) Electrons can also be gained. The energy change that occurs when a neutral atom gains an electron is the electron affinity. Trends are opposite the trends of Ionization Energy. Decrease across a period / Increase down a group

Trends in Electronegativity Valence electrons hold the atoms together in a compound, but when this happens it causes the electrons to gather closer to one atom than to the other. The electronegativity is the ability of an atom to attract the electrons. The most electronegative elements are in the upper right if you ignore the noble gases The least electronegative elements are in the lower left, the values only decrease slightly going down a group.

Trends in Chemical Reactivity (Reactivity = ability to form compounds via reactions) Noble Gases do NOT naturally react with anything Elements in Groups 1 and 17 (Alkali metals and Halogens) are HIGHLY reactive. They do not exist in nature as single atoms. Elements in Groups 2 and 16 are reactive, but not as reactive as groups 1 + 17. They can exist as pure elements but only after we artificially purify them. Transition metals can exist as pure elements (but often exist as compounds Metals are more reactive at the bottom of a group Non-metals are more reactive at the top of the group