Unit 6 The Periodic Table Properties of Elements

Unit 6: The Periodic Table Properties of Elements and Trends

Mendeleev The original periodic table was created by Dmitri Mendeleev He organized the table by atomic mass He predicted existence of 3 missing elements discovered later on

Modern Periodic Table The modern periodic table is organized by atomic number. Vertical columns called groups or families organize elements by reactive properties (valence e-) Horizontal rows called periods also represent the shells or energy levels The concept of periodic law is based on repeating patterns or trends in element properties. Henry Mosely

Major Types of Elements There are 3 main types of elements: 1) Metals 2) Metalloids 3) Non-Metals

Take Out Your Coloring Tools

Properties of Metals Solid at room temperature (except Hg) Luster – Reflects light to give a shine Malleable – Moldable or bendable Ductile – Can be stretched into thin wire Most have high melting points Good conductor of heat & electricity Have metallic bonds

Properties of Non-Metals Can be solid, liquid or gas at room temp Exist as Allotropes - different molecular forms of the same element (different properties, too) O 2 and O 3 Dull, lack luster Brittle Poor conductors

Metals Vs. Non-Metals lose electrons to form a full valence shell and become + ions (cations). Non-metals gain electrons to form a full valence shell and become – ions (anions). Ionic Compounds

Properties of Metalloids Solid at room temperature Semi-conductors Have varying properties that can be metallic or non-metallic 6 of them:

Groups on the Periodic Table In general, elements found within the same group have similar chemical and physical properties. This is mainly due to the fact that they have the same number of valence electrons.

Alkali Metals (Group 1) Metals (very reactive) Have one valence electron in outer shell Tend to lose that one electron to become a stable cation. Rarely found in nature. Exist as solids at room temperature, but are soft Silvery-white or grayish in color Examples:

Alkaline Earth Metals (Group 2) Metals (2 nd most reactive) Have two valence electrons Tend to lose the two electrons to become a stable cation Exist as solids at room temperature Silvery white or grayish white in color Examples:

Halogens (Group 17) Have 7 valence electrons Non-metals (most reactive) F, Cl (gases), Br (liquid), I (solid) Tend to gain one electron to form a more stable anion. Examples:

Noble Gases (Group 18) Very stable because of filled valence shell (8 electrons). All non-metal (gases at STP) Not Reactive Examples:

Transition Metals (Groups 3 -12) Also known as “heavy metals” Not as reactive metals Tend to be very dense (heavy for small amounts) Have multiple oxidation states (can form more than one ion) so we use roman numerals when naming compounds with them. Produce very colorful solutions when dissolved in water Examples:

Lanthanide & Actinide Series These are rare earth metals Many have applications in nuclear technology They actually fit in between groups 3 and 4 of the periodic table Above 92 they are mostly man-made and are unstable

Periodic Trends According to periodic law trends occur in the periodic table in the following Categories: - reactivity - metallic character - atomic radius* & ionic radius* - ionization energy* -electronegativity* *Sketch this on a page in your notebook

Period Trends in Atomic Radius – Is the distance from the center of the atom, “The Nucleus, ” to the furthest away electron shell Atomic radius decreases

Period Trends in Ionization Energy The energy required to remove an electron from an atom is known as the FIRST IONIZATION ENERGY. To remove a second electron requires what is called the SECOND IONIZATION ENERGY. Increasing IE

Cation Size When an electron is lost from an atom, the nucleus now has more pull on the outer energy level making the ion SMALLER. MINUS ONE ELECTRON

Anion Size When an electron is gained the nucleus now has less pull on the outer energy level making the ion LARGER. PLUS ONE ELECTRON

Electronegativity Can be defined as an atom’s ability to attract electrons. Expressed in arbitrary units on Table S of your Reference Tables. Non-metals tend to have higher electronegativities because they like to gain electrons. atoms electron s