VCE CHEMISTRY-UNIT 1 PART 1 DEVELOPMENT OF THE PERIODIC TABLE PERIODIC TABLE

MENDELEEV & THE PERIODIC TABLE Mendeleev discovered that the chemical properties of elements follow a periodic pattern as their relative atomic weights increase. This pattern is called periodicity. He noticed as the atomic weights increased across a period, the chemical properties gradually changed.

In 1869 he created the periodic table by: arranging elements with similar chemical properties into vertical columns (groups) arranging the elements in order of increasing relative atomic mass (periods) Development of the Periodic table

Mendeleev left gaps in the periodic table where he believed there should be elements that hadn’t yet been discovered. He was able to predict their chemical properties by the periodic pattern. Three predicted elements (Gallium, Germanium, and Scandium) were discovered 20 years later.

MENDELEEV’S PERIODIC TABLE Today’s table is ordered according to atomic number and electron configuration. HOWEVER.. since chemical properties result from electron configuration the modern table and Mendeleev’s table are the same

Mendeleev’s Periodic Law The properties of elements vary periodically with their atomic weights.

William Ramsay and the Noble Gases Ramsay was first to isolate the noble gases from the atmosphere (helium argon, neon, krypton, and xenon) Contributed to the discovery that helium is a product of the atomic disintegration of radium.

GLENN SEABORG AND THE TRANSURANIUM ELEMENTS Seabourg discovered many radioactive elements such as: Generally called the transuranium elements. californium, einsteinium, fermium, mendelevium nobelium plutonium americium curium berkelium

Task Write a short summary in your books of the contribution to the development of the Periodic Table by : Mendeleev Ramsay Seaborg

VCE CHEMISTRY-UNIT 1 Part 2: Understanding the Trends in the periodic table PERIODIC TABLE

Periodic Properties of the Elements We now know that the periodic properties are due to the electronic structure of atoms. Electronic structure explains the observed trends in –Atomic size –Metallic/non-metallic character –Chemical reactivity Al ZnGa CdIn Sn

RIM vs RAM RIM (I r ) uses the mass spectrometer to calculate the mass of the isotope RAM (A r ) is an average of the masses of all of the isotopes of that element naturally found in any mixture (p of text)

Group I elements Group I elements (minus hydrogen) are called the Alkali metals. They are very reactive elements. Eg: they reactive violently with water, increasingly down a group Hydrogen is placed in Group I as it has 1 electron in it’s outer shell but does not share similar properties

Group II elements Group II elements are called Alkaline Earth metals. They are not as reactive as Group I elements. They tend to form bases when mixed with H 2 O

Electron Shells The inner, completed electron shells are called core shells. The electrons in these shells are called core electrons. The outermost electron shell is called the valence shell. The electrons that sit in this shell are called valence electrons.

Core Charge The higher the core charge, the greater the ‘pull’ of the valence electrons to the nucleus – this decreases the radius A greater ‘pull’ draws the cloud of electrons further in, thus reducing the diameter.

Atomic Size The radius of an atom is found from the distance between nuclei in a molecule ci/chemistry/essentialchemis try/flash/atomic4.swf

ATOMIC SIZE -DOWN A GROUP EXPLANATION Number of shells ……………. Down a group Electrons are ……….…... away from the nucleus Size of atoms therefore ………….. down a group

Ionization Energy Ionization energy - The minimum amount of energy required to remove the outermost electron The greater the amount of electrons in the shell, the harder it is to remove one Shells are described in terms of the energy required to remove an electron

Electronegativity Definition The ability of an atom to attract electrons. The more electrons in the valence shell, the greater the ability to attract more electrons.

Metals and Nonmetals Metals are characterized by low ionization energy and high electrical conductivity; Non-metals by high ionization energy and poor electrical conductivity.