Periodic Table

I. History Stanislao Cannizzaro (1826-1910) (Italian Chemist) 1. reliable method to measure atomic masses B. Johann Dobereiner (1780-1949) (German Scientist) John Newlands (1837-1898 - English) 1. related atomic mass to properties 2. Newland’s Law of Octaves

John Newlands – Law of Octaves

John Newland’s Law of Octaves

Lothar Meyer (1835-1895 - German) 1. properties of elements show a repetitive pattern when they are arranged by atomic mass D.Dimitri Mendeleev (1834-1907- Russian) (father of modern periodic table) published system used today (1869) 2. elements arranged by increasing mass 3. left spaces for elements not yet discovered - predicted properties http://video.pbs.org/video/2365538111/

Dimitri Mendelev

Mendeleev’s Table His table re-organized

video

E. Henry Mosley (1887-1915) English 1.Arrange elements by increasing atomic number – this led to the-periodic law 2. Periodic Law - properties of elements are periodic functions of their atomic number

II. Arrangement of Elements Periodic Table – arrangement of elements in order of increasing atomic number so that elements with similar properties are in the same column 1. period – horizontal row (7) 2. group(family)- vertical columns (1-18) 3. periodicity – reoccurrence of similar properties of elements in groups

C. Special Groups on the Periodic Table Group # and Name 1 - Alkali Metals 2 - Alkaline Earth Metals 15 - Nitrogen Family - Oxygen 17 - Halogens 18 - Noble Gases

E. Metals – Metalloids - Nonmetals 1. Metals are on the left side – all are solids except mercury (Hg) a. elements near the left of a period are more metallic than those near the right b. elements near the top of a group are more metallic than those near the bottom 2. Metalloids – group of elements between metals and nonmetals(B,Si,Ge,As,Sb,Te) 3. Nonmetals are on the right side – all are solids or gases except bromine(Br) liquid

Metals – Metalloids - Nonmetals

PROPERTY METAL NONMETAL Luster high low Deformability malleable brittle and ductile Conductivity good poor Electron gain/lose lose gain Ion formed cation (+) anion(-) Ionization energy low high Electronegativity low high

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I. History of the Atomic Theory Modern Atomic Theory All matter is made up of small particles called atoms. Atoms of the same element have the same chemical properties while atoms of different elements have different properties (isotopes) Not all atoms of an element have the same mass, but they all have a definite average mass which is characteristic.

I. History of the Atomic Theory Atoms of different elements combine to form compounds and each element in the compound loses its characteristic properties. Atoms cannot be subdivided by chemical or physical changes. (nuclear reactions)

I. History of the Atomic Theory-time line 1803 1897 1909 1913 1935 Today solid particle electron proton e- orbit nucleus neutron Quantum Atom theory Dalton Thomson Rutherford Bohr Chadwick Schrodinger and others

II. History of the Atomic Structure J.J. Thomson (1887) Cathode Ray Tube Discovered matter contained negative charge Electron e -

Thomson’s Model Found the electron Said the atom was like plum pudding A bunch of positive stuff (pudding), with the electrons suspended (plums)

II. History of the Atomic Structure Robert Milikan (1909) Oil Drop Experiment Discovered mass and actual charge of electron Mass is 1/1840 the mass of a hydrogen atom e – has a mass of 9.11 x 10-28 g

II. History of the Atomic Structure So, at this point we know: Dalton’s Atomic Theory Electrons are negatively charged The mass of an electron is very small

The main points of Dalton's atomic theory were: Elements are made of extremely small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. Atoms cannot be subdivided, created, or destroyed. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged.

II. History of the Atomic Structure But: Atoms are neutral, so there must be a positive charge. Electron’s mass is so small…there must be more to an atom.

II. History of the Atomic Structure Ernest Rutherford (1909) Gold Foil Experiment Discovered the proton p+ When alpha (+2) particles hit screen, the screen lights up P.S. #p+ = atomic number

Florescent Screen Lead block Uranium Gold Foil

He expected: The alpha particles to pass straight through

He thought this would happen:

Because

Because, he thought the mass of the positive charge was evenly distributed in the atom

Here is what he observed:

So, he noticed: Most positive alpha particles pass right through Only a few were deflected

He reasoned: If a + particle hit a + point on the foil, it was repelled and deflected

He concluded: Atom is mostly empty space Has a small, dense positive center

II. History of the Atomic Structure At this point in 1909, we know: p+ = 1.67 x 10-24 g e- = 9.11 x 10-28 g The charges are balance! But, How are the electrons arranged? There is still mass that is unaccounted for

II. History of the Atomic Structure Niels Bohr (1913) Electrons orbit nucleus in predictable paths Much more on him later

II. History of the Atomic Structure Chadwick (1935) Discovers neutron in nucleus Neutron is neutral n0 Mass is 1.67 x 10-24 g

II. History of the Atomic Structure Charges balanced Mass accounted for But today, we subscribe to the Quantum Atom Theory to describe the atomic structure

II. History of the Atomic Structure Quantum Atom Theory The atom is mostly empty space Two regions: Nucleus- protons and neutrons Electron cloud- region where you have a 90% chance of finding an electron

III. Subatomic Particles Relative mass Actual mass (g) Name Symbol Charge Electron e- -1 9.11 x 10-28 Proton p+ +1 1amu 1.67 x 10-24 Neutron n0 1amu 1.67 x 10-24

Positive-negative attraction electric force between protons in one atom and electrons in another atom hold atoms together  chemical bond

III. Subatomic Particles Atomic number The number of protons in the nucleus of an atom Identifies the element No two elements have the same atomic number Mass number The number of protons plus neutrons in the nucleus of an atom Mass number is very close to the mass of an atom in amu (atomic mass units) Two atoms with the same atomic number but different mass number are called isotopes (mass number) – (atomic number) = #n 0 (number of neutrons) Electrons and Ions For neutral atoms, #e- = #p+ If there are more electrons, a negative ion forms If there are less electrons, a positive ion forms For now, we will work only with neutral atoms

III. Subatomic Particles Electrons and Ions For neutral atoms, #e- = #p+ If there are more electrons, a negative ion forms If there are less electrons, a positive ion forms For now, we will work only with neutral atoms

III. Subatomic Particles You can never change the number of protons and have the same element If you change the number of neutrons in an atom, you get An isotope If you change the number of electrons in an atom, you get An ion

X III. Subatomic Particles Notation Nuclear Notation- how we depict isotopes contains the symbol of the element, the mass number, and the atomic number Mass number # of P +N X Atomic number # of P

Na III. Subatomic Particles 23 11 Notation Nuclear Notatioin How many protons? How many neutrons? How many electrons? 23 Na 11

III. Subatomic Particles Hyphenation Notation Symbol or name of element – mass number Fluorine-19 Protons? Neutrons? Electrons? C-12

III. Subatomic Particles Average Atomic Mass Measured in grams (for a lot) Measured in amu (for a few)

Atomic Mass Mass of an atom Too small to measure in grams Use relative mass (amu) Almost the same as mass number Standard: 1 amu is defined as 1/12 the mass of one C-12 atom

III. Subatomic Particles Average Atomic Mass Weighted average mass of all known isotopes Weighted means that the frequency of an isotope is considered blackboard

1. Horizontal rows are called periods A. There are 7 periods

2. Vertical columns are called groups. a. Elements are placed in columns by similar properties. i. Also called families

3. The elements in the A groups are called the representative elements 8A0 1A 2A 3A 4A 5A 6A 7A

4. The group B are called the transition elements a. These are called the inner transition elements and they belong here- lanthanide actinide series

5. Group 1A are the alkali metals 6. Group 2A are the alkaline earth metals

7. Group 7A is called the Halogens 8. Group 8A are the noble gases

The part of the atom another atom sees is the electron cloud. More importantly the outside orbitals. The orbitals fill up in a regular pattern. The outside orbital electron configuration repeats. The properties of atoms repeat.

Classification of the elements A. Periodic Table- arrangements of elements so that elements with similar properties are in the same column 1. Valence electrons- electrons that exist on the highest principal energy level of an atom a. these are the e- that can be lost or gained in formation of a compound

2. energy level and period a. the energy level of valence electrons is indicated by which period it is found in. Ie: period 4 elements have valence electrons on the 4th energy level Lets work through a few.

3. Valence electrons and group number a. Representative element’s group number and number of valence electrons it contains also are related. Ie: Na has one e- on valence shell Lets work on some

D. Periodic Table Showing s,p,d,f Blocks