S MORRIS 2006 Where did it all begin

S. MORRIS 2006

Where did it all begin? The word “atom” comes from the Greek word “atomos” which means indivisible. The idea that all matter is made up of atoms was first proposed by the Greek philosopher Democritus in the 5 th century B. C.

HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (greek for indivisible)

HISTORY OF THE ATOM 1803 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. 2. All atoms of a given element are identical. The atoms of one element are different from those of any other element 3. Atoms of different elements can combine in simple whole number ratios to form compounds. 4. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2. 1

16 X + 8 Y 8 X 2 Y 2. 1

Law of Definite Proportions • The ratio of mass of elements in a compound is always the same • Every Water molecule will contain 16 g of oxygen and 2 g of hydrogen

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HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

HISTORY OF THE ATOM 1890 Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge like plums surrounded by pudding. PLUM PUDDING MODEL

HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them passed through. About 1 in 10, 000 hit

(1908 Nobel Prize in Chemistry) a particle velocity ~ 1. 4 x 107 m/s (~5% speed of light) 1. atoms positive charge is concentrated in the nucleus 2. proton (p) has opposite (+) charge of electron 3. mass of p is 1840 x mass of e- (1. 67 x 10 -24 g) 2. 2

Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m 2. 2

HISTORY OF THE ATOM helium nuclei gold foil helium nuclei They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.

HISTORY OF THE ATOM Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus. He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction However, this was not the end of the story.

HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 a + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1. 67 x 10 -24 g 2. 2

Subatomic Particles mass p = mass n = 1840 x mass e 2. 2

Bohr’s Atom electrons in orbits nucleus

Atoms • # of protons & neutrons • The basic unit of Matter • The smallest particle of an element that retains the properties of that element.

HELIUM ATOM Shell proton + - N N + electron What do these particles consist of? - neutron

ATOMIC STRUCTURE All About Atoms Particle Charge Mass Proton + charge 1 amu Neutron No charge 1 amu electron - charge 1/1836

ATOMIC STRUCTURE He 2 Atomic number Represents the number of protons in an atom Never changes 4 P+ equal to the number of e. Mass Number the number of protons and neutrons in an atom Neutrons equal mass # - atomic # number of electrons = number of protons

Ions Charged particles due to the loss or gain of electrons 2) Atoms that are called cations lose e- thus becoming positive Na 11 protons 11 electrons Na+ 11 protons 10 electrons 1) Atoms that are called anions gain e- thus become negative Cl 17 protons 17 electrons Cl- 17 protons 18 electrons 2. 5

Isotopes • Atoms with the same atomic number but different mass number • Atoms having the same number of protons but different numbers of neutrons • Average Atomic Mass of an element is the weighted average of an element’s naturally occurring isotopes

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Do You Understand Isotopes? Isotope Maker How many protons, neutrons, and electrons are in 146 C? 6 protons, 8 (14 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are in 116 C? 6 protons, 5 (11 - 6) neutrons, 6 electrons 2. 3

Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons # OF NEUTRONS = mass number – atomic number Mass Number A ZX Atomic Number 1 1 H 235 92 2 1 H U Element Symbol (D) 238 92 3 1 H (T) U 2. 3

Bohr Model of the Atom • Electrons travel around the nucleus in one of several orbits/shells/ nrg levels • Principal energy level is designated by a quantum number (n) • Quantum number is the same as period on the periodic table

Noble Gas Halogen Group Alkali Metal Alkali Earth Metal Period 2. 4

ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. • first shell maximum of 2 electrons • second shell maximum of 8 electrons • third shell max of 18 electrons • fourth shell max of 32 electrons

ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1. Electronic Configuration 2. Electron dot diagrams

ELECTRONIC CONFIGURATION With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example; Nitrogen configuration = 2 , 5 2 in 1 st shell 14 5 in 2 nd shell 2 + 5 = 7 7 N

ELECTRONIC CONFIGURATION Write the electronic configuration for the following elements; a) 20 40 Ca 11 b) 23 2, 8, 8, 2 d) 17 Cl 35 2, 8, 7 Na 8 c) 16 O 2, 8, 1 14 e) Si 28 2, 8, 4 2, 6 f) 5 B 11 2, 3

Valence • Valence shell is outermost occupied energy level and is the same as the period number in the periodic table • Valence electrons are the electrons in the outer energy level of an atom

DOT & CROSS DIAGRAMS With Dot & Cross diagrams elements and compounds are represented by Dots or Crosses to show electrons, and circles to show the shells. For example; X Nitrogen X X N XX X X N 7 14

DOT & CROSS DIAGRAMS Draw the Dot & Cross diagrams for the following elements; X 8 17 X a) O b) Cl 35 X 16 X X X X X Cl X X X O X X X X X

SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.