Atoms

Let’s Review! Matter is… Anything that has mass and takes up space All matter is made of elements – substances that cannot be broken down And elements are made of – ATOMS! The atom is the SMALLEST unit of matter

ATOMS ARE EVERYWHERE! They are in the air you breathe, the chair you’re sitting on, and the clothes you’re wearing.

Section 1 Atomic Theory

Democritus (400 B.C.) Ancient Greece Theorized that matter could not be divided infinitely, you had to reach a smallest piece Atomos: indivisible or can’t be cut (becomes atom) Since this was ancient Greece, he had no proof and few people believed it.

John Dalton (1808) People were more accepting of his atomic theory because he had evidence His theory: all matter is made of atoms (small solid spheres - Billiard Ball Model) atoms are indivisible and indestructible atoms of the same element are exactly alike atoms of different elements are different compounds are formed by joining two or more atoms

John Dalton (1808) Dalton’s theory of atoms supported the Law of Definite Proportions Law of Definite Proportions: A chemical compound is made up of the same percentage of elements. This means that elements combine is whole number ratios when they react Water is always made up of 2 parts hydrogen to 1 part oxygen giving us H2O This is true for all chemical compounds!

Sir J. J. Thomson (1897) Conducted experiments with cathode rays that showed atoms could be divided His experiments showed negatively charged particles that came from inside atoms These are electrons!

Sir J. J. Thomson (1897) A cathode ray has two metal plates at the end of a vacuum tube Cathode, which has a negative charge Anode, which has a positive charge When voltage is applied across the plates, a glowing beam comes from the cathode and strikes the anode Since it was a vacuum tube (a tube with all the air vacuumed out), he could see that the beam came from the negative end of the cathode

Sir J. J. Thomson (1897) Successfully separated negative particles, but could not separate the positive particle Proposed that electrons are spread throughout the atom like blueberries in a muffin Called the Plum Pudding Model

Ernest Rutherford (1911) Discovered the nucleus using the gold foil experiments Fired positively charged particles at a sheet of gold foil Most went through unaffected, some bounced away This suggested that an atom’s positive charge was concentrated at the center of the atom This is the NUCLEUS!

Ernest Rutherford Stated that electrons are scattered around the atom with mostly empty space between them and the nucleus Compared to the atom, the nucleus is very small, like a marble on a football field!

Niels Bohr (1913) Proposed that electrons are arranged in circular energy levels around the nucleus Like planets orbiting the sun When electrons gain energy, they “jump” from a lower level to a higher, a loss of energy causes it to “fall” from a higher level to a lower

Erwin Schrodinger (1926) & James Chadwick (1932) His model does not define the exact path of an electron, but predicts its location Shows the nucleus surrounded by an electron “cloud” Chadwick discovered the neutron, which has no electrical charge, and the same mass as the proton

Modern Atomic Theory Today we know that atoms have A nucleus with And Protons and Neutrons And Electrons that orbit the nucleus and behave like waves on a string

Section 2 Atomic Structure

Subatomic Particles Subatomic: lower (or smaller) than an atom Protons and electrons have an electrical charge Protons: positive charge Neutrons: no charge (neutral) Electrons: negative charge

Mass and Volume The nucleus makes up 99.99% of the mass of the atom. However, the nucleus is 1/100,000 of the volume of an atom. The volume is determined by the electron cloud.

Mass and Volume Since subatomic particles are so small they cannot be measured in grams They are measured in atomic mass units or amu 1 amu = 1.61x10-24 g 0.00000000000000000000000161! 1 g is about the mass of a paper clip!

protons p+ positive nucleus 1 amu neutrons no neutral electrons e- name symbol charge location mass protons p+ positive nucleus 1 amu neutrons no neutral electrons e- negative electron shell .0006 amu

Electrons Fill energy levels around the nucleus Each level only holds so many electrons Valence Electrons: electrons that fill the outermost level

Atomic Number Atomic Number: number of protons found in the nucleus The # of protons in an atom is unique to each element and is how we identify an element – it NEVER CHANGES In a neutral atom, the number of electrons equals the number of protons – meaning there will be NO overall charge on the atom

Atomic Mass Atomic Mass: total mass of all atom’s components Mass Number: total number of neutrons and protons in an atom subtract atomic # from atomic mass to find number of neutrons The number of neutrons in an atom can change, which means the atomic mass can change!

Let’s Try! Atomic # Mass Protons Neutrons Electrons 14 28 Atomic # 3 7 4 Atomic # Mass Protons Neutrons Electrons 9 19 10 Atomic # Mass Protons Neutrons Electrons 90 232 142

Isotopes Isotope: atoms of the same element with a different number of neutrons, and therefore different masses. Normally 1-2 stable isotopes for an element, and the atomic mass of the most common isotope is listed in your periodic table All others are unstable (they fall apart) through radioactive decay

Representing Isotopes Isotopes have the SAME chemical and physical properties. Isotopes can be represented in a number of ways called isotope notation Element symbol or name with atomic mass Element symbol with atomic number and atomic mass