Chapter 5: Atomic Structure

Front 1

In the early stages of its development, scientists/philosophers relied on _________.

Back 1

indirect evidence

Front 2

1. All matter is made up of tiny particles
2. Particles cannot be divided
3.He called these particles "atomos" meaning "invisible
460-370 BC

Back 2

Democritus

Front 3

1. Did NOT believe in atoms
2. Believed all matter is continuous
Four elements [earth, air, fire, water]
3. Accepted until 18th century when experimental evidence disproved his ideas
384-322 BC

Back 3

Aristotle

Front 4

1. Proposed his atomic theory in 1803
2. Experimental conclusions:
a. each element is composed of atoms
b.atoms of one element are all identical, atoms of different elements differ
c.atoms cannot be subdivided, created, or destroyed.
d. atoms combine to form compounds
e. in chemical reactions, atoms are combined, separated, or rearranged in simple, whole-number rations (Law of definite proportions).
1766-1844

Back 4

John Dalton

Front 5

Discovered electron in 1897, proposed "Plum Pudding Model"

Back 5

J.J. Thomson

Front 6

1. The nucleus of an atom has a concentrated (dense) mass of positive charge
2. Credited with discovering the existence of proton
1909-1911

Back 6

Ernest Rutherford

Front 7

Developed Bohr model of an atom; placed electrons in distinct energy levels around the nucleus
1913

Back 7

Niels Bohr

Front 8

Discovered the neutron
1932

Back 8

James Chadwick

Front 9

a. Energy levels proposed by Bohr are regions (or clouds) rather than orbits.
b. It is impossible to determine the exact location of an electron at any given time. However, its general location can be predicted.
c. 99.9% of an atoms mass is in the nucleus
d. Protons are positive; neutrons are neutral
e. Electrons have a negative charge and are located in energy levels (orbitals, regions) around the nucleus
f. An atom has the same # of protons and electrons when neutral in charge.

Back 9

Quantum Theory

Front 10

Actual Mass: 9.11 x10^ -28
Relative Mass: 0 (1/1840)
Charge?

Back 10

Electron (e-)

Front 11

Actual Mass: 1.673x10^ -24
Relative Mass: 1
Charge?

Back 11

Proton (p+)

Front 12

Actual Mass: 1.675x10^ -24
Relative Mass: 1
Charge?

Back 12

Neutron (n^o)

Front 13

Atomic # = ?

Back 13

# of protons, # of electrons in a neutral atom

Front 14

Atoms of the same element that have different masses
(Proton #s are the same, but neutron #s are different)

Back 14

Isotope

Front 15

Isotopes: 2 ways to write them

Back 15

Hyphen notation (word format)
symbol-mass # (protons + neutrons)
Symbol notation
mass #
symbol
atomic #

Front 16

(p+) + (n^o) = ?

Back 16

mass # (always a whole number)

Front 17

represented as an atomic mass unit or u. Determined using the isotope carbon-12. Carbon was assigned a mass of exactly 12 u's. The u's of an element are 1/12 the mass of a carbon-12 atom.

relative scale allows us to get rid of exponents

Back 17

relative atomic mass

Front 18

weighted average of all the isotopes of an element (g); the mass listed on periodic table

Back 18

average atomic mass

Front 19

Row # = ?

Back 19

Periods (# of energy levels)

Front 20

Column # = ?

Back 20

s & p blocks- give # of valence e-

Front 21

1)e- enter orbitals of lowest energy levels first
2) as the # of e- builds, new elements are built

Back 21

Aufbau Principle

Front 22

1) atomic orbitals contain a max of 2 e-
2) e- within an orbital must have opposite spin

Back 22

Pauli Exclusion Principle

Front 23

1) orbitals of equal energy (within same sublevel) are each occupied by one e- before any one orbital gets a sencond e-
2) all e- in singly occupied orbitals must have the same spin

Back 23

Hund's Rule

Front 24

American chemist who invented electron dot notation

Back 24

Gilbert Lewis

Front 25

exceptional electron configurations include ?

where are most located?

Back 25

Ag, Cr, Cu, Mo, Pd, Au