Integrated Science 11: Chapter 8 Test

Front 1

Atom

Back 1

a tiny unit of matter nature's basic particle

Front 2

Democritus

Back 2

believed matter was discontinuous

Front 3

Democritus

Back 3

believed in indivisible, smallest particles possible that he named atoms

Front 4

Aristotle

Back 4

believed matter was continuous

Front 5

False statement from Dalton's Atomic Theory

Back 5

Indivisible tiny particles called atoms make up all matter.

Front 6

False statement from Dalton's Atomic Theory

Back 6

All atoms of the same element are exactly alike in shape & mass.

Front 7

True statement from Dalton's Atomic Theory

Back 7

The atoms of different elements differ from one another.

Front 8

True statement from Dalton's Atomic Theory

Back 8

Atoms chemically combine in definite whole number ratios to form chemical compounds.

Front 9

True statement from Dalton's Atomic Theory

Back 9

Atoms are neither created or destroyed in chemical reactions.

Front 10

JJ Thompson

Back 10

discovered electron by using a cathode ray tube.

Front 11

Robert Millikan

Back 11

helped to discover the electron’s mass, along with the value of

Front 12

First suggestion made after electron discovery

Back 12

there must be something positive

Front 13

Second suggestion made after electron discovery

Back 13

there must be something more massive in the atom because the electron has such a small mass.

Front 14

JJ Thompson

Back 14

proposed “Plum Pudding” model.

Front 15

Description of the Plum Pudding model

Back 15

an atom was a blob of positively charged matter where the electrons were stuck like “raisins” in plum pudding.

Front 16

Rutherford

Back 16

shot alpha particles at a piece of gold foil.

Front 17

Rutherford

Back 17

concluded that an atom must have a tiny, but massive, positively charged nucleus surrounded by the electrons.

Front 18

Rutherford

Back 18

concluded from measurements that the atom is mostly empty space.

Front 19

Rutherford

Back 19

later discovered the proton (p+) as the particle of positive charge in the nucleus.

Front 20

Rutherford

Back 20

developed an atomic model to describe it structure, commonly called the solar system model.

Front 21

James Chadwick

Back 21

discovered the neutron (n0).

Front 22

Atomic Number

Back 22

the number of protons in an atom

Front 23

Isotope

Back 23

atoms of the same number of protons (same element) with a different number of neutrons.

Front 24

Carbon-12 isotope

Back 24

the standard which is assigned exactly 12.00 atomic mass units (u).

Front 25

Mass Number

Back 25

the number of protons and neutrons in an isotope.

Front 26

Atomic Weight

Back 26

a weighted average of the isotopes based on their mass compared to Carbon-12 and their relative abundance in nature

Front 27

Niels Bohr

Back 27

tied together the quantum idea with Rutherford’s model to develop his model of the hydrogen atom.

Front 28

Bohr’s Theory

Back 28

believed that electrons travel in specific, allowed orbits

Front 29

Bohr's Theory

Back 29

believed that when in these orbits, the electrons do not give off radiation

Front 30

Bohr's Theory & the Quantum Theory

Back 30

believed that an electron gains or loses energy only when moving from one orbit to the next

Front 31

Changing orbits is referred to as

Back 31

quantum leaps

Front 32

Ground State

Back 32

lowest energy state

Front 33

Excited States

Back 33

any states above ground

Front 34

Louis DeBroglie

Back 34

developed the theory that electrons travel as a wave

Front 35

Schrödinger

Back 35

derived an equation that mathematically proved the Quantum theory.

Front 36

Heisenberg Uncertainty Principle

Back 36

allowed the quantum theory to be widely accepted

Front 37

Orbitals

Back 37

a three-dimensional shell, a fuzzy region of space where it is probable that you will find the electron, a probability cloud

Front 38

Shapes and sizes of orbitals depend on

Back 38

the energy of the electrons contained inside them.

Front 39

Principle Quantum Number

Back 39

describes main energy level

Front 40

Principle Quantum

Back 40

designated using the following sybolism n = 1, 2, 3, etc.

Front 41

Angular Momentum Quantum Number

Back 41

defines energy sublevels within the main energy levels

Front 42

Angular Momentum Quantum Number

Back 42

indicates the shape of the orbital, identified with a letter

Front 43

s, p, d, f

Back 43

letters used to describe the shape of the orbitals

Front 44

Magnetic Quantum Number

Back 44

describes an orbital’s orientation in space

Front 45

Spin Quantum Number

Back 45

describes an electron’s spin orientation

Front 46

Pauli’s Exclusion Principle

Back 46

No two electrons in an atom can have the same four quantum numbers

Front 47

Electron Configuration

Back 47

the arrangement of electrons in orbitals

Front 48

Phosphorus - 1s22s22p63s23p3

Back 48

example of an electron configuration

Front 49

Periods

Back 49

Horizontal rows numbered 1-7

Front 50

Groups

Back 50

Also referred to as families, vertical columns

Front 51

Column 1

Back 51

alkali metals

Front 52

Column 2

Back 52

alkaline earth metals

Front 53

Columns 3 to 12

Back 53

transition metals (elements)

Front 54

Column 17

Back 54

halogens

Front 55

Column 18

Back 55

noble gases

Front 56

Columns 1, 2, 13 to 18

Back 56

representative elements

Front 57

Rare Earth Elements

Back 57

inner transition elements

Front 58

Two Series contained in the Rare Earth Elements

Back 58

Lanthanide Series & Actinide Series

Front 59

The columns of representative elements are numbered according to

Back 59

how many outer shell electrons there are.

Front 60

Electron dot notations are used to represent

Back 60

these outer shell electrons.

Front 61

Valence Electrons

Back 61

outer shell electrons used in bonding

Front 62

Ion

Back 62

an atom that gains or loses electrons

Front 63

Metals

Back 63

tend to lose electrons to form positive ions, cations

Front 64

Nonmetals

Back 64

tend to gain electrons to form negative ions, anions