Chemistry: Unit 3 Vocab

Front 1

Atomic Structure

Back 1

Refers to the identity & arrangement of smaller particles within atoms

Front 2

Nucleus

Back 2

small region, contains positively charged protons (p+) and neutral neutrons (n0)

Front 3

Electron cloud

Back 3

comparatively large in volume, contains negatively charged electrons (e- or -10e)

Front 4

Electrons

Back 4

discovery came from looking at relationship between electricity and matter by way of cathode ray tubes

Front 5

JJ Thomson

Back 5

found that all cathode rays were composed of identical, negatively charged particles – later termed electrons.

Front 6

JJ Thomson

Back 6

found that it had a very large charge when compared to its very tiny mass.

Front 7

Robert Millikan

Back 7

In 1909, continuing from Thomson's experiments, he found the mass of an electron to be 9.1 x 10-28 g by using the Oil Drop Experiment.

Front 8

One Inference about Atomic Structure Made from Millikan's work

Back 8

1.) Because atoms are electrically neutral, they must contain a positive charge to balance these electrons.

Front 9

One Inference about Atomic Structure Made from Millikan's work

Back 9

2.) Because the mass of an electron is so small, atoms must contain additional particles that account for most of their mass.

Front 10

Nucleus

Back 10

Found by Ernst Rutherford in 1911 (with Ernest Marsden & Hans Geiger)

Front 11

Nucleus

Back 11

the positively charged, dense central portion of the atom that contains nearly all of its mass but takes up only an insignificant fraction of its volume

Front 12

How the Size of Atoms are Measured

Back 12

Measurements are made by looking at radius, Measured from the middle of the nucleus to the edge of the electron cloud

Front 13

Units used for measuring atoms

Back 13

picometers, ranging from 40 – 270 pm

Front 14

Mass of p+

Back 14

1.673 x 10-24 g

Front 15

Mass of n0

Back 15

1.675 x 10-24 g

Front 16

Nuclear forces

Back 16

short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together

Front 17

Atomic Number

Back 17

the number of protons in the nucleus of each atom of that element

Front 18

Atomic Number

Back 18

It identifies an element.

Front 19

On periodic table, elements are placed in order of

Back 19

increasing atomic numbers.

Front 20

Isotopes (nuclide)

Back 20

atoms of the same element with different masses

Front 21

Protium

Back 21

1 p+, 1 e-, 0 n0

Front 22

Deuterium

Back 22

1 p+, 1 e-, 1 n0

Front 23

Tritium

Back 23

1 p+, 1 e-, 2 n0

Front 24

Mass Number

Back 24

total number of protons and neutrons in the nucleus of an isotope

Front 25

Hyphen notation

Back 25

name with mass number

Front 26

Nuclear symbol

Back 26

symbol for the element, with the mass number being superscript and atomic number being subscript

Front 27

Relative Atomic Mass

Back 27

the mass of an atom expressed in atomic mass units

Front 28

Atomic Mass Unit (amu)

Back 28

1/12 the mass of a carbon-12 atom or 1.6605402 x 10-24 g

Front 29

Carbon-12 isotope

Back 29

chosen as the standard for atomic mass and all other atoms are compared to this. It is assigned exactly 12 amu.

Front 30

Average Atomic Mass

Back 30

the weighted average of the atomic masses of the naturally occurring isotopes of an element

Front 31

Natural radioactivity

Back 31

the spontaneous emission of particles or energy from an atomic nucleus as it disintegrates

Front 32

Henri Becquerel

Back 32

discovered invisible radiation

Front 33

The Three Kinds of Radioactivity

Back 33

Alpha (α) particle: nucleus of a He atom, 2 protons & 2 neutrons
Beta (β) particle: high energy electron
Gamma (γ) ray: electromagnetic radiation of very short wavelength

Front 34

Radioactive decay

Back 34

the natural spontaneous disintegration or decomposition of a nucleus

Front 35

Nucleons

Back 35

another term for protons and neutrons

Front 36

Alpha (α) emission

Back 36

the release of an alpha particle from a disintegrating nucleus

Front 37

Beta (β) emission

Back 37

the release of a beta particle from a disintegrating nucleus

Front 38

Beta emission

Back 38

a neutron changes to a proton by emitting the electron

Front 39

Beta particles

Back 39

they are more penetrating and can travel several hundred centimeters in air

Front 40

Alpha particles

Back 40

can travel 2-12 cm through the air

Front 41

Gamma (γ) emission

Back 41

the release of a high-energy burst of electromagnetic radiation from an excited nucleus

Front 42

Gamma emission

Back 42

Nucleons don’t change in this type of emission

Front 43

Gamma rays

Back 43

most penetrating, can pass through a person - can be stopped by a piece of lead close to the source

Front 44

Gamma rays

Back 44

used to kill many parasites and bacteria in the foods we eat

Front 45

Radioactive Decay Series

Back 45

a process continuing through a series of decay reactions until a stable nucleus is achieved

Front 46

Half-Life

Back 46

the time required for 1/2 of the unstable nuclei to decay

Front 47

Instruments to measure radiation

Back 47

Photographic film & Geiger counters

Front 48

Unit of Radiation

Back 48

curie (Ci), commonly picocuries/ Liter (pCi/L)

Front 49

SI Unit for Radiation

Back 49

Becquerel (Bq), 1 nuclear disintegration / second

Front 50

Background radiation

Back 50

low amounts of radiation from natural sources

Front 51

Mass Defect

Back 51

The difference between the mass of the nucleons and the nucleus

Front 52

Binding energy

Back 52

energy released when nucleus is formed, or energy absorbed when breaking nucleus, can be calculated from the mass defect

Front 53

Nuclear fission

Back 53

splitting a massive nucleus into more stable, less massive nuclei with the release of energy

Front 54

Chain reaction

Back 54

a reaction where the products are able to produce more reactions in a self-sustaining series

Front 55

Critical mass

Back 55

the mass and concentration of nuclei that is sufficient to sustain a chain reaction

Front 56

Nuclear fusion

Back 56

less massive nuclei coming together to form more stable and more massive nuclei with the release of energy

Front 57

1 Problem using nuclear fusion

Back 57

1.) temperature - nuclei repel one another unless they have enough energy, approximately 100 million °C

Front 58

1 Problem using nuclear fusion

Back 58

2.) density - heavy hydrogen nuclei needed, 1 x 1014 /cm3

Front 59

1 Problem using nuclear fusion

Back 59

3.) time - must be sustained at a minimum of 10 atm at this density

Front 60

Characteristics of Bohr's Atomic Model

Back 60

specific allowed paths

Front 61

Characteristics of Bohr's Atomic Model

Back 61

lowest energy orbit was closest to the nucleus, electron must gain energy to move up to another orbit

Front 62

Characteristics of Bohr's Atomic Model

Back 62

each orbit or path had a fixed amount of energy

Front 63

Characteristics of Bohr's Atomic Model

Back 63

connected atom's electrons with line spectra, when electron moved down from higher to lower energy levels, emitted photon

Front 64

Problem with this Bohr's Atomic Model

Back 64

can only be proven with hydrogen atom

Front 65

Problem with this Bohr's Atomic Model

Back 65

did not explain chemical behavior of atoms

Front 66

Louis DeBroglie

Back 66

suggested that electrons have wave-like properties

Front 67

Heisenberg Uncertainty Principle

Back 67

it is not possible to measure the velocity, which is heavily dependent on the energy, and position of an electron at the same time.

Front 68

Edwin Schrödinger

Back 68

developed mathematical equation that described how electrons move around the nucleus as waves

Front 69

Schrodinger's Equation

Back 69

This equation worked for all atoms, not just Hydrogen as Bohr's did. This equation gives the probability of where to find the electron in the atoms.

Front 70

Schrodinger

Back 70

developed different possible cloud shapes that the electrons occupy.

Front 71

Quantum theory

Back 71

describes mathematically the wave properties of electrons and other small particles

Front 72

Orbitals

Back 72

a three dimensional region around the nucleus that indicates the probable location of an electron

Front 73

Shape & Size of the clouds depend on

Back 73

the energies of the electrons that occupy them.

Front 74

Quantum numbers

Back 74

numbers that specify the properties of atomic orbitals and the electrons in those orbitals

Front 75

The first 3 quantum numbers

Back 75

indicate the region occupied by the orbital.

Front 76

The 4th quantum number gives

Back 76

the fundamental state of the e-.

Front 77

Principal Quantum Number (n)

Back 77

indicates the main energy levels surrounding a nucleus (sometimes referred to as shells)

Front 78

As n ↑, distance from nucleus ↑ and...

Back 78

the energy ↑

Front 79

Angular Momentum (Orbital) Quantum Number (l)

Back 79

indicates the shape of an orbital

Front 80

s

Back 80

spherical shape,sharp

Front 81

p

Back 81

peanut shape, principal

Front 82

d

Back 82

double peanut shape, diffuse

Front 83

f

Back 83

flower shape, fundamental

Front 84

The number of possible shapes you can have in an energy level is equal to

Back 84

the principal quantum number.

Front 85

Magnetic Quantum Number (m)

Back 85

indicates the orientation of an orbital around a nucleus

Front 86

Orientations for s

Back 86

1 orientation

Front 87

Orientations for p

Back 87

3 orientations (px, py, pz)

Front 88

Orientations for d

Back 88

5 orientations

Front 89

Orientations for f

Back 89

7 orientations

Front 90

Number of orbitals =

Back 90

(principal quantum number)2

Front 91

Spin Quantum Number

Back 91

indicates two possible states of an electron in an orbital

Front 92

two possible values of Spin Quantum Number

Back 92

(+1/2, -1/2)

Front 93

the max number of electrons in each main energy level =

Back 93

2n2

Front 94

Electron Configurations

Back 94

the arrangement of electrons in atoms, where each element has its own distinct configuration

Front 95

Aufbau principle

Back 95

electron will occupy the lowest-energy orbital that can receive it

Front 96

Hund's Rule

Back 96

orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin

Front 97

Pauli Exclusion Principle

Back 97

no two electrons in the same atom can have the same set of four quantum numbers

Front 98

Orbital Notation

Back 98

Lines are labeled with the principal quantum number, n, and the subshell letter.

Front 99

Long Hand Electron-Configuration Notation

Back 99

This method eliminates the lines and arrows. To designate the number of electrons, a superscript is added.

Front 100

Short Hand Electron-Configuration Notation

Back 100

(also called Noble-Gas Notation) This is a method to condense all of the inner-shell electrons.

Front 101

Electron-Dot Notation

Back 101

This method only shows electrons in the outermost energy level. These electrons are also called... valence electrons.

Front 102

Mendeleev

Back 102

developed the first periodic table

Front 103

Mendeleev

Back 103

put in order of increasing atomic mass

Front 104

Mendeleev

Back 104

his table empty spaces left for undiscovered elements

Front 105

Mendeleev

Back 105

many accepted his table once predictions proven true

Front 106

Moseley

Back 106

his work lead to periodic table being re-organized

Front 107

Moseley

Back 107

he put in order of increasing atomic number

Front 108

Periodic Law

Back 108

the physical and chemical properties of the elements are periodic functions of their atomic numbers

Front 109

Periodic Table

Back 109

an arrangement of the elements in order of their atomic numbers so thatelements with similar properties fall in the same columns

Front 110

Main Group Elements

Back 110

s & p block elements

Front 111

Noble Gases

Back 111

normally unreactive (inert), except for He w/2 e-, all others have 8 e-,Lord Rayleigh & William Ramsay discovered many and placed new column on table for them

Front 112

Rare-Earth Elements

Back 112

separated from table to save space

Front 113

Lanthanide Series

Back 113

not really rare, similar properties to one another

Front 114

Actinide Series

Back 114

all reactive & unstable, first 4 found on Earth – others known as laboratory-made elements

Front 115

Alkali Metals

Back 115

most active metals

Front 116

Alkali Metals

Back 116

only 1 e- in outer shell

Front 117

Alkali Metals

Back 117

rarely found in nature as free elements

Front 118

Alkali Metals

Back 118

stored under oil or kerosene

Front 119

Alkali Metals

Back 119

soft, silvery-white, shiny

Front 120

Alkaline-Earth Metals

Back 120

active metal and rarely found free

Front 121

Alkaline-Earth Metals

Back 121

harder, stronger, and higher melting points than alkalis

Front 122

Transition Metals

Back 122

includes Lanthanide & Actinide Series

Front 123

Transition Metals

Back 123

can lose e- below the outermost shell

Front 124

Transistion Metals

Back 124

have tendency to share electrons

Front 125

Halogens

Back 125

most active nonmetals

Front 126

Halogens

Back 126

most commonly combine with a metal to form salts

Front 127

Atomic Radii

Back 127

size of an atom, determined by the electron cloud

Front 128

Atomic Radii

Back 128

half the distance between the nucleus of identical atoms joined in a molecule

Front 129

Reason for trend across periodic table for atomic radii

Back 129

More positive nucleus

Front 130

Trends for atomic radii

Back 130

Across: decrease
Down: increase

Front 131

Ionization Energy

Back 131

the energy that must be supplied when an electron is removed from an atom

Front 132

Ionization energy can be represented as

Back 132

A + energy → A+ + e-

Front 133

Ion

Back 133

an atom or group of atoms that has a positive or negative charge

Front 134

A+

Back 134

cation

Front 135

A-

Back 135

anion

Front 136

Ionization

Back 136

is a process that results in the formation of an ion.

Front 137

First Ionization Energy (IE1)

Back 137

specifically, the energy required to remove one electron from an atom

Front 138

Units for Ionization Energy

Back 138

kJ/mol

Front 139

Noble gases

Back 139

have the highest IE

Front 140

Second IE

Back 140

second electron removed, already positive so many times more difficult.

Front 141

Third IE

Back 141

third electron removed

Front 142

Trends for Ionization Energies

Back 142

Across: increase
Down: decrease

Front 143

Electron Affinity

Back 143

the energy change that occurs when an electron is acquired (or gains) by a neutral atom

Front 144

A + e- → A- + energy

Back 144

Representation when atoms readily take electrons.

Front 145

A + e- + energy → A-

Back 145

Representation when atoms are forced to gain electrons. These tend to be unstable and lose these e- spontaneously.

Front 146

Halogens

Back 146

have the highest electron affinity.

Front 147

Alkaline-earth metals and noble gases

Back 147

are the least likely to gain electrons.

Front 148

Trends for Electron Affinity

Back 148

Across: increase
Down: decrease

Front 149

Ionic Radii

Back 149

the radii once an ion is formed

Front 150

Formation of a cation

Back 150

decrease in the radius

Front 151

Cations

Back 151

are smaller than their neutral counterparts.

Front 152

Formation of an anion

Back 152

increase in the radius

Front 153

Anions

Back 153

are larger than their neutral counterparts.

Front 154

Electronegativity

Back 154

a measure of the ability of an atom in a chemical compound to attract electrons

Front 155

Valence Electrons

Back 155

electrons available to be lost, gained, or shared in the formation of chemical compounds

Front 156

Linus Pauling

Back 156

developed a scale for electronegativity

Front 157

Fluorine

Back 157

is most electronegative element, and arbitrarily assigned 4.0 on scale.

Front 158

Trends for Electronegativity

Back 158

Across: increase
Down: decrease

Front 159

Democritus

Back 159

coined the term atom for nature's basic particle

Front 160

Aristotle

Back 160

did not believe in idea of atoms

Front 161

Aristotle

Back 161

thought matter was continuous

Front 162

Aristotle

Back 162

Thought matter was made out of the four elements - Earth, Wind, Fire, Water

Front 163

alchemy

Back 163

the practical pursuit of the transmutation (transformation) of elements into one another

Front 164

Law of Multiple Proportions

Back 164

if two or more different compounds are composed of the saqme two elements, the masses of the second element combined with a certain mass of the first element can be expressed as ratios of small whole numbers

Front 165

John Dalton

Back 165

expanded on Democritus' work

Front 166

John Dalton

Back 166

proposed an atomic theory in 1801

Front 167

Wave/Particle Duality

Back 167

when material has certain wave-like properties and certain particle-like properties

Front 168

Electromagnetic Radiation

Back 168

a form of energy that exhibits wavelike behavior as it travels through space

Front 169

Electromagnetic Spectrum

Back 169

all electromagnetic radiation arranged according to increasing wavelength

Front 170

Wavelength

Back 170

the distance between corresponding points on adjacent waves

Front 171

Frequency

Back 171

the number of waves that passes a given point in a specific amount of time, usually one second

Front 172

Photoelectric effect

Back 172

a phenomena where light at a particular frequency shines on certain metals, and photoelectrons are emitted

Front 173

Quanta

Back 173

energy in specific amounts

Front 174

Quantum of light (energy)

Back 174

a finite quantity of energy that can be gained or lost by an atom

Front 175

Photon

Back 175

term used to describe an individual quantum of light

Front 176

Emission spectrum

Back 176

a spectrum of lines that represent the emission of photons with certain energies

Front 177

Balmer Series

Back 177

line emissions in the visible spectrum

Front 178

Lymen Series

Back 178

line emissions in the ultraviolet spectrum

Front 179

Paschen Series

Back 179

line emissions in the infrared spectrum

Front 180

Spectroscopy

Back 180

a process where a spectroscope is used to separate the light given off from substances into the line spectra