The Nuclear Atom

Front 1

Radioisotope

Back 1

Radioactive isotopes wherein the nucleus of an isotope is unstable, and therefore it decays and emits radiation in the form of gamma rays, alpha particles or beta particles

Front 2

Gamma Rays

Back 2

High density photons that are high in kinetic energy and require dense material to stop

Front 3

Beta Particle

Back 3

Electron Emitted From Electron Decay, Less Dense and High Energy

Front 4

Alpha Particle

Back 4

2 neutrons and 2 protons, dense and easy to stop

Front 5

Iodine-131

Back 5

Thyroid absorbs I-131, Iodine emits radiation in the form of gamma rays, a single photon emission computed tomography scanner detects radiation, 3-D maps thyroid showcasing any cancers, tumours or abnormalities

Front 6

Iodine-125

Back 6

Used to treat prostate cancer and brain tumours, emits gamma rays.

Front 7

Cobalt-60 (Co-60)

Back 7

Common in radiotherapy for cancer treatment

Front 8

Positron Emission Tomogrpahy

Back 8

Tissues absorb radioactive isotopes of elements and scanners recreate images and models of them

Front 9

Gamma Rays, X-Rays, UV

Back 9

Shorter Wavelength, High Frequency, High Energy

Front 10

Infrared, Microwaves, Radiowaves,

Back 10

Long Wavelength, Low Energy, Lower Frequency

Front 11

Lambda

Back 11

Wavelength

Front 12

E

Back 12

Energy

Front 13

V (nu)

Back 13

Frequency of wavelength per unit of time

Front 14

C

Back 14

Speed of light (3.00 x 10^8 ms^-1)

Front 15

Formula for c

Back 15

Lambda x frequency

Front 16

When is the relationship between frequency and wavelength

Back 16

Inversely proportional: High frequency means shorter wavelength

Front 17

What is the relationship between energy and wavelength

Back 17

Inversely proportional: High energy, low wavelength

Front 18

Continuous Spectrum

Back 18

Seen as white light, includes all the wavelengths on the visible spectrum

Front 19

Emission Spectrum

Back 19

Contains all the wavelengths of light that are not absorbed (and are thus emitted) by an atom

Front 20

Absorption Spectrum

Back 20

All wavelengths of light that are absorbed by an element

Front 21

Formula for Energy

Back 21

E=hv or hc/lambda

Front 22

What is Planck’s Constant

Back 22

6.63 x 10^-34 Js

Front 23

Convergence

Back 23

As the energy levels increase in energy they get closer together

Front 24

What’s the limit of convergence

Back 24

Continuum, beyond this continuum, once electrons pass this continuum, they have too much energy to be affected the gravitational force of the nucleus and can leave the atom as a free electron

Front 25

Balmer Series (3 to 2)

Back 25

656nm, red

Front 26

Balmer Series (4 to 2)

Back 26

486nm, blue-green

Front 27

Balmer Series, 5 to 2

Back 27

434nm, blue

Front 28

Balmer Series, 6 to 2

Back 28

410nm, 6 to 2

Front 29

Paschen Series

Back 29

The Paschen series is the set of emission lines from atomic hydrogen gas, due to electrons descending from an electron shell of number n greater than 3 down to that of n = 3: Wavelength emitted is infrared

Front 30

Lyman Series

Back 30

In physics and chemistry, the Lyman series is a hydrogen spectral series of transitions and resulting ultraviolet emission lines of the hydrogen atom as an electron goes from n ≥ 2 to n = 1: Energy emitted is in the UV spectrum

Front 31

What is the quantum mechanical model?

Back 31

A model of the atom that describes the possibility of placing electrons within an atom by describing the principal energy level, energy level, orbital (arbitrary level), and spin of electrons. Derived from the Schrodinger equation.

Front 32

What is Heisenburg’s Uncertainty Principle

Back 32

The idea that all substances behave like a wave and a particle, and that the values for certain pairs of physical quantities of a particle, such as position, x, and momentum, p, can not be predicted from initial conditions and will always be uncertain.

Front 33

Schrodinger Equation

Back 33

The Schrödinger equation is a linear partial differential equation that governs the wave function of a quantum-mechanical system: allows for the prediction of electrons in orbitals.

Front 34

Energy Levels

Back 34

In chemistry, the principal energy level of an electron refers to the shell or orbital in which the electron is located relative to the atom's nucleus. This level is denoted by the principal quantum number n.

Front 35

Atomic Orbitals (Sublevels)

Back 35

Atomic orbitals, as determined by Schrödinger’s equation, describe the probability of finding a given electron of an atom in a given region of space. Any orbital can hold 2 electrons with complementary spins. They are denoted by the letter L.

Front 36

The energy of sublevels, suborbitals

Back 36

S

Front 37

Magnetic Quantum Number

Back 37

Specifies the orientation in space of an orbital of a given energy (n) and shape (l). Denoted by +/-

Front 38

Hund’s Rule

Back 38

Hund's rule states that: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied.

Front 39

Pauli’s Exclusion Principle

Back 39

Pauli's Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers. In other words, (1) no more than two electrons can occupy the same orbital and (2) two electrons in the same orbital must have opposite spins

Front 40

Aufbau Principle

Back 40

It states that electrons are filled into atomic orbitals in the increasing order of orbital energy level. According to the Aufbau principle, the available atomic orbitals with the lowest energy levels are occupied before those with higher energy levels.

Front 41

AMU

Back 41

1.66 x 10^-24 g

Front 42

t1/2 of C-14

Back 42

5730 years

Front 43

Mass Spectrometer

Back 43

Apparatus used to find the mass and abundance of isotopes

Front 44

Group 15 Elements

Back 44

Pnictogens

Front 45

Period

Back 45

A period in the periodic table is a row of chemical elements. All elements in a row have the same number of electron shells, highest occupied energy level and same principle quantum number

Front 46

Main Group Elements

Back 46

1-2,13-18

Front 47

Zeff

Back 47

Effective Nuclear Charge: net positive charge experienced by valence electrons

Front 48

Atomic Radius

Back 48

The atomic radius of a chemical element is a measure of the size of its atom, usually the mean or typical distance from the center of the nucleus to the outermost isolated electron.

Front 49

Ionization Energy

Back 49

Ionization Energy is the minimum energy required to remove one mole of the most loosely bound electron of an isolated gaseous atom, positive ion, or molecule.

Front 50

Electron Affinity

Back 50

The electron affinity refers to the amount of energy released when a mole of electrons electron attaches to a neutral atom or molecule in the gaseous state to form an anion.

Front 51

Electronegativity

Back 51

Electronegativity, symbolized as χ, is the tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond.