Studied by 3 people

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1

JJ Thomson

Used a cathode ray to discover the electron, plum pudding model,

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2

Rutherford

Gold foil experiment, discovered nucleus. Believed that an atom's volume was mostly empty space and the nucleus contributed to virually all of the mass but none of the volume

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3

What led Thomson to propose the negative particle was smaller than atoms

the charge to mass ratio for ht electron was 2000 greater than a positie hydrogen ion. which was the smallest known particle known at the time

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4

charge to mass ratio

Thomson measured this ratio for cathode rays and in doing so discovered the electron

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5

Greater mass in the charge to mass ratio

deflected the least

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6

Less mass inthe charge to mass ratio

deflected the most

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7

Cathod Ray Experiment

showed that all atoms contain tiny negatively charged subatomic particles or electrons. Atoms bent towards the positive plate

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8

atomic number

Number of protons and electrons in the atom

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9

mass number

the total number of protons and neutrons in the nucleus of an atom. Also the weighted average of naturally occurring isotopes.

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10

neutron number

mass number - atomic number

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11

anion

A negatively charged ion

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12

Cation

A positively charged ion

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13

mass spectrometry

Charged particles movingthrouhg a magnetic field are deflected form their orginal path based on their charge-to-mass ratio. Isotopes seperates, a detector records separations, percent abundanc and mass of each isotope.

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14

Problem with Rutherford's model

An electron orbiting is an accelerting charge, and accelerating charges radiates energy. If it looses energy, electrons collapse into the nucleus and matter wouldn't exist.

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15

Max Planck

Proposed that energy could be shown to behave like partickles in energy packets called quanta. energy could onlu be absorbed or emititted in who numbers of quanta

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16

Plank's Quantum Theory

Energy is quantized and can be emitted or absorbed only in whole number of discreet packets.

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17

Bright line spectrum

colors produced when electrons fall to a lower energy level and release energy in the form of light.

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18

Bohr's Postulates

Electrons in an atom exist in stationary states, each state was given an interger number called a quantum number

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19

As long as an electron moved in an allowed orbit, the electron (therefore the atom) did not radiote or absorb energy

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20

The electron could only mve from one allowed orbit to another if it absorbed or emmited an amount of energy exactly equal to the energy diffference between the two orbits.

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21

ground state

the lowest energy state of an atom(n=1)

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22

excited state

A state in which an atom has a higher potential energy than it has in its ground state(n=2,3...)

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23

Why is the bright line spectrum not continuous

Becaue the amount of energy associated with each eletron orbit is fixed, the difference in energy between those orbits are already fixed. An electron cannot orbit between two energy levels

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24

How did Bohrs Theory save Rutherford nuclea atomic Model

As long as an electron orbits the nucleus in an allowed orbit or stationary state, no energy is lost but that electron does not collapse.

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25

How does the work of Planck and Einstein contribute to Bohr;s theory about electron behaviour

if energy was quantized and could thus exist in only certain amounts and not others, perhaps orbiting electrons in atoms could only possess certain amounts of energy and not others corresponding to orbits of only certain sizes.

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26

Einstein

Used quantum theory to explain photoelectric effect

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27

De Broglie

Scientist who suggested that all moving particles had a wave motion associated with them. Wave particle duality

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28

De broglie's hypothesis

The only allowed orbits for electrons are those whose size and therefore energy allows for a standing electron wave to be maintained

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29

Planck's constant

Damn Tiny

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30

De Broglie's equation

Wavelength of a particle. As wavelength increase, mass decreases. As mass increases, wavelength decreases. Inversely proportional

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31

Why did Bohr's atomic model need to be replaced?

Worked for hydrogen but it was unable to predict the line spectra emitted by multi electron atoms

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32

What concept about the atomic structure was Bohr's theory responsible for

Energies of electrons are quantized and explained eemmission line spectra were produced

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33

Why can wave nature be ignored when looking at macroscopic particles

Wave nature can be ignored when looking at macroscopic particles because the value of Planck's constant is so incredibly tiny, virtually any particle big enough for us to see has a wavelength so incredibly small that it can't even be measured, making it irrelevant.

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34

Heisenberg

Uncertainty Principle, it is impossible to know exactly both the velocity and the position of a particle at the same time. The more certain of one measure, the less certain you are of the other (the more certain of position, the less certain of where it is travelling)

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35

Why is it imposible to find and predict electron behaviour.

In order to see an electrion we need a form of illumination radition to bounce off of the electron. One photo strikign the electron is enough to blow it away which is why it is impossible

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36

Schrodinger

Equation gives the porbablit of find those electrons within a region of space around the nucleus

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37

Atomic Orbitals

regions in 3D space around a nucleus where electrons of particular energy are most likely to be found

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38

First part of Quantum Theory

1.The energies of electron in atoms are quantized because of their wave nature. This related to the idea that only certain allowed energy states associated with standing electron waves can exist (Bohr, De Broglie)

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39

Second Part of Quantum Theory

2.The Heisenerg uncertainty principle states that it is impossible to simultaneously both where an electron is and where it's going

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40

Third Part of Quantum Theory

3.Atomic orbitals are those regions in 3D space around a nucleus where electrons with a particular energy ar emost likely to be found

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41

Contrast Bohr's electron orbit with the quantum mechanical electron orbital

Boh's orbit was a definate path followed by an electron particle as it circled the nucleus. An orbital is a region around the nucleus where an electron with a particular amound of energy is most likely to be found.

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42

How ca de Broglie's theory be seen as the opposite of Planck's quantum theory

Planck insisted that energy, previously considered to be only a wave phenomenon, was quantized and thus

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43

possessed particle properties. De Broglie suggested that electrons, previously considered as being only particles,

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44

possessed wave properties.

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45

Contribution by Bohr

Niels Bohr was the first to establish the quantized energy states associated with electrons in atoms by

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46

applying Planck's quantum theory to atomic structure.

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47

Contribution by De Broglie

Louis de Broglie was the first to suggest that particles could possess a wave nature which supplied a reason for the allowed energy states of electrons in the Bohr model of the atom

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48

Contribution by Schrodinger

Was the first to mathematically treat the eectron as a wave as his wave equation supplied a region of 3D space around an atom's nucleus where electrons of particular energies were most likely to be located.

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49

Contribution by Heisenberg

Heisenberg's uncertainty principle forces us to accept that there is a limit to what can know for sure about

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50

matter at its most basic level, and forces us to accept probablity when describing electron behaviour

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51

quantum numbers

specify the properties of atomic orbitals and the properties of electrons in orbitals

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52

Principle quantum number

symbolized by n, indicates the main energy level occupied by the electron amd relative sizes of atomic orbitals

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53

second quantum number

L, related to the shpae of an atomic orbital. S, P, D, F

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54

magnetic quantum number

m, tells us the oreintation in space of a give atomic orbital. 3D Cartesian coordinate, x,y,z etc

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55

Total number of orbitals per energy level

n^2

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56

Node

A point of zero amplitude on a standing wave

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57

spin quantum number

It tells us the two possible elctron spins +1/2 or -1/2

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58

Pauli Exclusion Principle

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

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59

The maximum number of electrons that can exist in any energy level n is given by 2n^2

S<P<D<F

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60

Aufbau Principle

When filling orbitals, the lowest energy orbitals availbe are always filled first

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61

Hund's Rule

When orbitals of equal energy are being filled, the most stable configuration is the one iwth the maximum number of unpaire electrons with the same spin

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62

Isoelectronic

two species have the same number and configuration of electrons

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63

If hydrogen electron exists in a spherial orbital, why doesn't this mea the electron move around hte nucleus in a circle

The 1s orbital represents in 3D spcae where it is most likely related

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