"AP Chemistry - Electronic Structure of Atoms and Quantum Theory Review"

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58 Terms

1

Wavelength

Distance between two points on a wave. The unit of measurement is METERS (but can be measured in cm and nm)

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2

Frequency

Number of complete waves that pass a point each second. The unit of measurement is Hz (1/s)/

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3

Speed of light = wavelength x frequency

c = λv

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4

Inverse

Frequency and wavelength are what proportion?

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5

Long wavelength

Low energy and frequency gives what type of wavelength?

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6

Short Wavelegth

High energy and frequency gives what type of wavelength?

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7

Speed of light constant

3.00 x 10^8 m/s

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8

Wavelength and frequency

Each type of energy has a different what?

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9

Energy = Planck's constant x frequency

E = hv

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10

Planck's constant

6.626 x 10^-34 J/s

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11

Photoelectric effect

When light shines on a metal surface causes the electrons to be ejected from the metal.

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12

Photons

Einstein said that light travels in energy packets or what?

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13

Planck's theory

Energy absorbed/released from atoms in certain minimum amounts (quanta). This is assumed.

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14

Electromagnetic Spectrum

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15

Continuous Spectrum

Spectrums from light containing all wavelengths.

<p>Spectrums from light containing all wavelengths.</p>
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16

Line Spectrum

A spectrum containing radiation of only specific wavelengths.

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17

Ground state

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

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18

Excited state

When the electron is in a higher energy level.

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19

Orbitals

The allowed wave functions of the hydrogen atom.

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20

Electron configuration

A simple way of writing down the locations of all of the electrons in an atom. It describes the distance, energy, and size of orbitals.

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21

Principal quantum number (n)

Represents the distance from the nucleus. Energy levels. (1-7)

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22

Angular quantum number (l)

Represents the shape of the orbital. Sublevels. (0=s, 1=p, 2=d, 3=f).

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23

Magnetic quantum number (ml)

Represents the orientation in space. Axis orientation. (-3, -2, -1, 0, 1, 2, 3)

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24

Spin quantum number (Ms)

Represents the electron spin. Arrow direction. (+1/2 or -1/2).

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25

Aufbau Principle

Electrons fill the orbitals in order of increasing energy.

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26

Spinning change.

What produces a magnetic field?

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27

Diamagnetic

Occurs when all e- are paired.

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28

Paramagnetic

Has one or more unpaired e-.

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29

Pauli Exclusion Principle

Two electrons occupying the same orbital must have opposite spins.

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30

Hund's Rule

Single electrons with the same spin will occupy each orbital with equal energy before pairing.

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31

Orbital Notation

This image represents?

<p>This image represents?</p>
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32

Electron Configuration

This image represents?

<p>This image represents?</p>
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33

Noble Gas Notation

This image represents?

<p>This image represents?</p>
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34

Dimitri Mendeleev

The first scientist to create a periodic table of the elements in order of increasing atomic mass.

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35

Henry Mosley

Rearranged the periodic table to be in order of increasing atomic number.

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36

Periodic Law

When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.

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37

Nanometer (nm)

10^-9

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38

Megameter (Mm)

10^6

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39

Heisenberg Uncertainty Principle

It is impossible for us to simultaneously know both the exact momentum and location. It can only discuss the probability of an electron's position in space.

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40

Higher energy

Lower energy further from nucleus equals what?

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41

De Brogile

Suggested that electrons could behave as waves with a characteristic wavelength.

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42

Effective nuclear charge

The attractive force felt from the nucleus by electrons It depends on the distance from the nucleus and the number of inner electrons. (Result of Coulomb's law)

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43

Down a group

Increase of shielding occurs ONLY when?

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44

Shielding effect

The decrease in attraction between an electron and the nucleus of an atom.
-Atomic radius increases.
-Ionization Energy decreases.
-Electronegativity decreases.
-Zeff (effective nuclear charge) decreases.
-Harder to remove electrons.

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45

Atomic Radius

Half the distance between the nuclei in covalently bonded atoms.
-Decreases left to right
-Increases downward.

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46

Ionization Energy


The energy required to remove electrons.
-Increases left to right.
-Decreases downward.

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47

Electronegativity

Measure of an atom's bond to another atom.
-Increases left to right.
-Decreases downward.

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48

Electron Affinity

The energy change that results in the addition of electrons
-Increases left to right.
-Decreases downward.

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49

The force of attraction of the nucleus is greater than the ground state.

Why is there a huge jump between n=1 and n=2?

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50

Isoelectronic

When a group of ions have the same number of electrons.

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51

Metallic charcter

An element that has the physical and chemical properties of metals.
-Decreases left to right.
-Increases downward.

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52

The attraction of an extra proton.

The distance from the S to P orbital out weighs what?

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53

Repulsion of paired electrons.

As the e- moves to enter an orbital already having 1e- they
pair with opposite spins so the IE drops.
-Outweighs the increased attraction from the nucleus so less energy to remove the e-.

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54

Electron givers

The most reactive metals are the largest since they are the best what?

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55

Electron takers

The most reactive nonmetals are the smallest ones, the best what?

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56

Paramagnetism

The magnetic state of an atom with one or more unpaired electrons.

<p>The magnetic state of an atom with one or more unpaired electrons.</p>
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57

Diamagnetism

The magnetic state of an atom with paired electrons with no magnetic attraction.

<p>The magnetic state of an atom with paired electrons with no magnetic attraction.</p>
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58

Coulomb’s Law

The strength of the interaction between two electrical charges depends on the magnitudes of the charges and on the distance between them.

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