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Chemistry 4U > Unt 1 - Atomic Structure > Flashcards

Unt 1 - Atomic Structure Flashcards

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1
Q

What year did Dalton create the Billard ball model?

A

1809

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2
Q

What was the limitation of Dalton’s Billard ball model

A

It could not explain why atoms of elements combine in the way they do

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3
Q

What was Thompson’s model and what year was it created

A

1887,
A chocolate chip cookie
Electrons in a positive sphere

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4
Q

What and who disproved Thompson’?

A

Rutherford with the gold foil experiment

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5
Q

What year was Rutherford’s diagram created and what did it look like?

A

1909
Positive nucleus, electrons floating around

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6
Q

What were the limitations on Rutherford’s theory?

A

-A nucleus of only positive charges would repel
- didn’t account for entire weight of an atom
- Electrons in motion should continuously give off radiation (light) and lose energy, atoms would implode

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7
Q

What did Bohr’s diagram look like and what year was it created?

A

1913,
Orbits of electrons and energy levels

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8
Q

Limitations of Bohr diagram

A

-Could only explain single electron atoms (H, He, Li etc)

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9
Q

Why did Bohr disprove Rutherford?

A

Energy emitted should be continuous but it is not
Within energy level electrons do not emit energy
Electrons change level by absorbing or emitting energy

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10
Q

Who discovered neutrons?

A

James Chadwick, 1932

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11
Q

Who discovered protons?

A

Rutherford, 1914

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12
Q

Who discovered isotopes?

A

Frederick Soddy, 1913

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13
Q

What is a photon?

A

A particle of light
What a quantum of light energy is called

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14
Q

What is a Quantum?

A

A packet of energy
Plank

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15
Q

What is a continuous spectrum?

A

The spectrum that consists of a continuum of wavelengths

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16
Q

What is Emission spectrum?

A

The spectrum of frequencies of electromagnetic radiation emitted due to an electron making a transition from high energy state to a lower energy state

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17
Q

What is an absorption spectrum?

A

The spectrum of frequencies of electromagnetic radiation absorbed due to an electron making a transition from a low energy state to a higher energy state

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18
Q

In an orbit, are electrons a fixed or variable difference from the nucleus?

A

Fixed distance

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19
Q

Are orbitals 2D or 3D?

A

3D

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20
Q

How many electrons are there per orbital?

A

2

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21
Q

How many electrons in an orbit?

A

2, 8 or 18

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22
Q

An orbital with 1 orbital orientation, up to 2 electrons that is in all energy levels is called…

A

S-shape

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23
Q

An orbital with 3 orbital orientations, up to 6 electrons that is in energy level 2 and up is called…

A

P-shape

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24
Q

An orbital with 5 orbital orientations, up to 10 electrons and in energy level 3 and up is called…

A

D-shape

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25
An orbital with 7 orbital orientations, up to 14 electrons and is in energy level 4 and up is called…
F-shape
26
What is the right order when writing an energy level diagram?
1s, 2s, 2p, 3s, 3p, 4s, 4d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 5f, 6d
27
What can the row an element is in on the periodic table tell you about its energy level diagram?
Energy level
28
What form of diagram/ configuration do you draw the direction of spin each electron has?
Energy level diagram
29
What is this an example of? 1s^2, 2s^2, 2p^6, 3s^2, 3p^1
Complete electron configuration
30
What is this an example of? [Ne] 3s^2, 3p^1
Condensed electron configuration
31
What does VSEPR stand for?
Valence Shell Electron Pairs Repel
32
What is VSEPR theory?
Pairs of electrons in valence shell stay as far away as possible
33
What is a Hybrid orbital?
Atomic orbitals overlap to form new orbital containing a pair of electrons of opposite spin
34
Which VSPER shapes are always non-polar?
Linear Linear Trigonal Planar Tetrahedral
35
What VSEPR shapes may be polar?
Bent/Angular Bent/Angular Pyramidal
36
What EN difference does a bond need to be polar?
0.4-1.8
37
EN difference of a covalent bond?
<0.4
38
EN difference of an ionic bond
>1.8
39
EN difference of an ionic bond
>1.8
40
Give 2examples if intramolecular forces
Ionic and Covalent
41
What is an intramolecular force?
Forces between atoms and ions within a molecule
42
Which is stronger, intermolecular forces or intramolecular forces?
Intramolecular
43
What is another name for intermolecular forces?
Van der Waals forces
44
What is an intermolecular force?
Attractive forces between molecules
45
What are weaker, intermolecular forces or intramolecular forces?
Intermolecular forces
46
What are 3 examples of intermolecular forces?
H-bonds, dipole-dipole, London dispersion forces
47
Does it take less energy to melt a substance or break it down into separate elements?
Melt a substance
48
What is a dipole-dipole force
Attraction between oppositely charged ends of polar molecules
49
Which intermolecular force is the weakest?
London dispersion
50
What is a London dispersion force?
Nucleus attracts electrons from neighbouring molecules Present in all molecules, important in non-polar molecules At moderate distance attraction can be stronger than the repulsion of electrons
51
The higher amount of electrons the ________ the London force
Stronger
52
The _________ amount of electrons the weaker the London force
Lower
53
What size of molecules have low boiling points
Small neutral molecules
54
Do big molecules have high or low boiling points?
High
55
Molecules with the approximate same London forces have the same amount of _______
Electrons
56
The more polar a molecule and the stronger the dipole-dipole attraction the ___________
Higher the boiling point
57
What is the strongest intermolecular force?
Hydrogen bonds
58
What is a hydrogen bond?
Strong dipole-dipole force
59
H-bonds are _____%-______% as strong as covalent bonds
5%-10%
60
What elements does hydrogen bond with in H-bonds?
Fluorine Oxygen Nitrogen
61
Rank the intermolecular forces from weakest to strongest
London dispersion Dipole-dipole H-bonds
62
The higher the force the higher the __________+__________
Melting point and boiling point
63
What is an ion dipole?
Ion and polar molecule Ex. NaCl separating into ions when surrounded by water
64
Ion induced dipole, Ex O2 sticks to hemoglobin in bloodstream
Ion close to non polar molecule induced temp dipole
65
Dipole induced dipole
charge on polar molecule induces dipole on non-polar molecule
66
which intermolecular force has the highest vapour pressure and Wich has the lowest
HIghest: london Lowest: H-bonds
67
What is an ionic crystal?
A non-metal + metal with very strong ionic bonds in all three planes
68
What are the characteristics of ionic crystals?
* High melting point and boiling point * An electrolyte in solution * High solubility * Hard and brittle
69
How do metallic crystals differ from ionic crystals?
They have weaker electrostatic attraction and positive core with mobile electrons
70
List the properties of metallic crystals.
* Variable melting and boiling points * Electrolyte * Not soluble in water * Lustrous * Malleable * Ductile * Hard * Dense * Conducts heat
71
What defines a molecular crystal?
Non-metals with weak intermolecular attraction
72
What are the properties of molecular crystals?
* Low melting and boiling points * Not an electrolyte * Soft
73
What characterizes a covalent network crystal?
Covalent bonds between molecules
74
List the properties of covalent network crystals.
* Very high melting and boiling points * Not an electrolyte (except graphite) * Not soluble in water * Hard * Brittle
75
What is an amorphous solid?
A solid with various bonding that melts over a range of temperatures
76
What are the characteristics of amorphous solids?
* An electrolyte for metallic variants * Very soluble in water * Composed of molecules not arranged in an orderly crystalline structure
77
What is notable about graphite's bonds compared to diamonds?
The bonds in graphite are stronger than those in diamonds
78
What are the properties of graphite?
* Conductivity * Resistant to chemical reactions
79
How much stronger is graphene compared to diamond?
40x stronger than diamond
80
What makes graphene superior to graphite?
Significantly superior properties
81
What experiment disproved Thompson's model of the atom?
Gold foil experiment ## Footnote Conducted by Rutherford in 1909.
82
Who conducted the gold foil experiment?
Rutherford ## Footnote The experiment took place in 1909.
83
What particles were used in the gold foil experiment?
Alpha particles ## Footnote These particles were shot at thin gold foil.
84
What unexpected result occurred during the gold foil experiment?
1 in 1800 alpha particles deflected significantly ## Footnote This was not expected based on Thompson's model.
85
What conclusion was drawn from the gold foil experiment?
Atom is mostly empty space with a concentrated charge at the center, the nucleus ## Footnote This led to a new understanding of atomic structure.
86
What limitation of the gold foil experiment is mentioned?
Only positive nucleus would repel ## Footnote It did not account for the entire weight of an atom.
87
What is another limitation regarding electrons mentioned in the context of the gold foil experiment?
Electrons in motion should continuously give off radiation (light) and lose energy so atoms would implode ## Footnote This was not addressed in the experiment.
88
What concept did de Broglie introduce regarding electrons?
Electrons as waves ## Footnote De Broglie's theory states that electrons do not follow specific orbits but exhibit wave-like behavior.
89
What condition must the orbital circumference meet according to de Broglie's theory?
It needs to fit whole waves, not half ones ## Footnote This means that only orbits of certain sizes are allowed.
90
What does Heisenberg's uncertainty principle state?
It is impossible to predict both the position and momentum of an electron with certainty ## Footnote This principle arises due to the wave nature of matter.
91
What significant equation did Schrödinger derive?
An equation that tells the probability that an electron is at a particular point ## Footnote This equation is fundamental in quantum mechanics.
92
What replaced Bohr's orbits in Schrödinger's model?
Orbitals ## Footnote Orbitals represent regions of probable location for electrons rather than fixed paths.
93
How are electron clouds represented in quantum mechanics?
They can be drawn to show regions where electrons are likely to be found ## Footnote This visualization helps in understanding the behavior of electrons.
94
What is a characteristic of Bohr's orbits?
They are 2D ## Footnote In Bohr's model, electrons are fixed at a distance from the nucleus.
95
What is the maximum number of electrons that can occupy a Bohr orbit?
2, 8 or 18 electrons per orbit ## Footnote These numbers represent the capacity of different orbits in Bohr's model.
96
What is a characteristic of orbitals in quantum mechanics?
They are 3D ## Footnote Orbitals allow for variable distances of electrons from the nucleus.
97
What is the maximum number of electrons that can occupy a single orbital?
2 electrons per orbital ## Footnote This is based on the Pauli exclusion principle.

Chemistry 4U (4 decks)

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