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Chem U2 Atomic Structure and Properties > Early Models of the Atom (2.1) > Flashcards

Early Models of the Atom (2.1) Flashcards

1
Q

Which law did Antoine Lavoisier introduce?

A

Law of Conservation of Mass

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2
Q

Which law did Joseph Proust introduce?

A

Law of Definite Proportions

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3
Q

Which law did John Dalton introduce?

A

Law of Multiple Proportions

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4
Q

What were the 4 key points of Dalton’s atomic theory?

A
  1. Elements are made of atoms
  2. Atoms of a given element are identical; different elements have different atoms
  3. Compounds form when atoms join together; a given compound always has the same numbers and types of atoms
  4. Reactions involve the rearrangement of atoms (but don’t change the atoms themselves)
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5
Q

Who created the billiard ball model?

A

John Dalton

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6
Q

What is the charge-to-mass ratio of an electron?

A

charge on the electron (e) in coulombs (C) / mass in grams = -1.76 x 10⁸ C/g

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7
Q

Which model did J.J. Thompson create and how does it work?

A

Plum pudding model: entire atom is positively charged with negative electrons embedded in it

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8
Q

What did Ernest Rutherford conclude when he tested Thomson’s model?

A

The center of the atom must be positively charged and dense; used the term “protons” for the positively charged particles; most of the atom is empty space

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9
Q

Which model did Rutherford create and how does it work?

A

Nuclear atom model: an atom with a dense positive nucleus containing protons with electrons moving around the nucleus

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10
Q

Who suggested the existence of neutrons and why?

A

Rutherford; found that protons and electrons could only account for about half the mass of any atom

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11
Q

Who proved the existence of neutrons?

A

James Chadwick

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12
Q

Who first explained the idea of isotopes and what are they?

A

James Chadwick; atoms of the same element with different masses

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13
Q

Describe the nucleus

A

Very dense; contains nearly all of the atom’s mass

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14
Q

In standard atomic notation, what does each letter stand for?

A

A - the number of protons + neutrons (mass number)
Z - the number of protons (atomic number)
X - element symbol

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15
Q

What changes and what stays the same in isotopes?

A

The number of electrons and neutrons CAN change; the number of protons NEVER changes for any element

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16
Q

Describe how protons behave

A

Protons are tightly packed inside the nucleus - these particles repel each other

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17
Q

Describe how neutrons behave

A

Neutrons separate the protons reducing repulsion and stabilizing the nucleus

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18
Q

What force attracts protons and neutrons to each other?

A

Strong nuclear force

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19
Q

What is needed to make an atom stable?

A

The right balance of protons and neutrons

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20
Q

What happens when protons and neutrons are not properly balanced?

A

The nucleus becomes unstable and decays into a more stable nucleus

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21
Q

Define radioisotopes

A

Isotopes with unstable nuclei that decay into different (often more stable) isotopes

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22
Q

Define isotopic abundance

A

The percentage of a given isotope in a sample of an element

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23
Q

What is a mass spectrometer?

A

A measuring instrument used to determine the mass and abundance of isotopes

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24
Q

What do we call the data produced by mass spectrometers?

A

Mass spectrum

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25
Define nuclear radiation
Energy or small particles emitted as an isotope decays
26
What are the 3 common types of nuclear radiation?
1. Alpha particles 2. Beta particles 3. Gamma radiation
27
Define alpha particle
Same as He-4 atom (2 protons and 2 neutrons) with +2 charge; low energy - can be blocked by paper
28
Define beta particle
A negatively charged electron that can pass through paper but not metals
29
Define gamma radiation
High energy electromagnetic radiation; travels at speed of light and can penetrate most things except for lead
30
Define atomic mass unit
The mass of an element shown in the periodic table (atomic mass units, u); one atomic mass unit is equal to one twelfth the mass of a carbon-12 atom
31
How do you calculate average atomic mass?
Sum the known isotopes for an element multiplied by their percentage abundance
32
Define electromagnetic radiation
One of the many ways energy travels; travels at the speed of light through a vacuum
33
Define wavelength
The distance between peaks or troughs in a wave
34
What is frequency?
The number of waves (cycles) per second that pass a given point per second
35
Which wavelength radiation has the highest frequency?
Short-wavelength radiation
36
What's the speed of light formula?
(wavelength in m)(frequency in Hz) = C (speed of light)
37
What did Max Planck discover about matter?
Matter could not absorb or emit any quantity of energy
38
What did Planck mean by "energy is quantized"?
It could only be absorbed or emitted in whole number multiples of hv
39
What do the letters in hv stand for?
``` v = frequency h = Planck's constant 6.626x10^-34 J-s ```
40
What is the formula for the change in energy for a system?
``` Delta E (Change in energy) = nhv (n = an integer, h = Planck's constant, v = frequency) ```
41
Define continuous spectrum
The spectrum observed when light is passed through a prism (a rainbow)
42
Which light contains all wavelenghts?
White light
43
Define line spectrum
The spectrum observed when the emission spectrum of an element is passed through a prism (only a few discrete lines are seen)
44
What is the quantum model and who created it?
Niels Bohr proposed that the electron in hydrogen moves around the nucleus on fixed orbits (shells)
45
What did Bohr conclude about the movement of electrons?
They could only orbit in fixed paths or shells around the nucleus
46
What is Coulomb's Law?
The force between two charged particles is proportional to the magnitude of the charges and inversely proportional to the square of the distance between them
47
What is the Heisenberg Uncertainty Principle?
There is a limit to precisely knowing the position and momentum of a particle
48
What did De Broglie reason about the electron?
The electron, long considered to be a particle, showed wave properties too
49
In Schrodinger's equation, what does the H, w and E stand for?
H: the operator W: wave function E: total energy of atom
50
The more accurately we know a particle's position...
The less accurately we know momentum
51
What are probability distributions?
The probability of finding an electron in a given space
52
Where does one have the greatest probability of finding an electron and why?
Closest to the nucleus, because the further away the electron is, the more space it can occupy
53
What shape are s orbitals? How many electrons can they hold?
Spherical; can hold a maximum of 2 electrons
54
What shape are p orbitals and how many electrons can they hold in total?
Dumbbells/figure 8s; have 3 lobes (x, y, z) each holding 2 electrons for a total of 6 electrons
55
What shape are d orbitals and how many electrons can they hold in total?
Dumbbells/figure 8s; have 5 lobes each holding 2 electrons for a total of 10 electrons
56
What shape are f orbitals and how many electrons can they hold in total?
Dumbbells/figure 8s; have 7 lobes each holding 2 electrons for a total of 14 electrons
57
What is the aufbau principle?
In the ground state of an atom or ion, electrons fill orbitals of the lowest energy levels before occupying higher levels
58
Give the first twelve in the filling order of orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s
59
What is the electron correlation problem?
Since the exact pathways of electrons cannot be calculated, their exact repulsions cannot be calculated either
60
What is the shielding effect?
Inner shell electrons shield outer shell electrons from the pull of the nucleus
61
What is the order of the energies of orbitals (lowest to highest)
s < p < d < f
62
What is the Pauli Exclusion Principle?
Only 2 electrons with opposite spin can occupy an orbital
63
What is Hund's Rule?
Ensure each orbital has a single electron before doubly filling
64
How is our modern periodic table arranged?
According to atomic number
65
What are three characteristics of metals?
1. Conduct heat and electricity 2. Are malleable 3. Are ductile
66
Do metals lose or gain electrons? What ions do they form?
Metals tend to lose electrons to form positive ions (cations)
67
Do nonmetals lose or gain electrons? What ions do they form?
Nonmetals tend to gain electrons forming negative ions (anions)
68
State the periodic law
When the elements are arranged by atomic number, their chemical and physical properties recur periodically
69
What is a period and how does it work?
A row in the periodic table; as you go across a period, electrons are being added to the same shell until it is full at the noble gases
70
What is a group and how does it work?
A column in the periodic table; all elements within the same group have similar properties; electrons in the same group have the same valence electron configuration
71
What are alkali metals?
Soft, shiny metals; good conductors of heat and electricity; malleable and ductile; form 1+ cations when reacting with nonmetals; very reactive
72
What are alkaline earth metals?
Shiny metals; good conductors of heat and electricity; malleable and ductile; form 2+ cations; reactive, but less than alkali
73
What are transition metals?
Groups 3-11; very hard metals with high melting points; can form a variety of compounds with other elements due to their complex electron arrangement; often used as catalysts
74
What are metalloids?
The "staircase elements"; have properties of both metals and nonmetals; are shiny, but poor conductors of electricity
75
What are nonmetals?
All gases that are not noble gases or halogens; not as reactive as halogens, more reactive than noble gases
76
What are halogens?
Very reactive (dangerous) nonmetals; form 1- anions; poisonous; solids, liquids or gases, tend to form salts; not shiny and do not conduct electricity
77
What are noble gases?
Extremely non reactive; have a full outer shell
78
What is condensed electron configuration?
Place the atomic symbol for the last noble gas in square brackets, then add the valence electrons for the element
79
What are isoelectronic elements/ions?
Elements/ions that have the same electron configuration
80
What are inner electrons called?
Core electrons
81
Give an example of how (n+1)s orbitals fill before the nd orbitals
4s orbitals fill before 3d orbitals since 4s electrons can get closer to the nucleus providing more stability to the atom
82
What is true of groups 1-8 of the periodic table, when it comes to electron configurations?
Elements have the same number of valence electrons as their group number; they also have the same valence electron configurations
83
What is the atomic radius and what units is it measured in?
The distance from the centre of an atom to the boundary within which the electrons spend most of their time; measured in picometers
84
What trend is visible with atomic radius?
Atomic radius increases down a group and decreases across a period
85
What is effective nuclear charge? Provide a formula
The nuclear charge experienced by the outermost electrons of an atom, as a result of the shielding by the inner shell electrons; ENC = # protons - # inner electrons
86
What is the ionic radius?
The size of an ion as measured from the center of an ion to its outermost electrons
87
What is ionization energy?
The amount of energy (kJ/mol) required to remove the outermost electron from an atom or ion in the gaseous state
88
What trend is observed with ionization energy?
Ionization energy increases across a period, and decreases down a group
89
What is electron affinity?
The energy change that occurs when an electron is added to a neutral atom
90
What unit is electron affinity measured in?
kJ/mol
91
What trend is found with electron affinity?
Electron affinity increases across a period and decreases down a group
92
What is electronegativity?
The relative ability of an atom to attract shared electrons
93
What trend is found with electronegativity?
Electronegativity increases across a period and decreases down a group

Chem U2 Atomic Structure and Properties (4 decks)

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