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Chem 150 > Exam 1 > Flashcards

Exam 1 Flashcards

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1
Q

Where is mass number located in standard form of chemical symbol?

A

on the top

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2
Q

Mass number

A

gives number of protons and neutrons in atom

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3
Q

Where is the atomic number located in standard form of chemical symbol?

A

on the bottom

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4
Q

What is the atomic number?

A

Number of protons

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5
Q

What is atomic mass?

A

Weighted average of all the naturally occurring isotopes of that element

Units: amu,

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6
Q

Calculating atomic mass

A

Sum of (mass #)(% abundance)

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7
Q

Molar Mass

A

same as the atomic mass but in g/mol

1:1 ratio with atomic mass

Units: g/mol

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8
Q

5 stages of mass spectrometry

A

1) Vaporization
2) Ionization (knock off electrons to make a charged particle)
3) Acceleration
4) Deflection (uses a magnet)
5) Detection (mass/charge gets detected)

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9
Q

Evidence we say when using mass spectrum

A

based on peak heights in the mass spectrum, the relative intensity …

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10
Q

Reasoning we say when using mass spectrum

A

write out possible combinations in chart

multiple/determine probability of that combo

order DOES matter!!

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11
Q

Difference between percent abundance and “intensity %” or “relative intensity”

A

find out!!

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12
Q

Which properties of a wave determine the wave’s energy?

A

Intensity (amplitude) AND wavelength (frequency)

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13
Q

How does the photoelectric effect support the particle-like nature of light? with reasoning

A

Only photons at or above the threshold frequency will eject electrons

Reasoning: Planck states that the energy of the photon is reliant on frequency which matches up that only frequency affects it (not intensity)

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14
Q

How does the photoelectric effect not support the wave-like nature of light? with reasoning

A

Increasing the light intensity does not eject electrons if below threshold frequency

Reasoning: a wave with higher amplitude should eject more electrons since it has more energy, but it does not

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15
Q

If light is a particle how can you increase the intensity?

A

Increase the number of photons

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16
Q

Electromagnetic spectrum

A

continuum of radiation

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17
Q

Wavelength

A

distance between identical points on successive waves

Units: m, nm, Angstrom

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18
Q

Frequency

A

number of waves that pass through a particular point in a second

Units: Hz, 1/s

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19
Q

Amplitude (intensity)

A

distance from the midline of a wave to the top of the peak or bottom of the trough

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20
Q

Double slit experiment

A

Since an interference pattern occurred and light did not just travel as two lines through slits, it shows that light is a wave

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21
Q

List the types of waves from shortest wavelength to longest wavelength

A 
gamma
x-ray
ultraviolet
visible (purple-red)
infrared
microwave
radiowaves
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22
Q

What wavelengths are the visible spectrum?

A

290 nm - 800 nm

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23
Q

Blackbody radiation

A

when a solid is heated, it emits electromagnetic radiation in a wide range of wavelengths

radiates a continuous spectrum of all wavelengths

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24
Q

What does Planck propose?

A

Radiant energy can be emitted or absorbed only in discrete quantites

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25
Quantum
smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation
26
What is energy always emitted in?
Whole number multiples of h*v
27
Threshold frequency
electrons are held in a metal by attractive forces and to overcome this force and break them a higher frequency is needed frequency is proportional to energy in a particle (h*v), so this makes sense
28
Wave
a disturbance that travels through a medium, transporting energy from one location to another location without transporting matter
29
What happens when electrons move up in energy level?
Energy is absorbed
30
What happens when electrons move down in energy level?
Energy (photon) is emitted
31
Relationship between energy of an electron and its distance away from the nucleus
Energy is higher further away because it has to overcome electromagnetic attraction to nucleus also potential energy decreases, so kinetic energy increases
32
Atomic line spectra
the emission spectra of atoms in the gas phase do not show a continuous spread of wavelengths instead, they should bright lines in distinct parts of the visible spectrum
33
line spectrum
emission of light only at specific wavelengths
34
What does Bohr postulate?
Hydrogen's electron is only allowed to occupy certain orbits of specific energies Energies of electrons are quantized
35
Ground state
lowest energy state on an electron n=1
36
Excited state
As n increases, the stability of the electron decreases
37
What does the brightness of a spectral line depend on?
how many photons of the same wavelength are emitted
38
Emission spectra
the spectrum of frequencies emitted due to an atom making the transition from a higher energy state to a lower energy state
39
Radial probability distribution
tells you the probability that an electron in that orbital will be at a given distance from the nucleus
40
Penetration power
the ability of an orbital to attract an electron penetration power is highest closest to the nucleus because it is closer to positive charge since penetration power is high, shielding power is low, but its ability to shield others is high
41
Shielding power
refers to the core electrons repelling the outer rings nucleus has "less grip" on outer electrons and are shielded from them
42
Order of electron penetration from least to greatest
s, p, d, f Note: this is also the amount of shielding done
43
Periodic trends due to penetration and shielding
1) atomic radius decreases from left to right and top to bottom 2) ionization energy increases from left to right and decreases top to bottom 3) electronegativity increases left to right ad decreases top to bottom
44
What does the de Broglie equation tell us?
Electrons have wave + particle properties wavelength = h/mv more that electrons act as standing waves
45
Why do we not care about the waves of macroscopic items?
There mass is so big and de Broglie tells us that mass is inversley proportional to wavelength, so the wavelength would be really small
46
Difference between an orbit and an orbital
an orbit is a defined path around a center (larger) object that occurs due to gravitational or electromagnetic attraction an orbital is a 3D space where an electron is most likely to be due to electron density and probability
47
What does the Heisenberg Uncertainty Principle tell us?
it states that since the electron acts as a wave, and a wave is always moving, it is impossible for us to know the velocity and position of an electron at the same time also tells us the more certain we know position, then the less certain we know velocity and vice versa
48
Standing wave
A wave is fixed at two places this tells us that only certain wavelengths are allowed for a standing wave and energies are quantized standing waves also produce nodes
49
Types of nodes
Radial (goes in a circle) Planar (cuts through like a plane)
50
How do we determine types of nodes?
n-1 = how many total nodes n- l - 1 = how many radial nodes subtract radial from total to find amount of planar nodes
51
Principle Quantum Number (n)
tells us the main energy level an electron is in as n increases, distance from the nucleus increases, and therefore energy increases
52
What is another name for the principle quantum number?
Schell
53
Angular momentum number (l)
tells us the SHAPE of the orbital l is dependent on n. Can have integers from 0 up to n-1
54
What is l-value of a s orbital?
0
55
What is l-value of a p orbital?
1
56
What is l-value of a d orbital?
2
57
What is l-value of a f orbital?
3
58
What is another name for the angular momentum number?
Sometimes referred to as a sub-shell
59
Magnetic quantum number (ml)
tells us ORIENTATION is related to l ml = -l to +1 the number of possible mls, tells us the number of possible orientations for the orbital
60
Spin quantum number
can either be +1/2 or -1/2 if you have 2 electrons in one orbital, one has to spin + and one has to spin -
61
What does Ψ mean?
It gives us a mathematical shape of a wave It is a function of x. We could plot it
62
What does Ψ^2 mean?
It gives us the probability of finding an electron at any given point
63
Pauli Exclusion Principle
no 2 electrons in an atom can have the same 4 quantum numbers this means if 2 electrons are occupying the same orbital, they must be spinning in opposite directions
64
Aufbau Principle
process by which the periodic table can be built up by successively adding one proton to the nucleus and one electron to the most appropriate atomic orbit
65
Hund's Rule
most stable arrangement of electrons in orbitals of equal energy is the one which the number of electrons with the SAME spin is maximized
66
Diamagnetic
when all electrons are paired (2 in each orbital) repelled by a magnetic field
67
Paramagnetic
when 1 or more electrons are unpaired attracted by a magnetic field
68
3 Rules for writing electron configuration
1. Electrons will reside in orbitals of the lowest possible energy (most stable) 2. Each orbital can accommodate a max of 2 electrons 3. Electrons will not pair in equal energy orbitals if an empty orbital is available
69
Mole
SI unit for the amount of a substance
70
Mixture
combination of 2 or more substances in which each substance retains its distinct identity can be separated without changing the identity of the substances
71
Homogeneous mixture (and example)
the composition of the mixture is uniform throughout Ex: salt water
72
Heterogeneous mixture
the composition is not uniform
73
Difference between a mixture and a compound
compounds are chemically combined compounds cannot be taken apart without changing the identity of the substance
74
Avogadro's number
6.022 x 10^23
75
Difference between atomic mass and mass number
mass number gives the protons and neutrons in an atom atomic mass is a weighted average of all the naturally occurring isotopes of an element
76
Colors from high energy to low energy in the visible spectrum
violet, blue, green, yellow, orange, red
77
How does intensity affect the electrons that are ejected from a metal surface?
Intensity does not matter until the threshold frequency is reached
78
How does thinking of light as a particle connect to the Bohr model of the atom?
Bohr treats the electron as a particle that orbits in defined paths around a larger central mass this is a very classical mechanics way of thinking about the electron
79
What does it mean for an orbit to be quantized?
The electron can only be in discrete energy levels. It is these discrete orbits. It cannot be inbetween orbits.
80
Why does hydrogen have a line spectrum and not a continuous spectrum?
The energy that hydrogen can emit is quantized since the orbits are quantized
81
Exceptions to Hund's rule
1. Chromium [Ar] 4s1 3d5 2. Copper [Ar] 4s1 3d10 Lathanum and actinide series at bottom sometimes go to 5d orbital rather than 4f
82
Isoelectronic series
a group of atoms/ions that have the same number of electrons
83
How to determine which member of an isoelectronic series is larger?
Size is determined by number of protons More protons = smaller since the protons positive charge pulls the electrons closer
84
How is it possible that the hydrogen atom (which has 1 electron) has 4 lines in the visible spectrum?
never looking at just one atom of hydrogen looking at many hydrogen atoms that each have electrons bouncing from energy level to energy level

Chem 150 (10 decks)

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