exam 2 Flashcards
Describe the relationship between the frequency, wavelength, and velocity (speed) of a wave.
v= f x h
v = velocity, f = frequency and h = wavelength
if velocity is constant, then wavelength and frequency have an inverse relationship (if one increases, the other decreases)
Draw and compare two waves of different frequency
TO DRAW: For different frequencies, draw one wave with tightly packed peaks and troughs (high frequency) and another with widely spaced peaks and troughs (low frequency).
A wave with a higher frequency will have more cycles (oscillations) in a given period of time. The wavelength will be shorter, and the wave will appear more “compressed” with more peaks and troughs within the same space.
A wave with a lower frequency will have fewer cycles in the same time period, so the wavelength will be longer, and the wave will appear more “stretched out” with fewer peaks and troughs.
Draw and compare two waves of different wavelength
To Draw: you could draw one wave with long stretches between peaks (long wavelength) and another with short stretches (short wavelength).
The wave with the longer wavelength will have fewer cycles in a given distance, and the wave with the shorter wavelength will have more cycles in the same distance.
smaller = more compressed - taller
longer = shorter
Draw and compare two waves of different amplitude.
TO DRAW: you could draw one wave with tall peaks and deep troughs (high amplitude) and another with smaller peaks and shallower troughs (low amplitude).
Higher Amplitude: A wave with a higher amplitude will have higher peaks and lower troughs, meaning it carries more energy and has greater intensity.
Lower Amplitude: A wave with a lower amplitude will have smaller peaks and troughs, meaning it carries less energy and has lower intensity.
The wave with the higher amplitude will appear taller, with more extreme positive and negative values, while the wave with the lower amplitude will appear smaller and less intense.
what rank is the electromagnetic radiations: gamma ray, X-ray, UV, visible, IR, microwaves, radiowaves: following in terms of energy, wavelength, or frequency.
energy (highest to lowest): gamma rays, xrays, uv, visable, IR, microwaves, radiowaves
frequency(highest to lowest): gamma rays, xrays, UV, visable, IR, microwaves, radiowaves
wavelength (shortest to longest):gamma, xrays, uv, visable, IR, microwaves, radiowaves
identify experimental evidence for why electromagnetic radiation is a wave.
-interference patterns found during the double split experiment
-diffraction shown in double split experiment
-polarization
-photoelectric effect discovered by Albert Einstein in 1905
-refraction
define electromagnetic radiation
a form of energy that propagates through space in the form of a wave - has both electric and magnetic field components.
explain Thomas Young’s double slit experiment (1801)
When light (or other electromagnetic waves) passes through two slits, it produces an interference pattern on a screen placed behind the slits.
define interference pattern and the difference between constructive and destructive interference.
alternating bright and dark bands
constructive - bright bands
destructive - dark bands
define diffraction
when waves encounter a slit or “obstacle” the wave bends around the edges of the obstacle (or through the slit) creating patterns that are characteristics of wave behavior.
define polarization , how are waves polarized?
the orientation of the oscillations (movement back and forth) of the electric field in an electromagnetic wave
passing through certain materials which blocks orientations of the waves electric field
explains how the wave moves (its vibrations) once it is polarized it only moves in one direction.
what is the photoelectric effect?
discovered by Albert Einstein in 1905 which supports the idea of light having particle like properties. According to the property of a wave, the intensity (or amplitude) of light would affect the emission of electrons on a metal surface but the experiment showed only light of a certain FREQUENCY could eject electrons, regardless the intensity.
define “refraction”
light bends as it passes through different media (air to water) which happens due the change in speed of light as it passes from one medium to another (can be applied to all waves using Snell’s Law).
What is the balanced equation for the reaction of butane (C4H10) with oxygen (O2) to produce carbon dioxide (CO2) and water (H2O)? (This is called a combustion reaction.)
2C4h10+1302 –> 8CO2 +10H20
HOW TO FIND MOLAR MASS FROM MOLES ( suppose you have g of a substance, moles of substance)
MASS/MOLES = MOLES
10g/.5 = 20 g
To find the molar mass from moles, you simply divide the mass (in grams) by the number of moles. This will give you the molar mass in g/mol, which is the mass of one mole of the substance.
determine how much carbon dioxide (in grams) would be produced if 10 g of butane (c4h10) were burned in excess oxygen (02) .
ANSWER: ~ 30.3g
STEPS:
1. find balanced equation
2. find moles of butane (10g butane /molar mass of c4h10 = 58.12 g/mol ~0.172 butane mol)
3. convert butane to co2 –> use ratio ( 4 mol co2 : 1 mol butane) and multiply it by the moles of butane 4 mol co2/1 mol butane x .172 mol butane = 0.688 mol co2
4. Convert moles of co2 to grams (0.688 mol co2 x 44.01 g/mol co2 (molar mass of co2) ~ 30.3g)
calculate percent yield!
actual yield/theoretical yield x 100
ex:/ If produced 5.4 g carbon dioxide, what would be the percent yield?
5.4g/30.3g x100 ~ 17.8% (using the butane example).
f you had 10 g of butane and 10 g of oxygen, what is the maximum amount of carbon dioxide (in grams) that could be produced? First describe the method you would use to determine this, and then do the calculation.
*find limiting reactant
10g butane ~ 30 g CO2
10g O2 ~ 4.22 g CO2 where oxygen is the limiting reactant because it determines the amount of CO2 that can be produced because it has the least amount.
Explain what the mole is, and why we need to use it to convert from the atomic molecular level to the macroscopic (real world) scale. Explain how you would use it to do a stoichiometry calculation. Be sure to include the rationale for each step in the calculation.
mole is a very large number that is used to scale small things (atoms/molecules) up to a quantity that we can see and measure. A mole is a counting unit, like a dozen or a pair. A dozen means 12 things while a mole means 6.022 x 1023 things.
A mole cennects atoms (molecular level) to grams (macroscopic level).
5.0 moles of sulfur atoms and 10 moles of oxygen molecules are combined to form the maximum amount of sulfur trioxide. How many moles of which atoms/molecules remain in the reaction flask after the reaction has produced the maximum amount of products?
- balance eqaution
- identify limiting reactant using ratio of balanced eqaution (3 mol o2: 2 mol s)
5 mol S x 3/2 = 7.5 mol Oz
10 m O2 x2/3 = 6.6 mol S so sulfur is the limiting reactant. - use ratio of sulfur sulfur oxide (SO3) (2 mol S:2 mol SO3) 2/2 x 5 mol S = 5 mol SO3
- find remainder - aka 2.5 mol O2 10-7.5 = 2.5 moles O2
7a. What is the balanced equation for the reaction of acetylene (C2H2) with oxygen (O2) to produce carbon dioxide (CO2) and water (H2O)? (This is called a combustion reaction.)
2 , 5 –> 4 , 2
7b. How many moles of oxygen are required to burn 1.0 mole of acetylene?
2.5 mol o2
7c. How many grams of carbon dioxide are produced by the combustion of 1.3 grams of acetylene? Assume there is excess oxygen.
What abt with 3.5 g oxygen?
4.40 g co2
3.85 g co2 (3.5 g o only produces approximately .05 mol of o2, where 1.3 g of acetylene requires .125 mol o2).
what is evidence for light behaving as a particle?
photo electric effect - light shined on a metal surface can eject electrons at the right frequency.
planks hypothesis - energy can only be emitted or absorbed at fixed amounts
how to calculate the energy of photons of a given frequency or wavelength
- when given frequency use:
e = f x h
where e is the energy of the photon (J)
h is planks constant (6.626 x 10^-34 Js)
and f is frequency (Hz) - when given wavelength use:
c = (wavelength (m) x (frequency)
where c is speed of light (3.00 x 10^8 m/s)
what is planck’s constant?
6.626 x 10^-34 Js
Explain why the existence of photons (quantized light energy) explains the photoelectric effect.
Planck’s Law shows that there are packets of energy (e) that have their own unique frequency. Therefore, where the photon strikes an electron with sufficient energy, it can eject an electron that is directly related to the frequency emitted.
what is the speed of light?
3.00 x 10^8 m/s
how does doubling the frequency affect wavelength?
Doubling the frequency will cause the wavelength to shorten by half.
What is the wavelength of light with a frequency of 7.26 × 10^14 Hz?
Answer: 4.13 x 10^-7 m
c/f = wavelength where c is 3.00 x 10^8 m/s / 7.26 x 10^14 (1/s) = 4.13 x 10^-7 m which is 413 nm
how many nm are in a meter?
1 m = 10^9 nm
What is the frequency of radiation with a wavelength of 442 nm?
442 nm x 1 m / 10^9 nm –> 3.00 x 10^8 m/s / 442 x 10^-9 m
= 6.79 x 10^14 Hz (1/s)
What is the energy of a photon with a frequency of 7.26 × 10^14 Hz.
multiply by planks constant:
(6.626 x 10^-34 Js)x(7.26 x 10^14 Hz (1/s))
= 4.81 x 10^-19 J
What is the wavelength of a photon of energy 2.4 × 10^–16 J?
- find frequency using e = f x h
- find wavelength
answer: 8.29 x 10^-10 m or 0.829 nm
The energy required to break 1 mol of C–C bonds (that is 6.022 × 10^23 C–C bonds) is 348 kJ/mol. What would be the minimum frequency of a single photon that would break a single C–C bond?
- find energy required to break a single cc bond:
348 x 10^3 J (348 kj)/6.022 x10^23 = 5.78 x10^-19 J - Use planks eqaution to find frequency!
5.78x10^-19 J / 6.626 x 10^-34 Js = 8.72 x 10^14 Hz(1/s)
how many kj are in a J?
1 kj = 1 x 10^3 J
Make an argument that light (electromagnetic radiation) is a wave.
claim: light is a wave
evidence: When light shines through two slits a diffraction pattern is visible.
reasoning: This is because the waves that are propagating from each slit meet “in phase” and “out of phase” causing an interference pattern.
what does it mean for a wave to be “in phase” or “out of phase”?
When the waves meet “in phase”, constructive interference occurs causing the wave intensity to increase. This is seen by the bright places in the pattern.
When the waves meet “out of phase” (throughs facing opposite direction) , destructive interference occurs. The waves cancel each other producing the dark places in the pattern. This is how diffraction patterns indicate that light is a wave.
explain what is happening when an element produces an absorption spectrum
An electron in an n=1 state absorbs a photon of light, causing the electron to transition to a higher energy level. The photon must have the exact energy as the difference between levels n=1 and n=2.
explain what is happening when an element produces an emission spectrum
An electron in the n=2 state releases a photon of light as the electron transitions to a lower energy level. The photon released has the exact energy as the difference between levels n=2 and n=1.
what is the order from highest to lowest frequency of the colors?
highest - violet (shortest wavelength)
indigo
blue
green
yellow
orange
red - lowest (longest wavelength)
compare the possible electron transitions that would produce the red emission line and the yellow line in the spectrum
There is a bigger difference in energy between levels n=3 and n=1 so a higher energy photon is emitted during that transition. A photon of yellow light is higher in energy than a photon of red light.
why are all emission/absorption spectra’s different?
The spacing between energy levels is different for different elements.
When an electron transitions to a lower level, it emits a photon with an energy that is equal to the energy difference between the two energy levels. This energy difference is not the same for different elements, so the energy, frequency, and wavelength of the photons emitted is not the same.
Explain why we use energy diagrams to illustrate electron transitions, rather than using a Bohr model of the atom.
The Bohr model only works for hydrogen and implies knowledge of both the energy and position of electrons. Since we cannot know both, we use energy diagrams.
These diagrams show that an electron can absorb a photon of light and move to a higher energy level or release (emit) a photon light to move to a lower energy level. These energy levels are not circular orbits and are not a specified distance from the nucleus as Bohr proposed in his model of the atom.
what is an energy diagram?
These diagrams show that an electron can absorb a photon of light and move to a higher energy level or release (emit) a photon light to move to a lower energy level. These energy levels are not circular orbits and are not a specified distance from the nucleus as Bohr proposed in his model of the atom.
Explain how (and why) different atoms emit different wavelengths of light.
Atoms emit different wavelengths of light due to the unique structure and energy levels of the electrons within each atom
Compare and contrast atomic emission and absorption spectra and how they are generated.
Atomic Emission Spectra and Atomic Absorption Spectra both arise from the interaction of atoms with light, but they are generated in different ways and provide different information about the atoms involved.
what is emission spectra?
what is absorption spectra
occur when atoms absorb specific wavelengths of light, which causes electrons to be excited. These spectra show dark lines where light is absorbed (mostly color, with black lines). They are often used for measuring the concentration of elements in a sample.
what is emission spectra?
result from electrons falling from high energy levels to lower energy levels, emitting light in the process. These spectra show bright lines (mostly black with highlights of color) that can be used to identify elements and study their energy levels.
Use spectra to identify the presence of elements (by comparison, not calculation).
Hydrogen: It has well-known spectral lines, such as the red line at 656 nm in the Balmer series.
Sodium: It shows a characteristic yellow doublet at 589.0 nm and 589.6 nm.
Copper: It has distinct lines such as 510.6 nm and 578.2 nm in its emission spectrum.
Describe experimental evidence for the wave nature of electrons.
Electron Diffraction (Davisson and Germer Experiment, 1927) - In this experiment, electrons were fired at a crystal of nickel, and the scattered electrons were detected.
de Broglie’s Hypothesis
Electron Interference (Young’s Double-Slit Experiment with Electrons) - . In this experiment, a beam of electrons was directed through two slits, and the pattern of electrons on a screen behind the slits was recorded.
Wave-Particle Duality (de Broglie and Schrödinger)
Make an argument for why spectra are direct evidence for the existence of quantized energy levels in an atom.
Claim: Spectra are direct evidence for the existence of quantized energy levels in an atom because the discrete lines in atomic spectra correspond to specific energy transitions between quantized electron orbits within the atom.
Reasoning:
Discrete Lines in Spectra, Energy Level Transitions, Bohr’s Model of an Atom, The Rydberg Formula
When atoms absorb or emit light, they do so at specific, distinct wavelengths. In an atomic emission spectrum, we see a series of bright lines, and in an absorption spectrum, we see dark lines. These lines are not continuous but occur at specific wavelengths. The fact that the spectrum consists of only discrete lines and not a smooth continuum of colors is a key piece of evidence for quantized energy levels.
These spectral lines arise when an electron in an atom moves between different energy levels. For example, when an electron absorbs energy, it moves from a lower to a higher energy level. When it falls back down, it emits energy in the form of light. The energy of the emitted or absorbed light corresponds to the difference in energy between two of these discrete energy levels.
If the energy levels were not quantized, meaning they could vary continuously, the emitted or absorbed light would also be spread out across a continuous spectrum, without the sharp, distinct lines we observe. The fact that only specific wavelengths of light are emitted or absorbed supports the idea that only certain energy transitions are allowed, and that these transitions correspond to specific, quantized differences between energy levels.
Niels Bohr’s model of the hydrogen atom explains this phenomenon. According to Bohr, electrons can only occupy certain quantized orbits around the nucleus, and they can only absorb or emit energy in discrete amounts when they jump from one orbit to another. This directly leads to the observation of discrete spectral lines. Bohr’s model was later refined, but the basic idea that energy levels are quantized remains central to our understanding of atomic structure.
The Rydberg Formula:
The wavelengths of the lines in the hydrogen spectrum can be predicted using the Rydberg formula, which relates the wavelengths of the emitted light to the integer values of the energy levels involved in the transition. The fact that these wavelengths can be predicted with a mathematical formula based on discrete energy levels is further evidence that atoms have quantized energy states.
Conclusion:
Spectra provide direct evidence for the existence of quantized energy levels because the discrete lines in an atomic spectrum correspond to specific energy transitions between fixed energy levels in the atom. If energy levels were not quantized, we would not observe these sharp spectral lines, but rather a continuous range of wavelengths. Therefore, the structure of atomic spectra supports the idea that electrons in atoms occupy quantized energy levels.
de Broglie’s Hypothesis
In 1924, Louis de Broglie proposed that all matter (not just light) could have both particle and wave characteristics. He suggested that an electron could be treated as a wave with a wavelength (λ) given by the equation:
Schrödinger’s Equation
The wave-like nature of electrons was formalized by Erwin Schrödinger in 1926, who developed a mathematical model (the Schrödinger equation) to describe the behavior of quantum particles. The equation treats electrons as wavefunctions, which describe the probability distribution of where an electron might be found.
explain how to calculate the wavelength of particles (given mass and velocity).
use de Broglie’s equation, which relates the wavelength (𝜆) of a particle to its momentum. in the equation : wavelength = h/p where h is Planck’s constant and p is momentum of particle.
p = m x v so 𝜆 = h/mv
Describe how the wave properties of the electron are taken into account in the current model of the atom.
The current model of the atom, known as the quantum mechanical model (or the Schrödinger model) incorporates the wave properties of electrons to explain atomic structure and behavior.
This model is based on the principles of quantum mechanics, and it accounts for the wave-particle duality of electrons, which was established by de Broglie and confirmed through experiments like electron diffraction.
Describe an atomic orbital and what it represents.
An atomic orbital is a mathematical function that describes the wave-like behavior of an electron in an atom. It represents a region of space where the probability of finding an electron is high. Essentially, it shows where an electron is most likely to be found around the nucleus of an atom.
The atomic orbital is derived from the wave function (ψ) of an electron, which is a solution to Schrödinger’s equation.
describe an “s” orbital
spherical, no angular nodes (only one s orbital)
describe a “p” orbital
dumbell shaped, figure 8 shaped
consists of px, py, pz
Describe how the model of the atom changed from Dalton through Thompson, Rutherford, and Bohr to Schrodinger. Explain why each model changed and point out the problems with the previous model.
Dalton’s model portrayed atoms as solid, indivisible spheres, but was replaced by Thomson’s “plum pudding” model, which introduced the idea of electrons embedded in a positive charge, after discovering the electron. Rutherford’s gold foil experiment revealed that atoms have a dense nucleus, prompting Bohr to propose quantized orbits for electrons, but Schrodinger’s wave model eventually replaced Bohr’s fixed orbits with probabilistic electron clouds, solving problems of electron stability and orbit quantization.
how does ionization energy and atomic radii change throughout the periodic table?
they have an inverse relationship!
As you move down a column, ionization energy decreases and atomic radii increases due to more electron shells and more electron-electron repulsion.
If you move across the rows, ionization energy increases and atomic radii decreases to a strong effective nuclear force.
Explain the concept of effective nuclear charge
Effective nuclear charge (𝑍eff) is the net positive charge an electron feels from the nucleus, accounting for both the actual nuclear charge and the shielding from inner electrons. It increases across a period because additional protons are added, while shielding remains relatively constant, making electrons more tightly bound.
to find subtract (Pro+)-(core e-) = Zeff
how does effective nuclear charge affect atomic radii?
With a larger charge, there is a stronger electrostatic attraction force between the nucleus and the electrons, causing them to be pulled tightly in.
how does effective nuclear charge affect ionization energies?
With a stronger effective nuclear charge, there will need to be a stronger ionization energy to break the force.
Apply Coulomb’s law to explain periodic trends in atomic radii and ionization energies.
Coulomb’s law explains that as the effective nuclear charge increases across a period, the attraction between the nucleus and electrons strengthens, causing atomic radii to decrease.
describe how ionization energies support the idea of quantized energy levels in atoms.
Ionization energies show that electrons are removed in specific steps, with large jumps in energy required to remove electrons from different energy levels, supporting the idea of quantized energy levels. These energy jumps reflect the distinct energy levels within the atom, where each level has a specific, measurable amount of energy associated with it.
explain how to apply Coulomb’s law to explain the relative sizes of neutral atoms and ions.
Coulomb’s law explains that the greater the effective nuclear charge, the stronger the attraction between the nucleus and electrons, making the atom smaller. In ions, when electrons are gained or lost, the balance between nuclear charge and electron repulsion changes, causing the ion to be either larger (when gaining electrons) or smaller (when losing electrons).
how to Predict the relative sizes of isoelectronic atoms and ions.
what are the three quantum numbers that classify orbitals?
principal quantum number (n)
angular momentum quantum number (l)
magnetic quantum number (m).
what are the four most common orbitals?
s
p
d
f
how many mol of H are in CH2CL2
2x 6.022 x10^23 = 1.204 x 10^24
You start a reaction with 53 grams of H2 and excess O2 and you end up producing 400 grams of H2O. What is the percent yield for this reaction?
YOU MUST CALCULATE THEORETICAL YEILD FIRST !!!!!
(53G/2.018 = 26.3 mol H2 x 1mol H20 x 18.016 g/mol H2O = 473.8 g –> (400g/473.8 g x 100 = 84%.
find wavelength given energy of specific proton.
wv = (hc)/E
where h = 6.626 x 10^-34
c = 3 x10^8
- Hydrogen produces emission lines (colored lines on the spectrum) with the following wavelengths: 410 nm, 434 nm, 486 nm and 656 nm. Where would you expect to see the absorption lines (dark lines) for hydrogen on its absorption spectrum and why?
a. The absorption lines for hydrogen would appear at exactly the same wavelengths, because the energies of the photons emitted and absorbed by hydrogen are the same.
- What is the evidence that supports the claim that light is a particle?
c. When light shines on metal there is a threshold frequency, below which no electrons are ejected from a metal.
An MRI (magnetic resonance imaging) machine operates at a frequency of 30 kilohertz (kHz), what wavelength radiation does this correspond to?
convert 30 kHz to Hz!
30 x 10^3 = 30,000 Hz
- Why must we consider the wave properties of an electron, but not the wave properties of macroscopic objects (such as humans)?
The wavelength of the electron is similar in size to the atom and affects its properties, whereas the wavelength of the macroscopic object is much smaller than the object and does not affect its properties.
Write the electron configuration for fluorine (F) and the electron configuration for chlorine (Cl). What is different about the valence electrons in these two atoms?
II and IV, valence electrons have different energies and different size orbitals
Explain what happens at the atomic level to produce a line on an emission spectrum.
When an electron transitions from a higher energy level to a lower energy level it emits a photon.
Explain why the lines are different for each element.
Each element has different allowed energies for its electrons (ie different energy levels). Since the differences in energy between those levels are different, the energies of the photons released will also be different. Different energy photons have different colors.
If a spectrum was taken on the emissions from a star 300 light years away that contained Hg, would it look the same or different? Explain.
The lines would have the same pattern (same fingerprint) but would be slightly red-shifted because of the Doppler effect (the waves of light are stretched out as that star and Earth move away from each other).
d. Below is an example of an absorption spectrum. Why does it look different than the emission spectra? What happens at the atomic level to produce a line?
The absorption spectrum looks different because it shows the colors of light that are absorbed by the element (the black likes in the absorption spectrum) rather than the colors of light being emitted by the element (the colored lines in the emission spectrum).
A photon of light is absorbed by the atom, causing the electron to transition from a lower energy level to a higher energy level. This results in a line on the absorption spectrum.
Identify the aspects of atomic structure that determine the electrostatic forces that control the size of an atom. (evidence)
Effective nuclear charge
Shell/principal quantum number of valence shell electrons
Which e/m radiation has wavelengths on the order of the sizes of atoms?
x-rays
Which has the longest wavelength?
xrays, visable, or infrared?
infrared
Which has the highest frequency?
xrays, visable, and infrared
xrays
what is visible light?
white light (the rainbow), emitted from the sun
de Broglie statement
de Broglie: all matter has wave properties and, therefore, a wavelength λ.
how many kg are in a J?
1 J = 1 kg m2 s–2
If an electron is removed from an atom, what is the charge of the ion?
positive charged cation
which has larger radius Li or Li+
Li (Lithium) has a larger radius than Li⁺ (Lithium ion) because Li+ loses an electron. With the same amount of protons but less electrons causes the attractive force to be stronger.
