Apologia Chemistry Module 07 Review

Front 1

Some 2,400 years ago Democritus first theorized about atoms... What was his theory?

Back 1

Democritus thought that each type of atom had its own unique shape: some were triangles; some were pentagons; and some were squares. These shapes, he thought, fitted together like jigsaw puzzle pieces to form molecules

Front 2

In the second half of the 1800s, William Crookes performed experiments with atoms. What did he do and what did he conclude from his first experiment?

Back 2

Crookes studied how certain gases behaved when exposed to electricity. He studied this by filling a glass tube with a tiny amount of gas. He would then hook the tube up to a battery, allowing electricity to pass through it. Crookes noticed that when the battery was hooked up, the end of the tube would glow a faint greenish-yellow color. When the battery was not hooked up, there would be no glow.

He concluded from this experiment that particles were being produced in the gas by the battery. These particles, he theorized, traveled through the gas and hit the end of the tube, causing the glow that he saw.

Front 3

Why did Crookes call the particles in his tube "cathode rays?"

Back 3

Since these particles always seemed to travel from the post on the battery called the “cathode,” Crookes called them “cathode rays.”

Front 4

What did Crookes do to try to confirm his theory that the cathode rays were actually particles? Did he prove his theory?

Back 4

In order to try to confirm his theory that these cathode rays were actually particles, he put an obstacle between the cathode and the end of the tube. He reasoned that if these cathode rays were, in fact, particles, they would be blocked by the obstacle and would not hit the end of the tube. Since they could not hit the end of the tube, they would not cause a glow.

Yes, his experiment proved it by creating a "shadow" on the end of the cathode ray tube.

Front 5

What did English chemist, J. J. Thomson try and determine and disprove?

Back 5

He tried to determine more about the nature of these cathode rays. Thomson's work ended up demonstrating that one of John Dalton's main assumptions concerning atoms was wrong.

Dalton had developed a theory that assumed that atoms were indivisible. He reasoned that since atoms were the smallest unit of matter, they could not be broken down into smaller pieces.

However, Thomson showed that regardless of what type of gas he put in the Crookes tube, cathode rays were always formed. He further showed that these particles were precisely the same, whether he filled the Crookes tube with helium, argon, nitrogen, or any other gas. This led Thomson to believe that these cathode rays were a part of every atom and that his experiments were breaking them away from the atom.

Thomson's idea that cathode rays were actually small parts of the atom was further boosted by his discovery that these rays were electrically charged. (negatively charged)

Front 6

Scientists know how electrical charges behave and how they generate electricity, and thus they have a clear understanding of what an electrical charge is and why it exists! True, or False?

Back 6

False. Unfortunately, scientists do not have a clear understanding of what an electrical charge is or why it exists. We know how electrical charges behave and how they generate electricity, but we really don't know much about what they are or where they come from.

Front 7

What are the two types of electrical charge?

Back 7

positive and negative

Front 8

If something has no charge, what is it called?

Back 8

neutral

Front 9

What makes something have a positive or negative charge?

Back 9

We don't know what makes something negative or positive in charge, but we know that these charges do indeed exist.

Front 10

The vast majority of matter on earth has which type of electrical charge?

Back 10

neutral.... or no charge!

Front 11

What is the first rule of electrical charge?

Back 11

Like charges repel one another.

Front 12

If you have as many positive electrical charges as there are negative ones, what is the end result?

Back 12

Just as -1 + 1 = 0, the negative and positive charges cancel each other out to give no charge at all.

Front 13

What is the second rule of electrical charge?

Back 13

Opposite charges attract each other.

Front 14

A certain substance starts with no overall electrical charge because it had an equal number of positive and negative charges in it. However, it picks up a few extra negative charges from a different source, so what is the end result electrical charge?

Back 14

If there is more of one charge than the other, the substance takes on that electrical charge. (-2 + 1 = -1).

So in this case, it would have a negative charge.

Front 15

What part did the battery in Crookes' experiment play in terms of electrical charge?

Back 15

When the Crookes tube was filled with gas, the gas was always electrically neutral. Thus, the gas had as many negative charges as positive charges. However, when the battery was hooked up to the Crookes tube, negatively charged particles were produced. Thomson reasoned that those negatively charged particles must have been taken from the atoms of gas.

Front 16

True or False? Atoms can be broken down into smaller components.

Back 16

True. If negative charges could be pulled from an electrically neutral atom, then the atom must be composed of smaller components. Some of those components must be negatively charged, and the other components must be positively charged. Since the gas atoms started out neutral, there must be an equal number of positively and negatively charged components in atoms. Crookes and Thomson showed that these charged components could be separated; thus, atoms could be broken down into smaller components.

Front 17

Thomson later did a famous experiment which actually measured the ratio of the charge and mass of negatively charged particles. What did it prove?

Back 17

He showed that no matter what atoms they were pulled from, these particles always had the same ratio of charge to mass. He thus considered these particles to be something common to all atoms, and he called them electrons.

Front 18

By 1932, scientists had determined that atoms were made up of three components. What are they?

Back 18

protons, neutrons, and electrons

Front 19

What makes one atom different from another atom?

Back 19

What makes one atom different from another atom is the number of protons, neutrons, and electrons contained in it.

Front 20

All atoms have an equal number of electrons and protons. True? Or False?

Back 20

True.

Front 21

How many protons and electrons are in the following atom?

Chromium (Cr)

Back 21

On the periodic chart, Cr has an atomic number of 24. Thus, a Cr atom has 24 protons and 24 electrons.

Front 22

How many protons and electrons are in the following atom?

Beryllium (Be)

Back 22

On the periodic chart, Be has an atomic number of 4. Thus, a Be atom has 4 protons and 4 electrons.

Front 23

How many protons and electrons are in the following atom?

Lanthanum (La)

Back 23

On the periodic chart, La has an atomic number of 57. Thus, a La atom has 57 protons and 57 electrons.

Front 24

Each box in the periodic table represents a single type of atom. True or False?

Back 24

False. Each box on the periodic chart represents several different atoms, but all of those atoms are part of the same element. (isotopes)

Every element is composed of several isotopes. This means that every box on the periodic chart actually represents several different isotopes.

Front 25

Isotopes behave identically in their chemistry.... So what is the main difference between them?

Back 25

the main difference between them is their mass.

Front 26

How do you write out the full name of an element and include the mass number? (use carbon as your example)

Back 26

When writing out the full name of the element (carbon), you can include the mass number by adding a hyphen followed by the mass number. Thus, we could call this isotope “carbon-12.” On the other hand, if you just want to write the symbol for carbon (C), then the mass number must appear as a superscript which precedes the symbol. Thus, the symbol for this isotope would be 12*C. (the symbol * equals superscript)

Front 27

If several atoms all have the same number of protons (the same atomic number) but different numbers of neutrons, they are all isotopes.
True? Or False?

Back 27

True.

Front 28

What is the name and symbol of an atom which is made up of 10 protons, 10 electrons, and 11 neutrons?

Back 28

Since the atom has 10 protons, its atomic number is 10. Looking at the periodic chart, neon (Ne) is the element with that atomic number. To give the full name, we must also include its mass number. Since it has 10 protons and 11 neutrons, its mass number is 10 + 11 = 21. Therefore, the name of the atom is neon-21, and its symbol is 21*Ne.

Front 29

How many protons, electrons, and neutrons make up a 129*I atom?

Back 29

To determine the number of protons and electrons, all we have to do is find I (Iodine) on the periodic chart. We see that its atomic number is 53. That means that it has 53 electrons and 53 protons. In addition, we are given its mass number, which represents the total number of protons and neutrons. If the atom has 53 protons, and its total number of protons and neutrons is 129, then the number of neutrons it has is 129 - 53 = 76 neutrons.

76 neutrons!

Front 30

Give the number of protons, electrons, and neutrons that make up the following atom:

40*Ar

Back 30

Looking at the chart, Ar has an atomic number of 18. This means it has 18 protons and 18 electrons. Its mass number, according to the problem, is 40. If it has 40 total protons + neutrons, and it has 18 protons, then it has 40 - 18 = 22 neutrons.

Front 31

Give the number of protons, electrons, and neutrons that make up the following atom:

Chlorine-37

Back 31

Looking at the chart, chlorine (Cl) has an atomic number of 17. This means it has 17 protons and 17 electrons. Its mass number, according to the problem, is 37. If it has 37 total protons + neutrons, and it has 17 protons, then it has 37 - 17 = 20 neutrons.

Front 32

Give the number of protons, electrons, and neutrons that make up the following atom:

139*La

Back 32

Looking at the chart, La has an atomic number of 57. This means it has 57 protons and 57 electrons. Its mass number, according to the problem, is 139. If it has 139 total protons + neutrons, and it has 57 protons, then it has 139 - 57 = 82 neutrons.

Front 33

An atom has 34 protons, 34 electrons, and 41 neutrons. What is its full name and symbol?

Back 33

First, since it has 34 protons, we know its atomic number is 34. The element on the chart which has an atomic number of 34 is selenium (Se). The problem states that this atom also has 41 neutrons; thus, its mass number is 34 + 41 = 75. Therefore, the atom is selenium-75, which has a symbol of 75*Se.

Front 34

Why is isotopic enrichment so difficult to perform?

Back 34

The only real difference between 235*U and 238*U is the mass. Chemically, they behave in exactly the same way. This makes it nearly impossible to separate the two for isotopic enrichment.

Front 35

What was the “Manhattan Project” and what time period was it important?

Back 35

During World War II, scientists in a group called the “Manhattan Project” were able to figure out a way to partially separate 238*U from 235*U. The process is incredibly inefficient, but it is able to take natural uranium and enrich the 235*U content from 0.7% to a little more than 7%, which is enough to fuel a nuclear bomb.

Front 36

What was the Plum Pudding Model of the Atom?

Back 36

After electrons were first discovered (and before neutrons were discovered), chemists thought that an atom looked a lot like a dish of plum pudding. In essence, chemists thought that the atom was made up of a pudding-like substance which had positive electrical charge. Electrons, they thought, were suspended in this “pudding” like plums, making the entire atom electrically neutral.

Front 37

If the model is inconsistent with the experimental data, then the model is obviously incorrect. If, however, the model is consistent with the data, then we know that the model is right. That is why models are so important to science. True or False?

Back 37

False!! If the model is consistent with the data, then we don't necessarily KNOW that the model is right, but at least we know that it MIGHT not be wrong.

Front 38

Who determined that the Plum Pudding Model of the Atom was wrong? How did he/she determine that?

Back 38

Ernest Rutherford (Thomson's student) designed the perfect experiment to test the validity of the plum pudding model. In earlier experiments, Rutherford had identified a type of radiation called “alpha particles.” He determined that these alpha particles have positive electrical charges and are emitted by certain isotopes that were called “radioactive.” Rutherford decided that what he needed to do was to shoot some of the alpha particles at a thin metal foil. He reasoned that if the plum pudding model of the atom were correct, the alpha particles should pass through the foil without changing their direction much. After all, the positive alpha particles would be attracted by the negative “plums” (the electrons) but repelled by the positive “pudding.” Since the plums were randomly distributed around the pudding, the alpha particles would be attracted and repelled in many different directions, and the result would be that they would travel straight through the atom.

Rutherford surrounded the target with a zinc sulfide screen. When an alpha particle hits zinc sulfide, a bright green glow is emitted. Thus, Rutherford could see where the alpha particles went after they collided with the gold foil by simply looking at where the zinc sulfide screen was glowing. Wherever it was glowing very brightly, many alpha particles were hitting it. Wherever it was glowing faintly, only a few alpha particles were hitting it.

Rutherford ran his experiment and, to his astonishment, the zinc sulfide screen glowed everywhere! To be sure, the glow was the brightest directly in front of the target, but alpha particles were hitting the zinc sulfide screen at all other positions as well. Rutherford noticed that the farther away from the front of the target, the dimmer the glow was. Nevertheless, alpha particles were colliding with the target and bouncing in every conceivable direction, thus proving the plum pudding model of the atom was wrong.

Front 39

Describe Rutherford's model of the atom.

Back 39

In Rutherford's model, the protons were clustered together at the center of the atom, called the nucleus while the electrons orbited around the nucleus in the same way that the planets in our solar system orbit around the sun. This similarity between his view of the atom and our solar system is why Rutherford called it the “planetary model” of the atom.

Front 40

What is missing from Rutherford's planetary model of the atom, and why?

Back 40

Rutherford's model does not contain neutrons. This is because neutrons had not been discovered yet.

Front 41

Rutherford's model was not consistent with a current theory in physics. What was that theory?

Back 41

Physicists during Rutherford's time had noticed that whenever electrically charged particles moved in circles, they emitted light. When they emitted this light, they would lose energy and thus slow down. According to this theory, then, the Rutherford model of the atom would never be stable. This theory would predict that the electrons moving in a circle would continue to emit light until they lost all of their energy. This theory would additionally predict that the electrons, as they lost energy, would spiral towards the protons in the atom until they collided with them.

Front 42

Light is not matter. True or False?

Back 42

True. Since light neither has mass nor takes up space, it is not matter.

Front 43

There are currently two models for light. What are they?

Back 43

the particle model and the wave model

Front 44

Currently, theory states that light is both a particle and a wave. How can this be?

Back 44

The situation that light is in determines whether it will behave as a particle or a wave.

(particle/wave duality theory)

Front 45

Since wavelength is a measurement of distance, its units will be meters, centimeters, etc. The wavelength of a wave is usually given what symbol?

Back 45

The wavelength of a wave is usually given the symbol “λ” (the lower-case Greek letter “lambda”)

Front 46

What symbol is amplitude usually given?

Back 46

Amplitude is usually given the symbol “A”

Front 47

The wavelength of light corresponds to what?

Back 47

The wavelength of light corresponds to its color.

Apparently white light, once it passes through a prism, turns into a rainbow of colors. This is because a prism breaks light into its different wavelengths. Each color you see represents a range of wavelengths

Front 48

Why do the colors of a rainbow always appear in a certain order?

Back 48

The wavelength of light determines how the light passes through a prism. Since wavelength means color, the colors of light in a rainbow always appear in the same place, because they always pass through the water droplets (prism) in the air in exactly the same way.

Front 49

List the colors of the visible spectrum in order, starting with the color with the largest wavelength....

Back 49

Red
Orange
Yellow
Green
Blue
Indigo
Violet

Front 50

What does the amplitude of light mean?

Back 50

Well, simply put, the amplitude of a light wave tells us how bright the light wave is. A dim light bulb produces light waves that have a small amplitude, while a bright light bulb produces light with a large amplitude. Thus, if I had two light bulbs of the same color, but one was dimmer than the other, I could say that the light bulbs were both producing light of the same wavelength but of different amplitudes. Similarly, if I had two light bulbs of different color but the same brightness, I could say that the light bulbs produced light waves of equal amplitude but of different wavelengths

Front 51

Which travels faster: light or sound?

Back 51

Light... think: lightening strikes vs. thundercrash

Front 52

Regardless of the color or the brightness of light, its speed in air is always........ what?

Back 52

Regardless of the color or the brightness of light, its speed in air is always 3.0 x 108 m/s.

Front 53

c = 3.0 x 108 m/s

What does the "c" mean?

Back 53

the c = the speed of light, which is a physical constant

Front 54

f = c/λ

What does this mean?

Back 54

In this equation, “f” stands for frequency, “c” is the speed of light, and “λ” stands for wavelength. Notice how the units work out. Speed is in m/s, while wavelength is in m. When I divide the speed by the wavelength, meters cancel, and I am left with 1/s, which is also called “Hertz.”

Front 55

(c = 3.0 x 108 m/s, “nano” = 10-9)

What color is light that has a frequency of 6.4 x 10*14 Hz?

Back 55

this light is blue light.

(look in the book at OYO question after The Nature of Light for explanation)

Front 56

(c = 3.0 x 108 m/s, “nano” = 10-9)

If red light has wavelengths from 7.00 x 10*2 nm to 6.50 x 10*2 nm, what is the frequency range for red light?

Back 56

So the frequency range is between 4.3 x 10*14 Hz and 4.6 x 10*14 Hz.

(look in the book at OYO question after The Nature of Light for explanation)

Front 57

When wavelength is large, frequency is large. When wavelength is small, frequency is small. True or False?

Back 57

FALSE. Just the opposite. When wavelength is large, frequency is small. When wavelength is small, frequency is large.

Front 58

Memorize this:

As a light wave's frequency increases, its energy increases.
As its frequency decreases, its energy decreases.

Back 58

Memorize this: (part 2)

As a light wave's wavelength increases, its energy decreases.
As its wavelength decreases, its energy increases.

Front 59

The only difference between the light waves from radio signals or our televisions and the ones we see with our eyes is the wavelength. True or False?

Back 59

True.

Front 60

Microwave ovens use light in order to cook food. True or False?

Back 60

True.

Front 61

When visible light strikes your eye, it deposits _____________ there. What is this used for?

Back 61

When visible light strikes your eye, it deposits its energy there. Your eye uses that energy to transmit signals to the brain, and that's what causes you to see.

Front 62

What is so bad about ultraviolet light?

Back 62

It turns out that light with wavelengths shorter than visible light has enough energy to kill living tissue. Thus, when ultraviolet light strikes your skin, it can kill some of your cells. If your skin is exposed to too much ultraviolet light, a large number of cells will die, and you will get a burn on your skin, which we call a sunburn.

Front 63

What part does the ozone layer play in regards to ultraviolet light?

Back 63

The ozone layer is part of the filtering system that shields us from this destructive type of light. Although it is incredibly efficient, the ozone layer does allow some ultraviolet light through, and that is why you can get sunburned if you stay out in the sun too long without protection.

Front 64

Even though it is fun to see inside your body via x-rays, why is it not a good idea to take too many x-rays of your body???

Back 64

Gamma rays and X-rays are more energetic than ultraviolet light, so they are even more dangerous to living tissue. Although these X-rays will kill some living tissue, that risk is worth the benefit of being able to diagnose internal problems without surgery. Thus, as long as you do not get X-rays frequently, your risk associated with the nature of the light is low, and the benefit you get from the diagnosis of internal problems is high.

Front 65

As a light wave's frequency increases, its energy does as well. The mathematical equation that governs this relationship is:

Back 65

E = h · f

In this equation, “E” is the energy of the light wave and “f” is the frequency. The symbol “h” refers to another physical constant, called Planck's constant. This constant allows us to relate energy and frequency. Its value is 6.63 x 10 -34 J/Hz.

Front 66

How do the cones in your eyes work?

Back 66

The cones in your eye respond to certain specific energies of light. Some cones are sensitive only to low-energy visible light (red), while others are sensitive to medium-energy visible light (green), while still others are sensitive to the highest energy light (blue).

When colored light hits these cells, they will only send signals to the brain if the light that they are sensitive to is hitting them. Thus, if blue light hits your eyes, the cones that respond to blue light will send signals to your brain, but the cones that respond to green and red light will not.

As a result, your brain interprets the signals coming from your eyes and puts a blue image in your head.

Front 67

If you are looking at a mixture of light (say... purple), then how do your cones work then?

Back 67

If a mixture of colors hits your eyes, the cones send signals to your brain in proportion to the amount of light to which they are sensitive. (the purple hue is more blue than red for example)

All of the colors that you see are really the result of the brain receiving signals from three types of cone cells and adding those three primary colors (red, green, and blue) with the weight indicated by the amount of signal coming from each type of cell.

Front 68

Cones get tired pretty quickly, and when they have sent the same signal to the brain for a period of several seconds, they eventually just shut off. True or false?

Back 68

True.

Front 69

The way we perceive color, then, is based on the energy of the light that hits our eyes. What color of light contains all energies?

Back 69

White light.

Front 70

If you need to determine the atoms in a substance, one way to do it is to heat the substance up. What does this do?

Back 70

If you need to determine the atoms in a substance, one way to do it is to heat the substance up and determine the wavelengths of light that are emitted. Since each element in the substance will emit its own unique wavelengths of light, you should be able to match the elements with their respective wavelengths and thus determine the elemental composition of the substance.

Front 71

When you add energy to an atom by heating them or passing electricity through them, what does this show us?

Back 71

The atoms emitted light, some of which was in the visible spectrum. The interesting thing, however, was that each type of atom seemed to emit its own unique color or colors of visible light. For example, when mercury atoms were heated, a pale blue glow would occur. On the other hand, when electricity was passed through neon atoms, a bright pink light was emitted.

Front 72

Most colors that we see are the result of a range of wavelengths. How are the wavelengths of atoms different?

Back 72

atoms seem to emit individual wavelengths of light. This is very unusual. Most colors that we see are the result of a range of wavelengths, not a handful of individual ones.

Front 73

Rutherford assumed that all of the electrons orbited around the protons in a single orbit, and he had no knowledge of neutrons. How did Bohr's model of an atom differ from Rutherford's model?

Back 73

Bohr suggested, instead, that there were several possible orbits that the electrons could be in. Furthermore, he suggested, the electrons could jump from one orbit to another. In addition, by the time Bohr had come up with his model of the atom, the neutron had been discovered by Chadwick. Bohr put the neutrons and protons together in the center of the atom and called it the nucleus.

Front 74

How do the electrons jump from one orbit close to the nucleus to another farther away?

Back 74

the orbits that are farther away from the nucleus require more energy. Thus, if an electron wants to jump from an orbit that is close to the nucleus (orbit #1) to an orbit that is far away from the nucleus (orbit #3), it can do so by absorbing some energy. This is what happens when you heat a substance or pass electricity through it. The atoms in the substance absorb energy, allowing the electrons to move from an orbit close to the nucleus to an orbit far away from the nucleus.

Front 75

If an electron would like to move from an orbit far away from the nucleus to an orbit close to the nucleus, what must happen?

Back 75

The electron must somehow get rid of its extra energy. In other words, while in the orbit that is far from the nucleus, the electron has more energy than it would in an orbit that is close to the nucleus. In order to get to that closer orbit, the electron must find a way to get rid of that excess energy. The way it does this is to emit light.

Front 76

When you heat up a substance and its atoms emit light, you are witnessing a two-step process. What are they?

Back 76

1. Electrons in the atoms first gain energy (because of the heat) and move into orbits that are far away from the nucleus.

2. Then, after they sit in the far orbits for a short time, they decide they would rather be back in their original orbit, so they emit light in order to get rid of the excess energy and go back to their original orbit.

Front 77

Bohr assumed that an atom's electrons could only be in certain orbits, which was a quantum assumption. What was his other quantum assumption based upon the first?

Back 77

Since each orbit has a specific energy associated with it, this means that an electron can only have certain energies. If it is in orbit #1, it has a certain energy (E^1). If it is in orbit #2 it has another, higher energy (E^2). But, since the electron can never be in between orbit #1 and orbit #2, it can also never have any of the energies between E^1 and E^2.

Bohr said there were only certain energies that the electrons in the atom could have. This, another quantum assumption in Bohr's model, is actually the one we now see as critical to understanding atomic structure.

Front 78

Summarize the three main concepts of Bohr's model of the atom that we still follow and use today.

Back 78

1. Chemists still believe that electrons orbit around the nucleus and that there are many, many different orbits that electrons can continually jump between.

2. We still say the electrons need energy to go into orbits that are far away from the nucleus, and they need to release energy in the form of light in order to get back to orbits close to the nucleus.

3. Bohr's quantum assumption regarding energy is still followed. We assume that there are only certain energies that an electron in an atom can have. Thus, in one orbit, an electron might have an energy of E1, and in the next higher orbit have an energy of E2, but the electron can have no energy in between E1 and E2.

Front 79

The quantum mechanical model mainly differs from the Bohr model in what way?

Back 79

It differs in the type of orbits that the electrons can occupy.

In the Bohr model, electrons orbit the nucleus in circular orbits. In the quantum mechanical model, this is not the case.

The way we theorize that electrons orbit the nucleus today is a little more complex. We assume that electrons do not orbit in fixed circles, but instead orbit in “clouds” we call orbitals. When orbiting in a fixed circle, the electron is always the same distance from the nucleus. In today's model of the atom, electrons can, at different times, be at different distances from the nucleus but still be in the same orbital. However, if you were to watch the electron for quite a while, you would see that it stayed within a certain boundary.

Front 80

The other difference between the Bohr model and the quantum mechanical model is the shape of the orbitals. Explain.

Back 80

In the Bohr model, all electrons had to orbit in circles. In the quantum mechanical model, there are several differently shaped orbitals that electrons can use to orbit around the nucleus. The shapes of these orbitals can become very complex.

(We are only going to deal with the first three shapes...)

Front 81

Name the first three types of electron orbitals.

Back 81

s orbital

p orbital

d orbital

Front 82

Describe the features of the s orbital.

Back 82

The simplest type of orbital an electron can occupy is spherically shaped, with the nucleus at the center.

An electron that occupies an s orbital can be anywhere inside the sphere, but cannot venture outside of the sphere.

Electrons can orbit far away from the nucleus or close to the nucleus depending on their energy. If an electron orbits far away from the nucleus in a spherical pattern, it will occupy a large s orbital. If it orbits close to the nucleus in a spherical pattern, it will occupy a small s orbital.

The way we delineate small s orbitals from large ones is the number that appears next to the orbital letter. A 1s orbital is always smaller than a 2s orbital.

The farther away from the nucleus the electron is, the more energy it must have. Thus, the electrons in a 1s orbital have less energy than the electrons in a 2s orbital.

Since the size of the orbital determines the energy of the electrons, the number that appears next to the orbital letter is often called the “energy level” of the atom.

Front 83

Describe the features of the p orbital.

Back 83

This orbital is dumbbell-shaped with the nucleus in the center of the dumbbell.

Just as there are different sizes of s orbitals that correspond to different energies of electrons, there are also different sizes of p orbitals that correspond to different electron energies.

One interesting thing about p orbitals is that there are none of them on the first energy level of the atom. Thus, there is no such thing as a 1p orbital. The first p orbital is the 2p orbital.


Another interesting thing about p orbitals is that for every energy level (or size of the orbital), there are actually three different p orbitals. They are all shaped the same, but they are oriented differently in space. In other words, each of the three p orbitals is oriented along a different axis in three-dimensional space.

Front 84

True or False? There can be three different p orbitals for each energy level.

Back 84

True. There are three 2p orbitals, three 3p orbitals that are all bigger than the 2p orbitals, three 4p orbitals that are all bigger than the 3p orbitals, and so on.

Front 85

Explain the features of the d shaped orbital.

Back 85

The shapes of the d orbitals are very complex. What you need to know about d orbitals is that there are five different d orbitals for each energy level. Also, the lowest-energy (smallest) d orbitals are located on the third energy level; thus, the first set of d orbitals is the 3d orbitals.

Front 86

Within the same atomic energy level, the different shaped orbitals require different amounts of energy. List them in order from least amount of energy required to the most required.

Back 86

It is rather easy for an electron to whirl around the nucleus in a sphere, so s orbitals are the lowest-energy of all orbitals. On the other hand, a dumbbell is a little more complex, so p orbitals require more energetic electrons than s orbitals. Finally, since d orbitals have even more complex shapes, they require more energy than either s or p orbitals.

So the order would be:

s orbitals
p orbitals
d orbitals

Front 87

Each individual orbital can only hold two electrons. After two electrons, the orbital is full, and any new electrons must find another orbital to fill.

How does this rule relate to the p and d orbitals?

Back 87

Since there are three different p orbitals on an energy level, six electrons could fit into all three p orbitals on that energy level (two per orbital).

In the same way, since there are five different d orbitals on an energy level, then up to 10 electrons can go into the d orbitals of a given energy level.

Front 88

If an electron orbits as close as it possibly can to the nucleus, then it must do so in which type of orbital?

Back 88

If an electron is orbiting as close as it can to the nucleus, we say that it is in the first energy level. On the first energy level, only one type of orbital exists: the s orbital. Thus, if an electron orbits as close as it possibly can to the nucleus, then it must do so in a spherically shaped orbit (1s orbital).

Front 89

If the electron goes a little farther away from the nucleus and goes up to the next energy level... which orbitals can it occupy?

Back 89

If the electron goes up to the next energy level (#2), it is a little farther away from the nucleus and has the option of orbiting the nucleus in a spherical orbit (2s orbital) or in one of any three dumbbell-shaped orbits (2p orbitals).

Front 90

If the electron goes away from the nucleus even further into the third atomic energy level, which orbitals can it occupy?

Back 90

Then it has the choice of a spherically shaped orbit (3s orbital), one of three dumbbell-shaped orbits (3p orbitals), or one of five more complex orbits (3d orbitals).

Front 91

Summarize the quantum mechanical model of the atom.

Back 91

1. The atom has a nucleus of tightly packed protons and neutrons.

2. Its electrons whirl around that nucleus in several different energy levels (1, 2, 3...).

3. Depending on the energy level, the electrons have several different shapes of orbitals (s, p, d, and other orbitals) that they can use to travel around the nucleus.

4. Each specific energy level (1, 2, 3...) has its own energy requirements for the electrons.

5. In addition, each orbital shape has additional energy requirements for the electrons.

6. Based on the energy that an electron has, it will go into the energy level and orbital whose energy requirements it best meets.

7. Each orbital has a maximum capacity of two electrons.

Front 92

Remember when the textbook discussed the light that was emitted from excited atoms? It said that when an atom is heated, its electrons absorb some of that kinetic energy and use it to travel to an orbital farther away from the nucleus. The electron then emits light to go back down to its original orbital. Why did it do that? Once it jumped to the higher-energy orbital, why didn't it just stay there?

Back 92

It didn't stay there because once it was in a higher-energy orbital, it was not in its lowest possible energy state. Thus, the electron released energy in the form of light in order to get back to its lowest energy state.

***All forms of matter try to stay in their lowest possible energy state.***

Front 93

The first energy level should fill up before electrons go into the second energy level. In the same way, the second energy level should fill all of its orbitals before electrons go into the third energy level. This is all true. However, at the third energy level, things are a bit different... How so?

Back 93

The d orbitals in the third energy level take a lot of energy. In fact, the 3d orbitals require more energetic electrons than the 4s orbitals. Thus, if I had enough electrons to fill up the 3s and 3p orbitals, I would then start filling the 4s orbitals before I go back and fill up the 3d orbitals.

Front 94

As atoms get more and more electrons, the order of orbitals gets even more complicated. In order to be able to determine the proper electron configuration of any atom, then, we need to have some way of determining the order in which orbitals fill up. So, how is this done?

Back 94

By using the periodic chart.

The periodic chart gets its weird shape because of electron configurations. By looking at the chart, you can easily determine an atom's electron configuration;

Front 95

True or false? There is actually an f orbital....

Back 95

True!!!

Front 96

To use the arrangement of the periodic chart to determine electron configurations, what do you do? Also, how is the d orbital different?

Back 96

You have to look at the periodic chart and find the atom whose electron configuration you wish to determine. Then, starting with hydrogen, walk through the chart, assigning electrons in each orbital one box at a time. When you get to the box that represents the atom you are interested in, you're done.

In the d orbital block, the row that the elements are on is actually one number higher than the energy level that the electrons are in. So, when filling up d orbitals, subtract one from the row number in order to get the energy level.

Front 97

What do the superscripts in the electron configuration represent?

Back 97

the number of electrons in each orbital

Front 98

How do you check your electron configuration using the superscripts?

Back 98

The sum of the superscripts must equal the number of electrons in the atom.

Front 99

Write the electron configuration for phosphorus.

Back 99

In order to get to P (#15), we have to go through several boxes. We must walk through all of row 1, which has two boxes in the s orbital block. That gives us 1s*2. We must also go through all of row 2, which has two boxes in the s orbital block and six boxes in the p orbital block. This gives us 2s*2 2p*6. We also must go through the s orbital block in row 3. This gives us 3s*2. Finally, to get to P, we go through three boxes in the row 3, p orbital block. This gives us 3p*3. The electron configuration, then, is:


1s*2 2s*2 2p*6 3s*2 3p*3
(when writing this, normally you would not have spaces between .... but since I can't do superscripts it looks too confusing. Thus, spaces!)

The superscripts add up to the number of electrons in a P atom (15), so our answer checks out.

Front 100

Write the electron configuration for arsenic (As).

Back 100

To get to As, we have to walk through several boxes on the chart. We must walk through all of row 1, which has two boxes in the s orbital block. That gives us 1s*2. We must also go through all of row 2, which has two boxes in the s orbital block and six boxes in the p orbital block. This gives us 2s*2 2p*6. Row 3 has two boxes in the s orbital group and six in the p orbital box, so that gives us 3s*2 3p*6. Row 4 has two boxes in the s orbital block and 10 in the d orbital block that we must go through. Remembering that for d orbitals we must subtract one from the row number, this gives us 4s*2 3d*10. Finally, to get to As, we must walk through three boxes in the row 4, p orbital block. That gives us 4p*3. Therefore, the final electron configuration is:


1s*2 2s*2 2p*6 3s*2 3p*6 4s*2 3d*10 4p*3

Our superscripts add up to 33, so our answer checks out.

Front 101

What is the electron configuration of ruthenium (Ru)?

Back 101

To get to Ru, we have to walk through several boxes on the chart. We must walk through all of row 1, which has two boxes in the s orbital block. That gives us 1s*2. We must also go through all of row 2, which has two boxes in the s orbital block and six boxes in the p orbital block. This gives us 2s*2 2p*6. Row 3 has two boxes in the s orbital group and six in the p orbital box, so that gives us 3s*2 3p*6. Row 4 has two boxes in the s orbital block, 10 in the d orbital block, and six in the p orbital block. Remembering that for d orbitals we must subtract one from the row number, this gives us 4s*2 3d*10 4p*6. In row 5, we must go through the two boxes in the s orbital block, so that gives us 5s*2. In row 5, we also have to go through six d orbital boxes to get to Ru. Once again, we subtract one from the row number with d orbitals, so this gives us 4d*6. In the end, then, our electron configuration is:


1s*2 2s*2 2p*6 3s*2 3p*6 4s*2 3d*10 4p*6 5s*2 4d*6

Our superscripts add up to 44, so our answer checks out.

Front 102

What is the abbreviated electron configuration for I?

Back 102

To abbreviate an electron configuration, we just find the nearest 8A element that has a lower atomic number than the atom we are interested in. In this case, Kr fits the bill. We can then say the only differences between I and Kr are that there are two more boxes in the row 5, s orbital block, 10 boxes in the row 5, d orbital block, and five boxes in the row 5, p orbital block. Thus, we can write the electron configuration as:


[Kr]5s*2 4d*10 5p*5

Front 103

Write the electron configuration for Ga in abbreviated form.

Back 103

The nearest 8A element that has a lower atomic number than Ga is Ar. The only differences between Ga and Ar are that there are two boxes in the row 4, s orbital group, 10 boxes in the row 4, d orbital group, and one box in the row 4, p orbital group. Thus, the abbreviated electron configuration for Ga is:


[Ar]4s*2 3d*10 4p*1

Front 104

The atom is mostly made up of empty space. True? Or False?

Back 104

True. Since matter is made up entirely of atoms, this also means that matter is mostly empty space.

Front 105

True or False? Protons are more massive than neutrons.

Back 105

False. It turns out that the neutron is just slightly more massive than the proton. It's a good thing, too. Nuclear chemistry experiments indicate that the neutron must be more massive than the proton, or it would not exist!

Front 106

The mass of the electron is quite crucial for the existence of life. Why?

Back 106

The electron mass is just right to produce electron orbitals that are neither too big nor too small. The balance between ionic and covalent molecules is perfect for the formation and existence of life, due to the mass of the electron. Estimates indicate that a variation of as little as 2% in the mass of the electron would make it impossible for life as we know it to exist!

Front 107

Protons and electrons have perfectly balanced charges. That's a good thing, too, because calculations indicate that if the charges were out of balance by as little as 0.00000001%, all of the atoms in our bodies would instantaneously explode!


What is really amazing about the perfect charge balance between the proton and the electron is that these two particles are incredibly different. How are they different?

Back 107

The proton is about 2,000 times heavier than the electron. In addition, the proton seems to be made up of three smaller particles we call “quarks,” while as far as we can tell, the electron is not made up of any smaller particles. Nevertheless, despite the fact that these two particles are quite different, they have perfectly balanced charges!

Front 108

What was Rutherford's main contribution to the study of atomic structure?

Back 108

Rutherford experimentally determined that the plum pudding model of the atom was incorrect and
proposed his own model of the atom, which we call the planetary model. It became the foundation
upon which Bohr’s model was built.

Front 109

What do we call the experimental apparatus that William Crookes used in his experiments, and what did he discover with it?

Back 109

It is called a Crookes Tube, and he used it to discover cathode rays, which were later determined to be electrons.

Front 110

If two electrically charged particles repel one another, what can we conclude about their charges?

Back 110

Like charges repel each other. Thus, the particles must have the same type of charge.

Front 111

If a substance has 15 positive charges and 11 negative charges, is it positively charged, negatively
charged, or neutral?

Back 111

If a substance has an imbalance of charges, it takes on the charge that is more plentiful. Thus, this
substance will be positively charged.

Front 112

Where do the protons and neutrons exist in an atom?

Back 112

Protons and neutrons are tightly packed together in the nucleus of the atom.

Front 113

Why is it so hard to separate one isotope from another?

Back 113

Isotopes behave identically in terms of chemistry. It is therefore nearly impossible to separate them

Front 114

What were the differences between the plum pudding model of the atom and the planetary model of the atom?

Back 114

The plum pudding model of the atom had the positive and negative charges equally disbursed
throughout the entire atom. The planetary model, on the other hand, concentrated the positive charges
at the center of the atom and had the negative charges whirling around on the outside.

Front 115

If you have an orange light bulb and a violet one, which emits waves with the largest wavelength?
Which emits light of higher frequency? Which emits the higher energy light? (Do NOT look at the
figures in the book. Just remember the name that you wer told to remember from the book)

Back 115

Remember ROY G. BIV. This is the order of visible light wavelengths from the largest to the
smallest. Thus, the orange light bulb has larger wavelengths. When wavelength is large, however,
frequency is small; thus, the violet light has the highest frequency. The higher the frequency, the
higher the energy, so the violet light also has the highest energy.

Front 116

Two light bulbs emit light of the same color. One, however, is much brighter than the other. What
can you say about the wavelengths and amplitudes of the waves being emitted by each bulb?

Back 116

The wavelengths emitted by the lights are the same, but the brighter bulb emits waves of larger
amplitude.

Front 117

If an atom absorbs energy, what happens to its electrons? If the atom is emitting light, what are its
electrons doing?

Back 117

When atoms absorb energy, their electrons jump to higher energy orbitals. When they emit light, the electrons are dropping down into lower energy orbitals.

Front 118

What is the concept of ground state and why is it so important in chemistry?

Back 118

. The ground state of any substance is its lowest possible energy state. This is important in
chemistry because all matter strives to reach its ground state.

Front 119

The three fundamental particles that make up the atom are the proton, neutron, and electron. Order them in terms of decreasing mass.

Back 119

. The neutron is the heaviest, the proton is next, and the electron is the lightest. The proton and
neutron differ only slightly in mass, but the proton is 2,000 times heavier than the electron.

Front 120

Give the number of protons, electrons, and neutrons in the following atom...

90*Zr

Back 120

Looking at the chart, Zr has an atomic number of 40. This means it has 40 protons and 40 electrons. Its mass number, according to the problem, is 90. If it has 90 total protons + neutrons,
and it has 40 protons, then it has 90 - 40 = 50 neutrons.

Front 121

Give the number of protons, electrons, and neutrons in the following atom...

202*Hg

Back 121

Looking at the chart, mercury (Hg) has an atomic number of 80. This means it has 80 protons and 80 electrons. Its mass number, according to the problem, is 202. If it has 202 total protons + neutrons, and it has 80 protons, then it has 202 - 80 = 122 neutrons.

Front 122

Give the number of protons, electrons, and neutrons in the following atom...

58*Ni

Back 122

Looking at the chart, Ni has an atomic number of 28. This means it has 28 protons and 28 electrons. Its mass number, according to the problem, is 58. If it has 58 total protons + neutrons,
and it has 28 protons, then it has 58 - 28 = 30 neutrons.

Front 123

Give the number of protons, electrons, and neutrons in the following atom:

222*Rn

Back 123

Looking at the chart, Rn has an atomic number of 86. This means it has 86 protons and 86 electrons. Its mass number, according to the problem, is 222. If it has 222 total protons + neutrons and it has 86 protons, then it has 222 - 86 = 136 neutrons.

Front 124

Which of the following atoms are isotopes?

22*Na, 22*Ne, 23*Na, 22*Mg, 24*Na

Back 124

Isotopes have the name number of protons (thus the same atomic number and the same element symbol) but different numbers of neutrons (thus different mass numbers). Therefore 22Na, 23Na, and
24Na are isotopes. They all have the same number of protons (11), but they have 11, 12, and 13 neutrons, respectively. Remember, isotope is a relational term. It tells you how atoms relate to one another. There is not one “normal” atom with the rest being isotopes. Any group of atoms that all
have the same number of protons but different numbers of neutrons are isotopes.

Front 125

What is the symbol of the atom made up of 39 protons, 45 neutrons, and 39 electrons?

Back 125

If it has 39 protons and electrons, its atomic number is 39. The symbol that has atomic number 39 is Y. The mass number is the number of protons plus the number of neutrons, or 39 + 45 = 84. Thus, the symbol is 84*Y.

Front 126

One electron is in a 2p orbital while another is in a 3s orbital. Which has the higher energy? What shape is each electron's orbit?

Back 126

The electron in the 3s orbital has the higher energy, because energy level 3 is higher in energy than energy level 2. The electron in the 2p orbital is orbiting the nucleus in a dumbbell shape, while the 3s electron is orbiting the nucleus in a spherical shape.

Front 127

Give the full electron configuration of the following atom:

a. Ti

Back 127

To get to element Ti, we must go through row 1, which has two boxes in the s orbital block (1s2). We then go through all of row 2, which has 2 boxes in the s orbital block and 6 boxes in the p orbital
block (2s*2 2p*6). We also go through row 3, which has two boxes in the s orbital block and 6 in the p orbital block (3s*2 3p*6). We then go to the fourth row, where we pass through both boxes in the s orbital block (4s*2). Finally, we go through 2 boxes in the d orbital block. Since we subtract one from the row
number for d orbitals, this gives us 3d*2. Thus, our final electron configuration is:
1s*2 2s*2 2p*6 3s*2 3p*6 4s*2 3d*2

Front 128

Give the full electron configuration of the following atom:

S

Back 128

To get to element S, we must go through row 1, which has two boxes in the s orbital block (1s*2). We then go through all of row 2, which has 2 boxes in the s orbital block and 6 boxes in the p orbital
block (2s*2 2p*6). We also go through both boxes in the s orbital block of row 3, (3s*2). Finally, we go through 4 boxes in the p orbital block of row 3, giving us 3p*4. Thus, our final electron configuration
is:
1s*2 2s*2 2p*6 3s*2 3p*4

Front 129

Give the full electron configuration of the following atom:

Rb

Back 129

To get to element Rb, we must go through row 1, which has two boxes in the s orbital block (1s*2). We then go through all of row 2, which has 2 boxes in the s orbital block and 6 boxes in the p orbital
block (2s*2 2p*6). We also go through row 3, which has two boxes in the s orbital block and 6 in the p orbital block (3s*2 3p*6). We then go to the fourth row, where we pass through both boxes in the s orbital block, all 10 boxes in the d orbital block, and all 6 boxes in the p orbital block. Since we subtract one from the row number for d orbitals, this gives us 4s*2 3d*10 4p*6. Finally, we end up in the first box of the row 5, s orbital block. Thus, our final electron configuration is:
1s*2 2s*2 2p*6 3s*2 3p*6 4s*2 3d*10 4p*6 5s*1

Front 130

Give the abbreviated electron configuration for the following atom:

V

Back 130

The nearest 8A element that has a lower atomic number than V is Ar. The only difference between V and Ar is that there are 2 boxes in the row 4, s orbital group and 3 boxes in the row 4, d orbital group. Therefore, the abbreviated electron configuration for V is:

[Ar]4s*2 3d*3

Front 131

Give the abbreviated electron configuration for the following atom:

Sn

Back 131

The nearest 8A element that has a lower atomic number than Sn is Kr. The only difference between Sn and Kr is that there are 2 boxes in the row 5, s orbital group, 10 boxes in the row 5, d orbital group, and 2 boxes in the row 5, p orbital group. Thus, the abbreviated electron configuration for Sn is:


[Kr]5s*2 4d*10 5p*2

Front 132

Give the abbreviated electron configuration for the following atom:

In

Back 132

The nearest 8A element that has a lower atomic number than In is Kr. The only difference between In and Kr is that there are 2 boxes in the row 5, s orbital group, 10 boxes in the row 5, d orbital group, and 1 box in the row 5, p orbital group. Thus, the abbreviated electron configuration for In is:

[Kr]5s*2 4d*10 5p*1

Front 133

What is wrong with the following electron configuration?

1s*2 2s*2 2p*6 3s*2 3p*7 4s*2 3d*10

Back 133

You cannot have 7 electrons in p orbitals. Since there are three p orbitals per energy level and each can contain 2, the most you can ever have is 6.

Front 134

What is wrong with the following electron configuration?

1s*2 2s*2 2p*6 3s*2 3p*6 3d*10 4s*2 4p*5

Back 134

The order that the orbitals were filled is wrong. 3d fills up after 4s.