Module 2 Notes

1. What makes an investigation credible and reliable?  (Various lessons in M1 and M2)  Be sure you understand what science is considered scientific. It is based on empirical evidence, is peer-reviewed across the globe to reduce bias and is constantly reviewed, using the scientific method, to collect new evidence over a significant period.  Credible sources that report valid data and peer reviewed are necessary.
   
   

2. List the 5 statements of Dalton’s Atomic Theory. (2.01)  

   1. All matter is composed of small, indivisible particles called atoms.

   2. All atoms of a given element are identical in mass and properties.

   3. Atoms cannot be subdivided, created or destroyed in chemical reactions. 

   4. Compounds are formed by a combination of two or more atoms in definite arrangements in the ratio of small whole numbers. 

   5. In chemical reactions, atoms are combined, separated, or rearranged. 

3. Describe Crookes Experiment. (2.01) Note:  Be sure that you can describe all scientists and their experiments from the module.  Crookes conducted experiments with cathode ray tubes.  In the tubes, one end had a negative charge while the other end had a positive charge.  When an electrical supply that is attached is turned on, a glow forms between the two electrodes.  In Crookes experiment, he placed an object in the path of the ray.  He noticed a shadow on the positive electrode.  He was able to conclude that ‘something’ had traveled from the negative end to the positive end.  He hypothesized that the glow was made up of small particles.

4. Describe how Thomson’s work expanded upon Crookes.  (2.01) Thomson conducted experiments with the cathode ray tube as well.  Thomson's work expanded upon Crookes' research on cathode rays by providing a more detailed understanding of the properties of these rays and their behavior in a vacuum tube. He was able to calculate the charge ratio for the negatively charged particles.  He determined that the negative particles were much smaller than the atom.  We now call these particles the electrons. Thomson believed that the positive and negative pieces were evenly distributed throughout the atom.  He was trying to explain how this would look when he referred to the atom as “plum pudding.”  His model of the atom became known as the “plum pudding model.”

5. Explain Rutherford’s Gold Foil experiment and relate it to Thomson’s work. (2.01) Rutherford actually worked under Thomson.  He designed a test to further support Thomson’s model of the atom.  Rutherford fired alpha particles, which are positively charged small radioactive particles at a very thin piece of gold foil.  He expected the alpha particles to pass straight through with no change of direction.  What he observed though was something different.  He actually saw that occasionally, the alpha particles were deflected.  In fact, some of them came back at him!  While this may not seem to be that big of a deal to you, imagine you had fired a BB gun at a piece of toilet paper.  You would expect the BB to go straight through, right?  Well, what if you shot the BB gun, and a few of the BB’s went off on angles, and a few came back and hit you.  Rutherford’s experiment, while created to support Thomson’s work, actually led us to create a new and enhanced model of the atom.  Rutherford found that the atom is mostly empty space; there are sections, which he called the nucleus, that are very dense and positively charged.  He discovered that protons were in this nucleus. 

6. Describe Chadwick’s work in Chemistry.  (2.01)  Chadwick worked in Rutherford’s lab in trying to figure out why the masses of nuclei were larger than the mass of their protons.  Rutherford proposed that there was a third subatomic particle, the neutron, which had approximately the same mass as a proton, but no charge.  Chadwick continued to experiment to substantiate Rutherford’s neutron idea.  He was successful in 1932. 

7. Describe the current model of the atom. (2.01) The current model of the atom states that the atom is mostly empty space.  There is a very small, dense nucleus in which protons and neutrons reside.  The electrons are in 3-dimensional orbitals surrounding the nucleus.  The proton is positively charged, the neutron has no electrical charge, and the electron is negatively charged. The proton and neutron have the same relative mass and are roughly 2000 times more massive than the electron.  When compared to the proton and neutron, the electron basically has no mass.  It is so incredibly small, it is irrelevant.  The mass of the electron in an atom is like placing a hair on an elephant! It just isn’t significant.

8. How do particles become ions? (2.01)  Ions are created when the number of electrons change.  If an atom loses electrons, it becomes more positive.  If a particle gains electrons, it becomes negative. 

9. The table shows the number of charged subatomic particles in an ion. A positively charged substance is brought near the ion. What will most likely happen? (2.01)

This ion has 4 positive particles and 2 negatively charged particles.  Right now, it has more positives than negatives.  If a positively charged particle came near it, it would repel.

10. What are groups? (2.02)  Groups are vertical columns on the periodic table.  They are also known as a family.  Elements within the same group tend to have the same properties. 

11. Describe physical properties of metals and non-metals and their locations on the periodic table. (2.02)

Metals are groups 1-12, metalloids are a bit more challenging as it encompasses parts of groups 3-12.  Non-metals are on the right side.   Solids are on the left and a little bit on the right, gasses are on the far right side.  Metalloids have properties of both metals and nonmetals.  Be able to describe each group on the periodic table (see the periodic table interactive at the  bottom of lesson 2.06, page 3)

12. Describe physical properties of metals and nonmetals and explain how they exist in nature (2.02)

13. Which elements are radioactive? (2.02)  Elements on the bottom of the periodic table

14. What do we know about an atom if we know the atomic number? (2.02)  The atomic number tells us the identity of the atom as well as the number of protons.  In an electrically neutral atom, the number of protons equals the number of electrons. 

15. What does the atomic mass tell us about an atom?  (2.02)  The atomic mass tells us the total number of protons and neutrons in the nucleus of an atom.  We can write this mathematically:  mass = protons + neutrons 

16. How many protons, neutrons and electrons are in an atom of Carbon-12? Explain how you determined them. (2.02)  Carbon has an atomic number of 6; atomic mass of approximately 12.  The atomic number tells us that carbon has 6 protons and 6 electrons.  The mass is the total number of protons and neutrons.  Therefore, we can use this simple algebra formula:   12 = p + n.  We know that carbon has 6 protons:  12 = 6 + n.  Solving for ‘n’, we find there are 6 neutrons as well.

17.  How can an atom be electrically neutral? (2.02) To be electrically neutral, the number of protons (positive) must equal the number of electrons (negative).

18. Describe the quantum model.  (2.03) The quantum model of the atom introduces the concept of the orbital - they are complex, 3-dimensional shapes (sometimes called electron clouds), in which there is likely to be an electron. So, this model is based on probability rather than certainty

19. The diagram shows four different locations in an atom.Which locations are likely to have subatomic particles that are negatively charged? (2.03)

Regardless of the picture, the nucleus contains protons and neutrons.  The electrons are always outside the nucleus.  In this particular diagram, we would expect to find negatively charged electrons in areas 2 and 4 only. 

20. Argon (Ar) has an atomic number of 18. What is the electron configuration of argon? (2.03)  Using the embedded video on page 4 as well as figures from the S1 formula sheet, the electron configuration is:   1s² 2s² 2p⁶ 3s² 3p⁶

21. How many total electrons are in each sublevel (s, p, d, f) (2.03)? 

22. What do electromagnetic waves of light consist of? (2.04/2.05)  Waves of light (electromagnetic radiation) consist of quanta of energy (photon). We call that a photon of energy. One photon is a quantized amount of energy. In other words, a photon has a fixed amount of energy.  

23. Draw and label all parts of an electromagnetic wave (2.04).  

24. Compared to X-Rays, an electromagnetic wave that has a shorter wavelength will also have what type of frequency and speed (higher/lower for each) (2.04)  Speed and frequency have a direct relationship.  If the frequency increases, the speed increases.  Thus if one decreases the other will decrease.  Wavelength has an inverse relationship to them.  If wavelength increases, the speed & frequency will both decrease.  If the wavelength decreases, the speed & frequency will both increase.

To answer this question, if an EM wave has a shorter wavelength, it will have a higher speed and frequency.

25. What type of electromagnetic radiation has a higher frequency than microwave radiation? (2.04)  Here is a diagram of the EM spectrum in order from left (high frequency, high energy, short wavelength) to right (low frequency, low energy, long wavelength)  All the EM waves that are to the left of microwaves wil have a higher frequency: infrared, visible light, ultraviolet light, X-rays, Gamma rays.

26. How are energy, wavelength and frequency related? (2.04/2.05) Energy and frequency are directly related.  As one increases, so does the other.  Wavelength is inversely related to energy and frequency.  If the energy and frequency increase, the wavelength will decrease.  Likewise, if the energy and frequency decrease, the wavelength will increase.

27. Contrast ground state electrons and excited state electrons.  (2.05) Ground state electrons are those that are in their normal, or natural state.  Refer to question 20. That is technically the “ground state” electron configuration for argon.  The excited state is when an element absorbs energy, thus causing their outermost electrons to temporarily transition to higher energy levels.  

28. Explain what happens when an atom such as lithium is placed directly into a flame. (2.05)  When we place an atom into a flame, the valence electrons will gain energy.  This will cause them to jump into higher energy orbitals.  This new state is referred to as the excited state.  The excited state is pretty unstable, so the electrons will eventually fall back down into lower energy levels.  As the atoms fall down to lower energy levels, they will release photons of energy that are equal to the difference between the two orbitals.  We see this as visible light. The color of light will change depending on the amount of energy released.  

29. What is a photon? (2.05) A photon is a very basic particle which is pure energy, a quantum of energy, specifically. A quantum of energy is a specific packet of energy. 

30. When is a photon (energy) absorbed? (2.05) A photon (of energy) is absorbed when an electron(s) transitions into a higher energy level (higher energy)

31. Describe the process of atoms gaining energy and releasing energy. (2.05) Note:  Be sure you understand the photon's role.   When elements are heated or energized, their electrons absorb energy and transition to a higher energy level. This is their excited state. Eventually, electrons fall back to a lower orbit or their ground state; but when they do, they release energy in the form of light. This light appears in different colors, depending on the frequency and wavelength of the light released. Overall, photons, with a fixed amount of energy is absorbed when we go from lower energy levels to higher (n=2 to n=3 for example)  And, these photons are released when we go from higher energy levels to lower levels (n=3 to n=2 for example)

Loosely-based example: Think of a person who does not normally consume caffeine.  One day, they drink a pot of coffee.  They are probably going to be “excited” and run around, right? Over time, the energy they have from the caffeine eventually is released and they go back to their normal state (ground state).

32. An electron moved from shell n = 3 to shell n = 4. What most likely happened during the transition? (2.05) We know from the simulation on page 3 of lesson 2.03, that n represents the energy level. Energy level 1, or ‘n=1’ is closest to the nucleus and has the least amount of energy.  As we increase the energy level, the energy increases.  In order for an electron to go from n = 3 to n = 4, it must absorb photons of energy and jumps up (becomes excited) to the next energy level.  

Essay format questions:

33. Describe Hund’s Rule.  Create the orbital notation for the element Argon. (2.03)

Hund’s rule says two things.  Be able to describe in your own words for the DBA and exams.  Hund’s rule states that every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin

We know from questions # 20, the electron configuration of Argon is:  1s² 2s² 2p⁶ 3s² 3p⁶

To put this in “orbital notation” we will include information about each electron and how it spins. We can use up and down arrows to represent the spin.  We know that 2 electrons maximum can occupy one orbital.  In that orbital, the electrons will spin in opposite directions. It does not matter if it is:  up down   or    down up.  Just as long as they are opposite in one orbital.

 ↑↓ ↑↓ ↑↓   ↑↓   ↑↓   ↑↓ ↑↓   ↑↓   ↑↓

1s²       2s²       2p⁶             3s²        3p⁶

We could annotate this as:  

1s: up arrow & down arrow;  

2s: up arrow & down arrow;  

2p: up arrow & down arrow,   up arrow & down arrow,    up arrow & down arrow; 

3s: up arrow & down arrow;  

3p: up arrow & down arrow,   up arrow & down arrow,    up arrow & down arrow; 

34. Explain if the following is a reasonable ground state electron configuration: 1s² 2s³ 2p⁶  (2.03)  The order in which electrons fill up energy levels is:  1s,  2s,  2p,  3s,  3p,  4s,  3d.  

Of each energy level (1, 2, 3, 4), there are a different number of sublevels.

n = 1 has only only sublevel (s)

n = 2 has two sublevels (s, p)

n = 3 has only three sublevel (s, p, d)

n = 4 has four sublevels (s, p, d, f)             (Do you see the pattern?)

For EACH sublevel, there are different number of orbitals:

s = 1 orbital

p = 3 orbitals

d = 5 orbitals

f = 7 orbitals      (Do you see the pattern?)

For EACH orbital, there is always a maximum number of electrons.

So, looking at the sublevels,

s = 1 orbital……. Thus a maximum of 2 electrons

p = 3 orbitals……. Thus a maximum of 6 electrons

d = 5 orbitals……. Thus a maximum of 10 electrons

f = 7 orbitals……. Thus a maximum of 14 electrons      (Do you see the pattern?)

Let’s go back to the question:   1s² 2s³ 2p⁶  

 First, look at energy levels, do they make sense?  It starts at 1 and goes into 2. Good.  Let’s look at energy level and sublevel:  It goes 1s to 2s to 2p.  Check!  That is good.  Now, let’s look at the electrons (superscript) to ensure that is correct:  1s² - good, then 2s³ - NOT GOOD.  Do you see why?

The “s” sublevel always has only ONE orbital.  EACH orbital (regardless of energy or sublevel) has a maximum of 2 electrons. So, for this configuration to be correct, we would have to change the 2s³  to 2s².  The rest is great!

35. Describe principal quantum number.  (2.03) The principal quantum number is the energy level! It ranges from 1 - 7.   It is what tells us how close the electrons are from the nucleus and their energy level. We use the letter n to represent it.

n = 1 is closest to the nucleus; lowest energy

n = 2 is a little further aways from the nucleus, and has a bit more energy than n = 1. 

Honors Extension (MC only on exam):

36. Generally speaking, which group would have the highest electronegativity value in a period? (2.06, pg 4) Using the table from the S1 Formula Sheet as well as clicking on the diagram of pg 4 of the lesson, we can see that electronegativity increases left to right.  We generally do not assign values to noble gasses as they rarely form chemical bonds.  Due to this, we would expect the halogens to have the highest EN values. In fact, fluorine has the highest EN value! It really wants an electron.

37. Where are the most reactive metals located? (2.06, pg 5) Metals are generally found on the left side of the periodic table.  Of those, as you go down a group/column, the chemical reactivity increases. The most reactive metals are the bottom left on the periodic table.

38. Which elements have the highest electron affinity? (2.06, pg 5) Electron affinity is a trend that is a bit more complicated and has more exceptions.  It is the change of energy when an electron is added to an atom. In science, when we release energy, the value will have a negative value (we will see this again later). When we absorb energy, the value will be positive. Due to this, electron affinity can be a bit confusing.  When thinking of these values, think of the number line. Those closest to 0 gain/release little energy.  Electron affinity is how much energy is gained/released. Look at the number.  Then, the sign (+ or -) indicates if it absorbed or released energy. 

• Left to Right:  Generally speaking, the electron affinity values will become more negative. This means that the elements on the right side tend to release a lot of energy (Cl has the highest number: 349. It is negative value as it releases energy, -349)

• Down a Group:  Generally speaking, the electron affinity values will become more positive.  The electrons are being added to higher energy levels, which are further from the nucleus. Loosely bound electrons in higher energy levels do not release as much energy.  Elements at the bottom of the periodic table will release less energy.  Iodine is -295 (towards the bottom) whereas Cl is -349 (towards the top).  Comparing these two, Cl has the largest number: 349 compared to 295.  That is a greater change of energy, greater electron affinity.  Chlorine’s affinity is -349, which means energy is released.  Ionine’s value is -295, which means it releases energy, but not as much as chlorine. 

The elements with the highest affinity would be those toward the top-right of the periodic table. Keep in mind, the electron affinity trend has a lot of exceptions.  

39. How do the radii of Ca and Ca²⁺ compare? (2.06) In order to determine this, we need to figure out if we are losing or gaining electrons.  Ca has 20 electrons, the outermost two electrons are in the 4s sublevel.  When it forms the cation, it will lose 2 electrons.  Those two electrons in the outermost 4s energy level.  The ion, Ca²⁺ will become smaller as the outer shell, 4s, is now empty.  

• Ca:    1s² 2s² 2p⁶ 3s² 3p⁶4s²

• Ca²⁺:  1s² 2s² 2p⁶ 3s² 3p⁶

The overall trend is that if we lose electrons, the ion tends to be smaller than its respective atom. If we gain electrons, the ion tends to be larger than its respective atom.

40. What is the general trend for electronegativity across a period and down a group, and explain each of these trends in terms of attraction, repulsion, and effective nuclear charge. (2.06) Note: Be able to describe and give the trend for all periodic trends (Ionization energy, electronegativity, atomic radius, ionic radius, electron affinity).  Note: Be able to describe and give the trend for all periodic trends (Ionization energy, electronegativity, atomic radius, ionic radius, electron affinity). 

 As you go left to right across the periodic table the effective nuclear charge increases.  This means that the valence electrons are being held more tightly to the nucleus.  Therefore, as you go left to right across the table, electronegativity increases. The greater an atom attracts its own valence electrons, the greater it is able to attract another atom’s electrons in a chemical bond. This also causes the radius to decrease.

As you go down a group, the valence electrons are further from the nucleus, in higher energy levels.  The electronegativity will decrease as you do down a group due to the increased distance.  This also causes the atomic radius to increase.