Name: ______Period: ______
Ch. 11 Notes Chemistry
Modern Atomic Theory Bly
11.1 Atoms and Energy
A. Rutherford’s Atom –
1. The nucleus is small and dense.
2. Protons have a +1 charge and electrons have a charge of -1
3. Neutrons have no charge
4. Rutherford didn’t know how electrons moved around the nucleus.
5. Other scientists wanted to know why electrons didn’t crash into the nucleus
B. Energy and Light
1. Electromagnetic radiation is moving energy with properties of waves.
1. The wavelength is the distance from peak to peak.
2. The frequency is the number of peaks that pass per second.
3. Different types of light have different wavelengths.
4. Examples: Microwave radiation is absorbed by food and the food heats up.
5. Fire gives off infrared radiation, which feels warm.
6. Dual nature of light - light can be though of as a wave or a particle.
7. Light is made up of a stream of photons.
8. Photons are packets of energy.
9. Short wave-length light has high-energy photons.
10. Long wave-length light has low-energy photons.
11. Which has higher energy, red or blue light? (Blue).
C. Emission of Energy by Atoms
1. Atoms can absorb energy and change to the excited state.
2. Atoms release energy by emitting photons of light. Emit = release
3. The energy of the photon = the energy given off by the atom.
4. Draw a diagram of an atom…
a. moving from the ground state to the excited state
b. moving from the excited state to the ground state
Electromagnetic radiation tutorial questions:
1. Go to Bly’s Chemistry Moodle site (http://moodle2.ais.edu.vn/) and under “Chapter 11,” click on “Electromagnetic Radiation Tutorial.”
2. Section 3: What is the wavelength of the wave shown? Draw it and label the wavelength.
3. Section 4: This shows a graph of the wavelength (black), the frequency (blue) and the speed of light. Drag the red arrow to change the wavelength. How does this change the frequency? How does it change the speed of light?
4. Section 8: You don’t have to do the math. Press the stop button and measure the wavelength in nm (nanometers 10-9).
a. Question 1: wavelength =
b. Question 2: wavelength =
c. Question 3: wavelength =
11.2 The Hydrogen Atom
A. The Energy Levels of Hydrogen (p366)
1. The ground state is the lowest energy state of an atom.
2. When atoms have excess energy, they are in the excited state.
a. Note: Atoms don’t like being in the excited state, so they always move back to the ground state immediately.
3. When hydrogen is heated, we can see four different colors of light.
a. Only specific wavelengths of light are emitted. 410nm (purple), 434nm(blue), 486nm (green) and 656nm (red).
b. Hydrogen atoms have specific energy levels.
c. When a hydrogen atom changes to a lower energy state, a photon is emitted.
d. In atoms, energy levels are quantitized. This means that only certain energy values allowed.
4. Draw an energy diagram with one ground state and two excited states. Draw an arrow to show an electron moving to the ground state.
a. Energy levels are like a staircase.
b. An atom cannot have an energy level between the stairs.
c. Draw a picture of a staircase with energy levels:
B. The Bohr Model of the Atom
1. Electrons move in orbits around the nucleus, like planets around the sun.
2. Each orbit is at a specific energy level.
3. When an electron absorbs energy, it “jumps” to a higher energy level.
4. When an electron “jumps” to a lower-energy orbit, it emits a photon.
5. Draw the Bohr model with three energy levels (orbits).
a. Label the energy levels as E1, E2, E3.
b. Label the ground state (E1)
c. Label the excited states.
d. Draw an arrow to show an electron moving from E2 to E1. Draw the photon that the electron releases.
e. Repeat “d” for an electron moving from E3 à E2 and E3 à E1.
6. The Bohr model can correctly predict the energy of electrons in hydrogen, but not for other atoms.
7. The theory is not correct because electrons do not move in circular orbits.
C. The Wave Mechanical Model of the Atom – by Schrodinger and De Brogile
Story time: On a school day, Bly is usually in his classroom. Sometimes he leaves to get a drink of water, and sometimes he goes across the street to buy coffee. However, 90% of the time he is in his classroom. If you could freeze time, you would see Bly in his classroom 90% of the time. But you wouldn’t know where he is going or what he was doing because he is frozen in time. 10% of the time, he would not be in the classroom. Imagine if you took a picture every five minutes of Bly’s classroom. This is how the wave mechanical model of the atom shows the electron.
Now imagine that you buy a glow-in-the-dark rabbit, and bring it home. Most of the time (90% of the time), it sits on the couch in front of the TV, but sometimes it goes to eat some food from its bowl in the kitchen, and sometimes it goes to visit you in your room. If you turned out the lights and took a picture every minute for an hour, what would the picture look like? Use a small circle for the rabbit. Draw it:
Your bedroom TV Couch Food
1. This picture is how the wave mechanical model looks at electrons.
2. Schrodinger thought that electrons have both wave and particle characteristics.
3. Electrons move around the nucleus in orbitals, not orbits.
4. The orbital shows the probability of finding the electron at a certain distance from the nucleus.
5. This model does not describe how an electron moves.
6. It only shows where the electron will probably be.
11.3 Atomic Orbitals
A. The Hydrogen Orbitals
Hydrogen Energy Levels
1. Principal energy level à sublevel à orbital.
2. Principal energy levels have whole numbers (n = 1, 2, 3, etc.).
3. In each principle energy level, there are sublevels (s, p, d and f)
4. In each sublevel, there are orbitals. Orbitals contain the electrons.
5. The n =1 energy level has one sublevel, 1s
6. n =2 has two sublevels: 2s and 2p
7. 2s has one orbital, a sphere. Draw it:
8. 2p has three orbitals, py, px and pz. Each p orbital looks like two balloons.
Draw 2p:
9. 3d has four balloons. Draw 3d:
10. At higher energy levels (n), the average distance of the electron from the nucleus increases.
11. Electrons in 3p are farther away than in 2p.
Draw 2p: Draw 3p:
12. Question: Hydrogen has only one electron, so why does it have more than one orbital?
Answer: Orbitals can be empty. When hydrogen moves to the excited state, the electron “jumps” to a [ higher / lower ] (circle one) orbital.
13. Types of sublevels
a. s has one orbital
b. p has three orbitals
c. d has five orbitals
d. f has seven orbitals
14. Fill in this chart:
Principal energy level / Sublevels / # of orbitals / # of electronsn = 1 / 1s / 1
n = 2 / 2s, 2p / 8
3s, 3p, 3d / 9 / 18
n = 4 / 16
1. Pauli Exclusion Principle: At most, two electrons can be in the same orbital. One has an “up” spin, and one has a “down” spin.
11.4 Electron Configurations and Atomic Properties (p377)
1. The electron configuration of hydrogen is “1s1”
2. This means there is one electron in the 1s orbital.
3. Draw the orbital diagram (box diagram) for hydrogen:
4. Rules for electron configurations:
1. Electrons occupy the lowest orbital energies first.
2. Pauli Exclusion Principle: At most, two electrons can occupy the same orbital. One has an “up” spin, and one has a “down” spin.
3. Electrons are selfish. They want their own orbital. An electron will sit with another electron in an orbital unless there are no other free orbital spaces.
Examples: Element: Li
Orbital diagram:
1s / 2s / 2p / 2p / 2p / 3sElectron configuration:______
Element: N
Orbital diagram:
1s / 2s / 2p / 2p / 2p / 3sElectron configuration:______
Element: N
Orbital diagram:
1s / 2s / 2p / 2p / 2p / 3sElectron configuration:______
Condensed electron configurations:
1. Mg can be written as [Ne]3s2
2. You can use a noble gas symbol to summarize electron configurations. The complete electron configuration of Mg is ______. The configuration of Ne is ______. Do you see that Mg just adds the 3s2 to Ne?
3. Valence electrons are the electrons in the highest principal energy level of an atom. They are involved in bonding with other atoms.
4. Nitrogen has electrons in n = 1 and 2. Level 2 is the valence level. 1s22s22p3
5. Nitrogen has 5 valence electrons.
6. Write the electron configuration of fluorine (F). Underline the valence electrons.
7. Write the electron configuration of Mg and underline the valence electrons.
8. The core electrons are the inner electrons, and are not involved in bonding.
9. Question: What elements have the same number of valence electrons as N? P, As, Sb… Where are they on the periodic table? Group 5
10. Electrons with the same number of valence electrons have similar chemical properties.
B. Electron Configurations and the Periodic Table.
Label each sublevel (s, p, d, f) on this periodic table. Don’t forget the helium (He) is in a weird place.
You can use the periodic table to find out the order of filling orbitals. You don’t need “the field” (your orbital chart) to do it.
Li
N
Close books, using periodic table, give the complete electron configurations for:
Ca ______
Al ______
Zr ______
Give the condensed electron configurations for:
I ______
Cs ______
Exceptions to the rule: Cr and Cu don’t fill orbitals in this order. The electron configuration of Cr is [Ar]4s13d5. You would expect [Ar]4s23d4, right? Wrong. Why? In this case, the energy levels 4s and 3d are very close, and it takes some extra energy to pair with the 4s electron. The electron will occupy the 3d orbital because it is at a lower energy level than pairing with the 4s electron. ***You don’t need to memorize the electron configurations of Cr and Cu, but be able to explain why they are exceptions.
C. Atomic Properties and the Periodic Table
Representative Elements - groups 1A-8A
Ionization – ionization is when an atom loses an electron.
- Remember: ______lose electrons.
- ______gain electrons.
- going down a group, atoms are more likely to lose an electron. Cs is more likely to lose an electron than Li.
Atomic size – decreases going up and right.
- decreases going up because electrons are closer at lower principal energy levels.
- decreases going right because the + charge in the nucleus is stronger.
Ionization energy - the energy required to remove an electron from an atom in the gas phase.
- increases up a group
- increases to the right
Draw a periodic table. Write arrows that show how atomic size and ionization energy change going up and to the right on the table.
Guided Notes Ch 11 13 Bly