Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 1 of 6

Ch. 3 Objectives Key Obj. 1-Laws and Theories • Law of Definite Proportions: a chemical compound contains the same elements

in exactly the same ratio by mass no matter the size of the compound. • Ex1: Imagine a room full of NaCl vs. a tablespoon of NaCl. Does the amount of

NaCl change what it is made of? No • Ex2: The mass ratio of a compound (AB) is 1:3. (25% A, 75% B)

• If you have 4 g of AB, the mass of A is 1 g and the mass of B is 3 g. • If you have 100 g of AB, the mass of A is 25 g and the mass of B is 75 g.

• Law of Multiple Proportions: when two elements combine to form two or more

new compounds, the mass of each of the elements that combine in a way that their ratios are whole numbers. • Ex1: H2O vs. H2O2

• Example: a) 34g + 168 g =202 g c) 34g + 336g =370 g b) 68g + 168g = 202 g d) 68g + 336g = 404 g

• Which ones show the law of conservation of mass? A, C, D because they add up (left side=right side)

• Which ones show law of definite proportions? A and D because they combine in the same ratio (D is twice as much as A)

• John Dalton's Atomic Theory • John Dalton (1766-1844) took indirect evidence, took

other people’s ideas and came up with an atomic theory. He was the first to relate atoms (can’t see/measure) to mass (can measure)!

• Dalton's Atomic Theory (1808) 1. All matter is composed of extremely small particles

called atoms. 2. Atoms of a given element are identical in mass, size, and other properties. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements can combine in simple whole-number ratios to

form compounds. 5. In chemical reactions, atoms are separated,

combined, or rearranged. • Which points are incorrect? 2 & 3

• Dalton thought an atom was like a bowling ball. Obj. 2-The Electron-JJ Thomson

Name Symbol Charge Mass Location Electron e-1 -1 0 amu Electron Cloud

Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 2 of 6 • History

• Electrical currents passed through different gases at low pressures in cathode ray tubes (CRT).

• Things Discovered by JJ Thomson (by accident): 1) Different gases give off different colors. 2) When he placed a paddle wheel in the tube, it would spin. He

discovered the electrons have mass. 3) The rays deflected away from a negative plate. JJ discovered

electrons have a negative charge. • JJ Thomson's Chocolate Chip Cookie Model

Obj. 3-The Proton-Rutherford

Name Symbol Charge Mass Location Proton p+1 +1 1 amu nucleus

• History (1908-1909) • Ernest Rutherford performed the “Gold Foil

Experiment” • He shot + alpha particles through gold foil.

• Rutherford observed that most alpha particles went straight through and only a few deflected.

• Rutherford’s Conclusions: 1) Since most particles went straight through, the atom is mostly empty space. 2) Since a few deflected, the + alpha particles came close to small areas of positive charge, called protons.

• Sun-Planet Model Obj. 4-The Neutron

Name Symbol Charge Mass Location Neutron n0 0 1 amu nucleus

• History • James Chadwick found that adding the masses of the electrons and protons in

an atom did not equal the total mass of the atom. • Experiments proved that there had to be another heavy particle without a

charge, called the neutron.

Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 3 of 6 Obj. 5-The Modern Atom • Atomic Number: number of protons

• Tells which element it is • Whole number on the periodic table

• Mass Number: number of protons and neutrons • The number of neutrons can vary for an element. • NOT on the periodic table

• Isotopes: atoms of the same element that have the same number of protons, but a different number of neutrons. Therefore, different masses. • Two Ways to Write Isotopes:

Nuclear Symbol Hyphen Notation Format

Atomic #Mass #Symbolcharge

Name of Element-Mass #

Example

€

12H

Hydrogen-2

• Where would you have to look to find the atomic number if you are given the hyphen notation? The Periodic Table

Hyphen Notation Nuclear Symbol

n0 p+ e- Atomic # Mass #

Lead-204 82204Pb

122 82 82 82 204

Tellurium-101 52101Te2−

49 52 54 52 101

Manganese-48 2548Mn2+

23 25 24 25 48

Potassium-42 1942K

23 19 19 19 42

Obj. 6-Counting Atoms • Atomic Mass: mass of an atom based on the average of all isotope masses

• units are atomic mass unit (amu) • decimal number on the periodic table

• Mole: measure of the amount of substance • Avogadro's Number: 6.02 x 1023 or 602,000,000,000,000,000,000,000 • Avogadro's Number is the number of atoms in one mole

charge = protons - electronsNumber of Neutrons = mass # − atomic #

Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 4 of 6 Obj. 7-Wave Characteristics • Wavelength (λ): distance between

two corresponding points in a wave • Units: meter (m)

• Frequency (f): # of waves that pass a specific point in one second

• Units: Hertz (Hz), 1s

• Speed of light (c): 3 x 108 m/s • As wavelength decreases, frequency increases. This is an indirect relationship! • Electromagnetic Radiation: form of energy

that exhibits wave-like behavior as it travels through space (p. 92) • Travels only at the speed of light: 3 x 108 m/s • Only can see a small part of the whole

electromagnetic spectrum • Spectroscopy: way to identify substances

using spectra • Each element has a unique spectrum (fingerprints) • Max Planck (1900) studied the non-continuous spectrum and

discovered: 1. Energy emitted from the substances was not continuous. He

called these packets of energy photons! 2. Electrons like to be stable (lazy). Once they gain energy,

electrons want to lose it. 3. Electrons quickly release the gained energy in certain packets of energy

(photons), which we see as light, to become more stable. 4. As energy increases, the frequency increases. (Direct relationship!)

Obj. 8-Bohr’s Theory (1913) • Neils Bohr’s atomic model attempts to explain the non-continuous spectrum. • Bohr's Atomic Model (1913)

• Electrons can only be a certain distance away from the nucleus in a circular path, called an orbit. (Like a ladder, cannot be between orbits.)

• As electrons gain energy, they move further away from the nucleus. • To jump from one energy level to another, an electron has to gain or lose a

certain amount of energy (photons).

c = λ ⋅ fc = 3 x 108 m/s λ = wavelength (m) f = frequency (Hz, 1/s)

Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 5 of 6 • How did Bohr figure this out?

• Passed electricity through a tube of Hydrogen • Separated the light that was emitted from the

tube and noticed four bands of color (not continuous like a rainbow)

• What was Bohr’s explanation? 1) Since H only has one electron, each band of

color represented a different energy level jump. 2) Electrons would gain energy to achieve an

excited state, and lose the energy to drop down to a ground state.

3) The colors that are observed represent the energy that is being released by the electron when it moves to a lower energy level (orbit).

4) Since spectrum was not continuous, Bohr determined that the electron cannot exist everywhere, only in certain orbits.

5) 4 colors of light represented the 4 wavelengths, 4 frequencies, and 4 different amounts of energy.

Obj. 9-Quantum Theory (1926) • Developed by Edwin Schrodinger • He used some of Bohr’s ideas, but also described how electrons behave in their

“orbit”. His model is mathematical equations. • The Quantum Theory (1926)

1. Electrons are more like waves than a particle.

2. Electrons exit in 3-D regions called orbitals (electron clouds). Orbitals are regions where an electron will most likely be located

3. The size and shape of the orbital depend on the number of electrons. Obj. 10-Bohr’s vs. Quantum Theory

Similarities Differences 1. Electrons orbit the nucleus 2. When electrons gain energy, they move

away from the nucleus (excited state) 3. When electrons lose energy, they move

toward the nucleus (ground state) 4. A certain amount of energy is needed for

an electron to move from one energy level to another (photon)

1. Bohr’s model treats electrons as a particle, but the Quantum model treats electrons as a wave.

2. Bohr’s model uses orbits (2D), but the Quantum model uses orbitals (3D).

Osterberg-Chemistry Name_____________________________ Ch. 3-Atoms and Moles Date ______________ Hour__________ Page 6 of 6

How many electrons are in the orbital

the shape of the orbital

energy level (how far out)

�

↑↓1s

↑↓2s

↑2p

↑2p 2p

�

↑↓1s

↑↓2s

↑↓2p

↑↓2p

↑↓2p

↑↓3s

↑↓3p

↑3p

↑3p

�

↑↓1s

↑↓2s

↑↓2p

↑↓2p

↑↓2p

↑↓3s

↑↓3p

↑↓3p

↑↓3p

↑↓4s

↑3d

↑3d

↑3d

↑3d 3d

Obj. 11-Electron Configurations • Electron Configurations are a way to show how e- are arranged in atoms • Pauli Exclusion Principle: No more than two electrons can be in one orbital Orbital s p d f # of types 1 3 5 7 # of electrons 2 6 10 14 • Aufbau principle: electrons fill lowest energy level first • What do the numbers and letters tell us?

3s2

• Electron Configurations Using the Diagonal Rule • Examples: Remember # of Protons = # of electrons if neutral

• C (6 e-) 1s22s22p2 • S (16 e-) 1s22s22p63s23p4 • Cr (24 e-) 1s22s22p63s23p64s23d10 Try Mn, Sn, and Zn!

Obj.12àOrbital Filling Notation • Goes one step further and shows spin of electrons and how orbitals are filled • Hund’s principle: orbitals of same energy level are occupied by one e- before

being paired up because like charges repel O (6 e-) S (16 e-) Cr (24 e-) Obj. 12àElectron Dot Notation • Valence electrons: electrons in the outermost energy level; can only have 8 max

Li 1s22s1 N 1s22s22p3 Be 1s22s2 O 1s22s22p4 B 1s22s22p1 F 1s22s22p5 C 1s22s22p2 Ne 1s22s22p6