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Published byRobyn Strudwick Modified over 3 years ago

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**Chemistry Warm Up Some Dimensional Analysis Review.**

PLEASE SHOW YOUR WORK USING CONVERSION FACTORS AND DIMENSIONAL ANALYSIS If 6.02 x 1023 atoms of carbon have a mass of 12.0 grams, what the mass of x 1023 atoms of carbon atoms. Hint: set up the equality that you know. Make two conversion factors and use one to solve the problem. Check your work using dimensional analysis. 2. How many atoms are there in sample of carbon that weighs 30.0grams? 3. How many atoms are there in a sample that weighs 3.60 x 102 grams?

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**Chemistry Warm Up: Periodic Table Scavenger Hunt**

The periodic table is arranged by atomic number, not by atomic mass. Find a sequence of three elements that are arranged by atomic number but not by atomic mass. 2. Find three elements whose symbols don’t seem to have anything to do with their names. Write the name and the symbol for each. 3. There are two rows at the bottom of the periodic table. Use the atomic number to figure out where they fit in to the periodic table. 4. What would the periodic table look like if those two rows were inserted in order of their atomic number? Make a sketch.

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**Chapter5.1 Models of the Atom**

California State Science Standards Chemistry 1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept: g.* Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table. i.* Students know the experimental basis for the development of the quantum theory of atomic structure and the historical importance of the Bohr model of the atom.

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**Chapter5.1 Models of the Atom**

Dalton- Indivisible Atom J.J.Thomson discovers subatomic particle “Plum pudding,” model

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**Development of Atomic Models Rutherford’s Model**

Dense central Nucleus Electrons orbit like planets Atom mostly empty space Does not explain chemical behavior of atoms

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**The Bohr Model Electrons orbit the nucleus Specific circular orbits**

Quantum = energy to move from one level to another

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**The Bohr Model Energy level like rungs of the ladder**

The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder A quantum of energy is the amount of energy required to move an electron from one energy level to another

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The Bohr Model Energy level of an electron analogous to the rungs of a ladder But, the rungs on this ladder are not evenly spaced!

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**Quantum Mechanical Model**

Energy quantized; comes in chunks. A quantum is the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy. 1926 Erwin Schrodinger equation described the energy and position of electrons in an atom

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**Quantum Mechanical Model**

•Things that are very small behave differently from things big enough to see. •The quantum mechanical model is a mathematical solution •It is not like anything you can see.

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**Quantum Mechanical Model**

•Has energy levels for electrons. •Orbits are not circular. •It can only tell us the probability of finding an electron a certain distance from the nucleus.

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Atomic Orbitals •Energy levels (n=1, n=2…) •Energy sublevels = different shapes •The first energy level has one sublevel: 1s orbital -spherical

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Atomic Orbitals •The second energy level has two sublevels, 2s and 2p There are 3 p-orbitals

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**•The third energy level has three sublevels, 3s**

Atomic Orbitals •The third energy level has three sublevels, 3s 3p And 5 3d orbitals py

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Atomic Orbitals •The forth energy level has four sublevels, 4s 4p 4d orbitals And seven 4f orbitals

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**Atomic Orbitals The principal quantum number (energy level) equals the number sublevels**

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**5.2 Electron Arrangement in Atoms**

Electron Configuration Electrons and nucleus interact to produce most stable arrangement= Lowest energy configuration

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**Aufbau Principle Electrons fill the lowest energy orbitals first**

3 rules: Aufbau Principle Electrons fill the lowest energy orbitals first Hydrogen has 1 electron 1s1

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3 rules: Pauli Exclusion Principal- two electrons per orbital (one spin up, one spin down) Boron has 5 electrons 1s2 2s2 2p1

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3 rules: Hund’s rule- In orbitals with equal energy levels, arrange spin to maximize electrons with the same spin 1s22s22p3 Nitrogen has 7 electrons Hund’s Rule: Separate the three 2p elecrons into the three available 2p orbitals to maximize the electrons with the same spin.

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**Conceptual Problem p135 1s2 2s2 2p6 3s2 3p3**

Electron Configuration for Phosphorus (atomic # = 15) 1s2 2s2 2p6 3s2 3p3

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**Practice Problem 8a p135 1s2 2s2 2p2**

Electron Configuration for Carbon (atomic number = 6) 1s2 2s2 2p2

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**Practice Problem 8b p135 1s2 2s2 2p6 3s2 3p6**

Electron Configuration for Argon (atomic # = 18) 1s2 2s2 2p6 3s2 3p6

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**Practice Problem 8c p135 1s2 2s2 2p6 3s2 3p6 3d8 4s2**

Electron Configuration for Nickel (atomic # = 28) 1s2 2s2 2p6 3s2 3p6 3d8 4s2

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**Practice Problem 9a p135 1s2 2s2 2p1 1**

Electron Configuration for Boron (atomic # = 5) 1s2 2s2 2p1 How many unpaired electrons? 1

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**Practice Problem 8c p135 1s2 2s2 2p6 3s2 3p2 2**

Electron Configuration for Silicon (atomic # = 14) 1s2 2s2 2p6 3s2 3p2 How many unpaired electrons? 2

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**Exceptions to the Aubau Rule**

Copper atomic number=29 1s2 2s2 2p6 3s2 3p6 3d9 4s2 This is the expected electron configuration

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**Exceptions to the Aubau Rule**

Copper atomic number=29 1s2 2s2 2p6 3s2 3p6 3d10 4s1 Half-filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel This is the actual electron configuration.

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**Exceptions to the Aubau Rule**

Chromium atomic number=24 1s2 2s2 2p6 3s2 3p6 3d4 4s2 This is the expected electron configuration

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**Exceptions to the Aubau Rule**

Chromium atomic number=24 1s2 2s2 2p6 3s2 3p6 3d5 4s1 Half-filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel This is the actual electron configuration.

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**5.3 Physics and the Quantum Mechanical Model**

Or, “How do they get all those colors of neon lights?”

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Goals •Describe the relationship between wavelength and frequency of light •Identify the source of atomic emission spectra •Explain how frequency of emitted light are related to changes in electron energies •Distinguish between quantum mechanics and classical mechanics

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**Quick review of wave terminology**

Amplitude = height of wave Wavelength = distance between crests Frequency = number of crests to pass a point per unit of time

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**Light waves Amplitude = height of wave**

Wavelength = distance between crests Frequency = number of crests to pass a point per unit of time For light, the product of frequency and wavelength = speed of light, c Frequency • Wavelength = 3.00 x 108 So, as the frequency of light increases, the wavelength decreases

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**Electromagnetic Spectrum**

Visible light is only part of the electromagnetic spectrum:

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Wavelength of Light p140 Sample Problem: What is the wavelength of yellow light from a sodium lamp if the frequency is 5.10 x 1014 Hz (Hz = s-1) Wavelength • frequency = 3.00x108m/s Wavelength = 3.00x10-8 m/s / frequency Wavelength = 3.00x108m/s / 5.10x1014 s-1 Wavelenght = 5.88 x 10-7 m

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Wavelength of Light p140 #14:What is the wavelength of radiation if the frequency is 1.50x1013 Hz (Hz = s-1)? Is this longer or shorter than the wavelenght of red light? Wavelength • frequency = 3.00x108m/s Wavelength = 3.00x108 m/s / frequency Wavelength = 3.00x108m/s / 1.50x1013 s-1 Wavelength = 2.00 x 10-5 m Longer than red lightwhich if between 10-6 and 10-7 m

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Wavelength of Light p140 #15: What is the frequency of radiation if the wavelength is 5.00x10-8 Hz (Hz = s-1) In what range of the electromagnetic specrum is this? Wavelength • frequency = 3.00x108m/s frequency = 3.00x108 m/s / wavelength frequency = 3.00x108m/s / 5.00x10-8m Frequency = 6.00 x 1015 s-1 ultraviolet

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**Atomic Spectra When atoms absorb energy,**

Electrons move to higher energy levels. When electrons return to the lower energy level, they emit light Each energy level produces a certain frequency of light resulting in an emission spectrum

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**Atomic Spectra Emission spectra are like a fingerprint of the element**

We know what stars are made of by comparing their emission spectra to that of elements we find on earth

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**Explanation of Atomic Spectra**

Emission spectra like a fingerprint of the element We know what stars are made of by comparing their emission spectra to that of elements we find on earth

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