# Electrons in Atoms. What do you know about a wave?

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Electrons in Atoms

What do you know about a wave?

Electromagnetic spectrum Visible light: all that can be seen by the human eye *** Blue flames have a shorter wavelength than yellow flames ***Blue flames are higher in energy and hotter! The rest can’t be seen, but is still considered light

Graphing Direct and Inverse Proportions Direct Proportion: Inverse Proportion: Y X Y X

Frequency and Wavelength

Relationships Between Energy, Wavelength and Frequency of Waves: Energy and Frequency: Wavelength and Frequency: Energy and Wavelength: E ν λ ν λ E Direct Inverse

Problems: 1. If the wavelength of a wave doubles, and its energy and speed don’t change, what will happen to its frequency? 2λ → ½ υ The frequency (υ) will be cut in half. 2. If the frequency of a wave triples, what happens to the energy needed for the wave? 3υ → 3E You will need THREE times the energy 3. A student modifies the wavelength of a wave by changing the energy. If the energy is cut in half, what must have happened to the wavelength? ½ E → 2λ

Light has two components: Wavelength: in (m)Frequency: in (1/s) λ = ν = Energy: in (J) E = hυ or c = speed of light (3 x 10 8 m/s) h = Planck’s constant (6.6 x 10 -34 Js) 1 x 10 9 nm = 1 m Or 1nm = 1.0 × 10 -9 m λ ↑ ν ↓E↓ λ ↓ ν ↑E↑

Energy WS problems 1. Find the energy of a wave if its frequency is 2.2 x 10 16 1/sec E = hυ E = (6.6 x 10 -34 Js) x (2.2 x 10 16 1/sec) E = 1.45 x 10 -17 J

6. Find the wavelength of light if its energy is 2.8 x 10 -19 J. What color light would you see? Use E = hυ to find υ: 2.8 x 10 -19 J = (6.6 x 10 -34 Js) x υυ = 4.24 x 10 14 1/s Then solve for wavelength: = 7.07 x 10 -7 m = 707 nm Red color = =

#6 (using one equation) E= =

Review: Neil Bohr Confined electrons to energy levels Quantum leaps: Energy (light) being released when e- jump from excited state to ground state.

Atomic emission spectrum:  Each element has a unique line-emission spectrum Emission spectrum of Hydrogen and Iron:

Drawing Bohr’s atom: Energy levels# of e- 12 28 318 432

Present Day Model Electrons are located in an orbital Orbitals- regions around a nucleus that correspond to specific energy levels Orbitals are also called electron clouds Don’t know the exact position of the electron- creates fuzzy image Electron cloud

Quantum Numbers Quantum Mechanical Model - present day model of the atom Electrons do NOT orbit nucleus in circular pattern How can we keep track of the electrons? Four Quantum numbers - defines region in which electrons can be found

The fixed energies an electron can have are called energy levels. electron’s energy & distance from the nucleus ↑ {analogy is floors in apartment building} 1. Energy Level + n= 1, 1 st energy level, 1 st floor n= 2, 2 nd energy level, 2 nd floor Think of the atom as an apartment building with each floor representing an energy level

2. Sublevel Represent the sublevels of the main energy level {analogy is apartments on a floor} Type or shape of orbital: s, p, d, f

S - orbital Size:1S < 2S < 3S

p - orbital

d - orbital

3. Orbital subset of the sublevels probability map of finding an electron. Indicates numbers and orientations of orbitals around nucleus {analogy is rooms in an apartment} # of orbitals includes : one s orbital three p orbitals five d orbitals seven f orbitals

4. Spin indicates the orientation of an electron’s magnetic field relative to an outside magnetic field A single orbital can hold a maximum of 2 electrons, which must have opposite spins {analogy is roommates in a room}

S: 1 orbital Max e- = 2 P:3 diff. orbitals Max e- = 6 D:5 diff. orbitals Max e- = 10 f: 7 diff. orbitals Max e- = 14

Aufbau principle states that electrons fill orbitals that have the lowest energy first (e.g. 1s before 2s). Overlapping Orbitals

Electron Configuration arrangement of the electrons around the nucleus Based on the quantum model of the atom 1 H = 1s 1 # of electrons Sublevel 1 st energy level n = 1 1H1H Orbital Notation

Hund’s Rule states that orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin.

Pauli exclusion principle Maximum 2 electrons can occupy a single orbital and they must have opposite spins (  or  ). No 2 atoms in the same atom can have the same 4 quantum numbers Electrons in orbitals can be represented by arrows in boxes

Give the full electron configuration for... Li Na K 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The electrons in the outer most shell are called the valence electrons. Electrons in the inner shells are core electrons.

Electron Configuration Example: sulfur has sixteen electrons. Its electron configuration is written as 1s 2 2s 2 2p 6 3s 2 3p 4 Shorthand Notation: It can also be written as follows by using the previous noble gas:[Ne]3s 2 3p 4

Write the electron configuration for an atom whose atomic number is 20. atomic number = number of protons = number of electrons = 20 According to the aufbau principle, the order of orbital filling is 1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Or refer to the periodic table 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Abbreviated as follows: [Ar]4s 2

Try These 1. Try writing out the electron configuration for Potassium. 2. Write the electron Configuration for an atom with an atomic number 20. 3. Write an electron configuration for an atom of an element whose atomic number is 8. 4. Write an electron configuration for an atom of an element whose atomic number is 53.

Answers 1. K = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 2. Ca = [Ar]4s 2 3. O = 1s 2 2s 2 2p 4 4. I = [Kr] 5s 2 4d 10 5p 5

What is the correct electron configuration of a sulfur atom? A. 1s 2 2s 2 2p 4 3s 2 3p 6 B. 1s 2 2s 2 2p 6 3s 2 3p 3 C. 1s 2 2s 2 2p 6 3s 2 3p 4 D. 1s 2 2s 2 2p 6 3s 6 3p 2

Drawing Molecules using Lewis Electron-Dot Structures Lewis structure: shows how the valence electrons are arranged among the atoms in the molecule Valence electrons are represented by dots Nuclei and electrons of the inner energy level (if any) are represented by the symbol of the element

Exceptional Electron Configurations Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 The correct electron configurations are as follows: Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 These arrangements give chromium a half-filled d sublevel and copper a filled d sublevel. Some actual electron configurations differ from those assigned using the aufbau principle because although half-filled sublevels are not as stable as filled sublevels, they are more stable than other configurations.

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How are the quantum mechanical model and the Bohr model alike? How are they different? Like the Bohr model, the quantum mechanical model restricts the energy of electrons to certain values. Unlike the Bohr model, the quantum mechanical model does not specify an exact path the electron takes around the nucleus.

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Calculate the maximum number of electrons in the 5 th principal energy level (n = 5). The maximum number of electrons that can occupy a principal energy level is given by the formula 2n 2. If n = 5, 2n 2 = 50.

How does Electron Configuration relate to the Periodic Table?

2 6 10 14

Explain why the correct electron configuration of oxygen is 1s 2 2s 2 2p 4 and not 1s 2 2s 2 2p 3 3s 1. The 2p orbitals are lower in energy than the 3s orbital, so they will be completely filled before any electrons will be found in the 3s orbital.

Frequency and Wavlength  Light has two components: wavelength (λ – read lambda ) and frequency (ν – read nu )

Energy Levels in Atoms The rungs on this ladder are somewhat like the energy levels in Bohr’s model of the atom. A person on a ladder cannot stand between the rungs. Similarly, the electrons in an atom cannot exist between energy levels.

Energy Levels in Atoms The rungs on this ladder are somewhat like the energy levels in Bohr’s model of the atom. The energy levels in atoms are unequally spaced, like the rungs in this unusual ladder. The higher energy levels are closer together.

A prism separates light into the colors it contains. White light produces a rainbow of colors. Light and Atomic Emission Spectra Light bulb SlitPrism Screen

Light from a helium lamp produces discrete lines. Light and Atomic Emission Spectra SlitPrism Screen Helium lamp Spectrum: wavelengths of visible light that are separated when a beam of light passes through a prism; range of wavelengths of electromagnetic radiation

Hydrogen’s Line-Emission Spectrum Spectrum: wavelengths of visible light that are separated when a beam of light passes through a prism; range of wavelengths of electromagnetic radiation

An Explanation of Atomic Spectra The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels.

s 1 orientaion Max e- = 2 p 3 orientaion 3 diff. orbitals Max e- = 6 s orbitals are spherical p orbitals are dumbbell-shaped

f 7 orientaion 7 diff. orbitals Max e- = 14 D 5 orientaion 5 diff. orbitals Max e- = 10

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The Heisenberg Uncertainty Principle The Heisenberg uncertainty principle states that it is impossible to know both the velocity and the position of a particle at the same time. This limitation is critical when dealing with small particles such as electrons. But it does not matter for ordinary-sized objects such as cars or airplanes. Quantum Mechanics

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. To locate an electron, you might strike it with a photon. The electron has such a small mass that striking it with a photon affects its motion in a way that cannot be predicted accurately. The very act of measuring the position of the electron changes its velocity, making its velocity uncertain. Quantum Mechanics Before collision: A photon strikes an electron during an attempt to observe the electron’s position. After collision: The impact changes the electron’s velocity, making it uncertain.

Amplitude: the height of a wave’s crest Wavelength (,lambda): distance between adjacent crests of a wave Frequency (, nu): #of wave cycles to pass a given point per unit of time.

Electromagnetic spectrum Speed of light: 3  10 8 m/s in a vacuum. E = h ν

Frequency and Wavelength  Light has two components: wavelength (λ – read lambda ) and frequency (ν – read nu )  Frequency and wavelength are inversely proportional ( ν = 1/ λ ) E = h ν λ ↑ ν ↓E↓ λ ↓ ν ↑E ↑

 Each element has a unique line-emission spectrum Emission spectrum of Hydrogen and Iron: Atomic emission spectrum: the pattern formed when light passes through a prism or diffraction grating to separate it into the different frequencies of light it contains

1 st level = 1 orbital, shape is spherical (s) 1”s” orbital- (shape- spherical) (1s) 2 nd level = 4 orbitals 1 “s” orbital- (shape- spherical) (2s) 3 “p” orbitals- (shape- dumbbell) (2p) 3 rd level= 9 orbitals 1 “s” orbital- (shape- spherical) (3s) 3 “p” orbitals- (shape- dumbbell) (3p) 5 “d” orbitals- shape- most are four leaf clover (3d) 4 th level=16 orbital 1”s” orbital- (shape- spherical) (4s) 3 “p” orbitals- (shape- dumbbell) (4p) 5 “d”orbitals- shape- most are four leaf clover (4d) 7 “f” orbitals- shapes are complicated (4f)

Atomic Orbitals Summary of Principal Energy Levels and Sublevels Principal energy level Number of sublevels Type of sublevel Maximum number of electrons n = 111s (1 orbital)2 n = 222s (1 orbital), 2p (3 orbitals)8 n = 33 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals) 18 n = 44 4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals) 32 The numbers and types of atomic orbitals depend on the principal energy level.

Do Now Ho does E, λ, ν change as the wave goes from A to B? A. → B.

Do Now #1 Draw a picture of a wave with high frequency and low frequency. Label wavelength Circle the wave with high energy high frequencylow frequency. λ λ

The visible red light has a wavelength of about 650 nm. At sunrise and sunset, red or orange colors are present because the wavelengths associated with these colors are less efficiently scattered by the atmosphere than the shorter wavelength colors (e.g., blue and purple). A large amount of blue and violet light has been removed as a result of scattering and the longwave colors, such as red and orange, are more readily seen.