Chapter 5 - Electrons in Atoms

THE NUCLEAR ATOM AND UNANSWERED QUESTIONS 1. In the early 1900’s scientists were able to determine that an element’s chemical behavior is related to the arrangement of electrons in its atoms

WAVE NATURE OF LIGHT 1. Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space Examples: a) visible light (ROYGBIV) b) microwaves c) x-rays d) radio waves

WAVE NATURE OF LIGHT 2. Primary characteristics of all waves include wavelength, frequency, amplitude, and speed SEE Fig. 5-2, Page 119 3. Wavelength () is the shortest distance between equivalent points on a continuous wave is usually expressed in meters, centimeters, or nanometers

WAVE NATURE OF LIGHT 4. Frequency (  ) is the number of waves that pass a given point per second 5. The SI unit for frequency is the hertz (Hz) which equals one wave per second Example: 652 Hz = 652 waves/second = 652/s = 652s-1

WAVE NATURE OF LIGHT 6. The amplitude of a wave is the wave height from the origin to a crest, or from the origin to a trough 7. All electromagnetic waves travel at the speed of 3.00 X 108 m/s in a vacuum

WAVE NATURE OF LIGHT 8. The speed of light ( c ) is the product of its wavelength (  ) and its frequency (  ) c =   9. Although the speed of all electromagnetic waves is the same, waves may have different wavelengths and frequencies SEE Fig. 5-3, Page 119

WAVE NATURE OF LIGHT 10. White light is a combination of all the colors of the visible spectrum R O Y G B I V e r e r l n i d a l e u d o n l e e i l g o n g e e w o t SEE Fig. 5-4, Page 1`20

WAVE NATURE OF LIGHT 11. The electromagnetic spectrum encompasses all forms of electromagnetic radiation SEE Fig. 5-5, Page 120 Example Problem 5-1, Page 121 Practice Problems, Page 121

Practice Problems 1. A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? 2. What is the wavelength of X-rays having a frequency of 4.80 X 1017 Hz? 3. An FM radio station broadcasts at a frequency of 98.5 MHz. What is the wavelength of the station’s broadcast signal?

PARTICLE NATURE OF LIGHT 1. The temperature of an object is a measure of the average kinetic energy of its particles 2. As an object gets hotter it possesses a greater amount of energy, and emits different colors of light. SEE Fig. 5-6, Page 122

PARTICLE NATURE OF LIGHT 3. These different colors correspond to different frequencies and wavelengths 4. Max Planck determined that matter can gain or lose energy only in small, specific amounts called quanta

PARTICLE NATURE OF LIGHT 5. A quantum is the minimum amount of energy that can be gained or lost by an atom 6. The energy of a quantum is related to the frequency of the emitted radiation by the equation Equantum = h  where E is energy, h is Planck’s constant, and  is frequency

PARTICLE NATURE OF LIGHT 7. Planck’s constant has a value of 6.626 X 10-34 J•s 8. The photoelectric effect is a phenomenon in which photoelectrons are emitted from a metal’s surface when light of a certain frequency shines on the surface SEE Fig. 5-7, Page 123 Example: Calculator in Fig. 5-8, Page 123

PARTICLE NATURE OF LIGHT 9. While a beam of light has many wave characteristics, it can also be thought of as a stream of tiny particles, or bundles of energy called photons 10. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy

PARTICLE NATURE OF LIGHT 11. Albert Einstein calculated that a photon’s energy depends on its frequency Ephoton = h  Example Problem 5-2, Page 124 Practice Problems, Page 124

ATOMIC EMISSION SPECTRA 1. The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element 2. The atomic emission spectrum is characteristic of the element and is used to identify the element. Section 5.1 Assessment, Page 126

BOHR MODEL OF THE ATOM 1. Niels Bohr proposed a quantum model for atoms which included: a) predicted the frequencies of the lines in hydrogen’s atomic emissions spectrum b) proposed that the hydrogen atom has only certain allowable energy states

BOHR MODEL OF THE ATOM c) related the hydrogen atom’s energy states to the motion of the electron within the atom d) the single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits

e) the smaller an electron’s orbit, the lower BOHR MODEL OF THE ATOM e) the smaller an electron’s orbit, the lower the atom’s energy state, or energy level f) assigned a quantum number, (n), to each orbit and calculated the orbits radius SEE Table 5-1, Page 127

BOHR MODEL OF THE ATOM 2. The lowest allowable energy state of an atom is called its ground state 3. In the ground state, an atom does not radiate energy 4. When energy is added to an atom by an outside force, electrons can move up to a higher-energy orbit (called the excited state)

BOHR MODEL OF THE ATOM 5. Electrons can also drop to a lower-energy level and emit energy in the form of photons 6. To calculate the energy (E) of a photon we can use the following equation: ΔE = Ehigher-energy orbit - Elower-energy orbit = Ephoton = h 

THE QUANTUM MECHANICAL MODEL OF THE ATOM 1. Louis de Broglie proposed that particles of matter can behave like waves 2. de Broglie derived an equation for the wavelength (  ) of a particle of mass (m) moving at a velocity ( v )  = h mv

3. The de Broglie equation predicts that all moving particles have wave characteristics Example: 910 kg traveling at 25 m/s m = 910 kg v = 25 m/s h = 6.626 X 10-34 J  = ?  = h = 6.626 X 10-34 J = 2.9 X 10-38 m mv (910kg)(25m/s)

1. Werner Heisenberg concluded that it is THE HEISENBERG UNCERTAINTY PRINCIPLE 1. Werner Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object 2. Consider the energy of a photon: A high-energy photon of electromagnetic radiation has about the same energy as an electron. The interaction between the two particles changes both the wavelength of the photon and the position and velocity of the electron SEE Fig. 5-13, Page 132

3. The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time. Example: The effect of a photon emitted by a flashlight on a helium balloon is so small that it is virtually impossible to measure 4. Erwin Schrodinger derived an equation that treated the hydrogen atom’s electron as a wave

6. The region around the nucleus of an atom THE HEISENBERG UNCERTAINTY PRINCIPLE 5. The atomic model proposed by Schrodinger is called the quantum mechanical model of the atom 6. The region around the nucleus of an atom where electrons are located is called an atomic orbital SEE Fig. 5-13, Page 132

HYDROGEN’S ATOMIC ORBITALS 1. Principle Quantum Numbers ( n ) indicate the relative sizes and energies of atomic orbitals 2. As n increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases

HYDROGEN’S ATOMIC ORBITALS 3. n specifies the atom’s major energy levels, called principle energy levels 4. Up to seven energy levels have been detected 5. Principle energy levels contain energy sublevels

HYDROGEN’S ATOMIC ORBITALS 6. The quantum number assigned to an energy level also indicates the number of sublevels contained within the energy level Examples: First energy level (n = 1) has one sublevel Second energy level (n = 2) has two sublevels Third energy level (n = 3) has three sublevels

n = 3 has a 3s, 3p, and 3d sublevel HYDROGEN’S ATOMIC ORBITALS 7. Sublevels are labeled s, p, d, or f according to the shape of the atoms orbitals Examples: n = 1 has a 1s sublevel n = 2 has a 2s and 2p sublevel n = 3 has a 3s, 3p, and 3d sublevel n = 4 has a 4s, 4p, 4d, and 4f sublevel SEE Table 5-2, Page 134 Section 5.2 Assessment, Page 134

1. The arrangement of electrons in an atom is GROUND STATE ELECTRON CONFIGURATION 1. The arrangement of electrons in an atom is called the atom’s electron configuration 2. Electrons in an atom tend to assume the arrangement that gives the atom the lowest possible energy 3. The most stable, lowest-energy arrangement of the electrons in an atom of each element is called the element’s ground-state electron configuration

4. The aufbau principle states that each GROUND STATE ELECTRON CONFIGURATION 4. The aufbau principle states that each electron occupies the lowest energy orbital available 5. The sequence of atomic orbitals from lowest to highest energy (aufbau diagram) is shown in Figure 5-17, Page 135

Aufbau Diagram 7p

6. Key features of the aufbau principle include: GROUND STATE ELECTRON CONFIGURATION 6. Key features of the aufbau principle include: a) all orbitals related to an energy sublevel are of equal energy b) in a multi-electron atom, energy sublevels within a principal energy level have different energies

GROUND STATE ELECTRON CONFIGURATION c) in order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, and f d) orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level

The Pauli Exclusion Principle 7. The Pauli Exclusion principle states that a maximum of two electrons may occupy a single orbital, and the electrons must have opposite spins Each type of subshell contains a different number of orbitals.  Each orbital can hold at most 2 electrons.  www.wwnorton.com/chemistry/tutorials/ch3.html

The Pauli Exclusion Principle The following table shows how many electrons each type of subshell can hold. Subshell Type # of Orbitals Maximum # of Electrons s 1 2 p 3 6 d 5 10 f 7 14

Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy Levels Principle energy level (n) Type of sublevel Number of orbitals per type Number of orbitals per level(n2) Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7   

8. The Pauli Exclusion principle states that a GROUND STATE ELECTRON CONFIGURATION 7. Electrons can spin in one of two directions. () represents one direction and () represents the opposite direction 8. The Pauli Exclusion principle states that a maximum of two electrons may occupy a single orbital, and the electrons must have opposite spins

GROUND STATE ELECTRON CONFIGURATION 9. An atomic orbital containing paired electrons with opposite spin are written as  10. Hund’s Rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals

ORBITAL DIAGRAMS AND ELECTRON CONFIGURATION NOTATIONS 1. Two methods of representing an atom’s electron configuration are: a) orbital diagram b) electron configuration notation SEE Table 5-3, Page 137

2. Using Sodium (Na) as an example: Na 1s22s22p63s1       1s 2s 2p 3s 3. Electron configurations for noble gases use bracketed symbols Example: a) Helium = [ He] b) Neon = [ Ne ]

ORBITAL DIAGRAMS AND ELECTRON CONFIGURATION NOTATIONS 4. The electron configuration for an element can be represented using the noble-gas notation for the noble gas in the previous period and the electron configuration for the energy level being filled Example: Sodium (Na) = [Ne]3s1

Topic for 11.03.05 Ch. 5 Electrons in Atoms Discuss Valence electrons and Electron-dot structure

ORBITAL DIAGRAMS AND ELECTRON CONFIGURATION NOTATIONS 5. SEE Table 5-4 for electron configurations for elements in period three 6. A memory aid called a sublevel diagram is shown in Fig. 5-19, Page 138 Example Problem 5-3, Page 139 Sample Problems, Page 139

VALENCE ELECTRONS 1. Valence electrons are defined as electrons in the atom’s outermost orbital 2. Number of valence electrons is the same as an elements group number (A-families)

VALENCE ELECTRONS 3. Number of energy levels is the same as an elements period number 4. An atom’s electron-dot structure consists of the elements symbol surrounded by dots representing each of the atom’s valence electrons SEE Table 5-5, Page 140

Ch 5 test preparation: Example Problem 5-4, Page 141 Practice Problems, Page 141 Section 5.3 Assessment, Page 141 Chapter 5 Assessment, Pages 146 - 148 Standardized Test Practice - Ch. 5, Page 149