AP Chemistry - Chapter 7 Notes - Atomic Structure and Periodicity

Sep 22, 2026

Uploaded by Himbo

AP Chemistry Chapter 7 Notes cover electromagnetic radiation and the development of atomic theory through quantum ideas. It explains wavelength, frequency, and wave speed, and derives relationships among them. The notes introduce Planck’s quantum theory, Einstein’s photon concept, and de Broglie’s wave-particle duality, then summarize hydrogen’s emission spectra and the Bohr model’s energy levels and limitations. Finally, it presents the quantum mechanical model, Schrödinger’s equation, uncertainty principle, wave-function interpretation, and quantum numbers.

LanguageEnglish

CategoryExam

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Pages9

Reading Time23 min

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7.1 Electromagnetic Radiation

A. Types of EM Radiation

B. Properties of EM Waves

7.2 The Nature of Matter

A. Max Planck and Quantum Theory

B. Einstein and the Particle Nature of Matter

C. de Broglie and the Dual Nature of Light

7.3 The Atomic Spectrum of Hydrogen

A. Continuous spectra

B. Bright line spectra

7.4 Bohr Model (Neils Bohr, 1913)

A. Quantum Model

B. Calculating the energy of the emitted photon

C. Energy Change in Hydrogen atoms

D. Shortcomings of the Bohr Model

7.5 The Quantum Mechanical Model of the Atom

A. The Electron as a standing wave

B. The Schrödinger Equation

C. Heisenberg Uncertainty Principle

D. Physical Meaning of a Wave Function

7.6 Quantum Numbers

A. Principal Quantum Number (n)

B. Angular Momentum Quantum (l)

C. Magnetic Quantum Number (ml)

D. Spin Quantum Number

AP Chemistry A. Allan Chapter 7 Notes-Atomic Structure and Periodicity 7.1 Electromagnetic Radiation A. Types of EM Radiation (wavelengths in meters) 10-12 10-10 10-8 4 to 7x10-7 10-4 10-2 1 102 104 gamma xrays UV visible IR micro Radio waves FM short AM Wavelength increases Frequency decreases Energy decreases Speed is constant = 2.9979 x 108 m/sec B. Properties of EM Waves 1. Wavelength (λ) a. Distance between two consecutive peaks or troughs in a wave b. Measured in meters (SI system) 2. Frequency (ν ) a. Number of waves that pass a given point per second b. Measured in hertz (sec-1) 3. Speed ( c ) a. Measured in meters/sec 4. Relationship of properties a. λ ν = c 7.2 The Nature of Matter A. Max Planck and Quantum Theory 1. Energy is gained or lost in whole number multiples of the quantity hv Frequency = v Planck's constant = h = 6.626 x 10-34 J•S ∆E = nhv 2. Energy is transferred to matter in packets of energy, each called a quantum B. Einstein and the Particle Nature of Matter 1. EM radiation is a stream of particles- "photons"Ephoton = hv = ~~ ~~hc~~λ ~~ 2. Energy and mass are inter-related E 2 = mc C. de Broglie and the Dual Nature of Light 1. Light travels through space as a wave 2. Light transmits energy as a particle 3. Particle's have wavelength, exhibited by diffraction patterns λ = hmv~~ ~~ a. large particles have very short wavelengths b. All matter exhibits both particle and wave properties 7.3 The Atomic Spectrum of Hydrogen A. Continuous spectra 1. Contains all wavelengths of light B. Bright line spectra | | | | | | | | | | | --- | --- | --- | --- | --- | --- | --- | --- | --- | 410nm 434nm 486nm 656nm 1. Excited electrons in an atom return to lower energy states 2. Energy is emitted in the form of a photon of definite wavelength 3. Definite change in energy corresponds to: a. Definite frequency b. Definite wavelength ∆E = hv = ~~ ~~hc~~λ ~~ 4. Only certain energies are possible within any atom 7.4 Bohr Model (Neils Bohr, 1913) A. Quantum Model 1. The electron moves around the nucleus only in certain allowed circular orbits 2. Bright line spectra confirms that only certain energies exist in the atom, and atom emits photons with definite wavelengths when the electron returns to a lower energy state 3. Energy levels available to the electron in the hydrogen atom −18 􀂧􀂨 Z 2 􀂷􀂸 E = − 2 .178 x 10 J 􀂨 2 􀂸 􀂩 n 􀂹 n = an integer Z = nuclear charge J = energy in joules B. Calculating the energy of the emitted photon 1. Calculate electron energy in outer level 2. Calculate electron energy in inner level 3. Calculate the change in energy (∆E) ∆ E = energy of final state-energy of initial state 4. Use the equation: λ = cE~~ ~~ to calculate the wavelength of the emitted photon C. Energy Change in Hydrogen atoms 1. Calculate energy change between any two energy levels ∆E = − 2 . 178 x 10−18 J 􀂧􀂩 ~~ ~~ninal~~ ~~ − ~~ ~~nitial~~ ~~ 􀂷􀂹 D. Shortcomings of the Bohr Model 1. Bohr's model does not work for atoms other than hydrogen 2. Electron's do not move in circular orbits 7.5 The Quantum Mechanical Model of the Atom A. The Electron as a standing wave 1. Standing waves do not propagate through space 2. Standing waves are fixed at both ends 3. Only certain size orbits can contain whole numbers of half wavelengths a. fits the observation of fixed energy quantities - "quanta" B. The Schrödinger Equation (Erwin Schrödinger) 1. For the motion of one particle, in the along the x axis in space: h2 d 2ψ − 2 2 + V ψ = Eψ 8 π m dx 2. Solution of the equation has demonstrated that E (energy) must occur in integer multiples 3. Your book has done you the favor of greatly simplifying this and presents the general equation: ∧ Hψ = Eψ ∧ H = a set of mathematical functions called an "operator" ψ = a wave function. Specific wave functions are ca

FAQs

What are the main properties of electromagnetic waves in Chapter 7 Notes?

The notes define wavelength, frequency, and speed, explain how each is measured, and give the relationship λν = c.

How do Planck’s and Einstein’s ideas describe energy exchange in matter?

Planck’s model states energy is gained or lost in whole-number multiples of hv, and Einstein describes electromagnetic radiation as a stream of photons with Ephoton = hv.

What does the Bohr model say about electron orbits and why is it limited?

The model restricts electrons to certain allowed circular orbits and explains bright-line spectra using definite energy levels. It does not work for atoms other than hydrogen and electrons are not actually moving in circular orbits.

Sep 22, 2026

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