1. The Atom and Unanswered Questions
Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it.
The model doesn’t explain how the electrons were arranged around the nucleus.
The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.

2. The Atom and Unanswered Questions
In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame.

3. The Wave Nature of Light
Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space.
All waves can be described by several characteristics.

4. The Wave Nature of Light (cont.)
The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.
The frequency (ν) is the number of waves that pass a given point per second.
The amplitude is the wave’s height from the origin to a crest.

5. The Wave Nature of Light (cont.)

6. The Wave Nature of Light (cont.)
The speed of light (3.00  108 m/s) is the product of it’s wavelength and frequency c = λν.

7. The Wave Nature of Light (cont.)
Sunlight contains a continuous range of wavelengths and frequencies.
A prism separates sunlight into a continuous spectrum of colors.
The electromagnetic spectrum includes all forms of electromagnetic radiation.

8. The Wave Nature of Light (cont.)

9. The Particle Nature of Light
The wave model of light cannot explain all of light’s characteristics.
Matter can gain or lose energy only in small, specific amounts called quanta.
A quantum is the minimum amount of energy that can be gained or lost by an atom.
Planck’s constant has a value of  10–34 J ● s.

10. The Particle Nature of Light (cont.)
The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.

11. The Particle Nature of Light (cont.)
Albert Einstein proposed in 1905 that light has a dual nature.
A beam of light has wavelike and particlelike properties.
A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.
Ephoton = h Ephoton represents energy. h is Planck's constant.  represents frequency.

12. Atomic Emission Spectra (cont.)

13. Atomic Emission Spectra (cont.)
The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.
Each element’s atomic emission spectrum is unique.

14. Bohr’s Model Of The Atom
Electrons move around the nucleus in fixed paths or orbits much like the planets move around the sun
Orbit positions, labeled with the integer n, have specific potential energy
The lowest energy state of an atom is called the ground state (an electron with n = 1 for a hydrogen atom)

15. Atomic Emission Spectra
The excited atoms emit light to release energy.

16. Absorption And Emission
Electrons that absorb energy are raised to a higher energy level
A particular frequency of light is emitted when an electron falls to a lower energy level
FIG. 7.9 Absorption of energy and emission of light by the hydrogen atom. When the atom absorbs energy, the electron is raised to a higher energy
level. When the electron falls to a lower energy level, light of a particular energy and frequency is emitted. 16 16

17. Bohr's Model of the Atom
Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.
The lowest allowable energy state of an atom is called its ground state.
When an atom gains energy, it is in an excited state.

18. Bohr's Model of the Atom (cont.)
Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

19. Bohr's Model of the Atom (cont.)
Each orbit was given a number, called the quantum number.

20. Bohr's Model of the Atom (cont.)
Hydrogen’s single electron is in the n = 1 orbit in the ground state.
When energy is added, the electron moves to the n = 2 orbit.

21. Bohr's Model of the Atom (cont.)
Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines.
The behavior of electrons is still not fully understood, but it is known they do not move around the nucleus in circular orbits.

22. Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals.

23. The Quantum Mechanical Model of the Atom
Louis de Broglie (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors.

24. The Quantum Mechanical Model of the Atom (cont.)
The figure illustrates that electrons orbit the nucleus only in whole-number wavelengths.

25. The Quantum Mechanical Model of the Atom (cont.)
The de Broglie equation predicts that all moving particles have wave characteristics.
 represents wavelengths h is Planck's constant. m represents mass of the particle.  represents frequency.

26. The Quantum Mechanical Model of the Atom (cont.)
Heisenberg showed it is impossible to take any measurement of an object without disturbing it.
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.
The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

27. The Quantum Mechanical Model of the Atom (cont.)
Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.
Schrödinger’s equation applied equally well to elements other than hydrogen.

28. The Quantum Mechanical Model of the Atom (cont.)
The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.

29. Hydrogen Atomic Orbitals
Principal quantum number (n) indicates the relative size and energy of atomic orbitals.
n specifies the atom’s major energy levels, called the principal energy levels.

30. Quantum Numbers
Are a shorthand to describe characteristics of an electron’s position and to predict its behavior
n = principal quantum number. All orbitals with the same principle quantum number are in the same shell
l = secondary quantum number which divides the orbitals in a shell into smaller groups called subshells
ml = magnetic quantum number which divides the subshells into individual orbitals
Chem FAQs: What is the principal quantum number?
What is the secondary quantum number?
What is the magnetic quantum number?
Shell (a/k/a Energy level). subshells (a/k/a sublevel).
n Allowed values: the set of positive integers.
l Allowed values: from 0 to (n – 1).
a/k/a “azimuthal quantum number” and the “orbital angular momentum number”.
ml Allowed values: integers from –l to +l. 30 30

31. Quantum Numbers: What Do They Mean?
n = roughly describes a distance of the electrons from the nucleus.
designated by integers: 1, 2, 3, 4, 5, 6, …
l = describes the shape of the orbitals.
designated with numbers : 0, 1, 2, 3, 4, 5……
or with letters: s, p, d, f, g, h
ml =describes the spatial orientation of the orbital.
designated by numbers specific to the particular orbital
range from –l to +l 31

32. “s” Orbitals And Nodes
Orbitals get larger as the principle quantum number n increases
Nodes, or regions of zero electron density, appear beginning with the 2s orbital .
FIG Size variations among s orbitals. The orbitals become larger as the principal quantum number, n, becomes larger. 32

33. “p” Orbitals
Possess a nodal plane that separates the “lobes” of high probability
Dot-density diagrams of the cross section of the probability distribution of a single (a) 2p and (b) 3p orbital showing the nodal plane and the size difference
FIG Distribution of electron density in p orbitals. (a) Dot-density diagram that represents a cross section of the probability distribution in a 2p orbital. There is a nodal plane between the two lobes of the orbital. (b) Cross section of a 3p orbital. Note the nodes in the electron density that are in addition to the nodal plane passing through the nucleus. 33

34. Hydrogen Atomic Orbitals (cont.)
Each energy sublevel relates to orbitals of different shape.

35. Hydrogen Atomic Orbitals (cont.)

36. Ground-State Electron Configuration
The arrangement of electrons in the atom is called the electron configuration.
The aufbau principle states that each electron occupies the lowest energy orbital available.

37. Electrons Behave Like Tiny Magnets
Electrons within atoms interact with a magnet field in one of two ways:
clockwise (spin up)
anti-clockwise (spin down)
This gives rise to the spin quantum number, ms
allowed values: + 1/2 or –1/2
+1/2 (spin “up”) or –1/2 (spin “down”).
Chem FAQs: What is the significance of the spin quantum number ms? 37

38. Ground-State Electron Configuration (cont.)

39. Electron Configuration (cont.)
The Pauli exclusion principle states that a maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins.
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 energy level orbitals.

40. Ground-State Electron Configuration (cont.)

41. Ground-State Electron Configuration (cont.)
Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements.

42. Ground-State Electron Configuration (cont.)
The electron configurations (for chromium, copper, and several other elements) reflect the increased stability of half-filled and filled sets of s and d orbitals.

43. Valence Electrons
Valence electrons are defined as electrons in the atom’s outermost orbitals—those associated with the atom’s highest principal energy level.
Electron-dot structure consists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.

44. Valence Electrons (cont.)