Quiz Review

Light and quantized Energy Section 1 Electrons in Atoms Chapter 5

Unanswered Questions What have we learned? Subatomic particles exist The nucleus of an atom contains all of the atoms positive charge (Protons) and mass (Neutrons + Protons) Atoms are surrounded by fast, moving electrons

But there were some unanswered questions: How the atoms electrons were arranged in space Why the electrons were not pulled into the positively charged nucleus. Why there were differences and similarities in the chemical behavior among various elements

Unraveling of the puzzle: Unanswered Questions Unraveling of the puzzle: Scientists observed that certain elements emitted visible light when heated in a flame. Analysis of this light revealed that an elements chemical behavior is related to the arrangement of the electrons in its atoms. To understand this we have to first understand the wave nature of light

The Wave Nature of Light Electromagnetic Radiation: Is a form of energy that exhibits wave like behavior as it travels through space Visible light is a type of electromagnetic radiation Other examples include those from: Microwaves X-rays Signals sent from the Radio & TV station to your home

The Wave Nature of Light

The Wave Nature of Light Characteristics of Waves Wavelength (λ): the shortest distance between equivalent points on a continuous wave. Usually expressed in meters (m) Frequency (ν): the number of waves that pass a given point per second. Expressed as: Hertz (Hz) = One wave per second Waves = Waves/ second (1/s) or (s-1) Example: 652 Hz = 652 waves/second = 652/s = 652 s-1

The Wave Nature of Light Characteristics of Waves Amplitude: Wave’s height from the origin to a crest or origin to trough Electromagnetic Wave Relationship: Speed of Electromagnetic waves = 3.00x108 m/s Speed of light (c) = 3.00x108 m/s Equation: c = λv where: c = 3.00x108 m/s λ = wave length ν = frequency

The Wave Nature of Light Waves can have different wavelengths & frequencies Wavelengths and frequency are inversely related

Electromagnetic Spectrum

Practice Problems 1—4 page 140

Practice Problems 1—4 page 140

Practice Problems 1—4 page 140

Practice Problems 1—4 page 140

The Particle Nature of Light The wave nature of light does not explain certain phenomenon Light emission of certain metals Photoelectric effect Studied by German physicist Max Planck (1858-1947) in 1900

The Particle Nature of Light Planck’s conclusions: Matter can gain or lose energy only in small, specific amounts called quanta. Quantum: minimum amount of energy that can be gained or lost by an atom Planck’s constant (h) = 6.626 × 10-34 J ● s

The Photoelectric Effect Photoelectric effect: electrons are emitted from the surface of a metal when light of a certain frequency shines on the metal

The Photoelectric Effect Albert Einstein proposed in 1905 that light has a dual nature. “A beam of light has wavelike and particle-like properties.” Light particle: photon a particle of electromagnetic radiation with no mass that carries a quantum of energy

The Photoelectric Effect

Practice Problems 5—7 page 143

Atomic Emission Spectrum

Atomic Emission Spectrum Atomic emission spectrum of an element: the set of frequencies of the electromagnetic waves emitted by the atoms of the element

Atomic Emission Spectrum Cont’d Unique for each element and can be used to identify an element or determine whether an element is part of an unknown compound.

Atomic Emission Spectrum Cont’d The atomic emission spectrum is not continuous (made up of certain frequencies of light)

Quantum Theory and the Atom Section 2 Electrons in Atoms Chapter 5

Bohr’s model of the atom Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena (photoelectric effect) but not why atomic emission spectra of elements were discontinuous rather continuous. In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question.

Bohr’s model of the atom Elements have certain allowable energy states Ground State: Lowest energy state of an atom Excited Sate: The state of an atom when it gains energy

Bohr’s model of the atom Cont’d An atoms energy state is related to the location of the electron. Electron moves around the nucleus only in certain allowed circular orbits The smaller the electrons orbit the lower the atoms energy state or energy level The larger the electrons orbit the higher the atoms energy state or energy level

Bohr’s model of the atom Cont’d The energy state of an electron is related to its quantum number The quantum number is number of the orbit in an atom

Ex.Hydrogen has one electron which exists in the ground state or quantum # of n= 1 When energy is absorbed the electron moves to a higher energy orbit e.g. n=2 (the excited state)

The electron of the atom can return from: The excited state n=2 to its Ground state n=1 by emitting (releasing) energy as a photon

This energy difference can be calculated using the formula Δ E = E higher-energy orbit – E lower-energy orbit = E photon = hv

The Quantum Mechanical Model of the Atom Bohr’s model was flawed: Electrons do not move in a circular orbit Heisenberg uncertainty principle: It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time

The Quantum Mechanical Model of the Atom The quantum mechanical model makes no attempt to describe the electrons path around nucleus Boundary encompasses the 95% probability that the electron is located there at any given time

The Quantum Mechanical Model of the Atom Uses 4 numbers to “address” an electron in an atom. We will only work with the principal quantum # Principal quantum number, n Energy Sublevels is equal to the number of the principal quantum number

The Quantum Mechanical Model of the Atom Shape of Orbitals Sublevels are labelled s, p, d or f depending on the shape of the orbital s- Spherical p – Dumbbell shaped d – Same shape different planes f – Same shape different planes

The Quantum Mechanical Model of the Atom s Orbitals Have only 1 orbital Increase in size with increasing quantum number

The Quantum Mechanical Model of the Atom p Orbitals Have 3 orbitals Oriented along three axes: x, y & z

The Quantum Mechanical Model of the Atom d Orbitals Have 5orbitals Oriented along different planes: xy, xz, yz, x2-y2 Z2 – shaped differently oriented differently

The Quantum Mechanical Model of the Atom f Orbitals Have 7 orbitals Complex multilobed shapes

The Quantum Mechanical Model of the Atom Shape of Orbitals Each orbital contains at most 2 electrons Principal Quantum # Sublevels (type of Orbitals) # of orbitals related to sublevel n = 1 1s 1 n=2 2s 2p 3 n=3 3s 3p 3d 5 n=4 3f 7

The Quantum Mechanical Model of the Atom Uses four numbers to “address” an electron in an atom Principal quantum number, n Energy Sublevels is equal to the number of the principal quantum number Angular momentum quantum number, l Magnetic quantum number, m Spin quantum number, + ½ or – ½

Electron Configuration Section 3 Electrons in Atoms Chapter 5

Ground state electron configuration Electron Configuration: Arrangement of electrons in an atom The Aufbau principle: Each electron occupies the lowest energy orbital available because they are more stable at this level. Hund’s rule: 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

Ground state electron configuration The Pauli exclusion principle: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

Using the Aufbau principle: Each electron occupies the lowest energy orbital available because they are more stable at this level. Order of filling orbitals

Order of filling orbitals

The Pauli exclusion principle: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

Hunts Rule: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

Ground state electron configuration

Ground state electron configuration

Writing the electron configuration Neon Electron Configuration: 1s2 2s2 2p6

Ground state electron configuration Only works perfectly for all elements up to and including Vanadium, atomic number 23

Practice Problem 21-25 p 160

Ground state electron configuration Noble Gas notation: A method of representing electron configurations of noble gases.

Ground state electron configuration Noble Gases: Have eight electrons in their outermost orbital Usually unstable Noble Gas Notation: Uses bracketed symbols For example: Helium (He)

Ground state electron configuration For example:

Ground state electron configuration For example:

Ground state electron configuration Valence electrons Defined as atom’s outermost orbitals Atoms in the outermost orbitals are associated with the atom’s highest principal energy level For example:

Ground state electron configuration Electron-dot structure Devised by American Chemist G.N. Lewis Aka Lewis Dot structures Used by chemists to represent Valence Electrons Writing the Lewis Dot Structure: Element’s Symbol: Represents the atomic nucleus and inner level electrons.

Ground state electron configuration Dots: Represent Valence electrons Placed one at a time around the four sides of the symbol and then paired up until all are shown Example:

Practice Problems Pg. 162 #26-28

Practice Problems 26-28 p 162

Practice Problems 26-28 p 162

Practice Problems 26-28 p 162

Practice Problems 26-28 p 162

Practice Problems 26-28 p 162