Presentation on theme: "Unit 6 Chapter 6 and 7.1-7.6. Why are our bodies so large compared to an atom? Why is the atom so small? Think about it!"— Presentation transcript:
Unit 6 Chapter 6 and
Why are our bodies so large compared to an atom? Why is the atom so small? Think about it!
Behavior of electrons and atoms Wave Nature of Light Models of the Atom –Bohr Model –Quantum Mechanical Model Atomic Orbitals Electron Configurations Periodic Properties of Elements
Electronic Structure of the Atom Elements in the same group exhibit similar chemical and physical properties. –Alkali Metals: soft very reactive metal –Noble Gases gases inert (unreactive) Why???
Electronic Structure of the Atom When atoms react, their electrons interact. The properties of elements depend on their electronic structure. –the arrangement of electrons in an atom number of electrons distribution of electrons around the atom energies of the electrons
Electronic Structure of the Atom Understanding the nature of electrons and the electronic structure of atoms is the key to understanding the reactivity of elements and the reactions they undergo. Much of our knowledge of the electronic structure of atoms came from studying the ways elements absorb or emit light.
The Wave Nature of Light So, to understand electronic structure, we must learn about light. Light is a type of electromagnetic radiation –a form of energy with both electrical and magnetic components
Waves The nature of electromagnetic radiation: wavelike characteristics (like water waves). Describing waves: The distance between corresponding points on adjacent waves is the wavelength ( ).
Waves The number of complete wavelengths passing a given point per unit of time is the frequency ( ). For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.
Electromagnetic Radiation All radiation moves at same speed Speed of light c = 3.00 x 10 8 m/s c = wavelength in nm or m frequency in ‘per second’ (s -1 ) High frequency small/short wavelength Long wavelength low frequency
Think What is the frequency of green light with a wavelength of 520 nm? c = = c/ KGOU broadcast at a frequency of MHz (megahertz, 1 MHz = 10 6 s -1 ). What is the wavelength of this radiation? c = = c/
The electromagnetic spectrum: Wavelengths of rays: atomic nuclei Wavelengths of radio waves: football field
The Wave Nature of Light Different types of electromagnetic radiation have different properties because they have different u and l. –Gamma rays wavelength similar to diameter of atomic nuclei –Hazardous –Radio waves wavelength can be longer than a football field
The Nature of Energy The wave nature of light does not explain how an object can glow when its temperature increases. Max Planck explained it by assuming that energy comes in packets called quanta. Quantum: the smallest quantity of energy that can be emitted or absorbed as electromagnetic energy
Quantized Energy and Photons Planck proposed that the energy of a single quantum is directly proportional to its frequency: E = hu where E = energy u = frequency h = Planck’s constant (6.63x Joule-second or J-s)
Quantized Energy and Photons According to Planck’s theory, energy is always emitted or absorbed in whole number multiples of h u (i.e h u, 2h u, 3h u ) According to Planck’s theory, the energy levels that are allowed are ‘quantized.’ –restricted to certain quantities or values
Quantized Energy and Photons In order to understand quantized energy levels, compare walking up (or down) a ramp versus walking up (or down) stairs: –Ramp: continuous change in height –Stairs: quantized changed in height You can only stop on the stairs, not between them
Quantized Energy and Photons If Planck’s quantum theory is correct, why don’t we notice its effects in our daily lives? Planck’s constant is very small (6.63 x J-s). –A quantum of energy (E = hu) is very small. Gaining or losing such a small amount of energy is: –insignificant on macroscopic objects –very significant on the atomic level
1905: Einstein used Planck’s quantum theory to explain photoelectric effect. “photocells”
Photoelectric Effect –Light shining on a clean metal surface causes the surface to emit electrons. –The light must have a minimum frequency in order for electrons to be emitted.
Quantized Energy and Photons Einstein explained these results by assuming that the light striking the metal is a stream of tiny energy packets of radiant energy (photons). –The energy of each photon is proportional to its frequency. E = hu
Quantized Energy and Photons When a photon strikes a metal surface: –Energy is transferred to the electrons in the metal If the energy is great enough, the electron can overcome the attractive forces holding it to the metal. Any extra energy above the amount required to “free” the electron simply increases the kinetic energy of the electron.
Think! A laser emits light with a frequency,, of 4.69 x s -1. What is the energy of one photon of the radiation from this laser? E = h h = 6.63 x J-s E = 6.63 x J-s x 4.69 x s -1 = 3.11 x J
More practice What is the energy of one photon of yellow light with a wavelength of 589 nm? E = h and c = = c/ So… E = h c/ h = 6.63 x J-s and c = 3.00 x 10 8 m/s E = 3.37 x J
Quantized Energy and Photons Einstein’s explanation of the photoelectric effect led to a dilemma. –Is light a wave or does it consist of particles? Currently, light is considered to have both wave-like and particle-like properties. Matter also has this same dual nature.
Models of Atomic Structure Scientists initially thought of the atom as a “microscopic solar system.” –electrons orbiting the nucleus Unit 2 suggested that the atom has a tiny positively charged nucleus with a diffuse “cloud” of electrons surrounding it. –need better understanding of the nature of this “cloud” of electrons.
Atomic Models Two models are used to explain the behavior and reactivity of atoms and ions. Bohr Model And Quantum Mechanical Model
The Nature of Energy Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.
The Nature of Energy For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed.
Bohr Model Bohr developed an atomic model that explained the line spectrum observed for the hydrogen atom. High voltage H2H2 When an electrical current is passed thru a sample of H 2 (g), energy is transferred to the H 2 molecules. The molecules are broken up. The H atoms absorb energy and “jump” to a higher energy level.
The Bohr Model of the Atom High voltage H2H2 The H atoms “relax” back to their original energy level by giving off the absorbed energy as electromagnetic radiation.
The Bohr Model of the Atom High voltage H2H2 The light is analyzed in a spectrometer by separating it into its different colors.
The Bohr Model of the Atom High voltage H2H2 The separated colors are recorded as spectral lines. Atomic spectrum
The Bohr Model of the Atom The spectrum of atomic hydrogen consists of a series of discrete lines such as the ones shown previously. Why would an atom emit only certain frequencies of light and not all of them?
The Bohr Model of the Atom According to the Bohr Model of the atom: Electrons move in circular orbits around the nucleus. Energy is quantized: - only orbits of certain radii corresponding to certain definite energies are allowed - an electron in a permitted orbit has a specific energy (an “allowed energy state”)
The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 1.Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
The Nature of Energy 2.Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
The Nature of Energy 3.Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h
The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: E = −R H ( ) 1nf21nf2 1ni21ni2 - where R H is the Rydberg constant, 2.18 10 −18 J, and n i and n f are the initial and final energy levels of the electron. n = principal quantum number
The Bohr Model of the Atom Each orbit in an atom corresponds to a different value of n. –As n increases, the radius of the orbit increases (i.e. the orbit and any electrons occupying it are further from the nucleus) –n=1 is the closest to the nucleus Ångstroms for the hydrogen atom
The Bohr Model of the Atom The energy of the orbit is lowest for n=1 and increases with increasing n. –Lower energy = more stable –Lower energy = more preferred state
The Bohr Model of the Atom The lowest energy state of an atom is called the ground state. –n = 1 for the electron in a H atom When an electron has “jumped” to a higher energy orbit (i.e. n = 2, 3, 4…) it is considered to be in an excited state.
The Bohr Model of the Atom To explain the line spectrum for hydrogen, Bohr assumed that an electron can “jump” from one allowed energy state to another. –Energy absorbed e - “jumps” to higher energy state –e - “relaxes” back to a lower energy state energy is emitted
The Bohr Model of the Atom n=1 n=2 n=3 n=4 energy
The Bohr Model of the Atom Since the energies of the orbits in an atom are quantized, transitions from one allowed orbit to another involves only specific amounts of energy. D E = E f - E i
The Bohr Model of the Atom Since E = hu, the energy of the light emitted can have only specific values. –Therefore the u of the light can have only specific values as well. –So, the line spectrum for each element will be unique and will depend on the “allowed” energy levels in that element.
The Bohr Model of the Atom The Bohr model effectively explains the line spectra of atoms and ions with a single electron –H, He +, Li 2+ Another model is needed to explain the reactivity and behavior of more complex atoms or ions –Quantum mechanical model