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## Unit 2 contd

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**Unit 2 contd**Chapter 6 and 7.1-7.6**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. study light!**The Wave Nature of Light**• So, to understand electronic structure, we must learn about light. • Light is a type ofelectromagnetic radiation • 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 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 through vacuum at same speed: Speed of light c = 3.00 x 108 m/s c = ln • = wavelength in nm or m n = 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 = ln n = c/l • KGOU broadcast at a frequency of 106.3 MHz (megahertz, 1 MHz = 106 s-1). What is the wavelength of this radiation? c = ln l = c/n**The electromagnetic spectrum:**Wavelengths of g 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 n 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**• Red-hot object is cooler than white-hot object….why? T and l of radiation • Max Planck explained it by assuming that energy comes in packets called quanta. • Quantum:the smallest quantity of energy that can be emitted (released) or absorbed as electromagnetic energy**Quantized Energy and Photons**• Planck (1900) proposed that the energy of a single quantum is directly proportional to its frequency: E = hn where E = energy n = frequency h = Planck’s constant (6.63x10-34 Joule second or Js)**Quantized Energy and Photons**• According to Planck’s theory,energy is always emitted or absorbed in whole number multiples of hn(i.e hn, 2hn, 3hn) • 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 10-34 J-s). • A quantum of energy (E = hn) 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 = hn**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, n, of 4.69 x 1014 s-1. What is the energy of one photon of the radiation from this laser? E = hn h =6.63 x 10-34 Js E = 6.63 x 10-34 Js x 4.69 x 1014 s-1 = 3.11 x 10-19 J**More practice**What is the energy of one photon of yellow light with a wavelength of 589 nm? E = hn and c = ln n = c/l So… E = h c/l h = 6.63 x 10-34 Js and c = 3.00 x 108 m/s E = 3.37 x 10-19 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.**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 AtomNiels Bohr: 1922 Nobel Prize in**physics for model of hydrogen atom According to the Bohr Model of the atom: • Electrons move in circular orbitsaround 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 Bohr Model of the Atom**n=4 n=3 n=2 energy n=1 nucleus**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 • 0.529 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**• 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.**Bohr Model of the Atom**• Niels Bohr (Nobel Prize 1922): • electron in a hydrogen atom • Classical physics do not apply in the atom. • 2. Electrons orbit the nucleus, but only orbits of certain energies are allowed. • 3. Electrons can change from one allowed orbit to another, but that the change will either require energy or produce energy. The energy involved is often in the form of light. n=4 n=3 n=2 Increasing energy n=1 An electron in the lowest energy level (ground state) of H.**Sodium Lamp**The characteristic yellow light in a sodium lamp is the result of electrons in the high-energy “3p” orbital falling back to the lower-energy “3s” orbital. What does that mean?**The Bohr Model of the Atom**• The Bohr model effectively explains the line spectra of atoms and ions with a single electron • H, He+, Li2+ • Another model is needed to explain the reactivity and behavior of more complex atoms or ions • Quantum mechanical model**Heisenberg (Nobel Prize 1932)**• If electron has properties of both a particle and a wave (dual nature), it is impossible to know the exact position and momentum (mass times speed) simultaneously. Uncertainty principle Bottom line: Electrons don’t move in well-defined circular orbits around the nucleus. Only important for masses as small as an electron!!**Erwin Schrödinger developed a mathematical treatment into**which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics.**Quantum Numbers**• Solving the Schrödinger wave equation gives a set of “wave functions”, called orbitals, and their corresponding energies. • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers. = a way of expressing the probability of finding an electron at a particular location in space**Orbitals**• An orbital: • describes a specific distribution of electron density in space • has a characteristic energy • has a characteristic shape • is described by three quantum numbers: n, l, ml • can hold a maximum of 2 electrons • Note:A fourth quantum number (ms) is needed to completely describe each electron in an orbital**Quantum Numbers**• 1) Principal quantum number (n): • integral values n = 1, 2, 3, 4, .. • indicates the average distance of the electron from the nucleus • as n increases, the average distance from the nucleus increases n=4 n=3 n=2 Increasing energy n=1**Quantum Numbers**2) Angular momentum quantum number (l) • integral values l = 0, 1, 2,….(n-1) Example: If n = 4, then l = 0, 1, 2, or 3. • defines the shape of the orbital • The value for l from a particular orbital is usually designated by the letters s, p, d, f, and g: 01 2 3 4 s p d f g Value of l Letter used**Translate Quantum Numbers into Orbitals**• An orbital with quantum numbers of n = 3 and l = 2 would be a 3d orbital • An orbital with quantum numbers of n = 4 and l = 1 would be a 4p orbital**Quantum Numbers**3) Magnetic Quantum Number (ml) • The magnetic quantum number describes the three-dimensional orientation of the orbital. • Allowed values of mlare integers ranging from - l to l: − l ≤ ml≤ l. • Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.**Orbitals**• Orbitals with the same value of n form a shell. • Different orbital types within a shell are subshells. −l ≤ ml ≤ l 0…(n-1) Can draw orbital diagrams**s Orbitals**n = 1, 2, 3, 4,…… l = 0, 1, 2….(n-1) where 0=s, 1=p, 2=d, 3=f ml = - l ….0…+ l • The value of lfor s orbitals is 0 and therefore ml=0 • They are spherical in shape. • The radius of the sphere increases with the value of n. • found in all shells of an atom**s Orbitals**Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.