Learning Standards• Atomic Structure Broad Concept: Atomic models are used to explain atoms and help us understand the interaction of elements and compounds observed on a macroscopic scale.,
–Recognize discoveries from Dalton (atomictheory), Thomson (the electron), Rutherford (thenucleus), and Bohr (planetary model of atom) andunderstand how these discoveries lead to themodern theory.–Describe Rutherford’s “gold foil” experiment thatled to the discovery of the nuclear atom. Identify themajor components (protons, neutrons, and electrons)of the nuclear atom and explain how they interact.–Write the electron configurations for the firsttwenty elements of the periodic table.
Atomic theory proposed byJohn DaltonAll matter is composed of atomsAtoms cannot be made or destroyedAll atoms of the same element are identicalDifferent elements have different types ofatomsChemical reactions occur when atoms arerearrangedCompounds are formed from atoms of the
Atom Definition Atom What is inside the atom?
Discovery of Protons• Eugene Goldstein noted streams of positively charged particles in cathode rays in 1886. – Particles move in opposite direction of cathode rays. – Called “Canal Rays” because they passed through holes (channels or canals) drilled through the negative electrode.
Canal rays must be positive. Goldstein postulated the existence of a positive fundamental particle called the “proton”.
Thomson’s Experiment And Discovery of Electrons - Voltage source +Passing an electric current makes a beamappear to move from the negative to thepositive end.
Thomson’s Experiment Voltage source + -By adding an electric field he foundthat the moving pieces were negative.
The electron was discovered in1897 by Thomson. He imaginedthe atom as a “raisin pudding” withelectrons stuck in a cake ofpositive charge.
J.J. Thomson’s Model of Atom • Plum Pudding Model, 1896 • Thought an atom was like plum pudding – Dough was cloud – Raisins were electrons – Didn’t know about neutrons at this time
Rutherford’s experiment and discovery of nucleus• English physicist Ernest Rutherford (1911)• Shot alpha particles at fluorescent screen.• When an alpha particle hits a fluorescent screen, it glows.
FluorescentLead Uranium Screenblock Gold Foil
What he expectedHe ExpectedThe alpha particles to pass through withoutchanging direction very much.
What he got
He thought the mass wasevenly distributed in the atom
The Nuclear AtomSince someparticles weredeflected atlarge angles,Thomson’smodel couldnot be correct.
How Rutherfordexplained results….. • Atom is mostly empty space. • Small dense, positive piece at center is (NUCLEUS)
Rutherford’s Findings Most of the particles passed rightthrough A few particles were deflected VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: a) The nucleus is small b) The nucleus is dense c) The nucleus is positively charged
The model created by Rutherford hadstill some serious discordance.According to the classic science,electron moving around the nucleusshould emit an electromagnetic wave.Electron should than move not by thecircle but helical and finally collide withthe nucleus. But atom is stable.
Rutherford also realized that thenucleus must contain both neutral andpositively charged particles. Theneutron was then discovered in 1932by Chadwick.
Atomic number and Mass number :-• a) Atomic number (Z) :-• The atomic number of an element is the number of protons present in the• nucleus of the atom of the element.• All the atoms of an element have the same atomic number.• Eg :- Hydrogen – Atomic number = 1 (1 proton)• Helium - Atomic number = 2 (2 protons)• Lithium - Atomic number = 3 (3 protons)• b) Mass number (A) :-• The mass number of an element is the sum of the number of protons and• neutrons (nucleons) present in the nucleus of an atom of the element.• The mass of an atom is mainly the mass of the protons and neutrons in the nucleus of the atom.• Eg :- Carbon – Mass number = 12 (6 protons + 6 neutrons) Mass = 12u• Aluminium – Mass number = 27 (13 protons + 14 neutrons) Mass = 27u• Sulphur – Mass number = 32 (16 protons + 16 neutrons) Mass = 32u• In the notation of an atom the Mass number Symbol of 14• atomic number and mass number element E g :- N 7• are written as :- Atomic number
Isotopes• Isotopes are atoms of the same element having the same atomic numbers but different mass numbers.• Eg :- Hydrogen has three isotopes. They are1 Protium, Deuterium 3(D) and Tritium (T). 2• 1H 1H 1 H• Protium Deuterium Tritium• Carbon has two isotopes. They are :- 12 14• 6C 6C• Chlorine has two isotopes They are :- 35 37• 17 Cl Cl 17•
Isobars• Isobars are atoms of different elements having different atomic numbers but same mass numbers.• These pairs of elements have the same number of nucleons.• Eg :- Calcium (Ca) – atomic number - 20 and Argon (Ar) – atomic number 18 have different atomic numbers but have the same mass numbers – 40. 40 40 20 Ca 18 Ar• Iron (Fe) and Nickel (Ni) have different atomic numbers but have the same atomic mass numbers – 58.• 58 Fe 58 Ni 27 28•
Bohr’s Model of the Atom• Similar to Rutherford’s model• Thought atom was mostly empty space• Neils Bohr, 1913 – Nucleus in center is dense, positively charge – Electrons revolve around the nucleus.
Following Rutherford’splanetary model of the atom, itwas realized that the attractionbetween the electrons and theprotons should make the atomunstableBohr proposed a model inwhich the electrons wouldstably occupy fixed orbits, aslong as these orbits had specialquantized locations
Parts of an AtomEach element has a different number of protonsin its nucleus Protons have positive charge p Change the number of protons change elements This is called nuclear physicsThe element also has the same number ofelectrons Electrons have negative charge e Change the number of electrons ionize the element This is called chemistrySome elements also have neutrons Neutrons have no charge n They are in the nuclei of atoms
Subatomic particles ActualName Symbol Charge mass (g)Electron e- 9.11 x 10-28Proton p+ 1.67 x 10-24Neutron n0 1.67 x 10-24
Bohr’s model• Electrons move around the nucleus at stable orbits without emitting radiation.• Electron in one of these stable orbit has a definite energy.• Energy is radiated only when electrons make transitions from high energy orbit to a low energy orbit.
In the Bohr model, the electron canchange orbits, accompanied by theabsorption or emission of a photonof a specific color of light.
Wave Nature of Electromagnetic Radiation Waves have 3 primary characteristics: 1. Wavelength (λ): distance between two consecutive peaks in a wave. 2. Frequency (ν): number of waves (cycles) per second that pass a given point in space. 3. Speed: speed of light is 2.9979 * 108 m/s. We will use 3.00 x108 m/s.Wavelength and frequency can be interconvert and they have an inverse relationshipv = c/λ v = frequency (s1) λ = wavelength (m) c = speed of light (m s1)Wavelength is also given in nm (1 nm = 10-9 m) and Angstroms (Å) (1 Å = 10-10 m).The frequency value of s1 or 1/s is also called “hertz (Hz)” like KHz on the radio.
The Particle Nature of Light•Blackbody radiation•The photoelectric effect
Blackbody Radiation and the Quantization of Energy A BA. The interior of a cold ceramic-firing kiln approximates a blackbody, anobject that absorbs all radiation falling on it and appears black. A hotkiln emits light characteristic of blackbody radiation. B Planck’s formulagenerates a curve that fits perfectly the changes in energy and intensityof light emitted by blackbody at different wavelength for a giventemperature
Planck’s Formula• To find a physical explanation of blackbody Planck made a radical assumption that the hot, glowing object could emit (or absorb) only certain quantities of energy:• E = nhν• Where E is the energy of the radiation, ν is its frequency, n is a positive integer (1, 2, 3 and so on) called a quantum number and h is a proportionality constant now called Planck’s constant and has value = 6,626x10-34 J.s
The Photoelectric Effect and The Photon Theory of Light • Current flow when monochromatic light of sufficient energy shines on a metal plate • The photoelectric effect had certain features: the presence of a threshold frequency and the absence of a time lag • Carrying Planck’s idea of packeted energy, Einstein proposed that light itself is particulate, occurring as quanta of electromagnetic energy, called photon • In terms of Planck’s work we can say that each atom changes its energy whenever it absorbs or emits one photon, one “particle” of light, whose energy is fixed by its frequency • Ephoton = hν = ∆Eatom
Bohrs model for hydrogen atom • 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). • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be • Niels Bohr from the Planck’s assumption and radiated adopted atom. explained these phenomena in this way: 3.Energy is only absorbed or emitted in sucha way as to move an electron from one “allowed” energy state to another; the energy is defined by E = hν
Bohrs model for hydrogen atom • Lyman series The atom will remain in the excited state for a short time before emitting a photon and returning to a lower stationary state. All hydrogen atoms exist in n = 1 (invisible). • Balmer series When sunlight passes through the atmosphere, hydrogen atoms in water vapor absorb the wavelengths (visible). H atoms exist in n=2. Similarly it will fill in: • Paschen From n=4,5……… till n=3 • Brackutt from n=5,6……… till n=4 • Pfund from n=6,7……… till n=5
Bohrs model for hydrogen atomThe energy absorbed oremitted from the process ofelectron promotion ordemotion can be calculatedby the equation: 1 1 ∆E = −RH ( nf2 - n2 ) iwhere RH is the Rydbergconstant, 2.18 × 10−18 J, andni and nf are the initial andfinal energy levels of theelectron.
Limitation of Bohr’s Model• Bohr’s model only works for hydrogen atom and other one-electron (hydrogen-like) ionic species, such as He+, Li2+, etc.• For H-atom, electronic energy: En = -2.178 x 10-18 J(1/n2)• For other one-electron particle: En = -2.178 x 10-18 J(Z2/n2) – (Z = atomic number)• Bohr’s model cannot explain atomic spectra of atoms having more than one electron;• Bohr’s model also cannot explain why each line in the hydrogen spectrum appears as double-lines if the discharge tube is placed in magnetic field.• Perhaps his treatment of electron as having only particulate properties is insufficient.
Aufbau Principle • As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbital. • H : 1s1, He : 1s2, Li : 1s2 2s1, Be : 1s2 2s2 • B : 1s2 2s2 2p1, C : 1s2 2s2 2p2.
Hund’s Rule • The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. • N : 1s2 2s2 2p3, O : 1s2 2s2 2p4, • F : 1s2 2s2 2p5, Ne : 1s2 2s2 2p6, • Na : 1s2 2s2 2p63s1 OR [Ne] 3s1
Heisenberg Uncertainty Principlex = position hmv = momentum ∆ x ⋅ ∆ (m v ) ≥h = Planck’s constant 4πThe more accurately we know a particle’s position, theless accurately we can know its momentum. Both theposition and momentum of a particle can not bedetermined precisely at a given time. The uncertaintyprinciple implies that we cannot know the exact motion ofthe electron as it moves around the nucleus.
Quantum Numbers (QN)• The principal quantum number (n) is a positive integer (1, 2, 3 and so forth). It indicates the relative size of the orbital and therefore the relative distance from the nucleus of the peak in the radial probability distribution plot Principal QN (n = 1, 2, 3, . . .)• The angular momentum number (l) is an integer from 0 to n- 1. it is related to the shape of the orbital and is sometimes called orbital-shape quantum number Angular Momentum QN (l = 0 to n 1)• The magnetic quantum number (ml) is an integer from –l through 0 to +l. it prescribes the orientation of the orbital in the space around the nucleus and is sometimes called the orbital- orientation quantum number Magnetic QN (ml = l to l including 0)
Quantum Numbers (QN)• Spin Quantum Number, (s)This led to a fourth quantum number, the spin quantum number, ms. Electron Spin QN has only 2 allowed values: +1/2 and −1/2.
Pauli Exclusion Principle • No two electrons in the same atom can have exactly the same energy. • No two electrons in the same atom can have identical sets of quantum numbers. (n, l, ml, ms). • Therefore, an orbital can hold only two electrons, and they must have opposite spins.