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Atomic concepts Atomic concepts Presentation Transcript

  • ATOMIC CONCEPTS
  • EVOLUTION OF MODERN ATOMIC THEORY Chemistry and the Greeks  Matter composed of the four elements  Earth, Air, Fire, Water
  • DALTON’S ATOMIC THEORY (1803) Atom – The basic unit of matter.  All elements are composed of indivisible atoms.  All atoms of a given element are identical.  Atoms of different elements are different (Have different masses)  Compounds are formed by the combination of different elements. Dalton’s Atomic Model
  • J.J. THOMPSON AND THE ELECTRON Experimental studies of the atom soon showed that it is NOT indivisible but in fact made up of smaller particles.  J.J. Thompson used a cathode ray tube to discover one such particle.
  • J.J. THOMPSON AND THE ELECTRON Beam emitted by the cathode would respond in different ways when exposed to a magnetic field. When the positive end of When the negative end of a magnet was held near a magnet was held near the beam it would move the beam it would move toward the magnet. away from the magnet. (attracted) (repelled) Thompson surmised that the beam was composed of negatively charged particles which he called Electrons
  • J.J. THOMPSON AND THE ELECTRON Electron (e-) – negatively charged subatomic particles that part of an atom.  Thompson’s “Plum Pudding Model” (1897) of the atom visualized electrons as being embedded within the atom. The mass of the rest of the was evenly distributed and positively charged.
  • RUTHERFORD’S GOLD FOIL EXPERIMENT Used alpha particles directed at a thin piece of gold foil which led to the discovery of the Nucleus. Alpha particles are positively charged particles that are much smaller than the atom.
  • RUTHERFORD’S GOLD FOIL EXPERIMENT If Thompson’s plum pudding model was correct the alpha particles would simply pass through the foil with just a few being slightly deflected. Rutherford discovered that while most of the alpha particles did indeed pass through the foil, some were greatly deflected, and some even bounced back.
  • RUTHERFORD’S GOLD FOIL EXPERIMENT Rutherford theorized that the atom must then be composed of mostly empty space, with a dense positively charged core that he called the nucleus. Rutherford’s Atomic Model (1909)
  • PROTONS AND NEUTRONS Atoms electrically neutral so there must be particles to offset the electro-magnetic charge of the negative electrons. Protons (p+) – tiny positively charged particles found within the nucleus of the atom. Neutrons (n0) – tiny particles found within the nucleus of the atom having no electro-magnetic charge
  • MODERN ATOMIC THEORY (BOHR MODEL) Niels Bohr’s “Planetary Model” (1913) of the atom.  Nucleus (protons and neutrons) in the center.  Electrons shown in concentric circles or shells around the nucleus.  Designated by letters K, L, M, N, O, P, Q or the numbers 1 through 7.
  • MODERN ATOM THEORY(WAVE MECHANICAL MODEL) Dual Nature of Matter  Energy viewed as waves and matter as particles.  Electrons exhibit a dual nature in which they not only have mass but possess wavelike properties as well. Wave Mechanical Model  Dense centrally located positive nucleus  Electron no longer pictured in fixed orbits but as regions of differing energy levels where they are most likely to be found called Orbitals
  • SUBATOMIC PARTICLES An atom is the smallest unit of an element. It consists of three major particles.a.m.u. = atomic mass unita.m.u. = 1/12 the mass of a C-12 atom, or, 1.66x10-24 grams.
  • ATOMIC SYMBOLS Written in a shortened form as…. Atomic Mass rounded to the closest whole number.
  • PRACTICE Write a short form for each atomic symbol.
  • DIFFERENCES BETWEEN ATOMS Atomic Number: The number of protons in the nucleus of an atom.  It is also the number of electrons in an electrically neutral atom Atomic Mass (Mass Number):
  • SAMPLE PROBLEM Find the number of neutrons in an atom of ?
  • SAMPLE PROBLEM Find the number of neutrons in an atom of ?
  • DIFFERENCES BETWEEN ATOMS Why are there fractional mass numbers (atomic masses) on the periodic table?
  • DIFFERENCES BETWEEN ATOMS Why are there fractional mass numbers (atomic masses) on the periodic table? Answer: Because of the existence of Isotopes
  • ISOTOPES Atoms having the same number of protons but different number of neutrons.  Example; Average ≈ 22.98977 *Average is based on various isotopic masses and the relative abundances of each.
  • SAMPLE PROBLEM Atomic masses can be calculated from the mass and the abundance of naturally occurring isotopes. Carbon has two naturally occurring stable isotopes. Most carbon atoms (99.89%) are C-12, while the remaining 1.108% are C-13. What is the atomic number of carbon?
  • SAMPLE PROBLEM Atomic masses can be calculated from the mass and the abundance of naturally occurring isotopes. Carbon has two naturally occurring stable isotopes. Most carbon atoms (99.89%) are C-12, while the remaining 1.108% are C-13. What is the atomic number of carbon?
  • SAMPLE PROBLEM Element X has two naturally occurring isotopes. If 72.0% of the element has an isotopic mass of 84.9 amu and the 28.0% has an isotopic mass of 87.0 amu, the average atomic mass of element X is?
  • SAMPLE PROBLEM The average isotopic mass of chlorine is 35.5 amu. Which mixture of isotopes (shown as percents) produces this mass? 1. 50% C-12 and 50% C-13 2. 50% Cl-35 and 50% Cl-37 3. 75% Cl-35 and 25% Cl-37 4. 75% C-12 and 25% C-13
  • DIFFERENCES BETWEEN ATOMS
  • DIFFERENCES BETWEEN ATOMS
  • IONS Atoms of the same element having the same # of protons, but different # of electrons.  No longer electrically neutral, Ions are charged particles  Example;
  • IONS Atoms of the same element having the same # of protons, but different # of electrons.  No longer electrically neutral, Ions are charged particles  Example; 6 p+ 6 e-
  • IONS Atoms of the same element having the same # of protons, but different # of electrons.  No longer electrically neutral, Ions are charged particles  Example; 6 p+ 6 p+ 6 p+ 6 p+ 6 e- 10 e- 4 e- 2 e-
  • QUIZ Identify the number of protons, neutrons, and electrons for each element.
  • ATOMIC MODELS
  • ATOMIC MODELS Bohr Model  K-shell = max. of 2 e-  L-shell = max. of 8 e-  M-shell = max of 18 e-  N-shell = max. of 32 e-
  • BOHR MODEL
  • BOHR MODEL
  • BOHR MODEL
  • BOHR MODEL PRACTICE
  • BOHR MODEL PRACTICE
  • QUIZ Copy the picture and label the following, Atomic Mass, Atomic Number, Electron Configuration, Selected Oxidation State. Atomic ___________indicates the number of __________ within the nucleus of the atom. Atomic __________ is equal to the number of _________ plus the number of __________ within the nucleus of the atom. Selected Oxidation states indicate the most common __________ for a particular element. C-12, C-13 are examples of _____________.
  • ATOMIC MODELS –IMPORTANT DEFINITIONS Principal Energy Level: Region around the nucleus in which electron can be found.  Designated by letters K, L, M, N, O, P, Q or the numbers 1 through 7.  The closer to the nucleus the lower the energy. Quanta: Small amount of energy that an electron can release or absorb as it moves through principle energy levels. Ground State: All electrons fill lowest energy levels before higher energy levels are filled. Excited State: one or more electrons absorb energy (quanta) and occupy a higher principle energy level than Spectral Lines: As electrons at principle higher energy levels (excited state) fall back to their normal principle energy levels (ground state) they emit that extra energy in the form of light.  Visible Spectrum – ROY G BIV
  • EMISSION SPECTRA
  • EMISSION SPECTRA OF THE SUN
  • ORBITAL MODEL(WAVE-MECHANICAL MODEL) Principal Energy (Quantum) Level represents the level in which electrons are found.  These correlate with period number on the periodic table.  These are also your K, L, M, N…. Sublevels are represented by s, p , d, f  Number of sublevels = Principal Energy Level  Principal Energy Level 1 has one sub level (s)  Principal Energy Level 2 has two sub level (s,p) An Orbital is an exact region in which electrons within a principal energy level are most likely to be found.  The maximum number of electron in any orbital is 2
  • Note: The principal energy level is represented by n. Thenumber of Orbitals per level would be n2, and the maximumnumber of electrons per level would be 2n2.
  • SUBLEVELS
  • WRITING ELECTRON CONFIGURATIONS I n c r e a s i n g E n e r g y1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
  • WRITING ELECTRON CONFIGURATIONS
  • EXAMPLES OF ELECTRON CONFIGURATIONS Na S Kr
  • WRITE THE ELECTRON CONFIGURATIONS OFTHE FOLLOWING IONS Na+1 S-2 Kr+2
  • WRITING ELECTRON CONFIGURATIONS FROM THEINFORMATION ON THE PERIODIC TABLE OFELEMENTS.
  • WRITING ELECTRON CONFIGURATIONS FROM THEINFORMATION ON THE PERIODIC TABLE OFELEMENTS.
  • WRITING ELECTRON CONFIGURATIONS FROM THEINFORMATION ON THE PERIODIC TABLE OFELEMENTS.
  • EXCITED STATE ELECTRON CONFIGURATIONS A phosphorus atom has an electron configuration of 1s22s22p63s13p4 . Is the atom in its ground state, or is it in an excited state? P
  • EXCITED STATE ELECTRON CONFIGURATIONS Which electron configuration represents an atom in the excited state? 1. 1s22s22p63s2 2. 1s22s22p63s1 3. 1s22s22p6 4. 1s22s22p53s2
  • VALENCE ELECTRON & LEWIS DOT DIAGRAMS Valence electrons are the electrons that fill the outermost prinicpal energy level of an atom.  Example;  Mg 2-8-2 has 2 valence electrons. 2 2 6  Ne 1s 2s 2p has 8 valence electrons. Valence electrons are largely responsible for an elemement’s chemical and physical properties.
  • DO NOW: HOW MANY VALENCE ELECTRONS DOEACH OF THE FOLLOWING ATOMS HAVE? Na Al Cl Na+1 Al+3 Cl-1 Na Si S-2 Mg+2 Si-4 Ar
  • DO NOW: What is the most common isotope of the element Bromine?  How many protons, neutrons and electrons does it have?  How many valence electrons does it have? What are the most common ions for the element bromine? (Hint: There are three)  How many protons, neutrons and electrons does each ion have?  How many valence electrons does each ion have?
  • VALENCE ELECTRON & LEWIS DOT DIAGRAMS The term Kernel refers to all of the non-valence electrons as well as the nucleus (p+ & n0) of the atom. The Kernel is represented by the element’s symbol. Valence electrons are represented by dots. Na Ne N2-8-1 2-8-8 2-5
  • MORE ON ELECTRONS Similarly, shell electrons may be represented by arrows pointing in opposite directions (up & down) occupying their perspective orbitals. 1. Write the orbital notation for the outermost principal level for the following elements. Na 2-8-1 P 2-8-5 Cl 2-8-7
  • MORE ON ELECTRONS 2. Which is the correct orbital notation of a lithium atom in its ground state
  • MORE ON ELECTRONS 3. Which orbital notation correctly represents a noble gas in the ground state?
  • TABLE S Ionization Energy: the amount of energy needed to remove the most loosely held electron from the valence shell of an atom in the ground state.  Low Ionization energy = EASY to remove e-’s (Fr)  High Ionization Energy = DIFFICULT to remove e-’s (F)  The lowest Ionization Energies are found in the lower left corner of the periodic table (metals)
  • TABLE S Electronegativity: The affinity (attractiveness or pull for) of electrons by an atom.  High Electronegativity = Atoms most likely to gain e- (F)  Low Electronegativity = atoms most likely to lose e- (Fr)  The highest electronegativity (an arbitrary value of 4.0) can be found in the upper right corner of the periodic table (non- metals)  Does not include the Noble Gases (Group 18)
  • ELECTRONEGATIVITY As atom gain or lose e- they become Ions. 0 -1 F F 0 +1 Fr Fr
  • ELECTRONEGATIVITY As atom gain or lose e- they become Ions. High Ionization Energy (-) gains 0 High Electronegativity -1 F F one electron 2-7 2-8 Low Ionization Energy (+) loses 0 +1 Fr Low Electronegativity Fr one electron 2-8-18-32-18-8-1 2-8-18-32-18-8
  • ELECTRONEGATIVITY As atom gain or lose e- they become Ions. 0 -1 F F 2-7 2-8 1s22s22p5 1s22s22p6 0 Fr 2-8-18-32-18-8-1 1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p67s1 +1 Fr 2-8-18-32-18-8 1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p67s1
  • ELECTRONEGATIVITY The Noble Gases are not assigned electronegativity values. This is due to the fact that they tend not to gain or lose valence electrons because they have complete octets (complete outer principal energy levels). Helium Neon Argon Krypton Xenon Radon
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