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Chapter 10 Modern atomic theory

Chapter 10 Modern atomic theory. Atoms emit colors when electron enter the excited state and leap to a higher energy level. As they return to the ground state, they release that energy a a photon or light particle. Electrons can act as a particle or a wave.

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Chapter 10 Modern atomic theory

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  1. Chapter 10 Modern atomic theory

  2. Atoms emit colors when electron enter the excited state and leap to a higher energy level. • As they return to the ground state, they release that energy a a photon or light particle.

  3. Electrons can act as a particle or a wave. • This behavior led to the modern atomic theory of wave mechanics. • So, atomic structure went through four main stages to this point: • Plum pudding • Rutherford’s nuclear model • Bohr orbit model • Wave mechanic model

  4. Plum pudding • Lord Kelvin’s model • Protons are scattered in a negative “pudding” of electrons

  5. Nuclear Model • Rutherford’s model • Moves all protons to a nucleus • Electrons are scattered in a cloud around the nucleus

  6. Bohr Orbit • Bohr’s model • Keeps protons in a nucleus • Moves electrons into orbits around the positive charge • Works only for hydrogen

  7. Wave Mechanic Model • Also known as the orbital model • Keeps protons in the nucleus • Moves electrons in specific areas known as orbitals around the nucleus • Orbitals are different shapes • Electrons behave in orbitals according to the Pauli Exclusion Principle

  8. Pauli Exclusion Principle • Only two electrons per orbital • These electrons must spin in opposite directions • Electrons must add one at a time before pairing up

  9. Electron configurations • Electrons are ordered in wave mechanic model • Energy levels • Orbital • Orbital orientations allow for total electron count • Energy levels correspond to the row element is in the periodic table • Orbitals make a pattern in the table • Each orbital can be in different 3D spots, so you can have more than one orbital per energy level

  10. s orbital can hold 2 electrons • Sphere with one orientation • p orbital can hold a total of 6 electrons • Bow tie with three orientations • d orbital can hold a total of 10 electrons • Complex with 5 orientations • f orbital can hold a total of 14 electrons • Complex with 7 orientations

  11. Atomic electron configurations • Superscripts add up to element number • Must follow arrangement of orbitals on the periodic table • Ions have ec as well • Strive to be like noble gases • Either gain electrons to fill shell or lose electrons to fall back one energy level

  12. Hydrogen • 1s1 • Neon • 1s22s23p6 • Write electron configurations for • Ca (20) • I (53) • F-1 (9) • Mg+2 (12)

  13. Valence electrons • Electrons in outermost orbital or shell • Only electrons used in bonding and reactions • Electrons in highest energy level in electron configurations • Determine physical characteristics of atoms and lead to trends in the periodic table

  14. Trends in the periodic table • s,p,d,f orbital arrangement • Metal, non-metal, metalliods • Atomic size • Increase as you move down table • Decrease as you move across table • Ionization energy • Energy needed to remove electrons • Decrease as you move down table • Increase as you move across table • Electronegativity • Desire for an element to pull electrons toward itself • Decrease as you move down table • Increase as you move across table

  15. Chapter 11 bonding

  16. Electronegativity Values • All elements have an electronegativity value • The difference in these values determines the type of intramolecular bond (within the molecule) • Non-polar (0-0.7) • Polar (0.7-1.5) • 1.5-2.0 depends on the element as metal or non-metal • Metal/non-metal are ionic • Non-metal/non-metal are polar • Ionic (> 2.0) • All metal/non-metal bonds are ionic no matter what Δen

  17. Intramolecular Bonding • Non-polar • Equal sharing of electrons • Polar • Unequal sharing of electrons • Creates dipoles • Ionic • No sharing • Intramolecular bonding influences intermolecular bonding (between molecules)

  18. Intermolecular bonding • London dispersion • Non-polar • Dipole-dipole • Polar • Hydrogen bonding • Polar bonding between molecules that have H and O, N, and F • Ionic • Large networks created • Alternating positive and negative charges • Metallic • Sea of electrons • Conducts heat and electricity

  19. Representing bonding • Lewis structures are visual ways to represent bonds • Use only valence electrons • Shared and unshared pairs of electrons • Dashes represent one shared pair of electrons

  20. Molecular geometry • Bonding arrangement leads to the geometry of the molecules • Lewis structures can help predict molecular geometry • Three electron pair arrangements • Linear • Triangular • Tetrahedral • Use the chart to see molecular geometry

  21. Chapter 7 reactions

  22. Classes of Reactions • Precipitate • A reaction where a solid forms from the mixing of two clear liquids • Oxidation reduction • The movement of electrons • Oxidation is loss reduction is gain (OIL RIG) • Combustion reactions • Hydrocarbon reacts with O2 to make CO2 and H2O • Acid base reactions • HA plus BOH react to make a salt and water

  23. Types of reactions • Synthesis • A + B → AB • Decomposition • AB → A + B • Single displacement • AB + C → AC + B • Double displacement • AB + CD → AD + CB

  24. Predicting reaction products • Using the general equations on the previous slide, you must be able to predict the products of a reaction • Remember to make neutral compounds first. • For example: • C + O2→ • C + O2→ CO2 • HCl + NaOH→ • HCl + NaOH→ H2O + NaCl • K + CuSO4 → • K + CuSO4 → K2SO4 + Cu

  25. Precipitate reactions • Create a solid • Use the flow chart to identify products that should be solids

  26. NaS (aq) + CaNO3 (aq) → Na+1 (aq) + S-2 (aq) + Ca+2 (aq) + NO3-1 → NaS (aq) + CaNO3 (aq) → CaS (s) + NaNO3 (aq) NET: S-2 (aq) + Ca+2 (aq) → CaS (s)

  27. Oxidation Reduction • Electron transfer 2 Fe + 3 CuSO4→ Fe2(SO4)3+ 3 Cu • Fe begins at Fe0 and ends up as Fe+3 • Lost electrons • Cu begins as Cu+2 and ends up as Cu0 • Gains electrons

  28. Half Reactions 2 Fe + 3 CuSO4→ Fe2(SO4)3+ 3 Cu Fe0→ Fe+3 + 3e- Cu+2 + 2 e- → Cu0 2Fe0→2 Fe+3 + 6e- 3Cu+2 + 6e- →3 Cu0 2Fe0 + 3 Cu+2→ 2Fe+3 + 3 Cu0

  29. Solution reactions • Dissociation • Process where an ionic compound separates in water into ions • Dissociated ions are known as electrolytes • Electrolyte solutions can conduct electricity • Acids dissociate into H+1 ions and anions • HF → H+1 + F-1 • Bases dissociate into cations and OH-1 • NaOH → Na+1 + OH -1

  30. Chapter 12 gases

  31. Kinetic molecular theory of gases • Gases are made of tiny particles • Particles are very small and far apart so their size is zero • Particles are in constant, random motion and collide to create pressure • Particles do not repel or attract each other • Kinetic energy of the particle is directly related to the temperature

  32. Gas Laws

  33. Dalton’s law of partial pressure • Pt = P1 + P2 + P3……. • Key is to find the pressure with PV=nRT • Total pressure is the sum of the individual pressures

  34. Gas Stoichiometry • All gas laws and gas stoichiometry use Kelvin scale for temperature • 1 mole of any gas (the volume) is equal to 22.4 L • STP is standard temperature and pressure • 1 atm and 273 K • Problems must always focus on finding moles first, then volume • If at STP, use 22.4 L • If not at STP, use PV=nRT once you have moles (n)

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