Physical Science 100 Lecture Notes Chapters 21-25
Physical Science 100 Lecture Notes Chapters 21-25 Phy S 100
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This 10 page Class Notes was uploaded by Alisa Pierce on Saturday November 7, 2015. The Class Notes belongs to Phy S 100 at Brigham Young University taught by Professor Hirshmann in Summer 2015. Since its upload, it has received 17 views.
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Date Created: 11/07/15
Physical Science 100 Lecture Notes Chapters 2125 How to identify metals on the periodic table: To the left of the "staircase" o Nonmetals are on the right of the "staircase" Properties of metals Metals melt at high temperatures Metals are dense Metals have high boiling temperatures Conduct both heat and electricity Malleable flatten into thin sheets Opaque can't see through even thin sheets Reflective (Metallic Luster) shiny Four key factors that explain metallic bonding: Valence electrons are involved in bonding The electrons in metals have low ionization energies There are few valence electrons in metals compared to the number allowed in the outer shell Electrons want to be in the lowest energy orbital possible Many closed space molecular orbitals give rise to a continuous "band" of energies Metals have few valence electrons compared to number of orbitals Electrons want to be in low energy states "filled" levels Valence electrons in metals The molecular orbitals extend through an entire piece of metal (we say they are delocalized) There are lots of available energy levels for the electrons Very small amounts of energy can move electrons between orbitals The electrons are not tightly attached to any particular atom "sea of electrons" Properties that arise form the "energy band" and mobile electrons High melting temperatures nuclei are surrounded by an electron sea; melting this requires the breaking of strong attractive interactions Electrical conductivity mobile charge carriers are the electrons Thermal conductivity electrons can absorb/give up heat easily; transport it away because of electron mobility Malleability electrons serve as a kind of lubricant, allowing layers of nuclei to slide past one another Properties arising from the Energy Band and mobile electrons Opacity metals can absorb all colors of light also tied to existence of many available energy levels Luster and Reflectivity Alloy A combination of two or more metals to make a new metal Allows form because metals all have: o Few valence electrons o Low ionization energies Are the properties of alloys the same as pure metals? Alloys are not as good conductors of electricity and heat Semiconductors Not quite metal or nonmetal On "Staircase" of periodic table Have some properties of metals o Conduct electricity under certain conditions o Solids with high melting points Widely used in computers and other electronic devices o Particularly, Si and Ge The electrical conductivity of semiconductors and metals with temperature is very different. More electrons, different energies and closer distance to the nucleus...atomic orbitals interact differently Molecular orbitals of semiconductors have one major difference. o A tiny bit of energy is all that you need to move electrons around. The key difference between a metal and a semiconductor is the "band gap" o Band gap = a gap of energy. o No levels in gap Light Emitting Diodes an application of semiconductors Electrical energy (from a battery) can also kick electrons upstairs into conduction band. o When electrons fall back downstairs they emit photons o Photon color emitted by diode is related to energy of band gap If the band gap is really big, there is a lot of energy released Because blue has the highest energy, it has the highest energy. The highest band gap corresponds to the highest energy Ionic Compounds vs Metals Ionic Compounds o Network solids o High melting Temperatures o Brittle solids o Don't conduct heat and electricity in solid o Often colorless and usually transparent in big chunks (white when powdered) Metals o Network solids o High melting temperatures o Malleable o Good conductors of head and electricity in solid o Opaque Metals vs Nonmetals Metals o Large atoms o Few valence electrons o Low ionization energies Nonmetals o Small atoms o Many valence electrons o High ionization energies Energy can be lowered Metals lose valence electrons Nonmetals gain valence electrons Compare & Contrast: energy levels Ionic Compound energy levels o Few levels spaced very far apart Metal energy band o Many clearly spaced levels spread out over many nuclei Salts are generally transparent to light Electron orbitals are localized around individual ions which have few energy levels Examples of Ionic Compounds NaCl o Ions: same charges and similar sizes Al2O3 o Ions: different charges and sizes How the ionic model explains properties of salts (ionic compounds) High melting and boiling temperatures o Strong attractions between positive and negative ions o Attractive forces act over fairly large atomic distances Brittleness o Strong repulsion when ions with like charge come together; material shatters to relieve the stress Conductivity Ionic compounds don't conduct as a solid o Salt alone does not conduct electricity o Salt put in solution does conduct electricity They don't conduct when molten or dissolved o Dissolved salt did conduct electricity, but dissolved sugar did not Use Periodic table to make predictions Metals want to give up electrons Nonmetals want to take on electrons Unreactive noble gases don't form ions The octet rule Atoms will most likely form an ion that has the ns2np6 configuration of the closest noble gas atom. o Metals take on this configuration bylosing electrons o Nonmetals take on this configuration by gaining electrons Families Chlorine and Fluorine will form the same types of compounds since their valence electrons are the same number and same orbital type Ionic charges in compounds NaCl o Na^+1 and Cl^1 KBr o K^+1 and Br^1 MgF2 o Mg^+2 and F^1 Al2O3 o Al^+3 and O^2 Naming convention for salts The metal comes first with its name unchanged The nonmetals come second, with the suffix "ide" appended Predicting formulas for salts Find the number of electrons lost by the metals and gained by the nonmetals If they are equal, the atoms combine one to one If they are NOT equal, use the number lost/gained for the other atom's subscript Carbon Carbon shares electrons in covalent bonds Covalent materials Generally having melting and boiling points in the intermediate to low range Poor conductors of heat and electricity May be solids, liquids, and gases Exist as molecules (as opposed to network and extended solids) Energy configurations As with the other bond types, electrons want to be in LOW ENERGY configurations This generally means having a full shell of valence electrons Covalent bonds form Nonmetals do not want to give up electrons. Instead, they share them o Shared electrons are lower in energy o Valence shells are filled by sharing Both atoms have access to electrons after bonding Single covalent bond The sharing of 2 electrons creates a "single" covalent bond. Double Covalent bond Triple Covalent Bond The octet rule for nonmetals All nonmetal atoms except H and He have 8 valence slots. H and He have 2 valence slots. Molecular bonds reflect this Equal or Unequal sharing Identical atoms o share equally o Nonpolar bond o Ex: H2 Different atoms o May share unequally o Polar bonds o Ex: CH2O Atoms that attract electrons strongly are said to have "high electronegativity" Unequal sharing leads to "polar" molecules Electron density map for H2O slight concentration of electrons at Oxygen end For a molecule to be polar, o Bonds must have unequal sharing: "dipoles" (two poles, + and ) o Bond dipoles must not "cancel" Water is polar CO2 is not Polar bonding = asymmetrical arrangement Nonpolar bonding = symmetrical arrangement How molecules stick together Covalent bonds between atoms are very strong.. Question: which of the fundamental interactions is important for forces between molecules? Answer: the electromagnetic interaction Question: comparing the distances between two bound atoms in a molecule and the distances between molecules, which is greater? Answer: the distance between molecules Question: which force is greatest? Bonding forces between atoms within a molecule Forces within vs forces between Forces within a molecule compromise the types of bonding we have discussed; tend to be large. o Ex: covalent, ionic, metallic o Intramolecular forces Forces between molecules determine the "stickiness" of molecules and their attraction to other, nearby molecules; tend to be small o Helps determine the state of a substance o Intermolecular forces o Dispersion forces, dipole to dipole interactions, hydrogen bonding, covalent bonding Intermolecular forces "Between" different molecules Caused by permanent or temporary charges on molecules Much weaker than covalent bonding interactions Wide range of strengths explains wide range of boiling, melting points of covalent materials When these forces are greater than the forces involved in random motions, the molecules stick together Dispersion forces, dipole to dipole interactions, hydrogen bonding, covalent bonding Dipole forces o Forces exists between dipolar molecules, those with distinct positive and negative ends o So ice melts at a higher temperature than "dry ice" DipoleDipole interactions o Formed in molecules with bonds between atoms of different electronegativities Hydrogen bonding o A type of dipoledipole interaction; unusually strong o Limited to H bound to N, O, or F o Among strongest intermolecular interactions o Happens because H is small and has only 1 electron, and the atom it is bound to is quite electronegative. H is essentially a "bare" proton o A type of dipole bond o Water is the most important example of hydrogen bonding o This also causes ice to be LESS dance than water at 4 degrees C. o The packing of the molecules (due to hydrogen bonding) in the solid form leaves large empty regions or "voids" that account for the low density. Dispersion forces (or van der Waals forces) o A weak attraction between unpolarized molecules arising from the random motion of the electrons o Always present, very weak. o Electron cloud can fluctuate o A "temporary" dipole forms. o "Induces" a new dipole in other nearby atoms or molecules. o Evidence for such forces Noble gases can be liquified Bigger electron cloud, the stronger the forces Higher boiling temperatures of gases The electromagnetic attraction between molecules can be very different Not all molecules with polar bonds (unequal sharing) become polar molecules. Water Properties o Fairly reactive, "universal" solvent Nitrogen Properties o Chemically unreactive o Colorless o Triple bond o No lowlying molecular orbitals o Boiling point 77 K Properties of covalent materials Glucose (sugar) o Crystalline, molecular solid o Sticky Fatty Acids + Glycerol = Acetic Acid Fatty Acids = major component of fats and oils Saturated fatty acids o No kinks Unsaturated fatty acids o One kink (mono) o More than 1 kink (poly) Kinks occur at double bonds Molecules without kinks can snuggle closer together o Result: more and stronger dispersion forces between tails o Stronger hydrogen bonding between CO2H groups on different molecules o Strong forces mean high melting temperatures Transfatty acids: doing away with the kinks Cis vs Trans Double Bonds o Cis double bond gives kink o Trans double bond has no kink Not naturally occurring in foods Good Fats vs Bad Fats Good fats o Kinky unsaturated fats Low melting points Don't clog your arteries Bad fats o Unkinky fats o Saturated & trans fats High melting points Silicate Minerals Molecular Ions IF atoms have too few or too many electrons, they can borrow or give up electrons Covalent bonding within the ion Stronger covalent bonds if number of electrons doesn’t match total nuclear charge, resulting molecule is charged These charged molecules assemble together in crystal lattice like ionic materials The electromagnetic force acts on protons The nuclear strong force acts on Definitions: Nucleon : a proton or a neutron Atomic number : number of protons Mass number : number of nucleons Isotope: an atom with a particular atomic number (protons) but a different mass number o Different isotopes have different numbers of neutrons
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