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Physical Science 100 Lecture Notes Chapters 21-25

by: Alisa Pierce

Physical Science 100 Lecture Notes Chapters 21-25 Phy S 100

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Alisa Pierce

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In-class lecture notes for physical science 100, covering chapters 21-25
Physical Science
Professor Hirshmann
Class Notes
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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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