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82 Cards in this Set
- Front
- Back
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Chemical bonds |
Electrostatic forces holding groups of atoms or ions together |
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Electrostatic forces holding groups of atoms or ions together |
Chemical bonds |
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Categories of chemical bonds |
Covalent bonds- shared electron pair Ionic bonds - transferred electrons Metallic bonds - cations in sea of electrons |
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where do you see bonding capacity in Lewis symbols?
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unpaired dots
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ionic bonds
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electrostatic attraction of closely packed, oppositely charged ions
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electrostatic attraction of closely packed, oppositely charged ions |
ionic bonds
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cation with ionic bonds
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loss of electron(s) by an atom with low ionization energy (metals)
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anion with ionic bonds
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gain of electron(s) by atom with high electron affinity (nonmetals)
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how are ions attracted?
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electrostatically attracted attraction of oppositely charged ions for one another is ionic bond |
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energy to remove an electron
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ionization energy
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ionization energy
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energy to remove an electron
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electron affinity
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energy to add an electron
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energy to add an electron
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electron affinity
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energy involved in ionic bonding process
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combination of ionization energy and electron affinity is still endothermic (process requires energy) but when 2 ions bond, more than enough energy is released, making the overall process exothermic |
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lattice energy |
energy required to completely separate 1 mole of solid ionic compound into gaseous ions |
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energy required to completely separate 1 mole of solid ionic compound into gaseous ions |
lattice energy
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lattice energy equation
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E = k ( (Q1Q2)/r )
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lattice energy relationship to charges of ions
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lattice energy increases as charges increase larger charge = ions more strongly attracted = larger lattice energy |
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lattice energy relationship with distance between ions
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lattice energy increases as distance between ions decreases larger ion = weaker attraction = smaller lattice energy |
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lattice energy relationship to melting and boiling point
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larger lattice energy = higher melting point
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which is more important to lattice energy? ion charges or distance between ions? |
ion charges
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ionic compounds relationship to melting/boiling points
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ionic compounds have high melting/boiling points, bc breaking down crystal should require a lot of energy attractions between ions are strong (larger the lattice energy) |
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general number for melting point of ionic compounds
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MP generally greater than 300 degrees Celcius
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ionic compounds states at room temp
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all ionic solids are solid at room temp
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ionic solids relationship to hardness
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ionic solids relatively hard compared to most molecular solids
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ionic solids relationship to electricity
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ionic solids do not conduct electricity (no movement of electrons) |
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ionic compounds relationship to electricity
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ionic compounds conduct electricity in the liquid state or when dissolved in water
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ionic solids and strength
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ionic solids are brittle, shatter when struck
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covalent bond
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sharing of valence electrons
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sharing of valence electrons
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covalent bond
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bond length
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distance between atoms when energy is at a minimum
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distance between atoms when energy is at a minimum |
bond length
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coordinate covalent bond
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both electrons are donated by one of the atoms
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both electrons are donated by one of the atoms |
coordinate covalent bond
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polar covalent bond
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bonding electrons spend more time near one of the 2 atoms involved
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bonding electrons spend more time near one of the 2 atoms involved |
polar covalent bond
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two alike atoms share bonding atoms equally
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nonpolar covalent bond
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nonpolar covalent bond
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two alike atoms share bonding atoms equally
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how to show polar covalent bonds in Lewis structures
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add partially positive and partially negative symbols
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electronegativity
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a measure of the ability of an atom in a molecule to draw bonding electrons to itself
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a measure of the ability of an atom in a molecule to draw bonding electrons to itself
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electronegativity
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electronegativity trend
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similar to ionization energy and electron affinity
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how to measure polarity of a bond
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absolute value of the difference in electronegativity of 2 bonded atoms (rough measure of polarity)
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if difference between bonded atoms is 0:
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bond is a pure covalent
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if difference between bonded atoms is 0.1 to 0.4:
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bond is nonpolar covalent
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if difference between bonded atoms is 0.5 to 1.9:
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bond is polar covalent
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if difference between bonded atoms is 2+ :
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bond is ionic
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how to know which atom goes in center for lewis structure
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should be the least electronegative atom in center
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allotropes
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different molecular forms of an element ex) carbon: graphite/diamond, oxygen: O2/O3(ozone) |
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delocalized bonding
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a bonding pair of electrons is spread over a number of atoms rather than be localized between two atoms
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a bonding pair of electrons is spread over a number of atoms rather than be localized between two atoms |
delocalized bonding
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bond length differences for single and double bonds
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single bond would have longer bond length than double bond (but different for resonance structures) |
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formal charge
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measure of the # of electrons formally assigned to an atom in a molecule
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measure of the # of electrons formally assigned to an atom in a molecule |
formal charge
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formal charge formula
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# valence electrons - ( # nonbonding pairs - 1/2 # bonding electrons )
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how to choose between two similar formal charges
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negatives should be on more electronegative atoms choose Lewis formulas without like charges on adjacent atoms |
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names for different exceptions to octet rule
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electron deficient free radicals expanded valence shell |
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what to do with odd-electron molecules
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put electrons on more electronegative atoms first
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what are free radicals?
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unpaired electrons (very reactive)
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when can we expand valence shell
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with Z greater than 12 (3rd row) using empty d orbitals |
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when do expanded valence shells occur
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in molecules having strongly electronegative elements (F, O, Cl) when expanded shell decreases formal charge on central atom |
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bond length
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distance between nuclei in a bond
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distance between nuclei in a bond
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bond length
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what does bond length depend on
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identity of the atoms number of bonds b/w them |
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bond order
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number of bonds between 2 atoms # of pairs of electrons in a bond |
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# of bond b/w 2 atoms # of pairs of electrons in a bond |
bond order
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bond order for single, double, and triple bonds
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1 for single bond, 2 for double bond, 3 for triple bond
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bond length and bond order
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as bond order increases, bond gets shorter and stronger
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bond length of single vs double bond
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shorter bond is double bond, longer bond is single bond
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BOND ORDER IN REALITY
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I DON'T UNDERSTAND THE SLIDE ON THE SECOND TO LAST PAGE THAT SAYS BOND LENGTH VS BOND ORDER
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bond energy
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energy needed to break 1 mole of covalent bonds in gas phase breaking bonds consumes energy; forming bonds releases energy |
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energy needed to break 1 mole of covalent bonds in gas phase
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bodn energy
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using bond energies to estimate delta H rxn equation
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delta H rxn = ( sum delta H bond breaking) + ( sum delta H bond forming) (bond forming will end up being negative) |
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what allows metals to lose electrons easily
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they have a low ionization energy
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simplest theory of metallic bonding
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metal atoms release their valence electrons to be shared by all to atoms/ions in the metal an organization of metal cation islands in a sea of electrons electrons delocalized throughout metal structure |
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what does metallic bonding result from
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attraction of cations for the delocalized electrons
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metallic solids relationship to electricity
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conducts electricity well as temp increases, electrical conductivity of metals decrease |
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metallic solids relationship to conducting heat
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do conduct heat well
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metallic solids relationship to reflecting light
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do reflect light
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metallic solids relationship to strength
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malleable and ductile
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metallic solids relationship to melting/boiling points
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generally have high MP and BP all but Hg solid at room temp |
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melting points for metals periodic trend
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generally increase left to right across period
generally decrease down column |
