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