Chemistry Exam 2 Test Answers – Flashcards
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Unlock answersIonic Compound Characteristics |
not malleable high melting/boiling point don't conduct electricity except when dissolved (free ions) |
Covalent Compound Characteristics |
stronger than ionic don't conduct electricity low melting/boiling point not solid at room temp distinct molecules |
Atomic Radius Trends |
increases going down a group decreases from L--> R |
Atomic Radius Exceptions |
Ga < Al Ga has higher Zeff |
Ionization Energy Trends |
increases left to right decreases going down a group |
Ionization Energy Exceptions |
B<Be and Al<Mg N>O, P>S, As>Se
|
How to find group number through IE |
count number of IEs before HUGE JUMP in value that equals valence e- that equals group number |
Electron Affinity |
amount of energy to add e- positive or negative |
when gain e-, EA is |
negative |
when lose e-, EA is |
positive |
Electron Affinity Trends |
Halogens: -(favorable)EA Noble gases: +(unfavorable)EA C has -EA N has +EA
|
trend in gaining or losing e- |
McNa |
oxide |
binary compound with oxygen |
metal oxide in water |
basic |
nonmetal oxide |
acidic in water |
Trends In Ions |
Anions: radius increase, IE decrease Cations: radius decrease, IE increase |
Isoelectronic series |
series of atoms and ions with exact same number total e- can only have 1 neutral atom |
Add 1st e- = -EA |
exothermic |
Add 2nd e- = +EA |
endothermic |
Ionic bonds |
transfer e- metals and nonmetals |
covalent bonds |
share e- nonmetals |
metallic bonds |
pooling of e- metals |
duet rule |
H, He, Li, Be |
electrostatic attractions |
cation and anion |
ionic compounds must be in terms of charge |
neutral |
electrostatic energy equation |
Electrostatic energy = (charge cation)(charge anion) = ∆H˚lattice (cation radius)+(anion radius) |
lattic energy trends |
ion size increases, lattice energy decreases ion charge increase, lattice energy increases |
charge is a bigger influence than radius |
diatomic molecules |
H2, N2, O2, Halogens |
Hess's Law |
change in energy depends on start and end, not path |
How do you find the highest lattice energy when comparing compounds? |
highest charge smallest radius |
bond energy |
energy needed to break a covalent bond |
bonds strength and bond energy |
strong bond = high BE weak bond = low BE |
breaking bonds gives off energy |
A ; B(g) ; A(g) + B(g)∆H˚bond breaking ;= BEAB;;;;;;;;;; always; 0 |
making bonds takes energy |
A(g) + B(g) ; A ; B(g)∆H˚bond formation ;= -BEAB;;;;;;; always; 0 |
bond length |
how close we can get the nuclei sum of the covalent radii |
bond order |
what type of bond we have |
bond length and bond strength |
A;B
|
energies in molecules |
bond: leads to energy of reactions (PE) translational: moving around in space (KE) rotational: spin (KE) vibrational: bond wiggling (KE) each molecule has all at once, all the KEs ;move all KE at same temp, bond energy leads to energy changes in reaction |
melting points and boiling points break... |
lattice interactions |
why do ionic compounds have high melting and boiling points? |
the melting and boiling points do not affect covalent bonds |
what is an exception to covalent bonds and low melting/boiling point? |
diamond is a very hard network of carbons hard substance high melting point covalently bonded |
change in energy equation |
∆E = Ein- Eout |
electronegativity |
measurement of an atom's tendency to pull on e- in a bond |
EN trends |
increases towards F EN for H=P B<H<C F is the highest at 4.0
|
polar covalent |
∆EN > 0 bonds between two different nonmetals |
nonpolar covalent |
∆EN = 0 bond between two atoms of the same nonmetal H-P bond nonpolar |
ionic character goes up as ΔEN increases |
H-P H-C H-N H-O H-F ---------------------------→ -------------------------→
∆EN ioniccharacter |
greater EN value |
greater pull on e- |
greater the difference between the ENs |
greater the ionic character |
EN trend |
increases L-->R decreases down a group |
Metalllic Bond Characteristics |
no set # of atoms in a metal sample metals deform instead of shatter--malleable conducts electricity and heat (solid and liquid state) most are solids moderate-high melting point much higher boiling point |
ammonium |
NH4+ |
hydronium |
H3O+ |
acetate |
CH3COO- (C2H3O2-) |
cyanide |
CN- |
hydroxide |
OH- |
hypochlorite |
ClO- |
chlorite |
ClO2- |
chlorate |
ClO3- |
perchlorate |
ClO4- |
nitrite |
NO2- |
nitrate |
NO3- |
permanganate |
MnO4- |
carbonate |
CO32- |
hydrogen carbonate (bicarbonate) |
HCO3- |
chromate |
CrO42- |
dichromate |
Cr2O72- |
peroxide |
O22- |
hydrogen phosphate |
HPO42- |
dihydrogen phosphate |
H2PO4- |
sulfite |
SO32- |
sulfate |
SO42- |
hydrogen sulfate (bisulfate) |
HSO4- |
the ion with the most O atoms |
per ate |
the ion with one fewer O atoms |
-ate |
the ion with two fewer O atoms |
-ite |
the ion with the least (three fewer) O atoms |
hypo ite |
1 |
mono- |
2 |
di- |
3 |
tri- |
4 |
tetra- |
5 |
penta- |
6 |
hexa- |
7 |
hepta- |
8 |
octa- |
9 |
nona- |
10 |
deca- |
methane |
CH4 |
ethane |
C2H6 |
how do you find number of H atoms for alkanes? |
double number C atoms and add 2 |
propane |
C3H8 |
butane |
C4H10 |
pentane |
C5H12 |
hexane |
C6H14 |
heptane |
octane |
C8H18 |
nonane |
C9H20 |
decane |
C10H22 |
how do you decide which lewis structure contributes more? (think formal charge) |
want lowest magnitude of formal charge most negative formal charge on most electronegative atom |
there is no double bond with B |
empirical formula |
relative numbers of atoms with smallest ratio possible |
molecular formula |
shows actual number of each type of atom |
structural formula |
shows how atoms are connected |
molecular weight |
sum of atomic masses of every atom in one molecule g/mol amu/molecule |
molecular mass |
mass of molecular formula |
empirical mass |
mass of empirical formula |
in ionic compound, molecular mass=empirical mass |
Ionic Nomenclature--Main Group Metals |
Give name of metal (cation) Name of nonmetal with -ide suffix
Ex: LiBr --Lithium Bromide |
Ionic Nomenclature--Transition Metals
|
Give name of metal (cation) Add charge in parentheses and in roman numerals Add name of nonmetal with -ide suffix (anion)
Ex: FeCl3--Iron (III) chloride |
Polyatomic Ion Nomenclature |
Use name No suffixes
Ex: NaNO3--Sodium nitrate |
oxyanions |
anions containing oxygen |
MgSO4 • 7H2O |
the molecule is hydrated |
Covalent Nomenclature--Binary Compounds |
Name of 1st element--lower group number, or higher period number (H is never first) Name 2nd element with -ide suffix (O with hallogen, name hallogen 1st) Indicate number of atoms with prefix (never used mono- with 1st element)
Ex: P2Cl5-- diphosphorous pentachloride |
Binary Compounds--Common Names |
H2O: water NH3: ammonia CnH2n+2: alkane |
How to draw Lewis structures |
Place elements relative to each other (pick central atom--lowest group # or highest period #, noble gas) Count valence e- Add single bonds between central atoms and terminal atoms Calculate bonding pairs Fill out octets with lone pairs Count valence e- used (if equals # valence e-, then we're done) ; |
number bonding pairs equation |
[8(# atoms)-# valence e-]/2 |
resonance structure |
actual structure = average of all resonance structures |
resonance |
involves placement of double and triple bonds |
bond order in resonance |
(bonding pairs)/(bonds) ; ; Ex:O3=3/2 |
delocalized e- |
e- not stuck in between two atoms, free to roam across molecule |
formal charge equation |
count bonding pairs as 1 lone pairs separately decide charge if more or less than number valence e- |
Exceptions to Octet Rule |
not enough e- (Be or B) odd number e- (at the end, take away from least electronegative element--NO2 is weird) too many e- (expand octets--in row 3 or lower, extra e- on central atom)
|
puttting (+) formal charge on something very electronegative is bad |
table showing usual bonding and lone pairs for C, N, O, Halogens, and H |
|
molecular geometry |
arrangement of e- groups around central atom |
e- group |
bond or lone pair each count as one
|
molecular shape |
dependent on atoms (terminal) around central atoms where the atoms can be based on number of e- groups |
VSEPR |
valence shell e- pair repulsion e- groups arrange themselves around atoms to maximize distance between them |
angles in geometry |
linear: 180° trigonal planar: 120° trigonal bipyramid: 90°, 120°, 180° (t-shaped: 90°, 180° linear: 180°) octahedral: 90°, 180°
|
geometry=shape |
when all e- are bonding groups |
A: central atom X: terminal atom E: lone pair |
only count lone pairs on central atom
|
with two central atoms, talk about shape/geometry separately for central atoms |
to make polar molecule |
break symmetry using lone pairs or changing the identity of terminal atoms |
isomer |
two different molecules with the same formular |
P less electronegative than N, P=H |
need polar bonds before you can have polar molecule |
to determine number of e- groups |
(bonds + lone pairs) around central atom |
hybrid orbital formation |
start with an s orbital need as many hybrid orbitals as e- groups end with as many hybrid orbitals as starting atoms mix in p orbitals to get proper number add d orbitals when necessary form σ bonds or hold lone pairs π bonds with unhybridized p orbitals |
σ bond |
head-to-head overlap of hybrid orbitals first bond between any 2 atoms |
what dictates hybridization? |
geometry and shape |
π bonds |
side-to-side overlap of unhybridized p orbitals any multiple bonds cannot be rotate each p orbital can only form 1 π bond, not 2
|
how many σ bonds and π bonds in a single bond? double bond? triple bond? |
1 σ, 0 π 1 σ, 1 π 1 σ, 2 π |