Chemistry Exam 2 Test Answers – Flashcards
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Unlock answers| Ionic 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 |
