Chem Flashcards
1
Q
P4O10
A
tetraphosphorous decoxide
2
Q
SO3
A
sulfur trioxide
3
Q
N2O
A
dinitrogen monoxide
4
Q
S2O3
A
disulfur trioxide
5
Q
H2O
A
dihydrogen monoxide
6
Q
HBr(aq)
A
hydrobromic acid
7
Q
HCl(aq)
A
hydrochloric acid
8
Q
HF(aq)
A
hydrofluoric acid
9
Q
H2SO4(aq)
A
sulfuric acid
10
Q
H2SO3(aq)
A
sulfurous acid
11
Q
H2PO4(aq)
A
phosphoric acid
12
Q
H2PO3(aq)
A
phosphorous acid
13
Q
HNO2(aq)
A
nitrous acid
14
Q
What is the empirical formula for P5H10?
A
PH2
15
Q
What is the empirical formula for C6H12O6 (glucose)?
A
CH2O
16
Q
What is the empirical formula for C2H2
A
CH
17
Q
Calculate the percent composition of C in C6H12O6 (glucose).
A
40% carbon
18
Q
Calculate the percent composition of Cl in CCl2F2.
A
58.64% chlorine
19
Q
Chemical formula that is a simplified version of the given molecular formula
A
Empirical formula
20
Q
Chemical formula that is the actual molecular structure of the given compound.
A
Molecular formula
21
Q
A chemical compound always contains the same elements in exactly the
same proportions by weight or mass
A
Law of Definite Proportions
22
Q
Mass cannot be created or destroyed in ordinary chemical and physical
changes.
Ex – the total mass of reactants will equal the total mass of the products
A
Law of Conservation of Mass
23
Q
A pure sample of sodium fluoride (NaF) contains 35 g of sodium. How many grams of fluorine are
present in this sample?
A
29 grams
24
Q
When two elements combine to form two or more compounds, the mass
of one element that combines with a given mass of the other is in the ratio of small whole numbers
Ex – CO and CO2 (both contain the same two elements but are combined differently)
A
Law of Multiple Proportions
25
nickel (III) sulfide
Ni2S3
26
A compound formed by combining a metal and non metal
Ionic compound
27
manganese (II) phosphate
Mn3(PO4)2
28
Silver acetate
AgC2H3O2
29
Magnesium sulfate heptahydrate (Epsom salt)
H14MgO11S
30
A water molecule that binds with a polyatomic ion
Hydrate
31
Potassium carbonate
K2CO3
32
# of atoms bonded to central atom - # of lone pairs on central atom
steric number (SN)
33
# of bonds/# of domains
Bond order
34
Strength/energy required to break electron bond between atoms
Bond energy
35
Bond length is _____ to bond energy
Inversely proportional
36
Electrons will arrange themselves in a way to minimize repulsive forces between them
VSEPR Theory
37
Exceptions to the octet rule
1) Molecules with an odd number of electrons
2) Molecules in which one or more atoms has > or < 8 electrons
38
180° bond angle, steric number 2
Linear
39
120° bond angle, steric number 3
Trigonal planar
40
109.5° bond angle, steric number 4
Tetrahedral
41
90°, 120° bond angle, steric number 5
Trigonal bipyramidal
42
90° bond angle, steric number 6
Octahedral
43
Total electrons - (lone pairs + number of bonds)
Formal charge
44
Formal charge of each atom in this molecule
Left Nitrogen: neutral
Central Nitrogen: +
Oxygen: -
45
Formal charge of each atom in this CO2 molecule
Left oxygen: -
Central carbon: neutral
Right oxygen: +
46
The separation of electrical charge created when atoms of different electronegativities form a
covalent bond, specifically a polar covalent bond.
Dipole
47
Criteria for a molecule to be polar
1) Molecule must possess polar covalent bonds
2) Molecule must have a net-non-zero dipole moment
48
Indication of direction which points towards the more electronegative atom in a bond
Dipole moment
49
T/F all non-polar molecules have a zero dipole moment
True
50
Even though the individual bonds in CO2 are polar, the offset of the dipoles makes CO2 a _______ molecule.
Non-polar
51
The total number of molecular orbitals produced is always equal to the total number of atomic orbitals contributed by the atoms that have combined.
The bonding molecular orbital is lower in energy than the parent orbital, and the antibonding orbital is higher in energy than the parent orbital.
Electrons of the molecule are assigned to orbitals of successively higher energy.
Atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy.
Molecular orbital theory
52
First bond formed, electrons located between two nuclei
Sigma orbital
53
Electrons located above and below bonding axis
Pi orbital
54
Constructive combination of atomic orbitals
Bonding molecular orbital
55
Destructive combination of atomic orbitals
Antibonding molecular orbital
56
1. Valence electrons are located in quantum-mechanical atomic orbitals (s, p, d, f, or hybrid
combinations of these).
2. The overlap of two half-filled orbitals and spin-pairing of the two valence electrons
results in a
3. The geometry of the overlapping orbitals determines the shape of the molecule.
Valence bond theory
57
What are the hybridization types for the following steric numbers
2, 3, 4, 5, 6
SN Hybridization
2 sp
3 sp2
4 sp3
5 sp3d
6 sp3d2
58
A covalent bond results when ________
electrons are shared between a pair of atoms
59
On April 19, 1995, the Murrah Federal Building in Oklahoma City was destroyed, killing 168 people with a simple but powerful bomb made from 4800 lb of ammonium nitrate. What is the formula for ammonium nitrate?
NH4NO3
60
What is the percent composition of fluorine in the compound barium fluoride?
The molar mass of the compound is 175.34 g/mol.
21.67%
61
Answer the following questions about the polyatomic ion nitrate (molar mass 62.00 g/mol)
What is the charge of the nitrate polyatomic ion?
Does the polyatomic ion have multiple resonance structures?
If you had 24 g of nitrate, how many oxygen atoms would you have?
-1
Yes
7.0 x 10^23
62
In which bond does the H atom have the highest electron density?
H — Li
63
In which bond does the Cl atom have the highest electron density?
H — Cl
64
How many shared electron pairs are there in the Lewis structure of C2H4Cl2?
7
65
Determine the molecular geometry of N2O.
Linear
66
The electron group geometry of AsH3 is ________, and its molecular geometry is ________
Tetrahedral, trigonal pyramidal
67
Glucose (C6H12O6) is oxidized by molecular oxygen and forms carbon dioxide and water. How many O2 molecules are needed for each molecule of glucose that is oxidized?
6
68
6Li(s) + N2(g) → 2Li3N(s)
How many moles of lithium are needed to produce 0.45 mol of Li3N when the reaction is carried out in the presence of excess nitrogen?
1.4
69
Water is formed in the redox reaction between hydrogen and oxygen in the balanced chemical equation below. If 50.0 g of oxygen are reacted with excess hydrogen, how much water is produced?
2 H2 (g) + O2 (g) ⟶ 2 H2O (l)
56.3g
70
What mass (in grams) of barium hydroxide is needed to neutralize 40.0 L of 1.60 M hydrochloric acid?
5,483.68 g
71
What happens when a solute dissolves?
It separates into individual particles within a solvent
72
How many g of lithium sulfate are required to make
750 mL solution with a concentration of 3.5 M?
288.58g
73
Soluble salts, strong base/acid
Strong electrolyte
74
Weak acids/bases
Weak electrolyte
75
Dissolves but does not dissociate (typically covalent bonds)
Non-electrolyte
76
Specific heat formula
Q (heat energy) = Mass (m) * heat capacity (c) * ΔT (change in temperature)
Answer is given in J/kgC°
77
First Law of Thermodynamics
Energy can change forms, but can be neither created nor destroyed.
78
HF, AgCl, and PbCl2 are all examples of…
Weak electrolytes
79
Se, As, Ge, Ga, Cu, Ca, K are atoms listed in size from ___________ to ___________.
smallest to largest
80
F>O>C>Be>Li>Na are in order from ________ positive to ________ positive for ionization energy
most, least
81
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2 is the electron configuration for what element?
Ti
82
In general, atomic radii ________ down a group and ________ across a period
increase, decrease
83
The change in energy for the following reaction is referred to as the _________ for oxygen:
O(g) + e^- ==> O^-
electron affinity
84
The central atom in PH3 is surrounded by
three single bonds and one lone pair of electrons
85
Which member of the halogen family has the smallest first ionization energy?
Astatine
86
What is the formal charge on each atom in chloroform, CHCl3? (C = ___, H = ___, each Cl = ___)
0, 0, 0
87
__________________ is a measure of the ability of an atom in a molecule to attract electrons to itself.
Electronegativity
88
What is the molecular shape of ClO3F as redicted by the VSEPR theory?
tetrahedral
89
How many electrons are in a triple bond?
6
90
What types of elements undergo covalent bonding?
two nonmetals
91
What is the molecular shape of NOCl (N is the central atom) as predicted by the VSEPR theory?
bent
92
T/F: AsCl5 will have a(n) expanded octet.
True
93
What is the molecular shape of the BBr3 molecule?
trigonal planar
94
T/F: Valence electrons are the closest to the nucleus.
False
95
T/F: Valence electrons are the most accessible
True
96
T/F: Valence electrons participate in bonding
True
97
T/F: Valence electrons determine the chemical properties of an element
True
98
How many pure atomic orbitals of each type must be hybridized to form a set of sp^(3)d orbitals?
1s, 3p, 1d
99
Is CaCO3 soluble or insoluble?
insoluble
100
Is ZnSO4 soluble or insoluble?
soluble
101
Is Hg(NO3)2 soluble or insoluble?
soluble
102
Is NH4ClO soluble or insoluble?
soluble
103
What precipitate will form when aqueous (NH4)2S reacts with aqueous Cu(NO3)2?
CuS
104
Molecules or ions with the same number of valence electrons are called:
isoelectric
105
Which level is associated with each quantum number
l = 0
l = 1
l = 2
l = 3
s
p
d
f
106
The most stable arrangement of electrons in orbitals of equal energy is the one in which the number of electrons with the same spin is maximized and there is the lowest amount of energy
Hund's Rule
107
Electrons fill lower energy levels first
Aufbau Principle
108
Exceptions to the octet rule:
Boron and Beryllium can be central atoms, but they do not need octets, although they can have octets under certain conditions
Compounds with odd numbers of electrons may have central atoms without octets, but try to get it as close as possible
Central atoms that are nonmetals from period three or higher can have more than eight electrons around them (expanded octets) because they can use their d orbital electrons as extra valence electrons
109
Elements that will form an expanded octet
sulfur, phosphorus, silicon, and chlorine
110
What is the difference between electronic geometry and molecular geometry?
Electron geometry is the geometry of a compound including its lone pair electrons. Molecular geometry does not include lone pair electrons
111
Molecules with an ____ number of electrons are highly reactive.
odd
112
A species that can accept a pair of electrons
Lewis acid
113
A species that has a line pair of electrons that it can donate
Lewis base
114
How do you determine the central atom in the Lewis Structure of a compound?
The central atom is always the least electronegative element
115
The most electronegative atom is:
Fluorine, which is why it will never form a double bond
116
The ability of an atom in a bond to draw shared electrons towards itself
electronegativity
117
The energy required to convert a mole of an ionic solid to its constituent ions in the gas phase
Lattice energy
118
The simplest organic compounds which only contain only hydrogens and carbons
hydrocarbons
119
The electrostatic attraction that holds oppositely charged ions together in a compound
ionic bond
120
Formed between two nonmetals
covalent bond
121
Partial negative and partial positive attraction
hydrogen bond
122
Periodic trend of nuclear charge
increases left to right
123
Periodic trend of effective nuclear charge
increases left to right
124
Periodic trend of atomic radius
decreases left to right, increases top to bottom
125
Periodic trend of ionization energy
increases with effective nuclear charge
126
Periodic trend of electron affinity
increases left to right
127
Good conductor of heat and electricity
metals
128
Poor conductor of heat and electricity
nonmetals
129
Elements with properties that are intermediate between metals and nonmetals
metalloids
130
Why is hydrogen placed in group 1A even though it really belongs in a group by itself
it can either gain or lose an electron depending on the element it bonds with
131
Represented by "n"
Designates the size of the orbital (1, 2, 3, 4, 5)
Principal quantum number
132
Represented by "l"
Describes the shape of the orbital
Angular momentum quantum number
133
Represented by "ml"
Describes the orientation of the orbital in space
NOTE: Can only be one of the possible values from "l"
Magnetic quantum number
134
Represented by "ms"
Describes the spin on an electron
Electron spin quantum number
135
What does the Pauli exclusion principle state?
No two electrons in an atom can have the same four quantum numbers
136
Polyatomic ions that contain one or more oxygen atoms and one other atom of another element as the central atom
oxyanion
137
Produce hydrogen ions and the corresponding oxoanions when dissolved in water
oxyacids
138
1 carbon bonded with hydrogen ions
methane
139
2 carbons bonded with hydrogen ions
ethane
140
3 carbons bonded with hydrogen ions
propane
141
4 carbons bonded with hydrogen ions
butane
142
5 carbons bonded with hydrogen ions
pentane
143
6 carbons bonded with hydrogen ions
hexane
144
7 carbons bonded with hydrogen ions
heptane
145
8 carbons bonded with hydrogen ions
octane
146
NO2-
nitrite
147
NO3-
nitrate
148
(SO3)2-
sulfite
149
(SO4)2-
sulfate
150
(PO3)3-
phosphite
151
(PO4)3-
phosphate
152
(CO3)2-
carbonate
153
OH-
hydroxide
154
ClO-
hypochlorite
155
(ClO2)-
chlorite
156
(ClO3)-
chlorate
157
(ClO4)-
perchlorate
158
(MnO4)-
permanganate
159
(C2H3O2)-
acetate
160
HCO3-
hydrogen carbonate
161
NH4+
ammonium
162
IO3-
iodate
163
A less than full charge on part of a molecule, created by the unequal sharing of electrons
partial charge
