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Nicholas > Chem > Flashcards

Chem Flashcards

1
Q

P4O10

A

tetraphosphorous decoxide

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2
Q

SO3

A

sulfur trioxide

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3
Q

N2O

A

dinitrogen monoxide

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4
Q

S2O3

A

disulfur trioxide

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5
Q

H2O

A

dihydrogen monoxide

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6
Q

HBr(aq)

A

hydrobromic acid

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7
Q

HCl(aq)

A

hydrochloric acid

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8
Q

HF(aq)

A

hydrofluoric acid

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9
Q

H2SO4(aq)

A

sulfuric acid

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10
Q

H2SO3(aq)

A

sulfurous acid

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11
Q

H2PO4(aq)

A

phosphoric acid

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12
Q

H2PO3(aq)

A

phosphorous acid

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13
Q

HNO2(aq)

A

nitrous acid

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14
Q

What is the empirical formula for P5H10?

A

PH2

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15
Q

What is the empirical formula for C6H12O6 (glucose)?

A

CH2O

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16
Q

What is the empirical formula for C2H2

A

CH

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17
Q

Calculate the percent composition of C in C6H12O6 (glucose).

A

40% carbon

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18
Q

Calculate the percent composition of Cl in CCl2F2.

A

58.64% chlorine

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19
Q

Chemical formula that is a simplified version of the given molecular formula

A

Empirical formula

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20
Q

Chemical formula that is the actual molecular structure of the given compound.

A

Molecular formula

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21
Q

A chemical compound always contains the same elements in exactly the
same proportions by weight or mass

A

Law of Definite Proportions

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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

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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

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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

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25
Q

nickel (III) sulfide

A

Ni2S3

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26
Q

A compound formed by combining a metal and non metal

A

Ionic compound

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27
Q

manganese (II) phosphate

A

Mn3(PO4)2

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28
Q

Silver acetate

A

AgC2H3O2

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29
Q

Magnesium sulfate heptahydrate (Epsom salt)

A

H14MgO11S

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30
Q

A water molecule that binds with a polyatomic ion

A

Hydrate

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31
Q

Potassium carbonate

A

K2CO3

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32
Q

of atoms bonded to central atom - # of lone pairs on central atom

A

steric number (SN)

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33
Q

of bonds/# of domains

A

Bond order

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34
Q

Strength/energy required to break electron bond between atoms

A

Bond energy

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35
Q

Bond length is _____ to bond energy

A

Inversely proportional

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36
Q

Electrons will arrange themselves in a way to minimize repulsive forces between them

A

VSEPR Theory

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37
Q

Exceptions to the octet rule

A

1) Molecules with an odd number of electrons
2) Molecules in which one or more atoms has > or < 8 electrons

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38
Q

180° bond angle, steric number 2

A

Linear

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39
Q

120° bond angle, steric number 3

A

Trigonal planar

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40
Q

109.5° bond angle, steric number 4

A

Tetrahedral

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41
Q

90°, 120° bond angle, steric number 5

A

Trigonal bipyramidal

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42
Q

90° bond angle, steric number 6

A

Octahedral

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43
Q

Total electrons - (lone pairs + number of bonds)

A

Formal charge

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44
Q

Formal charge of each atom in this molecule

A

Left Nitrogen: neutral
Central Nitrogen: +
Oxygen: -

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45
Q

Formal charge of each atom in this CO2 molecule

A

Left oxygen: -
Central carbon: neutral
Right oxygen: +

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46
Q

The separation of electrical charge created when atoms of different electronegativities form a
covalent bond, specifically a polar covalent bond.

A

Dipole

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47
Q

Criteria for a molecule to be polar

A

1) Molecule must possess polar covalent bonds
2) Molecule must have a net-non-zero dipole moment

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48
Q

Indication of direction which points towards the more electronegative atom in a bond

A

Dipole moment

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49
Q

T/F all non-polar molecules have a zero dipole moment

A

True

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50
Q

Even though the individual bonds in CO2 are polar, the offset of the dipoles makes CO2 a _______ molecule.

A

Non-polar

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51
Q

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.

A

Molecular orbital theory

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52
Q

First bond formed, electrons located between two nuclei

A

Sigma orbital

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53
Q

Electrons located above and below bonding axis

A

Pi orbital

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54
Q

Constructive combination of atomic orbitals

A

Bonding molecular orbital

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55
Q

Destructive combination of atomic orbitals

A

Antibonding molecular orbital

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56
Q
  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.
A

Valence bond theory

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57
Q

What are the hybridization types for the following steric numbers
2, 3, 4, 5, 6

A

SN Hybridization
2 sp
3 sp2
4 sp3
5 sp3d
6 sp3d2

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58
Q

A covalent bond results when ________

A

electrons are shared between a pair of atoms

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59
Q

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?

A

NH4NO3

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60
Q

What is the percent composition of fluorine in the compound barium fluoride?

The molar mass of the compound is 175.34 g/mol.

A

21.67%

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61
Q

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?

A

-1

Yes

7.0 x 10^23

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62
Q

In which bond does the H atom have the highest electron density?

A

H — Li

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63
Q

In which bond does the Cl atom have the highest electron density?

A

H — Cl

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64
Q

How many shared electron pairs are there in the Lewis structure of C2H4Cl2?

A

7

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65
Q

Determine the molecular geometry of N2O.

A

Linear

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66
Q

The electron group geometry of AsH3 is ________, and its molecular geometry is ________

A

Tetrahedral, trigonal pyramidal

67
Q

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?

A

6

68
Q

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?

A

1.4

69
Q

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)

A

56.3g

70
Q

What mass (in grams) of barium hydroxide is needed to neutralize 40.0 L of 1.60 M hydrochloric acid?

A

5,483.68 g

71
Q

What happens when a solute dissolves?

A

It separates into individual particles within a solvent

72
Q

How many g of lithium sulfate are required to make
750 mL solution with a concentration of 3.5 M?

A

288.58g

73
Q

Soluble salts, strong base/acid

A

Strong electrolyte

74
Q

Weak acids/bases

A

Weak electrolyte

75
Q

Dissolves but does not dissociate (typically covalent bonds)

A

Non-electrolyte

76
Q

Specific heat formula

A

Q (heat energy) = Mass (m) * heat capacity (c) * ΔT (change in temperature)

Answer is given in J/kgC°

77
Q

First Law of Thermodynamics

A

Energy can change forms, but can be neither created nor destroyed.

78
Q

HF, AgCl, and PbCl2 are all examples of…

A

Weak electrolytes

79
Q

Se, As, Ge, Ga, Cu, Ca, K are atoms listed in size from ___________ to ___________.

A

smallest to largest

80
Q

F>O>C>Be>Li>Na are in order from ________ positive to ________ positive for ionization energy

A

most, least

81
Q

1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2 is the electron configuration for what element?

A

Ti

82
Q

In general, atomic radii ________ down a group and ________ across a period

A

increase, decrease

83
Q

The change in energy for the following reaction is referred to as the _________ for oxygen:
O(g) + e^- ==> O^-

A

electron affinity

84
Q

The central atom in PH3 is surrounded by

A

three single bonds and one lone pair of electrons

85
Q

Which member of the halogen family has the smallest first ionization energy?

A

Astatine

86
Q

What is the formal charge on each atom in chloroform, CHCl3? (C = ___, H = ___, each Cl = ___)

A

0, 0, 0

87
Q

__________________ is a measure of the ability of an atom in a molecule to attract electrons to itself.

A

Electronegativity

88
Q

What is the molecular shape of ClO3F as redicted by the VSEPR theory?

A

tetrahedral

89
Q

How many electrons are in a triple bond?

A

6

90
Q

What types of elements undergo covalent bonding?

A

two nonmetals

91
Q

What is the molecular shape of NOCl (N is the central atom) as predicted by the VSEPR theory?

A

bent

92
Q

T/F: AsCl5 will have a(n) expanded octet.

A

True

93
Q

What is the molecular shape of the BBr3 molecule?

A

trigonal planar

94
Q

T/F: Valence electrons are the closest to the nucleus.

A

False

95
Q

T/F: Valence electrons are the most accessible

A

True

96
Q

T/F: Valence electrons participate in bonding

A

True

97
Q

T/F: Valence electrons determine the chemical properties of an element

A

True

98
Q

How many pure atomic orbitals of each type must be hybridized to form a set of sp^(3)d orbitals?

A

1s, 3p, 1d

99
Q

Is CaCO3 soluble or insoluble?

A

insoluble

100
Q

Is ZnSO4 soluble or insoluble?

A

soluble

101
Q

Is Hg(NO3)2 soluble or insoluble?

A

soluble

102
Q

Is NH4ClO soluble or insoluble?

A

soluble

103
Q

What precipitate will form when aqueous (NH4)2S reacts with aqueous Cu(NO3)2?

A

CuS

104
Q

Molecules or ions with the same number of valence electrons are called:

A

isoelectric

105
Q

Which level is associated with each quantum number
l = 0
l = 1
l = 2
l = 3

A

s
p
d
f

106
Q

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

A

Hund’s Rule

107
Q

Electrons fill lower energy levels first

A

Aufbau Principle

108
Q

Exceptions to the octet rule:

A

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
Q

Elements that will form an expanded octet

A

sulfur, phosphorus, silicon, and chlorine

110
Q

What is the difference between electronic geometry and molecular geometry?

A

Electron geometry is the geometry of a compound including its lone pair electrons. Molecular geometry does not include lone pair electrons

111
Q

Molecules with an ____ number of electrons are highly reactive.

A

odd

112
Q

A species that can accept a pair of electrons

A

Lewis acid

113
Q

A species that has a line pair of electrons that it can donate

A

Lewis base

114
Q

How do you determine the central atom in the Lewis Structure of a compound?

A

The central atom is always the least electronegative element

115
Q

The most electronegative atom is:

A

Fluorine, which is why it will never form a double bond

116
Q

The ability of an atom in a bond to draw shared electrons towards itself

A

electronegativity

117
Q

The energy required to convert a mole of an ionic solid to its constituent ions in the gas phase

A

Lattice energy

118
Q

The simplest organic compounds which only contain only hydrogens and carbons

A

hydrocarbons

119
Q

The electrostatic attraction that holds oppositely charged ions together in a compound

A

ionic bond

120
Q

Formed between two nonmetals

A

covalent bond

121
Q

Partial negative and partial positive attraction

A

hydrogen bond

122
Q

Periodic trend of nuclear charge

A

increases left to right

123
Q

Periodic trend of effective nuclear charge

A

increases left to right

124
Q

Periodic trend of atomic radius

A

decreases left to right, increases top to bottom

125
Q

Periodic trend of ionization energy

A

increases with effective nuclear charge

126
Q

Periodic trend of electron affinity

A

increases left to right

127
Q

Good conductor of heat and electricity

A

metals

128
Q

Poor conductor of heat and electricity

A

nonmetals

129
Q

Elements with properties that are intermediate between metals and nonmetals

A

metalloids

130
Q

Why is hydrogen placed in group 1A even though it really belongs in a group by itself

A

it can either gain or lose an electron depending on the element it bonds with

131
Q

Represented by “n”
Designates the size of the orbital (1, 2, 3, 4, 5)

A

Principal quantum number

132
Q

Represented by “l”
Describes the shape of the orbital

A

Angular momentum quantum number

133
Q

Represented by “ml”
Describes the orientation of the orbital in space
NOTE: Can only be one of the possible values from “l”

A

Magnetic quantum number

134
Q

Represented by “ms”
Describes the spin on an electron

A

Electron spin quantum number

135
Q

What does the Pauli exclusion principle state?

A

No two electrons in an atom can have the same four quantum numbers

136
Q

Polyatomic ions that contain one or more oxygen atoms and one other atom of another element as the central atom

A

oxyanion

137
Q

Produce hydrogen ions and the corresponding oxoanions when dissolved in water

A

oxyacids

138
Q

1 carbon bonded with hydrogen ions

A

methane

139
Q

2 carbons bonded with hydrogen ions

A

ethane

140
Q

3 carbons bonded with hydrogen ions

A

propane

141
Q

4 carbons bonded with hydrogen ions

A

butane

142
Q

5 carbons bonded with hydrogen ions

A

pentane

143
Q

6 carbons bonded with hydrogen ions

A

hexane

144
Q

7 carbons bonded with hydrogen ions

A

heptane

145
Q

8 carbons bonded with hydrogen ions

A

octane

146
Q

NO2-

A

nitrite

147
Q

NO3-

A

nitrate

148
Q

(SO3)2-

A

sulfite

149
Q

(SO4)2-

A

sulfate

150
Q

(PO3)3-

A

phosphite

151
Q

(PO4)3-

A

phosphate

152
Q

(CO3)2-

A

carbonate

153
Q

OH-

A

hydroxide

154
Q

ClO-

A

hypochlorite

155
Q

(ClO2)-

A

chlorite

156
Q

(ClO3)-

A

chlorate

157
Q

(ClO4)-

A

perchlorate

158
Q

(MnO4)-

A

permanganate

159
Q

(C2H3O2)-

A

acetate

160
Q

HCO3-

A

hydrogen carbonate

161
Q

NH4+

A

ammonium

162
Q

IO3-

A

iodate

163
Q

A less than full charge on part of a molecule, created by the unequal sharing of electrons

A

partial charge

Nicholas (6 decks)

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