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CHM 1040 > final > Flashcards

final Flashcards

1
Q

commonly used prefixes in metric system

A

prefix symbol meaning power of 10
mega M 1,000,000 10^6
kilo k 1000 10^3
deci d 0.1 10^-1
centi c 0.01 10^-2
milli m 0.001 10^-3
micro μ 0.000001 10^-6
nano n 0.000000001 10^-9

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2
Q

rules for sig figs

A

1) all nonzero integers ALWAYS count for significance ex: 3456 has 4 sig figs
2) zeros (3 classes of zeros)
a) leading zeros:NEVER count as sig figs
ex: 0.048 has 2 sig figs
b) captive zeros: ALWAYS count as sig figs ex: 16.07 has 4 sig figs
c) trailing zeros: only significant when # HAS A DECIMAL POINT
ex: 9.300 has 4 sig figs; 0.004020 has 4 sig figs; 150 has 2 sig figs

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3
Q

sig figs for x and ÷

A

the # of sig figs for answer is LEAST amount of sig figs u have in the problem

ex: 1.342 x 5.5 = 7.4
* when solving many calculations for 1 problem, don’t convert sig figs until LAST STEP

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4
Q

sig figs for + and -

A

only count decimal places

ex: 23.445 + 7.83 = 31.275 = 31.28
ex: 101 + 1.0 = 102
- always solve from left to right

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5
Q

density

A

density =mass/volume or D=m/V

  • mass of substance per unit volume of substance
  • common units: g/mL or g/cm^3
  • when mass increases, density increases (directly proportional)
  • when mass is constant and volume decreases, density increases (inversely proportional)
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6
Q

physical properties

A
  • characteristic that are directly observable and unique to a substance
  • ex: odor, volume, color, state (s,l,g,p), density, boiling pt, melting pt
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7
Q

chemical properties

A
  • a substance’s ability to make new substances
  • characteristics that determine how the composition of matter changes as a result of contact w/ other matter/influence of e
  • characteristics that describe behavior of matter
  • ex: flammability, rusting of steel, toxicity, enthalpy, chemical stability, reactivity, digestion of food
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8
Q

ethyl alcohol at 78ºC boiling point is a __ property

A

physical

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9
Q

hardness of a rock is a __ property

A

physical

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10
Q

sugar fermenting to form ethyl alcohol is a __ property

A

chemical

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11
Q

physical change

A
  • change in 1 or more properties of a substance and not in its chemical composition
  • ex: boiling pt or freezing water
  • 3 states of water: in all phases, water mols are still intact; motion of mols and distance between them change
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12
Q

chemical change

A
  • given substance becomes a new substance w/ diff properties and diff composition
  • ex: Bunsen burner, methane reacts w/O2 to make CO2 and H2O, baking a cake, bleaching teeth (rxn happening), digesting food
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13
Q

crushing salt is a __ change

A

physical

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14
Q

burning wood is a __ change

A

chemical

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15
Q

dissolving sugar in water is a __ change

A

physical

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16
Q

melting a popsicle is a __ change

A

physical

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17
Q

protons and neutrons have the same __

A

mass

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18
Q

why do diff atoms have diff chemical properties?

A

the chemistry of an atom arises from its e-

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19
Q

isotopes

A

diff # of neutrons

  • atoms w/ the same # of protons but diff # of neutrons
  • show almost identical chem. properties
  • chemistry of an atom is due to its e-
  • in nature, elements are usually found as a mixture of isotopes
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20
Q

isotope symbol

A 
A
   X
Z
X= the chemical symbol of element 
A= mass # (# of protons + # of neutrons)
Z= atomic # (# of protons)
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21
Q

groups/families

A

elements in same VERTICAL columns and have similar chemical properties

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22
Q

periods

A

HORIZONTAL rows of elements

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23
Q

metals

A

LEFT of staircase

-most elements are these

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24
Q

non-metals

A

RIGHT of staircase

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25
metalloids
ON the staircase | -have some metallic and some non-metallic properties
26
what are the metalloids?
``` Boring Silly Germs Are Ants Telling Politics Boron (B) Silicon (Si) Germanium (Ge) Arsenic (As) Antimony (Sb) Tellurium (Te) Polonium (Po) ```
27
what are the four physical properties of metals?
1) efficient conduction of heat and electricity 2) malleability (aluminum foil) - they can be hammered into thin sheets 3) ductility - they can be pulled into wires - can be molded 4) lustrous appearance (shiny)
28
physical properties of non-metals
- lack properties of metals - exhibit more variation in properties - can be g, l, s @ room temp
29
physical properties of metalloids
exhibit a mixture of metallic and non-metallic properties
30
7 diatomic molecules
I Bring Clay For Our New Home - Iodine (I2): lustrous, dark purple solid - Bromine (Br2): reddish-brown liquid - Chlorine (Cl2): pale green gas - Fluorine (F2): pale yellow gas - Oxygen (O2): pale blue gas - Nitrogen (N2): colorless gas - Hydrogen (H2): colorless gas
31
ions
- atom's not neutral - atom w a charge - elements become ions - imbalance of e- - atoms can form ions by gaining/losing e-
32
cations
metals tend to LOSE 1 or more e- to form (+) ions
33
anions
non-metals tend to GAIN 1 or more e- to form (-) ions | -name changes to end in -ide
34
ion charges and the periodic table chart
Group or Family Charge Alkali Metals (1A) 1+ Alkaline Earth Metals (2A) 2+ Halogens (7A) 1- Noble Gasses (8A) 0 ``` *group 1: +1 group 2: +2 group 3: +3 group 6: -2 group 7: -1 group 8: 0 ```
35
what are the 3 steps of the scientific method?
1) make an observation 2) form a hypothesis 3) perform experiment
36
theory vs law
- theory: answers "why?", leads to more questions | - the law: "this is what happened", doesn't lead to more questions
37
which of the 3 subatomic particles are the smallest?
electrons
38
__ contribute to the mass of an atom
protons
39
__ contribute to the size of an atom
electrons
40
__ dictates chemistry of an atom
electrons
41
naming compounds: binary compounds
- composed of 2 elements - divided into broad classes - compounds that contain a metal & non-metal (ionic) - compounds that contain 2 non-metals (covalent)
42
nomenclature
naming
43
naming compounds: binary ionic compounds
- contains (+) cations and (-) anions - type 1 - type 2
44
binary ionic compounds: type 1
-metal and non-metal -compounds -fixed charges -metal present forms only 1 cation consists of: -alkali metals -alkaline earth metals -Al3+ -Ga3+ -In3+ -Zn2+ -Ag+
45
charges of transition metals
- Al3+ - Ga3+ - In3+ - Zn2+ - Ag+
46
binary ionic compounds: type 2
- metal and non-metal - compounds - no fixed charge - need roman numerals (indicates charge of metal cation) - charge of metal ion must be specified - metal present can form 2 or more cations w diff charges - consists of: transitional metals
47
naming type 1 binary ionic compounds
- cation is always named first (element name) | - anion named second (end in -ide)
48
examples of type 1 binary ionic compounds
``` NaCl: sodium chloride CaS: calcium sulfide Kl: potassium iodide SrI2: strontium iodide ZnS: zinc sulfide CaBr2: calcium bromide aluminum sulfide: Al2S3 Rb2O: rubidium oxide ```
49
naming type 2 binary ionic compounds
- cation always named 1st - anion named 2nd (-ide) - charge of cation is specified by roman numeral
50
examples of type 2 binary ionic compounds
``` CuBr: copper (I) bromide FeS: iron (II) sulfide PbO2: lead (IV) oxide MnI2: manganese (II) iodide CoCl3: cobalt (III) chloride CuI: copper (I) iodide tin (IV) bromide: SnBr4 ```
51
name CrO2
chromium (IV) oxide
52
name chromium (II) fluoride
CrF2
53
what is the name of SrB2?
strontium bromide
54
what is the name of K2S?
potassium sulfide
55
what is the correct name of the compound that results from the most stable ion for sulfur & the metal ion that contains 24 e-?
iron (II) sulfide
56
binary compounds: type 3
- non-metal and non-metal | - greek prefixes to denote the # of atoms of each element
57
naming type 3 binary compounds
- 1st element named first & full element name is used - 2nd element is named as though it were an anion - prefixes are used to denote # of atoms present - prefix mono- is never used for naming the 1st element
58
prefixes used to indicate numbers in chemical name
``` prefix number mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 nona- 9* deca- 10* ```
59
examples of type 3 binary covalent compounds
``` CO2: carbon dioxide SF6: sulfur hexafluoride N2O4: dinitrogen tetroxide CO: carbon monoxide NO2: nitrogen dioxide selenium hexafluoride: SeF6 PCl5: phosphorus pentachloride dinitrogen monoxide: N2O ```
60
name SeO2
selenium dioxide
61
polyatomic ions
charged entities composed of several atoms bound together - have special names * *MUST MEMORIZE**
62
examples of polyatomic ions
``` NaOH: sodium hydroxide Mg(NO3)2: magnesium nitrate (NH4)2SO4: ammonium sulfate Fe3(PO4)2: iron (II) phosphate Ca(HCO3)2: calcium bicarbonate potassium permanganate: KMnO4 Fe(OH)2: iron (II) hydroxide antimony (III) oxide: Sb2O3 ```
63
rules for naming acids
- if the anion DOESN'T HAVE an oxygen, the acid is named w/ the prefix hydro- and the suffix -ic attached to root of element name - if anion DOES HAVE an oxygen, the acid name is formed from the root element name w/ the suffix -ic or -ous after it * *ALL END IN "acid"** - ate = -ic - ite = -ous
64
examples of naming acids
``` HCl: hydrochloric acid HCN: hydrocyanic acid H2S: hydrosulfuric acid HF: hydrofluoric acid H3PO4: phosphoric acid H2S: hydrosulfuric acid nitric acid: HNO3 H2SO4: sulfuric acid HC2H3O2: acetic acid ```
65
names of some acids with O
``` ACID NAME HNO3 nitric acid HNO2 nitrous acid H2SO4 sulfuric acid H2SO3. sulfurous acid H3PO4 phosphoric acid HC2H3O2 acetic acid ```
66
names of acids without O
``` ACID NAME HF hydrofluoric acid HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid HCN hydrocyanic acid H2S hydrosulfuric acid ```
67
NH4 +
ammonium
68
NO2 -
nitrite
69
NO3 -
nitrate
70
SO3 2-
sulfite
71
SO4 2-
sulfate
72
HSO4 -
hydrogen sulfate/bisulfate
73
OH -
hydroxide
74
CN -
cyanide
75
PO4 3-
phosphate
76
HPO4 2-
hydrogen phosphate
77
H2PO4 -
dihydrogen phosphate
78
CO3 2-
carbonate
79
HCO3 -
hydrogen carbonate/ bicarbonate
80
ClO -
hypochlorite
81
ClO2 -
chlorite
82
ClO3 -
chlorate
83
ClO4 -
perchlorate
84
C2H3O2 -
acetate
85
MnO4 -
permanganate
86
Cr2O7 2-
dichromate
87
O2 2-
peroxide
88
clues that a chemical rxn has occurred
1) the colors change 2) a solid forms 3) bubbles form 4) heat and/or a flame is produced, or heat is absorbed
89
balancing a chemical equation
- coefficients balance rxn - sum of coefficients * diatomic mols ∆ eq 1) atoms are neither created nor destroyed 2) can never ∆ subscripts in a complete rxn 3) reactants -> products
90
how to write + balance equations
- read description of chm rxn - identify reactants, products, and their states - write apprp formulas - write unbalanced eq that summarizes the info from previous step - balance the equation by inspection, starting w/ most complicated mol - should finish w/ the same # of molecules on both sides of arrow
91
precipitate
the solid formed during a precipitation rxn | *exchange cations & anions
92
what are the 3 types of equations for rxns in aqueous solutions?
- molecular - complete ionic - net ionic
93
molecular equation
- shows complete formulas of all reactants and products - doesn't give a very clear picture of what actually occurs in solution - consists of mols are a whole
94
complete ionic equation
- all substances that are strong electrolytes are represented as ions - spectator ions
95
spectator ions
ions that don't participate directly in a rxn in sol | * the ions that cancel out
96
net ionic equation
- includes only those components that are directly involved in the rxn - spectator ions not included - order doesn't matter - should always be balanced
97
Write the correct molecular equation, the complete ionic equation, and the net ionic equation for the reaction between cobalt(II) chloride and sodium hydroxide
molecular: CoCl2 (aq) + 2 NaOH (aq) -> Co(OH)2 (s) +2 NaCl (aq) complete ionic: Co2+ (aq) + 2 Cl- (aq) + 2 Na+ (aq) + 2 OH- (aq) -> Co(OH)2 (s) + 2 Na+ (aq) + 2 Cl- (aq) net ionic: Co2+ (aq) + 2 OH- (aq) -> Co(OH)2 (s)
98
what are the 7 strong acids?
1) H2SO4 = sulfuric acid 2) HCl = hydrochloric acid 3) HBr = hydrobromic acid 4) HI = hydroiodic acid 5) HNO3 = nitric acid 6) HClO4 = perchloric acid 7) HClO3 = chloric acid
99
strong acid
a strong electrolyte that produces H+ ions (protons) when it is dissolved in water
100
acids give __ to bases; bases accept __ and give __ to acids
H+; H+; OH-
101
is an acid and base is a good relationship?
yes
102
strong base
a substance that produces hydroxide ions (OH-) in water
103
what are the strong bases?
Group 1 Group 2 + OH-
104
the net ionic equation for the rxn of a strong acid and a strong base is always the production of __
water
105
Avogadro's number
6.022 x 10^23
106
molar mass
mass in grams of a substance unit: g/mol - coefficients don't count, only subscripts
107
conversion trick
``` K H D b D C M king henry died by drinking chocolate milk km, hm, Dm, base unit (m, L, g), cm, mm kilo: 1000 hecto: 100 deca: 10 U: unit deci: 0.1 centi: 0.01 milli: 0.001 *micro: 0.000001 or 10^-6 *nano: 0.000000001 or 10^-9 ```
108
mass percent of an element in a compound
mass of element present in 1 mole of compound / mass of 1 mole of compound
109
empirical formula
formula of a compound that expresses the smallest whole-number ratio of the atoms present -simplest whole-number ratio
110
how to find the empirical formula
assume 100 g and change % to g convert g to mol divide by lowest number round any decimals
111
molecular formula
the exact formula of the molecules present in a substance (empirical formula)n *n is an integer *always bigger than empirical formula
112
how to find the molecular formula
- use the molar mass given and divide by the molar mass from ur new empirical formula - use the number to multiply the empirical formula
113
do diatomics or sig figs matter in empirical/molecular formulas?
no
114
A gaseous compound containing carbon and hydrogen was analyzed and found to consist of 83.65% carbon by mass. Determine the empirical formula of the compound.
C3H7
115
A gaseous compound containing carbon and hydrogen was analyzed and found to consist of 83.65% carbon by mass. The molar mass of the compound is 86.2 g/mol. The empirical formula was determined to be C3H7. What is the molecular formula of the compound?
C6H14 molar mass of C3H7 = 43.086 g/mol 86.2 g/mol / 43.086 g/mol = 2 (C3H7) x 2 = C6H14
116
When 1.00 g of metallic chromium is heated with elemental chlorine gas, 3.045 g of a chromium chloride salt results. Calculate the empirical formula of the compound.
CrCl3 1.00g Cr 3.045 g CrCl 3.045 - 1.00 = 2.045 Cl Cr 1.00/51.996 = 0.0192 / 0.0192 = 1 Cl 2.045/35.45 = 0.0577 / 0.0192 = 3 CrCl3
117
CrO4 2-
chromate
118
mole ratio
coefficient of unknown / coefficient of known
119
Consider the following balanced equation:Na2SiF6(s) + 4Na(s) → Si(s) + 6NaF(s) How many moles of NaF will be produced if 3.50 moles of Na is reacted with excess Na2SiF6?
5.25 moles NaF
120
Propane, C3H8, is a common fuel used for heating in rural areas. Predict the number of moles of CO2 formed when 3.74 moles of propane is burned in excess oxygen. oxygen. C3H8+ 5O2→ 3CO2+ 4H2O
11.2 moles
121
steps for calculating the masses of reactants and products in chm rxns
- balance the eq for the rxn - convert the masses of reactants/products to moles - use balanced eq to set up appropriate mole - use mole ratio(s) to calculate the # of moles of desired reactant or product - convert moles back to mass
122
For the following unbalanced equation: Cr(s) + O2(g) → Cr2O3(s) How many grams of chromium(III) oxide can be produced from 15.0 g of solid chromium and excess oxygen?
21.9 g Cr2O3
123
limiting reactants
- determine which reactant is limiting to calculate the amnt of products tht will be formed - methane and water will react to form products according to the eq: CH4 + H2O -> 3 H2 + CO a) H2O molecules are used up first, leaving 2 unreacted CH4 molecules - amnt of products that can form is limited by the water - water is limiting reactant - methane is in excess
124
when solving for the amount of excess left over we must find out how much was __ - how much we __ __
needed; started with
125
percent yield
actual yield / theoretical yield x 100%
126
find the percent yield of a product if 1.50g of SO3 is made from 1.00g of O2 + excess sulfur 2 S + 3 O2 -> 2 SO3
89.8%
127
when solving the molar mass of diatomic molecules, always
x 2
128
energy
the ability to do work and produce heat | - needed to oppose natural attractive forces
129
law of conservation of energy
- energy can be converted from one form to another - can neither be created nor destroyed - total e content of the universe is constant
130
system
part of the universe on which we wish to focus attention
131
surroundings
include every this else in the universe
132
endothermic process
- heat flow is INTO a system - ABSORBS e from surroundings (+) q
133
exothermic process
- e flows OUT of the system | (-) q
134
energy gained by surroundings must be __ to the energy lost by the system
equal
135
feels warmer is (exothermic/endothermic)
exothermic | - giving off heat
136
feels colder is (exothermic/endothermic)
endothermic | - heat is trapped
137
is freezing water endothermic or exothermic?
exothermic
138
Endo or exo: your hand gets cold when u touch ice
exothermic
139
endo or exo: the ice gets warmer when u touch it
endothermic
140
endo or exo: water boils in a kettle being heated on a stove
endothermic
141
endo or exo: water vapor condenses on a cold pipe
exothermic
142
endo or exo: ice cream melts
endothermic
143
endo or exo: methane is burning in a Bunsen burner in a lab
endothermic
144
the more ordered something is, the __ energy it has
less
145
internal energy (E)
- E of a system is the sum of the KE and PE of all the particles in the system - ∆ in the E of a system is ∆E = q + w - sign reflects the system's POV - endothermic process: (+) q - exothermic process: (-) q
146
internal energy formula
∆E = q + w
147
system does work on the surroundings, w is __
(-)
148
surroundings do work on the system, w is
(+)
149
1 calorie = __ joules
4.184 J
150
specific heat capacity
the e required to ∆ the temp of 1 gram of a substance by 1ºC
151
formula to calculate energy required for a rxn
``` Q = mc∆T *Q= energy (heat) required (J) m= mass (g) c= specific heat capacity (J/g ºC) ∆T= change in temp (ºC) (Tfinal - Tinitial) ```
152
calculate the amnt of heat e (in joules) needed to raise the temp of 6.25g of water from 21.0ºC to 39.0ºC *specific heat capacity of water: 4.184 J/g ºC
471 J
153
A 100.0 g sample of water at 90 °C is added to a 100.0 g sample of water at 10 °C the final temp of the water is likely to be: a) between 50ºC and 90ºC b) 50ºC c) between 10ºC and 50ºC
50 ºC
154
A 100.0 g sample of water at 90.0 °C is added to a 500.0 g sample of water at 10.0 °C 1) The final temperature of the water is likely to be: a) between 50ºC and 90ºC b) 50ºC c) between 10ºC and 50ºC 2) calculate the final temp of water
1) between 10ºC and 50ºC | 2) 23.3ºC
155
You have a Styrofoam cup with 50.0 g of water at 10.0 C. You add a 50.0 g iron ball at 90.0 C to the water. (sH2O= 4.18 J/g °C and sFe= 0.45 J/g °C) 1) The final temperature of the water is likely to be: a) between 50ºC and 90ºC b) 50ºC c) between 10ºC and 50ºC 2) calculate the final temp of the water.
1) between 10ºC and 50ºC | 2) 18ºC
156
change in enthalpy (∆H)
- state function - ∆H= q at constant pressure - ∆Hp = heat
157
Consider the rxn: S(s) + O2(g) → SO2(g)ΔH = –296 kJ per mole of SO2 formed Calculate the quantity of heat released when 2.10 g of sulfur is burned in oxygen at constant pressure
-19.4 kJ
158
Consider the combustion of propane: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)ΔH= –2221 kJ Assume that all of the heat comes from the combustion of propane. Calculate ΔH when 5.00 g of propane is burned in excess oxygen at constant pressure
-252 kJ
159
Hess's Law
- in going from a particular set of reactants to a particular set of products, the ∆H is the same whether the rxn takes lace in one step or in a series of steps
160
characteristics of enthalpy changes
- if a rxn is reversed, the sign of ∆H is also reversed - magnitude of ∆H is directly proportional to the quantities of reactants and products in a rxn - if the coefficients in a balanced rxn are multiplied by an integer, the value of ∆H is multiplied by that integer
161
1) NH3 (g) -> 1/2N2 (g) + 3/2H2 (g) ∆H= 46 kJ 2) 2 H2 (g) + O2 (g) -> 2 H2O (g) ∆H= -484 kJ calculate the ∆H for the rxn: 2 N2 (g) + 6 H2O(g) -> 3 O2 (g) + 4 NH3 (g)
1268 kJ
162
``` Calculate ΔH for the reaction: SO2+ ½O2→SO3 Given: 1) S + O2→ SO2ΔH = –297 kJ 2) 2S + 3O2→2SO3ΔH = –792 kJ ```
-99 kJ
163
Hess's Law: like terms
- same sides : ADD | - opposite sides: SUBTRACT
164
``` Given: 1) C(s) + O2(g) -> CO2(g) ∆H = -393 kJ 2) 2CO(g) + O2(g) -> 2CO2(g) ∆H = -566 kJ Calculate the ∆H for the rxn: C(s) + 1/2O2(g) ->CO(g) ```
-110 kJ
165
orbitals
- nothing like orbits! - probability of finding an e- within a certain space around the nucleus was predicted - the wave mechanical model gives no info abt when the e- occupies a certain pt in space/ how it moves - don't have sharp boundaries - chemists define an orbital's size as the sphere that contains 90% of the total e- probability
166
hydrogen energy levels
- called principal energy levels - they are indicated by whole #s - what goes up must go down * start at ground state and go up to n=4
167
e- configuration tree
``` 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p ```
168
e- configuration
``` s= 2 p= 6 d= 10 f= 14 *divide numbers above by 2 to determine the # of boxes for the orbital diagram ```
169
where can u find the number of e- for an atom?
periodic table
170
e- configuration of Li atom
1s1 2s1
171
e- configuration of an O atom
1s1 2s2 2p4
172
e- configuration of S
1s2 2s2 2p6 3s2 3p4
173
what is the e- configuration of an H atom?
1s1
174
as we go across a period from left to right, the ionization e ___
increases
175
as we go down a group from top to bottom, the ionization e ___
decreases
176
as we go across a period from left to right, the atomic radius __
decreases
177
as we go down a group from top to bottom, the atomic radius ___
increases
178
as we go down a group, the size of the atom __
increases
179
core e-
the inner e- that aren't involved in binding atoms to each other
180
valence e-
e- in the outermost or the highest principal e level of an atom - 1s2 2s2 2p6, where valence e- equal 8 - the elements in the same group have the same valence e- configuration - elements w the same valence e- arrangement show very similar chm behavior
181
the elements in the same group have the same __
valence e- configuration
182
ionic bonding
- metal + a non-metal | - e- are transferred from an atom that can lose e- relatively easily to an atom tht has a high affinity for e-
183
covalent bonding
aka non polar covalent - equal sharing of e- - e- are shared by nuclei - same element
184
polar covalent bond
- unequal sharing of e- between 2 atoms - one atom attracts the e- more than the other atom - results in development of bond polarity (partial (+) and (-) charge)
185
as we go across a period, electronegativity __
increases
186
as we go down a group, electronegativity ___
decreases
187
what is the most electronegative element?
fluorine
188
what is the least electronegative element?
cesium and francium
189
rank the following from smallest to largest atomic radius: | Ar, S 2-, Ca 2+, K+, Cl-
Ca2+ < K+ < Ar < Cl- < S 2-
190
the cation is always __ than the anion
smaller
191
which atom or ion has the smallest radius? a) O 2+ b) O + c) O d) O 2-
O 2+
192
the polarity of a bond depends on the difference between the __ values of the atoms forming in the bond
electronegativity
193
arrange the following bonds from most polar to least polar a) N-F O-F C-F b) C-F N-O Si-F c) Cl-Cl B-Cl S-Cl
a) C-F, N-F, O-F b) Si-F, C-F, N-O c) B-Cl, S-Cl, Cl-Cl
194
which of the following bonds would be the least polar yet still be considered polar covalent? Mg-O, C-O, O-O, Si-O, N-O
N-O
195
which of the following bonds would be the most polar w/o being considered ionic? Mg-O, C-O, O-O, Si-O, N-O
Si-O
196
bond polarity: dipole moment
- property of a molecule whose charge distribution can be represented by a center of a positive charge and a center of (-) charge - dipole character of a molecule is represented by an arrow aka vector - arrow points to (-) charge center and its tail indicates the (+) center of charge * tug of war
197
how to find the number of core and valence e-
look at the group # to determine the # of valence e- and subtract the atomic number by that to get the # of core e-
198
cations are always __ than their parent atoms
smaller
199
anions are always __ than their parent atoms
larger
200
lewis structures
- representation of molecules - shows how valence e- are arranged among atoms in a molecule - octet rule - duet rule
201
writing lewis structures
- bonding pairs are shared between 2 atoms | - unshared pairs aka lone pairs are not shared and not involved in bonding
202
octet rule
outermost shell has 8 e-
203
steps for writing lewis structures
- sum the valence e- from all atoms - use 1 pair of e- to form a bond between each pair of bound atoms - arrange the remaining e- to satisfy the octet rule for each second-row element & duet rule for hydrogen
204
in lewis structures, everything must have __ rule
octet
205
exceptions to the octet rule
- boron can have 6 e- | - sulfur can have 10 e-
206
draw a lewis structure for BF3
24 e-
207
linear structure
atoms arranged in a line
208
trigonal planar structure
atoms arranged in a triangle
209
diatomics have __ polarity
low
210
dipole moments only occur in __ bonds
covalent
211
in dipole moments, the farther apart the atoms are, the more __
polar
212
find the molecular structure of NH3
tetrahedral
213
as we go left to right, the atomic size __
decreases
214
how to find the amount of excess left over in rxn
1) moles of LR x mole ratio (moles of ER ÷ moles of LR) = moles used of ER 2) moles started with ER - moles used ER = moles ER left over 3) convert moles to grams by x molar mass of ER
215
exothermic has a __ w
(+)
216
endothermic has a __ w
(-)
217
system does work, w is __
(-)
218
surroundings do work, w is __
(+)
219
gas
- Uniformly fills any container - Mixes completely with any other gas - Exerts pressure on its surroundings - has small forces between them
220
Unit millimeters of mercury (mm Hg) is often called the
torr
221
pressure formula
pressure = force / area
222
1 atm = __ mm Hg
760
223
1 atm = __ torr
760
224
1 atm = __ Pa
101,325
225
The pressure of the air in a tire is measured to be 28 psi. Represent this pressure in both torr and pascals.
1. 9 atm x 760 torr / 1 atm = 1.4 x 10^3 torr | 1. 9 atm x 101325 Pa/ atm = 1.9 x 10^5 Pa
226
The vapor pressure over a beaker of hot water is measured as 656 torr. What is this pressure in atmospheres?
656 torr x 1 atm / 760 torr = 0.863 atm
227
Robert Boyle's Experiment
Study of the relationship between the pressure of the trapped gas and its volume - when u increase P, you decrease V - ex: pillow
228
Boyle's Law Graph
L shaped curve - pressure and volume - inversely proportional
229
Boyle's Law
* P1V1 = P2V2 - Pressure and volume are inversely related or inversely proportional - PV = k - k is a constant for a given amnt of gas at a specific temp - V= k/p - T and moles are constant
230
A sample of helium gas occupies 12.4 L at 23 °C and 0.956 atm. What volume will it occupy at 1.20 atm assuming that the temperature stays constant?
9.88 L
231
absolute zero
temperature of −273 °C beyond which matter cannot be cooled
232
Charles's Law Graph
V on y axis and T on x axis - direct relationship - straight up
233
Charles’s Law
* V1/T1 = V2/T2 - Direct proportionality between volume and temperature(in Kelvin) represented by the equation known as Charles’s law - V = bT (b is a proportionality constant) - K = °C + 273 * 0 K is called absolute zero - constant pressure and moles
234
__ K is called absolute zero
0
235
Suppose a balloon containing 1.30 L of air at 24.7 °C is placed into a beaker containing liquid nitrogen at –78.5 °C. What will the volume of the sample of air become (at constant pressure)?
0.849 L
236
Avogadro’s Law
* n1/V1 = n2/V2 - Volume and numbers of moles are directly related (constant T and P) - V = an (a is a proportionality constant) - constant temp and pressure
237
Avogadro's Law Graph
- direct relationship - straight up - volume on y axis and moles on x axis
238
If 2.45 mol of argon gas occupies a volume of 89.0 L, what volume will 2.10 mol of argon occupy under the same conditions of temperature and pressure?
76.3 L
239
Ideal Gas Law
* PV = nRT - R = 0.08206 L atm / mol K - We can bring all the laws together under one comprehensive law: a) charles’s law: V =bT (constant P and n) b) Avogadro’s law: V =an (constant Tand P) c) Boyle’s law: V= k / P (constant T and n)
240
what is the universal gas constant?
0.08206 L atm / mol K
241
An automobile tire at 23 °C with an internal volume of 25.0 L is filled with air to a total pressure of 3.18 atm. Determine the number of moles of air in the tire.
3.27 mol
242
What is the pressure in a 304.0-L tank that contains 5.670 kg of helium at 25 °C?
114 atm
243
At what temperature (in °C) does 121 mL of CO2 at 27 °C and 1.05 atm occupy a volume of 293 mL at a pressure of 1.40 atm?
696 °C
244
Dalton’s Law of Partial Pressures
- For a mixture of gases in a container a) P total = P1 + P2 + P3 + . . . - the total pressure exerted is the sum of the pressures that each gas would exert if it were alone - The pressure of the gas is affected by the number of moles of particles present - The pressure is independent of the nature of the particles
245
Consider the following apparatus containing helium in both sides at 45 °C. Initially, the valve is closed. - After the valve is opened, what is the pressure of the helium gas?
2.25 atm
246
27. 4 L of oxygen gas at 25.0 °C and 1.30 atm and 8.50 L of helium gas at 25.0 °C and 2.00 atm were pumped into a tank with a volume of 5.81 L at 25 °C. - Calculate the new partial pressure of oxygen - Calculate the new partial pressure of helium - Calculate the new total pressure of both gases
new partial pressure of oxygen = 6.13 atm new partial pressure of helium = 2.93 atm new total pressure of both gases = 9.06 atm
247
Molar Volume of an Ideal Gas
- For 1 mole of an ideal gas at 0 °C and 1 atm, the volume of the gas is 22.42 L - STP
248
standard temperature and pressure (STP)
- 0 °C and 1 atm | - any gas with 0 °C and 1 atm will have a volume of 22.4 L at STP
249
A sample of oxygen gas has a volume of 2.50 L at STP. How many grams of O2 are present?
3.57 g
250
Consider the following reaction: Zn(s) + 2HCl(aq ) -> ZnCl(aq) + H(g) If 15.00 g of solid zinc reacts with 100.0 mL of 4.00 M hydrochloric acid, what volume of hydrogen gas will be produced at 25 °C and 1.00 atm?
4.89 L
251
Intramolecular Forces
* WITHIN - Take place within the molecule - Molecules are formed by sharing electrons between the atoms - Hold the atoms of a molecule together - 2 types: ionic and covalent *ionic is stronger
252
intermolecular forces
* BETWEEN - Forces that occur between molecules - 3 types: D-D, H-bonding, L-D
253
__ forces are stronger than __ forces
Intramolecular; intermolecular
254
what are the 2 types of intramolecular forces?
ionic and covalent | - ionic is stronger
255
what are the 3 types of intermolecular forces?
dipole-dipole, Hydrogen bonding, and London dispersion - H-bonding is strongest - London dispersion is weakest
256
dipole-dipole forces
- dipole moment - Molecules with dipole moments can attract each other electrostatically - They line up so that the positive and negative ends are close to each other - Only about 1% as strong as covalent or ionic bonds
257
hydrogen bonding
* Hydrogens have FON * strongest - Strong dipole–dipole forces occur between molecules when hydrogen is bound to a highly electronegative atom - Nitrogen, oxygen, or fluorine (FON) - Affects physical properties (Boiling point) a) stronger the force the higher the boiling pt
258
London Dispersion Forces
* weakest * occurs in ALL molecules - e- constantly moving - Forces that exist among noble gas atoms and non-polar molecules - Instantaneous dipole that occurs temporarily in a given atom induces a similar dipole in a neighboring atom - Significant in large atoms/molecules - Occurs in all molecules, including non-polar ones
259
London Dispersion Forces: Non-polar molecules
- also interact by developing instantaneous dipoles | - Become stronger as the size of atoms or molecules increases
260
melting and boiling points
Stronger the intermolecular forces, higher the melting and boiling points - H-bonding has highest melting and boiling points
261
Which molecule is capable of forming stronger intermolecular forces? H2O or N2?
H2O
262
Which gas would behave more ideally at the same conditions of P and T? CO or N2?
N2
263
Consider the following compounds: NH3, CH4, H2 | How many of the compounds above exhibit London dispersion forces?
3
264
the stronger the force the __ the vapor pressure
lower
265
London dispersion forces have __ vapor pressures
high
266
Which of the following would be expected to have the highest vapor pressure at room temperature? a) CH3CH2CH2OH b) CH3CH2CH2NH2 c) CH3CH2CH2CH3 d) CH3CH2CH3
d) CH3CH2CH3
267
Crystalline Solids
Substances with a regular arrangement of their components form crystalline solids
268
types of crystalline solids
- ionic solids - molecular solids - atomic solids
269
ionic solids
- components are ions | - Ions at the points of the lattice that describes the structure of the solid
270
molecular solids
- components are molecules | - Discrete covalently bonded molecules at each of its lattice points
271
atomic solids
- components are atoms | - Atoms at the lattice points that describe the structure of the solid
272
examples of 3 types of crystalline solids
1) diamond -atomic 2) NaCl -ionic 3) H2O (ice) -molecular
273
solubility of ionic substances
- Ionic substances break up into individual cations and anions when dissolved in water - Polar water molecules interact with the positive and negative ions of a salt - Ethanol is soluble in water because of its polar O—H bond - ion breaks into cation and anion
274
solubility of polar substances
Why is solid sugar soluble in water? | because of the O-H polar bond
275
Why is solid sugar soluble in water?
because of the O-H polar bond
276
How Substances Dissolve
- A “hole”must be made in the water structure for each solute particle - The lost water–water interactions must be replaced by water–solute interactions - “Like dissolves like”
277
like dissolves __
like
278
polar molecules dissolve in __
polar solvents
279
``` Which of the following solutes will generally not dissolve in the specified solvent? Choose the bestanswer. Assume all of the compounds are in the liquid state. a) CCl4 mixed with water or H2O b )NH3 mixed with water or H2O c) CH3OH mixed with water or H2O d) N2mixed with methane or CH4 ```
a) CCl4 mixed with water or H2O
280
Saturated solution
contains as much solute as will dissolve at that temperature
281
Unsaturated solution
is a solution that has not reached the limit of solute that will dissolve in it
282
Relatively large amount of solute is dissolved in a __ solution
concentrated
283
Relatively small amount of solute is dissolved in a __ solution
dilute
284
mass percent
mass percent = mass of solute ÷ mass of solution x 100% | mass percent = grams of solute ÷ (grams of solute + grams of solvent) x 100%
285
What is the percent-by-mass concentration of glucose in a solution made by dissolving 5.5 g of glucose in 78.2 g of water?
6.6%
286
calculating molarity
- Molarity, M, equals the number of moles of solute per volume of solution in liters M = moles of solute / liters of solution
287
you have 1.00 mol of sugar in 125.0 mL of solution. Calculate the concentration in units of molarity.
8.00 M
288
concentration of ions
- concentration of ions stays consistent For a 0.25 M CaCl2 solution:CaCl2→ Ca2+ + 2Cl- - Ca2+: 1 × 0.25 M = 0.25 M Ca2+ - Cl- : 2 × 0.25 M = 0.50 M Cl–
289
Standard Solution
A solution whose concentration is accurately known | - curve is straight up
290
Steps to Make a Standard Solution
1. Weigh out a sample of solute 2. Transfer it to a volumetric flask 3. Add enough solvent to bring the volume up to the mark on the neck of the flask
291
dilution
M1 x V1 = M2 x V2 - The process of adding water to a concentrated or stock solution to achieve a solution of desired concentration - Dilution with water does not alter the numbers of moles of solute present - Moles of solute before dilution equals moles of solute after dilution * changing denominator of molarity
292
Steps to Dilute a Solution
1. Transfer a measured amount of original solution to a flask containing some water 2. Add water to the flask by swirling to bring the volume up to the calibration mark 3. Mix by inverting the flask
293
What is the minimum volume of a 2.00-M NaOH solution needed to make 150.0 mL of a 0.800-M NaOH solution?
60.0 mL
294
What is the minimum volume of a 5.68-M NaOH solution needed to make 323.0 mL of a 2.96-M NaOH solution?
168 mL | * 155 mL water added (323 - 168 = 155)
295
Steps for Solving Stoichiometric Problems Involving Solutions
1. Write the balanced equation for the reaction - For reactions involving ions, it is best to write the net ionic equation 2. Calculate the moles of reactants 3. Determine which reactant is limiting 4. Calculate the moles of other reactants or products, as required 5. Convert to grams or other units, if required
296
10. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change) - What precipitate will form? - What mass of precipitate will form?
- lead(II) phosphate, Pb3(PO4)2 | - 1.1 g Pb3(PO4)
297
10. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change) - What is the concentration of nitrate ions left in the solution after the reaction is complete?
0.27 M
298
10. 0 mL of a 0.30-M sodium phosphate solution reacts with 20.0 mL of a 0.20-M lead(II) nitrate solution (assume no volume change) - What is the concentration of phosphate ions left in the solution after the reaction is complete?
0.011 M
299
dipole moment
molecules w polar bonds often behave in an electric field as if they had a center of positive charge (+) end attracts (-) end
300
non polar
equal sharing of e- (diatomics)
301
polar
non-equal sharing of e- (HCl)
302
Conjugate base
everything that remains of the acid molecule after a proton is lost
303
Conjugate acid
formed when a proton is transferred to the base
304
Conjugate acid–base pair
Consists of two substances related to each other by the donating and accepting of a single proton
305
acids in water
Water acts as a base accepting a proton from the acid | - Forms hydronium ion (H3O+)
306
``` Which of the following pairs represents a conjugate acid–base pair? a) HCl, HNO3 b )H3O+, OH– c) H2SO4, SO42– d) HCN, CN– ```
d) HCN, CN–
307
strong acid
- Completely ionized or completely dissociated HA(aq) + H2O(l) → H3O+(aq) + A–(aq) * breaks apart 100% - forward rxn predominates
308
weak acid
- Most of the acid molecules remain intact HA(aq) + H2O(l) ← H3O+(aq) + A–(aq) - A- is a much stronger base than H2O - strength of the conjugate base compared with that of water
309
Strong acids contain relatively __ conjugate bases
weak
310
weak acids have a __ base
strong
311
``` Consider a 1.0 M solution of HCl • Order the following from strongest to weakest base and explain your answer – H2O(l) – A–(aq) (from weak acid HA) – Cl–(aq) ```
A-, H2O, Cl-
312
water is _
amphoteric | - can be both an acid (H3O+) and a base (OH-)
313
at 25ºC, Kw =
Kw = [H+] [OH-] = 1.0 x 10^-14 | *no matter what the solution contains
314
[H+] = [OH–]
neutral solution
315
H+] > [OH–]
acidic solution
316
[H+] < [OH–]
basic solution
317
in an acidic, basic, or neutral solution Kw always equals
Kw = [H+][OH–] = 1.0 × 10^-14
318
an acid/base reaction is also called a
neutralization reaction
319
pH =
− log [H+]
320
as [H+] increases, pH __
decreases
321
pH range
7 is neutral > 7 is basic < 7 is acidic *The lower the pH, them ore acidic the solution
322
pH scale
scale is from 0-14 - acidic: 0-6.9 - neutral: 7 - basic: 7.1-14
323
pOH =
- log [OH-]
324
[H+] =
10^-pH
325
[OH-] =
10^-pOH
326
pH + pOH =
14
327
[H+] [OH-] =
1.0 x 10^ -14
328
calculate the pH for 1.0×10^-4 M H+
4.00 | pH = –log[H+] = –log(1.0 × 10–4 M) = 4.00
329
calculate the pH for 0.040 M OH–
12.60
330
The pH of a solution is 5.85. What is the [H+] for this solution?
[H+] = 1.4 × 10^-6 M
331
calculate the pOH for 0.040 M OH–
1.40
332
The pH of a solution is 5.85. What is the [OH–] for this solution?
7.1 × 10^-9 M
333
To convert from moles to atoms, ___
multiply by Avogadro's number
334
To convert from atoms to moles, __
divide by Avogadro's number (or multiply by its reciprocal)
335
To convert from grams to moles, __
divide using molar mass
336
to covert from moles to grams, __
multiply by molar mass
337
HA (aq) + H2O (l) -> H3O+ (aq) + A- (aq) | label acid, base, conjugate acid, and conjugate base
acid: HA base: H2O conjugate acid: H3O+ conjugate base: A-
338
group 1 element charges | +1
``` alkali metals Highly Nasty Kids Rub Cats Li (lithium) Na (sodium) K (potassium) Rb (rubidium) Cs (cesium) ```
339
group 2 element charges | +2
``` alkaline earth metals Beer Mugs Can Serve Bar Rats beryllium (Be) magnesium (Mg) calcium (Ca) strontium (Sr) barium (Ba) radium (Ra) ```
340
group 3 element charges | +3
``` Bears Always Give In boron (B) aluminum (Al) gallium (Ga) indium (In) ```
341
group 6 element charges | -2
``` Old Soldiers Seem Tense oxygen (O) sulfur (S) seaborgium (Se) tellurium (Te) ```
342
group 7 element charges | -1
``` halogens “Floor Cleaner Broken?” I Asked fluorine (F) chlorine (Cl) bromine (Br) iodine (I) ```
343
8 strong bases
Group 1 and 2 + OH- 1) NaOH 2) KOH 3) LiOH 4) RbOH 5) CsOH 6) Ca(OH)2 7) Sr(OH)2 8) Ba(OH)2

CHM 1040 (6 decks)

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