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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
Q

metalloids

A

ON the staircase

-have some metallic and some non-metallic properties

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26
Q

what are the metalloids?

A 
Boring Silly Germs Are Ants Telling Politics
Boron (B)
Silicon (Si)
Germanium (Ge)
Arsenic (As)
Antimony (Sb)
Tellurium (Te)
Polonium (Po)
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27
Q

what are the four physical properties of metals?

A

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)

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28
Q

physical properties of non-metals

A
  • lack properties of metals
  • exhibit more variation in properties
  • can be g, l, s @ room temp
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29
Q

physical properties of metalloids

A

exhibit a mixture of metallic and non-metallic properties

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30
Q

7 diatomic molecules

A

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
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31
Q

ions

A
  • atom’s not neutral
  • atom w a charge
  • elements become ions
  • imbalance of e-
  • atoms can form ions by gaining/losing e-
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32
Q

cations

A

metals tend to LOSE 1 or more e- to form (+) ions

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33
Q

anions

A

non-metals tend to GAIN 1 or more e- to form (-) ions

-name changes to end in -ide

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34
Q

ion charges and the periodic table chart

A

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
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35
Q

what are the 3 steps of the scientific method?

A

1) make an observation
2) form a hypothesis
3) perform experiment

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36
Q

theory vs law

A
  • theory: answers “why?”, leads to more questions

- the law: “this is what happened”, doesn’t lead to more questions

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37
Q

which of the 3 subatomic particles are the smallest?

A

electrons

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38
Q

__ contribute to the mass of an atom

A

protons

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39
Q

__ contribute to the size of an atom

A

electrons

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40
Q

__ dictates chemistry of an atom

A

electrons

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41
Q

naming compounds: binary compounds

A
  • composed of 2 elements
  • divided into broad classes
  • compounds that contain a metal & non-metal (ionic)
  • compounds that contain 2 non-metals (covalent)
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42
Q

nomenclature

A

naming

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43
Q

naming compounds: binary ionic compounds

A
  • contains (+) cations and (-) anions
  • type 1
  • type 2
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44
Q

binary ionic compounds: type 1

A

-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+

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45
Q

charges of transition metals

A
  • Al3+
  • Ga3+
  • In3+
  • Zn2+
  • Ag+
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46
Q

binary ionic compounds: type 2

A
  • 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
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47
Q

naming type 1 binary ionic compounds

A
  • cation is always named first (element name)

- anion named second (end in -ide)

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48
Q

examples of type 1 binary ionic compounds

A 
NaCl: sodium chloride
CaS: calcium sulfide
Kl: potassium iodide
SrI2: strontium iodide 
ZnS: zinc sulfide
CaBr2: calcium bromide
aluminum sulfide: Al2S3
Rb2O: rubidium oxide
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49
Q

naming type 2 binary ionic compounds

A
  • cation always named 1st
  • anion named 2nd (-ide)
  • charge of cation is specified by roman numeral
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50
Q

examples of type 2 binary ionic compounds

A 
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
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51
Q

name CrO2

A

chromium (IV) oxide

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52
Q

name chromium (II) fluoride

A

CrF2

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53
Q

what is the name of SrB2?

A

strontium bromide

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54
Q

what is the name of K2S?

A

potassium sulfide

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55
Q

what is the correct name of the compound that results from the most stable ion for sulfur & the metal ion that contains 24 e-?

A

iron (II) sulfide

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56
Q

binary compounds: type 3

A
  • non-metal and non-metal

- greek prefixes to denote the # of atoms of each element

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57
Q

naming type 3 binary compounds

A
  • 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
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58
Q

prefixes used to indicate numbers in chemical name

A 
prefix      number
mono-         1
di-               2
tri-               3
tetra-          4
penta-        5
hexa-         6
hepta-        7
octa-          8
nona-         9*
deca-         10*
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59
Q

examples of type 3 binary covalent compounds

A 
CO2: carbon dioxide 
SF6: sulfur hexafluoride
N2O4: dinitrogen tetroxide
CO: carbon monoxide
NO2: nitrogen dioxide
selenium hexafluoride: SeF6
PCl5: phosphorus pentachloride
dinitrogen monoxide: N2O
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60
Q

name SeO2

A

selenium dioxide

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61
Q

polyatomic ions

A

charged entities composed of several atoms bound together

  • have special names
  • *MUST MEMORIZE**
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62
Q

examples of polyatomic ions

A 
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
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63
Q

rules for naming acids

A
  • 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
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64
Q

examples of naming acids

A 
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
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65
Q

names of some acids with O

A 
ACID                 NAME
HNO3           nitric acid
HNO2          nitrous acid
H2SO4         sulfuric acid
H2SO3.         sulfurous acid 
H3PO4      phosphoric acid
HC2H3O2    acetic acid
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66
Q

names of acids without O

A 
ACID    NAME
HF       hydrofluoric acid
HCl      hydrochloric acid
HBr      hydrobromic acid
HI         hydroiodic acid 
HCN     hydrocyanic acid
H2S      hydrosulfuric acid
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67
Q

NH4 +

A

ammonium

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68
Q

NO2 -

A

nitrite

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69
Q

NO3 -

A

nitrate

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70
Q

SO3 2-

A

sulfite

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71
Q

SO4 2-

A

sulfate

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72
Q

HSO4 -

A

hydrogen sulfate/bisulfate

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73
Q

OH -

A

hydroxide

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74
Q

CN -

A

cyanide

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75
Q

PO4 3-

A

phosphate

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76
Q

HPO4 2-

A

hydrogen phosphate

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77
Q

H2PO4 -

A

dihydrogen phosphate

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78
Q

CO3 2-

A

carbonate

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79
Q

HCO3 -

A

hydrogen carbonate/ bicarbonate

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80
Q

ClO -

A

hypochlorite

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81
Q

ClO2 -

A

chlorite

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82
Q

ClO3 -

A

chlorate

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83
Q

ClO4 -

A

perchlorate

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84
Q

C2H3O2 -

A

acetate

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85
Q

MnO4 -

A

permanganate

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86
Q

Cr2O7 2-

A

dichromate

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87
Q

O2 2-

A

peroxide

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88
Q

clues that a chemical rxn has occurred

A

1) the colors change
2) a solid forms
3) bubbles form
4) heat and/or a flame is produced, or heat is absorbed

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89
Q

balancing a chemical equation

A
  • 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
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90
Q

how to write + balance equations

A
  • 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
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91
Q

precipitate

A

the solid formed during a precipitation rxn

*exchange cations & anions

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92
Q

what are the 3 types of equations for rxns in aqueous solutions?

A
  • molecular
  • complete ionic
  • net ionic
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93
Q

molecular equation

A
  • 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
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94
Q

complete ionic equation

A
  • all substances that are strong electrolytes are represented as ions
  • spectator ions
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95
Q

spectator ions

A

ions that don’t participate directly in a rxn in sol

* the ions that cancel out

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96
Q

net ionic equation

A
  • includes only those components that are directly involved in the rxn
  • spectator ions not included
  • order doesn’t matter
  • should always be balanced
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97
Q

Write the correct molecular equation, the complete ionic equation, and the net ionic equation for the reaction between cobalt(II) chloride and sodium hydroxide

A

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)

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98
Q

what are the 7 strong acids?

A

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

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99
Q

strong acid

A

a strong electrolyte that produces H+ ions (protons) when it is dissolved in water

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100
Q

acids give __ to bases; bases accept __ and give __ to acids

A

H+; H+; OH-

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101
Q

is an acid and base is a good relationship?

A

yes

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102
Q

strong base

A

a substance that produces hydroxide ions (OH-) in water

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103
Q

what are the strong bases?

A

Group 1
Group 2
+
OH-

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104
Q

the net ionic equation for the rxn of a strong acid and a strong base is always the production of __

A

water

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105
Q

Avogadro’s number

A

6.022 x 10^23

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106
Q

molar mass

A

mass in grams of a substance

unit: g/mol
- coefficients don’t count, only subscripts

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107
Q

conversion trick

A 
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
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108
Q

mass percent of an element in a compound

A

mass of element present in 1 mole of compound / mass of 1 mole of compound

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109
Q

empirical formula

A

formula of a compound that expresses the smallest whole-number ratio of the atoms present
-simplest whole-number ratio

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110
Q

how to find the empirical formula

A

assume 100 g and change % to g
convert g to mol
divide by lowest number
round any decimals

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111
Q

molecular formula

A

the exact formula of the molecules present in a substance
(empirical formula)n
*n is an integer
*always bigger than empirical formula

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112
Q

how to find the molecular formula

A
  • use the molar mass given and divide by the molar mass from ur new empirical formula
  • use the number to multiply the empirical formula
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113
Q

do diatomics or sig figs matter in empirical/molecular formulas?

A

no

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114
Q

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.

A

C3H7

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115
Q

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?

A

C6H14

molar mass of C3H7 = 43.086 g/mol

86.2 g/mol / 43.086 g/mol = 2

(C3H7) x 2 = C6H14

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116
Q

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.

A

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

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117
Q

CrO4 2-

A

chromate

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118
Q

mole ratio

A

coefficient of unknown / coefficient of known

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119
Q

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?

A

5.25 moles NaF

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120
Q

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

A

11.2 moles

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121
Q

steps for calculating the masses of reactants and products in chm rxns

A
  • 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
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122
Q

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?

A

21.9 g Cr2O3

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123
Q

limiting reactants

A
  • 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
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124
Q

when solving for the amount of excess left over we must find out how much was __ - how much we __ __

A

needed; started with

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125
Q

percent yield

A

actual yield / theoretical yield x 100%

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126
Q

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

A

89.8%

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127
Q

when solving the molar mass of diatomic molecules, always

A

x 2

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128
Q

energy

A

the ability to do work and produce heat

- needed to oppose natural attractive forces

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129
Q

law of conservation of energy

A
  • energy can be converted from one form to another
  • can neither be created nor destroyed
  • total e content of the universe is constant
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130
Q

system

A

part of the universe on which we wish to focus attention

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131
Q

surroundings

A

include every this else in the universe

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132
Q

endothermic process

A
  • heat flow is INTO a system
  • ABSORBS e from surroundings
    (+) q
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133
Q

exothermic process

A
  • e flows OUT of the system

(-) q

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134
Q

energy gained by surroundings must be __ to the energy lost by the system

A

equal

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135
Q

feels warmer is (exothermic/endothermic)

A

exothermic

- giving off heat

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136
Q

feels colder is (exothermic/endothermic)

A

endothermic

- heat is trapped

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137
Q

is freezing water endothermic or exothermic?

A

exothermic

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138
Q

Endo or exo: your hand gets cold when u touch ice

A

exothermic

139
Q

endo or exo: the ice gets warmer when u touch it

A

endothermic

140
Q

endo or exo: water boils in a kettle being heated on a stove

A

endothermic

141
Q

endo or exo: water vapor condenses on a cold pipe

A

exothermic

142
Q

endo or exo: ice cream melts

A

endothermic

143
Q

endo or exo: methane is burning in a Bunsen burner in a lab

A

endothermic

144
Q

the more ordered something is, the __ energy it has

A

less

145
Q

internal energy (E)

A
  • 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
Q

internal energy formula

A

∆E = q + w

147
Q

system does work on the surroundings, w is __

A

(-)

148
Q

surroundings do work on the system, w is

A

(+)

149
Q

1 calorie = __ joules

A

4.184 J

150
Q

specific heat capacity

A

the e required to ∆ the temp of 1 gram of a substance by 1ºC

151
Q

formula to calculate energy required for a rxn

A 
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
Q

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

A

471 J

153
Q

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

A

50 ºC

154
Q

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

A

1) between 10ºC and 50ºC

2) 23.3ºC

155
Q

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.

A

1) between 10ºC and 50ºC

2) 18ºC

156
Q

change in enthalpy (∆H)

A
  • state function
  • ∆H= q at constant pressure
  • ∆Hp = heat
157
Q

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

A

-19.4 kJ

158
Q

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

A

-252 kJ

159
Q

Hess’s Law

A
  • 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
Q

characteristics of enthalpy changes

A
  • 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
Q

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)

A

1268 kJ

162
Q 
Calculate ΔH for the reaction: 
SO2+ ½O2→SO3
Given: 
1) S + O2→ SO2ΔH = –297 kJ                 
2) 2S + 3O2→2SO3ΔH = –792 kJ
A

-99 kJ

163
Q

Hess’s Law: like terms

A
  • same sides : ADD

- opposite sides: SUBTRACT

164
Q 
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)
A

-110 kJ

165
Q

orbitals

A
  • 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
Q

hydrogen energy levels

A
  • 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
Q

e- configuration tree

A 
1s
2s   2p
3s   3p  3d
4s   4p  4d    4f
5s   5p  5d    5f
6s   6p  6d
7s   7p
168
Q

e- configuration

A 
s= 2
p= 6
d= 10
f= 14
*divide numbers above by 2 to determine the # of boxes for the orbital diagram
169
Q

where can u find the number of e- for an atom?

A

periodic table

170
Q

e- configuration of Li atom

A

1s1 2s1

171
Q

e- configuration of an O atom

A

1s1 2s2 2p4

172
Q

e- configuration of S

A

1s2 2s2 2p6 3s2 3p4

173
Q

what is the e- configuration of an H atom?

A

1s1

174
Q

as we go across a period from left to right, the ionization e ___

A

increases

175
Q

as we go down a group from top to bottom, the ionization e ___

A

decreases

176
Q

as we go across a period from left to right, the atomic radius __

A

decreases

177
Q

as we go down a group from top to bottom, the atomic radius ___

A

increases

178
Q

as we go down a group, the size of the atom __

A

increases

179
Q

core e-

A

the inner e- that aren’t involved in binding atoms to each other

180
Q

valence e-

A

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
Q

the elements in the same group have the same __

A

valence e- configuration

182
Q

ionic bonding

A
  • 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
Q

covalent bonding

A

aka non polar covalent

  • equal sharing of e-
  • e- are shared by nuclei
  • same element
184
Q

polar covalent bond

A
  • 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
Q

as we go across a period, electronegativity __

A

increases

186
Q

as we go down a group, electronegativity ___

A

decreases

187
Q

what is the most electronegative element?

A

fluorine

188
Q

what is the least electronegative element?

A

cesium and francium

189
Q

rank the following from smallest to largest atomic radius:

Ar, S 2-, Ca 2+, K+, Cl-

A

Ca2+ < K+ < Ar < Cl- < S 2-

190
Q

the cation is always __ than the anion

A

smaller

191
Q

which atom or ion has the smallest radius?

a) O 2+
b) O +
c) O
d) O 2-

A

O 2+

192
Q

the polarity of a bond depends on the difference between the __ values of the atoms forming in the bond

A

electronegativity

193
Q

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

a) C-F, N-F, O-F
b) Si-F, C-F, N-O
c) B-Cl, S-Cl, Cl-Cl

194
Q

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

A

N-O

195
Q

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

A

Si-O

196
Q

bond polarity: dipole moment

A
  • 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
Q

how to find the number of core and valence e-

A

look at the group # to determine the # of valence e- and subtract the atomic number by that to get the # of core e-

198
Q

cations are always __ than their parent atoms

A

smaller

199
Q

anions are always __ than their parent atoms

A

larger

200
Q

lewis structures

A
  • representation of molecules
  • shows how valence e- are arranged among atoms in a molecule
  • octet rule
  • duet rule
201
Q

writing lewis structures

A
  • bonding pairs are shared between 2 atoms

- unshared pairs aka lone pairs are not shared and not involved in bonding

202
Q

octet rule

A

outermost shell has 8 e-

203
Q

steps for writing lewis structures

A
  • 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
Q

in lewis structures, everything must have __ rule

A

octet

205
Q

exceptions to the octet rule

A
  • boron can have 6 e-

- sulfur can have 10 e-

206
Q

draw a lewis structure for BF3

A

24 e-

207
Q

linear structure

A

atoms arranged in a line

208
Q

trigonal planar structure

A

atoms arranged in a triangle

209
Q

diatomics have __ polarity

A

low

210
Q

dipole moments only occur in __ bonds

A

covalent

211
Q

in dipole moments, the farther apart the atoms are, the more __

A

polar

212
Q

find the molecular structure of NH3

A

tetrahedral

213
Q

as we go left to right, the atomic size __

A

decreases

214
Q

how to find the amount of excess left over in rxn

A

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
Q

exothermic has a __ w

A

(+)

216
Q

endothermic has a __ w

A

(-)

217
Q

system does work, w is __

A

(-)

218
Q

surroundings do work, w is __

A

(+)

219
Q

gas

A
  • Uniformly fills any container
  • Mixes completely with any other gas
  • Exerts pressure on its surroundings
  • has small forces between them
220
Q

Unit millimeters of mercury (mm Hg) is often called the

A

torr

221
Q

pressure formula

A

pressure = force / area

222
Q

1 atm = __ mm Hg

A

760

223
Q

1 atm = __ torr

A

760

224
Q

1 atm = __ Pa

A

101,325

225
Q

The pressure of the air in a tire is measured to be 28 psi. Represent this pressure in both torr and pascals.

A
  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
Q

The vapor pressure over a beaker of hot water is measured as 656 torr. What is this pressure in atmospheres?

A

656 torr x 1 atm / 760 torr = 0.863 atm

227
Q

Robert Boyle’s Experiment

A

Study of the relationship between the pressure of the trapped gas and its volume

  • when u increase P, you decrease V
  • ex: pillow
228
Q

Boyle’s Law Graph

A

L shaped curve

  • pressure and volume
  • inversely proportional
229
Q

Boyle’s Law

A
  • 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
Q

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?

A

9.88 L

231
Q

absolute zero

A

temperature of −273 °C beyond which matter cannot be cooled

232
Q

Charles’s Law Graph

A

V on y axis and T on x axis

  • direct relationship
  • straight up
233
Q

Charles’s Law

A
  • 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
Q

__ K is called absolute zero

A

0

235
Q

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)?

A

0.849 L

236
Q

Avogadro’s Law

A
  • 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
Q

Avogadro’s Law Graph

A
  • direct relationship
  • straight up
  • volume on y axis and moles on x axis
238
Q

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?

A

76.3 L

239
Q

Ideal Gas Law

A
  • 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
Q

what is the universal gas constant?

A

0.08206 L atm / mol K

241
Q

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.

A

3.27 mol

242
Q

What is the pressure in a 304.0-L tank that contains 5.670 kg of helium at 25 °C?

A

114 atm

243
Q

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?

A

696 °C

244
Q

Dalton’s Law of Partial Pressures

A
  • 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
Q

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?

A

2.25 atm

246
Q
  1. 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
A

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
Q

Molar Volume of an Ideal Gas

A
  • For 1 mole of an ideal gas at 0 °C and 1 atm, the volume of the gas is 22.42 L
  • STP
248
Q

standard temperature and pressure (STP)

A
  • 0 °C and 1 atm

- any gas with 0 °C and 1 atm will have a volume of 22.4 L at STP

249
Q

A sample of oxygen gas has a volume of 2.50 L at STP. How many grams of O2 are present?

A

3.57 g

250
Q

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?

A

4.89 L

251
Q

Intramolecular Forces

A
  • 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
Q

intermolecular forces

A
  • BETWEEN
  • Forces that occur between molecules
  • 3 types: D-D, H-bonding, L-D
253
Q

__ forces are stronger than __ forces

A

Intramolecular; intermolecular

254
Q

what are the 2 types of intramolecular forces?

A

ionic and covalent

- ionic is stronger

255
Q

what are the 3 types of intermolecular forces?

A

dipole-dipole, Hydrogen bonding, and London dispersion

  • H-bonding is strongest
  • London dispersion is weakest
256
Q

dipole-dipole forces

A
  • 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
Q

hydrogen bonding

A
  • 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
Q

London Dispersion Forces

A
  • 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
Q

London Dispersion Forces: Non-polar molecules

A
  • also interact by developing instantaneous dipoles

- Become stronger as the size of atoms or molecules increases

260
Q

melting and boiling points

A

Stronger the intermolecular forces, higher the melting and boiling points
- H-bonding has highest melting and boiling points

261
Q

Which molecule is capable of forming stronger intermolecular forces? H2O or N2?

A

H2O

262
Q

Which gas would behave more ideally at the same conditions of P and T? CO or N2?

A

N2

263
Q

Consider the following compounds: NH3, CH4, H2

How many of the compounds above exhibit London dispersion forces?

A

3

264
Q

the stronger the force the __ the vapor pressure

A

lower

265
Q

London dispersion forces have __ vapor pressures

A

high

266
Q

Which of the following would be expected to have the highest vapor pressure at room temperature?

a) CH3CH2CH2OH
b) CH3CH2CH2NH2
c) CH3CH2CH2CH3
d) CH3CH2CH3

A

d) CH3CH2CH3

267
Q

Crystalline Solids

A

Substances with a regular arrangement of their components form crystalline solids

268
Q

types of crystalline solids

A
  • ionic solids
  • molecular solids
  • atomic solids
269
Q

ionic solids

A
  • components are ions

- Ions at the points of the lattice that describes the structure of the solid

270
Q

molecular solids

A
  • components are molecules

- Discrete covalently bonded molecules at each of its lattice points

271
Q

atomic solids

A
  • components are atoms

- Atoms at the lattice points that describe the structure of the solid

272
Q

examples of 3 types of crystalline solids

A

1) diamond -atomic
2) NaCl -ionic
3) H2O (ice) -molecular

273
Q

solubility of ionic substances

A
  • 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
Q

solubility of polar substances

A

Why is solid sugar soluble in water?

because of the O-H polar bond

275
Q

Why is solid sugar soluble in water?

A

because of the O-H polar bond

276
Q

How Substances Dissolve

A
  • 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
Q

like dissolves __

A

like

278
Q

polar molecules dissolve in __

A

polar solvents

279
Q 
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

a) CCl4 mixed with water or H2O

280
Q

Saturated solution

A

contains as much solute as will dissolve at that temperature

281
Q

Unsaturated solution

A

is a solution that has not reached the limit of solute that will dissolve in it

282
Q

Relatively large amount of solute is dissolved in a __ solution

A

concentrated

283
Q

Relatively small amount of solute is dissolved in a __ solution

A

dilute

284
Q

mass percent

A

mass percent = mass of solute ÷ mass of solution x 100%

mass percent = grams of solute ÷ (grams of solute + grams of solvent) x 100%

285
Q

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?

A

6.6%

286
Q

calculating molarity

A
  • Molarity, M, equals the number of moles of solute per volume of solution in liters
    M = moles of solute / liters of solution
287
Q

you have 1.00 mol of sugar in 125.0 mL of solution. Calculate the concentration in units of molarity.

A

8.00 M

288
Q

concentration of ions

A
  • 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
Q

Standard Solution

A

A solution whose concentration is accurately known

- curve is straight up

290
Q

Steps to Make a Standard Solution

A
  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
Q

dilution

A

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
Q

Steps to Dilute a Solution

A
  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
Q

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?

A

60.0 mL

294
Q

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?

A

168 mL

* 155 mL water added (323 - 168 = 155)

295
Q

Steps for Solving Stoichiometric Problems Involving Solutions

A
  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
Q
  1. 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?
A
  • lead(II) phosphate, Pb3(PO4)2

- 1.1 g Pb3(PO4)

297
Q
  1. 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?
A

0.27 M

298
Q
  1. 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?
A

0.011 M

299
Q

dipole moment

A

molecules w polar bonds often behave in an electric field as if they had a center of positive charge
(+) end attracts (-) end

300
Q

non polar

A

equal sharing of e- (diatomics)

301
Q

polar

A

non-equal sharing of e- (HCl)

302
Q

Conjugate base

A

everything that remains of the acid molecule after a proton is lost

303
Q

Conjugate acid

A

formed when a proton is transferred to the base

304
Q

Conjugate acid–base pair

A

Consists of two substances related to each other by the donating and accepting of a single proton

305
Q

acids in water

A

Water acts as a base accepting a proton from the acid

- Forms hydronium ion (H3O+)

306
Q 
Which of the following pairs represents a conjugate acid–base pair? 
a) HCl, HNO3
b )H3O+, OH–
c) H2SO4, SO42–
d) HCN, CN–
A

d) HCN, CN–

307
Q

strong acid

A
  • Completely ionized or completely dissociated HA(aq) + H2O(l) → H3O+(aq) + A–(aq)
  • breaks apart 100%
  • forward rxn predominates
308
Q

weak acid

A
  • 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
Q

Strong acids contain relatively __ conjugate bases

A

weak

310
Q

weak acids have a __ base

A

strong

311
Q 
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

A-, H2O, Cl-

312
Q

water is _

A

amphoteric

- can be both an acid (H3O+) and a base (OH-)

313
Q

at 25ºC, Kw =

A

Kw = [H+] [OH-] = 1.0 x 10^-14

*no matter what the solution contains

314
Q

[H+] = [OH–]

A

neutral solution

315
Q

H+] > [OH–]

A

acidic solution

316
Q

[H+] < [OH–]

A

basic solution

317
Q

in an acidic, basic, or neutral solution Kw always equals

A

Kw = [H+][OH–] = 1.0 × 10^-14

318
Q

an acid/base reaction is also called a

A

neutralization reaction

319
Q

pH =

A

− log [H+]

320
Q

as [H+] increases, pH __

A

decreases

321
Q

pH range

A

7 is neutral
> 7 is basic
< 7 is acidic
*The lower the pH, them ore acidic the solution

322
Q

pH scale

A

scale is from 0-14

  • acidic: 0-6.9
  • neutral: 7
  • basic: 7.1-14
323
Q

pOH =

A
  • log [OH-]
324
Q

[H+] =

A

10^-pH

325
Q

[OH-] =

A

10^-pOH

326
Q

pH + pOH =

A

14

327
Q

[H+] [OH-] =

A

1.0 x 10^ -14

328
Q

calculate the pH for 1.0×10^-4 M H+

A

4.00

pH = –log[H+] = –log(1.0 × 10–4 M) = 4.00

329
Q

calculate the pH for 0.040 M OH–

A

12.60

330
Q

The pH of a solution is 5.85. What is the [H+] for this solution?

A

[H+] = 1.4 × 10^-6 M

331
Q

calculate the pOH for 0.040 M OH–

A

1.40

332
Q

The pH of a solution is 5.85. What is the [OH–] for this solution?

A

7.1 × 10^-9 M

333
Q

To convert from moles to atoms, ___

A

multiply by Avogadro’s number

334
Q

To convert from atoms to moles, __

A

divide by Avogadro’s number (or multiply by its reciprocal)

335
Q

To convert from grams to moles, __

A

divide using molar mass

336
Q

to covert from moles to grams, __

A

multiply by molar mass

337
Q

HA (aq) + H2O (l) -> H3O+ (aq) + A- (aq)

label acid, base, conjugate acid, and conjugate base

A

acid: HA
base: H2O
conjugate acid: H3O+
conjugate base: A-

338
Q

group 1 element charges

+1

A 
alkali metals
Highly Nasty Kids Rub Cats 
Li (lithium)
Na (sodium)
K (potassium)
Rb (rubidium)
Cs (cesium)
339
Q

group 2 element charges

+2

A 
alkaline earth metals
Beer Mugs Can Serve Bar Rats
beryllium (Be)
magnesium (Mg)
calcium (Ca)
strontium (Sr)
barium (Ba)
radium (Ra)
340
Q

group 3 element charges

+3

A 
Bears Always Give In
boron (B)
aluminum (Al)
gallium (Ga)
indium (In)
341
Q

group 6 element charges

-2

A 
Old Soldiers Seem Tense
oxygen (O)
sulfur (S)
seaborgium (Se)
tellurium (Te)
342
Q

group 7 element charges

-1

A 
halogens 
“Floor Cleaner Broken?” I Asked
fluorine (F)
chlorine (Cl)
bromine (Br)
iodine (I)
343
Q

8 strong bases

A

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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