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CHE 105 Chemistry > Final Exam Review > Flashcards

Final Exam Review Flashcards

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1
Q

composition

A

what something is made of

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2
Q

structure

A

what something is made of and how the components are arranged

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3
Q

pure substances

A

elements and compounds

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4
Q

atom

A

the fundamental unit of matter

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5
Q

element

A

made of only one type of atom

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6
Q

compounds

A

composed of more than one element, bound in fixed ratios

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7
Q

molecules

A

groups of atoms that bind tightly together and behave as a single unit

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8
Q

mixture

A

contain more than one substance but not bound in a fixed ratio

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9
Q

homogeneous mixture

A

compounds mix evenly
ex. air

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10
Q

heterogeneous mixture

A

components do not mix evenly
ex. sand separates from water; cookie dough

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11
Q

solid

A

definite shape and volume

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12
Q

liquid

A

definite volume but no definite shape

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13
Q

gas

A

no definite shape or volume

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14
Q

chemical properties

A

can NOT be measured without changing the identity of the substance

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15
Q

chemical changes

A

chemical reactions; change the identity of the substance

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16
Q

physical properties

A

can be measured without changing the identity of the substance

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17
Q

physical changes

A

don’t change the identity of the substance

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18
Q

exothermic change

A

releases heat energy

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19
Q

endothermic change

A

absorbs heat energy

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20
Q

scientific method

A

make an observation, formulate new ideas, test ideas with experiments

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21
Q

hypothesis

A

a tentative explanation that has not been tested

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22
Q

theory

A

how or why it happens (supported by experimental evidence)

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23
Q

law

A

what happens (statement that describes observations that are true)

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24
Q

scientific notation moving decimal right

A

negative exponent

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25
Q

scientific notation moving decimal left

A

positive exponent

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26
Q

scientific notation multiplication

A

multiply coefficients and add the exponents

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27
Q

scientific notation division

A

divide coefficients and subtract exponents

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28
Q

accuracy vs. precision

A

accuracy: how reliable are the measurements

precision: how closely are the measurements made / how finely are the measurements

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29
Q

significant digits

A

-nonzeros
-zeros between nonzeros
-zeros after the decimal point

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30
Q

not significant digits

A

zeros to the left of all nonzeros

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31
Q

how to convert between units

A

1) write out the two conversion factors
2) set up the problem to cancel units

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32
Q

micro

A

10^-6

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33
Q

milli

A

10^-3

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34
Q

centi

A

10^-2

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35
Q

kilo

A

10^3

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36
Q

1 mL =

A

1 cm^3

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37
Q

1 L =

A

1 dm^3

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38
Q

1 m^3 =

A

1,000 dm^3

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39
Q

boiling and freezing points of water in Celsius

A

100 and 0

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40
Q

density =

A

mass / volume

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41
Q

Democritus

A

first proposed atoms

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42
Q

Lavoisier

A

Law of Conservation of Mass

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43
Q

Dalton

A

atomic theory

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44
Q

Mendeleev

A

periodic table

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45
Q

Thomson

A

discovered the electron

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46
Q

Rutherford

A

showed that atoms are mostly empty space with a small, dense nucleus

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47
Q

Bohr

A

proposed the Bohr model

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48
Q

Law of Conservation of Mass

A

in chemical reactions, matter is neither created or destroyed

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49
Q

proton mass, charge, and location

A

m: 1.0073
c: +1
l: nucleus

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50
Q

neutron mass, charge, and location

A

m: 1.0087
c: no charge
l: nucleus

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51
Q

electron mass, charge, and location

A

m: 0.0005
c: -1
l: around the nucleus

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52
Q

atomic number

A

-number of protons in an atom
-number of electrons in a neutral atom

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53
Q

mass number

A

the number of protons and neutrons

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54
Q

isotopes

A

have the same atomic number but a different mass

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55
Q

writing atomic symbols

A

top left: mass number = p+n
bottom left: atomic number = p

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56
Q

how to calculate average atomic mass

A

(value A x fraction A) + (value B x fraction B)

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57
Q

group 1A (column 1)

A

alkali metals

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58
Q

group 2A (column 2)

A

alkaline earth metals

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59
Q

group 7A (column 17)

A

halogens

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60
Q

group 8 A (column 18)

A

noble gases

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61
Q

transition elements

A

columns 3-12

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62
Q

pure metals

A

columns 1-2

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63
Q

nonmetals

A

upper right side (plus H)

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64
Q

metalloids

A

stairstep between metals and nonmetals

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65
Q

inner transition elements

A

two rows on the bottom of the periodic table

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66
Q

line spectra of atoms

A

occurs when gases are passed through a prism; each element produces a unique line spectrum

white light passed through a prism results in the full spectrum of the rainbow

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67
Q

how do atoms absorb or emit light energy

A

by electrons rising or falling to different energy levels

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68
Q

how many electrons fit in level 1

A

2

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69
Q

how many electrons fit in level 2

A

8

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70
Q

how many electrons fit in level 3

A

18

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71
Q

how many electrons fit in level 4

A

32

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72
Q

quantum mechanics

A

describes electrons most probable locations and energies

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73
Q

s subshell

A

1 orbital
2 electrons

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74
Q

p subshell

A

3 orbitals
6 electrons

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75
Q

d subshell

A

5 orbitals
10 electrons

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76
Q

f subshell

A

7 orbitals
14 electrons

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77
Q

electron capacity =

A

of orbitals x 2

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78
Q

significant figures addition and subtraction

A

round to the last decimal place of the least precise starting measurement

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79
Q

significant figures multiplying and dividing

A

same number of digits that are in the least precise measurement

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80
Q

E = ?

A

Planck’s constant x speed of light / wavelength

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81
Q

how to determine the number of protons and electrons in an atom using the periodic table

A

atomic number

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82
Q

how to determine the number of valence electrons using the periodic table

A

group number (group = vertical)

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83
Q

electron configuration and the periodic table:
1. row indicates…
2. column gives…

A
  1. row: highest occupied electron energy level
  2. column: outermost electron configuration
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84
Q

cation

A

positive charge

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85
Q

anion

A

negative charge

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86
Q

isoelectronic

A

same number of electrons

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87
Q

s, p, d, f blocks
sdpf

A

s: column 1-2
p: everything right of the transition metals
d: transitions metals
f: lower periods

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88
Q

ionic bonds occur between…

A

oppositely charged ions

positive charges must equal the negative charges

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89
Q

how are ionic compounds named

A

cation anion
ex. sodium chloride, magnesium sulfate, copper (II) chloride

90
Q

ionic lattice

A

an array of positive and negative ions

91
Q

ionic compounds

A
  1. oppositely-charged ions
  2. don’t form molecules
  3. formula unit or empirical formula
  4. ionic lattice
92
Q

covalent compounds

A
  1. share electrons
  2. between nonmetal atoms
  3. usually form molecules
  4. molecular formula
93
Q

naming binary covalent compounds

A

the atom to the left is named first…the atom to the right is named as the anion

94
Q

binary covalent compounds prefixes

A

1 - mono
2 - di
3 - tri
4 - tetra
5 - penta

95
Q

electrolyte solutions

A

ionic compounds dissolved in water

96
Q

acids

A

covalent compounds that produce H+ ions in aqueous solution

97
Q

decomposition

A

one forms two or more
2H2O -> 2H2 + O2

98
Q

single displacement

A

one element replaces another
Zn + CuCl2 -> Cu
usually a metal involved

99
Q

synthesis

A

two form one
H2 + Cl2 -> 2HCl

100
Q

double displacement

A

two ions replace each other
(the anions swap positions)

101
Q

metal + nonmetal

A

ionic compound
(metal cation + nonmetal anion)

102
Q

oxidation

A

loss of electrons
(getting more positive)

103
Q

reduction

A

gain of electrons
(if the element gains an electron, it is reduced; negative = reduced)

104
Q

combustion reactions

A

gas combines with elements or compounds to produce oxides

ex. Sn + O2 -> SnO2
tin (IV) oxide - ionic
ex. C+ O2 -> CO2
carbon dioxide - covalent

105
Q

combustion of hydrocarbons

A

hydrocarbon + oxygen -> carbon dioxide + water
(must contain oxygen and will always create carbon dioxide and water)

106
Q

molecular equation reaction in aqueous solutions

A

shows ions together as compounds
KBr (s) -> KBr (aq)

107
Q

ionic equation reaction in aqueous solutions

A

shows dissociated ions as separate species; ionic compounds dissociate when dissolved in water
ex. KBr (s) -> K+ (aq) + Br- (aq)

108
Q

precipitation reaction

A

two aqueous solutions produce an insoluble product known as the precipitate (double displacement)

109
Q

acid base neutralization

A

acid + base -> water + salt
ex. H+ (aq) + OH- (aq) -> H2O (l)

110
Q

oxidation-reduction reaction

A

electron transfer between metal and nonmetal

111
Q

complete ionic equation

A

shows all ions present

112
Q

net ionic equation

A

omit spectator ions; only shows ions that react

113
Q

how to calculate formula mass of a compound

A

add the atomic masses of each atom in the formula unit

114
Q

percent composition of one element

A

(mass of one element / mass of compound) x 100%

115
Q

conversion factor between grams and moles

A

molar mass

116
Q

conversion factor between moles and particle/atoms

A

Avogadro’s number
6.022x10^23

117
Q

conversion factor between the moles of two different substances

A

mole ratio from balanced equation

118
Q

limiting reagent

A

completely consumed; the reagent that forms the least amount of product

119
Q

excess reagent

A

not completely consumed; reagent will be left over after the reaction is complete

120
Q

theoretical yield

A

the amount of a product that can form, based on the balanced equation

121
Q

actual yield

A

the amount actually obtained

122
Q

percent yield =

A

(actual yield / theoretical yield) x 100%

PAT x 100%

123
Q

percent yield

A

the percentage of the theoretical yield that was obtained

124
Q

chemical and physical occurrences that can lead to an actual yield that is less than the theoretical yield

A

-spillage
-contamination
-material sticks to container wall
-unwanted side products
-product lost during purification

125
Q

heat

A

transfer of kinetic energy

126
Q

work

A

transfer from one form to another

127
Q

delta E =

A

q + w
heat energy + work

128
Q

endothermic change

A

-energy on the reactants side
-absorbs heat energy

129
Q

exothermic change

A

-energy on the products side
-releases heat energy

130
Q

system does work on surroundings

A

releases energy (-q)

131
Q

surroundings does work on system

A

absorbs energy (+q)

132
Q

delta E system =

A

-delta E surroundings

133
Q

specific heat

A

the amount of heat required to raise the temp of 1g of material by 1 degree C

134
Q

heat

A

total kinetic energy transferred from one substance to another

135
Q

temperature

A

the average kinetic energy of the particles in a substance

136
Q

specific heat equation

A

q = (m)(s)(delta T)

137
Q

coffee cup calorimetry

A

q water = -q metal
q = (m)(s)(delta T)

138
Q

(mw)(sw)(delta Tw) =

A

-(mm)(sm)(delta Tm)

139
Q

bomb calorimetry

A

measures energy content (heats of reaction) in food and fuels

140
Q

bomb calorimetry equation

A

q = (c)(delta T)
c = heat capacity

141
Q

heat problems using the reaction enthalpy

A

stoichiometry with the delta H value

142
Q

drawing Lewis structures

A

1) add up all valence electrons
2) frame the structure (leftmost = central atom)
3) fill octets on outer atoms first
4) fill the octet on the central atom (any remaining e- on central atom; use double/triple bonds if needed)

143
Q

octet rule

A

8 electrons in the valence shell

144
Q

formal charge

A

FC = (valence e-) - (# of covalent bonds) - (# of unshared e-)

145
Q

relating the formal charge on each atom to the overall charge of the ion or molecule

A

a negative overall charge can come from a negative formal charge

146
Q

two electron sets

A

linear

147
Q

three electron sets

A

trigonal planar

148
Q

four electron sets

A

tetrahedral

149
Q

2 bonded groups, 0 lone pairs

A

linear

150
Q

2 bonded groups, 1 lone pair

A

bent

151
Q

2 bonded groups, 2 lone pairs

A

bent

152
Q

3 bonded groups, 0 lone pairs

A

trigonal planar

153
Q

3 bonded groups, 1 lone pair

A

trigonal pyramidal

154
Q

4 bonded groups, 0 lone pairs

A

tetrahedral

155
Q

using arrows to show polar covalent bonds

A

arrow points towards the most electronegative atom

delta negative near more electronegative atom

156
Q

electronegativity of covalent

A

less than 0.5

157
Q

electronegativity of polar covalent

A

0.5 to 2.0

158
Q

electronegativity of ionic

A

greater than 2.0

159
Q

molecular dipole

A

an overall polarity in a molecule

160
Q

identifying molecules with net dipoles

A

1) analyze the electronegativity of the individual bonds within the molecule
2) determine if the individual bond dipoles cancel each other out
3) if they do not cancel out, the molecule has a net dipole and is considered polar

161
Q

predicting the direction of the dipoles

A

always point the arrow towards the more electronegative atom in a bond

162
Q

behavior of particles in solids

A

particles are close together and held in a fixed place

163
Q

behavior of particles in liquids

A

particles are close together but move freely past each other

164
Q

behavior of particles in gases

A

particles are far apart and have very little interaction

165
Q

properties of ionic substances (interactions, melting points, boiling points)

A
  • high MP
  • high BP
  • lattice
  • brittle
  • poor EC
166
Q

properties of metallic substances (interactions, melting points, boiling points)

A
  • lattices
  • electrons move easily between atoms
  • malleable and ductile
  • high MP
  • high BP
  • high EC
  • metallic luster
167
Q

properties of molecular substances (interactions, melting points, boiling points)

A
  • covalent bond
  • low MP
  • poor EC
  • flammable
  • variable solubility
168
Q

properties of network covalent substances (interactions, melting points, boiling points)

A

covalent networks: lattices of covalent bonds that form giant molecules

  • high MP
  • high BP
  • high hardness
  • poor EC
  • insoluble
  • brittle
169
Q

hydrogen bonding

A

strong intermolecular force between molecules containing H-F, H-O, or H-N bonds

170
Q

dipole-dipole forces

A

attractions between polar covalent molecules

171
Q

London dispersion forces

A

weak intermolecular forces that result from instantaneous dipoles

must have only covalent bonds to be dispersion forces

172
Q

instantaneous dipoles

A

not a permanent dipole

173
Q

describing the types of forces present based on a formula or Lewis structure

A

Hydrogen Bonding: H-F, H-O, or H-N bonds

Dipole-Dipole: molecules with a net dipole

London Dispersion: all covalent molecules

174
Q

qualitatively describing the behaviors of gases

A
  • particles are spaced far apart
  • particles have no attraction for each other
175
Q

describing the relationships between P, V, T, and n using the combined gas laws and the ideal gas laws

A

PV = nRT

P: pressure
V: volume
n: number of moles
T: temp in K
R: 0.0821

176
Q

solving stoichiometry problems involving the gas laws

A

PV = nRT is the conversion factor

177
Q

mass %

A

(mass of solute / mass of solution) x 100%

% (m/m)

178
Q

volume %

A

(volume of solute / volume of solution) x 100%

% (v/v)

179
Q

mass / volume %

A

(mass of solute / volume of solution) x 100%

% (m/v)

180
Q

solute

A

the substance that dissolves

181
Q

solvent

A

the major component of the solution

182
Q

parts per million

A

(mass of solute / mass of solution) x 10^6

183
Q

parts per billion

A

(mass of solute / mass of solution) x 10^9

184
Q

molarity

A

moles per liter

185
Q

M =

A

moles / V

moles solute / liter solution

186
Q

moles (molarity equation) =

A

MV

187
Q

(Mi) (Vi) =

A

(Mf) (Vf)

188
Q

preparing dilute solutions

A

1) add the desired amount of solute to a volumetric flask
2) partially fill the flask, and mix
3) dilute to the correct volume

189
Q

ionic compounds __________ in water

A

dissociate

190
Q

ionic equation

A

shows dissociated ions as separate species

191
Q

molar concentration of each ion in a solution

A

coefficient x molarity

192
Q

three colligative properties

A

1) freezing point depression
2) boiling point elevation
3) osmotic pressure

193
Q

freezing point depression

A

increase solute concentration = lower freezing point

more moles = lower freezing point

194
Q

boiling point elevation

A

antifreeze raises boiling point and lowers freezing point

195
Q

osmotic pressure

A

water moves toward the more concentrated solution

196
Q

hypertonic

A

in solution of high solute concentration

197
Q

hypotonic

A

in solution of low solute concentration

198
Q

conversion factor between solution volume and moles

A

molarity

199
Q

chapter 11 volume questions

A

1) balanced equation
2) determine mole ratio
3) calculate moles of first solution
4) calculate moles of the second solution
5) calculate the volume of the second solution by (moles / Molarity)
6) convert the volume to mL

200
Q

acids Arrhenius definition

A

produce H+ ions in water

201
Q

bases Arrhenius definition

A

produce OH- ions in water

202
Q

acids Bronsted-Lowry definition

A

compounds that donate H+ ions

proton (H+) donor

203
Q

bases Bronsted-Lowry definition

A

compounds that accept H+ ions

proton (H+) acceptor

204
Q

strong acid

A

completely ionize in water
ex. HCl, HBr, HI, H2SO4

205
Q

weak acid

A

partially ionize in water
ex. HF + H2O <-> H3O + F

206
Q

a conjugate base is identical to the acid, but…

A

without the H+

207
Q

a conjugate acid is identical to the base, but with an…

A

extra H+

208
Q

acid + base =

A

water + salt

209
Q

reactions of acids with metals

A

most metals react with acid to produce metal cations and hydrogen gas

210
Q

nonmetal oxides often react with water to form…

A

acids

211
Q

[H+] [OH-] =

A

1x10^-14

212
Q

pH + pOH =

A

14

213
Q

pH =

A

-log [H+]

214
Q

pOH =

A

-log [OH-]

215
Q

atomic and mass numbers have to be _______ on both sides of the equation

A

the same

216
Q

three major types of radioactive decay

A

alpha, beta, gamma

217
Q

alpha decay

A

emission of a helium nucleus (add Helium on products side)

218
Q

beta decay

A
  • converts a neutron into a proton and electron
  • emission of an electron
219
Q

predicting the material remaining using half-lives

A

start -> divide by two -> divide previous number by two -> continue

220
Q

isotopes

A

have the same atomic numbers but different mass numbers

CHE 105 Chemistry (5 decks)

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