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chemistry > regents concepts and skills > Flashcards

regents concepts and skills Flashcards

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AP® Chemistry flashcards Periodic Table flashcards
1
Q

matter is classified as

A

a pure substance or as a mixture of substances

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2
Q

three phases of matter

A

solid, liquid, gas

structure shows phase

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3
Q

a pure substance

A

has a constant composition and constant properties

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4
Q

elements cannot be

A

broken down by chemical change

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5
Q

mixtures can be

A

separated by physical means

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6
Q

when different substances are mixed

A

a homogenous or heterogeneous mixture is formed

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7
Q

proportions of a mixture can be varied, but

A

each component retains its original properties

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8
Q

different properties of substances…

A

permit physical separation of the components

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9
Q

forms of energy

A

chemical, electrical, electromagnetic, thermal, mechanical, nuclear

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10
Q

physical change results in

A

rearrangement of existing particles in a substance with the same properties

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11
Q

chemical change results in

A

formation of different substances with changed properties

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12
Q

the modern model of the atom

A

has evolved over a long period of time through the work of many scientists

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13
Q

each atom has a

A

nucleus with an overall positive charge surrounded by one or more negatively charged electrons

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14
Q

subatomic particles in the nucleus are

A

neutrons (n) and protons (+)

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15
Q

proton

A

positive

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16
Q

neutron

A

neutral

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17
Q

electron

A

negative

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18
Q

number of protons equals

A

number of electrons in an atom

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19
Q

protons and neutrons are

A

about equal in mass

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20
Q

electrons are (mass)

A

smaller than protons/neutrons

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21
Q

isotopes

A

same number of protons, but different number of neutrons

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22
Q

atomic number

A

number of protons

identifies element

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23
Q

mass number

A

sum of protons and neutrons

identifies an isotope

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24
Q

when an atom gains electrons

A

it becomes negative

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25
Q

when an atom loses electrons

A

it becomes positive

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26
Q

chemical compound can be represented by

A

a specific chemical formula and assigned name based on the IUPAC system

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27
Q

types of chemical formulas

A

empirical
molecular
structural

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28
Q

in all chemical reactions there is a conservation of

A

mass, energy, and charge

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29
Q

types of reactions include

A

synthesis, decomposition, single-replacement, double-replacement

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30
Q

the mass of each proton and each neutron

A

is appx. equal to one amu (atomic mass unit)

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31
Q

mass of an atom is very nearly qual to

A

its mass number

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32
Q

the average atomic mass of an element is

A

the weighted average of the masses of naturally occurring isotopes

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33
Q

empirical formula is

A

the simplest whole number ratio of atoms in an element

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34
Q

empirical formula is different from

A

the molecular formula (actual ratio of atoms in a molecule of that compound)

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35
Q

formula mass of a substance

A

sum of the atomic masses of it’s atoms

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36
Q

molar mass (gram formula mass)

A

the mass of one mole of that substance

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37
Q

one mole

A

avogadro’s number

6.02 x 10^23

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38
Q

percent composition by mass of each element in a compound

A

can be calculated mathematically

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39
Q

a balanced chemical equation represents

A

conservation of atoms

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40
Q

coefficients in a balanced chem equation can be used to determine

A

mole ratios in reaction

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41
Q

heat is a transfer of energy from

A

a body of higher temp to a body of lower temp

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42
Q

thermal energy is associated with

A

the random motion of atoms and molecules

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43
Q

chemical and physical reactions can be

A

exothermic or endothermic

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44
Q

energy released or absorbed is equal to

A

the difference between potential energy of products and potential energy of reactants

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45
Q

energy released or absorbed by a chemical reaction can be represented by

A

a potential energy diagram

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46
Q

entropy

A

a measure of randomness or disorder of a system

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47
Q

greater entropy

A

greater disorder

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48
Q

systems in nature tend to undergo changes toward

A

lower energy, higher entropy

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49
Q

temperature is the measurement of

A

average kinetic energy of the particles in a gas

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50
Q

temperature is not

A

a form of energy

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51
Q

the concept of an ideal gas

A

a model to explain the behavior of gases

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52
Q

a real gas is most like an ideal gas when

A

the real gas is at low pressure and high temperature

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53
Q

kinetic molecular theory (KMT) for an ideal gas is that the particles are/have

A

in random, constant straight-line motion
separated by great distances relative to size
no attractive forces between them
collisions that may result in the transfer of energy between particles (but total energy of system remains the same)

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54
Q

kinetic molar theory describes

A

relationships of pressure, volume, temperature, velocity, and frequency

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55
Q

equal volumes of different gases at the same temp and pressure contain

A

equal numbers of particles

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56
Q

concepts of kinetic and potential energy can be used to explain

A

processes such as fusion (melting), solidification (freezing), vaporization (boiling, evaporation), condensation, sublimation, deposition

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57
Q

stability of an isotope

A

ratio of neutrons and protons in nucleus

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58
Q

most nuclei are stable, but

A

some are unstable and decay spontaneously, editing radiation

59
Q

each radioactive isotope has

A

a specific mode and rate of decay

60
Q

transmutation

A

a change in the nucleus of an atom that converts the atom to another

61
Q

transmutation can occur

A

naturally or can be induced by bombarding the nucleus with high-energy particles

62
Q

spontaneous decay can involve the release of

A

alpha particles
beta particles
positrons
gamma radiation

63
Q

different emissions differ in

A

mass, charge, ionizing power, penetrating power

64
Q

nuclear reactions include

A

natural transmutation
artificial transmutation
fission
fusion

65
Q

both benefits and risks are associated with

A

fission and fusion reactions

66
Q

nuclear reaction can be represented by equations that include

A

symbols that represent atomic nuclei
subatomic particles
emissions such as alpha, beta, gamma

67
Q

the energy released in a nuclear reaction (fission or fusion) comes from

A

the fractional amount of mass converted into energy

68
Q

the energy released during nuclear reaction is much greater than

A

the energy released during chemical reactions

69
Q

risks with the use of radioactive isotopes

A

biological exposure
longterm storage and disposal
mutations
nuclear accidents

70
Q

benefits of radioactive isotopes (in research, medicine, and industry)

A 
radioactive dating
tracing chemical/biological processes
industrial measurements
nuclear power
detection and treatment of disease (such as cancer)
71
Q

each electron in an atom has its own

A

distinct amount of energy

72
Q

when an electron gains a specific amount of energy it is in the

A

excited state

73
Q

electron going from higher to lower energy state

A

energy is emitted (can identify the element)

74
Q

electron orbitals

A

ground state

most probable electron location

75
Q

valence electrons

A

the outer most electrons

76
Q

valence electrons affect

A

chemical properties of the element

77
Q

lewis dot structure

A

shows valence electrons

78
Q

placement on periodic table

A

indicate physical and chemical properties

79
Q

elements on table are arranged in

A

increasing atomic number

80
Q

elements can be classified by their properties, such as

A

metals, nonmetals, metalloids (B, Si, Ge, As, Sb), noble gases

81
Q

elements can be differentiated by physical properties, such as

A

density, conductivity, malleability, solubility, and hardness

82
Q

elements can be differentiated by heir chemical properties by

A

how they act in a chemical reaction

83
Q

groups 1, 2, and 13-18

A

each group has same # of valence electrons

have similar properties, but not helium

84
Q

as the elements go down in groups

A

atomic radius increases
electronegativity decreases
first ionization energy decreases
metallic properties decrease

85
Q

when atom gains 1+ electrons

A

radius increases

becomes more negative

86
Q

when an atom loses 1+ electrons

A

radius decreases

becomes more positive

87
Q

degree of polarity is determined by

A

electronegativity difference between two bonded atoms

88
Q

when a bond is broken

A

energy is absorbed

89
Q

when a bond is formed

A

energy is released

90
Q

atoms attain a stable valence electron configuration by

A

bonding with another atom

91
Q

noble gases have

A

stable valence electron configurations and generally do not bond

92
Q

major categories of compounds

A

ionic

molecular

93
Q

chemical bonds are formed when valence electrons

A

transfer from one atom to another (ionic)
share between atoms (covalent)
mobile within a metal (metallic)

94
Q

electronegativity indicates

A

how strongly an atom attracts electrons in a chemical bond

95
Q

electronegativity values are assigned

A

according to arbitrary scales

96
Q

metals tend to react with

A

nonmetals to form ionic

97
Q

nonmetals tend to react with

A

other nonmetals to form molecular (covalent) compounds

98
Q

ionic compounds containing polyatomic ions have

A

both ionic and covalent bonding

99
Q

intermolecular forces created by an unequal distribution of charge result in

A

varying degrees of attraction between molecules

100
Q

the polarity of a molecule is determined by the

A

shape of the molecule and the distribution of charge around

101
Q

symmetrical molecules

A

nonpolar

102
Q

asymmetrical molecules

A

polar

103
Q

diatomics

A

H, N, F, O, I, Cl, Br

104
Q

hydrogen bonding is

A

a strong intermolecular force

105
Q

organic compounds contain

A

carbon atoms that bond to one another

106
Q

how to name compounds

A

IUPAC

107
Q

hydrocarbons

A

hydrogen and carbon only

108
Q

saturated hydrocarbons

A

only single carbon-carbon bonds

109
Q

unsaturated hydrocarbons

A

at least one multiple carbon-carbon bond

110
Q

categories of organic compounds

A 
organic acid
alcohol
esters
aldehydes
ketones
ethers
halides
amines
amides
amino acid
111
Q

types of organic reactions

A 
addition
substitution
polymerization
esterification
fermentation
saponification
combustion
112
Q

isomers of organic acids

A

have the same molecular formula, but different structures and properties

113
Q

a solution is a

A

homogenous mixture

114
Q

solubility is dependent on

A

temperature, pressure, chemical natures of solute and solvent

115
Q

concentration can be expressed as

A

molarity
percent by mass
parts per million

116
Q

salt

A

freezing point decreases

boiling point increases

117
Q

greater concentration

A

the greater the effect

118
Q

electrolyte

A

releases mobile ions in aq solution

allows solution to conduct electricity

119
Q

rate of chemical reaction depends on

A 
temp
concentration
nature of reactants
surface area
presence of a catalyst
120
Q

Le Châtelier’s principle is used to

A

predict effects of stress on a system at equilibrium

121
Q

catalyst provides

A

alternative reaction pathway that has lower activation energy

122
Q

collision theory

A

a reaction is most likely to occur if reactants of particles collide w/ proper energy and orientation

123
Q

some chemical and physical changes can reach

A

equilibrium

124
Q

at equilibrium the rate of forward reaction equals

A

the rate of reverse reaction

125
Q

the measurable quantities at equilibrium

A

remain constant

126
Q

acidity or alkalinity are measured by

A

pH level

127
Q

on the pH scale, each decrease of one unit of pH represents

A

a tenfold increase in hydronium ion concentration

128
Q

in the process of neutralization, an arrhenius base and an arrhenius aside react to form

A

salt and water

129
Q

arrhenius bases yield

A

hydroxide ions as the only negative ions in an aq solution

130
Q

arrhenius acids yield

A

hydrogen ions as the only positive ions in an aq solution

131
Q

arrhenius bases and acids are

A

electrolytes

132
Q

titration

A

lab process to determine concentration of an unknown

133
Q

voltaic cells

A

spontaneously coverts chem. energy to elect. energy

134
Q

electrolytic cells

A

require elec. energy to produce chem. change

electrolysis

135
Q

redox reaction involves

A

transfer of electrons

136
Q

reduction

A

gain of electrons

137
Q

half-reaction can be written

A

to represent reduction

to represent oxidation

138
Q

oxidation

A

loss of electrons

139
Q

in a redox reaction

A

of electrons lost = # of electrons gained

140
Q

oxidation states can be assigned to

A

atoms or ions

141
Q

change in oxidation # indicates

A

that oxidation or reduction has occurred

142
Q

an electrochemical cell can be

A

voltaic of electrolytic

143
Q

in electrochemical cell

A

oxidation occurs at the anode

reduction occurs at the cathode

chemistry (13 decks)

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