Chemistry Flashcards
1
Q
the study of the composition, structure, and properties of matter
A
Chemistry
2
Q
anything that occupies space and has mass
A
Matter
3
Q
the amount of matter in an object (expressed in g, kg, tons, pounds)
A
Mass
4
Q
the gravitational force on an object (expressed in Newton)
A
Weight
5
Q
formula for weight
A
Weight = mass x gravity (W=mg); gravity (Earth) = 9.8 m/s^2
6
Q
exact space occupied by an object; Impenetrability property: no two objects can occupy the same space at the same time
A
Volume
7
Q
the amount of matter in a given volume of material (expressed in g/cm^3 or kg/m^3)
A
Density
8
Q
formula for density
A
Density = mass/volume (D=m/V)
9
Q
matter is composed of particles called atoms
A
Atomic Theory
10
Q
the basic building blocks that makeup matter; smallest particle
A
Atom
11
Q
Atom is from a Greek word ______ that means _____
A
atomos = indivisible
12
Q
What are atoms made up of?
A
protons (+), electrons (-), neutrons (no charge)
13
Q
What subatomic particles are on the inside of the nucleus?
A
protons and neutrons
14
Q
What subatomic particle is residing around the nucleus?
A
electrons
15
Q
atomic number: number of ______
A
protons
16
Q
atomic model that stated the atom is hard and indestructible
A
Billiard Ball Model
17
Q
Who is the scientist behind Billiard Ball Model?
A
John Dalton
18
Q
this atomic model shows electrons embedded in a positively-charged sphere
A
Plum Pudding Model
19
Q
Who is the scientist behind Plum Pudding Model?
A
JJ Thompson
20
Q
In this atomic model, the mass and all of the positive charge of an atom is concentrated in the nucleus
A
Nuclear Model
21
Q
Who is the scientist behind the Nuclear model?
A
Ernest Rutherford
22
Q
In this atomic model, electrons travel around the nucleus in a circular orbit; their energy is proportional to their distance form the nucleus
A
Planetary Model
23
Q
Who is the scientist behind the Planetary model?
A
Niels Bohr
24
Q
In this atomic model, the electron is a wave; found in orbitals
A
Quantum Model
25
Who is the scientist behind the Quantum model?
several scientists but Erwin Schrodinger is the most relevant
26
form of matter with a constant composition
Pure substances
27
How can elements and compounds be pure substances?
because their composition is always fixed, and their constituents can only be separated using chemical or electrochemical reactions
28
simplest form of matter, cannot be broken down into smaller components
Elements
29
How many elements were currently known to man?
118 elements
30
Who created the Periodic Table of Elements, and when did he create it?
Dmitri Mendeleev in 1869
31
What is the most common and lightest element in the universe?
Hydrogen
32
What is the second most abundant element, although it is very rare on Earth?
Helium
33
What is the most common element on Earth by its mass?
Oxygen
34
What is the first man-made element?
Technetium
35
charged atom
Ion
36
negatively-charged; gain electrons; electrons>protons
Anion
37
positively-charged; loss of electrons; electrons
Cation
38
energy required to remove an electron from an atom
Ionization energy
39
energy released when an electron is added to an atom
Electron Affinity
40
ability to attract electrons
Electronegativity
41
lustrous, malleable (can be pounded into thin sheets), good conductor of heat and electricity
Metals
42
What are some of the metal elements?
aluminum, iron, gold, cooper, mercury, lead
43
poor conductor of heat and electricity, brittle
Non-metals
44
What are some of the non-metal elements
carbon, oxygen, hydrogen, phosphorus
45
have some characteristics of either metals or non-metals
Metalloid
46
What are some examples of metalloid elements?
silicone, boron, arsenic
47
highly reactive metals; have one excess electron which they tend to lose, thus they usually have a charge of +1; usually from compounds with halogens
Alkali metals
48
highly reactive non-metals; lack one electron on their outer shell which they try to acquire from other atoms, thus they usually have a charge of -1; usually forms compounds with alkali metals
Halogens
49
inert gases; unreactive, very stable elements owing to their full outer shell of eight electrons
Noble gases
50
elements to the right of metalloids (except H)
Non-metals
51
elements to the left of metalloids (except H
Metals
52
atoms of the same element that differ in the number of neutrons
Isotopes
53
describes how electrons are distributed in its atomic orbitals
Electron Configuration
54
the number of "excess" electrons in an atom
Valence electrons
55
every element has the same electronic configuration as the element before it in the periodic table, plus one extra
Aufbau principle
56
electrons tend to stay unpaired in orbitals with equal energies
Hund's rule
57
members of a family of an element that all have the same number of protons but different numbers of neutrons
Isotopes
58
two or more elements with atoms bonded together
Compounds
59
How can compounds be separated?
Chemical Methods
60
carbon-containing compounds (e.g. sugar, acetone, methane)
Organic compounds
61
compounds with no carbons (e.g. salt, lye, water)
Inorganic compounds
62
a characteristic sour taste
Acids
63
Acids litmus test
blue to red
64
acids pH level
less than 7; the lower the pH, the more acidic
65
What do acids yield?
hydrogen ions
66
has a litmus test of red to blue
bases
67
pH level of bases
between 7-14
68
What do bases yield?
hydroxide ions
69
no litmus test changes
neutrals
70
sodium and chlorine compound
table salt (NaCl)
71
one carbon atom and two oxygen atoms
Carbon dioxide (CO2)
72
6 carbon atoms, 12 hydrogen atoms, 6 oxygen atoms
Sugar/glucose (C6H12O6)
73
formed by a cation and an anion wherein there is a transfer of electrons so that each atom has the same number of electrons as the nearest noble gas (so it is more stable)
Ionic Compound
74
the compound formed by atoms that share electrons so that each atom has the same number of electrons as the nearest noble gas
Covalent Compound
75
two types of chemical bonding
Intramolecular and Intermolecular
76
attraction between atoms in a molecule
Intramolecular
77
2 types of intramolecular
Covalent bond and Ionic bond
78
attraction between molecules
Intermolecular
79
3 types of intermolecular
London dispersion forces, dipole-dipole forces, hydrogen bond
80
sharing of electrons between two non-metals
Covalent bond
81
transfer of electrons; between metals and non-metals/cations and anions
Ionic bond
82
between two temporarily polar molecules
London dispersion forces
83
or polar covalent molecules
Dipole-dipole forces
84
between the H atom of one molecule and an N,F, or O atom of another molecule
Hydrogen bond
85
VSEPR
Valence Shell Electron Pair Repulsion
86
shapes of molecules depend on the fact that electron pairs tend to move as far apart from one another as possible
VSEPR
87
having a partial positive and partial negative ends, or poles
Polar molecule
88
when a liquid is soluble in another liquid; the liquid present larger quantity is solvent
Miscible
89
combination of two pure substances
Mixtures
90
appears as one uniform material with uniform characteristics
Homogeneous mixture
91
universal solvent
water
92
the ______ the solvent, the more solute would be dissolves
warmer
93
two parts of a solution
solvent and solute
94
the part of a solution that dissolves
solvent
95
the part of a solution that is dissolved
solute
96
types of solutions
unsaturated, saturated, supersaturated
97
type of solution that can still dissolve more solute
unsaturated
98
type of solution that already contains the maximum amount of solute
saturated
99
type of solution that is used to pressure or heat to dissolve more than the usual amount of solute
supersaturated
100
does not have uniform particles, therefore the particles could be identified individually
Heterogeneous mixture
101
its particles are larger than those in a solution, so the particles often settle in the bottom of the solvent
suspension
102
in between being a homogeneous and heterogeneous mixture
Colloids
103
the scattering of light exhibited by a colloid mixture
The Tyndall Effect
104
Ideal Gas Law formula
pV = nRT
105
decreasing the volume increases collisions and increases pressure
Boyle's Law
106
Boyle's law formula
p1V1 = p2V2
107
increasing the temperature increases volume
Charles' Law
108
Charles' law formula
V1/T1 = V2/T2
109
increasing the temperature increases pressure
Gay-Lussac's Law
110
Gay-Lussac's law formula
p1/T1 = p2/T2
111
Combined Gas law formula
p1V1/T1 = p2V2/T2
112
Avogadro's law formula
V1/n1 = V2/n2
113
the number of units in one mole of any substance
Avogadro's number: 6.02214076 x 10^23
114
measure the dimensions then compute using formulas
Regular solids
115
water displacement
Irregular solids
116
use a graduated cylinder
Liquids
117
no definite volume, pressure is usually measured instead
Gases
118
particles in gas are infinitely small and are in constant random motion; gases do not experience intermolecular forces
The Kinetic Molecular Theory of Gases
119
according to the Kinetic Molecular Theory of Gases, gas molecules undergo to this as they bounce perfectly to one another
perfectly elastic collisions
120
the kinetic energies of gas molecules are directly proportional to their what?
temperatures (the hotter, the faster the molecules move)
121
Law of Thermodynamics where energy can neither be created nor destroyed; it can only change forms
1st Law (Conservation of Energy)
122
Law of Thermodynamics where entropy of an isolated system always increases; entropy, or S, is the degree of disorder
2nd Law
123
Law of Thermodynamics where entropy at -273.15K (absolute zero temperature) is zero; there can be no temperature lower than absolute zero
3rd Law
124
enthalpy is the heat transferred between system and surroundings at constant pressure
Hess' Law
125
the amount of energy needed to raise the temperature of a substance by one degree; in J/kg-C; decreases as mass and temperature to be attained increases
Specific heat
126
hotness or coldness
temperature
127
total kinetic energy of molecules
heat
128
temperature where liquid turns into gas; requires the addition of heat
Boiling point
129
boiling point of water
100 degrees Celsius
130
the temperature where solid becomes liquid; this change from solid to liquid requires heat
Melting Point
131
melting point of water
0 degrees Celsius
132
two or more compounds combine to form one compound
Combination reaction
133
the opposite of a combination reaction - a complex molecule breaks down to make simpler ones
Decomposition reaction
134
two solutions of soluble salts are mixed resulting in an insoluble solid forming
Precipitation reaction
135
A + B -> AB
combination reaction
136
AB -> A + B
decomposition reaction
137
A + Soluble salt B -> Precipitate + Soluble salt C
precipitation reaction
138
an acid and a base react with each other; generally, the product of this reaction is salt and water
Neutralization reaction
139
Acid + Base -> Salt + Water
neutralization reaction
140
oxygen combines with a compound to form carbon dioxide and water; these reactions are exothermic meaning they give off heat
combustion reaction
141
A + O2 -> H2O + CO2
combustion reaction
142
one element takes place with another element in the compound
displacement reaction
143
A + BC -> AC + B
displacement reaction
144
factors that increase rates of reactions
higher temperatures, higher concentration of reactants, larger surface area of reactants, addition of catalyst (substance that increases the rate of reaction without itself being consumed)
145
factors that affect solubility
temperature, surface area of solutes, pressure
146
mole fraction
moles solute/moles solution
147
mass percentage
mass solute/mass solute x solvent x 100% (not sure)
148
mass-volume percentage
(mass of solute/volume of solvent) x 100%
149
volume-volume percentage
volume of solute/volume of solvent x 100%
150
Molarity (M)
moles solute/volume of solution (L)
151
Molality (m)
moles solute/mass of solvent (kg)
