TEST 1 Flashcards
1
Q
1 cm3 =
A
1mL
2
Q
1mL=
A
1 cm3
3
Q
1 angstrom
A
10^-10 m
4
Q
10^-10 m
A
1 angstrom
5
Q
Conversion between Celsius and Kelvin
A
1 kelvin degree change = 1 degree change celsius BUT must add 273.15 when converting from celsius to kelvin. T(K)=T(C)+273.15
6
Q
qualitative vs quantatative
A
qualitative observations (color, hot or cold, large or small) and quantitative measurements
7
Q
SI unit length
A
meter: m
8
Q
SI unit mass
A
kilogram: kg
9
Q
SI unit time
A
second: s
10
Q
SI unit temperature
A
Kelvin: K
11
Q
SI unit amount of substance
A
Mole: mol
12
Q
mass
A
the measure of the quantity of matter within an object
13
Q
giga
A
G (10^9); billion
14
Q
Mega
A
M (10^6); million
15
Q
Kilo
A
k (10^3) thousand
16
Q
Deci
A
d (10^-1) tenth
17
Q
Centi
A
c (10^-2) one hundredth
18
Q
Milli
A
m (10^-3) one thousandth
19
Q
Micro
A
Mu (10^-6) one millionth
20
Q
Nano
A
n (10^-9) one billionth
21
Q
Pico
A
p (10^-12)
22
Q
Femto
A
f (10^-15)
23
Q
distance between O-H
A
0.94 Angstroms=9.4 x 10^-11 m
24
Q
area
A
m^2
25
speed
m/s
26
energy
Joules
27
1 J=
1 kgm^2/s^2
28
volume
L; m^3; 1 cm^3=1 mL
29
density
g/cm^3 = mass/volume
30
intensive property
property that does not depend on the quantity of matter (i.e. temp, pressure)
31
extensive property
property that does depend on the quantity of matter (i.e. mass, volume)
32
matter
anything that takes up space and has mass
33
weight
the force gravity exerts on an object; directly proportional to the mass of the object
34
law of conservation of matter
there is no detectable change in the total quantity of matter present when matter converts from one type to another (a chemical change) or changes among solid, liquid, or gaseous states (a physical change).
35
classifying matter
mixtures (composed of two or more types of matter that can be present in varying amounts and can be separated by physical changes; can be divided into homogenous (also called a solution, exhibits a uniform composition and appears visually the same throughout.) and heterogeneous (A mixture with a composition that varies from point to point)) vs pure substances (have a constant composition. All specimens of a pure substance have exactly the same makeup and properties; can be divided into elements (cannot be broken down by chem changes) vs compounds (can be broken down by chem changes))
36
atom
the smallest particle of an element that retains the properties of that element and can enter into a chemical combination; matter is made up of atoms
37
molecule
A molecule consists of two or more atoms joined by strong forces called chemical bonds. The atoms in a molecule move around as a unit; ring. A molecule may consist of two or more identical atoms or two or more different atoms. if the 2 atoms are identical, it is an element and if they are different, it is a compound. A molecule is the smallest unit of a molecular compound
38
diatomic molecules
hydrogen (H2), oxygen (O2), phosphorus (P4), sulfur (S8)
39
A physical property is
a characteristic of matter that is not associated with a change in its chemical composition. (i.e. density, color, hardness, melting and boiling points, and electrical conductivity.) Some physical properties, such as density and color, can be observed without changing the physical state of the matter.
40
A physical change is
a change in the state or properties of matter without any accompanying change in the chemical identities of the substances contained in the matter. (i.e. when wax melts, when sugar dissolves in coffee, and when steam condenses into liquid water (Figure 1.18). Other examples of physical changes include magnetizing and demagnetizing metals (as is done with common antitheft security tags) and grinding solids into powders (which can sometimes yield noticeable changes in color). In each of these examples, there is a change in the physical state, form, or properties of the substance, but no change in its chemical composition.)
41
chemical property
The change of one type of matter into another type (or the inability to change) (i.e. include flammability, toxicity, acidity, and many other types of reactivity.)
42
A chemical change
always produces one or more types of matter that differ from the matter present before the change. (i.e. formation of rust, an explosion of nitroglycerin, reactions that are performed in a lab (such as copper reacting with nitric acid), all forms of combustion (burning), and food being cooked, digested, or rotting)
43
When adding or subtracting numbers,
round the result to the same number of decimal places as the number with the least number of decimal places`
44
When multiplying or dividing numbers,
round the result to the same number of digits as the number with the least number of significant figures
45
If the digit to be dropped (the one immediately to the right of the digit to be retained) is less than 5,
“round down” and leave the retained digit unchanged;
46
If the digit to be dropped (the one immediately to the right of the digit to be retained) is more than 5,
“round up” and increase the retained digit by 1.
47
If the dropped digit is 5, and it’s either the last digit in the number or it’s followed only by zeros,
round up or down, whichever yields an even value for the retained digit. If any nonzero digits follow the dropped 5, round up.
48
precise vs accurate
Measurements are said to be precise if they yield very similar results when repeated in the same manner. A measurement is considered accurate if it yields a result that is very close to the true or accepted value. Precise values agree with each other; accurate values agree with a true value.
49
1 inch
2.54 cm
50
2.54 cm
1 inch
51
1 L
1000 mL
52
1000 mL
1 L
53
freezing/boiling points Celsius and Fahrenheit
On the Celsius scale, 0 °C is defined as the freezing temperature of the water and 100 °C as the boiling temperature of the water. The space between the two temperatures is divided into 100 equal intervals, which we call degrees. On the Fahrenheit scale, the freezing point of water is defined as 32 °F and the boiling temperature as 212 °F. The space between these two points on a Fahrenheit thermometer is divided into 180 equal parts (degrees).
54
dalton's atomic theory
1. Each element is composed of tiny, indestructible (F) particles called atoms.
2. All atoms of a given element have the same mass (F) and other properties that distinguish them from the atoms of other elements.
3. Atoms combine in simple, whole-number ratios to form compounds.
4. Atoms of one element cannot change into atoms of another element. (F) In a chemical reaction, atoms only change the way that they are bound together with other atoms.
55
radioactivity
a phenomenon in which an atom breaks down, provided evidence that atoms are made of smaller particles.
56
mass and charge comparison of protons, electrons and neutrons
* The mass of a neutron is similar to that of a proton.
| * The proton and electron have equal and opposite charge.
57
atomic number
number of protons; symbolized by Z; ever element has a unique atomic number
58
mass number
= # protons + # neutrons
59
isotopes
Atoms of the same element with a different number of neutrons; Isotopes have the same atomic number and different mass numbers
60
hydrogen isotope
PROTIUM: 11H; 1 proton, 0 neutrons
DEUTERIUM: 21H; 1 proton, 1 neutron
TRITIUM: 31H; 1 proton, 2 neutrons
61
the majority of an atom's mass is located in
The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume.
62
the majority of an atom's mass is located in
The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume.
63
ion
atom with unequal protons and electrons. Atomic charge = number of protons − number of electrons
64
anion
An atom that gains one or more electrons will exhibit a negative charge and is called an anion.
65
cation
Positively charged atoms called cations are formed when an atom loses one or more electrons.
66
cation
Positively charged atoms called cations are formed when an atom loses one or more electrons.
67
average mass
the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes.
The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance.
68
mass spectrometer
a sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms (or molecules) to become electrically charged, typically by losing one or more electrons. These cations then pass through a (variable) electric or magnetic field that deflects each cation’s path to an extent that depends on both its mass and charge (similar to how the path of a large steel ball-bearing rolling past a magnet is deflected to a lesser extent than that of a small steel BB). The ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios (a mass spectrum) is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio.
69
molecular formula
A molecular formula is a representation of a molecule that uses chemical symbols to indicate the types of atoms followed by subscripts to show the number of atoms of each type in the molecule. (A subscript is used only when more than one atom of a given type is present.) Molecular formulas are also used as abbreviations for the names of compounds. Also, a molecular formula is always a whole-number multiple of an empirical formula.
70
structural formula
The structural formula for a compound gives the same information as its molecular formula (the types and numbers of atoms in the molecule) but also shows how the atoms are connected in the molecule. A ball-and-stick model shows the geometric arrangement of the atoms with atomic sizes not to scale, and a space-filling model shows the relative sizes of the atoms.
71
Do a subscript following a symbol and a number in front of a symbol represent the same thing?
a subscript following a symbol and a number in front of a symbol do not represent the same thing; for example, H2 and 2H represent distinctly different species. H2 is a molecular formula; it represents a diatomic molecule of hydrogen, consisting of two atoms of the element that are chemically bonded together. The expression 2H, on the other hand, indicates two separate hydrogen atoms that are not combined as a unit. The expression 2H2 represents two molecules of diatomic hydrogen
72
empirical formula
indicates the types of atoms present and the simplest whole-number ratio of the number of atoms (or ions) in the compound
73
isomers
compounds with the same chemical formula but different molecular structures; This small difference in the arrangement of the atoms has a major effect on their respective chemical properties.
74
structural isomers
compounds in which the molecules differ in how the atoms are connected to each other.
75
structural isomers
compounds in which the molecules differ in how the atoms are connected to each other.
76
spatial isomers
the relative orientations of the atoms in space can be different.
77
the periodic law
the properties of the elements are periodic functions of their atomic numbers
78
A modern periodic table arranges the elements
in increasing order of their atomic numbers and groups atoms with similar properties in the same vertical column
79
metals
elements that are shiny, ductile/malleable, good conductors of heat and electricity. Conduct heat and electricity
Solids at room temperature (except
Mercury, Hg, which is liquid)
80
nonmetals
elements that appear dull, do not conduct.
| • At room temperature, Solids, Gases, (H2, O2, N2, Cl2, F2 and Group 8A) and one liquid (Br2)
81
metalloids
elements that conduct heat and electricity moderately well, and possess some properties of metals and some properties of nonmetals. B, Si, Ge, As, Sb,Te. solids, good semiconductors, ie: poor conductors at low temperatures and moderately good at high temperatures. When doped (a small amount of impurity added) they become very good conductors.
82
alkali metals
the elements in group 1 (the first column) form compounds that consist of one atom of the element and one atom of hydrogen. These elements (except hydrogen) are known as alkali metals. Very reactive, never found free in nature. React with water to produce basic solutions.
83
alkaline earth metals
The elements in group 2 (the second column) form compounds consisting of one atom of the element and two atoms of hydrogen, Very reactive, never found free in nature. React with water to produce basic
solutions. Mg and Ca are the 5th and 7th most abundant metals in the earth’s crust which is why there called alkaline EARTH metals
84
pnictogens
group 15
85
chalcogens
(from the Greek word, khalkos, for copper) because most copper ores contain these elements. group 16. sulfur (most stable w/8), selenium, and tellurium.
86
halogens
group 17. diatomic. [chlorine will always mean Cl2, etc etc.] called halogen which means salt-forming. all halogens will react with metals to form salts. (i.e. 2Na(s) + Cl2(g) → 2NaCl)
87
Group 8A: Noble, Inert, Rare
| Gases
group 18, Found in small amounts on earth and in the atmosphere.
• Argon (lazy one), Neon (new one), Krypton (hidden one), Xenon (strange one), Rn, He
• Least re-active of all elements.
88
atomic mass given in square brackets
This is done for elements that consist entirely of unstable, radioactive isotopes. An average atomic weight cannot be determined for these elements because their radioisotopes may vary significantly in relative abundance, depending on the source, or may not even exist in nature. The number in square brackets is the atomic mass number (an approximate atomic mass) of the most stable isotope of that element.
89
hydrogen
The most abundant element in the universe, making up about 3/4 of all matter.
• Typically exists on earth as a compound, not as an element.
• Note: Despite its position, does not share properties with Group 1A.
90
group 3A- aluminum
the most abundant of all metals, and most abundant in earth’s crust (8.3% by mass).
Lightweight and resists corrosion.
Reactive. It occurs naturally as compounds in rock and soil, called Minerals, and in gems.
91
group 3A- Gallium
one of the few metals that can be liquid close to room temperature.
92
group 4A: Carbon
basis of all organic compounds
93
organic compounds
used to be defined as a compound associated with a living system. but someone was able to synthesize urea in a lab so since then has been redefined as carbon-based compounds
94
organic compounds
used to be defined as a compound associated with a living system. but someone was able to synthesize urea in a lab so since then has been redefined as carbon-based compounds
95
DNA
carbon-based compound. rings contain carbon and nitrogen
96
allotropes
some nonmetals, like carbon, have allotropes- diff structural forms of the same element with diff properties.
97
diamond
allotrope of carbon made of pure carbon. each carbon bonded to 3 other carbons in 3d space. super strong.
98
graphite
comes from layered sheets of hexagons, slippery texture. makes up pencil lead. layers slip off when u write. also made up of pure carbon
99
buckyball
a molecule of carbon with 60 carbon atoms.
100
group 5A; Nitrogen
* Almost 80% of earth’s atmosphere consists of N2
| * Nitrogen is a major component of DNA and Proteins.
101
Group 5A; Bismuth
Bismuth (Z=83) is the heaviest element which is not radioactive.
102
radioactivity
certain elements are inherently unstable and spontaneously break down. past bismuth, every element is radioactive to a certain extent. some slower than others. how radioactive an element is has to do with its size. larger--> more radioactive. tritium and carbon14 are radioactive too but past bismuth all are reactive and before some are and some arent
103
group 6A: oxygen
20% atmosphere is O2
Very reactive - most chemical energy comes from reactions in which something combines with oxygen! (i.e. we eat food bc want the chemical energy in food. what rxn produces that energy? breaks down into glucose which interacts with oxygen to produce energy. or in a car fuel reacts with oxygen to produce energy) one of the nonmetals that has allotropes. monoatomic oxygen is v reactive.
104
ozone
allotrope of oxygen. we have 2 types of ozone in atmosphere. stratospheric ozone which protects us from UV radiation and tropospheric ozone which is harmful to the respiratory system.
105
Polyatomic Elements
White Phosphorous, one allotrope of this element, has the formula P4; S8 Sulfur
106
Two rows at the bottom of the table are the
Lanthanides [from lanthanum (Z = 57) to hafnium (Z = 72)] and the Actinides [from actinium (Z = 89) to rutherfordium (Z = 104)].
107
transition elements
middle of periodic table,
108
color of iron
orange
109
color of nickel
green
110
color of Copper
blue
111
color of Zn
clear
112
period
row on periodic table
113
group
a column on the periodic table
114
j.j. Thompson experiment and his discovery
determined an electron is a subatomic particle and discovered the mass-charge relationship by charging one end of a cathode ray tube and noting the ray emitted by the cathode (negative plate) across the tube and by putting + and - plates above and below, thereby causing the ray to vee. mass determined type of magnetic field needed and charge determined electric field needed.
115
Millikin's oil drop experiment
determined charge of a single electron. allowed oil to drop into a chamber with positively and negatively charged plates above and below. depending on # of charges on them, droplets would slow, stop or start to rise in opposite direction.
116
forces involved in milikan's eperiment
1. gravity causes the droplet to fall Fg= (mass)(gravity constant)
2. electrostatic force causes the droplet to fall. electrostatic force= (electric field)(charge)
by equating the 2, Millikin learned the charge on each droplet.
117
period
row on periodic table
118
group
a column on the periodic table
119
j.j. Thompson experiment and his discovery
determined an electron is a subatomic particle and discovered the mass-charge relationship by charging one end of a cathode ray tube and noting the ray emitted by the cathode (negative plate) across the tube and by putting + and - plates above and below, thereby causing the ray to vee. mass determined type of magnetic field needed and charge determined electric field needed.
120
Millikin's oil drop experiment
determined charge of a single electron. allowed oil to drop into a chamber with positively and negatively charged plates above and below. depending on # of charges on them, droplets would slow, stop or start to rise in opposite direction.
121
forces involved in Millikan's experiment
1. gravity causes the droplet to fall Fg= (mass)(gravity constant)
2. electrostatic force causes the droplet to fall. electrostatic force= (electric field)(charge)
by equating the 2, Millikin learned the charge on each droplet.
122
dissolving implies
physical change- a homogenous mixture
123
dissolving when a gas released implies
CHEM CHANGE
124
is solubility a physical or chemical property
physical
125
is melting/boiling point a physical or chemical property
physical
126
sig figs multiplication and division
same amount of sig figs as whichever value has fewer
127
sig figs addition and subtraction
same amount of places as whichever value has fewer
128
how many sig figs does an exact # have
infinite amount of sig figs
129
rounding when # to right of last SF is 5 and there are no other nonzero numbers after it
round up
130
rounding when # to right of last SF is 5 and there are other nonzero numbers after it
round up or down in order to keep the last sig fig an even number
131
are properties of a compound = to the properties of the elements it is composed of?
no
132
are all compounds composed of molecules?
no. but molecular compounds are
133
is air homogenous or heterogenous
homogenous bc can be liquified and separated bc nitrogen and oxygen have different boiling points
134
can a pure substance be broken down physically
no
135
chemical change
a pure substance undergoes a chemical reaction that transforms it into a different pure substance
136
pure substance
has a fixed composition and CANNOT be broken down by physical changes
137
mixture
does NOT have a fixed composition and CAN be broken down by physical changes
138
heterogenous
no uniform composition, its distributed component are larger than molecules; can be separated physically
139
homogenous
uniform throughout, distributed on a molecular/particulate level; can be separated physically
140
naming molecular compounds
put corresponding prefix in front of each element.
141
1
mono
142
2
di
143
3
tri
144
4
tetra
145
5
penta
146
6
hexa
147
7
hepta
148
8
octa
149
9
nona
150
10
deca
151
metals ___ electrons to form ___
Metals lose electrons to form cations
152
nonmetals ___ electrons to form ___
nonmetals gain electrons to form anions
153
For the main group elements, cations and anions form to
achieve the same number of electrons as the nearest noble gas atom.
154
Possible charges on Cr cation
Cr3+ or Cr6+ Chromium (III) or Chromium (VI) Ion
155
Possible charge on Mn cation
Mn2+ Manganese (II) Ion
156
Possible charge on Fe
Fe2+ or Fe3+ Iron (II) or Iron (III) Ion
157
Possible charge on Cu
Cu2+ or + Copper Ion
158
Possible charge on Ag
Ag+ Silver Ion
159
Possible charge on Hg
Hg2+ or Hg22+ Mercury(II) or Mercury(I) Ion
160
Possible charge on Zn
Zn2+ Zinc Ion
161
Possible charge on Pb
Pb2+ Lead Ion
162
naming metal ions
The name of the element followed by the word “ion”. But when an element can form more than one type of ion, we also add in Roman numerals to state the charge.
163
naming nonmetal ions
add ide
164
ammonium
one of the few common polyatomic
| cations. the rest of the polyatomic ions we will be discussing are anions
165
naming oxyanions
The one with less oxygen atoms is ‘ite’ and the one with more oxygen atoms is ‘ate’. and If there are more than two ions in the series, then the prefixes hypo-, meaning less than, and per-, meaning more than, are used.
166
oxyanion
polyatomic ions that contain one or more oxygen atoms.
167
ionic compounds
A compound that contains ions and is held together by ionic bonds; generally a metal + a nonmetal; solids that typically melt at high temperatures and boil at even higher temperatures; In solid form, an ionic compound is not electrically conductive because its ions are unable to flow (“electricity” is the flow of charged particles). When molten, however, it can conduct electricity because its ions are able to move freely through the liquid
168
N2O
nitrous oxide
169
NO
Nitrogen monoxide
170
Naming binary acids (comprised of hydrogen and one other nonmetallic element)
The word “hydrogen” is changed to the prefix hydro-
The other nonmetallic element name is modified by adding the suffix -ic
The word “acid” is added as a second word
171
naming oxyacids (compounds that contain hydrogen, oxygen, and at least one other element)
Omit “hydrogen”
Start with the root name of the anion
Replace –ate with –ic, or –ite with –ous
Add “acid”
172
name H2CO3
carbonic acid
173
empirical formulas are derived from experimentally measured element masses by:
1. Deriving the number of moles of each element from its mass
2. Dividing each element’s molar amount by the smallest molar amount to yield subscripts for a tentative empirical formula
3. Multiplying all coefficients by an integer, if necessary, to ensure that the smallest whole-number ratio of subscripts is obtained
174
how to calculate % composition
mass of element / mass of compound * 100%
175
Deriving Empirical Formulas from Percent Composition
the percent composition can be used to calculate the masses of elements present in any convenient mass of compound; these masses can then be used to derive the empirical formula in the usual fashion.
176
Derivation of Molecular Formulas
molar mass/empirical mass and then you multiply the resulting number by the subscripts in the empirical formula. that this same approach may be used when the molar mass (g/mol) instead of the molecular mass (amu) is used. In this case, one mole of empirical formula units and molecules is considered, as opposed to single units and molecules.
177
concentration
The relative amount of a given solution component
178
solvent
Often, though not always, a solution contains one component with a concentration that is significantly greater than that of all other components: the solvent. The solvent may be viewed as the medium in which the other components are dispersed, or dissolved.
179
aqueous solution
A solution in which water is the solvent is called an aqueous solution.
180
Solute
A solute is a component of a solution that is typically present at a much lower concentration than the solvent. Solute concentrations are often described with qualitative terms such as dilute (of relatively low concentration) and concentrated (of relatively high concentration).
181
molarity=
the number of moles of solute in exactly 1 liter (1 L) of the solution:
182
dilution
Dilution is the process whereby the concentration of a solution is lessened by the addition of solvent. For example, a glass of iced tea becomes increasingly diluted as the ice melts. The water from the melting ice increases the volume of the solvent (water) and the overall volume of the solution (iced tea), thereby reducing the relative concentrations of the solutes that give the beverage its taste (Figure 3.16).
Dilution is also a common means of preparing solutions of a desired concentration. By adding solvent to a measured portion of a more concentrated stock solution, a solution of lesser concentration may be prepared. For example, commercial pesticides are typically sold as solutions in which the active ingredients are far more concentrated than is appropriate for their application. Before they can be used on crops, the pesticides must be diluted. This is also a very common practice for the preparation of a number of common laboratory reagents.
183
According to the definition of molarity, the number of moles of solute in a solution (n) is equal to
the product of the solution’s molarity (M) and its volume in liters (L)
184
the dilution equation.
n=ML
n1=M1L1
n2=M2L2
M1L1=M2L2
C1V1=C2V2
