Unit 1 Structure of Matter
AP Chemistry
I.
Measurement- (1.4 -1.6)
A.
Metric System- developed in 1791 in France during the French
Revolution to standardize all measurements.
1. Based on properties of natural objects, size of Earth, weight of
water, speed of light, etc.
2. In 1960 an international agreement was reached specifying a
particular choice of metric units for use in scientific
measurements. These preferred units are called SI base units,
after the French Système International d'Unités.
3. Prefixes used to determine magnitude of particular unit.
4. Conversions are all base 10.
B.
Uncertainty in measurements- all measurements involve a certain
amount of error or uncertainty.
1. Due to errors of measuring device or operator error.
2. Uncertainty must be minimized.
3. Precision- how closely grouped a series of measurements are
to each other. Use % deviation.
4. Accuracy- how close to the actual or accepted value a series
of measurements are. Use % error.
5. Percent error- used when you are comparing your result to a
known or accepted value.
%𝑒𝑟𝑟𝑜𝑟 =
6.
C.
D.
E.
F.
|𝑦𝑜𝑢𝑟 𝑟𝑒𝑠𝑢𝑙𝑡 − 𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒|
× 100
𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒
Percent deviation- used to compare how closely grouped your
results are to the mean. Steps:
∑ 𝑑𝑎𝑡𝑎 𝑝𝑜𝑖𝑛𝑡𝑠
𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑑𝑎𝑡𝑎 𝑝𝑜𝑖𝑛𝑡𝑠(𝑁)
∑|𝑒𝑥𝑝.𝑣𝑎𝑙𝑢𝑒−𝑚𝑒𝑎𝑛|

calculate mean=

calculate average deviation=

%deviation= 100(
𝑎𝑣𝑒. 𝑑𝑒𝑣𝑖𝑎𝑡𝑖𝑜𝑛
𝑚𝑒𝑎𝑛
𝑁
)
Significant Figures- All measurements are approximations—no
measuring device can give perfect measurements without
experimental uncertainty. Significant figures are important because
they tell us how good the data we are using is. Sig. Fig’s. indicate
the level of certainty of data
Rules:
1. Ignore leading zeros.
2. Ignore trailing zeros, unless they come after a decimal point.
3. Everything else is significant.
4. exact numbers (metric conversions, counting or written
numbers) have infinite number of sig. fig’s.
Sig. Fig’s. in calculations
1. limited by least accurate measurement
2. x, : answer has the same number of sig figs as the
measurement with the fewest
3. +, : result should be equal to the smallest number of decimal
places in the original measurements.
Conversions (dimensional analysis)- allows one to change from
one unit to another. Units are always used in all calculations.
They are multiplied, divided, and canceled like any other algebraic
quantity.
1. Set up equality in fraction form.
2. The equalities are then lined up sequentially and units used on
the top and bottom of neighboring fractions are alternated so
that units cancel.
5.00 𝑖𝑛 ×
2.54 𝑐𝑚
= 12.7 𝑐𝑚
1.00 𝑖𝑛
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G. How to measure amount of matter- (3.4-3.5) moles are used to
measure amount.
1. A mole is a unit used in chemistry to measure amount of
atoms, molecules, ions, etc.
a. 1 mole = 6.022 x 1023 particles
b. Also, 1 mole of atoms has a mass equal to the average
mass in grams of an element on the PT.
c. Molar mass (MM) is the sum of masses of atoms in a
chemical formula. Units are:
 Al: 27.0 g/mol
 H2O: 18.0 g/mol
𝐠𝐫𝐚𝐦𝐬
𝐦𝐨𝐥𝐞
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑖𝑛 1 𝑚𝑜𝑙 𝑜𝑓 𝑐𝑚𝑝𝑑
d. %composition=
𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑚𝑜𝑙 𝑜𝑓 𝑐𝑚𝑝𝑑
× 100%
e.
2.
empirical formulas- smallest whole # ratio of elements in a
compound.
f. molecular formulas- actual # of elements in a compound.
Density- amount of mass in a specific volume
a. 𝑑 =
b.
II.
𝑚
𝑉
density of H2O= 1.00
𝑔
𝑚𝐿
= 1.00
𝑔
𝑘𝑔
= 1000 3
𝑚
𝑐𝑚3
History of the Atom- (2.1-2.7)
A.
John Dalton (1808) proposed his Atomic Theory:
1. Elements made of tiny particles, called atoms
2. Atoms of a given element are identical; the atoms of
different elements are different.
3. In chemical reactions, atoms are combined, separated, or
rearranged. Atoms cannot be subdivided, created, or
destroyed. He assumed that the atom was the ultimate
particle.
4. Atoms of one element may combine with atoms of other
elements, usually in small whole number ratios, to form
compounds.
B.
J.J. Thomson- In 1897, J.J. Thomson discovered the electron,
the first subatomic particle. He also was the first to attempt to
incorporate the electron into a structure for the atom.
1.
He used a cathode ray tube to make his conclusions. He
observed:
a. The cathode rays traveled in straight lines.
b. Cathode rays were deflected by a magnetic field
c. The rays were deflected away from a negatively
charged object.
d. All metals produce these rays.
2.
He proposed that atoms consist of small, negative
electrons embedded in a massive, positive sphere.
Measured charge-to-mass ratio of e-.
3.
C.
Ernest Rutherford- New Zealand chemist (1871-1937)
1. Gold foil experiment- alpha particle (2 protons and 2
neutrons bonded together) bombardment of gold foil.
Observations:
a. most went straight through unaffected
b. small number had small deflections
c. rarely they would come straight back.
Conclusions:
a. nucleus is positively charged
b. mass of atom is located in nucleus
c. atom is mostly empty space.
2
III.
D.
Robert Millikan- 1909 Oil Drop Experiment: determines the size
of the charge on an electron.
1. What Millikan did was to put a charge on a tiny drop of oil,
and measure how strong an applied electric field had to be
in order to stop the oil drop from falling.
2. He noticed that the charge was always a multiple of
-1.6 x 10 -19 C ( coulombs are a quantity of electrical
charge), the charge on a single electron.
E.
Laws that led to the modern atomic theory- conservation of
mass, definite proportions and multiple proportions.
1. Conservation of mass- in a chemical reaction, atoms
(and therefore mass) are never lost or gained only
rearranged.
2. Definite Proportions- in a pure compound the proportions
of atoms by mass are always the same.
3. Multiple Proportions- If two elements A and B form more
than one compound, the masses of B that can combine
with a given mass of A are in a ratio of small whole #’s.
Parts of the Atom (2.3-2.5)
A.
Proton- symbol ( 1 H or p+)
1. Positively charged particle, (+1).
2. Part of the dense nucleus along with neutrons
3. Mass of 1.0073 amu per proton, about 2000 times more
massive than an electron.
4. Along with neutrons in the nucleus make up most of the mass
of the atom
5. Along with neutrons in the nucleus make up a small part of the
atoms overall volume.
6. Scientists have agreed to identify elements by atomic
number, which is the number of protons each atom has.
Symbol for atomic number is (Z).
1
B.
Neutron- symbol ( 0 n or n0)
1. Electrically neutral, 0 charge.
2. About the same mass as a proton, 1.0087 amu.
3. Found in the nucleus.
4. Number of neutrons determines the isotope.
C.
Electron- symbol ( -1 e or e–)
1. Electrons occupy 3D regions of space called orbitals that
surround the nucleus.
2. Negatively charged (-1 charge),
3. 1/2000 the mass of a proton, 5.5 x 10-4 amu.
4. loss, gain, and sharing of electrons important in many
chemical reactions.
D.
Ion- ions are atoms that have lost or gained electrons.
–
# electrons ¹ # protons
–
e- > p+: (–) charged (anion): Xn–
e- < p+: (+) charged (cations): Xn+
E.
1
Particle
Location
Relative
Charge
Mass
(amu’s)
Symbol
proton
nucleus
+1
1.0073
p+ or 1 H
neutron
nucleus
0
1.0087
electron
orbital
-1
.00055
1
0
0
-1
1
n or n0
e or e-
0
Isotopes- atoms of the same element that contain different
numbers of neutrons.
1. Mass Spectrometer is used to differentiate isotopes.
a. particles are turned into positive ions, accelerated, and
then deflected by a magnetic field.
b. the resulting path of ions depends on their
mass/charge ratio (m/Z).
c. large m/Z value deflected least.
Mass Spectrograph
Two peaks showing two
isotopes of Boron
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2.
3.
4.
5.
For any element, there is no set number of neutrons in the
nucleus. For example, most hydrogen atoms (atomic
number 1) have no neutrons, a small percentage, have one
neutron, and a smaller percentage have 2.
We identify isotopes by their mass number (A), which is
the total number of protons and neutrons.
The atomic number (Z) is the number of protons, it defines
the atom.
The total mass of an atom is called its atomic mass. This
is the sum of the masses of all the atom’s components.
a.
average atomic mass- is the average mass of all
isotopes of an element as they occur in nature.
b. The unit to measure atomic masses is the atomic
mass unit (amu).
–
1 amu = 1.66x10-24 g.
–
1 amu = 1/12 the mass of a C-12 atom.
6.
A
Notation for isotopes Z X , for example: 79 Au
201
7. 100mav = %1m1 + %2m2 + ...
IV.
Classifying Matter (1.2-1.3)
A.
V.
Matter can be classified according to it’s state and/or it’s
composition.
1. Three states of matter- solid, liquid, gas.
a. solid- atoms packed very close in fixed locations. Atoms
vibrate, have fixed volume, rigid shape. Can be crystalline
or amporphous.
b. liquid- packed close, but are free to move relative to each
other. Fixed volume, but not shape.
c. gas- atoms are very far apart, assume shape and volume
of container, compressible.
2. Composition- see diagram.
Radioactive decay (21.2, 21.4)
A.
Radioactive decay
1. Decay is the release of radiation by a radioactive isotope.
2. The nuclei are unstable and emit radiation
a. The “tug-of-war” between the attraction of the strong
nuclear force and the repulsion of the electromagnetic
force between protons has interesting implications for the
stability of a nucleus.
b. Atoms outside the zone of stability tend to decay and
release radiation, until they get back to the “belt of
stability”
c. Eventually, a point is reached beyond which there are no
stable nuclei: the bismuth nucleus with 83 protons and
126 neutrons is the largest stable nucleus.
B.
1.
2.
3.
Half-Life
Time it takes for half of a given amount of a radioactive
isotope to undergo decay.
rate of decay is proportional to #of nuclei present:
rate = kNt (k: rate constant, Nt is amount at time “t”)
time for half of remaining atoms to decay (t½) is constant:
𝒕𝟏/𝟐 =
4. 𝑙𝑛
𝑁𝑡
𝑁0
𝒍𝒏𝟐
𝒌
=
𝟎.𝟔𝟗𝟑
𝒌
= −𝑘𝑡 or 𝑁𝑡 = 𝑁0 𝑒−𝑘𝑡 𝑁𝑡 =
𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑟𝑎𝑑𝑖𝑜𝑎𝑐𝑡𝑖𝑣𝑒 𝑛𝑢𝑐𝑙𝑒𝑖 𝑎𝑡 𝑡𝑖𝑚𝑒 𝑡
𝑁0 = 𝑖𝑛𝑖𝑡𝑖𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑟𝑎𝑑𝑖𝑜𝑎𝑐𝑡𝑖𝑣𝑒 𝑛𝑢𝑐𝑙𝑒𝑖
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VI.
Atomic spectroscopy and the Bohr model(6.1, 6.3, 6.4, )
A new model of the atom evolved out of the similarities
discovered between the behavior of light & electrons.
Analysis of the light revealed that an elements chemical
behavior is related to the arrangement of it’s electrons.
A.
Wave Nature of Light. Light is a form of
electromagnetic radiation with three characteristics:
1. wavelength- measured in meters or nanometers (m
or nm) is the distance between two consecutive
crests.
2. frequency – measured in hertz (Hz) is the number
of wavelengths that pass a certain point per second.
3. speed- how fast a wave is moving through space.
All EM radiation travels at 3.0 x 108 m/s.
4. Because light moves at a constant speed there is a
relationship between frequency and wavelength.
𝒄=𝝀∙𝝂
5.
𝜈 = frequency in Hertz
c = speed of light 3.0 x 108m/s
λ = wavelength in meters
Light energy comes in packets, called photons.
𝑬𝒑𝒉𝒐𝒕𝒐𝒏 = 𝒉 ∙ 𝝂
𝒉 = 𝑷𝒍𝒂𝒏𝒄𝒌′ 𝒔 𝒄𝒐𝒏𝒔𝒕𝒂𝒏𝒕 = 𝟔. 𝟔𝟐𝟔 × 𝟏𝟎−𝟑𝟒 𝑱 ∙ 𝒔
substitute �=�∙�
𝑬𝒑𝒉𝒐𝒕𝒐𝒏 =
𝒉∙𝒄
𝝀
B.
By passing light through a prism, the color components
of the light can be separated.
1. A continuous spectrum shows all the wavelengths
of light that are being emitted by white light. (think of
a rainbow)
2. An emission spectrum shows the specific
frequencies of light emitted by a specific atom that is
being excited.
3. Atoms can be identified by the light they emit, by
their unique emission spectrum.
C.
The Danish scientist Niels Bohr (1885-1962) explained
the formation of emission spectra (for hydrogen only):
1. Potential energy of an electron depends on its
distance from the nucleus.
2. When an atom absorbs a photon of light, it is
absorbing energy.
a. Absorption of a photon causes a low potential
energy electron in an atom to become a high
potential energy electron.
b. When a high potential energy electron loses some
of its energy, the electron moves closer to the
nucleus and the energy lost is emitted as a
photon.
3.
Since light energy is quantized, the energy of an
electron must also be quantized. In other words, an
electron cannot have just any amount of potential
energy.
a. Within the atom there must be a number of
distinct energy levels, analogous to steps on a
staircase.
b. Where you are at on the “staircase” is restricted
to where the stairs are. Similarly, there are
only a limited number of permitted energy
levels in an atom. An electron cannot exist
between levels.
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4.
Equation to calculate the energy that an electron would have
at any energy level:
−𝟐. 𝟏𝟖𝒙𝟏𝟎−𝟏𝟖 𝑱
𝑬𝒏 =
𝒏𝟐
a.
b.
c.
n is the energy level in question and the negative sign
means that the lower energies correspond to states with
larger negative numbers for energy values – be careful!
ground state (n = 1) electron has lowest (most negative)
energy
excited state (n > 1), electron energy increases until
ionized (E = 0 J)
∆Eelectron = En-final – En-initial
∆Eelectron > 0 when increasing n
∆Eelectron < 0 when decreasing n
|∆Eelectron| = Ephoton
Bohr developed a conceptual model in which an electron moving around
the nucleus is restricted to certain distances from the nucleus, these
distances are determined by the amount of energy the electron has. This is
called the planetary orbital model.
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