Unit 1 Structure of Matter
Honors Chemistry
I.
Measurement- (2.1 -2.3)
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.
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 =
C.
D.
E.
F.
|𝑦𝑜𝑢𝑟 𝑟𝑒𝑠𝑢𝑙𝑡 − 𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒|
× 100
𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒
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
1
G.
How to measure amount of matter- (11.1-11.4) 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
d.
𝐠𝐫𝐚𝐦𝐬
𝐦𝐨𝐥𝐞
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑖𝑛 1 𝑚𝑜𝑙 𝑜𝑓 𝑐𝑚𝑝𝑑
%composition =
𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑚𝑜𝑙 𝑜𝑓 𝑐𝑚𝑝𝑑
× 100%
e.
2.
II.
𝑚
𝑉
a.
𝑑=
b.
density of H2O= 1.00
𝑔
𝑚𝐿
= 1.00
𝑔
𝑘𝑔
= 1000 3
𝑚
𝑐𝑚3
Classifying Matter (3.1-3.4)
A.
B.
III.
empirical formulas- smallest whole # ratio of elements in a
compound.
f. molecular formulas- actual ratio of atoms in a compound.
g. conversions: moles ⇆ grams (use MM and dimensional
analysis.)
Density- amount of mass in a specific volume
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.
Separation Techniques- matter can be separated based on physical and
chemical properties.
1. Filtration- separation based on particle size; heterogeneous
mixture containing a solid phase.
2. Distillation- separation based on boiling point; homogenous
mixture, aka a solution.
3. Paper/Column Chromotography- separation based on molecule
attraction for the stationary phase; components of a solution are
separated from each other using the property of differential
migration (different rates of flow).
Solutions
History of the AtomA.
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.
2
b.
c.
2.
3.
IV.
Cathode rays were deflected by a magnetic field
The rays were deflected away from a negatively charged
object.
d. All metals produce these rays.
He proposed that atoms consist of small, negative electrons
embedded in a massive, positive sphere.
Measured charge-to-mass ratio of e-.
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.
D.
Robert Millikan- 1909 Oil Drop Experiment: determines the size of the
charge on an electron.
2. 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.
3. 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 (4.3)
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, zero 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.
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-
1
0
3
4.
D.
E.
loss, gain, and sharing of electrons important in many chemical
reactions.
Ion- ions are atoms that have lost or gained electrons.
–
# electrons ¹ # protons
–
e- > p+: (–) charged (anion):
Xn–
e- < p+: (+) charged (cations): Xn+
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.
2. For any element, there is no set number of neutrons in the
nucleus. For example, most hydrogen atoms (atomic #1) have no
neutrons, a small percentage, have one neutron, and a smaller
percentage have 2.
3. We identify isotopes by their mass number (A), which is the total
number of protons and neutrons.
4. The atomic number (Z) is the number of protons, it defines the
atom.
5. The total mass of an atom is called its atomic mass. This is the
sum of the masses of all the atom’s components.
Mass Spectrograph
Two peaks showing two
isotopes of Boron
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.
IV.
A
6.
Notation for isotopes Z X , for example: 79 Au
7.
100mav = %1m1 + %2m2 + ...
201
Radioactive decay (25.2, 25.3)
A.
B.
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.
Half-Life
1. Time it takes for half of a given amount of a radioactive isotope to
undergo decay.
2. rate of decay is proportional to #of nuclei present:
3. time for half of remaining atoms to decay (t½) is constant:
4. 𝑎𝑚𝑜𝑢𝑛𝑡 𝑟𝑒𝑚𝑎𝑖𝑛𝑖𝑛𝑔 = (𝑖𝑛𝑖𝑡𝑖𝑎𝑙 𝑎𝑚𝑜𝑢𝑛𝑡)(1⁄2)𝑛
𝑛 = #𝑜𝑓 ℎ𝑎𝑙𝑓 𝑙𝑖𝑣𝑒𝑠 𝑡ℎ𝑎𝑡 ℎ𝑎𝑣𝑒 𝑝𝑎𝑠𝑠𝑒𝑑
4
VI.
Atomic spectroscopy and the Bohr model- (5.1, 5.2)
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 �=�∙�
𝒉∙𝒄
𝑬𝒑𝒉𝒐𝒕𝒐𝒏 = 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.
5
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.
6