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Regents Chemistry
Mrs. Ingersoll
Unit 1 – Introduction to Matter
and Atomic Structure
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Matter
has mass.......
.....and volume
can be divided into substances + mixtures
Substance
matter with constant properties and composition,
regardless of its source – it is the same throughout
can be divided into elements and compounds
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Element
A substance that can not be broken down
to simpler substances
Made of only one kind of atom
Examples:
gold
carbon
oxygen
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Compounds:
--have two or more kinds of atoms held together by chemical bonds
Examples
water
carbon dioxide sodium chloride
glucose
H2O
CO2
NaCl
C6H12O6
-- formed by a chemical change
-- properties of the original elements are no longer apparent
-- can be decomposed by a chemical reaction
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Mixtures
--variable properties and composition
can be: homogeneous
Or
heterogeneous
-- evenly mixed
-- not evenly mixed
Salt water, Kool-aid
Any solution
Chocolate chip cookies, Raisin bran
--form without a chemical change
--properties of ingredients still apparent
--can be separated without a chemical reaction by using the
physical properties of the substances in the mixture
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Mixtures: Separation Techniques
1. Filtration – based on
differences in particle size,
or solubility
2. Distillation – based on
differences in boiling point
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3. Chromatography – based on differences in solubility and
attraction to another substance, ex. paper
Differences in density, freezing point, and other properties can
also be used to separate mixtures
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SUMMARY
Element
Substances
Compound
Mixture
element
+
element
element
+
compound
compound
+
compound
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THE STRUCTURE OF THE ATOM
Democritus, 400 BC
--Matter is made of small, indivisible
particles – “atomos”
In the 1700’s, scientists making accurate measurements
discovered several new laws --
1. Law of Conservation of Matter
--matter is not created or destroyed
during chemical or physical changes
2. Law of Definite Composition
-- compounds have an unvarying
composition
3. Law of Multiple Proportions
-- elements combine in simple ratios
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John Dalton, 1807 -- Atomic Theory
1. Each element is composed of one kind of atom
2. Atoms of different elements have different
masses and properties
3. Atoms can combine and recombine with each
other; but they don’t change from one kind to another; they aren’t
created or destroyed
4. In compounds, atoms combine in simple, constant ratios
Dalton thought that atoms were hard,
unbreakable spheres
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J.J. Thomson, 1897
--experimented with cathode ray tubes:
vacuum tubes with high voltage
passing through
--found that cathode rays could be:
1. deflected by magnets
2. bent toward positive electric fields
CONCLUSIONS:
1. cathode rays were streams of negatively charged particles
 electrons
2. atoms are made of sub-atomic particles with opposite
electrical charge
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Ernest Rutherford, 1910:
Gold Foil Experiment
-- took alpha particles (2 protons + 2 neutrons)
and shot them toward a thin gold foil
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Rutherford’s Results:
Most of the particles went straight
through; a few were deflected
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Interpretation of Rutherford’s ExperimentHis observations and conclusions:
1. Since most of the particles went straight through,
the atom is mostly empty space.
2. Some of the alpha particles were deflected so
the atom contains a nucleus that is:
a. very small
b. dense (a large mass in a small volume)
c. positively charged
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The nucleus contains most of the atom’s mass
and all its positive charge
proton
mass
1u
neutron
1u
0
Atomic number
Atomic mass
(mass number)
charge
+1
= number of protons
determines the element
= number of protons + neutrons
atomic mass
atomic number
10.81
B
5
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Not all atoms of the same element are identical
p
p
n
Atomic number: 1
Atomic mass:
1
H-1
n
p
n
1
1
2
H-2
3
H-3
These are called isotopes of hydrogen
Isotopes: atoms of the same element with different masses
-- same p, different n
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Isotopes of hydrogen have the same chemical properties
All are:
-- flammable gases
--able to react with oxygen to form water
They differ in mass and nuclear properties
All elements exist as mixtures of isotopes -the atomic mass of an element is the mass of an average atom
A carbon-12 atom has a mass of 12.0000000...
yet the atomic mass of the element carbon is 12.011
because a few carbon-13 and carbon-14 atoms are mixed in
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Sample Problem
Element X consists of 35% X-21 and 65% X-23
What is the average atomic mass of X?
Solution
Assume you have 100 atoms: 35 of them would be X-21
65 of them would be X-23
Find the total mass of all 100 atoms:
35 x 21 = 735
65 x 23 = 1495
2230 u = mass of 100 atoms
2230 / 100 = 22.30 u
The average is always between the isotope masses
and always closer to the one that is most abundant
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Charged Atoms
+ charge in nucleus, – charge outside
Add + and – to find the total charge
If the charge = 0 ( p = e ), it is a neutral atom
If the charge = 0 ( p = e ), it is an ion
Describing Atoms:
1. Atomic Number = number of protons
2. Atomic Mass (or mass number) = number of protons + neutrons
3. Electrical Charge = protons(+) + electrons(-)
a neutral atom has a charge of 0
(protons = electrons)
a positive ion (+ charge) has more protons than electrons
a negative ion (- charge) has more electrons than protons
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ee-
Examples:
atomic number = _________
3
atomic mass = ________
7
= proton
charge = _______
+1
= neutron
element = ______
Li
10
How many neutrons are in an atom of F? _______
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How many electrons are in a neutral atom of boron? ______
Suppose an atom has 14 protons, 16 neutrons, and 16 electrons
Si
-- what element is it? __________
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-- atomic number = _______
-- atomic mass = ______
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-2
-- charge = _______
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HOW TO READ
THE PERIODIC TABLE OF THE ELEMENTS
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ELECTRONS
Niels Bohr, 1913
-- said that electrons orbit the nucleus in
set paths, like planets around the sun:
or
These orbits were designated:
K
L
M
N
O
closest to nucleus
lowest potential energy
P
Q
farthest from nucleus
highest potential energy
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Bohr said that e- could be in any of these 7 orbits, but NOT
in between.
Add energy
e- move to higher energy levels, absorbing
the energy
Later, the e- falls back to a lower energy level, and the energy
is released in the form of electromagnetic radiation
e
e
e
energy absorbed
energy released
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Electromagnetic Radiation
-- energy that travels as waves
-- can travel through a vacuum
-- moves at 3.0 x 108 m/s, or “c” (speed of light)
λ -- wavelength
-- the number of waves that move past a fixed point in
one second is called the frequency, “f”
-- frequency is measured in Hertz (Hz), or waves per second
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-- As λ decreases, f increases
-- c = λ f
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In 1900, Max Planck discovered:
the energy carried by electromagnetic energy
is directly proportional to its frequency
(high frequency means high energy)
E= h f
energy = (planck’s constant)(frequency)
--so, as
λ
,
f
,
E
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Back to atoms ....
When all the e- in an atom are in the lowest possible energy orbits,
the atom is in the ground state – its normal condition
Add energy, and the e- move to higher energy levels, and the atom
is in the excited state – an unstable condition
Energy is released when the e- return to the ground state
L K jump--- short, little energy released
light emitted has low f, high λ
red
M
K
L
M
N
N
K jump--- medium energy released
light emitted has med f, med λ
green
K jump--- long, lots of energy released
light emitted has high f, low λ
violet
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Each electron jump releases one photon of a certain λ
This produces the spectrum of the element
Since each element starts with a different electron arrangement,
each element has a different spectrum
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Bohr’s model worked well for hydrogen, but not for more
complex elements. Other scientists in the 1920’s came up with
several modifications to his theory.
The Current Understanding of Electron Arrangement:
The Wave-Mechanical model
We can’t know an electron’s distance from the nucleus, as
Bohr thought, but we can calculate how likely it is to be found
at certain distances
K
L
electron probability
electron probability
probability
probability
probability
distance from nucleus
distance from nucleus
distance from nucleus
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The wave-mechanical model says that
electrons are in an electron cloud –
there are NO orbits
Instead, we have Principal Energy Levels
or Shells:
K
1
L
2
M
3
N
4
O
5
P
6
Q
7
Each level can hold a given number of electrons:
PEL
No. of e1
2
2
8
3
18
4 or higher
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2
If the PEL = n, the number of electrons = 2n
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The Principal Energy Levels are divided into orbitals
(NOT ORBITS) of different energy
An orbital is the region where the electron is most likely to
be found
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Rules for electrons in atoms:
1. No 2 electrons in an atom have exactly the same energy
2. Octet Rule – the last shell can never hold more than 8 e-
Example:
Atomic No. 18 Ar: 2 – 8 – 8
Atomic No. 19
K:
Atomic No. 19 K:
2–8–9
2–8–8–1
3rd shell can hold 10 more e
violates octet rule
4th shell begins to fill
So, a shell may not be filled with electrons before the
next shell begins to be filled
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-
15.9994
Electron configurations on the periodic table are for:
Neutral atoms-- Ex. oxygen is 2–6
-2
O
8
2–6
when its charge is –2, the electrons are… 2–8
In the ground state -- Ex. oxygen is 2–6
oxygen is 2—5—1 in the excited state
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Atoms are most stable when their outermost energy level
has as many e- as possible
(8 e- ; or 2 e- for the 1st shell)
This situation is called a stable octet
Elements that have a stable octet when neutral are called the
noble gases (group 18)
These elements undergo few, if any, chemical reactions
Other elements change during chemical reactions to obtain a
stable octet
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Electrons in the outermost energy level are called
valence electrons
Ex. Atomic No. 13 Al 2 – 8 – 3
Lewis electron-dot structures:
-- use one dot for each valence electron
Ex.
Al
Ne
Li
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