Download Honors Unit 2a Atomic Structure I

Number of Protons
Atomic Number
Always an integer!
Number of Protons + Neutrons
Mass Number
Always an integer!
Left Superscript = mass number
12C
6
Left Subscript = atomic number
12C
6
35
80Br
35
Atomic Number = ?
20
20Ne
10
Mass Number = ?
27
27Al
13
Mass Number = ?
20
40Ca
20
Atomic Number = ?
Neutral atoms have the same
number of electrons and protons.
# of electrons in a neutral
atom?
Atoms of the same element with a
different # of neutrons
Same # of protons, different # of neutrons
Same atomic #, different mass #
12
6C
14
and 6C
Isotope
Charge = +1, mass = 1 amu,
location = inside nucleus
Characteristics of Proton
Charge = 0, mass = 1 amu,
location = inside nucleus
Characteristics of Neutron
Charge = -1, mass = 1/1836 amu or
0.0005 amu, location = outside
nucleus
Characteristics of Electron
Summary of facts for
subatomic particles
Relative Mass (amu)
Relative Charge
Location
Proton
1.007276 or  1
+1
Nucleus
Neutron
1.008665 or  1
0
Nucleus
Electron
.00054858 or 
0.0005 or  0
-1
Outside
Nucleus
An atom that has gained or lost
electrons & so carries charge
Ion
An atom that has LOST
electrons
Positive Ion
An atom that has GAINED
electrons
Negative Ion
# protons - # electrons
Charge
Protons & Neutrons
Nucleons
Smallest bit of an element that
retains the properties of the
element.
atom
Smallest bit of an element that can
participate in a chemical reaction.
atom
Mass number – atomic number
# of neutrons
Subtract the atomic number FROM the mass number!
8 neutrons
6 protons
6 electrons
14C
6
# of neutrons = ?
# of protons = ?
# of electrons = ?
5 neutrons
4 protons
4 electrons
9Be
4
# of neutrons = ?
# of protons = ?
# of electrons = ?
22 neutrons
18 protons
18 electrons
40Ar
18
# of neutrons = ?
# of protons = ?
# of electrons = ?
8 neutrons
7 protons
7 electrons
15N
7
# of neutrons = ?
# of protons = ?
# of electrons = ?
Right superscript = charge
2+
24Mg
12
8 neutrons
7 protons
10 electrons (gained 3)
-3
15N
7
# of neutrons = ?
# of protons = ?
# of electrons = ?
10 neutrons
9 protons
10 electrons (gained 1)
-1
19F
9
# of neutrons = ?
# of protons = ?
# of electrons = ?
8 neutrons
8 protons
10 electrons (gained 2)
-2
16O
8
# of neutrons = ?
# of protons = ?
# of electrons = ?
12 neutrons
11 protons
10 electrons (lost 1)
+1
23Na
11
# of neutrons = ?
# of protons = ?
# of electrons = ?
12 neutrons
12 protons
10 electrons (lost 2)
+2
24Mg
12
# of neutrons = ?
# of protons = ?
# of electrons = ?
14 neutrons
13 protons
10 electrons (lost 3)
+3
27Al
13
# of neutrons = ?
# of protons = ?
# of electrons = ?
•Charge on the nucleus only. Does not
include the electrons.
•Always positive.
•Equals the number of protons.
•Equals the atomic number.
Nuclear Charge
Positive ion
Cation
Negative ion
Anion
Billiard Ball Model
Dalton’s Model
Solid
Indivisible
Homogeneous
1. All matter is composed of atoms.
2. Atoms of a given element are identical, atoms
of different elements are different.*
3. Atoms cannot be subdivided, created, or
destroyed.*
4. Atoms of different elements combine in small
whole number ratios to make compounds.
5. In chemical reactions, atoms are rearranged.
Dalton’s model
Plum Pudding Model
Thomson’s Model
Solid
Divisible
Inhomogeneous: contain charges!
Electrons are particles!
- +
+ -+ +
+ -
Deflection of cathode ray
Deflection
in
magnetic
field
No deflection in
field free region
Deflection in
electrostatic
field
Thomson gets credit for discovering
electron because he got the first “numbers”
– he found the charge-to-mass ratio of the
electron.
Thomson’s model
Nuclear Model
Rutherford’s Model
Mostly empty space
Divisible
Inhomogeneous
Contains a small,
dense positive nucleus
-
+
-
Nuclear Model
Rutherford’s model
Shot αlpha particles at gold foil.
1.
Most went through, so most of the atom
is empty space.
2. Some deflected back by small dense
positive nucleus.
Rutherford’s Experiment
Rutherford’s Experiment
A very
small
percent of
the alpha
particles
deflected
back:
Evidence
for a small,
dense,
positive
nucleus.
Most of
the alpha
particles
went
through
so most
of the
atom is
empty
space
Shell Model
Bohr’s Model
Shell Model
Bohr’s Model
Electron is still a particle.
Quantized energy levels.
Electrons move on 3-D spherical orbits.
In the NYS Reference Tables!
Bohr configurations are “irregular” because the
Bohr model is incorrect. You cannot predict
them for the larger atoms, even if you know the
maximum capacities of each orbit.
Bohr Configurations
Sulfur: 2-8-6
Valence electrons
are in outermost
orbit
16 p
Bohr Diagram
Maximum Capacity of Orbits
Orbit, n
Maximum Capacity
1
2
2
8
3
18
4
32
n
2n2
Sulfur: 2-8-6
3 occupied
levels but only
two completely
occupied levels.
Read question
with care!
Bohr Configuration
Wave Mechanical Model
Schrodinger’s Model
Electron is treated as a wave.
Electron Energy is Quantized.
Most probable location = orbitals.
Wave Mechanical Model
Schrodinger’s Model
• Use dots or x’s to represent the valence
electrons.
•
• The symbol represents the nucleus and all
the inner shell electrons – this is the
kernel of the atom.
• In NYS, the # of dots has to match the
# of valence electrons.
Lewis Dot Diagrams for Atoms
Lewis Dot Diagrams
The mass of the entire atom: includes
protons, neutrons, electrons.
Expressed relative to the mass of a
Carbon-12 atom.
atomic mass
1 atomic mass unit  1/12 the mass
of a C-12 atom.
or
The C-12 atom has a mass of
12.000 . . . atomic mass units.
atomic mass unit
Table of Isotopic Masses:
Mass of one specific isotope
Notice that these are decimals!
Why is C-12 exactly
12.0000000…?
Because C-12 is the standard!
Isotopic Mass
Note: we use these rarely!
The weighted average of the masses of the
naturally occurring isotopes of an element.
What are these masses
in the periodic table?
Average atomic mass
Warning: chemists get sloppy & call this atomic mass.
1) Convert % abundance to decimal
format.
2) Multiply abundance factor by
appropriate mass.
3) Sum
Average atomic mass
1) Final answer must be between
the highest & lowest masses.
2) Final answer will be closest to
mass of most abundant isotope.
Quick check on average
atomic mass calculation.
1) 75% = .75 and 25% = .25
Convert % to decimal
2) (.75) X 35 = 26.25 and (.25) X 37 = 9.25
Multiply each abundance factor by appropriate mass
3) 26.25 + 9.25 = 35.5 = avg. atomic mass of Cl
Add up all the terms
4) Ans is between 35 & 37, but closer to 35.
Quick check. Report to tenths place.
Calculate the average atomic mass of Cl.
Note: The NYS Regents
make severe
approximations to the
isotopic masses! So no
worries about sig figs!
Isotope
Percent Abundance
Cl-35
75%
Cl-37
25%
Mass is neither created nor destroyed
Total Mass Before = Total Mass After
Total Mass Reactants = Total Mass Products
Law of Conservation of Mass
for ordinary chemical and physical change
O2 + 2H2  2H2O
32 g + X g = 36 g
X=4g
32 grams of oxygen reacts with X grams
of hydrogen yielding 36 g of water.
Recall: in
this kind of
problem you
do NOT use
the
coefficients!
Law of Conservation of Mass
for ordinary chemical and physical change
Reactants  Products
aA + bB  cC + dD
A & B are reactants.
C & D are products.
Chemical Equations
A chemical compound contains the
same elements in exactly the same
proportions by mass regardless of
sample size or source.
Law of Definite Proportions
NaCl is
39.3% Na and 60.7% Cl
no matter how big the sample or where
it is from
Law of Definite Proportions
When two or more different compounds are
composed of the same two elements, then the
ratio of the mass of the second element combined
with a certain fixed mass of the first element is
always a ratio of small whole numbers.
Law of Multiple Proportions
Normalized Data!
Grams Mn
Grams O
Compound A
17.16 g
5.00 g
Compound B
12.87 g
5.00 g
Compound C
11.44 g
5.00 g
Take ratios of the Mn masses!
A/B = 17.16/12.87 = 1.33 = 4/3
A/C = 17.16/11.44 = 1.5 = 3/2
B/C = 12.87/11.44 = 1.125 = 9/8
This is the fixed
mass, so you can
forget about it!
Law of Multiple Proportions
Both atomic mass units and the
mole are based on C-12.
Relative atomic mass
Dual Perspective
Microscopic vs. Macroscopic
Relative atomic mass
Take the relative atomic mass from the
periodic table and
1. Stick “atomic mass unit” after it to get
the average mass of one atom
or
2. Stick “gram” after it to get the mass of
one mole of that element.
Relative atomic mass
•
Measure of the amount of substance in terms
of the number of particles.
•
The amount of any substance that contains as
many particles as there are atoms in 12 grams
of pure 12C.
Mole
Dual Perspective
The average Li
atom has a mass
of 6.941 atomic
mass units.
A mole of Li
atoms has a mass
of 6.941 grams.
6.02 X 1023
Avogadro’s Number
mole
Mass of one mole of a substance. For
elements, the molar mass is the
relative atomic mass expressed in
grams.
Molar Mass
# of moles = Mass of sample
Molar Mass
# of moles from Table T.
6.02 X 1023
1 Mole
3.01 X 1023
0.5 Mole
1.50 X 1023
0.25 Mole
12.04 X 1023
Or
1.204 X 1024
2.0 Mole