unit 2. atoms are measured in picometers, 10 -12 meters hydrogen atom, 32 pm radius nucleus tiny...
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Major Subatomic Particles
•Atoms are measured in picometers, 10-12 meters Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble• Radius of the nucleus is on the order of 10-15 m
• Density within the atom is near 1014 g/cm3
Name Symbol Charge Relative Mass (amu)
Actual Mass (g)
Electron e- -1 1/1840 9.11x10-28
Proton p+ +1 1 1.67x10-24
Neutron no 0 1 1.67x10-24
Elemental Classification•Atomic Number (Z) = number of protons (p+) in the
nucleus Determines the type of atom
• Li atoms always have 3 protons in the nucleus, Hg always 80
• Mass Number (A) = number of protons + neutrons [Sum of p+ and nº]
Electrons have a negligible contribution to overall mass
• In a neutral atom there is the same number of electrons (e-) and protons (atomic number)
Nuclear Symbols•Every element is given a corresponding symbol
which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and
atomic number
E A
Z
elemental symbol
mass number
atomic number
ATOMIC NUMBER AND MASS NUMBER
the number of protons in an atom
the number of protons and neutrons in an atomH
e2
4
Atomic Number
Mass Number
Number of electrons = Number of protons
in a neutral atom 5
•Find the number of protons number of neutrons number of electrons atomic number mass number
W184 74
F199 Br80
35
IonsCation is a positively charged particle.
Electrons have been removed from the element to form the + charge.
ex: Na has 11 e-, Na+ has 10 e-
Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge.
ex: F has 9 e-, F- has 10 e-
Isotopes•Atoms of the same element can have different
numbers of neutrons and therefore have different mass numbers
• The atoms of the same element that differ in the number of neutrons are called isotopes of that
element
• When naming, write the mass number after the name of the element
H11Hydrogen-1
H21
Hydrogen-2
H31Hydrogen-3
Calculating AveragesAverage = (% as decimal) x (mass1) + (% as decimal)
x (mass2) + (% as decimal) x (mass3) + …
Problem:
Silver has two naturally occurring isotopes, 107Ag with a mass of 106.90509 u and abundance of 51.84 % ,and 109Ag with a mass of 108.90476 u and abundance of 48.16 % What is the average atomic mass?
Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u)
= 107.87 amu
• If not told otherwise, the mass of the isotope is the mass number in ‘u’
• The average atomic masses are not whole numbers because they are an average mass value
• Remember, the atomic masses are the decimal numbers on the periodic table
Average Atomic Masses
• Calculate the atomic mass of copper if copper has two isotopes 69.1% has a mass of 62.93 amu The rest (30.9%) has a mass of 64.93 amu
• Magnesium has three isotopes 78.99% magnesium 24 with a mass of 23.9850 amu 10.00% magnesium 25 with a mass of 24.9858 amu The rest magnesium 26 with a mass of 25.9826 amu What is the atomic mass of magnesium?
More Practice Calculating Averages
BohrProposed electrons (e-) orbit around the nucleus
in circular pathsSaid e- in a particular path have a fixed energy
(energy levels)e- can go from any energy level to another by
gaining or losing a specific amount of energy = a “quantum of energy”
When e- absorbs a quantum of energy, it goes from it’s ground state (where it’s normally found) to an excited state
The excited state is at a higher energy level
Bohr postulated that: Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is
away from the nucleus An atom with maximum number of
electrons in the outermost orbital energy level is stable (unreactive)
Think of Noble gases
Atomic Line Emission Spectra and Niels BohrAtomic Line Emission Spectra and Niels Bohr
Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms.Problem is that the model only works for HydrogenNiels Bohr
(1885-1962)
Line Emission Spectra of Excited AtomsLine Emission Spectra of Excited AtomsExcited atoms emit light of only
certain wavelengthsThe wavelengths of emitted light
depend on the element.
Drawback to BohrBohr’s theory did
not explain or show the shape or the path traveled by the electrons.
His theory could only explain hydrogen and not the more complex atoms
Energy level populations (Science10)
Electrons found per energy level of the atom.
The first energy level holds 2 electronsThe second energy level holds 8 electronsThe third energy level holds 18 electrons
The Quantum Mechanical ModelEnergy is quantized. It comes in chunks.A quanta is the amount of energy needed to
move from one energy level to another.Since the energy of an atom is never “in
between” there must be a quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom – an ORBITAL
Orbits (Bohr) vs Orbitals (Quantum Mechanics)
Bohr said electrons travel in an orbit – can predict exact location of electron at any point in time.
Schrodinger used mathematics (calculus) to find the region in space where an electron will be found 90% of the time - this region is called an orbital. There are 4 main types of orbitals – s, p, d, and f.
Modern View of the Atom The modern view of the atom
suggests that the atom is more like a cloud.
Atomic orbitals around the nucleus define the places where electrons are most likely to be found.
23
s orbitals
1 s orbital forevery energy level
1s 2s 3sSpherical shapedEach s orbital can hold 2 electronsCalled the 1s, 2s, 3s, etc.. orbitals
p orbitalsStart at the second energy level 3 different directions3 different shapesEach orbital can hold 2 electrons
The d sublevel contains 5 d orbitalsThe d sublevel starts in the 3rd energy
level 5 different shapes (orbitals)Each orbital can hold 2 electrons
The f sublevel has 7 f orbitalsThe f sublevel starts in the fourth energy levelThe f sublevel has seven different shapes
(orbitals)2 electrons per orbital
Electron ConfigurationWe use e- configuration as a shorthand
to show how e- are arranged around a nucleus
Example: Carbon is …
1s2 2s2 2p2
Electron ConfigurationsThe way electrons are arranged in atoms.Aufbau principle- electrons enter the
lowest energy first.This causes difficulties because of the
overlap of orbitals of different energies.Pauli Exclusion Principle- at most 2
electrons per orbital - different spinsHund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair up until they have to .
Summary
Sublevel
# of shapes
(Orbitals)
Max number
of e-
Starts at
energy level
s 1 2 1
p 3 6 2
d 5 10 3
f 7 14 4
Electron Arrangement
1st Rule: The Aufbau Principlee- fill orbitals of the lowest energy firstWe can use the periodic table to help
us!
AbbreviationsWe can abbreviate electron
configurations using the Noble GasesEx: Sulfur
1s2 2s2 2p6 3s2 3p4
[Ne] 3s2 3p4
Ex: Lead 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
6s2 4f14 5d10 6p2
[Xe] 6s2 4f14 5d10 6p2
2nd Rule: Pauli Exclusion Principle
Each orbital orientation can hold up to 2 e-
e- must have opposite spins (up/clockwise or down/counter clockwise)
Therefore:s has up to 2 e- (1 orientation)p has up to 6 e- (3 orientations)d has up to 10 e- (5 orientations)f has up to 14 e- (7 orientations)
We can use the 2nd rule to draw Orbital Diagrams
3rd Rule: Hund’s Rulee- will not pair up until each orbital
orientation has 1 e- in itThe first e- in a pair will spin up, the
second will spin downExample: Oxygen is 1s2 2s2 2p4
1s 2s 2p
Orbital NotationOrbital Notation shows us visually the
arrangement and spin of electronsExample: Carbon is 1s2 2s2 2p2
1s 2s 2p
Energy Level DiagramsEnergy Level
Diagrams give us the same information as orbital diagrams, plus they show us the different energy levels of each orbital
Example: Carbon is 1s2 2s2 2p2 1s
2s
2p
Phosphorous, 15 e- to place
The first to electrons go into the 1s orbital
Notice the opposite spins
only 13 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
The next electrons go into the 2s orbital
only 11 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 2p orbital
• only 5 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 3s orbital
• only 3 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The last three electrons go into the 3p orbitals.
• They each go into separate shapes
• 3 unpaired electrons
• 1s22s22p63s23p3
Orbitals fill in order Lowest energy to higher energy.Adding electrons can change the energy of
the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order
Write these electron configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4
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