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ATOMS & THE PERIODIC
TABLE
Subatomic Particles
• Protons and electrons are the only particles that have a charge.
• Protons and neutrons have essentially the same mass.
• The mass of an electron is so small we ignore it.
Atoms and the Periodic Table
Subatomic Particle
Location Charge Size Mass
Protonp+
Nucleus Positive1+
About same as neutronD = 10-15 m
1 amu1.6726 x 10-24 g
Neutronn0
Nucleus No chargeNeutral
0
About same as protonD = 10-15 m
1 amu1.6749 x 10-24 g
Electrone-
Orbital cloud
Negative1-
Tiny compared to
proton & neutron
D = 10-18 m
1/1840 amu9.11 x 10-28 g
Size of Atoms
Particle Diameter (meters)atom 10-10
nucleus 10-14
proton 10-15
neutron 10-15
electron 10-18
1 Angstrom = 10-10 m
History of the Atomic ModelDemocritus – 400 BC
atoms make up all substances
John Dalton – 1766-1844atom is a solid hard sphere
Joseph John Thomson – 1856-1940discovered the electron in 1897plum pudding model of atompositive sphere with negative e- embedded
• Lived from (1766-1844)• All elements are composed of atoms• Atoms of the same element are identical.
Each element has unique properties .• Atoms of different elements can be
chemically combined in simple whole number ratios to form compounds.
• Law of Conservation of Matter/Mass
The Electron
• Streams of negatively charged particles were found from cathode tubes.
• J. J. Thompson is credited with their discovery (1897).
The Atom, circa 1900:
• “Plum pudding” model, put forward by Thompson.
• Positive sphere of matter with negative electrons imbedded in it.
Ernest Rutherford – 1871-1937gold foil experimentmost of atom empty spacepositive nucleus contains most of the massdiscovered protons in 1919
James Chadwick – 1891-1974discovered the neutron in 1932.
Niels Bohr – 1885-1962 (Danish)electrons move around the nucleus in fixed
orbits that have a set amount of energy
Electron Cloud Model of Atom1. came to be used to estimate the positions of electrons in an atom2. uncertainty principle, which states that it is not possible to obtain precise values of both position and momentum of a particle at the
same time3. probability of finding an electron
Protons• The number of protons distinguished 1 atom
from another• Most atoms are very stable• It takes a lot of energy to add or remove a
proton from an atom• Atomic number = number of protons• The Periodic Table is arranged by number of
protons
Symbols of Elements
Elements are symbolized by one or two letters.
Atomic Number
Number of protons = The atomic number
Atomic Mass
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.
Mass Number
• The number of protons plus neutrons in an atom.
• Always a whole number.• Written or indicated like this:
H C Cu Si K
Calculating Number of Neutrons
• Subtract:
Mass Number - Atomic Number = Neutrons
Mass Number - # of protons = Neutrons
Notes on Finding Numbers of Protons, Neutrons, Electrons
Isotopes• Atoms of the same element with different
numbers of neutrons
Isotopes
• To distinguish between isotopes of an element
• Ex: Neon has 3 main isotopes
Neon Protons Neutrons Mass number
Ne-20 10 10 20Ne-21 10 11 21Ne-22 10 12 22
Average Atomic Mass
• Atomic mass unit – 1/12 the mass of a C-12 atom
• To calculate avg. atomic mass
Cesium Natural % Abundance
Mass Number
Avg. Atomic Mass
Cs-133 75% 133 99.75
Cs-132 20% 132 26.4
Cs-134 5% 134 6.7
Example – Avg. Atomic Mass
Chemical Goals
• To be Chemically StableUnstable atoms are radioactive: their nuclei change or decay by spitting out radiation, in the form of particles or electromagnetic waves.
• To be Electronically neutralTo have no charge on the atom. To have the same # of protons as electrons.
Why does an atom stay together?
The strong nuclear force keeps protons and neutrons together in the nucleus in spite of the repulsion of the protons for each other.
The strong nuclear force acts only over a very short distance. It doesn’t work outside the nucleus.
The strong nuclear force is stronger than the electromagnetic force.
Valence Electrons• Electrons in outer energy level.• Can only have 8 or less. This is the Octet Rule.• These electrons are the ones involved in
bonding with other atoms and the ones with the most energy.
• Looking at the Periodic Table, you can tell the number of valence electrons for the A Families from the Roman Numeral designation
Electrons
1. e- located far from nucleus in a series of energy levels.
2. Each e- has a certain amount of energy.
3. The further the e- gets from the nucleus the more energy it has. Valence e- have the most energy. Those closest to the nucleus have the least amount of energy.
Energy Levels
1. Each energy level can only hold a certain number of e-
2. Electrons always fill low energy orbitals (closest to
the nucleus) before filling higher energy ones.
3. The high the energy level occupied by the e-, the easier it is for the e- to escape from the atom.
4. Quantum of energy – amount of energy needed to move an e- from its present energy level to the next one
5. Ground State – the lowest energy level for
an e-.
Electron Placement on the Periodic Table
2 e-
8 e-
18 e-
32 e-
Type of Sublevel
Number of orbitals in Sublevel
Number of electrons in
orbital
s 1 2
p 3 6
d 5 10
f 7 14
Principal Energy Level
(n)
Number of Sublevels
Type of Sublevel(Atomic Orbitals)
Number of Electrons
2n2
1 1 s 2
2 2 s, p s-2, p-6 = 8
3 3 s, p, d s-2, p-6,d-10 = 18
4 4 s, p, d, f s-2, p-6,d-10, f-14
= 32
Electron Configuration Chart
Aufbau /Orbital Diagram
EX: Electron Configuration for Potassium 19 electrons
EX: Electron Configuration for Arsenic 33 electrons
EX: Electron Configuration for Silver 47 electrons
Pauli Exclusion Principle
• An orbital can hold 0, 1, or 2 electrons and if there are 2 electrons in the orbital they must have opposite spin.
Aufbau Principle
• Rules for Orbital Filling• Lower-energy orbitals fill first.• An orbital can hold only 2 e- with opposite
spins.• The most stable arrangement for e- is one
with the maximum number of unpaired e-. This minimized e- to e- repulsion and stabilizes the atom. – Hund’s Rule
Hund’s Rule
• When filling up sublevels other than s, electrons are placed in individual orbitals before they are paired up.
Increasing Energy in Electron Sublevels
THE DIAGONAL RULE MUST GO IN THIS ORDER: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d and 7p. These orbitals will account for all the elements now known. This diagonal rule can help account for the octet rule, too.
The Energy Flow
• The order of increasing energy of the orbitals is then read off by following these arrows, starting at the top of the first line and then proceeding on to the second, third, fourth lines, and so on. This diagram predicts the following order of increasing energy for atomic orbitals.
• 1s < 2s < 2p < 3s < 3p <4s < 3d <4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p < 8s ...
Bohr Model of Atomsonly represents energy levels, not orbitals
Lithium
Nobel GasesNeon
Argon
Krypton
Drawing AtomsThe Bohr Model
e- 3 p+
11 p+
Lewis StructuresElectron Dot Diagrams
• Describes e- arrangement in atoms • Describes bond arrangement in molecules.• Uses dots to represent valence e- around an
atom• EX:
Li Ne O Si
Dimitri Mendeleev• In the late 1860's, Mendeleev
began working on his great achievement: the periodic table of the elements. By arranging all of the 63 elements then known by their atomic weights, he managed to organize them into groups possessing similar properties. Where a gap existed in the table, he predicted a new element would one day be found and deduced its properties. And he was right. Three of those elements were found during his lifetime: gallium, scandium, and germanium.
Mendeleev’s Periodic Table
Moseley’s Periodic Table
In 1913 Henry Moseley came up with this Periodic Table. The elements are arranged by increasing atomic number.
• A group (also known as a family) is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table.
• Elements in a group have similar configurations of the outermost electron of their atoms – same number of valence e-
• This gives the groups of elements similar physical and chemical characteristics.
• With each group across a period, the elements have one more proton and electron and become less metallic.
• Rows of elements are called periods. The period number of an element signifies the highest unexcited energy level for an electron in that element.
Physical Properties
Metals• Good electrical conductors
and heat conductors. • Malleable - can be beaten
into thin sheets. • Ductile - can be stretched
into wire. • Possess metallic luster. • Opaque as thin sheet. • Solid at room temperature
(except Hg).
Nonmetals• Poor conductors of heat
and electricity. • Brittle - if a solid. • Nonductile. • Do not possess metallic
luster. • Transparent as a thin sheet. • Solids, liquids or gases at
room temperature.
Chemical Properties
Metals• Usually have 1-3 electrons
in their outer shell. • Lose their valence electrons
easily. • Have lower
electronegativities.
Nonmetals• Usually have 4-8 electrons
in their outer shell. • Gain or share valence
electrons easily. • Have higher
electronegativities.
Metalloids
• Electronegativities and ionization energies between those of metals and nonmetals
• Possess some characteristics of metals/some of nonmetals
• Reactivity depends on properties of other elements in reaction
• Often make good semiconductors• Boron, Silicon, Germanium, Arsenic, Antimony,
Tellurium, Polonium
Ionic BondingBetween Metals & Nonmetals
Metals• Make Ionic Compounds• Lose Electrons• Have positive oxidation
numbers• Are first in a formula
Ex: Na2S
• When naming, write the name just as it is
• Ex: Sodium sulfide
Nonmetals• Make Ionic Compounds• Gain Electrons• Have negative oxidation
number• Are second in the formula
Ex: MgO• When naming, drop the
ending of the name and add IDE
• Ex: Magnesium oxide
More Ionic Bonding
• Strong attractions or forces between atoms in these compounds
• High melting and boiling points, good conductors
• Have a lattice structure
Sodium chloride Lattice
Covalent BondingCovalent Molecules
• Between 2 or more nonmetals• Share electrons• Still try to obey the Octet Rule• Weak bonds between molecules but strong
bonds between atoms• Low melting and boiling points, usually non
conductors• Simple molecules or giant structures can form
More Covalent Bonding
• Diamond and graphite are held together by this type of bond (allotropes)
• Ex: H2O, H2O2, CH4, CO2
• Ex: Diatomics – H2, N2, O2, F2, Cl2, Br2, I2
Water Molecule
Moles•The mole is the SI unit for amount of substance
•A mole (abbreviated mol) is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.
Avogadro’s Number
• 6.02 x 1023 – is the number of particles in exactly one mole of a pure substance.
• Conversion factor!• 6.02 x 1023 particles = 1 mole
If I have 3.45 moles of hydrogen, how many particles do I have?
6.02 x 1023 = 1 Mole