atoms & the periodic table. 2 chapter outline what is atom? chemical properties of atoms: the...
TRANSCRIPT
Atoms & the Periodic Table
2
Chapter Outline
• What is Atom?
• Chemical properties of Atoms: the Periodicity
• Isotopes
• Electrons in Atom: Quantum physics’ view
• Valence electrons and the Periodic Table• Periodic trend: Atomic Radius, Metallic
Character, Ionization Energy
3
Experiencing Atoms• Atoms: incredibly small, yet compose everything• atoms are the pieces of Elements• properties of the atoms determine the properties
of the elements
4
Within an Atom• Atoms = (Protons +
NeutronsNeutrons) + ElectronsElectrons• The nucleus (Protons +
NeutronsNeutrons) is only about 10-13 cm in diameter yet with most of the mass of the atom
• The electrons move outside the nucleus with an average distance of about 10-8 cm
• The atom is electrically neutral : #proton (#p) = #electron (#e) Nucleus
Proton
Neutron
Electron
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Subatomic
Particle Mass
g
Mass
amu
Location
in atom
Charge Symb
ol
Proton 1.67 x 10-24 1 nucleus +1 p, p+, H+
Electron 9 x 10-28 ~0 empty space -1 e, e-
Neutron 1.67 x 10-24 1 nucleus 0 n, n0
Comparison among Proton, Electron, Neutron
6
Elements• each Element has a unique number of protons
in its nucleus
• Atomic number: the number of Protons in the nucleus of an atomthe elements are arranged on the Periodic Table
in order of their atomic numbers
• Name and Symbol of an Elementsymbol either one or two letters
one capital letter or one capital letter + one lower case
8
The Size of Atoms
Atomic Mass Unit (amu):
1 amu = 1.66 10-24 g
Hydrogen the smallest atom
mass of H atom= 1.67 x 10-24g ~ 1 amu
volume of H atom = 2.1 x 10-25cm3
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Isotopes
The same element could have atoms with different masses
Examples:
• 2 isotopes of chlorine atoms in nature: one weighs about 35 amu (Cl-35); another weighs about 37 amu (Cl-37)
• Carbon-12 (C-12) is much more abundant than C-13.
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Isotopesall isotopes of an element: • chemically identical
undergo the exact same chemical reactions• the same number of protons• different masses due to different numbers of
neutrons. Example: C-14 atom has eight neutrons; C-12 atom has six neutrons.
• identified by their mass numbersprotons + neutrons
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• Atomic Number (Z)Number of protons
• Mass Number (A)Protons + Neutrons
Abundance = relative amount found in a sample
Example: Cl-35 (75%) vs. Cl-37 (25%)
Isotopes
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Isotopic Symbol
• Cl-35 has a mass number = 35, 17 protons and 18 neutrons (35 - 17). The symbol for this isotope would be
Atomic SymbolA = mass numberZ = atomic number#neutrons = A - Z
AXZ
Cl3517
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Mass number = Atomic number (# protons, or #p) + #neutrons
U = uraniumAtomic Number = 92
#p = atomic number = 92#e = #p = 92
Mass Number = #p + #n238 = 92 + #n
146 = #n
Isotopic symbol element atomic number #p #e #n
Example:How many protons, neutrons, and electrons in an atom of
23892 U
#proton = 92#neutron = 146#electron = 92
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Mass Number is Not the Sameas Atomic Mass
• Atomic mass (or Atomic Weight) is an experimental number determined from all naturally occurring isotopes
• Mass number refers to the number of protons + neutrons in one isotopenatural or man-made
When given the relative abundance of all isotopes, we can find the Atomic mass
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The Modern Periodic Table
• Elements with similar chemical and physical properties are in the same column
• columns are called Groups or Familiesdesignated by a number and letter at top
• rows are called Periods
• each period shows the pattern of properties repeated in the next period
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The Modern Periodic Table
• Main Group Main Group = Representative Elements = ‘A’ groups
• Transition Metals Transition Metals = ‘B’ groupsAka Transition elements
• Inner Transition Elements = Bottom rows = Rare Earth Elementsmetalsreally belong in Period 6 & 7
Main group vs. Transition metals, Inner transition metals
= Metal
= Metalloid
= NonmetalIA
IIIB
IIIAIIA
VIIB VIIIB IB IIB
VIIIA
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Metals: Physical vs. Chemical Properties
• solids at room temperature, except Hg
• reflective surfaceshiny
• conduct heat, electricity• Malleable (can be shaped)
• Tend to Lose electrons and form Cations in reactions. Na Na+ + e -
• about 75% of the elements are metals
• lower left on the table
19
Nonmetals: Physical vs. Chemical Properties
• Elements found in all 3 states
• poor conductors of heat or electricity
• solids are brittle• Tend to gain electrons in
reactions to become anions: Cl + e - Cl-
• upper right on the tableexcept H
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Metalloids: between Metals and Nonmetals
• show some properties of metals and some of nonmetals
• also known as semiconductors Properties of Silicon
shinyconducts electricity
does not conduct heat wellbrittle
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= Alkali Metals
= Alkaline Earth Metals
= Noble Gases
= Halogens
= Lanthanides
= Actinides
= Transition Metals
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= Rare Earth Metals
= Transuranium element
= Transition Metals
U
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Important Element - Hydrogen• nonmetal• colorless, diatomic gas H2
very low melting point & density
• reacts with Nonmetals to form molecular compoundsHCl is acidic gasH2O is a liquid
• reacts with Metals to form hydridesmetal hydrides react with water to form H2
Nickel-metal hydride (NiMH) used in rechargeable battery
• HX dissolves in water to form acids
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Important Element - CarbonThree forms of pure carbon:•Diamond: hardest substance in nature•Graphite: soft and slippery solid•Buckminsterfullerene: a molecule made of 60(Images from public domain wikipedia.com)• carbon atoms in a sphere
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Carbon as backbone for Organic/Biochemical Molecules
Carbon atoms capable of forming robust bonds with many other elements and themselves.
Examples:•Small molecules: Butane, Sugar, Fatty acid, Vitamins•Big molecules (Polymers): Starch, Kevlar, Teflon, Protein, and DNA
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Group IA: Alkali Metals
• Usually Hydrogen is included• All metals: soft, low melting
points• Flame tests Li = red, Na =
yellow, K = violet
Chemical Property: • Very reactive.
React with water to form basic (alkaline) solutions and H2.
releases a lot of heat• Tend to form water soluble
compounds, such as table salt and baking soda.
colorless solutions
lithium
sodium
potassium
rubidium
cesium
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Group IIA: Alkali Earth Metals
Physical properties: harder, higher melting, and denser than alkali metals
• flame tests Ca = red, Sr = red, Ba = yellow-green
Chemical properties:• reactive, but less than
corresponding alkali metal• form stable, insoluble oxides.
oxides are basic = alkaline earth• reactivity with water to form H2,
Be = none; Mg = steam; Ca, Sr, Ba = cold water
magnesium
calcium
beryllium
strontium
barium
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Group VIIA: Halogens• nonmetals
• F2 & Cl2 gases; Br2 liquid; I2 solid
• all diatomic
• very reactive
• Cl2, Br2 react slowly with water
Cl2 + H2O HCl + HOCl
(chlorine)
• react with metals to form ionic compounds
• HX all acidsHF weak < HCl < HBr < HI
bromine
iodine
chlorine
fluorine
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Group VIIIA: Noble Gases
• all gases at room temperature, very low melting and
boiling points
• very unreactive, practically inert
• very hard to remove electron from or give an electron to
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Quantum Physicists including Schrödinger:
• Electrons move very fast around the nucleus
• Electrons show up with a particular probability at certain location of the atom
• Orbital: A region where the electrons show up a very high probability when it has a particular amount of energygenerally set at 90 or 95%
Atomic Orbitals
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Electron Shells
• Electron Shell: the main energy level for the orbital. Principal quantum number n = 1, 2, …
For a chlorine atom, three shells of electrons:•The innermost shell (n = 1, 2 electrons) has the lowest energy•The outmost shell (n = 3, 7 electrons) has the highest energy
Cl (17p+ & 17e-)
17 p+
2e-
8e-
7e-
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Each Shell have Subshells
• Each Electron Shell has one or more Subshells (s, p, d, f) the number of subshells = the Principal quantum number n
n = 1, one subshell (1s);
n = 2, two subshells (2s, 2p)
n = 3, three subshells (3s, 3p, 3d)
n = 4, three subshells (4s, 4p, 4d, 4f)
• each Subshell has orbitals with a particular shape the shape represents the probability map
90% probability of finding electron in that region
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Shapes of Subshells
p Orbitals: px , py , pzs Orbital
d Orbitals
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f orbitals
34Tro: Chemistry: A Molecular Approach, 2/e
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Shapes of f orbitals: 4f orbitals(downloaded from public domain)
The coloration corresponds to the sign of function.
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Shells & Subshells
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How does the 1s Subshell Differ from the 2s Subshell?
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Subshells and Orbitals• Among the subshells of a principal shell, slightly different
energies for multielectron atoms, the subshells have different
energies: s < p < d < f
• each subshell contains one or more Orbitalss : 1 orbitalp : 3 orbitalsd : 5 orbitalsf : 7 orbitals
within one subshell, different orbitals have the same energy. Example: 2px, 2py and 2pz
39
Electron ConfigurationsDefinition: The distribution of electrons into the
various energy shells (n = 1,2,3,…) and subshells (s, p, d, f) in an atom in its ground state
• Each energy shell and subshell has a maximum number of electrons it can holdSubshell s = 2, p = 6, d = 10, f = 14Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e
• Electrons fill in the energy shells and subshells in order of energy, from low energy upAufbau Principal (“Construction” in German)
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Ene
rgy
1s
7s
2s
2p
3s
3p3d
6s6p
6d
4s
4p4d
4f
5s
5p
5d5f
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Order of Subshell Fillingin Ground State Electron Configurations
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
1. Diagramputting each energy shell ona row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right)
2. draw arrows throughthe diagonals, looping back to the next diagonaleach time
42
Spinning Electron(s) in Orbital• Experiments showed Electrons spin on an axis
generating their own magnetic field
Pauli Exclusion Principle • each Orbital may have a maximum of 2 electrons,
with opposite spin• Two electrons sharing the same orbital must have
Opposite spinsso there magnetic fields will cancelanalogous to two bar magnets in parallel: only opposite
alignment could stabilize each other.
43
Orbital Diagrams• often an orbital as a square • the electrons in that orbital as arrows
the direction of the arrow represents the spin of the electron
orbital with1 electron
unoccupiedorbital
orbital with2 electrons
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Filling the Orbitals in a Subshellwith Electrons
• Energy shells fill from lowest energy to high1 → 2 → 3 → 4
• Subshells fill from lowest energy to highs → p → d → f
• Orbitals of the same subshell have the same energy. Three 2p orbitals; Five 3d orbitals
Electrons prefer “spreading out” in orbitals of same subshell before they pair up in orbitals.Hund’s RuleExample: 2p3 _ _ _ instead of ____
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Electron Configuration of Atoms in their Ground State
• Electron configuration: a listing of the subshells in order of filling with the number of electrons in that subshell written as a superscript
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• a shorthand way : use the symbol of the previous noble gas in [] for the inner electrons, then just write the last set
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1
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1s22s22p63s2 = [Ne]3s2
1s 2s 2p 3s 3p
Example: Ground State Orbital Diagram and Electron Configuration of Magnesium
47
Practice: Write Electron Configuration for the following atoms at the Ground state
• Calcium• Sulfur• Sodium• Chlorine
Important: you are required to be able to write electron configuration for first three rows of elements
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Valence Electron vs. Core Electron
Valence Electron: the electrons in all the subshells with the highest principal energy shell
• Example: electrons in bold Mg = [Ne]3s2 O = [He]2s22p4
Br = [Ar]4s23d104p5
• Core electrons: electrons in lower energy shells • Chemists have observed that one of the most important
factors in the way an atom behaves, both chemically and physically, is the Number of Valence electrons
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Valence Electrons
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
• the highest principal energy shell that contains electrons is the 5th : 1 valence electron + 36 core electrons
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• the highest principal energy shell that contains electrons is the 4th : 8 valence electrons + 28 core electrons
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Electrons Configurations andthe Periodic Table
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Electron Configurations fromthe Periodic Table
Example: Be 2s2 B 2s22p1 C 2s22p2 N 2s22p3 O 2s22p4
• Elements in the same period (row) have Valence Electrons in the same principal energy shell.
• #Valence electrons increases by one from left to right
• Elements in the same group have the same #valence electron and they are same kind of subshell
Example: IIA: Be 2s2 Ca 3s2 Sr 4s2 Ba 5s2
• VIIA: F 2s22p5 Cl 3s23p5 Br 4s24p5 I 5s25p5
52
Electron Configuration & the Periodic Table
• Elements in the same Group have similar chemical and physical properties their valence shell electron configuration is the same
• No. Valence electrons for the main group elements is the same as the Group Number
Example: • Group IA: ns1 ; • Group IIIA: ns2np1 • Group VIIA: ns2np5
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s1
s2
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
p1 p2 p3 p4 p5 s2
p6
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1234567
Electron Configuration & the Periodic Table
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Electron Configuration from the Periodic Table
• Inner electron configuration = Noble gas of the preceding period
• Outer electron configuration: from the preceding Noble gas the next period (Subshells) Elementthe valence energy shell = the period numberthe d block is always one energy shell below the
period number and the f is two energy shells below
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Electron configuration & Chemical Reactivity
• Chemical properties of the elements are largely determined by No. Valence electrons
• Why elements in groups? Since elements in the same column have the same #valence electrons, they show similar properties
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Electron Configuration: Noble Gas
• Noble gases have 8 valence electronsexcept for He, which has only 2
electrons• Noble gases are especially nonreactive
He and Ne are practically inert
The reason: the electron configuration of the noble gases is especially stable
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Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A)
• have one more electron than the previous noble gas, [NG]ns1
• tend to lose their extra ONE electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ chargeNa Na+ Li Li+
58
Everyone Wants to Be Like a Noble Gas!Halogens (Group 7A)
• one fewer electron than the next noble gas: [NG]ns2np5
• Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns2np5 + 1e [NG]ns2np6
forming an anion with charge 1-: Cl Cl-
• Reactions with Nonmetals: tend to share electrons so that each attains the electron configuration of a noble gas
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Everyone Wants to Be Like a Noble Gas!Summary
• Alkali Metals as a group are the most reactive metals they react with many things and do so rapidly
• Halogens are the most reactive group of nonmetals• one reason for their high reactivity: they are only
ONE electron away from having a very stable electron configuration the same as a noble gas
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Stable Electron ConfigurationAnd Ion Charge
• Metals: Cations by losing enough electrons
to get the same electron configuration as the previous noble gas
• Nonmetals: Anionsby gaining enough
electrons to get the same electron configuration as the next noble gas
Atom Atom’s Electron Config
Ion Ion’s Electron Config
Na [Ne]3s1 Na+ [Ne]
Mg [Ne]3s2 Mg2+ [Ne]
Al [Ne]3s23p1 Al3+ [Ne]
O [He]2s2p4 O2- [Ne]
F [He]2s22p5 F- [Ne]
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Trends in Atomic SizeIncreases down a group
valence shell farther from nucleuseffective nuclear charge fairly close
Decreases across a period (left to right)adding electrons to same valence shelleffective nuclear charge increasesvalence shell held closer
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Trends in Atomic Size
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Metallic Character• Metals
malleable & ductileshiny, lusterous, reflect lightconduct heat and electricitymost oxides basic and ionic form cations in solution lose electrons in reactions – oxidized
• Nonmetalsbrittle in solid statedullelectrical and thermal insulatorsmost oxides are acidic and molecular form anions and polyatomic anionsgain electrons in reactions - reduced
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Trends in Metallic Character
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Electron Configuration Affects the Size of Atoms and Metallic Character:
Within a Group• Within the same Group, from top to
bottom:As quantum number n increases for the valence
electron(s) valence electron(s) further away from the
nucleus Larger Atomic Radius weaker Coulombic force (electrostatic force)
withholding valence electrons electrons easier to be lost Stronger metallic character
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4 p+
2e-
2e-
12 p+
2e-
8e-
2e-
Be (4p+ & 4e-)
Example: Group IIA
Mg (12p+ & 12e-)
Ca (20p+ & 20e-)
20 p+
2e-
8e-
2e-
8e-
2r
qqkF
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Electron Configuration Affects the Size of Atoms and Metallic Character:
Over the Period
• Within the same Period (row), from left to right:Same quantum number n for the valence
electron(s)As Nucleus has increasing number of protons (p+) Stronger Coulombic force (electrostatic force)
withholding valence electrons Valence Electrons closer the nuclues Smaller Atomic RadiusValence electrons harder to be lost Weaker metallic character
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Li (3p+ & 3e-)
Example: Period 2
Be (4p+ & 4e-) B (5p+ & 5e-)
6+2e-4e-
C (6p+ & 6e-)
8+2e-6e-
O (8p+ & 8e-)
10+2e-8e-
Ne (10p+ & 10e-)
2e-1e-
3+2e-2e-
4+2e-3e-
5+
2r
qqkF
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Ionization Energy (IE)
In an atom, electrons (“-” charge) are attracted to the nucleus (“+” charge). Energy is required to remove the electron from an atom.
Na + energy Na+ + e-
Neutral atom IE Cation
Higher IE corresponds to lower Metallic property.
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Trends in Ionization EnergyDecreases down a group
valence shell farther from nucleuseffective nuclear charge fairly close
Increases across a period (left to right)adding electrons to same valence shelleffective nuclear charge increasesvalence shell held closer
71
Practice: Rank elements K, Mg, S, F:A. increasing metallic characterB. increasing atomic radiiC. increasing ionization energy
F, S, Mg, K
F, S, Mg, K
K, Mg, S, F