Transcript Atoms, Molecules, and Ions
Atoms, Molecules, and Ions
Chemistry Timeline #1
B.C.
400 B.C. Demokritos and Leucippos use the term "atomos” 2000 years of Alchemy 1500's Georg Bauer: systematic metallurgy Paracelsus: medicinal application of minerals 1600's Robert Boyle:
The Skeptical Chemist.
elements Quantitative experimentation, identification of 1700s' Georg Stahl: Phlogiston Theory Joseph Priestly: Discovery of oxygen Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook
Chemistry Timeline #2
1800's Joseph Proust: The law of definite proportion (composition) John Dalton: The Atomic Theory, The law of multiple proportions Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules Amadeo Avogadro: Molar volumes of gases Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements Dmitri Mendeleyev: The periodic table J.J. Thomson: discovery of the electron Henri Becquerel: Discovery of radioactivity 1900's Robert Millikan: Charge and mass of the electron Ernest Rutherford: Existence of the nucleus, and its relative size Meitner & Fermi: Sustained nuclear fission Ernest Lawrence: The cyclotron and trans-uranium elements
Laws
• Conservation of Mass • Law of Definite Proportion – – compounds have a constant composition.
– They react in specific ratios by mass.
• Multiple Proportions – When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.
Proof
• Mercury has two oxides. – One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass.
• Show that these compounds follow the law of multiple proportion.
• Speculate on the formula of the two oxides.
Dalton’s Atomic Theory (1808) John Dalton
small particles called atoms
All matter is composed of extremely Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Atoms cannot be subdivided, created, or destroyed
Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element.
Modern Atomic Theory #2
Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.
Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded
Particle Electron Proton
Atomic Particles
Charge -1 Mass (kg) 9.109 x 10 -31 +1 1.673 x 10 -27 Location Electron cloud Nucleus Neutron 0 1.675 x 10 -27 Nucleus
The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons) Electrons are found outside of the nucleus (the electron cloud) Most of the volume of the atom is empty space “q” is a particle called a “quark”
About Quarks…
Protons and neutrons are NOT fundamental particles.
Protons are made of two “up” quarks and one “down” quark.
Neutrons are made of one “up” quark and two “down” quarks.
Quarks are held together by “gluons”
Isotopes
Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2
Atomic Masses
Atomic mass is the average of all the naturally isotopes of that element.
Carbon = 12.011
Isotope Carbon-12 Symbol 12 C Composition of the nucleus 6 protons 6 neutrons % in nature 98.89% Carbon-13 13 C 1.11% Carbon-14 14 C 6 protons 7 neutrons 6 protons 8 neutrons <0.01%
Molecules
Two or more atoms of the same or different elements, covalently bonded together.
Molecules are discrete structures, and their formulas represent each atom present in the molecule.
Benzene, C 6 H 6
Covalent Network Substances
Covalent network substances have covalently bonded atoms, but do not have discrete formulas.
Why Not??
Graphite Diamond
Ions
Cation: A positive ion
• Mg
2+ , NH 4 +
Anion: A negative ion
Cl
-
, SO 4 2
Ionic Bonding: Force of attraction between oppositely charged ions.
Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions).
Periodic Table with Group Names
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions H + Li + Na + K +
Predicting Ionic Charges
Group 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+
B 3+
Predicting Ionic Charges
Al 3+ Ga 3+ Group 13: Loses 3 electrons to form 3+ ions
Predicting Ionic Charges
Caution! C 2 2 and C 4 are both called carbide Group 14: Loses 4 electrons or gains 4 electrons
Predicting Ionic Charges
N 3 P 3 As 3 Nitride Phosphide Arsenide Group 15: Gains 3 electrons to form 3- ions
Predicting Ionic Charges
O 2 S 2 Se 2 Oxide Sulfide Selenide Group 16: Gains 2 electrons to form 2- ions
F Cl Fluoride
Predicting Ionic Charges
Chloride Br I Bromide Iodide Group 17: Gains 1 electron to form 1- ions
Predicting Ionic Charges
Group 18: Stable Noble gases do not form ions!
Predicting Ionic Charges
Groups 3 - 12: Many transition elements have more than one possible oxidation state.
Iron(II) = Fe 2+ Iron(III) = Fe 3+
Predicting Ionic Charges
Groups 3 - 12: Some transition elements have only one possible oxidation state.
Zinc = Zn 2+ Silver = Ag + Cadmium = Cd 2+
Writing Ionic Compound Formulas
Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES!
2. Check to see if charges are balanced.
Ba
2+ 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
NO
3 -
2
Not balanced!
Writing Ionic Compound Formulas
Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES!
2. Check to see if charges are balanced.
NH
4 +
2
SO
4 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!
Naming Ionic Compounds
• 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca
2+ = calcium ion
• 3. Monatomic anion = root + -ide • Cl -
= chloride
• CaCl
2 = calcium chloride
Naming Ionic Compounds
(continued) Metals with multiple oxidation states
some metal forms more than one cation use Roman numeral in name PbCl 2 Pb 2+ is the lead(II) cation PbCl 2 = lead(II) chloride
Naming Binary Compounds
Compounds between two nonmetals First element in the formula is named first. Second element is named as if it were an anion. Use prefixes Only use mono on second element P 2 O 5 CO 2 = = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide
Acids
• Substances that produce H + when dissolved in water.
• All acids begin with H.
• Two types of acids: • Oxyacids • Non-oxyacids ions
Naming acids
• If the formula has oxygen in it • write the name of the anion, but change – ate to -ic acid – ite to -ous acid • Watch out for sulfuric and sulfurous • H 2 CrO 4 • HMnO 4 • HNO 2
Naming acids
• If the acid doesn’t have oxygen • add the prefix hydro • change the suffix -ide to -ic acid • HCl • H 2 S • HCN
Formulas for acids
• Hydrofluoric acid • Dichromic acid • Carbonic acid • Hydrophosphoric acid • Nitric acid • Perchloric acid • Phosphorous acid HF H 2 Cr 2 O 7 H 2 CO 3 H 3 P HNO 3 HClO 4 H 3 PO 3
Selenium would commonly form this ion:
1. Se 2+ 2. Se + 3. Se 2 4. Sl 2 5. S 2 6. Se 7. Se 3-
0% 0% 0% 0% 0% 0% 0% 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 Se2 + 13 Se+ 14 Se2 15 Sl 2 17 S2 18 Se 19 Se3 -
Cesium would commonly form this ion:
1. Ce 2+ 2. Cs + 3. Cs 2 4. C 5. Cs 6. Cs 2+ 7. Cm +
0% 0% 0% 0% 0% 0% 0% 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 C e2 + 13 C s+ 14 C s2 15 C 17 C s 18 C s2 + 19 C m+
1 21
This is the formula for zinc hydroxide:
1. ZnOH 2. ZnOH 2 3. Zn(OH) 2 4. ZnH 2 5. Zn(OH 2 ) 2 6. Zn 2 (OH) 2 7. Zn 2 H
0% 0% 0% 0% 0% 0% 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 0% Zn OH Zn OH 2 14 Zn (O H)2 15 Zn H2 16 Zn (O H2 )2 Zn 2(O H)2 19 Zn 2H 20
This is the formula for hydrochloric acid:
1. HCl 2. HClO 3. HClO 2 4. HClO 3 5. HClO 4 6. H 2 Cl 7. H 2 ClO
0% 0% 0% 0% 0% 0% 0% 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 H Cl 13 H Cl O 14 H Cl O2 15 H Cl O3 H Cl O4 18 H 2C l 19 H 2C lO
Iron would commonly form this ion:
1. Fe 2+ 2. Fe + 3. Fe 2 4. Fe 5. Ir 2+ 6. Ir + 7. Fe 3+
Fe2 + 0% Fe+ 0% Fe2 0% 0% Fe 0% Ir 2+ 0% Ir + 0% Fe3 + 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
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Which points of Dalton’s theory are not true based on current understanding of the atom?
1. All matter is composed of extremely small particles called atoms 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties 3. Atoms cannot be subdivided, created, or destroyed 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds 5. In chemical reactions, atoms are combined, separated, or rearranged
0% 0% 0% 0% 0% 1 . A ll ma tter ..
.
2 . A to ms o f a ..
.
3 . A to ms ca nn o.
..
4 . A to ms o f d i..
.
5 . I n ch emic al ...
1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
Quantum Mechanics
The Puzzle of the Atom
Protons and electrons are attracted to each other because of opposite charges
Electrically charged particles moving in a curved path give off energy
Despite these facts, atoms don’t collapse
Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!
Toupee?
The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.
Louis deBroglie
Electromagnetic radiation propagates through space as a wave moving at the speed of light.
C
c =
= speed of light, a constant (3.00 x 10 8 = frequency, in units of hertz (hz, sec -1 ) m/s) = wavelength, in meters
Types of electromagnetic radiation:
The energy (E ) of electromagnetic radiation is directly proportional to the frequency (
) of the radiation.
E = h
E
= Energy, in units of Joules (kg·m 2 /s 2 )
h
= Planck’s constant (6.626 x 10-34 J·s)
= frequency, in units of hertz (hz, sec -1 )
Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table
Relating Frequency, Wavelength and Energy
c
E
h
Common re-arrangements:
E
hc
hc E
Spectroscopic analysis of the visible spectrum… …produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen spectrum…
…produces a “bright line” spectrum
Electron transitions involve jumps of definite amounts of energy.
This produces bands of light with definite wavelengths.
Bohr Model Energy Levels
Electron Energy in Hydrogen
E electron
2 .
178
x
10 18
J
Z
2
n
2
Z n
= nuclear charge (atomic number) = energy level ***Equation works only for atoms or ions with 1 electron (H, He + , Li 2+ , etc).
Calculating Energy Change,
E, for Electron Transitions
E
2 .
178
x
10 18
J Z
2
n
2
final
-
Z
2 2
n initial
Energy must be absorbed from a photon ( +
E ) to move an electron away from the nucleus Energy (a photon) must be given off ( -
E ) when an electron moves toward the nucleus
Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
Principal quantum number (n) Angular momentum quantum number (l) Magnetic quantum number Spin quantum number (s) (m)
Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli
Principal Quantum Number
Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell:
2n
2
Angular Momentum Quantum Number
The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.
l =3 f
Magnetic Quantum Number
The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.
Assigning the Numbers
integers.
The three quantum numbers (n, l, and m) are The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc.
can be any integer between 0 and n - 1.
The angular momentum quantum number (l ) For n = 3, l can be either 0, 1, or 2.
The magnetic quantum number (m
l
) can be any integer between -l and +l.
For l = 2, m can be either -2, -1, 0, +1, +2.
Principle, angular momentum, and magnetic quantum numbers: n, l, and m
l
Spin Quantum Number
Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:
1 2 1 2
An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
Orbital shapes are defined as the surface that contains 90% of the total electron probability.
Schrodinger Wave Equation
8
h
2 2
d
2
m dx
2
V
E
Erwin Schrodinger Equation for probability of a single electron being found along a single axis (x-axis)
Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg You can find out where the electron is, but not where it is going.
OR… You can find out where the electron is going, but not where it is!
Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital.
Orbitals in outer energy levels DO penetrate into lower energy levels.
Penetration #1 This is a probability Distribution for a 3s orbital.
What parts of the diagram correspond to “nodes” – regions of zero probability?
Which of the orbital types in the 3 rd Does not seem to have a “node”?
energy level WHY NOT?
Penetration #2
The s orbital has a spherical shape centered around the origin of the three axes in space.
s orbital shape
P orbital shape
There are three peanut-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.
d orbital shapes Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of: “double peanut” …and a “peanut with a donut”!
Shape of f orbitals Things get even more complicated with the seven f orbitals that are found in the f sublevels beginning with n = 4. To remember the shapes, think of: Flower
Element Lithium Configuration notation 1s 2 2s 1 Beryllium 1s 2 2s 2 Boron Carbon 1s 2 2s 2 p 1 1s 2 2s 2 p 2 Nitrogen 1s 2 2s 2 p 3 Oxygen Fluorine Neon 1s 2 2s 2 p 4 1s 2 2s 2 p 5 1s 2 2s 2 p 6 Orbital notation Noble gas notation [He]2s 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 3 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 4 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 5 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 6 ____ ____ ____ ____ ____ 1s 2s 2p
Orbital filling table
Electron configuration of the elements of the first three series
Irregular confirmations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel
In Bohr’s atomic theory, when an electron moves from one energy level to another energy level more distant from the nucleus.
1 21
1. energy is emitted 2. energy is absorbed 3. no change in energy occurs 4. light is emitted 5. none of these
en er gy is emit en ted er n gy o is ch a bs an or ge bed in en er gy o cc ur li s gh t i s emit ted n on e of th es e 15 16 17 18 19 20 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 0% 0% 0% 0% 0%
Which form of electromagnetic radiation has the longest wavelengths?
1. gamma rays 2. microwaves 3. radio waves 4. infrared radiation 5. x-rays
0% 0% 0% 0% 0% ga mma ra ys mi cr ow av es ra di o w av in es fra re d ra di at io n x ra ys 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
How many electrons in an atom can have the quantum numbers n = 3, l = 2?
1. 2 2. 5 3. 10 4. 18 5. 6
2 0% 5 0% 0% 10 0% 18 0% 6 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
Which of the following combinations of quantum numbers is not allowed?
n
1. 1 2. 3 3. 2 4. 4 5. 4
l
1 0 1 3 2
m
0 0 –1 –2 0
s
½ –½ ½ –½ ½
0% 0% 0% 0% 0% 0% n lms 1 10 ½ 3 00 –½ 2 1– 1½ 4 3– 2– ½ 4 20 ½ 1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
The electron configuration of indium is
1. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 5d 10 2. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4d 10 4p 1 3. 1s 2 3s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 6 5s 2 5d 10 5p 1 4. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 5. none of these
1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 0% 0% 0% 0% 0% 16 2s 22 p6 3s 23 1 s2 p6 2s 4s 22 23 p6 17 d1 3s ...
23 1 s2 p6 3s 4s 22 23 p6 d1 3s ...
23 18 p6 s2 2s 4s 24 22 p6 d1 3s ...
23 p6 4s 19 23 d1 ...
n on e of th 20 es e
Ag has __ electrons in its d orbitals.
1 21 2 22 3 23 4 24 5 25 6 26 7 27 8 28 9 29 10 30 11 12 13 14 15 16 17 18 19 20
Periodicity
Determination of Atomic Radius: Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels
Table of Atomic Radii
Ionization Energy - the energy required to remove an electron from an atom
Increases for successive electrons taken from the same atom
Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels
Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Outer electrons are farther from the nucleus
Ionization of Magnesium Mg + 738 kJ
Mg + + 1451 kJ
Mg + Mg + e 2+ + e Mg 2+ + 7733 kJ
Mg 3+ + e -
Table of 1 st Ionization Energies
Another Way to Look at Ionization Energy
Yet Another Way to Look at Ionization Energy
Electron Affinity - the energy change associated with the addition of an electron
Affinity tends to increase across a period
Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals
Table of Electron Affinities