MOD2 - Modern Atomic Theory

atom

• Greek word “atomos” — indivisible

• Smallest particle of matter

1. Law of conservation of mass

2. Law of constant composition

3. Law of multiple proportions

The three atomic theories by John Dalton

Conservation of Mass

Dalton’s Atomic Theory

• by Antoine Lavoisier

• Mass is neither created nor destroyed

• Mass of product = Total mass of reactants

  ex. 2H2 + O2 —> 2H2O

Constant Composition

Dalton’s Atomic Theory

• aka “Law of Definite Composition”

• Relative number of atoms per element in the compound is same in any sample

  ex. Carbon dioxide in soda, dry ice, in blood

Multiple Proportions

Dalton’s

• Elements A and B of compound A+B have respective ratios w/ small whole numbers

• Different compounds from the same elements have different relative number of atoms

Cathode Ray Particles

The experiment that lead to the discovery of electrons.

subatomic particles

Many discoveries led to the fact that the atom was made of even smaller particles called ______

ex. Electrons and cathode rays, radioactivity, nucleus, protons, and neutrons

Millikan Oil-Drop Experiment

The experiment that lead to the discovery of electron charge.

Thomson’s Model of the Atom

Model of the Atom

Electrons are scattered in a positively charged space, its mass evenly distributed

Rutherford’s Gold Foil Experiment

The experiment that lead to the discovery of the positively-charged nucleus, making Thomson’s model irrelevent.

The model violates the laws of physics.

The problem of Rutherford’s Nuclear Model of the Atom

electromagnetic radiation

Moves as waves through space at the speed of light

1. Each element is composed of extremely small particles called atoms.

Dalton’s Atomic Theory

Recite the postulate associated w/ the illustration

2. All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements.

Dalton’s Atomic Theory

Recite the postulate associated w/ the illustration

3. Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

Dalton’s Atomic Theory

Recite the postulate associated w/ the illustration

4. Compounds are formed when atoms of more than one element combine: a given compound always has the same relative number and kind of atoms.

Dalton’s Atomic Theory

Recite the postulate associated w/ the illustration

J.J. Thomson

He discovered the electron

Rutherford’s Nuclear Model of the Atom

Model of the Atom

wavelength (λ)

Distance between corresponding points on adjacent waves

frequency

Number of complete waves passing any point per second

waves

All electromagnetic radiation travels as _____.

3.00 × 108 m/s

Speed of light (c) = _____

c = λν

Speed of light formula

Max Planck

He provided the explanation for thr behavior of energy.

quanta

Behavior of Energy

• Energy comes in packets called _____

• This also explains the photoelectric effect

Bohr’s Model

Behavior of Energy

• Energy of electrons are quantized

• Electrons orbit around the nucleus at discrete radii

• BUT this is only applicable to hydrogen

discrete—cannot have more or less

Meaning of quantized

Schrodinger’s Model

Behavior of Energy

• aka Quantum Model

• Used wave-particle duality to solve the behavior of atoms through equations called wave functions

• Solutions of the equations gave probabilities

• Widely accepted model today

2

Number of electrons that can fit in one orbital

Heisenberg’s Uncertainty Principle

The speed and position of the electron cannot be known simultaneously

Electrons (e-)

• Negatively charged unit of atom

• Charge of -1

• Exact location outside nucleus unknown

• Tinier than other subatomic particles

• Mass is negligible

Proton (p+)

Subatomic Particles in the Nucleus

• Positively charged unit of atom

• Charge of +1

• Does not move

• Characterizes an element

• Contributes to mass of atom

Neutron (n or n0)

Subatomic Particles in the Nucleus

• Neutral unit of atom

• Charge of 0

mass number; protons + neutrons

Atomic symbol A is called the ________ and indicates the number of _______.

atomic number; protons

Atomic symbol Z is called the ________ and indicates the number of _______.

protons (atomic number)

The number of electrons is equal to _____ unless it is an ion.

Isotopes

• Atoms of an element with varying mass number

• An atoms’s number of protons is always the same, but its number neutrons may vary

  ex. Carbon 12 and Carbon 14

Ions

• Charged atoms

• Indicated by a loss or gain of electrons

Cation (+)

• Atom with more protons than electrons

• Excess positive charge

Anion (-)

• Atom with more electrons than protons

• Excess negative charge

0

+3

-1

Fill in the missing charges

boom in the discovery of elements; Triadic; Every eighth element

History of the Periodic Table

• 18th-19th centuries: ___________

• Trends among elements:

  • _______ relationship according to atomic mass (ex. Li, Na, K)

  • __________ had similar properties (Law of Octaves)

1. First or first two letters of the element

   • O - Oxygen

   • Ca - Calcium

2. Latin name

   • Au - Gold (Aurum)

   • W - Tungsten (Wolfram)

3. Locations or people

The three bases for atomic symbol naming

amu - atomic mass unit

Unit of measurement of atomic weight

1 amu = 1.66054 × 10^-24

1 amu = _________

∑[(isotope mass)(percentage)]

Atomic weight formula for isotopes

Dmitri Mendeleev

• Father of the periodic table

• Organized the elements by atomic mass

• Observed patterns while arranging elements

Group

Period

Periodic Table

• Column: ______

• Rows: ______

reactivity

reactive

Periodicity

• Same group, same ______

• Groups on the left are more _______

Family Groups (done)

Family Groups

Just for review, explain the table in your own terms.

Photoelectric effect

• Each metal has its own energy at which it ejects electrons

• Higher energy, more electrons emitted

• Lower energy, electrons are not emitted

• Energy is proportional to frequency: E = hν

Planck’s constant = 6.626 × 10^-34 J•s

Planck’s constant = __________

Ground state

• Lowest energy level in an atom

• Nearest to the nucleus

• Most stable organization

• Where electrons are when no energy has been absorbed

Excited state

• Above ground state

• Where electrons are when the atom absorbs a quantum of energy

light

Transition of electron from higher to lower energy levels releases the absorbed energy in the form of _____.

Energy levels

Orbitals

Shells: _______

• s – 3 electrons ; n = period ; 1 orbital

• p – 6 electrons ; n = period ; 3 orbitals

• d – 10 electrons ; n = period ; 5 orbitals

• f – 14 electrons ; n = period ; 7 orbitals

Orbitals

Subshells:

• s – __ electrons ; n = __ ; __ orbital

• p – __ electrons ; n = __ ; __ orbitals

• d – __ electrons ; n = __ ; __ orbitals

• f – __ electrons ; n = __ ; __ orbitals

Principal Quantum Number (n)

Quantum Numbers

• aka the Period

• Describes the energy level on which the orbital resides

• Values are integers ≥ 1

• Correspond to the values in the Bohr model

Angular Momentum Quantum Number (l)

Quantum Numbers

• Defines the shape of the orbital.

• Values are integers ranging from 0 to (n − 1)

• Orbital group designates its different values

• Defines the shape of the orbitals

Magnetic Quantum Number (ml)

• Describes the three-dimensional orientation of the orbital

• Allowed values are integers ranging from −l to l including 0

• Where the last electron is located in Hund’s Rule

Magnetic Spin Quantum Number (ms)

• Describes its magnetic field, which affects its energy

• Reason why there are only two electrons allowed in an orbital

• either -1/2 or +1/2

• Which way the arrow faces in Hund’s Rule

  • -1/2 means arrow faces down

  • +1/2 means arrow faces up

the same energy level

the same sub shell

Energy Levels Multi-electron Atoms

• Energies differ due to repulsion of electrons

• Not all orbitals on ________ are degenerate: ns < np < nd < nf

• Orbitals in ________ are degenerate

No two electrons in the same atom can have the same set of four quantum numbers (all electrons in an atom are unique)

Pauli Exclusion definition

Electron configuration

Arrangement of electrons in the orbitals of an atom

Aufbau Principle

Electrons fill the orbital with the lowest energy first

maximized

Hund’s Rule

For degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is _______.

Valence electrons

Electrons in the outermost shell

Core electrons

electrons in the remaining shells (electrons that aren’t the valence electrons)

Effective Nuclear Charge

• Net electric field created by the nucleus and the electron density of the other electrons

• Simultaneous attraction of electrons to the nucleus and repulsion among electrons

• Smaller than the actual charge (Z)

right; down

Effective Nuclear Charge Trend:

inreases to the ______ of a period & ______ a group

Zeff = Z-S

Effective nuclear charge equation

As more electrons fill the orbital, the protons pull the electrons inward more, causing a decrease in the radius.

Explanation of atomic radius

to the bottom left

Atomic Radius Trend

smaller

Because the outermost electron is removed and repulsions between electrons are reduced

Ionic Radius

Cations are _____ than their parent atoms.

bigger

Because electrons are added and repulsions between electrons are increased

Ionic Radius

Anions are ____ than their parent atoms.

Ionization Energy

* More electrons in a subshell = Higher ionization energy = HARDER to remove an electron

Minimum energy required to be absorbed by atom to release an electron from the ground state of a gaseous atom or ion

to the top right

Ionization Energy Trend

to the top right

Ionization Energy Trend

First ionization energy

Energy required to remove the first electron

Second ionization energy

Energy required to remove the second electron

Electron Affinity

• Energy change accompanying the addition of an electron to a gaseous atom

• Exothermic in nature

  Cl + e- → Cl-

• 2A - s is full

• 5A - p is half full

• 8A - p is full

* Electron affinity value for these groups are positive

Groups excluded in electron affinity because their last subshell is full or half full (stable sila)

Cation

Metals tend to form this ion

Anion

Nonmetals tend to form this ion

Because they have the same number of valence electrons

Reason why elements in the same group have similar properties