Chapter 2

The Structure of the Atom and the Periodic Table

2.1 Composition of the Atom

• Atom - the basic structural unit of an element

• The smallest unit of an element that retains the chemical properties of that element

• Nucleus - small, dense, positively charged region in the center of the atom

protons - positively charged particles

neutrons - uncharged particles

Electrons, Protons and Neutrons• Atoms consist of three primary particles

• electrons• protons• neutrons

Characteristics of Atomic Particles

• Electrons are negatively charged particles located outside of the nucleus of an atom

• Protons and electrons have charges that are equal in magnitude but opposite in sign

• A neutral atom that has no electrical charge has the same number of protons and electrons

• Electrons move very rapidly in a relatively large volume of space while the nucleus is small a dense

Mass

Number

Atomic Number

Charge of particle

Symbol of the atom

Symbolic Representation of an Element

CAZ X

• Atomic number (Z) - the number of protons in the atom

• Mass number (A) - sum of the number of protons and neutrons

Atomic Calculations

number of protons + number of neutrons = mass number

number of neutrons = mass number - number of protons

In a neutral atom, the number of protons = number of electrons

Very Important: the number of protons does not change!

Name Charge Mass(amu) Mass (grams)

Electrons (e) -1 5.4 x 10-4 9.1095 x 10-28

Protons (p) +1 1.00 1.6725 X 10-

24

Neutrons (n) 0 1.00 1.6750 x 10-24

Selected Properties of the Three Basic Subatomic Particles

Calculate the number of protons, neutrons and electrons in each of the following:

B115

Fe5526

Determining the Composition of an Atom

4

Hydrogen(Hydrogen - 1)

Deuterium(Hydrogen - 2)

Tritium(Hydrogen - 3)

Isotopes of Hydrogen

• Isotopes - atoms of the same element having different masses contain same number of protons contain different numbers of neutrons

Isotopes

Isotopic Calculations

• Isotopes of the same element have identical chemical properties

• Some isotopes are radioactive

• Find chlorine on the periodic table

• What is the atomic number of chlorine?

• What is the mass given?

• This is not the mass number of an isotope

Atomic Mass• Find chlorine (Cl) on the table. What is this

number, 35.34?

• The atomic mass - the weighted average of the masses of all the isotopes that make up chlorine

• Chlorine consists of chlorine-35 and chlorine-37 in a 3:1 ratio

• Weighted average is an average corrected by the relative amounts of each isotope present in nature

Determining Atomic Mass

Calculate the atomic mass of naturally occurring chlorine if 75.77% of chlorine atoms are chlorine-35 and 24.23% of chlorine atoms are chlorine-37

Step 1: convert the percentage to a decimal fraction

0.7577 chlorine-35

0.2423 chlorine-37

Step 2: Multiply the decimal fraction by the mass of that isotope to obtain the isotope contribution to the atomic mass.

For chlorine-35:0.7577 x 35.00 amu = 26.52 amu

For chlorine-370.2423 x 37.00 amu = 8.965 amu

Step 3: sum these partial weights to get the weighted average atomic mass of chlorine:

26.52 amu + 8.965 amu = 35.49 amu

It’s just like calculating your grade!

• 20% of the quiz average• 60% of the exam average• 20% of the final exam

Score = 0.2*quiz ave. + 0.6*exam ave. + 0.2*final

Atomic Mass Determination• Nitrogen consists of two naturally occurring

isotopes99.63% nitrogen-14 with a mass of 14.003 amu

0.37% nitrogen-15 with a mass of 15.000 amu

• What is the atomic mass of nitrogen?

Ions

• Ions - electrically charged particles that result from a gain or loss of one or more electrons by the parent atom

• Cation - positively chargedresult from the loss of electrons23Na 23Na+ + 1e-

• Anion - negatively chargedresults from the gain of electrons19F + 1 e- 19F-

2.2 Development of Atomic Theory

• Dalton’s Atomic Theory - the first experimentally based theory of atomic structure of the atom.

Postulates of Dalton’s Atomic Theory

1. All matter consists of tiny particles called atoms

2. An atom cannot be created, divided, destroyed, or converted to any other type of atom

3. Atoms of a particular element have identical properties

4. Atoms of different elements have different properties

5. Atoms of different elements combine in simple whole-number ratios to produce compounds (stable aggregates of atoms)

6. Chemical change involves joining, separating, or rearranging atoms

Postulates 1, 4, 5 and 6 are still regarded as true.

• Electrons were the first subatomic particles to be discovered using the cathode ray tube

Indicated that the particles were negatively charged.

Evidence for Subatomic Particles: Electrons, Protons and Neutrons

Evidence for Protons and Neutrons

• Protons were the next particle to be discovered, by Rutherford Protons have the same size charge but opposite in sign

Proton is 1837 times as heavy as electron

• Neutrons Postulated to exist in 1920’s but not demonstrated to

exist until 1932 by Chadwick.

Almost the same mass as the proton (slightly heavier).

Thomson’s model of the atom

http://nobelprize.org/educational_games/physics/quantised_world/structure-images/fig2b.gif

The “plum pudding” model.

Evidence for the Nucleus• Earnest Rutherford’s “Gold Foil

Experiment” lead to the understanding of the nucleus

Most alpha particles pass through the foil without being deflected

Some particles were deflected, a few even directly back to the source

Rutherford’s Gold Foil Experiment

• Most of the atom is empty space• The majority of the mass is located in a

small, dense region

Rutherford’s model of the atom

http://www2.kutl.kyushu-u.ac.jp/seminar/MicroWorld1_E/Part2_E/P25_E/atom.gif

2.3 Light, Atomic Structure, and the Bohr Atom

• Rutherford’s atom – tiny, dense, positively charged nucleus of protons surrounded by electrons

• How do we describe the relationship of the electrons to each other and the nucleus?

• The problem... our classical understanding of physics didn’t work for the atom! This will take some explaining.

Light and Atomic Structure

• Spectroscopy - absorption or emission of light by atoms.

Used to understand the electronic structure.

• To understand the electronic structure, we must first understand light, Electromagnetic Radiation travels in waves from a source

speed of 3.0 x 108 m/s

Radio• Knew from radio

that if we accelerate an electron back and forth in a wire it will radiate a radio wave (electromagnetic radiation).

Wavelengths• Light is propagated (moves) as a collection

of sine waves• Wavelength is the distance between identical

points on successive waves• Each wavelength travels at the same velocity,

but has its own characteristic energy

high energyshort wavelength

low energylong wavelength

Electromagnetic Spectrum

Bohr Theory• Atoms can absorb and emit energy via

promotion of electrons to higher energy levels and relaxation to lower levels

• Energy that is emitted upon relaxation is observed as a single wavelength of light

• Spectral lines are a result of electron transitions between allowed levels in the atoms (in other words, allowed energies are “quantized” in that only certain quantities of energy are allowed.

Electrons exist in fixed energy levels surrounding the nucleus

Quantization of energy

Promotion of electron occurs as it absorbs energy

Excited State

Energy is released as the electron travels back to lower levels

Relaxation

The Bohr Atom

Electronic Transitions• Amount of energy absorbed in jumping

from one energy level to a higher energy level is a precise quantity

• Energy of that jump is the energy difference between the orbits involved

• Orbit - what Bohr called the fixed energy levels

• Ground state - the lowest possible energy state

Modern Atomic Theory

• Bohr’s model of the atom when applied to atoms with more than one electron failed to explain their line spectra

• One major change from Bohr’s model is that electrons do not move in orbits

• Atomic orbitals - regions in space with a high probability of finding an electron

• Electrons move rapidly within the orbital giving a high electron density

The Quantum Mechanical Atom

• Bohr’s model of the hydrogen atom didn’t clearly explain the electron structure of other atoms

Electrons in very specific locations, principal energy levels

Wave properties of electrons conflict with specific location

• Schröedinger developed equations that took into account the particle nature and the wave nature of the electrons

Schröedinger’s equations

• Equations that determine the probability of finding an electron in specific region in space, quantum mechanics

o Principal energy levels (n = 1,2,3…). n is also known as the principal quantum number.

But there is more to the structure of thearrangement of electrons found aroundthe nucleus of an atom, as we will see.

Energy Levels and Sublevels

PRINCIPAL ENERGY LEVELS

• n = 1, 2, 3, …

• The larger the value of n, the higher the energy level and the farther away from the nucleus the electrons are

• The number of sublevels in a principal energy level is equal to n in n=1, there is one sublevel

in n = 2, there are two sublevels

The angular momentum quantum number.

• An object, such as an electron, that moves around another object (the nucleus in our case) will have angular momentum.

• Given by l=0,1,...n-1

• l=0 is s subshell, l=1 is p subshell, etc.

• Subshells increase in energy:

s < p < d < f

• e.g., electrons in 3d subshell have more energy than electrons in the 3p subshell

Principle energylevel (n)

Possiblesubshells

1 1s

2 2s, 2p

3 3s, 3p, 3d

4 4s, 4p, 4d, 4f

Sublevels in Each Energy Level

OrbitalsOrbital - a specific region of a sublevel containing a

maximum of two electrons

• The number of orbitals in a subshell is given by the magnetic quantum number.

• m = -l,...,0,...,+l

• Orbitals are named by their sublevel and principal energy level

1s, 2s, 3s, 2p, etc.

• Each type of orbital has a characteristic shape

s is spherically symmetrical

p has a shape much like a dumbbell

Orbital Shapes

• s is spherically symmetrical

• Each p has a shape much like a dumbbell, differing in the direction extending into space

SubshellNumber of

orbitals

s 1

p 3

d 5

f 7

• An orbital can only hold two electrons!! Once it is filled it cannot accept more.

m = -l, ...0,...+l

Electron Configuration• Electron Configuration - the

arrangement of electrons in atomic orbitals

• Aufbau Principle - or building up principle helps determine the electron configuration Electrons fill the lowest-energy orbital that is

available first Remember s<p<d<f in energy When the orbital contains two electrons, the

electrons are said to be paired and the orbital is full

Rules for Writing Electron Configurations

• Obtain the total number of electrons in the atom from the atomic number

• Electrons in atoms occupy the lowest energy orbitals that are available – 1s first

• No more than 2 electrons in any orbital• Maximum number of electrons in any principal

energy level is 2(n)2

• Follow the periodic table!

Writing Electron Configurations

• HoHydrogen has

only 1 electrono It is in the

lowest energy level & lowest orbital

o Indicate number of electrons with a superscript

o 1s1

• Lio Lithium has 3

electronso First two have

configuration of Helium – 1s2

o 3rd is in the orbital of lowest energy in n=2

o 1s2 2s1

Classification of Elements According to the Type of

Subshells Being Filled

Electron Configuration Examples

• Give the complete electron configuration of each element

Be

N

Na

Cl

Ag

What noble gas configuration is this?

•Neon•Configuration is written: [Ne]3s23p1

Shorthand Electron Configurations

• Uses noble gas symbols to represent the inner shell and the outer shell or valance shell is written after

• Aluminum- full electron configuration is: 1s22s22p63s23p1

• Remember:

o How many subshells are in each principle energy level?

o There are n subshells in the n principle energy level.

o How many orbitals are in each subshell?

o s has 1, p has 3, d has 5, and f has 7

o How many electrons fit in each orbital?

o 2

Shorthand Electron Configuration Examples

N

S

Ti

Sn

2.4 The Periodic Law and the Periodic Table

• Dmitri Mendeleev and Lothar Meyer - two scientists working independently developed the precursor to our modern Periodic Table.

• They noticed that as you list elements in order of atomic mass, there is a distinct regular variation of their properties.

• Periodic Law - the physical and chemical properties of the elements are periodic functions of their atomic numbers.

Classification of the Elements

Parts of the Periodic Table

• Period – a horizontal row of elements in the periodic table. They contain 2, 8, 8, 18, 18, and 32 elements,

• Group – also called families are columns of elements in the periodic table.

• Elements in a particular group or family share many similarities, as in a human family.

Category Classification of Elements

• Metals - elements that tend to lose electrons during chemical change, forming positive ions.

• Nonmetals - a substance whose atoms tend to gain electrons during chemical change, forming negative ions.

• Metalloids - have properties intermediate between metals and nonmetals.

Classification of Elements Metals

• Metals: A substance whose atoms tend to lose

electrons during chemical change Elements found primarily in the left 2/3 of

the periodic table

• Properties: High thermal and electrical conductivities High malleability and ductility Metallic luster Solid at room temperature

Classification of Elements Nonmetals

• Nonmetals: oA substance whose atoms may gain

electrons, forming negative ionso Elements found in the right 1/3 of the

periodic table

• Properties:o Brittleo Powdery solids or gasesoOpposite of metal properties

2.5 Electron Arrangement and the Periodic Table

• The electron arrangement is the primary factor in understanding how atoms join together to form compounds

• Electron configuration - describes the arrangement of electrons in atoms

• Valence electrons - outermost electrons The electrons involved in chemical bonding

Valence Electrons• The number of valence electrons is

the group number for the representative elements

• The period number gives the energy level (n) of the valence shell for all elements

Valence Electrons and Energy Level

• How many valence electrons does Fluorine have?

7 valence electrons

• What is the energy level of these electrons?

Energy level is n = 2

Determining Electron ArrangementPractice

List the total number of electrons, total number of valence electrons, and energy level of the valence electrons for:

• Na

• Mg

• S

• Cl

• Ar

2.6 The Octet Rule

• The noble gases are extremely stable Called inert as they don’t readily bond to other

elements

• The stability is due to a full complement of valence electrons in the outermost s and p sublevels: 2 electrons in the 1s of Helium the s and p subshells full in the outermost shell of

the other noble gases (eight electrons)

Octet of Electrons

• Elements in families other than the noble gases are more reactiveo Strive to achieve a more stable electron

configurationo Change the number of electrons in the atom to

result in full s and p sublevels

• Stable electron configuration is called the “noble gas” configuration

The Octet Rule

• Octet Rule - elements usually react in such a way as to attain the electron configuration of the noble gas closest to them in the periodic tableo Elements on the right side of the table move right to the

next noble gaso Elements on the left side move “backwards” to the

noble gas of the previous row

• Atoms will gain, lose or share electrons in chemical reactions to attain this more stable energy state

NaSodium atom

11e-, 1 valence e-

[Ne]3s1

Na+ + e-

Sodium ion10e-

[Ne]

Ion Formation and the Octet Rule

• Metallic elements tend to form positively charged ions called cations

• Metals tend to lose all their valence electrons to obtain a configuration of the noble gas

AlAluminum atom13e-, 3 valence e-

[Ne]3s23p1

Al3+ + 3e-

Aluminum ion10e-

[Ne]

• All atoms of a group lose the same number of electrons

• Resulting ion has the same number of electrons as the nearest (previous) noble gas atom

Ion Formation and the Octet Rule

Using the Octet Rule

• The octet rule is very helpful in predicting the charges of ions in the representative elements

• Transition metals still tend to lose electrons to become cations but predicting the charge is not as easy

• Transition metals often form more than one stable ion Iron forming Fe2+ and Fe3+ is a common example

Examples Using the Octet Rule

• Give the charge of the most probable ion resulting from these elementsCa

Sr

S

P

• Which of the following pairs of atoms and ions are isoelectronic?Cl-, Ar

Na+, Ne

Mg2+, Na+

O2-, F-

+K3919

-23216S

+22412 Mg

Calculating Subatomic Particles in Ions

• How many protons, neutrons and electrons are in the following ions?

2.7 Trends in the Periodic Table

• Many atomic properties correlate with electronic structure and so also with their position in the periodic table atomic size ion size ionization energy electron affinity

Atomic Size• The size of an element increases moving

down from top to bottom of a group

• The valence shell is higher in energy and farther from the nucleus traveling down the group

• The size of an element decreases from left to right across a period

• The increase in magnitude of positive charge in nucleus pulls the electrons closer to the nucleus

Variation in Size of Atoms

Cation SizeCations are smaller than their parent atom• More protons than electrons creates an increased

nuclear charge• Extra protons pulls the remaining electrons

closer to the nucleus• Ions with multiple positive charges are even

smaller than the corresponding monopositive ions

Which would be smaller, Fe2+ or Fe3+? Fe3+

• When a cation is formed isoelectronic with a noble gas the valence shell is lost decreasing the diameter of the ion relative to the parent atom

Anion Size

Anions are larger than their parent atom.

• Anions have more electrons than protons

• Excess negative charge reduces the pull of the nucleus on each individual electron

• Ions with multiple negative charges are even larger than the corresponding monopositive ions

Relative Size of Select Ions and Their Parent Atoms

ionization energy + Na Na+ + e-

Ionization Energy

• Ionization energy - The energy required to remove an electron from an isolated atom

• The magnitude of ionization energy correlates with the strength of the attractive force between the nucleus and the outermost electron

• The lower the ionization energy, the easier it is to form a cation

Ionization Energy of Select Elements

Br + e- Br- + energy

Electron Affinity

• Electron Affinity - The energy released when a single electron is added to an isolated atom

• Electron affinity gives information about the ease of anion formation

Large electron affinity indicates an atom becomes more stable as it forms an anion

Periodic Trends in Electron Affinity